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Lesson 1Introduction

As a beginning water conditioning professional, you are sometimes calledupon to sell equipment in different communities. In many cases, thesecommunities get their water supplies from different sources. Perhaps onecommunity draws its water from a lake; another from a large communitywell system. It is superfluous to add here that these water sources maypossess extreme variations in quality.

If you sell to prospects in rural areas, it is not only possible but prob-able that you may have run into prospects who, although their wells areadjacent, obtain water that is remarkably different in quality.

Why are such contrasts in the quality of various water supplies pos-sible? What causes these wide variations in water?

The many variations and conditions of water are all symptoms of thefact that water is a solvent. Many scientists believe it is one of the finestsolvents available. But to understand more about water, one should un-derstand the life cycle of water…its remarkable and ceaseless travelsabove, below, and on the surface of the earth.

Water and Water Quality

Water…what a wonder! People use it in so many ways, it is perhaps un-realistic to expect it to meet all the demands they make of it. Still, withthe right treatment, water can and does meet all its obligations.

Public health authorities, industrial firms, commercial firms, hos-pitals and institutions, farmers, and homemakers…each has special re-quirements in terms of water quality. And when water quality fails tomeet these requirements, trouble begins.

Even the space age scientist gets into the act when he calls fordeionized water to clean the metal skins of his satellites. His requestfor deionized water stems from the fact that it prevents local “hot spot”corrosion and thus unwanted residue weight as his ships soar out intospace. More mundane uses are for final rinsing of automobiles, trucks,and aircraft.

Deionized water. Water from which most of the mineralshave been removed, normally by the ion exchange process.Usually two ion exchange resins are used. One resin takesout all positively charged ions (Ca++, Mg++, Na+, etc.) andreleases a chemically equivalent amount of hydrogen ions.The second resin takes out the negatively charged ions(HCO3

–, Cl–, SO4––, etc.) and releases an equivalent of hy-

droxide (OH–). The hydrogen and hydroxide ions introducedthen combine to form water. (H+ + OH– = HOH or H2O.) Alsosometimes called demineralized water.

Deionization is also called demineralization, but not with completeaccuracy, since the term deionization specifies the removal of dissolvedsubstances in ionic form. Water treatment processes, such as distilla-tion and reverse osmosis, also remove dissolved substances (dissolvedsolids) from water. These processes have the advantage of being ableto remove not only dissolved substances in ionic form from water, butalso substances such as sugar and other organic matter which dissolvein water but do not form ions. Thus, water may contain dissolved im-purities or contaminants which are ionized and those which are not.Dissolved substances which form ions make water a better electrolyte(conductor of electricity), and the amount of ionized substances pres-ent in water can be measured by its conductivity, and conversely, theirabsence by increased electrical resistance.

Deionization, distillation, and reverse osmosis are processes whichcause direct removal of impurities from water as contrasted with in-direct processes which involve conversion of water impurities to their

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insoluble form as precipitates, with subsequent direct removal byfiltration, such as the removal of dissolved iron by oxidation and sub-sequent filtration.

To provide the right water for any demand, whether it be that ofthe public health authority, the homemaker, or the space age scientist,two all-important factors must be considered:

1. Precisely what does analysis of the raw water supply indicate?2. To what end use will the water be put?

Analysis of a water may show that it contains (a) dissolved miner-als, (b) dissolved gases, (c) turbidity and sediment, (d) color and organicmatter, (e) taste and odor, and/or (f) microorganisms.

Microorganisms. These are extremely small animal or veg-etable organisms, especially any of the bacteria, protozoa,or viruses. Some of these microorganisms are so small thatthey cannot be seen under conventional microscopes.

Whether or not any of these impurities are harmful in a given sit-uation in turn depends on:

1. the nature and the amount of the impurities;2. the tolerance permissible for each of these impurities; and3. the end use of the water.

Water of a quality that may prove unacceptable or unsatisfactoryfor certain requirements may be quite satisfactory in other instances.To cite an example, water with 15 grains per gallon of hardness (257milligrams/liter) is objectionable for laundering and bathing. This samewater, however, is satisfactory for sprinkling the lawn.

Both the quality of a raw water and its end use must always be de-termined before it can be treated economically.

What precisely is this fascinating substance—water?Webster defines it briefly as: “The liquid which descends from the

clouds in rain, and which forms rivers, lakes, seas, etc. Pure ordinarywater (H2O) consists of hydrogen (11.188 percent) by weight and oxy-gen (88.812 percent). It has a slightly blue color and is very slightlycompressible. At its maximum density at 39.2°F or 4°C, it is the stan-dard for the specific gravities of solids and liquids. Its specific heat isthe basis for the calorie and the Btu units of heat. It freezes at 32°F or0°C.”

Note the term “pure water” in this definition. We often use theterm “pure water.” But, “pure water” (H2O) occurs so rarely that prac-tically speaking, it is a nonexistent liquid.

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Even the term “pure water” is somewhat ambiguous. It has differ-ent connotations to individuals in various fields. The bacteriologist,for example, is apt to regard “pure water” as a sterile liquid, that is, onewith no living bacteria in it. The chemist, on the other hand, mightwell classify water as “pure” when it possesses no mineral, gaseous, ororganic impurities. It is obvious that “pure water”, as described here,is likely to be found only in laboratories…and even there only underideal conditions.

The United States Environmental Protection Agency (EPA) providespractical standards for water in terms of its suitability for drinking (orpotability) in the Primary Drinking Water Regulations and for aestheticconsiderations in the Secondary Drinking Water Regulations.

In its Drinking Water Regulations, the USEPA takes into consider-ation adequate protection of water against the effects of contamina-tion, both through natural processes and through artificial treatment.The standards list requirements for bacterial count, physical, and chem-ical characteristics.

It is almost impossible to find a source of water that will meet basicrequirements for a public water supply without requiring some formof treatment. In general, the requirements for a public water supply areas follows:

1. That it shall contain no disease-producing organisms.2. That it be colorless and clear.3. That it be good-tasting, free from odors, and preferably cool.4. That it be noncorrosive.5. That it be free from objectionabla gases, such as hydrogen

sulfide, and objectionable staining minerals, such as iron andmanganese.

6. That it be plentiful and low in cost.

While the presence of coliform bacteria and toxic chemical con-tent in a water supply would cause a water to be classified as unsafe todrink, other factors such as taste, odor, color, and mineral content havea certain aesthetic effect which can cause a water to be rejected as a us-able supply.

Coliform bacteria. An organism of the bacteria family,harmless in itself, but since Escherichia coli (E. coli), a memberof this group, exists and grows as part of the normalmicrobe population in the digestive tract of warm bloodedanimals, it serves as a strong indicator of sewage contami-nation. Detecting its presence in a water supply is compar-

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atively easy. One caution: while E. coli is one of the most pro-lific and beneficial bacteria in the human body, helping tometabolize food in the intestine, it also has the ability tomutate. A new strain called E. coli 0157:H7, or 0157 forshort, can produce a strong toxin. Compared to ordinarystrains of the E. coli bacterium, 0157:H7 is rare; but it can becarried in food and water, and it can cause serious illnessand even death in about one percent of infected people.

A potable, or safe water, is not necessarily usable or useful for manypurposes. For this reason, it may require treatment of another sort torender it useful to the needs of the home or industry…or for use by thespace-age scientist, for example. In any event, no snap judgementshould be the basis for determining whether or not a certain water canmeet requirements for a certain use.

There are tremendous variations in the quality of water from areato area. Review of the maps at the end of this chapter gives some indi-cation of the variations. These, however, are only broad, general indi-cations of the differences. Even within a specified area, significantdifferences may be noted.

In some cases there are variations in the quality of water in a givenarea, even on a day-to-day basis. Why do such variations occur?

The answer can be traced to the fact that water is a solvent. Wateris aptly described as “the universal solvent.” Scientists generally agreethat it is one of the best solvents available.

As a result of its solvent action, water dissolves at least a portion ofeverything it touches. It dissolves metals, rocks, waste matter, gases,dust, and numerous other foreign substances and may contain appre-ciable amounts of these dissolved materials.

The dissolved mineral content of water ranges from 20 to 80 partsper million (milligrams per liter) in areas where there are only slightlysoluble granite formations. From this low level, it increases quite no-ticeably depending on area conditions.

The dissolved solids content of the oceans is in the 35,000 ppm(mg/L) range. It is estimated that there are enough dissolved solids inthe oceans to cover all the earth’s land surfaces to a depth of 112 feet.Each year inland waterways carry billions of tons more of dissolvedsolids into the oceans.

In any area, the dissolved solids content of a water supply may varysharply depending on whether the water is drawn from a deep well, alake, a river, or a pond.

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The Hydrologic Cycle

Water begins its never-ceasing cycle as vapor in the atmosphere. Thisvapor in the atmosphere, as well as the water in the lakes and oceans,provides protection against extremes of both heat and cold.

Hydrologic cycle. This term, or the more common one“water cycle,” refers to the complete sky-to-earth-to-sky cir-cuit pursued by water in nature. It includes water’s precipi-tation as rain, snow, hail, or dew; its journey over, around,and through obstacles above, on, and below the earth’s sur-face and its eventual evaporation and return to the atmos-phere. It is the largest water purification system known toman.

Scientists estimate that the sun converts matter into energy at therate of 250 million tons per minute. Even though the earth receivesonly a token portion of this heat energy (less than one two-billionthpart), everything here would burn to a crisp were it not for the fact thatthe water above and on the earth absorbs most of the heat.

In a large desert, for example, there is but a small amount of water.Consequently, there are wide extremes in the heat. The Sahara Deserttypifies this condition. There, under the sun’s penetrating rays, tem-peratures rise to 125°F during the daytime and fall below freezing atnight.

In the atmosphere, the various substances do not combine chem-ically. Instead, each retains its own characteristic properties.

The make-up of the atmosphere. The composition of thetroposphere (the layer closest to the earth) has been calcu-lated to be nitrogen, 78.09 percent; oxygen, 20.95 percent;argon, 0.93 percent; carbon dioxide, 0.03 percent; togetherwith minute amounts of neon, krypton, helium, hydrogen,xenon, and ozone. In addition to these gases, the atmos-phere contains varying percentages of water vapor. About9⁄10 of the mass of the atmosphere lies within ten miles ofthe earth’s surface. The above figures are composition byvolume.

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Meteoric Water

When millions of vapor particles unite, they form droplets of moisture.As these increase in size, they finally become heavy enough to fall toearth as precipitation in such varied forms as rain, snow, sleet, hail,and dew.

Meteoric water. A term applied to all moisture precipitat-ing from the atmosphere. Depending on conditions, it mayfall as rain, snow, sleet, or hail.

It is estimated that 16 million tons of precipitation in any of theseforms falls earthward each second. Through the process of evaporation,it is then drawn back into the atmosphere. In nature’s balanced oper-ations, evaporation equals precipitation.

As water falls to earth in this never-ceasing moisture circulatingsystem, it serves to cleanse both the air and the ground. No doubt youhave, many times, noted the fresh, clean smell of the air after a heavyrain. This is because the rain has absorbed suspended solid matter (dust,dirt, and soot), gases, odors and other impurities, polluting the air overthe area. While precipitation may remove large quantities of impuri-ties, it never succeeds in wholly eliminating them.

When precipitation continues for some time, the first amounts tofall are apt to contain a great deal more suspended particles and dis-solved solids than that which falls later. An analysis of the mineral con-tent of rainwater in a large city after four hours of precipitation andagain after 22 hours shows the following variations (expressed as partsper million as calcium carbonate):

After 4 hours After 22 hoursHardness 43 8Calcium 42 8Magnesium 1 —Sodium 11 —NH3 (ammonia) 3 5Bicarbonate 19 5Chloride 10 5Sulfate 27 3Nitrate 1 —

As you can see, even this sweeping of the atmosphere did not re-move all the dissolved solids in 22 hours of rainfall.

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Of all forms of precipitation, the snow falling high in the moun-tains contains the least amount of mineral content. This is due, how-ever, to the smaller amount of dust in the atmosphere at high altitudes.As a result, many mountain streams deriving their water from highfallen snow have extremely low dissolved mineral content.

Rainwater is also saturated with dissolved air (about 20 to 29 mil-liliters per liter from 60°F to 32°F). The amount of free carbon dioxidein rain varies from two to six parts per million. Any amount of free car-bon dioxide above 1 or 2 ppm comes not from the atmosphere itself,but from other sources such as chimneys or industrial fumes. Rainwateralso encounters sulfuric acid from the gases in burning coal over cities.In addition, it may pick up bacteria and the spores of microorganisms.

Milliliter. The milliliter is 1⁄1,000 of a liter.

Liter. This term is a unit of volume, slightly larger than aquart. One U.S. gallon is the equivalent of 3.785 liters.

How much water falls in a day in the United States? The UnitedStates Geological Survey has estimated that approximately 4,300 bil-lion gallons of water fall within the continental limits of this countryeach day.

The sun draws about 70 percent of this daily precipitation back upinto the atmosphere through the process of evaporation almost im-

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The Hydrologic Cycle

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mediately—certainly before it seeps into the soil or goes far in theprocess of run-off. As shown in the diagram, the upturned arrows in-dicate that the sun causes evaporation of water even while it is falling.The sun also draws water from the soil, from surface run-off, vegeta-tion, streams, lakes and oceans, and through the process of transpira-tion.

Transpiration. This term refers to the loss of water byplants through evaporation. Leaves, in particular, lose waterthrough this process. Transpiration involves considerablequantities of water. It is estimated that a birch tree will loseapproximately 500 quarts of water on a dry day. Excessivetranspiration (when the amount of water lost exceeds theamount absorbed by the roots) results in wilting of leaves.Plants also exude water in a liquid form in a process knownas guttation.

The 30 percent of precipitation which is not quickly evaporatedeither seeps deep into the soil or finds its way into lakes and rivers andeventually flows into the oceans. Surface topography, porosity of thesoil, the degree of its saturation at the time of a rainfall, surface vege-tation, and atmospheric conditions—these are all factors that help todetermine the distribution of water after precipitation. Water’s solventaction, which permits it to have a cleansing action on the atmosphere,continues after it reaches the earth. A certain percentage of precipita-tion becomes surface runoff. In this process, it acquires further amountsof hardness minerals in addition to quantities of clay, silt, and decayedanimal and vegetable matter. The violence of surface runoff often leadsto serious land erosion problems. Not only is much of this water lostto the oceans, but it also brings with it vast amounts of top soil.

The Grand Canyon of Arizona is a monumental example of theerosive action of water. Not all runoff is violent. Where heavy vegeta-tion and gently sloping grades permit, water enjoys an almost imper-ceptible rate of flow. Under such conditions, it can absorb objectionabletaste, odor, and color from available decaying plant and animal life.Only a portion of the total precipitation seeps into the soil.

Curiously, when water percolates into the ground, it loses some ofthe impurities it absorbed from the air and on the ground. But whilethe soil structure filters out certain impurities, it provides ampleopportunity for water to dissolve large amounts of earth minerals.These, of course, increase its hardness and iron content, among otherthings. As water seeps into the soil, it begins a journey that may carryit for quite some distance through underground interstices, crevices,

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and fissues. Further, the journey may require many years before thiswater is pumped to the surface.

Environmental Factors

A water supply is the product of its environment. The gases carbondioxide and oxygen enter the water from the atmosphere. The carbondioxide can unite with water to form carbonic acid.

In vegetated areas, oxygen in water is consumed and carbon diox-ide increased through decay of vegetation.

In limestone areas, the water containing carbonic acid reacts withlimestone and becomes hard. Calcium and magnesium bicarbonatesare formed.

In granite or sandy areas, the water retains its carbonic acid, butdoes not become hard, due to the absence of limestone.

In arid regions, oxygen from the atmosphere is not consumed toany degree, nor is carbon dioxide increased by decay. Where sand andgranite predominate, the water will be low in hardness and slightlyacid. In areas where calcium or magnesium chloride or sulfate are

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How Water Collects impurities

found, the water will become very hard. The hardness will be chieflynoncarbonate even though limestone is present.

This environmental background has important implications in re-gard to the corrosiveness of well or groundwater, as follows:

1. In vegetated areas where limestone and other hardness min-erals are present, the hard water will not be corrosive due toneutralization of carbonic acid and the virtual absence ofdissolved oxygen. When such water is softened by ion ex-change, the corrosion rate will remain low. The corrosive-ness of both the hard and softened water will increase if theyare aerated.

2. In vegetated areas where granite and sand predominate, thewater will be low in hardness and usually low in total dis-solved solids (conductivity). It can be corrosive, however,due to the presence of carbonic acid which can dissolve irondirectly. Such water supplies usually produce objectionable“red” (rusty) water, but corrosion is usually uniform ratherthan of the “pitting” type. Copper or other corrosion resist-ant materials will be much more satisfactory than galvanizedsteel in such supplies. Neutralization will control corrosion.

3. In arid regions where limestone and noncarbonate hardnessminerals are found, both the hard and softened water sup-plies will tend to be corrosive due to their dissolved oxygencontent and conductivity.

4. In arid regions where granite and sand predominate, watersupplies will be low in hardness and conductivity. They maystill be corrosive, however, due to their dissolved oxygencontent, but usually are less corrosive than water suppliesfrom arid regions which have a higher dissolved solids con-tent (as in #3 above).

It is apparent that the environment can be a guide to the corrosivenature of a water supply.

Note: Contrary to the prevailing notion that oxygen-depleting re-actions in the soil zone and in the aquifer rapidly reduce the dissolvedoxygen content of recharge water to detection limits, two to eight mil-ligrams per liter of dissolved oxygen have been found in water from avariety of deep aquifers in Nevada, Arizona, and the hot springs of theAppalachians and Arkansas, Science Magazine reports. The prevailingopinion is that the majority of dissolved oxygen in recharge water isconsumed in the soil and unsaturated zones by microbial respirationand the decomposition of organic matter, or rapidly, thereafter, in theaquifer by various mineral-water and organic-oxidated reactions. USGS

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researchers document the widespread presence of dissolved oxygen insignificant (two to eight mg/L) concentrations in water several thou-sand to more than 10,000 years old from deep aquifers in both arid andhumid climates, and at distances as great as 80 km from recharge areas.

More puzzling is the presence of dissolved oxygen in thoseArkansan and Appalachian hot springs in which water has passed prin-cipally through fractured siliceous rocks. Perhaps all pertinent reactions(organic or inorganic) involving dissolved oxygen have gone to com-pletion within the aquifer prior to entry of the extant groundwater, theUSGS scientists postulate.

Age of Groundwater

The period of time since groundwater fell as rain can now be estimatedby a technique based on the amount of tritium found in groundwater.This technique was developed by Dr. Willard Libby, who was one ofthe members of the Atomic Energy Commission, and some of his for-mer associates at the Institute for Nuclear Research at the University ofChicago.

Tritium is a radioactive isotope of hydrogen, believed to be formedin the atmosphere from the action of cosmic rays on ordinary hydro-gen. Thus, tritium is found in all atmospheric water, such as rain andsnow. As a radioactive material, tritium gradually decays or decomposesinto simpler substances and has a known “half-life” of 121⁄2 years. Thatis, one half of the radioactive form is dissipated in 121⁄2 years. An ad-ditional half is lost in the succeeding 121⁄2 years, and so on, until theamount remaining is too small to be measured.

Isotopes. Forms of atoms of an element which differ in theirmasses due to variations in the numbers of mass particles intheir nuclei. Hydrogen has three known isotopes: the mostcommon form has only a proton (a relative mass of one anda single positive electrical charge) in its nucleus; a secondisotope known as “deuterium” has one proton and a neu-tron (neutral in charge and also with a relative mass of one)in its nucleus, and thus a relative mass of two; a third iso-tope known as “tritium” has two neutrons and a single pro-ton in its nucleus, and thus has a relative mass of three.

As the approximate amount of tritium originally present in water asit fell as rain is known and the amount remaining can be measured,the length of time the water has been underground can be calculated

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unless the amount remaining is too small to be detected by the in-struments currently available.

Tests of this type on deep well water from several locations in Ne-braska indicated underground water ages of about 14 to 61 years; testson Illinois water gave ages of 50 to more than 100 years. (Beyond 100years, the tritium concentrations could not be measured accurately.)These ages are generally in keeping with the anticipated values when allthe hydrologic factors in each area are considered.

As the diagram shows, water must travel through various strata be-fore becoming groundwater. Below the surface, it moves first throughthe subsoil (the belt of soil water), the intermediate layer, the capillaryfringe, and finally into the groundwater table.

These layers vary in depth and are not too sharply defined. In fact,there is a gradual transition from one to another until the groundwaterlevel or zone of saturation is reached. Even after water moves into thetopsoil and subsoil, much of it may still return to the atmosphere eitherthrough evaporation or transpiration. Water is held in the subsoil by mo-lecular attraction. It is only after sufficient water has accumulated herethat it begins to seep downward under the pull of gravity. The subsoilmay extend down 50 feet. It supplies the water needed for the growthof vegetation. Consequently, it is extremely important to farmers.

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Groundwater Zones and BeltsZ

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Water in the intermediate belt is generally considered in “dead stor-age.” To all intents and purposes, it is suspended and does not flowinto wells. This belt varies from a hairsbreadth to several hundred feet.

Below lies the capillary fringe. Water in this fringe is continuouswith the water in the zone of saturation, but is held back by capillaryaction. The thickness of this capillary fringe depends on its composi-tion. In silty materials, it may extend down for several feet. In coarse,gravelly materials, it may go down less than an inch. Even in this cap-illary fringe, water will still not enter well systems. It is only when itreaches the zone of saturation that it may be drawn back up to the sur-face by wells.

Capillary action. Where water touches a solid, capillary ac-tion causes the water at that point to rise higher than thatportion of its surface not in contact with the solid. Capillaryaction is due to adhesion, cohesion, and surface tension.Capillarity is one of the causes of water’s rising in the soil asin the capillary fringe. Kerosene rising in the wick of an old-fashioned lamp is another example of this seeming contra-diction of the law of gravity.

This zone of saturation forms a huge natural reservoir that feedssprings and streams in addition to our wells. Its thickness varies fromtwo to hundreds of feet, depending on local geologic conditions. Theupper surface of the zone of saturation is neither stationary nor level.It possesses many surface irregularities and may range up or downmany feet over a period of years at any given location. The fluctuationsin its content depend on the amount of recharge and pumpage.

In general, the contours of the water table parallel the surface con-tours. However, the water table goes deeper under high elevations andrises nearer to the surface under lower elevations. At springs and flow-ing streams, the surface and water table elevations coincide. Below theeconomically important zone of saturation lies dense, solid rock. Whilethis rock is known to hold substantial amounts of internal water, thereis no practical way of bringing it to the surface.

Groundwater

Under most conditions, groundwater supplies are higher in mineralcontent than surface waters in the same area. This is due to their longerexposure to rock formations. Exceptions do occur, as when surfacewaters originate in a region of relatively soluble rock and later flow into

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an area of less soluble rock. In such cases, groundwaters in the latterarea may be lower in mineral content than that of waters nearer theground surface.

Meanwhile, as water seeps through the ground and adds to its min-eral content, much of its suspended matter, color, and bacterial con-tent are filtered out. Thus, a deep well is likely to provide water that isclear, colorless, and low in bacterial count. Of course, there are excep-tions. It might be expected that the deeper the wells go, the morehighly mineralized are their waters. This is the usual case. In some shal-low wells, however, the mineral absorption is far greater than for deepwells in the same general area.

Temperature of Groundwater

The temperature of water from wells is remarkably constant. In wellsthat are from 30 to 60 feet deep, water temperature is 2° to 3°F abovethe annual mean temperature of the locality. Water decreases in tem-perature about 1°F for each 64 feet of depth to the well.

In general, deep wells extend down through an impervious layerto reach an underlying supply. Shallow wells, in contrast, are sunk ineasily penetrated strata to a point where they are below the water table.In terms of depth, deep wells are classified as those extending below25 feet; those going less than this are considered shallow wells. Actu-ally, deep wells vary from 100 to 3,000 feet. The vast majority are inthe 100 to 1,000 foot range. Deep well water usually shows but slightchange in composition over a long period of time. In one study ofsome wells in Florida over a 24 year period, hardness ranged from ahigh of 342 to a low of 304 parts per million. Alkalinity went from ahigh of 168 to a low of 148.

Springs provide another source of groundwater. It is a popular be-lief that spring waters are clear, colorless, sparkling, and absolutely pure.While these facts hold true for many springs, others show a markeddegree of turbidity, especially after a heavy rainfall.

Spring waters further contain rather large amounts of dissolvedmineral matter and are hard. On the score of potability, no spring watershould be considered safe to drink unless it is given periodic bacterialexamination.

Other groundwaters could also include mine waters and connatewaters. Large quantities of water are found in many mines and mustbe removed by pumping. In some cases, mine waters are no differentthan other ground supplies. Generally, however, they have a high sul-furic acid and iron content. As a result, they may be extremely corro-

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sive. Connate water or oil field brines are the remains of ancient seasin which sedimentary rock was, at one time, deposited. These “fossilwaters,” as they are sometimes called, are generally highly saline. Inthe operation of oil fields, they have only nuisance value and presentserious disposal problems when brought to the surface.

While groundwater supplies have definite advantages, they alsopresent problems. Important disadvantages are:

1. The presence of hardness mineral compounds in largeramounts than in the surface waters of the same locality as arule.

2. Iron and manganese are present in many well supplies.3. Hydrogen sulfide is sometimes present.4. The cost of pumping well water is usually greater than that

for pumping surface water.5. The mineral content of several wells may differ widely even

though located close to each other.6. The supply may be uncertain.7. They may contain nitrate or detergent contamination. The

presence of nitrates or detergents in a groundwater supplycan indicate pollution from sewage.

Surface Waters

Lakes, rivers, reservoirs, ponds, etc., are termed surface waters. They re-ceive water directly from precipitation and surface runoff. These variousbodies of water also receive a portion of their total amount from un-derwater springs connected with the groundwater supply. The previ-ous diagram (Groundwater Zones and Belts) shows how the bed ofstreams extends below the groundwater level.

As we have seen, surface waters are generally lower in mineral con-tent. On the other hand, they possess far more contamination and areunsafe to use for human consumption unless properly treated.

Pollution of water comes from many sources. Municipalities andindustries sometimes discharge waste materials into bodies of waterthat are used as public sources of supply. This is a most serious sourceof contamination. Surface runoff also brings mud, leaves, decayed veg-etation together with human and animal wastes into streams and lakes.In turn, these organic wastes cause algae and bacteria to flourish.

There is a belief that rivers and streams purify themselves in thecourse of their flowing 20 miles. This action should not be taken forgranted, however.

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Organic pollution of water is reduced by nature in many ways:

1. Bacteria and algae consume large quantities of organic waste.Larger microorganisms devour the bacteria and algae. Inturn, the microorganisms provide food for fish and otherhigher forms of animal life.

2. Unless the rate of flow is too fast, mud and suspended mat-ter will naturally settle to the bottom and oxidation will ren-der organic matter harmless. Rough bed streams, riffles, andspillways speed this process.

3. Due to its ultraviolet rays, sunlight also has some germicidaleffect on the water. Sunlight is not constant due to cloudyweather and its unavailability at night.

Algae. Any of a group of one-celled or many-celled mi-croorganisms which are found in water and damp places.Algae contain chlorophyl and have no true root, stem, orleaf structure. Included among the algae are seaweed andpond scum.

Oxidation. This term will be discussed in Lesson 3.

Riffle. A riffle is a shoal, reef, or rocky obstruction in astream of water. As the water flows over and around the ob-struction, a stretch of shallow, rapid, or choppy water is pro-duced depending on the nature of the obstruction.

Rivers and streams also show great variations in their dissolvedmineral content. Tests taken over a period of a year at both the RockRiver and the Arkansas River showed that both had the same averagebicarbonate content of 207 ppm. In contrast, the Rock River had a totalchloride and sulfate content of 30 ppm, while the Arkansas River con-tained 613 ppm of these ions, mostly present in the form of hardnesscompounds.

In general, lakes and reservoirs (especially large ones) show fairlyconstant dissolved solids content. Because they are relatively morequiet than moving bodies of water, lakes and reservoirs are very effi-cient settling basins. The result is they possess less turbidity. Large bod-ies of water are frequently subject to seasonal changes that cause thewater to become quite turbid for a period of time. Our definition ofwater at the beginning of this lesson states that it achieves its maxi-mum density at 39.2°F. As it becomes chilled to this point in the fall

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or warmed to it in the spring, the denser water cannot stay at the top.As it sinks, it causes convection currents to be set up. Sometimes thesebecome so strong that they lead to a complete overturning of the waterand bring about the turbid condition. Heavy storms will also churn upa lake or reservoir and make it turbid.

Cistern Water

In some areas, there are attempts to collect rainwater in cisterns. Ingeneral, these cistern waters are harder and contain more total solidsthan rain. This is due to the accumulation of dirt and dust on the sur-faces that drain to the cisterns. One study shows that in 500 householdcisterns, hardness ranged from 35 to 150 ppm. Further, cistern watersoften have a high bacterial count and noticeable color. While in manycases the organisms found in cisterns are nonpathogenic, it is advis-able to chlorinate this water where it is used for drinking purposes.

Nonpathogenic. A pathogen is any microorganism or virusthat is disease-producing. Hence “nonpathogenic” is a termthat refers to the fact that a substance does not contain dis-ease-producing organisms.

Summary

The source of any water supply determines the kinds and amounts of itsimpurities. Groundwater obtained from deep wells usually contains highconcentrations of dissolved minerals. This water is usually clear and col-orless due to its filtration through rock and sand. It also may containvarious types of pollution, including detergents and industrial wastes.It is now known that such forms of pollution may travel quite some dis-tance in water. Shallow wells provide water with varying amounts ofmineral impurities. There is also the danger that water from such sourcesmay become contaminated with human and animal wastes.

Surface waters contain many impurities—silt, sand, and clay—which give them a muddy or cloudy appearance. If runoff to thempasses over agricultural land, it may also absorb chemical wastes andtoxic refuse from animals.

Where water flows sluggishly through swampland, it may acquireobjectionable taste, odor, and plant color. During periods of flooding,these swamps may discharge their decayed vegetation, color, andmicroorganisms into moving streams and rivers.

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Deep wells and large lakes alone provide water that is more or lessconsistent from season to season. Smaller bodies of water, shallowwells, and springs often reflect seasonal—even daily—variations in theirmineral content.

To understand why water from different sources varies in quality,it is necessary to know something about basic water chemistry.

When suspended in the atmosphere, water vapor approximates dis-tilled water. It is relatively free from impurities and remains thus aslong as it stays aloft. When water vapor condenses sufficiently to fallto earth, it comes into contact with gases in the surrounding air—car-bon dioxide, nitrogen, and oxygen. Atmospheric dust may also con-tain minute particles of silica, oxides of iron, and other materialstogether with dust, pollen, and some microorganisms.

In falling, moisture absorbs amounts of the atmospheric gases be-cause these are partially soluble in water. The colder the water, the moreof the surrounding gaseous content it dissolves.

If we chemically diagram the action of water as it dissolves someof the carbon dioxide in the air, it would look this way:

H2O + CO2 → H2CO3

Water Dissolves and Collects Carbon Dioxide to Produce Carbonic Acid

Normally when such water reaches the earth, it is slightly acid, cor-rosive, and relatively soft (though not as soft as man can make itthrough his skill in the treatment of water). After water reaches theground, it may pick up additional amounts of carbon dioxide from de-caying vegetable matter. Equipped with this booster action, it acquireseven greater potential for dissolving minerals and other impurities onor below the surface. Water at the surface is slightly acid. If, however, ithas the opportunity to seep into the soil and pass through a limestonestratum, the acid condition due to the carbon dioxide will be neutral-ized. At the same time, the water will pick up a large amount of mineralcontent. Chemically this can be diagrammed:

H2CO3 + CaCO3 → Ca(HCO3)2

Carbonic Acid Reacts with Insoluble Calcium Carbonate to ProduceSoluble Calcium Bicarbonate

Limestone, a common rock formation, contains varying portionsof both calcium and magnesium carbonates. These are the unseenhardness minerals which plague so many supplies. The basic reactionshown in the above diagram holds true for both minerals. Iron andmanganese are found in water supplies less frequently. But again, their

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basic chemical reaction in water is quite similar. Because it is a solvent,water also picks up the soluble chlorides, sulfates, and nitrates of cal-cium and magnesium. Similarly, it absorbs the carbonate, bicarbonate,chloride, sulfate, and nitrate compounds of sodium as well as quantitiesof silica. Close scrutiny of a water supply after exposure to many com-mon gases and minerals will give a good idea of the active solvent thatwater can be.

In Lessons 2 and 3 we will review the basic chemistry that goes onin the reactions of water molecules with the various compounds in it.

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Groundwater Hardness Areas in the United States

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Groundwater Quality Areas in the United States

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Groundwater Quality Areas in the United States(Continued)

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Lesson 2Introduction

Our study of the water cycle indicated that water vapor in the atmosphereis virtually free from impurities, but as it forms droplets and falls to earthin any of its various forms, it picks up and dissolves a variety of impuri-ties. Some of these impurities produce physical changes in the water. Oth-ers lead to chemical changes.

An example of a physical change is the presence of turbidity result-ing from dust picked up by the droplets falling to earth. A chemical changeresults from the act of calcium and magnesium salts dissolving in thewater to form hardness.

In order to understand more about these changes and how they takeplace, it is necessary to review some basic chemistry. This second lessonprovides a quick refresher course on this subject.

Please bear in mind that no effort has been made to treat this sub-ject in detail. If you wish to pursue the subject further, contact the WQAtechnical staff.

Introduction to Water Chemistry

Based on the discoveries of our satellites, it appears that water is aunique substance in our discovered universe. The presence of water onearth is in itself unique, for the planet earth has few natural liquids.Water is the prime resource of man’s food supply and his most impor-tant household and industrial tool. But most important is the fact thatwater is a major constituent of all living matter, comprising up to two-thirds of the human body. Next to the air we breathe, water is man-kind’s most important substance.

Scientists now agree that the entire universe is composed of justtwo ingredients—energy and matter. Under the proper conditions, bothenergy and matter are capable of change. Energy can be converted intomatter, and matter can produce energy. Such conversion suggestschange. And change, of course, continues all the time. Let’s briefly con-sider the types of change that occur.

Physical and Chemical Changes

Take a lump of coal, for example. First break it into a half dozen pieces.Next set fire to a few of the pieces.

In breaking up a large lump of coal in several or many smallerpieces, you have produced physical change. The size of the originallump has changed. But nothing has happened to the composition ofthe unburned segments.

Physical change. A physical change is one in which thereis no change in the molecules which make up a given sub-stance. Turning water into ice or vapor does not constitutea chemical change because the same molecules make up theliquid, solid, and vapor states of water. The only differencebetween ice, steam, and water is this: molecules in ice es-sentially have no freedom. They only vibrate within the crys-tal. The molecules in water are free to move within the limitsof the container, as limited by gravity. The molecules insteam are completely free to move within the container, ifany. They are essentially unaffected by gravity.

In contrast, the pieces of coal which you have burned have becomeashes. They look different; they feel different. This is because a chem-ical change has taken place.

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Chemical change. A chemical change occurs when newmolecules are formed as a result of the change.

When water turns to steam at 212°F, a physical change occurs. Onthe other hand, when propane gas is ignited, it turns to carbon dioxideand water vapor in a chemical change.

Elements

As we pointed out above, there are two prime ingredients—matter andenergy. Scientists further categorize matter into a group of basic sub-stances called elements. An element can be defined as a substancewhich cannot be decomposed chemically into a simpler substance.

The periodic table of all elements is printed in the front of this text.In 1871 Mendeleev, a Russian chemist, documented that the proper-ties and behaviors of elements are repeated in a periodic manner.Mendeleev’s periodic table gets its name from this fact that the prop-erties of elements are repeated periodically in going from left to rightacross a horizontal row or period of elements. The table is arranged sothat an element has properties similar to those of other elements aboveor below it in the table. Elements with similar chemical properties arecalled groups of elements and are contained in vertical columns in theperiodic table. For example, sodium (Na) behaves similarly to potas-sium (K), magnesium (Mg) behaves similarly to calcium (Ca), and chlo-rine (Cl) behaves similarly to bromine (Br).

Of course, nowadays, we are well aware of man’s ability to smashthe atom. In light of this, perhaps a more up-to-date definition mightbe: an element is a substance which cannot be broken into simpler sub-stances without disrupting the atom.

Since the distinguishing characteristic of an element is the fact thatit cannot be broken down into simpler substances, it is obvious theremust be other substances which can be so broken down. Such sub-stances are called compounds and mixtures.

Compounds and Mixtures

How can one distinguish between compounds and mixtures? A com-pound has a definite and unvarying composition.

Water is a typical compound. It is composed of two elements—hydrogen and oxygen—in definite proportions. Regardless of whereone finds water, it always consists of these two elements and always in

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the same proportion. Salt is another common compound. Whether itcomes from a salt mine or is produced in a laboratory, salt is a com-pound of the two elements—sodium and chlorine—in an unvaryingratio.

Water, as a typical compound, also suggests another characteris-tic of the compound, namely a unique “personality” of its own. Al-though made up of hydrogen and oxygen, water is quite differentfrom these two elements both physically and chemically. And so weshould add to our definitions: a compound has well defined charac-teristics of its own, usually entirely different from those of its compo-nent elements.

Further, water freezes at 32°F and boils at 212°F. This indicates an-other characteristic of the compound: a pure compound has a definitefreezing and a definite boiling point.

And finally, water, as a typical compound, is a uniform substanceno matter whether one is considering a drop, a glassful, or a lake. Thus,a compound is homogeneous.

Homogeneous. A substance which is uniform throughoutin its composition is said to be homogeneous. For example,homogenized milk is homogeneous because the globules offat have been broken down to smaller sizes and have beendistributed throughout the liquid.

In sharp contrast, a mixture will vary in the amounts of the ingre-dients it contains. A mixture of sand and salt, for example, may havejust a bit of salt and a large amount of sand. Or it may be a blend of alarge amount of salt and little sand. No exact ratios of substances arenecessary to constitute a mixture. At the same time, the ingredients ina mixture continue to maintain their essential properties. The salt stilltastes salty; the sand continues to be gritty. The properties of the mix-ture are simply the total of the separate properties of the salt and sand.In this salt-sand mixture, the original ingredients could be recoveredthrough some type of mechanical process. And finally, a mixture mayhave varying proportions of its ingredients in different parts of the sam-ple. There may be more salt than sand at the bottom and less at thetop of a mixture. In a word, mixtures are usually heterogeneous.

Heterogeneous. A mixture which is not uniform through-out is said to be heterogeneous. Unhomogenized milk maycontain large globules of fat which can separate and floatto the top of the container.

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Atoms and Molecules

Over the years, scientists have discovered in nature or produced in theirlaboratories more than 100 elements. Each of these basic materials hasits own peculiar atomic structure. The molecule of an element may con-sist of one atom or two or more similar atoms. These atoms are incon-ceivably small. The weight of an oxygen atom, for example, has beendetermined at 0.000,000,000,000,000,000,000,0266 gram. (There are453.6 grams in a pound.) Obviously, ordinary units of measurement areinsufficient for use in determining the weight of an atom or a molecule.

Molecule. Molecules consist of two or more atoms of oneor more elements. (Cl2) is a molecule of chlorine. (NaCl) is amolecule of salt.

Atom. The smallest possible unit of an element. The atom canonly be broken down into the fundamental particles of matter,such as the proton, electron, neutron, antiproton, and meson.

Gram. The gram is the basic unit of weight in the metricsystem. It is equal to 1⁄28 of an ounce (.0022046 pound or15.4324 grains troy). It is meant to be, and virtually is, theweight of distilled water at 4°C contained in a cube whoseedge is one hundredth of a meter (one cubic centimeter).One cubic centimeter equals a milliliter, and this is why aliter of water weighs one million milligrams and one mil-ligram per liter equals one part per million (ppm). 453.6 g= 1 pound.

Over a century ago, chemists developed a purely relative scale foratomic weights. At the time, they assigned a mass of 16 to the oxygenatom. They then expressed the atomic weight for all other elements interms relative to the weight of 16 assigned to the oxygen atom.

Note: As it occurs in nature, oxygen is a mixture of three iso-topes—oxygen-16, oxygen-17, and oxygen-18. Over the years,chemists developed a table of atomic weights, with a value of16 established for this natural mixture. However, because ofdifferent needs, physicists developed a separate table basedupon the single isotope, oxygen-16. Thus, minor but definitevariations existed between the tables. A new table, basedupon carbon-12, was adopted, which is now used by both. Theatomic weights used in this course are taken from this table.

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The hydrogen atom was discovered to weigh approximately 1⁄16 thatof oxygen. Its actual atomic weight is 1.00797.

The weight of molecules can be easily calculated if the formula fora given molecule is known. Simply add the atomic weights of the var-ious atoms which make up the molecule. Here are a few examples ofhow molecular weight is determined:

Substance/Formula Computation of Molecular Weight

Oxygen molecule/O2 2 × atomic weight = 2 × 16 = 32Hydrogen molecule/H2 2 × 1 = 2Water molecule/H2O (2 × 1) + 16 = 18 (approx.)

Very quickly now we will see the bearing these atomic weights haveon this subject of water chemistry.

Over the years, scientists have gathered a great mass of significantinformation on the nature of the atom, both through theoretical studyand testing. As a result of their genius, we now have this picture of theatoms and their structures:

The Nuclear Atom—Protons, Neutrons, and Electrons

The planets circling around the sun in their well-defined orbits illus-trate the structure and activity of the atom. The sun can be comparedto the nucleus of neutrons and protons comprising the center of theatom. The planets surrounding the sun behave in much the same wayas do the electrons in their orbits around the atom nucleus, except incontrast to planets, electrons have properties of both a particle and anenergy wave, i.e., a standing wave is always associated with an electronmoving in its orbit.

The hydrogen atom is the least complex of the atom structures.This atom has one proton which is surrounded by a single electron ran-domly moving or existing throughout its orbital space.

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Hydrogen

Here are the structures of a few more atoms:

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Another element with an extremely simple atomic structure is he-lium. The helium atom has a center consisting of two protons and twoneutrons. Around this core circle the two electrons.

Helium

Carbon

Nitrogen Oxygen

Note here there are six electrons, two in the inner orbit and four inthe outer one. This happens to be the carbon atom. The following areseveral other atoms in which electrons circle the nucleus in two orbits:

Here are a few of the atomic structures in which the electrons cir-cle the nucleus in three orbits:

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As the various atomic structures shown above indicate, the funda-mental particles of matter are the electron, neutron, and proton. Actu-ally, there are more than 30 basic particles known (to name but a few:meson, positrons, anti-protons). It is the electron, neutron, and proton,however, that directly establish the chemical characteristics of the atom.Each of these three particles has its own definite characteristics:

Approximate Mass ElectricalParticle (Its Atomic Weight) Charge

Electron 1⁄1,837 (or negligible) –1Proton 1 +1Neutron 1 0

Protons and electrons determine chemical properties of an atom.Neutrons merely add mass or weight to the atomic unit.

The relation of the electron, proton, and neutron to the basic atomcan best be illustrated with specific examples.

Table 1 at the end of this chapter gives fluorine an atomic numberof 9. This means the fluorine atom contains 9 protons in its nucleus.These 9 protons have a total charge of +9. Since the fluorine atom isneutral, it must contain an equal number of electrons which have sin-gle negative electrical charges.

Now refer back to Table 1 again. It lists the atomic weight of fluo-rine as 19. We know that the 9 protons and 9 electrons have a com-bined weight of just 9 units because the weight of the electrons isnegligible (see above).

Sodium Magnesium Chlorine

The fluorine atom must then have other constituents to make upthe difference in weight between 9 and 19 units. The missing weightis supplied by the 10 neutrons in the fluorine atom. These provide thenecessary atomic weight without affecting the nuclear charge (theatomic number).

Two general observations can be made here:

1. The number of protons in the nucleus of an atom equal thenumber of electrons outside the nucleus.

2. The number of neutrons in the nucleus equal the mass num-ber minus the atomic number.

Study of the atomic weights in Table 1 shown at the end of thischapter indicates that few of the elements have atomic weights whichare whole numbers. Magnesium, for example, has an atomic weight of24.31. Again, chlorine has a weight of 35.453—just about halfway be-tween two whole numbers.

Applying our knowledge of the atomic structure, we might con-clude that the nucleus of a magnesium atom contains 12 protons and121⁄3 neutrons.

Or, in the case of the chlorine atom, its nucleus should contain 17protons and about 181⁄2 neutrons.

However, a neutron is a fundamental particle. It obviously does notoccur as a half or third of a neutron. Study of this perplexing problemled to the discovery of isotopes.

By definition, isotopes are atoms with the same atomic number,consequently the same number of protons and electrons and the samechemical properties—but with differing atomic weights due to differingnumbers of neutrons in the nucleus. All known elements exist in twoor more isotopic forms.

Earlier we stated that magnesium has an atomic weight of 24.31.It is interesting to note that 78.60 percent of magnesium atoms havean atomic weight of 24; 10.11 percent of them have an atomic weightof 25; and 11.29 percent an atomic weight of 26. As you can see, thevast majority of them have the atomic weight of 24.

Similarly, chlorine has an atomic weight of 35.453. Seventy-fivepoint four percent of chlorine atoms have an atomic weight of 35; 24.6percent have an atomic weight of 37.

The isotopes of hydrogen are of such importance they have eachbeen given their own names. The hydrogen isotope with an atomicweight of 3 is referred to as tritium; the isotope with an atomic weightof 2 is deuterium; and the isotope of 1 is called protium.

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Common hydrogen is almost completely made up of atomswith 1 proton and 1 electron. A very few hydrogen atomswith 1 neutron, 1 proton, and 1 electron are mixed in. Theseforms always exist in normal hydrogen gas in the same ra-tios. Because of the greater weight of deuterium, the atomicweight of hydrogen is slightly more than 1.000.

Heavy water is H2O which has hydrogen isotopes in the deuteriumform. Result: pure heavy water has a molecular weight of 20 rather than18 (see Table 2 at the end of this chapter).

Let’s now backtrack for a minute. Earlier we briefly discussed thepatterns of electrons as they orbit around the nuclei of several atoms.Specifically mentioned were hydrogen, helium, carbon, nitrogen, oxy-gen, sodium, magnesium, and chlorine.

In the case of hydrogen and helium, the electrons move in a sin-gle orbit or shell around the nuclei. In contrast, the electrons of car-bon, nitrogen, and oxygen move in two shells or orbits around theirnuclei. And in the sodium, magnesium, and chlorine atoms, there arethree orbital paths.

Study of all the elements show that electrons are arranged in fromone to seven shells around the nucleus; the maximum number of elec-trons in each shell is strictly limited by the laws of physics. The differ-ent number of each shell is designated by the number, n.

The various shells or energy levels surrounding the nucleus (seeTable 2 at the end of this chapter) differ in their capacities to accom-modate electrons. The maximum “population” in the various electronshells has been determined as follows:

n=1: 2 electronsn=2: 8 electronsn=3: 18 electrons

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nucleus

n=1n=2n=3n=4n=5

A definite mathematical progression can be observed here indicat-ing the maximum number of electrons possible in any given electronshell of a normal atom. It can be stated thus:

“The greatest number of electrons that can be accommodatedin a specific electron shell is 2n2 where “n” is the order num-ber of the shell.”

Thus, in the first shell, “n” equals one. One squared still equals one.This squared number times two then equals two.

Moving out to the second shell, “n” now equals two. The squaringof two will give you an answer of four. Now two times four equalseight. And so on.

Under normal conditions, the single hydrogen electron movesaround its nucleus in the n=1 shell. Like all electrons, it obeys the lawof nature in remaining at the lowest possible energy level (in this casethe n=1 shell).

Electrons in Chemical Interaction

There are some definite and most interesting patterns to be notedamong the electrons of the various elements. Those with strong simi-larities of behavior have similar patterns of electrons.

Consider the group of elements known as halogens. Fluorine hastwo electrons in the n=1 shell, seven in the n=2 shell. Chlorine has twoin the n=1 shell, eight in the n=2 shell, and seven in the n=3 shell.Bromine has two, eight, and eighteen respectively in the first threeshells and seven in the n=4 shell. Iodine has two, eight, eighteen, eight-een, and seven. Notice that they all have seven electrons in their outershells.

Halogens. The halogens embrace four elements which areextremely active chemically. Included are fluorine, chlorine,bromine, and iodine. These elements never occur in naturein a free state, but are found in compounds with variousother elements.

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Here are the orbital patterns of electrons in several groups or “fam-ilies” of elements:

Alkali Metalsn=1 n=2 n=3 n=4 n=5 n=6

Lithium 2 1Sodium 2 8 1Potassium 2 8 8 1Rubidium 2 8 18 8 1Cesium 2 8 18 18 8 1

Alkaline Earth Metalsn=1 n=2 n=3 n=4 n=5 n=6

Beryllium 2 2Magnesium 2 8 2Calcium 2 8 8 2Strontium 2 8 18 8 2Barium 2 8 18 18 8 2

Halogensn=1 n=2 n=3 n=4 n=5 n=6

Fluorine 2 7Chlorine 2 8 7Bromine 2 8 18 7Iodine 2 8 18 18 7

Inert Gasesn=1 n=2 n=3 n=4 n=5 n=6

Helium 2Neon 2 8Argon 2 8 8Krypton 2 8 18 8Xenon 2 8 18 18 8

Notice the inert gases. They all possess an outer shell of eight elec-trons. The one exception is helium with two. These elements are alluniquely stable. They show no inclination to combine with other ele-ments in various compounds.

There is a theory that all elements have a chemical “drive” toachieve the inert gas type of structure. Sodium and potassium, for ex-ample, seem to strive to get rid of that one extra electron so that theywould then have eight electrons in their outer shells and so resemblethe inert gases nearest them.

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Note that the fluorine atom is one shy of having eight electrons inits outer orbit. Sodium, on the other hand, has one electron travelingalone in its outer orbit.

If fluorine and sodium are brought into contact under the properconditions, the sodium atom transfers its one excess electron to thefluorine atom which has a space for it.

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After the transfer takes place, there are no longer sodium and flu-orine atoms. Instead there is now the compound: sodium fluoride.

The sodium particle with its new + charge is now termed a sodiumion. The fluorine atom has now become a fluoride ion. Let’s analyzethis a bit further. Remember, atoms are neutral. Their protons (positivecharges) equal their electrons (negative charges). What happens whenthe lone sodium electron in the outer shall transfers to the fluorineatom? The particle that prior to the union was a sodium atom still hasits nuclear charge of +11 (11 protons). But now it has only 10 extra-nuclear electrons (10 negative charges). The result is now a net chargeof +1. (+11 – 10 = +1)

Similarly, the particle that prior to the union was a fluorine atomwith a nuclear charge of +9 now has an extranuclear charge of –10.Hence, it has a total negative charge of –1. (+9 – 10 = –1)

Ions

Virtually all inorganic reactions involved in geological and biologicalsystems are ionic in nature, and ions and their reactions are essentialto plant and animal life. Ions and ion exchange determine, in a largemeasure, the fertility of soils. Yet these same ions can seriously inter-fere with the beneficial use of water by man.

In 1884, the Swedish chemist Arrhenius studied the fact that certainsubstances which dissolve in distilled water make it a good conductor ofelectricity, while other substances do not increase its conductivity.

To explain the mechanism by which these substances conduct elec-trical current, he proposed his theory of ionization in 1887. He sug-gested that when certain substances (later called electrolytes) dissolvein water, they form electrically charged particles called “ions.” Further,because the solution, as a whole, is electrically neutral, he assumed theexistence of two types of ions, one charged positively and the othercharged negatively.

An ion can be defined as an electrically charged atom or group ofatoms in solution. Positively charged ions are called “cations” becausethey migrate to the cathode, or negative electrode, when a current ofelectricity is passed through the solution. Negatively charged ions arecalled “anions” because they migrate to the anode or positive electrode.The word “ion,” derived from the Greek, means “to go” or “to wander.”

When sodium chloride (table salt) is dissolved in water, each mol-ecule that is ionized produces two ions, a sodium ion with a single pos-itive charge and a chloride ion with a single negative charge.

Na+ + Cl– = NaCl

Sodium Ion + Chloride Ion = Sodium Chloride Molecule

Ions exist independently in solution and possess specific proper-ties, which may differ greatly from those of their atoms or molecules.For example, metallic sodium reacts violently with water, producinghydrogen gas and caustic soda, but sodium ions exist calmly in solu-tion. Chlorine gas is poisonous, but chloride ions are not. In fact, bothsodium and chloride ions are essential to life.

Many impurities exist as ions in natural waters. The most commonof these are:

Cations

Calcium . . . . . . . . . . . . . . . Ca++

Magnesium . . . . . . . . . . . . Mg++

Sodium . . . . . . . . . . . . . . . Na+

Iron . . . . . . . . . . . . . . . . . . Fe++

Manganese. . . . . . . . . . . . . Mn++

Anions

Bicarbonate . . . . . . . . . . . . HCO3–

Chloride. . . . . . . . . . . . . . . Cl–

Sulfate . . . . . . . . . . . . . . . . SO4– –

Nitrate . . . . . . . . . . . . . . . . NO3–

Carbonate . . . . . . . . . . . . . CO3– –

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These electrically charged dissolved particles make ordinary natu-ral water a good conductor of electricity. Conversely, pure water has ahigh electrical resistance, and resistance is frequently used as a meas-ure of its purity.

Since many of these ions interfere with the beneficial uses of water,a number of methods have been developed for the reduction or sub-stantial removal of ionic impurities from water.

There are a number of ways in which dissolved ionic impurities inwater may be reduced or substantially removed: (1) distillation, (2) pre-cipitation and separation, (3) ion exchange, and (4) membrane sepa-ration.

Ion. An ion is an electrically charged atom or group ofatoms. The electrical charge of an ion with a single nucleusis due to the gain or loss of one or more electrons. Several ofthese simple ions may combine into groups, which thenhave a charge which is the sum of the charges of the sim-ple ions in the group. The concept of valence or the com-bining power will be explained later. The loss or gain occursduring chemical reactions in which electrons are transferredfrom one atom to another.

Note the change in the name of the fluorine atom to flu-oride ion. Most ions with positive or plus charges bear thesame name as their corresponding atoms. But most nega-tive ions (when derived from single atoms) have the ending-ide added to the root of the word for the correspondingelement. For example, fluoride, bromide, iodide and sulfideions.

Thus, ions have properties quite different from their parent atoms.Sodium is a silver-white alkaline metallic element with a waxlike con-sistency. Fluorine is a corrosive greenish-yellow gas. When combinedthrough chemical reaction, their ions result in a white crystalline sub-stance of much the same appearance as common salt, except that it ismore of a powdered consistency.

In many compounds, the ions are equally attracted to each otherin all directions. While there is no pairing of ions as such, there is onepositive ion for each negative ion. In the case of sodium fluoride, thismeans one sodium ion for each fluoride ion.

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The above sketch gives an indication of one of many types ofarrangements of ions attracted in all directions to ions of the oppositecharge.

From the various examples cited on the last few pages, it is obvi-ous that atoms have varying amounts of electrons in their outer shellswhich can be transferred or shared.

Earlier, we saw that the sodium atom has one electron in its outershell available for transfer. We also saw that magnesium has two elec-trons available for this purpose. And, although it takes an almost im-possible amount of chemical energy to do it, aluminum has threeelectrons available to transfer. In the case of carbon, the four electronsin its outer orbit are shared rather than transferred to other atoms.

Valence

Chemists use the term valence when describing the combining capac-ities of the various atoms or ions. This word valence is defined: “Thevalence of an element in an ionic compound is equal to the numberof electrons that an element loses or gains during the formation of thecompound.”

This combining power is expressed as a small whole number andmay be + or –.

When an atom of an element loses one or more electrons to be-come a positive ion, the element is said to have a positive valence.When an atom picks up one or more electrons to form a negative ion,the element is classified as having a negative valence.

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Going back to our examples, sodium has a valence of +1; magne-sium +2; aluminum +3. Chlorine has a valence of –1, as does fluorine.

NOTE: In the case of elements in a covalent compound (wherepairs of electrons are shared), the valence is equal to the number ofshared pairs.

There are, in addition, a group of metals—the transition metals—which may transfer varying numbers of electrons under different cir-cumstances. These elements display what is referred to as variablevalences.

Iron can lose either two or three electrons to form either doubly ortriply charged iron ions (Fe++ or Fe+++). Copper, in some compounds,exists as +1; in others it has a double ++ charge.

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TABLE 1ELEMENTS—BASED ON CARBON-12

(1961 International Atomic Weights)

40

ATOMIC ATOMICELEMENT SYMBOL NUMBER WEIGHT

Aluminum Al 13 26.98Antimony Sb 51 121.75Argon A 18 39.948Arsenic As 38 74.92Barium Ba 56 137.34Beryllium Be 4 9.01Bismuth Bi 83 208.980Boron B 5 10.811Bromine Br 35 79.909Cadmium Cd 48 112.40Calcium Ca 20 40.08Carbon C 6 12.011Cerium Ce 58 140.12Cesium Cs 55 132.905Chlorine Cl 17 35.453Chromium Cr 24 51.996Cobalt Co 27 58.933Columbium Cb 41 92.906Copper Cu 29 63.54Dysprosium Dy 66 162.50Erbium Er 68 167.26Europium Eu 63 151.96Fluorine F 9 18.998Gadolinium Gd 64 157.25Gallium Ga 31 69.72Germanium Ge 32 72.59Gold Au 79 196.967Hafnium Hf 72 178.49Helium He 2 4.0026Holmium Ho 67 164.930Hydrogen H 1 1.00797Indium In 49 114.82Iodine I 53 126.90Iridium Ir 77 192.2Iron Fe 26 55.847Krypton Kr 36 83.80Lanthanum La 57 138.91Lead Pb 82 207.19Lithium Li 3 6.939Lutecium Lu 71 174.97Manganese Mn 25 54.938Magnesium Mg 12 24.312Mercury Hg 80 200.59

ATOMIC ATOMICELEMENT SYMBOL NUMBER WEIGHT

Molybdenum Mo 42 95.94Neodymium Nd 60 144.24Neon Ne 10 20.183Nickel Ni 28 58.71Nitrogen N 7 14.0067Osmium Os 76 190.2Oxygen O 8 15.9994Palladium Pd 46 106.4Phosphorus P 15 30.97Platinum Pt 78 195.09Potassium K 19 39.102Praseodymium Pr 59 140.91Protactinium Pa 91 231Radium Ra 88 226Radon Rn 86 222Rhenium Re 75 186.2Rhodium Rh 45 102.91Rubidium Rb 37 85.47Ruthenium Ru 44 101.07Samarium Sm 62 150.35Scandium Sc 21 44.956Selenium Se 34 78.96Silicon Si 14 28.086Silver Ag 47 107.870Sodium Na 11 22.9898Strontium Sr 38 87.62Sulfur S 16 32.064Tantalum Ta 73 180.948Tellurium Te 52 127.60Terbium Tb 65 158.924Thallium TI 81 204.37Thorium Th 90 232.038Thulium Tm 69 168.934Tin Sn 50 118.69Titanium Ti 22 47.90Tungsten W 74 183.85Uranium U 92 238.03Vanadium V 23 50.942Xenon Xe 54 131.30Ytterbium Yb 70 173.04Yttrium Y 39 88.905Zinc Zn 30 65.37Zirconium Zr 40 91.22

TABLE 218 COMMON CHEMICAL ELEMENTS

(This table shows the atomic number, atomic weight, andthe orbits of their electrons)

THE NUMBER OF ELECTRONS IN VARIOUS ORBITS

Name of Atomic Atomic The The Thethe Element Number Weight n=1 Shell n=2 Shell n=3 Shell

Hydrogen 1 1.00797 1

Helium 2 4.0026 2

Lithium 3 6.939 2 1

Beryllium 4 9.012 2 2

Boron 5 10.811 2 3

Carbon 6 12.011 2 4

Nitrogen 7 14.007 2 5

Oxygen 8 15.9994 2 6

Fluorine 9 18.9984 2 7

Neon 10 20.183 2 8

Sodium 11 22.9898 2 8 1

Magnesium 12 24.312 2 8 2

Aluminum 13 26.98 2 8 3

Silicon 14 28.086 2 8 4

Phosphorus 15 30.97 2 8 5

Sulfur 16 32.064 2 8 6

Chlorine 17 35.453 2 8 7

Argon 18 39.948 2 8 8

41

TABLE 3VALENCES OF SOME COMMON ELEMENTS

METALS

Fixed Valences Variable Valences

Sodium +1 Mercury +1 or +2Potassium +1 Copper +1 or +2Silver +1 Iron +2 or +3Magnesium +2 Manganese +2, 3, 4, 6 or 7Zinc +2Calcium +2Aluminum +3

NONMETALS

These have the following fixed valences in electrovalent compounds:

Chlorine –1Fluorine –1Bromine –1Sulfur –2

(These elements have a variety of valences in covalent compounds.)

42

43

Lesson 3Introduction

In Lesson No. 2 we discussed some basic facts about chemistry in an ef-fort to understand what is involved in the chemical actions and reactionsof molecules, atoms, ions, protons, neutrons and electrons.

Now let’s continue the discussion further with the thought of apply-ing this general knowledge to the specifics of water chemistry.

Water Chemistry

Water is an extremely stable compound. It does not break down underheat into oxygen and hydrogen molecules to any large extent. At atemperature of 3,600°F, less than two percent of water molecules dis-sociate into hydrogen and oxygen.

In addition, water ionizes to H+ and OH– ions very slightly. In apure water, only one water molecule out of 550,000,000 separatesinto its ions. Note: The relative concentrations of the OH– and H+ arevery important in water treatment and will be considered later in thechapter.

Thus as an extremely stable compound, water neither decomposesnor ionizes readily.

Another important fact about water is that it is not a good oxidiz-ing agent. Since water contains 88.81 percent oxygen, one might ex-pect it to be a good oxidizing agent. Due to its stability, however, waterdoes not release its oxygen except to extremely powerful oxidizingagents.

Oxidation and Reduction

Precisely what is meant by the terms “oxidizing agent” together withthe related term “reducing agent”?

Originally, the term oxidation signified a chemical reaction in-volving the addition of oxygen to a compound. In a similar way, theremoval of oxygen was defined as reduction.

Now the meaning of both terms has been broadened due to in-creased knowledge of the atom, its structure, and reactions.

In the broad sense, oxidation of an atom or ion refers to an in-crease in positive valence or a decrease in negative valence. Reductionhas exactly the opposite meaning—a decrease in positive valence or anincrease in negative valence.

The following formula illustrates this:

Cl2 + 2Fe++ → 2Fe+++ + 2Cl–

Elemental Chlorine Molecule Plus 2 Ferrous Iron Ions Reacts toProduce 2 Ferric Iron Ions Plus 2 Chloride Ions

What takes place? The elemental chlorine molecule is reduced (va-lence zero to –1) and the two ferrous ions are oxidized (valence +2 to +3).

44

In this chemical reaction each iron ion loses 1 electron. It now hasa valence of +3. As the substance that loses electrons, it is referred toas a reducing agent.

The elemental chlorine yields two chloride ions in this formula.These carry a negative charge and have a valence of –1. As the sub-stance that gains electrons, it is called the oxidizing agent.

Oxidation and reduction always occur together and in equalamounts.

Oxidation Valence Reduction

As this chart graphically illustrates, oxidation occurs whenthere is an increase in positive valence. As the positive va-lence decreases, or the negative valence increases, reduc-tion occurs.

A substance that loses electrons is oxidized and is referred to as areducing agent; a substance that receives electrons is reduced and iscalled an oxidizing agent. Oxidizing agents, such as the halogens chlo-rine, fluorine, and bromine, take electrons away from reducing agents,such as metals, that give up electrons.

Here are a few examples of oxidation and reduction that occur inwater:

1. Fe++ –e → Fe+++

Ferrous Less 1 Reacts to FerricIron Electron Produce Iron

2. S– – –2e → S°Sulfide Less 2 Reacts to ElementalIon Electrons Produce Sulfur

45

(In

crea

se)

(Decrease)

++++++

––––––

+++++

+

–––

–––

In Lesson No. 2 there were numerous sketches of electron patternsaround the nuclei of some elements and a few common compounds.We saw how the removal of an electron from an atom gives it a singlepositive charge. Likewise, the addition of an electron to an atom pro-vides a single negative charge. In all molecular structures, the plus orpositive charges must equal the minus or negative charges.

For a moment, let’s review the chemical structure of water. The hy-drogen atom, you will recall, looks this way—one proton with a singleelectron circling it. Note: For simplicity, electron pathways are dia-grammed in this text as being discrete circular pathways; however, theactual space and volume occupied by electrons are depicted by math-ematical models of all the space within each shell.

46

Similarly, elemental hydrogen gas (H2) contains two hydrogenatoms each sharing its one electron.

1 P

The oxygen atom has 8 protons and 8 neutrons in its nucleus plus 8electrons.

1 P 1 P

8 p8 n

And again elemental oxygen gas (O2) is made up of two oxygen atomssharing 4 electrons.

8 p8 n

8 p8 n

Bear in mind that hydrogen and oxygen atoms, as such, seldomoccur in their atomic states. But molecules composed of two atomsapiece of these gases do occur regularly.

When chemically combined, two hydrogen atoms and one oxygenatom produce the electrically neutral water molecule.

47

1 P 8 p8 n 1 P

1 P 8 p8 n 1 P

Water, you will recall, ionizes slightly. When this occurs, it forms:

(+) (–)

1 P8 p8 n 1 P

(+)(–)

hydrogen ion and hydroxide ion

When ionization occurs, one of the hydrogen ions breaks away. Itnow has only a proton, no electron. It carries a positive charge.

The remaining ion contains both hydrogen and oxygen. It is calleda hydroxide ion. Note that there are 8 protons in the oxygen portionand one in the hydrogen portion of this ion. At the same time, thereare 10 electrons (8 contributed by the oxygen ion and 1 each from thetwo hydrogen ions). This hydroxide ion contains one more electronthan proton. It, therefore, carries a negative charge.

hydroxide ion hydrogen ion(negative charge) (positive charge)

This can be diagrammed more simply:

48

+ Hydrogen + Hydrogen–

–

Oxygen

hydroxide ion hydrogen ion(negative charge) (positive charge)

Obviously, the left block is out of balance. The hydroxide ion withits negative charge and the hydrogen ion with its positive charge read-ily combine as the water molecule.

+ Hydrogen

+ Hydrogen

–

–

Oxygen

Hydroxide. The name given to the ion formed of an oxy-gen and a hydrogen atom. It has the formula OH–. Thiscomplex ion goes through many chemical reactions as aunit. In general, it performs as though it were a singlecharged atom.

Up to this point, we have used the terms positive and negative elec-trostatic charge to indicate the force that holds ions together.

In water chemistry, the terms cation (positive) and anion (nega-tive) are used. These terms were coined by Michael Faraday, a greatnineteenth century scientist. He applied them to the substances whichappeared at the cathodes and anodes during his experiments in elec-trolysis. (The cathode is the negative terminal that attracts the posi-tively charged cations; the anode is the positive terminal that attractsnegatively charged anions.)

A Word About Ions

In the process of solution, many molecules dissociate into two or moreions which are theoretically free to move about as independent particles.

An ion differs from an atom or a molecule in that it carries an elec-tric charge.

Ions are of two kinds: one (the cation) is electrically positive, andthe other (the anion) negative. In water, the sum of the positivecharges equals the sum of the negative charges; so the solution remainselectrically neutral.

The block building technique provides a graphic way to illustratehow the compounds in water chemistry are formed.

Hydrogen, you will remember, has a valence of +1. Any atom ormolecule that combines with hydrogen on a one-to-one basis has a va-lence of –1.

49

+ Hydrogen(H)

–Chloride(Cl)

Sodium +(Na)

– Chloride(Cl)

The resulting compound is hydrogen chloride or hydrochloric acid.Sodium is another cation with a valence of +1. It can combine with

the chloride ion on a one-to-one basis. The chloride ion possesses thecorrect –1 valence for uniting with sodium.

And, of course, the resulting compound is sodium chloride (com-mon salt).

Anion (pronounced an-i-on). A negatively charged ion insolution. It may be a single atom with a charge, such aschloride(Cl–), or a group of atoms with a charge such as sul-fate (SO4

– –).

Cation (pronounced cat-i-on). A positively charged ion insolution. It may be a single atom with a charge, such as cal-cium (Ca++) or a group of atoms such as (NH4

+).

Sodium hydroxide is a compound that results from the combiningof a simple ion and a complex ion.

Na + OH –

Simple Ion Complex Ion

Combined it becomes

50

Na +

H +

–

–

O

or even more simply

Na + – OH

One further example:

The bicarbonate ion contains one part hydrogen, one part carbon,and three parts oxygen.

As we have seen, the + and – charges must be equal to build a com-pound.

Magnesium has a valence of +2. This cation with its two positivecharges can be diagrammed thus:

H++ +

+

–

–

–– –

–+O C O = – HCO3

O

+

+

Magnesium(Mg)

– Hydroxide(OH)

+

+

Magnesium(Mg)

If a single hydroxide ion were to be combined with the magnesiumion, the result would be:

The resulting compound is a magnesium hydroxide molecule.

An anion with a valence of –2 will combine with a magnesium ionto form a compound. Thus…magnesium carbonate:

51

Obviously they are mismatched. Since the plus and minus valencecharges must match, two hydroxide ions are required.

– Hydroxide(OH)

– Hydroxide(OH)

+

+

Magnesium(Mg)

Carbonate(CO3)

+

+

–

–

Magnesium(Mg)

Aluminum has a valence of +3. It appears thus…

+

+

+

Aluminum(Al)

– OH

– OH

+

+

+

Aluminum(Al)

– OH

If combined with hydroxide, three of the latter molecules areneeded to produce the chemical compound aluminum hydroxide:

There are a number of cations and anions which are to be foundin water chemistry. These include:

Common AtomicElement Symbol Valences Weight

Aluminum Al +3 26.98Barium Ba +2 137.34Calcium Ca +2 40.08Carbon C +4 [–4,–2,+2] 12.011Chlorine C1 –1 [+1,+3,+5,+7] 35.453Copper Cu +1,+2 63.54Fluorine F –1 18.998Hydrogen H +1 1.008Iodine I –1 126.90Iron Fe +2,+3 55.85Magnesium Mg +2 24.312Manganese Mn +2,+4 [+3,+6,+7] 54.938Nitrogen N +3,+5 [+1,+2,+4] 14.0067Oxygen 0 –2 [–1] 15.9994Phosphorus P +3,+5 [+1,+2,+4] 30.974Potassium K +1 39.102Silicon Si +4 28.086Silver Ag +1 107.870Sodium Na +1 22.9898Sulfur S –2 [+4,+6] 32.064Zinc Zn +2 65.37

[Brackets indicate less common valences.]

Anions Symbol

Acetate C2H3O2–

Bicarbonate HCO3–

Carbonate CO3––

Hypochlorite ClO–

Nitrate NO3–

Nitrite NO2–

Hydroxide OH–

Permanganate MnO4–

Phosphate PO4–––

Sulfite SO3––

Sulfate SO4––

Thiosulfate S2O3––

52

The common cations and anions that form the basic units in waterchemistry are as follows:

53

+

+

+

Al

+

+

+

Fe

–

–

–

PO4

+

+

Ca

+

+

Mg

–

–

CO3

–

–

SO4

+

+

Fe

H +

Na +

CATIONS

Hydrogen

Sodium

Calcium

Magnesium

Aluminum

FerrousIron

FerricIron

Forms waterwith hydroxideand acids withother anions

Forms a basewith hydroxideand soluble saltswith other anions

These formhardnesscompounds,bases withhydroxide,and salts withother anions

– OH

– Cl

– HCO3

– NO3

ANIONS

Hydroxide

Chloride

Bicarbonate

Carbonate

Sulfate

Nitrate

Phosphate

Formswater withhydrogenand a basewith metals

These formhardnesscompounds,bases withhydroxide,and salts withother anions

Anionsthatformacidswith hydrogenand saltswith metalions

When various cations and anions found in water combine, the re-sulting compounds are classified as bases, acids, and salts. Some of theserelationships are indicated above.

What are these acids, bases, and salts, and how may they be defined?

Acids

Acids can be defined as compounds that release hydrogen ions (H+) ina solution. All acids contain hydrogen. In general, they have (1) a moreor less sour taste; (2) they change the color of indicators (acids turn lit-mus paper red); and (3) they react with bases to form a salt and water.

However, in addition to their properties held in common, acidshave other properties which may vary widely. These specific propertiesof each acid are due to the anion present and to the undissociated mol-ecules. Thus, the molecules of various acids are capable of releasing dif-ferent amounts of free hydrogen ions in solution. Examples of therelease of hydrogen ions (H+) in a solution:

Hydrochloric Acid (Strong Acid):

HCl → H+ + Cl–

Acetic Acid (Weak Acid):

C2H4O2→← H+ + C2H3O2

–

Acetic Acid (vinegar) is weakly ionized and only releasessmall amounts of free hydrogen ions into solution.

Strong acids and bases separate into their ions and stay separated.This is indicated by a single arrow pointing in one direction. Weakacids and bases are in a continuous process of ionization, but the freeions are also continuously recombining to form molecules. At anytime, only a small portion of the acid or base is present as ions. Thisequilibrium process is indicated by arrows pointing in both directions.

Bases

Bases are substances which can release hydroxyl (OH–) ions. Sodiumhydroxide (NaOH) and ammonium hydroxide (NH4OH) are examples.These hydroxides ionize as follows:

54

Sodium Hydroxide:

NaOH → Na+ + OH–

Ammonium Hydroxide:

NH4OH →← NH4+ + OH–

There are a great many bases. In general, their solutions (1) tastebitter rather than sour; (2) feel slippery; and (3) reverse the colorchanges produced by the acids in indicators. For example, they turnlitmus paper blue.

Ammonium Hydroxide is a weak base as it only releasessmall amounts of hydroxide ion (OH–) into solution.

As in the case of acids, each base has individual properties. Theseare due in each compound to the cation present and the nonionizedmolecules of the base. As in the case of acids, bases demonstrate vary-ing degrees of ionization. Those that ionize to a large extent are calledstrong. Those that ionize only slightly are weak.

Salts

Salts are formed from a metallic ion (positively charged cation) and anon-metallic ion (negatively charged anion).

Salts are of three types. They are classified as normal, acid, andbasic.

A normal salt is a compound produced by the union of the cationsof any base and the anions of any acid.

Normal salts possess no one ion characteristic, although they arenearly all strongly ionized. In regard to their solubility, they range be-tween wide limits. Some dissolve less than one milligram per liter.Others dissolve in much less than their own weight of water, e.g.sodium chloride.

HCL + NaOH → NaCl + H2O

An acid salt is one composed of both metallic and hydrogencations plus anions of an acid, e.g. sodium bisulfate and sodium bicar-bonate.

NaOH + H2SO4 → NaHSO4 + H2O

NaOH + H2CO3 → NaHCO3 + H2O

55

A basic salt is one composed of metallic cations, together with hydroxylanions of a base and anions of an acid, e.g. basic carbon carbonate.

2Cu(OH)2 + H2CO3 → Cu2(OH)2CO3 + 2H2O

Turning again to our building blocks, we can diagram many of thecommon compounds occurring in the field of water chemistry. Theseare shown below.

Hydrogen Compounds

56

H + – OH

– HCO3

– Cl

H2OWater

H +

HClHydrochloric acid.

It may be added in thetreatment of water.

H +

H2CO3Carbonic acid.

This boosts the solventaction of water.

H2SO4Sulfuric acid.

This may be used inwater treatment.

H2CO3When heated,

carbonic acid breaksdown into water and

carbon dioxide.

H3PO4Phosphoric acid.

This may also be usedin water treatment.

CO3

H +

H + –

–

SO4

H +

H + H +

H +

H +

–

–

–

– PO4

–

Except for water, all the above hydrogen compounds are acids. Coldwater absorbs carbon dioxide to produce carbonic acid. In the aboveblocks, carbonic acid is diagrammed as both a bicarbonate and a car-bonate of hydrogen. In carbonic acid, neither the bicarbonate nor car-bonate is readily ionized. When heated, the carbonic acid molecule maybreak down by evolving the carbon dioxide and leaving plain water.

Calcium Compounds

57

Ca(OH)2Calcium hydroxide.A compound used

in lime-sodasoftening.

CaCl2Calcium chloride.

A constituentof hardness.

CaCO3Calcium carbonate

forms sludge orsoft scale.

Ca3(PO4)2Tricalcium phosphate.

This is the sludgeformed by the action

of phosphate oncalcium hardness.

– OH

– OH

+

+

Ca

– Cl

– Cl

+

+

Ca

+

+

–

–

Ca

– HCO3

– HCO3

CO3

CaSO4Calcium sulfate

forms hard scale.

+

+

–

–

Ca SO4

+

+

Ca

Ca(HCO3)2Calcium bicarbonate.

Heat drives off thecarbon dioxide and precipitates

calcium carbonate.

–

–

–

PO4

–

–

–

PO4

+

+

Ca

+

+

Ca

+

+

Ca

These calcium compounds are frequently found in raw water or areadded in the treatment of water problems.

At one extreme of the solubility scale is calcium chloride which isvery soluble. Calcium sulfate is much less soluble, and calcium car-bonate is very slightly soluble in water (less than 20 ppm).

When water containing carbon dioxide (CO2) comes in contactwith insoluble deposits of calcium carbonate, a chemical reaction oc-curs. This results in calcium bicarbonate formations. These and themagnesium bicarbonates are extremely soluble in water and are usu-ally the principal hardness compounds found in raw water.

Magnesium Compounds

58

Mg(OH)2Magnesium hydroxide.

This is insolubleprecipitate formed inlime-soda treatment.

MgCl2Magnesium chloride.A soluble, corrosive

constituent ofhard water.

MgCO3Magnesium carbonate

is slightly soluble in water (less than 100 ppm).

– OH

– OH

+

+

Mg

– Cl

– Cl

+

+

Mg

+

+

–

–

Mg

– HCO3

– HCO3

CO3

MgSO4Magnesium sulfate

is extremely soluble.

+

+

–

–

Mg SO4

Mg(NO3)Magnesium nitrate

is a soluble compound.

+

+

Mg

– NO3

– NO3

+

+

Mg

Mg(HCO3)2Magnesium bicarbonate.

Heat drives off thecarbon dioxide to form magnesium carbonate.

Like the calcium compounds, the magnesium compounds rangefrom extremely soluble to insoluble states. The magnesium carbonatesare moderately soluble in pure water, up to 100 ppm.

Water containing more than four or five grains of the carbonateform of magnesium is considered unusual.

Hardness due to the dissolved calcium and magnesium bicarbon-ates in water is often referred to as temporary hardness due to the factthat these bicarbonates may revert back to an “insoluble” form or scalewhen water is heated.

These hardness minerals range from a low of one grain per gallonin some supplies to over 350 grains per gallon in others. Most waterscontain amounts ranging from three to 50 grains.

Sodium Compounds

59

Na + – OH

NaOHSodium hydroxide.

May be used inwater treatment.

Na2CO3Sodium carbonate

(soda ash) is used toneutralize acid waterand is also used in

lime-soda ashwater treatment.

Na + – Cl

NaClSodium chloride.This is used to regenerate ion

exchange softeners.

Na +

NaHCO3Sodium bicarbonate.

Frequently found in both hard and

softened waters, this contributes to alkalinity.

CO3

Na +

Na + –

–

Na2SO3Sodium sulfite.This compound

chemically absorbsoxygen in water.It is a chemicalreducing agent.

SO3

Na +

Na + –

–

Na2SO4Sodium sulfate.

This is sometimespresent in water, andis also a byproductof water treatment.

SO4

Na +

Na + –

–

– HCO3

Most of the sodium compounds possess the characteristic of highsolubility. They are rarely scale forming.

In addition to the sodium compounds found in raw water, severalof these compounds are extremely important in the treatment of water.

Sodium sulfate appears as an end product when oxygen is absorbedchemically with sodium sulfite.

A water analysis will show up many facts about a water. The min-eral and gaseous compounds in a supply can be quite numerous. Somehave already been mentioned, but we have, by no means, exhaustedthe list. Let’s briefly consider some of the other water contaminants.Each will be covered more thoroughly in further lessons.

Alkalinity (OH–, CO3– –, and HCO3

–)

60

–

–

CO3

OH – HCO3 –

Hydroxide

Carbonate

Bicarbonate

Alkalinity is caused by the presence of bicarbonates, carbonates,and hydroxides in water.

Excessive alkalinity produces a “soda” taste. It has a drying effecton the skin.

Carbon Dioxide (CO2)

This compound is present in most water supplies in concentrationsranging from zero to about 50 parts per million.

Surface waters from lakes and rivers usually contain more carbondioxide than do well waters. In some groundwater supplies, however,the amounts are so great that it bubbles out when pressure is released.

Carbon dioxide combines with water to form carbonic acid. Thisweak acid accelerates corrosion, particularly when water is heated.

Excessive concentrations of carbon dioxide are generally indicatedby low pH values.

+

+

+

+

C

–

–

O

–

–

O

Chloride (Cl–)

61

Cl –

F –

S

H +

H + –

–

In small amounts, chlorides are present in almost all natural watersupplies. Where concentrations are high, chlorides may be objection-able either because of taste, corrosive tendencies, or because of theiradverse effects in the softening process.

Fluoride (F–)

These can be detrimental or beneficial. It all depends on their con-centration in water.

In some areas, water contains up to 1 ppm of fluorides. Where thisoccurs, the fluorides act as a preventative against tooth decay. But flu-orides in excess of 2 ppm can also cause a dark brown stain or give achalky white appearance to teeth. Skeletal fluorosis, a serious cripplingbone disorder resembling osteoporosis, can develop from many years ofexposure to drinking water with more than 4 ppm of fluoride.

When fluorides are present in excessive concentrations, it is nec-essary to remove them.

Hydrogen Sulfide (H2S)

Waters containing hydrogen sulfide are commonly called “sulfurwaters.”

In high concentrations, this gas is flammable and poisonous. It iscorrosive to most metals, and it tarnishes silver readily.

The distinctive odor associated with hydrogen sulfide is apparentin concentrations as low as 0.5 ppm.

Iron (Fe++, Fe+++)

62

+

+

+

Fe

+

+

Fe

Ferrous

Ferric

+

+

+

Mn

+

+

Mn

Manganic

Manganous

Iron is found in groundwater supplies in varying amounts and invarious forms.

Most iron-bearing water contains less than 5 ppm of iron. Occa-sionally concentrations up to 60 ppm and higher are found.

Iron occurs, most frequently, in groundwater supplies as ferrous bi-carbonate. Sometimes it is found as ferrous sulfide and in acid minewaters as ferrous sulfate.

Iron is highly objectionable. It imparts a metallic taste to water.Much worse, it stains practically everything with which it comes in con-tact. Even in such small proportions as 0.3 ppm, iron causes staining.

When drawn, iron-bearing water is usually clear and colorless withdissolved ferrous (+2 valence) iron. Exposure to air, however, causes itto oxidize to ferric (+3 valence) iron, to cloud up and deposit a yel-lowish or reddish-brown sediment of ferric hydroxide. It is this sub-stance which causes staining.

Manganese (Mn++, Mn+++)

Manganese could be described as a “companionable” element.Where it is found in water, it is usually accompanied by iron. It is gen-erally rare, however.

Manganese produces objectionable dark brown or black stains evenwhere concentrations are as low as those above 0.05 ppm.

Manganese deposits will collect in plumbing, and tap water willfrequently contain a black sediment and turbidity.

Dissolved manganese oxidizes more slowly than iron. For this rea-son, it is more difficult to remove from water by oxidation and filtra-tion. Ion exchange is better than oxidation for manganese removal.

It usually occurs as manganous bicarbonate or manganous hy-droxide and less frequently as manganous sulfate.

Nitrite (NO2–) and Nitrate (NO3

–)

63

NO2 – NO3

–

NitrateNitrite

While some soils have natural nitrate content, nitrates frequentlyindicate polluted water due to the presence of organic matter.

Although most nitrate polluted water is found in shallow wells,deep wells may also be affected.

Concentrations as low as 10 to 20 ppm of nitrate nitrogen havebeen the cause of infants becoming ill and have even led to theirdeaths.

Silica (SiO2) (Silicon Dioxide)

This hard, glassy mineral substance is absorbed in water as it flowsover rock or percolates through rock formations.

Silica has little effect on the water used by the average family. Inthe industrial field, however, it does produce extremely hard scale inboilers and on turbine blades as an after-effect of steam.

Silica in amounts ranging from less than 1 ppm to over 100 ppmis found in all natural water supplies.

Sodium (Na+)

SiO2

Na +

Sodium is found to a greater or lesser degree in all water suppliesdepending on local conditions.

In low concentrations, sodium salts have little or no effect. In largeamounts, they tend to increase the corrosive action of water.

Sodium can be removed from water only through such processesas deionization, distillation, or reverse osmosis.

Quite similar are potassium salts. Both salts are extremely soluble inwater.

Sulfate (SO4– –)

64

SO4

–

–

CH4

Sulfates are found in almost all natural water supplies. The amountof sulfates varies according to soil characteristics in different areas.

Some industrial wastes are high in sulfates and increase the sulfatecontent of natural water supplies.

High sulfate concentrations pose special problems in the condi-tioning of water. High sulfates generally mean extreme hardness, highsodium salt concentrations, and high acidity.

Sulfates give water a medicinal taste and have a pronounced laxa-tive effect on those not accustomed to them.

Methane (Marsh Gas) (CH4)

This is occasionally found in well waters. In sufficient amounts itconstitutes a fire and explosion hazard. It is readily identified by itsflammable nature.

In addition to the various gases and minerals in water, there aremany types of bacteria and microorganisms. These include diatoms,molds, bacterial slimes, algae, iron and manganese bacteria, etc. Theirtreatment will be considered in Lessons 5 and 8.

pH

No study of water chemistry would be complete without reference topH. pH stands for “potential of hydrogen” and is a term used to meas-ure the intensity of the acidity or the alkalinity of water.

The pH measurement scale goes from 0 to 14, with 7 the neutralpoint.

65

109876543210 11 12 13 14

ACIDIC NEUTRAL BASIC (ALKALINE)

pH figures larger than 7 indicate alkaline solutions with the inten-sity of alkalinity increasing as the number becomes larger.

pH figures lower than 7 indicate acid solutions with the intensityof acidity increasing as the numbers get smaller.

pH. The hydrogen ion concentration or the potential ofhydrogen in water. The dictionary defines pH as “the sym-bol for the logarithm of the reciprocal of the hydrogen ionconcentration expressed in gram atoms per liter of a solu-tion, and used to indicate acidity or alkalinity.”

While pH is often defined as the measure of the acidity or alkalin-ity of a substance, the meaning should be more restrictive. Actually itapplies to a measurement of the intensity factors rather than the quan-tity factors of various degrees of acidity or alkalinity.

The number of free hydrogen ions in N1 hydrochloric acid is 75times the number of free hydrogen ions in N1 acetic acid. Hydrochlo-ric acid is a strong acid while acetic acid is weak.

Quantitatively they are of the same strength because as much al-kali is required to neutralize the one as the other.

There are two important things to bear in mind when considering pH:

1. It is always an intensity measure, not one of quantity. Inmuch the same way a thermometer will tell how cold aroom is but not how much warm air is necessary to heat it.

2. It is an exponential function. pH 10 is ten times as alkalineas pH 9 and 100 times as alkaline as pH 8. Similarly, a pH 2is 100 times more acid than pH 4 and 1,000 times more acidthan pH 5.

N1—a normal solution, (or 1 N) contains the equivalentweight of a compound in one liter of solution. For example,sodium chloride has an equivalent weight of 58.443. A so-lution containing 58.443 grams of sodium chloride per literis said to be “one normal,” or 1 N NaCl. A solution contain-ing 5.8443 grams per liter is 0.10N NaCl or N/10 NaCl.

(NaCl) indicates the normality of a solution of sodiumchloride. This terminology is commonly used in equationsto avoid excessive writing.

Chemists have discovered that the concentration of OH–

in water times the concentration of H+ always equals a con-stant, 1 × 10–14 or 1. As an equation, this is expressed:

(OH–) × (H+) = 1 × 10–14

As you know, the OH– makes a solution alkaline, and theH+ makes it acid. When they are equal in concentration, thesolution is neither acid or alkaline but neutral:

(OH–) = (H+)

as (OH–)(H+) = �101–14�

(OH–) = �101

–7� and (H+) = �101

–7�

Thus �10

1–7� × �

101

–7� = �101–14�

As the use of such figures as �101

–7� are difficult to use, chemistshave developed another short-cut and established the pHterminology. The pH number basically is the exponent inthe denominator of the fraction which expressed the con-centration of H+. In the examples, pH 6 or pH 5 means (H+)= �10

1–6� or �10

1–5�.

NOTE: The closer pH approaches 0, the higher the concen-tration of free (H+), the greater the acidity, and the lowerthe (OH–). Yet the solution still does contain some OH–.

66

The following chart will show the relative intensity of the acidity orbasicity of water.

ACIDITY BASICITY0 100,000,000,000,000 11.0 10,000,000,000,000 102.0 1,000,000,000,000 100

Acidic 3.0 100,000,000,000 1,0004.0 10,000,000,000 10,0005.0 1,000,000,000 100,0006.0 100,000,000 1,000,000

Neutral —- 7.0 10,000,000 10,000,0008.0 1,000,000 100,000,0009.0 100,000 1,000,000,000

10.0 10,000 10,000,000,000Basic 11.0 1,000 100,000,000,000

12.0 100 1,000,000,000,00013.0 10 10,000,000,000,00014.0 1 100,000,000,000,000

Summary

In this lesson, we have examined the subject of water and its chem-istry even more closely. We have found that water is stable even at ex-tremely low temperatures, but it is not a good oxidizing agent.

The term “oxidation” and the related term “reduction” are impor-tant. Originally, oxidation meant the addition of oxygen to a com-pound; reduction, the removal of oxygen from a compound.

In a broad sense, oxidation now means an increase in positive va-lence or a decrease in negative valence. Reduction which goes on si-multaneously refers to a decrease in positive valence and an increasein negative valence.

When chemically combined, two hydrogen atoms plus one oxy-gen atom produce an electrically neutral water molecule. When waterionizes—and it does so just slightly—it forms a hydrogen ion and a hy-droxide ion. Again, the hydroxide ion (negatively charged) and the hy-drogen ion (positively charged) readily recombine as a water molecule.Actually, the positive hydrogen ion is a cation; the negative hydroxideion is an anion. These can combine together or with other cations andanions in various compounds. Complex ions also combine in the samemanner.

67

A simple method of diagramming the combining of various ionsis the building block technique. This technique is used throughout thiscourse in water treatment.

The compounds resulting from this combining of cations and an-ions in water are classified as acids, salts, and bases.

Briefly, acids can be defined as compounds which release hydro-gen ions in solution. All acids contain hydrogen.

Bases are substances which can release hydroxide (OH–) ions. Saltsare substances containing both metallic (positively charged) ions andnonmetallic (negatively charged) ions. Salts are classified as normal,acid, and basic.

Among the various contaminants to be found in water are hard-ness compounds, chlorides, sulfates, fluorides, hydrogen sulfide, iron,manganese, sodium, silica, and others. These will be discussed, in de-tail, in succeeding lessons of this course.

68

69

Lesson 4Introduction

Invisible hardness minerals make water troublesome and difficult to use.When water contains these hardness minerals, it can plague both indus-try and the homemaker.

So frequently is hardness found in water, some people are apt to re-late the water conditioning industry solely to this problem. In a way, thisassociation of hardness and water conditioning on their part is a tributeto the industry for engineering products capable of transforming unsatis-factory, difficult-to-use waters into fully satisfactory, usable supplies.

Hard Water—Ranges and Problems

Hard water is a serious problem, and it is a common one. Water in 85percent of the United States is so hard it should be softened to be ofmaximum usefulness.

There are only a few areas where water is sufficiently soft to be sat-isfactory for most homemaking needs. No natural water supply is com-pletely free of hardness.

Communities that draw water directly from snow-filled mountainstreams enjoy nearly ideal water in terms of a low amount of hardness.

New York City with supplies of one to three grains of hardness pergallon has relatively soft water. Even here there are opportunities forsales of water conditioning equipment. There are industries which musthave water free of hardness materials. Some laundries in the area, forexample, have found that zero soft water provides substantial soap sav-ings.

Actually, the hardness of water supplies in this country ranges from1 to 350 gpg (17.1 to 5985 mg/L).

gpg (grains per gallon). This is the most common methodof designating the hardness of a water supply in our indus-try. Grains per gallon equals the number of grains of a givensubstance in one U.S. gallon of water. One grain equals 1⁄7000

pound, and one U.S. gallon of water weighs 8.33 pounds.

Hardness can also be expressed in terms of parts per mil-lion (ppm) or milligrams per liter (mg/L). However, becauseof high amounts of hardness in water, it is generally easierto express hardness in terms of grains per gallon. Conver-sion of parts per million or milligrams per liter into grainsper gallon is quite simple. Simply divide the parts per mil-lion (or milligrams per liter) by 17.1 to convert to grains pergallon.

Parts per million(milligrams per liter) = Grains per gallon

17.1

Further discussion of this topic will be included in LessonNo. 9 on Water Analysis Studies.

70

Here is what an analysis in grains per gallon or parts per mil-lion means to you, according to the U.S. Department of theInterior and Water Quality Association standards:

milligrams perliter (mg/L)

grains per parts pergallon (gpg) million (ppm)

less than 1.0 less than 17.1 soft1 to 3.5 17.1 to 60 slightly hard3.5 to 7.0 60 to 120 moderately hard7.0 to 10.5 120 to 180 hard10.5 and over 180 and over very hard

Most waters possess hardness minerals in amounts from 3 to 50gpg (51.3 to 855 mg/L). Unfortunately, where water is extremely hard,the problem is often compounded by the presence of other contami-nants such as iron and manganese.

Most people are quite aware that a water containing 15 to 30 grains(256.5 to 513 mg/L) of hardness minerals is definitely hard and diffi-cult to use.

On the other hand, many people will tolerate a 5 grain (85.5 mg/L)water that is very objectionable to anyone accustomed to using com-pletely soft water.

Why does hard water constitute a problem?Actually hardness is a source of many problems. One important

trouble area is the way hardness minerals react with soaps and deter-gents.

So important is this aspect of the hardness problem that hardnessis sometimes defined as “the effect of certain elements which combinewith soap to form an insoluble material known as curd.”

The list of elements that possess this property of hardness includeiron, copper, and manganese, all present normally in relatively smallquantities. More common, of course, are calcium and magnesium,which are usually present in significant amounts. Clothes washed inzero soft water are free of troublesome hard water soap curd.

For the homemaker, water hardness makes home cleaning opera-tions more difficult.

In the laundry, hard water leaves soap curd and detergent depositson fabrics. This dulls colors and gives a grey or yellow appearance towhite fabrics. Also, hard water soap curd clings to fabric fibers, causingthreads to become brittle and shortening the life of the material.

71

Hard water wastes soap and synthetic detergents.Hard water leaves unsightly soap scum rings in the bathtub.Hard water spots and streaks glassware and dishes.Hard water builds up scale deposits in all water-using appliances,

clogs hot water pipes.Even more important to consumers and families—hard water ham-

pers good grooming efforts.

Hard Water Scale

Scale is one of the most serious problems caused by hardness mineraldeposits. This particular byproduct of water hardness puts many water-using appliances out of service. It clogs hot water pipes and can sharplyreduce the heating efficiency of a boiler or water heater. When hardwater is heated, scale is formed. This is due to (1) the breakdown of cal-cium and magnesium bicarbonates, (2) their reversion to the highlyinsoluble carbonate forms, (3) their precipitation from the water, and(4) their concentration on the interior surfaces of the water heater.

Under certain conditions, the deposits form sludges. Both sludgesand scale can lead to a sharp reduction in operating efficiency.

Softened Water Energy Savings Study

In a recent University study, energy consumption of gas and electricwater heaters operated and tested on hard water supplies was measured

72

and compared to measured energy consumption of gas and electricwater heaters operated and tested on softened water supplies.

The gas heaters operated and tested on hard water consumed 29.57percent more Btu’s of energy than the gas heaters operated and testedon softened water for the same amount of energy delivered.

The electric heaters operated and tested on hard water consumed21.68 percent more Btu’s of energy than the electric water heatersoperated and tested on softened water for the same amount of energydelivered.

It is not necessary to heat water to a high temperature to producescale. Any increase above the original temperature of the water cancause lime scaling to occur.

Although no chemical reaction occurs which causes calcium sul-fate to deposit when the water is heated, this hardness mineral is un-usual as it is less soluble in hot water than in cold.

Hard water can also be troublesome in industry. In many industrialapplications, however, not only must hardness be removed from thewater, but all mineral content must be eliminated. Mineral depositscan cause serious difficulties in boilers, air conditioning systems, gaso-line, and diesel engine cooling systems.

Water is an excellent solvent. Lesson No. 1 outlines how water col-lects the various contaminants found in it.

As moisture falls through the atmosphere, it absorbs amounts of car-bon dioxide (CO2). It also collects amounts of this gas on and in theground from decaying vegetation. Since carbon dioxide is a product ofboth combustion and decay, it is present in practically all water supplies.

When carbon dioxide dissolves in water, some of it forms a weakacid called carbonic acid.

H2O + CO2 → H2CO3

This acid is responsible for dissolving limestone or carbonate de-posits in the earth. It also produces certain types of corrosion in waterand steam lines. The natural solvent action of water is enhanced by carbonic acid, making it even more effective in dissolving hardnessminerals.

Hardness minerals—calcium and magnesium—are in plentiful sup-ply. While they are not found in their elemental form in the earth, theyoccur in combination with other elements in an abundance of forms.Common calcium minerals include chalk, limestone, and marble.These substances are chiefly calcium carbonate (CaCO3) or mixtures ofcalcium and magnesium carbonates and other impurities. Gypsum iscalcium sulfate (CaSO4). In this compound, calcium is combined withsulfur and oxygen.

73

Epsom salt is magnesium sulfate (MgSO4).Ions of the following calcium and magnesium compounds are

found in water:calcium carbonatecalcium bicarbonatecalcium sulfatecalcium chloridemagnesium carbonatemagnesium bicarbonatemagnesium sulfatemagnesium chloride

The amounts of these various chemical compounds present inwater supplies depend on two factors:

1. The minerals present in the earth; and2. Their solubility in water related to carbon dioxide (CO2) con-

centration.

On a decreasing scale of solubility, calcium chloride, magnesiumchloride, and magnesium sulfate are extremely soluble. They may befound in water in almost unlimited amounts. Calcium sulfate is lesssoluble. At the other end of the solubility scale are the calcium andmagnesium carbonates which are very slightly soluble in pure water.The amounts of these last two compounds in water rarely exceeds twoand five grains respectively. Calcium and magnesium carbonates areseldom found in natural water supplies because of their very low sol-ubility.

While “insoluble” carbonates are rarely found, they are found intheir extremely soluble form in hard water as calcium and magnesiumbicarbonates. When water containing carbon dioxide comes into con-tact with calcium and magnesium carbonates in the ground, a chemi-cal reaction takes place. The “insoluble” carbonate forms of magnesiumand calcium are transformed into highly soluble bicarbonates. Theseare the principal hardness compounds found in water. It is interestingto note that these bicarbonate forms exist only in solution. If heat isapplied to water, the bicarbonates can release carbon dioxide and re-vert to their carbonate or “insoluble” state. For this reason, bicarbonatehardness is often referred to as temporary hardness.

Hardness caused by the presence of the soluble chlorides and sul-fates of calcium and magnesium are classified as “permanent” becausethese compounds cannot be removed from water through simple heat-ing.

When these chemical compounds are dissolved in water, their ionsare released.

74

Water Analysis—Hypothetical Combinations

Suppose for a minute we analyze the total mineral content of a typicalwater. This one has, as we shall see, nine grains of minerals per gallon.It could well be the water which Chicago, Detroit, Cleveland, or anyof a number of other cities draw from the Great Lakes.

CATIONS ANIONS

Ca 5.0 gpg* HCO3– 7.0 gpg*

Mg 2.5 gpg* SO4–– 1.0 gpg*

Na See note Cl– 1.0 gpg**as CaCO3

Diagrammed, these minerals would appear as shown on the chartbelow:

75

Diagram of Mineral Concentration of WaterWith 9 Grains Total Minerals

NOTE: Analysis for sodium is not usually made directly in a wateranalysis. Its concentration is estimated by the difference between thetotal of the anions and the total hardness.

Explanation

The bar at the left in the graph represents the cations or positive ionsof the various minerals in solution.

The bar at right represents the anions or negative ions.

Calcium(Ca++) Bicarbonate

(HCO3– )

Sodium(Na+)

Magnesium(Mg++)

Sulfate(SO4

– – )Chloride(Cl–)

Remember, that in all compounds, the sum of the positive chargesequals the sum of the negative charges. As a water analysis report sim-ply gives the total of various compounds, the same holds true.

Our sample shows positive ions as follows: 5.0 gpg calcium, 2.5 gpgmagnesium, 1.5 gpg sodium for a total of 9.0 gpg. The compensat-ing negative ions are: 7.0 gpg bicarbonate, 1.0 gpg sulfate, and 1.0 gpgchloride.

A chemist making an analysis of this 9 grain water could report itsdissolved minerals in the following manner:

HYPOTHETICAL COMBINATIONS

Calcium bicarbonate Ca(HCO3)2 5.0 gpg as CaCO3Magnesium bicarbonate Mg(HCO3)2 2.0 gpg as CaCO3Magnesium sulfate MgSO4 0.5 gpg as CaCO3Sodium sulfate Na2SO4 0.5 gpg as CaCO3Sodium chloride NaCl 1.0 gpg as CaCO3

Total minerals 9.0 gpg as CaCO3

These hypothetical combinations shown above are one of the waysof describing dissolved minerals in water.

Of course, all of the compounds listed would separate into ionswhen dissolved in water. Thus the various ions, not the complete com-pounds, would actually be present. However, if a chemist wanted toprepare a water sample having the same chemical characteristics as thesample which was analyzed, he could simply weigh out the amountsof the compounds listed and dissolve them in water.

When hypothetical combinations are calculated, the ions are com-bined in their order of increasing solubility. As calcium compounds aregenerally less soluble than other compounds, calcium is usually first onthe list of cations. Magnesium is second and sodium or potassium is last.

Similarly, the anions are used in the following order: hydroxides,carbonates, bicarbonates, sulfates, chlorides, and nitrates.

Note that all the various hardness mineral compounds listed aboveare expressed in grains per gallon as calcium carbonate (CaCO3).

In order to make the calculations as shown, the concentrations ofthe ions must be expressed in such a way that they can be added andsubtracted directly. This is similar to the conversion of 1⁄3 and 1⁄4 to 4⁄12

and 3⁄12 when these fractions are to be used in the same addition orsubtraction problem.

Calcium carbonate has a molecular weight very close to 100, (ac-tually 100.089) and an equivalent weight of 50 (50.045). It is possiblethat this is the reason for its selection as the basic compound, for it cer-tainly simplifies the calculations.

76

If it is stated that a water has invisible hardness minerals in theamount of 10 grains per gallon as CaCO3, this hardness may be due tocalcium or magnesium carbonates, bicarbonates, sulfates or chlorides,or any combination of these compounds. But in every case, the com-bined concentration is chemically equivalent to 10 grains per gallonof calcium carbonate, and the various calculations involved can bemade with ease.

The hardness as CaCO3 of the mineral compounds in water can bedetermined if the chemical analysis of the water is known. The con-centrations of each of the hardness-forming impurities are divided bythe equivalent weight of the compound and multiplied by the equiv-alent weight of CaCO3. Here are a few of these equivalent weights:

EquivalentHardness Producing Compound Weight_______________________________Magnesium sulfate MgSO4 60.187Calcium bicarbonate Ca(HCO3)2 81.057Calcium chloride CaCl2 55.493

To determine the equivalent weight of any mineral compound interms of calcium carbonate:

concentration equivalent weight of CaCO3of the mineral ×

equivalent weight of mineral

equals the concentration of that mineral as CaCO3.

For example:

equivalent wt CaCO3 concentration of10.0 gpg MgSO4 ×equivalent wt MgSO4

= MgSO4 as CaCO3

50.04510.0 × 60.187 = 8.3 gpg as CaCO3.

Traces of elements or compounds are not normally considered inthese calculations. Iron, for example, would not be included unlesspresent in extremely high concentrations.

In the example shown above, the calcium and bicarbonates arecombined first. The excess bicarbonates are then combined with themagnesium. The analysis still does not balance, and the remainingmagnesium is combined with part of the sulfates present. The remain-ing sulfates and all of the chlorides are expressed as sodium com-pounds. (Adding 5.0 gpg Ca++ as CaCO3 to 5.0 gpg HCO3

– as CaCO3produces 5.0 gpg Ca(HCO3)2 as CaCO3, not 10 gpg.)

77

Table of Equivalent Weights

CATIONS

Aluminum 8.994Ammonium 18.0386Calcium 20.040Hydrogen 1.00797Iron (ferrous) 27.924Iron (ferric) 18.614Magnesium 12.156Potassium 39.102Sodium 22.9898

ANIONS

Hydroxide 17.007Carbonate 30.005Bicarbonate 61.017Sulfate 48.031Chloride 35.453Nitrate 62.005Phosphate 31.657Fluoride 18.998Sulfide 16.032

COMPOUNDS

Aluminum sulfate 57.025Calcium carbonate 50.045Calcium bicarbonate 81.057Calcium sulfate 68.071Calcium chloride 55.493Calcium hydroxide 37.047Magnesium carbonate 42.161Magnesium bicarbonate 73.173Magnesium chloride 47.609Magnesium sulfate 60.187Sodium bicarbonate 84.007Sodium carbonate 52.995Sodium chloride 58.443Sodium sulfate 71.021

78

How Hard Water Causes Soap Curd Formation

If such a 7.5 grain water is used in the home for washing purposes, thepositive calcium and magnesium ions react with the soap used to pro-duce insoluble soap curds. Using the block technique, here is what hap-pens:

79

Na+ S.A.–

Na+ S.A.–

S.A.–

S.A.–

+

+

Ca

+

+

Ca

Sodium Soap

(S.A. stands for soap anion)

Plus Calcium Ion Produces Calcium Soap(insoluble curd)

It is these insoluble soap curds that produce the ring around thebathtub, the half-clean, dingy-looking clothes, etc.

Again, if such water is brought into the home or factory and heatedor boiled, the bicarbonate anions revert to their carbonates, which com-bined with calcium and magnesium, produce scale on the walls of thecontainer or in the piping system.

In a manner of speaking, scientists working out the problem ofeliminating hardness minerals from a water supply applied the causeof the problem in obtaining the solution.

They reasoned that if some ions not contributing to hardness couldbe introduced into a water supply as a replacement for the calcium andmagnesium ions, water could be made totally soft, totally useful. Theresult was the ion exchange process for treating hard water.

The Ion Exchange Principle

The idea of ion exchange is not new. Scientists have been aware of theprinciple for a long time. However, it has only been since the start ofthe present century that the principle has been put to practical use.One area in which it has been highly effective has been in the treat-ment of water for removal of hardness minerals and certain other con-taminants.

Recognized household water softening equipment now on the mar-ket makes use of the ion exchange principle. Equipment using thisprinciple contains a bed of permanent bead-like or granular softeningmaterial through which the water flows. As the water travels through

80

Hard water entering softener

LEGEND

Ion exchange material

Magnesium ions

Calcium ions

Sodium ions

the bed of ion exchange material, the hardness minerals are removed,leaving the water soft and more satisfactory for household use.

Bed. The granules or particles of ion exchange material ina softener are referred to as the bed.

The ion exchange material (usually resin beads or granules) con-sists of permanent insoluble anions, kept electrically neutral by re-placeable sodium cations. Hard water contaminated with calcium andmagnesium ions enters the exchange column or bed. As it flowsthrough it, the magnesium and calcium cations in the water are drawnto the anions of the ion exchanger. The ion exchanger has a greateraffinity for the calcium and magnesium ions than for the sodium ions.Therefore, the calcium and magnesium ions are absorbed, and a chem-ically equivalent number of sodium ions is released into the water.Thus, a water containing the ions of calcium bicarbonate when itenters, contains the ions of sodium bicarbonate as it leaves the ion ex-changer bed. In brief, harmless sodium ions have replaced the trouble-producing hardness ions.

Ion exchange occurs literally billions of times between the mate-rial in the exchange column and the minerals in the water as soften-ing proceeds. In Lesson No. 3, diagrams using the block techniqueillustrated a number of basic relationships.

Bicarbonate was diagrammed thus:

81

O C O H

O (–)

+

+

Ca

Condensed to a simpler form this equals:

Calcium was diagrammed…

HCO3 –

And calcium bicarbonate…

82

HCO3 –

HCO3 –

+

+

Ca

Water Softening

Now when this calcium bicarbonate in solution flows through the ex-change material in the softener, the chemical change which occurs isdiagrammed below.

After a vast number of hardness ions in the water has become af-fixed to the softening material through the attraction of positive andnegative charges, and most of the sodium ions have been released, theunit can no longer soften the water. It has become temporarily ex-hausted.

In actual practice, a small number of sodium ions remains on thesoftening material after the unit is exhausted. If no new chemical re-action is set into operation at this point, the incoming calcium bicar-bonate ions flow untouched through the unit.

Just one of the ions causing hardness is shown for sake of clarityin diagramming. Actually, most water supplies contain a number ofvarious hardness ions. The same process in each case applies equallyin their removal from the water.

(1) The calcium ions in the water enter the ion exchange column.Here, the waters pass through the bed of the softening material.

+

+

Ca

Na+

Ca++ NaR CaR

Na+

––

––

R

––

––

R

+

+

Ca

Calciumions. Hard

watercompound

enteringsoftener.

Sodium-cationexchanger

Calcium-cation exchanger(partially exhausted)

(1) (2) (3)

2Na+

Na+

Na+

(4)

(2) The softening material consists of fixed, irreplaceable anions.Affixed, that is, chemically bonded to them are mobile, replaceablecations of sodium.

(3) As the softening material anions have a greater affinity for thecalcium ions than for sodium ions, it attracts them. In the process thecalcium ions “knock” the sodium ions off the exchange material. Asthis continues, the exchange or softening material becomes loadedwith calcium ions. Note that two sodium ions are released for each ofthe calcium ions absorbed by the softener.

83

Calciumions plus

excess brinerinsed down

the drain.

Calcium in resin.(Unit is nowexhausted.)

Sodium ions.Strong brinesolution for

regeneratingsoftener.

Sodium-cationexchanger

+

+

Ca

Na+

Ca++2Na+ NaRCaR

Na+

––

––

R

––

––

R

+

+

Ca

Na+

Na+

(5) (6) (7) (8)

(4) Water that contains calcium ions as it enters the softener willhave a chemically equivalent amount of sodium ions in it on leavingthe softener.

(5) After a certain prescribed amount of water has gone throughthe unit, the calcium ions will replace all but a small percentage of thesodium ions in the softener. At this point, the softener is consideredexhausted and requires regeneration.

(6) Now a rich brine solution is introduced into the softener bybackwashing. Note: to recharge a softener, a concentrated solution ofthe regenerant (sodium chloride) is necessary to force the accumulatedcalcium ions free of the softening material.

(7) The softening resin is now regenerated with sodium ions.(8) The calcium ions and excess brine solution are rinsed away.

When this process is completed, the unit is again charged with sodiumions and is ready to continue the process of softening the water.

Recharging or Regeneration

Recharging or regeneration is necessary at this point. To do this, areverse ion exchange operation is now put into motion. In this reverse

process, it is necessary to bombard the exchange material with the orig-inal type of cations in a concentrated solution. The affinity of the ex-changer for the hardness ions is overcome by the use of a relativelystrong solution of sodium ions. Generally, sodium chloride in a con-centrated solution is used for this purpose. Equilibrium forces nowdrive sodium from its strong concentration in the water to replace andbalance with the hardness concentrations on the resin. What occursin all examples of ion exchange is a “swap” or balanced exchange ofions.

The calcium ions in the softening process are not destroyed. Theyhave merely been replaced in the water by a chemically equivalentamount of sodium ions. The same type of balanced exchange occurswith whatever other hardness minerals that are removed from water.

Domestic Water Conditioning Equipment

Most manufacturers make two basic types of water softeners. These canbe classified as fully automatic and demand initiated regeneration mod-els. These terms have been defined in the “Recommended IndustryStandard for Household Water Softeners,” (S-100), published by theWater Quality Association.

A fully automatic softener is usually equipped with a timer whichautomatically initiates every step in the regeneration process. The re-generating of the unit is usually done at night when water usage is at aminimum.

With a demand initiated regeneration unit, all operations, includ-ing bypass (of hard or soft water depending on design) and return toservice, are initiated and performed automatically in response to thedemand for treated water. Salt storage shall be sufficient for multipleregenerations.

Salt for regeneration may be put in the softener in any of severaldifferent ways depending on the type of equipment. With some mod-els, salt is placed directly in the unit through an opening at the top atthe proper time during the regeneration.

In other units, the salt brine form may be stored in a separateclosed tank. When needed, this brine is forced into the softener byfresh water under pressure. Most common in modern softeners, brineis stored in a nonpressure container and is drawn into the softener bysuction to feed the proper amount of brine into the softener resin tank.

Water softeners are rated in terms of grains of capacity. This ca-pacity refers to the ability of the unit to remove the stated number ofgrains of hardness from a supply of water. The capacity of a unit de-pends on the amount of ion exchanger in the softener, the amount of

84

salt used to regenerate it, plus a variety of other design factors such asregeneration flow rates, etc.

The ion exchange process outlined up to this point meets the re-quirements for the preparation of fully soft water for home needs.

Industrial demands, on the other hand, often call for water that iscompletely free of mineral contaminants. For this reason, the treatmentof the water must be carried further than is necessary for home needs.

Deionization

The process used for removal of all dissolved salts from water is referredto as deionization or demineralization. Deionization requires the flowof water through two ion exchange materials in order to effect the re-moval of all salt content.

Deionization. The terms demineralization and deionizationare used somewhat interchangeably by the industry. Whilethe term demineralization is generally better understood,deionization is especially apt.

The passage of water through the first exchange material removesthe calcium and magnesium ions just as in the normal softeningprocess. Unlike home equipment, deionization units also remove allother positive metallic ions in the process and replace them with hy-drogen ions instead of sodium ions.

As the metallic ions in the water affix themselves to the exchangematerial, the latter releases its hydrogen ions on a chemically equiva-lent basis. A sodium ion (Na+) displaces one hydrogen ion (H+) fromthe exchanger; a calcium ion (Ca++) displaces two hydrogen ions; a fer-ric ion (Fe+++) displaces three hydrogen ions, etc. (Recall that home sof-teners also release two sodium ions for every calcium or magnesiumion they attract.)

This exchange of the hydrogen ions for metallic ions on an equiv-alent basis is a chemical necessity that permits the exchange material tomaintain a balance of electrical charges.

Now, because of the relatively high concentration of hydrogenions, the solution is very acidic.

At this point, the deionization process is just half complete. Whilethe positive metallic ions have been removed, the water now containspositive hydrogen ions and the anions originally in the raw water.

The partially treated water now flows through a second unit, thistime an anion exchange material. This second exchange material

85

86

Water containing various mineralcontaminants entering here

LEGEND

Cations

Hydrogen ion

Anions

Hydroxide ions

H2O

Cation exchange unit

Anion exchange unit

normally consists of replaceable hydroxyl anions and fixed irreplace-able cations.

Now the negative ions in solution (the anions) are absorbed into theanion exchange material. Released in their place are hydroxyl anions.

All that emerges from such a two unit system is ion-free water. Itstill contains the positive hydrogen ions released in the initial exchangeplus the negative hydroxyl ions released in the second exchange.

What has become of these two ions? Through the magic of chem-istry, they have combined (positive to negative) to produce water mol-ecules (HOH or H2O) which are in no way different from the water inwhich they were produced.

The result of this two-stage ion exchange process is water that ismineral-free.

Equipment for use in the deionization process may be of severaltypes. Available are both multiple bed and single bed units. Multiplebed units have pairs of tanks, one for the cation exchanger, the otherfor the anion exchanger. Single bed units incorporate both the cationand anion exchangers, mixed in a single tank.

Deionized water has a wide range of uses in industry. Chemical pro-duction, pharmaceuticals, electroplating, microelectronics, computer chipmanufacturing, electric power generation turbines, and leather goods pro-cessing are among the many diversified applications for deionized water.

Efforts to produce mineral-free water through multiple distillationare not as successful and have proved to be extremely complex and re-quire elaborate and costly equipment.

Lime-Soda Ash Treatment

Lime-soda ash treatment for the reduction of hardness involves the ad-dition of slake lime [Ca(OH)2] to a hard water supply to remove thecarbonate hardness by precipitation with the precipitation being re-moved by filtration. Noncarbonate hardness is, in turn, reduced by theaddition of soda ash (Na2CO3) to form insoluble precipitate which isalso removed by filtration.

This particular method of removing hardness is sometimes used bymunicipal water plants to reduce the amount of calcium and magne-sium in a water supply. While it is quite effective in reducing hardness,it is not a complete removal treatment.

Often when a city has a raw water source that has 35 to 40 grainhard water, the local water system will use the lime-soda ash treatmentto reduce hardness to between 5 and 10 grains.

Lime-soda ash treatment is especially effective if a water containsbicarbonate (temporary) hardness. Where calcium and magnesium are

87

primarily in chloride or sulfate compounds, this treatment is notice-ably less effective.

Slaked lime is used to remove calcium bicarbonate from water. Inthe water to be treated, the slaked lime ions react with the calcium bi-carbonate to form the very slightly soluble calcium carbonate. This pre-cipitated material is usually removed by first settling and then filtering.

Ca(OH)2 + Ca(HCO3)2 → 2 CaCO3↓ +2 H2O

Calcium Hydroxide Plus Calcium Bicarbonate Reacts to Form Calcium Carbonate Plus Water

NOTE: The arrow pointing down indicates the formation of an in-soluble compound.

To remove the magnesium, additional lime is used. The reactionfor this process is:

Ca(OH)2 + Mg++ → Mg(OH)2↓ + Ca++

Calcium Hydroxide Plus Magnesium Ions React to Form Magnesium Hydroxide Plus Calcium Ions

This step has simply replaced the magnesium with calcium. If sodaash is then fed into the water, the calcium will precipitate as calciumcarbonate:

Ca++ + Na2CO3 → CaCO3↓ + Na+

Calcium Ions Plus Sodium Carbonate React to Form Calcium Carbonate Plus Sodium Ions

There are many variants possible under this general heading. Theirdiscussion here, however, is not essential to our course of study.

Lime-soda ash treatment becomes increasingly costly when thehardness of the water must be reduced to less than 5 grains. Munici-pally, the complete elimination of hardness is rarely attemped as lessthan 50 percent of a municipality’s water is used for home consump-tion. The use of soda ash for the reduction of noncarbonate hardnessincreases the sodium in the effluent water in the same proportion asion exchange softening.

The use of the lime-soda ash treatment is impractical for individ-ual home softening of supplies. For one thing, there are difficulties infeeding lime and soda ash into household water. Further, close controlof the operation is required both while the settling and filtering occurs.

88

An additional deterrent to home use of the lime-soda ash treatmentis the size of the equipment necessary, together with the high amountof lime sludge that is produced.

Partial List of Types of Commercial andIndustrial Firms Which Use Soft Water for

Various Operations

Adhesive manufacturersAluminum manufacturersAtomic energy plantsAutomotive industriesBeauty and barber shopsBlanket millsBottlersBuildings: public, private, office, etc.CafeteriasCarpet millsChemical plants, all typesCosmetic manufacturersCotton: mills, bleaching, spinners, etc.DairiesDisinfectant manufacturersDistilleriesDoctors’ offices and clinicsElectric light and power plantsElectrochemical plantsEnameling companiesExplosives manufacturersFiltration plantsFisheriesFood industriesGas engines: power plants, pumping stations, etc.Glass factories: plate, bottle, etc.Glue plantsGypsum products plantsHospitalsHotelsHydrogenation plantsIce plantsIon exchange manufacturersInk manufacturersInsecticide manufacturers

89

Partial List of Types of Commercial andIndustrial Firms Which Use Soft Water for

Various Operations (Continued)

InstitutionsInstrument manufacturersIron: mines, mills, foundries, etc.Jams, jellies and pectin manufacturersJute millsKnitting millsLaboratories: variousLaundriesLeather: tanneries and manufacturersLinen millsLocomotive worksMachine shops and manufacturersMercerizing plantsMotelsMunicipalities: water works, power plants, etc.Navy: ships, air bases, camps, hospitals, etc.Nickel: refineries and manufacturersNursing homesOil wells: drilling, flooding, repressuring, etc.Optical equipment manufacturersOrdinance plantsPaint and pigment manufacturersPaper mills: variousPipe line stationsPharmaceutical manufacturersPhotographic equipment manufacturersPumping stationsQuarriesRailways: electric, dieselRefrigeration plantsRestaurantsService stationsSewage disposal plantsShipsSoap factoriesSteel millsSwimming pools: variousSynthetic detergents plantsTanneries

90

Partial List of Types of Commercial andIndustrial Firms Which Use Soft Water for

Various Operations (Continued)

Textile plants: variousUranium refineriesWashing machine manufacturersWatch factoriesWater works: municipal and privateWood working plantsX-ray equipment manufacturersYarn mills

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93

Lesson 5Introduction

Chemically similar, iron and manganese are twin troublemakers both inthe home and in industry. Even where their concentrations in water arelow, they cannot be ignored. Iron presents the bigger problem only becauseof its greater prevalence.

The presence of iron alone or in combination with manganese leadsto staining problems. They are discussed together in this lesson.

Iron and Manganese

Iron Water Problems

There is little doubt when water contains iron. “Iron water” readilystains plumbing fixtures, porcelain, and cooking utensils. When usedin the laundry, it soon stains washables with reddish-brown discol-orations. An “iron water” also leaves its telltale marks on walls andfloors if used in doing home cleaning chores.

Iron water, if not treated, can lead to serious complications inmany industrial applications. In fact, there is hardly any wet processwork that can be carried on successfully with water that contains iron.

Iron imparts a disagreeable metallic taste to water. Even whenwater contains small amounts of iron, a disagreeable, somewhat as-tringent quality is apparent. Naturally, when iron is present in de-tectable amounts, it can ruin the flavor of tea, coffee, and alcoholicbeverages. Further, the combination of soluble iron and certain of theconstituents in the beverages gives them an unappetizing, inky blackappearance.

The USEPA Secondary Drinking Water Regulations recommend amaximum of 0.3 milligrams per liter (ppm) iron, as Fe, and a maxi-mum of 0.05 mg/L (ppm) manganese, as Mn, because of the stainingthat higher concentrations can cause.

In its insoluble forms, iron can form deposits in pressure tanks,pipe lines, water heaters, commodes, and in any other equipmentwhere water is used.

Iron problems, either alone or in combination with other trouble-some water conditions, are frequent due to the fact that about five per-cent of the earth’s crust is made up of iron. Though not found in a purestate, iron ores are abundant and widely distributed over the earth.

Dissolved concentrations of iron in excess of 60 mg/L are knownto exist. Usually, however, no more than 5 mg/L of iron are present ina water supply. Unfortunately, iron in water becomes a real source oftrouble in homes when as little as 0.3 mg/L is present, generally con-sidered to be the minimum staining level in homes. For many indus-trial needs, an even more critical tolerance of just 0.1 mg/L is necessary.

Water collects iron in several ways. Even as it falls through the air,water acquires small amounts of the oxides of iron found in the at-mospheric dust. Water, rich in carbon dioxide, readily dissolves ironfrom the earth’s plentiful deposits as it leaches these in its undergroundflow.

This bathtub shows the results of iron staining. Not an eye-appealing sight! Water containing soluble iron is clear and colorless

94

when drawn into the tub. When it comes into contact with the air, theiron forms a gelatinous precipitate. The stains that result are extremelyhard to remove. Witness this ruined bathtub!

States of Iron

Iron exists in three basic forms as elemental metallic iron, in ferrous(Fe++), and ferric (Fe+++) states. Ferrous iron usually occurs in waterdrawn from wells. It is present due to the solubility of ferrous bicar-bonate as a result of the action of carbon dioxide on iron deposits inthe ground.

Iron remains in this soluble ferrous state as long as the water re-mains underground, where molecular oxygen is scarce. Carbon diox-ide is commonly found together with high iron concentrations, butthis is not necessarily the case. When this iron-bearing water is firstbrought to the surface, it is usually clear and colorless with a distinctiron taste. After aeration or exposure to the air, the water develops amilk-like haze which soon turns reddish-brown in color.

Chemically what happens is this: upon exposure to the air, molec-ular oxygen begins to enter the water as carbon dioxide escapes. Theoxygen then oxidizes the ferrous ions (Fe++) changing them to ferricions (Fe+++). At this point, the ferric ions combine with free hydroxylions (OH–) to form the insoluble gelatinous compound ferric hydroxide[Fe(OH)3]. As the individual molecules join together, characteristic rustcolor (often called “red water” or “ rusty water”) appears. And finally,

95

a gelatinous precipitate of ferric hydroxide settles to the bottom of thecontainer. In this way, the soluble ferrous ions convert into the insol-uble ferric hydroxide state. Iron flavor noted in water containing fer-rous ions markedly decreases as the ferrous iron passes into the ferricform.

Actually, iron in natural water supplies may be present in a numberof forms including: (1) soluble ferrous ions; (2) ferric ions, soluble invery acid water; (3) ferric hydroxide, insoluble in neutral or alkalinewater; (4) ferric oxides, which show up as particles of rust from pipes;and (5) in combination with organic compounds or iron bacteria.

Iron is generally found in the ferrous state (colorless and soluble)in groundwater supplies. As iron oxidizes upon exposure to the air, itusually settles out. For this reason, it is rarely found in surface watersupplies.

When iron is found in surface supplies, the water may well be ex-tremely acid, or the iron may be combined in various complex mole-cules which resist oxidation. In some surface waters, iron may bepresent in an organic (chelated) form. Such water usually contains agreat deal of colored colloidal turbidity which does not settle and isdifficult to remove by filtration. Unfortunately, organic iron can bequite troublesome, although significant progress in the treatment ofthis type has been made.

Chelate. To combine into a complex molecule having greatstability due to the molecular arrangement.

Iron Bacteria

Iron bacteria frequently thrive in iron-bearing water. As they develop,these bacteria form reddish-brown growths which may clog pipes andreduce flow rates. A decaying mass of these iron bacteria can cause badtastes and odors in a water supply, and often cause extremely discol-ored water when the slimy growths break free in slugs at high flowrates. These iron bacteria can grow either in darkness or in light, butare most frequentlv noticed in toilet flush tanks. They require watercontaining an adequate supply of ferrous ions and free oxygen. Whilethey have been grown in cultures containing no iron, they thrive bestin iron-bearing waters. The most common names for the various typesof iron and manganese bacteria include: Crenothrix, Gallionella, andLeptothrix.

96

Crenothrix. This term is sometimes incorrectly used in re-ferring to all iron or manganese bacteria. Some 18 or morevarieties of iron and manganese bacteria have been classi-fied and studied over the years. Recent study, however, in-dicates that some of the varieties of bacteria are simplydifferent forms of the same bacteria. This study shows thatthe different forms develop differently under differing en-vironmental conditions. The most widely accepted classifi-cations include:

(a) Gallionella(b) Crenothrix(c) Leptothrix

One authority in the subject of iron bacteria states that,despite a long search for Crenothrix, no evidence of a sepa-rate bacteria of this type has been discovered. This same au-thority feels that all iron bacteria are, in fact, forms ofGallionella or Sphaerotilus which includes what are popularlycalled Crenothrix and Leptothrix.

As this brief discussion indicates, iron can be in water in a numberof forms, the cause of which can be quite varied. The chemistry of ironremoval is not difficult once the cause has been clearly determined.Corrective measures present difficulties in some instances only becauseit is not always easy to determine the cause of the problem and becausethe operation of certain types of water conditioning equipment maynot be well understood.

To determine the proper corrective steps requires a bit of sleuthing.The iron, as we have seen, may be in water either in a ferrous or ferricstate. Further, it may be the result of corrosion. The problem may, in alarge measure, be due to the presence of iron bacteria. Because theproblem of iron-bearing water is complex, it is difficult to establishrules for treatment. What must be done depends on the cause and thecharacteristics of the water quality. Various treatment methods are dis-cussed later in this lesson.

Manganese

Manganese is seldom found alone in a water supply. It is frequentlyfound in iron-bearing waters but is more rare than iron. Chemically itcan be considered a close relative of iron since it occurs in much the

97

same forms as iron. When manganese is present in water, it is every bitas annoying as iron, perhaps even more so. In low concentrations, itproduces extremely objectionable stains on everything with which itcomes in contact. Deposits collect in pipe lines, and tap water may con-tain black sediment and turbidity due to precipitated manganese.When fabrics are washed in manganese-bearing water, dark brown orblack stains are formed due to the oxidation of the manganese.

Remember, the USEPA Secondary Drinking Water Regulations rec-ommend a limit of 0.05 mg/L for manganese because of the stainingwhich may be caused. For many industrial purposes the manganesecontent should not exceed 0.01 to 0.02 mg/L. And in some cases eventhis is considered excessive.

Due to the fact that dissolved manganese oxidizes more slowlythan iron, it is generally more difficult to remove from water by oxi-dation and filtration.

Pure elemental manganese metal is gray tinged with pink, brittle,and somewhat harder than iron which it resembles. The pure metal isnot found in nature. However, this chemically active element is foundin many compounds. Deposits occur in certain portions of this coun-try as well as in many other parts of the world.

Manganese is present most frequently as a manganous ion (Mn++)in water. Salts of manganese are generally more soluble in acid than inalkaline water. In this way, they are similar to iron. The manganous ionis usually introduced to water through the solubility of manganous bi-carbonate.

Further, some surface waters and shallow wells contain organic orcolloidal manganese compounds. Manganese bacteria can also causeproblems similar to those caused by iron bacteria—clogging, staining, etc.

Colloidal. Containing or pertaining to colloids which are in-soluble particles 0.01 to 0.1 microns which is about thesame size as viruses. These particles are larger than mole-cules but small enough so that they remain suspended in aliquid without settling. A colloid does not affect the freez-ing point, boiling point or vapor tension of the liquid inwhich it is suspended.

Suspended insoluble manganic hydroxide, known as “black water,”while not rare, is less common. This is probably due to the fact that amuch higher pH is necessary to precipitate manganic hydroxide thanis necessary to the production of ferric hydroxide.

Manganous bicarbonate in solution is colorless. The result is thatunaerated deep well waters containing manganous ions are clear when

98

freshly drawn. Exposure to the air soon converts the clear, solublemanganous ions into the black insoluble substance that is manganesedioxide. Then the trouble begins. The reactions occurring whenmanganous ions are converted to manganese dioxide are as follows:

Reaction occurring in the oxidation of manganese

2Mn++ + O2 + 2H2O → 2MnO2 + 4H+

Manganous Ions Plus Oxygen Plus Water React to ProduceManganese Dioxide Plus Hydrogen Ions

Removal of Iron and Manganese from Water

As mentioned earlier, manganese and iron (especially the latter) pro-duce problems that may be of various types and may be due to variouscauses. Many types of treatment are effective for the removal of ironand manganese from water, but not all methods are equally effectiveunder all conditions.

Generally speaking, there are three basic methods of treating watercontaining these two contaminants. These are:

1. Ion exchange water softeners2. Oxidation and filtration

a. Iron filtersb. Feed oxidizing agent (chlorine, potassium perman-

ganate, or ozone) and filter3. Sequestration—use of such materials as polyphosphates

a. “Pot” feedersb. Solution feeders

Potassium Permanganate (KMnO4). A dark purple crys-talline solid, used as an oxidizing agent in a wide variety ofprocesses. As a solid it is quite stable, and may be kept in-definitely if stored in a closed container in a cool dry area.

Sequestration. The forming of a complex molecule with anion to prevent its normal chemical reaction. For example,polyphosphates sequester or “tie up” ferrous iron, and in-hibit or hamper normal oxidation.

Polyphosphates. A group of molecularly dehydrated phos-phates. In general these materials do not have crystallinestructure, and are commonly referred to as “glasses.” Thepolyphosphates are used in water conditioning for a variety

99

of problems. They may be used to chemically soften waterby “sequestering” hardness, to control iron and manganesestaining by the same process, and to control corrosion bydepositing a thin glass-like film on the interior of the waterlines, water heaters, etc. They are available in the very sol-uble sodium form and in slowly soluble calcium or magne-sium forms.

These methods are appropriate for use in treating waters that havean essentially neutral pH.

Where waters are acid, an especially useful technique is oxidationand filtration accomplished by feeding an oxidizing agent plus an al-kali into the water, then filtering.

All the methods of treatment to be described in the following pagesare for the treatment of iron problems in raw water. Not included inthis discussion are corrosion control techniques for the elimination ofiron problems beyond the point of treatment. The same methods applyequally to the control of manganese problems. The type and amount ofiron in a water supply may vary considerably. Further, the presence ofother water problems has to be taken into account.

Much depends on the source of an iron problem plus the designand flow variables of equipment used for treatment.

Ion Exchange

The use of a water softener unit is considered a satisfactory way to re-move limited amounts of iron from water supplies. No hard and fastrules can be given on the amount of iron that can be treated. The an-swer, in each case, depends upon the design of the softener as well as anumber of other variables. Softener manufacturers normally set the lim-its of tolerance for their equipment based on experience with the prod-uct design and water quality characteristics, such as pH, of the area.

Ion exchange materials remove ferrous ions just as they do calciumand magnesium ions.

Some experts feel that this type of equipment will effectively treatferrous iron in amounts comparable to the amounts of hardnesscations that can be removed from water. Others believe that the use ofsuch equipment should be considered only where amounts of iron aresmall, for example, less than 5 ppm.

Using our building block technique, let’s review here quickly whathappens when iron is removed from water through the use of the ionexchange process.

100

Explanation of the Ferrous Iron Ion Exchange Process

101

Where water contains ferrous ions (soluble iron) as well as hardness ions,all may be removed simultaneously through the sodium ion exchangeprocess. (The calcium and magnesium ions in hard water are not shownfor sake of simplicity.)

The reactions involved in the removal of iron during a softening cycle andin the salt regeneration cycle are exactly the same as in the softening process.

NOTE: When the regeneration frequency of a softener is calculatedand a significant amount of iron is present in the water, you must take thisfactor into account (usually add three to five grains per gallon of hardnessfor each part per million of iron). In no case should you allow the softenerto become completely exhausted. In some cases, only one-half to three-fourths of the normal hardness capacity should be used before regenera-tion is begun to prevent iron fixing as precipitated compounds and foulingof the ion exchange resin.

NOTE: Some manufacturers set separate capacity ratings for iron. Insuch cases, you then base the regeneration frequency on the lower num-ber of gallons calculated from the hardness and iron capacities. A rule ofthumb when using anion exchange water softener for removal of iron isto regenerate often and regenerate vigorously.

Ferrousions

Sodium cationexchanger

Partially exhausted cation exchanger

Iron isexchangedfor sodium

+

+

Fe

Na+

Fe++ Na2R FeR

Na+

––

––

R

––

––

R

+

+

Fe

(1) (2) (3)

2 Na+

Na+

Na+

(4)

+

+

Fe

Na+

Fe++2 Na+ Na2RFeR

Na+

––

––

R

––

––

R

+

+

Fe

Na+

Na+

Ferrousions

Exhausted cationexchanger

Sodium ionsfrom salt

Sodium cation exchanger (regenerated)

(5) (6) (7) (8)

If iron present in a water supply remained only in the soluble fer-rous state prior to and during its passage through the softener (as inthe case of the hardness producing magnesium and calcium ions),treatment could be more positive, and probably more satisfactory.

As it is, most iron-bearing waters contain amounts of insoluble fer-ric iron along with the soluble ferrous iron. This oxidation of the ironmay be caused by contact of the water with air during the pumpingprocess or while it is stored in pressure tanks with air cushions. Whenwater contains iron, it is essential to use a bladder-type pressure tankto prevent contact of air and water.

In any event, some of the iron is apt to have become insoluble be-fore it enters or while it is in the water softener. Now this insoluble fer-ric hydroxide can be removed by simple filtration as the softeningprocess goes on. But its removal presents one serious problem. Troublesets in when this ferric hydroxide begins to precipitate in the intersticesof the resin bed. Even worse, ferric hydroxide (a jelly-like substance)may coat the resinous beads. If present in sufficient quantity, this coat-ing may completely stop the interchange between the hardness min-erals and the sodium ions.

Interstices. Small spaces between the resin particles in theion exchange material.

Remember: The capacity and efficiency of a softener results fromthe water coming into intimate contact with the entire surface of allthe ion exchange particles in the resin bed.

Some of this ferric hydroxide is removed from the unit when thesoftener is backwashed. At the same time, some of the precipitated ironcollects to form large particles which may sift down to the bed supportduring the backwashing. Such a situation can lead to pressure dropcomplications.

The better the filtration effected with a unit, the more effectivelyit removes precipitated iron. Good filtration depends on the particlesize (grading of the resin), the media bed depth, media bed volume,and the operating rate of the flow through the bed. The smaller thecrevices between the particles of ion exchanger and the deeper the bed(within limits), the better the unit will filter precipitated iron, but thehigher will be the pressure drop.

Pressure Drop. A loss in water pressure between the inletand the outlet of a water softener unit.

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Resin Bed Cleaning

Assuming that all these conditions are at their best to provide for goodfiltration, the second requirement is the need to keep the resin bedclean. A clean ion exchange bed largely depends on frequent and thor-ough backwashing. The advent of the fully automatic units hasbrought a marked improvement in the treatment of certain iron-bearing waters because of their frequency of backwashing and regen-eration. (Note: Not all fully automatic softeners are designed for useon iron-bearing water.) If the regeneration program is properly regu-lated, precipitated iron may not have time to “set” permanently withinthe resin bed.

As a safety precaution in the maintenance of a clean ion exchangebed, the periodic application of a chemical cleaner, along with the saltat the time of regeneration, is helpful. Some softener manufacturersrecommend “preventative maintenance” with each regeneration. Toaccomplish this, they suggest a bed cleaning agent with the salt. Weakor mutated acids, sodium bisulfite, sodium hydrosulfite, and other ma-terials have been used to remove iron deposits. However, indiscrimi-nate use of some of these materials may cause other problems. In eachcase, the softener manufacturer should be consulted for his/her rec-ommendation. If a bed of a softener is fouled with large quantities ofiron, professional servicing of the unit is recommended.

How much iron will a softener remove? There are differences ofopinion here which you will have to evaluate on the basis of yourequipment and local conditions.

The ability of a softener to resist the fouling action of iron, espe-cially in its ferric hydroxide state, varies with the design, type of ex-change medium and operating conditions, including pH, turbidity,organic matter, and other factors in relation to the water.

When iron content exceeds the maximum established by the in-dividual softener manufacturer, another type of corrective treatmentmust be brought into use.

Oxidizing Filters

For medium concentrations of iron, the use of an oxidizing filter canbe a most effective means of treatment when the pH is 6.8 or above.When used, the oxidizing filter should be installed in the water lineahead of the softener. Oxidizing filters normally contain a base mate-rial that has been coated with manganese dioxide. This may be a man-ganese-treated greensand, a manufactured manganese material, naturalmanganese-bearing ores, and similar materials. These manganese

103

oxides convert the soluble ferrous iron in the water into ferric iron. Asthe ferric hydroxide forms, it is filtered from the water by the granularmaterial in the filter bed of the tank. Periodic backwashing is, of course,necessary to remove the precipitated iron from the unit. Less fre-quently, it is necessary to regenerate the filter bed.

When the iron accumulation in the filter becomes excessive, re-generation is necessary. First the unit is backwashed to remove precip-itated iron. An appreciable volume of water is needed for even thesmallest units to backwash the dense manganese material properly.This is a factor to consider carefully in any installation because of theimportance of obtaining thorough removal of the filtered iron duringthe cleansing backwash process.

The backwash rate of different filter materials varies considerably.In general, the rate is higher than for most softeners. Whenever thesefilters are selected, the water supply should be first checked to be cer-tain that adequate volume and flow rate of water is available for properbackwashing. This is essential for best results with any filter. The nextstep in the regeneration is to pass a solution of an oxidizing agent suchas potassium-permanganate through the filter bed. This reoxidizes andrestores the manganese dioxide coating on the base filter materials.After rinsing, it is again ready for use.

Oxidation and Filtration

All concentrations of iron may be oxidized by feeding solutions of oxi-dizing agents such as chlorine (in the form of household hypochloritebleach), ozone, hydrogen peroxide, or potassium permanganate, orozone, into the water. This method is particularly valuable when the ironis combined with organic matter, when iron bacteria are present, orwhen the iron concentration is too high for other treatment methods.

Note 1:

Bleach and permanganate solutions should be prepared weekly. Nei-ther is permanently stable in solution. As a result, a drift (the concen-tration of the solution gradually becomes weaker) of feed may occur.

Note 2:

To establish the strength of a bleach solution to be fed to a water sys-tem, the following factors must be known: (1) well pump capacity ingallons per hour; (2) chemical pump capacity in gallons per hour (it iswise to set adjustable pumps at the middle of their range, if possible.

104

This will permit adjustments in both directions.); (3) desired chlorinefeed in ppm; and (4) weekly water consumption in home.

Bleach containing 5.25% sodium hypochlorite has a strength ofapproximately 52,500 ppm chlorine. Thus:

Well pump capacity in gal/hr × desired chlorine feed × 128 oz/gal

Chemical pump setting in gal/hr × 52,500 ppm chlorine= liquid oz of 5.25% hypochlorite bleach per gallon of solution

For example, assume: A well pump capacity of 300 gal/hr; a chem-ical pump capacity of 30 gallons per day but set at 15 gallons per day(equal to 0.625 gal/hr); a desired chlorine feed of 4 ppm; 2,000 gallonsof water used per week. Then:

300 gal/hr × 4 ppm × 128 oz/gal0.625 gal/hr × 52,500 ppm

= �13523,,861030

� = 4.68 oz bleach/1 gallon of solution

Then:

Weekly water consumption × chemical pump setting in gal/hr

Well pump capacity in gal/hr

= gallons of solution required.

Thus:

= 6.67 × .625 = 4.16 gallons of solution.

As a reserve of solution is always advisable, in this case five gallons,should be made up. Thus, dilute (5 oz. bleach per gallon × 5 gallons)25 oz. of bleach to 5 gallons with soft water, where possible. Changethe chemical pump setting if necessary to adjust the concentration.

Note 3: Permanganate

Well pump cap. gal/hr × 133 advoir oz/gal × desired KMnO4 ppm

Chemical pump capacity in gal/hr × 1,000,000= oz of permanganate per gallon of solution

Water consumption in gal/wk × chemical pump capacity in gal/hr

Well pump capacity in gal/hr= gallons of solution for 1 week.

2,000 gallons per week × .625 gal/hr�����300 gal/hr

105

Example: Well pump capacity is 300 gal/hr; chemical pump set at15 gal/day or .625 gal/hr; desired feed of permanganate is 20 ppm; andweekly water consumption is 4,000 gallons. Thus:

300 gal/hr × 133 × 20 ppm =

798,000 =

1.28 oz KMnO4

6.25 gal/hr × 1,000,000 625,000 per 1 gal. of water

4,000 gal per week × .625 = 8.33 gallons needed

300 gal per hour

Thus, dissolve 11.5 oz (1.28 oz × 9) of potassium permanganate in ninegallons of water.

Note 4:

Solutions of sodium hypochlorite are stable only when very stronglyalkaline. In fact, in the process of manufacture, free chlorine combineswith hydroxide ions: Cl2 + 2OH– → H2O + Cl– + ClO–. When thehypochlorite solution is added to water, and the alkalinity neutralized,the chlorine is released. This action is speeded as the acidity is in-creased. Thus, the oxidizing power of the chlorine is released more rap-idly at low pH values.

Potassium permanganate is a strong oxidizing agent across a broadrange of pH values. In neutral or alkaline solutions, the permanganateis reduced to the insoluble manganese dioxide, MnO2. However, invery acid water with pH below 4, it may be reduced all the way to thesoluble manganous ion, Mn++. Thus, consideration should be given tothe pH of the water when the oxidizing agent is selected.

Note 5:

If organic or chelated iron is present, contact times for both chlorineand permanganate may have to be increased significantly. The preciseamount of time cannot, unfortunately, be calculated and may have tobe adjusted on the basis of experience.

A variety of chemical solution feeders are used for this purpose, in-cluding positive displacement pumps, eductors, and several types ofsuction devices.

Like the manganese dioxide in an iron filter, the chemical oxidiz-ing agents convert ferrous iron to the ferric state. The precipitated ironis then removed from the water by filtration.

When chlorine is used as the oxidizing agent, and sufficient con-tact time is made available, not only will the iron be removed, but the

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water will be disinfected. Activated carbon filters are commonly usedto remove the precipitated iron as well as any excess chlorine in thewater. When potassium permanganate is fed, an iron filter is usuallyused to remove the precipitated iron.

With this method of treatment, the oxidizing agent must beintroduced into the water ahead of a retention tank. Under normalconditions, the retention tank, or a similar vessel, will serve to providenecessary reaction (contact) time and to mix the oxidizing agent thor-oughly with the water.

Note 6:

Sufficient alkalinity must be present to assure precipitation of iron andmanganese. The precipitation of these troublesome metals forms anacid and unless sufficient alkalinity is present to neutralize the acid,complete precipitation is prevented. A minimum of 100 milligrams perliter (6 grains per gallon) of total alkalinity is recommended. Alkalin-ity can be added by feeding an alkali such as soda ash (Na2CO3) or bypreceding the iron/manganese treatment with a calcite neutralizer fil-ter. Also, the influent water pH must be greater than 6.5 regardless ofalkalinity, and pH greater than 7.0 is recommended.

When only simple ferrous iron is present in the water, it can be fil-tered immediately after it leaves the retention tank.

Where iron bacteria are present in the water, two approaches havebeen used. In one method, short contact times have been used, usu-ally with high concentrations of chlorine, to kill the bacteria. This ap-proach depends upon the filter to remove the dead bacteria with anyprecipitated iron present. It is probable that oxidation of the bacteriacontinues on the filter bed.

In the other approach, longer contact time is used, usually with rel-atively low concentrations of either chlorine or permanganate, to ob-tain more complete oxidation of the bacteria ahead of the filter. Reportsindicate that both methods have been widely and successfully used.

Water treatment authorities have found that one of the most ag-gravating forms of iron is organic (chelated) iron. Iron in this form doesnot respond to the more simple iron treatment methods, for it isbound into organic materials which both tie up the iron in a nonionicform, and are unusually resistant to oxidation. Thus, neither softenersnor iron filters will effectively remove this iron.

Strong oxidizing agents and long contact times are frequently theonly answer to the presence of organic iron. High chlorine concentra-tions have been effective in some cases, but some authorities point outthat it has been their experience that, in most waters, potassium per-manganate will far surpass chlorine in oxidizing the organic iron. The

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key to the process is undoubtedly in assuring sufficient contact timebetween the oxidizing agent and the organic matter to insure completereaction. Unfortunately, this cannot be estimated without tests on thewater, but in many cases 20 to 30 minutes of contact are necessary.Where the proper conditions are provided, this type of treatment pro-duces excellent results regardless of the quantity of iron in the water.

Depending upon the amount of sludge produced, however, caremust be exercised to insure frequent backwashes to keep the filter bedclean. In some cases, retention tanks or settling basins are provided toreduce the sludge load on the filters. More details on the operation ofa chemical feed pump will be given later in this lesson.

Sequestration

There are, in addition, several other techniques for the control of sol-uble iron in the water. These make use of polyphosphates to keep theiron in solution.

Polyphosphates do not remove iron from water. Rather, they sta-bilize and disperse the iron so that the water remains clear and doesnot produce iron stains.

However, polyphosphate treatment may not prevent iron from pre-cipitating when water is heated for a time, as in the hot water heater, incooking or in the brewing of tea or coffee, as heating can cause reversionto the orthophosphate which has no equivalent sequestering action.

Orthophosphate. A simple phosphate compound, in whichthere is only one phosphate group (PO4

–––) per molecule.In contrast, polyphosphates may contain many such groupsin a single molecule.

Polyphosphate treatment also exerts a dispersing action on old de-posits of iron in tanks and pipes. While this is advantageous to the re-moval of these deposits, an iron problem may be temporarilyintensified, and frequent flushing of pressure tanks and hot water stor-age tanks may be necessary until the old iron deposits are removed.

Polyphosphate treatment is not suited for treatment of iron in mu-nicipal supplies at the point of use. This is due to the fact that suchiron may be partially or completely precipitated and insoluble beforeit enters the home. Polyphosphates are not effective in the control ofprecipitated iron, organic iron, or iron bacteria.

Solutions of the very soluble sodium polyphosphates may be fedinto the water with the various small chemical solution feeders. In

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109

A Positive Displacement Pump

Solution Pot-Type Feeder

general, these units add solutions only when the well pump is operat-ing, so that adjustments can be made to provide fairly uniform dosage.The original cost of the solution feeder is somewhat of a disadvantage.However, this is sometimes offset by feeding several solutions to con-trol more than one water problem.

A simple phosphate cartridge filter or a phosphate pot-type feederis a less expensive original unit. These feeders utilize polyphosphateswhich commonly have calcium or magnesium incorporated into themto provide a product which dissolves slowly and evenly.

The feeder or filter is installed in the water line so that all or a partof the water supply passes through the filter or feeder tank. The waterpicks up some of the polyphosphate solution and carries it into thewater line.

Since water usage in the home is not constant, the feeding fromsuch filters and solution pots sometimes produces slugs or excessivedosages as well as very low dosages. To provide adequate feed at alltimes, it is usual practice to set the feed rate somewhat higher than theconcentration actually required. A pressure or retention/equalizationtank will help to equalize the fluctuations in feed rate which mightoccur under varying flow conditions.

In order to maintain water clarity and prevent the possibility ofiron stains, these polyphosphates must be introduced into the waterat a point where the iron is still present in dissolved form. This shouldbe before the pressure or retention/equalization tank and as close todischarge point as possible.

Discussion

All of the techniques mentioned up to this point are designed to treatiron in waters that have an essentially neutral pH, with the exceptionof the polyphosphates, which are effective in the pH range of 5.0 to 8.0.

Where waters are acid, the following methods of treatment are rec-ommended:

• The use of an ion exchange softener is acceptable in low pHwaters.

• Larger amounts of iron and acidity call for the use of a neu-tralizing calcite filter or alkalinity feed, a strong oxidant feed,and a filter.

One highly useful technique is the treatment of water with an al-kali such as soda ash to raise the pH of the acid water and chlorine toprecipitate the iron. With such treatment, a simple carbon filter cansometimes be used to remove the insoluble iron from the water.

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An alternate method is again the use of an alkali for treatment ofthe acidity and permanganate greensand for the iron.

When rusty water and staining persists in spite of adequate andproper treatment, it is likely that iron bacteria or corrosion are re-sponsible. Even where the iron content of water is slight, iron bacteriacan feed on the iron and store it in their sheaths. Further, iron bacte-ria may actually take iron from steel pipe. Occasionally a bacterialslime accumulation breaks up to cause discharge of extremely turbidwater.

The presence of iron bacteria can be determined microscopically.Also at times, slimy reddish brown growths will be apparent in theflush tanks of water closets.

Chlorine, in the form of a solution of sodium hypochlorite, is usu-ally used to kill the bacteria. However, a filter will be necessary to re-move these dead, iron-laden bacteria from the water.

The following procedure is effective in killing iron bacteria alreadyin the piping system. It may be used periodically if necessary, but isprimarily recommended when steps have been taken to prevent fur-ther entry of the bacteria:

1. Add approximately one-half gallon of sodium hypochlorite(household bleach) to the well. With some wells, it may bepossible to add the liquid through the breather pipe. Withothers, it may be necessary to remove the well seal. Flushseveral gallons of water into the well to rinse the bleachdown to the water level, and reseal the well.

2. One at a time, open each household tap, letting the waterrun until the odor of chlorine is apparent. Close that tap andrepeat with the next one. Do this until the entire plumbingsystem is filled with chlorinated water.

3. After about 12 hours contact time, the chlorinated watershould be thoroughly flushed from the entire system. Con-siderable red water will be flushed out. Particular attentionshould be given to the flushing of the water heater, anddraining of the heater is desirable.

Manganese Removal

Light concentrations of manganese can be removed with a water sof-tener. Higher concentrations may be removed with oxidizing filterswith considerable success. Very high manganese concentrations, orthose complicated by organic matter, etc., call for chemical oxidation,as with iron, plus filtration.

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Note that chlorine will not completely oxidize manganese unlessthe pH is above 9.5, whereas potassium permanganate is effective atpH values above 7.5. Thus, permanganate is the preferred oxidizingagent in most cases.

Chemical Solution Feeders

A wide variety of chemical solution feeders may be used to introducesolutions into the household water supply for treatment purposes.Chlorine or permanganate for the oxidation of iron and manganese;alkalis for the neutralization of acid water; and polyphosphates for thecontrol of soluble iron can all be used with these feeders. The followingmaterial is aimed at chemical feed pumps, but many of the same prin-

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Correct Chemical Feeder Installation

ciples will apply to eductor and suction devices, differing only in theinstallation and method of operation.

Chemical feed pumps are usually installed in connection with pri-vate well systems. Coordination of the feed pump with the well pumpprovides a most satisfactory method of proportioning a solution intothe water line. When the motor of one of these pumps is wired to op-erate with the well pump, the solution is fed at a quite consistent ratioto the water drawn.

Solutions should be introduced into the water line prior to aretention or storage tank. As most of these pumps use a reciprocatingpiston-like action, solutions are not fed continuously, but in intermit-tent shots as the pump cycles.

NEW ART TOCOME?

The retention tank can serve as an excellent mixing unit to evenout the intermittent feed provided by the pulsating action of the

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Incorrect Chemical Feeder Installation

pump. For good results, all the water must pass through the tank. Also,the inlet and the outlet must be placed so as to avoid a direct flow ofwater through the unit, if possible. To do this, connect the outlet andinlet at opposite sides of the tank and at different levels. It is advisableto install a drain valve in the tank, separate from the inlet and outlet,to permit periodic flushing of the tank for removal of sludge.

It is highly important for all of the water to pass through the re-tention tank as shown above. The installation shown on page 113 isnot acceptable because the pressure tank is on a blind leg. Here, waterenters and leaves the retention tank through the same pipe. Mixing ispoor. And, if taps are opened with the pump in operation, it is proba-ble that some untreated water will flow into the household water lines.

Recent tests show that properly installed retention tanks producemixing which approaches theoretically perfect curves. However, thesetests also show that because of the high propensity for short-circuitingwith flowing water currents, it is not wise to depend upon an unbaf-fled retention tank to provide much “contact time” if this is necessaryfor a chemical reaction to take place. The USEPA says that retentiontanks with no internal baffling may actually provide as little as ten per-cent of the contact time calculated by dividing volume by flow rate.

However, where contact time is to be provided, two simple auxil-iary devices may be used. One is a simple coil of plastic pipe. The otheris an ordinary tank loaded with coarse gravel. To calculate the contact

time provided, divide 38% of the actual water capacity of the gravel-filled tank, or 75% of the volume of the pipe, in gallons, by the maxi-mum flow rate of the system in gallons per minute. The results willgive the minimum contact time provided by these devices in minutes.The choice in each instance will depend upon the local costs for thesetwo devices.

In many cases where two chemicals are fed into the water, they canbe mixed and fed from a single container. When chemicals are notcompatible and might produce an undesirable reaction in the solutiontank, a feed pump with two separate pumping heads and two solutioncontainers are necessary. An alternate method is the use of two sepa-rate pumps.

Pumps with two heads are also used to increase capacity beyondthat possible with a single head unit. Two pumps or two heads can beused to pump double the volume of solution from a single tank.

Because of the versatility of chemical feed pumps, they have appli-cations for a wide variety of water conditioning problems. Their disad-vantages are: initial cost; the necessity for periodic preparation ofsolutions; and, the initial trial and error period commonly needed to es-tablish the proper feed rate, as feeds cannot always be calculated precisely.

Summary

Iron and manganese-bearing waters can be extremely troublesome. Ofthe two, iron is found more frequently. Generally, it occurs in con-centrations of less than 5 mg/L (ppm); however, amounts of more than60 mg/L (ppm) are on record. In the case of both iron and manganese,amounts as little as 0.3 mg/L (ppm) of iron and 0.05 mg/L of man-ganese can cause serious staining of fixtures and washable fabrics. Formany industrial uses, 0.1 mg/L (ppm) of iron and 0.01 mg/L (ppm) ofmanganese is critical.

Not only does iron occur naturally in water in ferrous and ferric forms,but it also may be due to corrosion or iron bacteria. The variety of ironproblems calls for careful consideration of the proper corrective treatment.

For both iron and manganese, numerous similar treatment tech-niques are available. Generally speaking, removal of iron and man-ganese is possible with these techniques:

1. Ion exchange;2. Oxidation and Filtration

a. Iron filtersb. Feed oxidizing agent (chlorine, ozone, or permanganate)

and filter;

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3. Sequestration, such as polyphosphatesa. Pot feeders or phosphate cartridge filterb. Solution feeders

If a water is acid as well as iron-bearing, a combination of a neu-tralizer filter or soda ash and an oxidizing agent should be fed, and afilter used to remove the precipitated iron. A review of this chapter willemphasize the proper method of treatment for the various types ofconditions encountered in iron-bearing waters.

Sometimes, iron problems persist despite the best possible treat-ment. Investigation may show that the problems are due to iron bac-teria. To remove iron bacteria, it is necessary to disinfect and flush outthe entire plumbing system with heavily chlorinated water.

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Lesson 6Introduction

Corrosion is a baffling problem. Quite unlike the hardness, iron andmanganese problems we have studied thus far, the cause and treatmentof corrosion problems require close study to insure good results fromtreatment.

Corrosion

Corrosion is a phenomenon associated with the behavior of a metal inits environment which, in this discussion, refers to the behavior of met-als used in plumbing systems in fresh water suitable for household use.Corrosion has also been properly described as the tendency of a metalto revert to its natural stable state as an ore. The process may involve asecond step in which an oxide, hydroxide, or carbonate of the metalmay form and deposit at the corrosion site. Thus, corrosion takes placewhen a metal dissolves in, or is disintegrated by, water. In certain cases,the second step corrosion products or deposits formed may be protec-tive against further corrosion.

Complexity is added by the fact that different metals have differ-ent tendencies to corrode or not to corrode in the same water. A cer-tain type of water may corrode one metal and not another, but thereverse may be true for another type of water.

Water itself can present many different environments. Surfacewater is quite different from groundwater. Both vary geographically,and surface water also varies seasonally. Many communities take theirwater supply from multiple sources, and certain communities even sup-ply mixed surface and groundwater. There is considerable variabilityin the water environment.

Household plumbing constitutes a second variable environment.The fact that a number of metals, alloys, and even nonmetals may com-monly be used in a single household plumbing system or water-usingappliance adds further variation and complexity.

Unfortunately, therefore, domestic water supplies vary in qualityand in their tendency to corrode, and the materials used to constructplumbing systems also differ in their tendencies to be corroded.

Up to this point, we have viewed corrosion in water as a chemicalphenomenon, subject to variations in chemical impurities (gases andminerals) present in the water and variations in the chemical compo-sition of plumbing materials. We have not yet considered the physicalfactors of water temperature and water flow velocity—both of whichcan strongly influence corrosion and erosion rates.

Lastly, it is generally accepted that corrosion of metals is electro-chemical, resulting from the flow of electric current between electrodes,which may be between different metals, or between anodic and ca-thodic areas on the surface of a single metal. Disintegration of themetal occurs at the anode areas. This brings our discussion of corro-sion to a third variable or complexity: the fact that metal surfaces them-selves are not homogeneous in their composition and, therefore, anumber of anodic and cathodic areas may be present on the surface ofthe same metal. Metallic impurities, accumulations of sediment or

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corrosion products, even adherent biological deposits may be directlyor indirectly related to the development of electrical corrosion circuitsor galvanic cells.

An excellent method for corrosion control in water heaters is ca-thodic protection which involves the use of a sacrificial anode, usuallycomposed of magnesium or aluminum. Chemical control of corrosionattempts to retard electrode reactions. Ever since the problems con-nected with corrosion of municipal distribution systems and house-hold plumbing systems in the United States began to be recognizedearly in this century, achievement of a cure has continued to remainelusive, as has the determination of all the exact causes of corrosionand the ability to predict or anticipate rates of corrosion and useful lifeexpectancy of distribution and plumbing systems.

While great strides have been made, chiefly in the use of corrosionresistant materials of construction and through the use of cathodic pro-tection for water heaters, each solution has encountered a few prob-lems in universal countrywide application. In fact, use of each materialor system has uncovered new types of corrosion problems which werelargely unpredictable and unanticipated.

Corrosion of plumbing systems in water supplies depends on somany interdependent variables that no simple equation or corrosionindex is capable of predicting the corrosion potential of a water sup-ply, and no generally applicable recipe for universal corrective treat-ment is available.

The Langelier Saturation Index

In 1936, W. F. Langelier pointed out in a discussion of the saturationindex (calcium carbonate solubility index) which bears his name “…theLangelier saturation index is an indication of directional tendency andof driving force, but it is in no way a measure of capacity. The capac-ity to coat the pipe will depend on the property of the water to resistchange in the value of the index following attack.” Since the Langelierindex relates only to the state of saturation of calcium carbonate at anygiven time, it is incorrect to classify a water with a negative Langelierindex as corrosive. It would be more accurate to state that a water sam-ple exhibiting a positive Langelier value will deposit calcium carbon-ate, and that a water sample exhibiting a negative value will notdeposit, but will dissolve calcium carbonate. Thus, these indices, basedon calcium carbonate solubility equations, are approximate measures ofsolubility, not corrosivity. This is verified by many water supplies thatdo not conform to the predictions of these indices. For example, thereare many water supplies exhibiting high positive index values that are

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corrosive and many water supplies with negative values with softenedas waters with no calcium carbonate, that are not corrosive; facts whichare in direct contradiction to the original hypothesis that positive indexvalues signify deposition of calcium carbonate in pipes and thus pro-tection, whereas negative index values signify lack of deposition, ordissolving, of calcium carbonate from pipes and thus potential for cor-rosion.

Even though no simple equation or corrosion index could predictthe corrosive potential of a water supply accurately, and though no uni-versally applicable method for corrosion control was available, a greatdeal of highly creditable work in determining the causes of corrosionin water supplies and the development of methods for corrosion controlhas been done over a period of more than sixty years by water chemistsand engineers. Such studies have resulted in the development of a num-ber of sound methods for the control of specific corrosion problems atthe point of use, even though these methods cannot be applied uni-versally for all types of corrosion in all water supplies. However, withproper knowledge of specific water characteristics, actual plumbing sys-tems, and other pertinent existing conditions, the point-of-use waterconditioning professional can successfully employ specific corrosioncontrol methods to correct specific corrosion problems.

Causes of Corrosion

Corrosion can be defined as the destructive disintegration—the “eat-ing away” of metals due to electrochemical action.

There are three basic types of corrosion. They are due to:

1. oxygen2. acids3. galvanic action

When the corrosion of metal occurs, the same type of reaction oc-curs with all three of these corrosion-producers. To illustrate: in thecase of iron, it reacts in the following manner:

Fe → Fe++ + 2e

Elemental Iron Metal Reacts to Produce Ferrous Iron Ion Plus 2 Electrons

What causes this reaction depends on the corrosion-producer ineach case. For example, when oxygen is the cause of corrosion, thisportion of the total reaction is:

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(A)2H2O +02 + 4e → 4OH–

Water Plus Elemental Oxygen Plus 4 Electrons Reacts to Produce Hydroxide Ions

Combining these two half-reactions, we get the following:

(B)2H2O + 2Fe + O2 → 2Fe++ + 4OH

Water Plus Elemental Iron Metal Plus Elemental Oxygen Reacts to Produce Ferrous Iron Ions Plus Hydroxide Ions

The basic iron metal is consumed at this point. The ferrous ironcan now react with additional oxygen to form ferric oxide or be car-ried away as ionic iron by water.

Iron Corrosion. The corrosion of iron can take place in theair as well as in water as long as moisture is available. Metalfence posts provide a striking example of the effect of bothoxygen and moisture in corrosion. Above the ground suchposts are exposed primarily to oxygen. Below the groundthey are exposed primarily to moisture. Close examinationof a rusted fence post shows that most of the corrosiontakes place right at the ground level where both oxygen andmoisture are present.

With acids as the cause of corrosion, the acid half balance of thereaction becomes:

(C)2H+ + 2e → H2

Hydrogen Ions Plus 2 Electrons React to Produce Hydrogen Gas

Coupling the two half reactions, the balanced equation in this casebecomes:

(D)2H+ + Fe → Fe++ + H2

Hydrogen Ions Plus Elemental Iron React to Produce Ferrous Iron Ions Plus Hydrogen

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When galvanic action is the cause, the half reaction is:

(E)Cu++ + 2e → Cu

Copper Ions Plus Electrons React to Produce Elemental Copper

This reaction will occur if copper ions are in solution. In othercases, hydrogen gas may be formed as given in example (C).

(F)Cu++ + Fe → Fe++ + Cu

Copper Ions Plus Iron React to Produce Ferrous Iron Ions Plus Elemental Copper

Beyond equation (F), final disposition of the ferrous iron (formedby acid or galvanic action) would depend on the pH and dissolved oxy-gen in water. A wide variety of potential conditions exists which couldlead to many different results.

We have limited our illustrations of corrosive activity to iron. Thisis not perhaps too strange, for we are all apt to think of corrosion pri-marily in relation to iron. In fact, most people consider iron—the mostwidely used of all metals—as the metal that corrodes most easily.

Some metals are more chemically active than others. Some reactmore readily in forming ions or compounds. These reactions are linkedto a metal’s readiness to release electrons in the formation of these com-pounds or ions. For example, some metals combine with oxygen or hy-drogen quite easily; others are comparatively inert. Potassium, at oneextreme, readily combines with oxygen. Gold, at the other extreme, isextremely stable.

A glance at the chart on page 121 shows that six metals are morechemically active than iron. In each case, they can be oxidized moreeasily. It may come as a surprise to note the positions of aluminum andzinc above iron on the list.

When aluminum or zinc are exposed to air, oxidation occurs. How-ever, the oxide that forms adheres to the metal underneath. In this way,the oxide acts as a protective coating to prevent further contact be-tween the bare metal and the air.

In contrast, when iron is oxidized, the rust that forms may flakeoff almost as rapidly as it occurs. The reason for this: the oxidation ofiron usually results in a rather porous material. In some cases, this de-posit flakes off or washes away. In so doing, it continually exposes moremetal. Even when iron oxide does not separate from the metal surface,it does not form a sufficiently solid coating to prevent continued

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corrosion. To control this corrosion of iron, coatings of other materi-als are usually applied to the surface.

From observation and previous reading, what are the factors thatproduce the corrosive action of water?

When water is acid, or even slightly alkaline, it has a tendency tobe corrosive. No doubt, you have seen examples of how strong acidshave rapidly dissolved metals. When a water is low in pH, the sametype of action occurs, although at a slower rate. An acid water may betraced to several different causes. For example, it may be acid due tothe presence of certain dissolved gases, such as carbon dioxide, or hy-drogen sulfide. The acidity may also be due to certain acid industrialwastes.

Mine waters frequently contain high concentrations of strongacids, and are probably the most corrosive of all “natural” water sup-plies. In addition, they are frequently heavily charged with iron. Thesemine waters can, however, be satisfactorily treated, providing their sul-fate content is not too great. When an acid water attacks the walls of ametal container, the entire metal surface usually corrodes rather evenly.An exception occurs where water flows in a steady, consistent patternthrough a container. When this happens, the water is likely to eat deepgrooves in the metal.

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Some Common Metals in Order of Their Decreasing Activity fromAnodic to Cathodic

AnodicPotassium (K)Calcium (Ca)Sodium (Na)Magnesium (Mg)Aluminum (Al)Zinc (Zn)Iron (Fe)Nickel (Ni)Tin (Sn)Lead (Pb)Copper(Cu)Mercury (Hg)Silver (Ag)Platinum (Pt)Gold (Au)Cathodic M

ore

Nob

leM

ore

Rea

ctiv

e

Curiously, byproducts of corrosion frequently act to protect met-als from further attack. One common byproduct is hydrogen gas. (Thereaction showing this hydrogen gas byproduct is illustrated in equa-tions C and D.) If the water is quiescent, the hydrogen gas acts as a pro-tective film to prevent further corrosion. Another byproduct is zinccarbonate. This is found when galvanized pipe corrodes. Other by-products vary depending on the type of metal. In many cases, theytend to act as a protective film. If these byproducts are swept away bythe flow of the water, there is nothing to protect against the damagingeffects of continuing corrosion.

The electrical conductivity of a water supply also affects its corro-sive action. It is well known that an electrical current can be producedby immersing plates of dissimilar metals in a solution which conductselectricity. Under such conditions, a definite and measurable amount ofelectricity will flow through a connection between the plates. This con-nection may be an external wire, or it may be a direct contact betweenthe metal plates.

Electrolytes. Substances which ionize in solution, that is,dissociate into ions. These solutions thus become capable ofconducting an electric current.

This sketch illustrates a most interesting experiment. Itshows how water can be a good electrical conductor undercertain conditions. When water is a good conductor, theelectric bulb will light up.

Place distilled water in the beaker. The bulb remains unlit.This indicates that distilled water itself is a nonconductor.Now place solutions of sugar or alcohol in the distilledwater. There is still no release of electricity. Such solutionsare classified as nonelectrolytes.

Finally, dissolve a compound such as common salt in thedistilled water. When this occurs, the bulb burns brightly.This indicates that this compound is an electrolyte.

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For example, if plates of zinc and copper are placed in a solutionwhich conducts electricity, with a wire connected between the plates,the following occurs:

1. The zinc will pass into a solution as zinc ions.2. A flow of current will run through the connecting wire.

A similar situation may occur in a household plumbing system. Itsometimes happens that zinc galvanized pipe comes in contact withcopper or brass (a copper alloy). Under these circumstances, that sec-ond condition—direct contact between the metal plates—exists. As thewater is an electrolytic solution (that is, capable of carrying electric cur-rent), the zinc in the galvanized pipe will pass into solution as zincions. The speed of this reaction increases with the conductivity of thewater. Over a period of time, this loss of zinc ions can be detected inthe deterioration of the pipe.

We can say then: When dissimilar metals are in contact in a solu-tion that can carry an electric current, two actions occur:

1. An electric current flows between the two metals.2. One of the metals gradually dissolves.

In many home water systems, galvanized pipe is used in conjunc-tion with brass valves or other fittings. At every joint between the dis-similar metals, an electric current is generated, and this causes corrosionof one of the metals. The rate at which this reaction occurs is largelydetermined by the conductivity of the water and the amount of brassin relation to zinc. In turn, this conductivity is determined by theamounts of various minerals in the water. Fortunately, most water sup-plies have low conductivity. Thus, the corrosion which occurs does notcause major problems. However, with other water supplies this type ofcorrosion increases the amount of iron in the water and leads to pre-mature failure of pipe lines and water heaters. Galvanic corrosion pro-duces pitting or deep etching in the less noble of the two metals.

Noble. When this term is applied to metals, it refers to theirability to resist corrosion. The more noble a metal, thegreater its power to resist corrosion.

Further, this type of corrosion occurs close to the point where theless noble connects with the more noble metal. It would seem that theuse of a single metal throughout a plumbing system might prevent thistype of corrosion.

While this practice does help, it will not guarantee the preventionof corrosion. Why? Because local impurities may exist on the metallic

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surface. There may be impurities in the zinc used to galvanize steelpipes. Again, there may be impurities due to either the metallic or thenonmetallic substances in the water itself. In any event, there is alwaysthe possibility of corrosion. The rate of corrosion due to electrolytic ac-tion depends on the dissolved mineral solids in the water. The more ofthese solids, the greater its ability to carry an electric current. Hencethe greater its corrosiveness.

Free oxygen in water can also cause corrosion. Just as metals ex-posed to air rust or tarnish, so metals in the plumbing system can beattacked by the oxygen in the water. The oxygen in water actually com-bines with the metal to form an “oxide.” The chemical reaction is thesame as occurs when moist air rusts unprotected metallic surfaces. Inwater or in moist air, reaction with oxygen consumes a portion of themetal present. Corrosion in water systems due to dissolved oxygencommonly takes the form of deep pitting rather than a general attackof the entire surface. The end result may not look as bad as when theentire surface is attacked. But a single small hole can make a compo-nent of a plumbing system useless.

Temperature increases affect the rate of corrosion, because an in-crease in temperature increases the rate of the electrochemical corro-sion reaction. As a rule, an increase in temperature results in a steppingup of the corrosive action of the water. In household water supplies,the use of greater volumes of hot water has intensified the corrosiveaction in hot water piping. Even where there is little or no corrosionin the cold water lines of a home, a high corrosion rate may exist inthe hot water system. Certainly corrosion is far more probable in thehot water lines than in the cold in almost any household installation.Studies show that the corrosion of steel may be stepped up three tofour times the normal rate when temperature of the water is increasedfrom 60°F to 140°F. Over 140°F, the rate may double with every 20 de-gree increase in temperature.

Due to the possible interplay of the three major causes of corro-sion, it is not always simple to pinpoint the source of corrosion troublein any given instance. Yet you need to have some idea of what’s caus-ing the trouble in order to correct a corrosion problem.

Many people mistakenly believe that maintaining a certain amountof hardness in water is the best way to prevent corrosion. The idea isthat hardness causes scale to form on pipe walls. This scale then servesto protect the plumbing system from the attacks of corrosion. Whilethis method can work under certain conditions if carefully controlled,the method is not reliable in the home and is more apt to provokeproblems than to lead to their control. Why? With the varying tem-peratures and flow rates in household plumbing, it is almost impossi-ble to maintain a uniform and impervious scale coating on pipes. It

126

would be difficult to establish and then maintain the right amount ofscale coating. If too much were permitted to build up, it could restrictthe rate of flow and hamper efficient heat transfer. Costs go up, effi-ciency goes down as scale forms. Thus, at best, hardness is a poor an-swer to corrosion.

What can be done about corrosion? There are several acceptedmethods for controlling this problem. These include neutralizing fil-ters, polyphosphates, silicates, oxygen scavengers, eliminating strayelectrical currents in the water pipes, use of nonmetallic plastic pipeand plumbing materials, and coatings. Each has a record of some suc-cess in minimizing the corrosive action of both hard and soft waters.What percentage of success can be expected from use of these tech-niques? The answer must be: “It depends…” This answer is due to thefact that corrosion has various causes. Further, the problem can be ag-gravated by other conditions. All have an effect on the possible successof available treatments. The choice of treatment in each instance de-pends on careful attention to the corrosion factor plus other problemspresent.

Neutralizing Filters

Tanks containing limestone chips (calcite) bring results by eliminatingacid conditions due to carbon dioxide (carbonic acid) or small amountsof mineral acids. In this treatment, the acid combines with the

127

carbonates in the limestone to form relatively inoffensive bicarbon-ates. Limestone filters are easy to use. They require little care. Thehomeowner need only add more limestone chips occasionally.

Two problems may occur with the use of these filters:

1. Neutralizing the acid may cause an unstable condition to de-velop so that scale will form quite easily. If there is muchacid content, the filtering operation may so unbalance thewater that a softener will be needed to remove the hardness.

2. Removal of an acid condition may intensify iron problems.If the water is iron-bearing, use of a limestone filter may pre-cipitate ferric iron in or even beyond the filter, and thus in-tensify staining problems. More frequent backwashing maybecome necessary so the filter does not clog with precipi-tated iron.

Phosphates

Orthophosphates, as well as silicates, react with small quantities ofmetal ions such as zinc, lead, copper, iron, manganese, calcium, mag-nesium, and silicon in water and with metals of the equipment itselfto form a thin, tight, noncorrosive, passivating film on pipe walls. Thisfilm is almost like a glass coating. These films or coatings keep wateraway from metal plumbing materials and, in this way, control corro-sion. Polyphosphates, such as sodium tripolyphosphate, sodium py-rophosphate, sodium hexametaphosphate, and a group known asbimetallic glassy phosphates also form soluble iron, manganese, cal-cium, magnesium, silicon, lead, and other metal complexes throughsequestration (see page 99).

Polyphosphates are produced by dehydrating or condensingorthophosphates. Polyphosphates sometimes tend to rehydrate andrevert to the orthophosphate form. The rate of reversion is acceleratedby high water temperature, low pH, time, and oxides of certain heavymetals, including iron. Different commercial formulations may includeselected agents, such as zinc, that have been shown to reduce thereversion tendency.

Polyphosphates in solution are anionic and may be removed fromwater supplies with anion exchange treatment. However, once reactedwith a metal (such as iron), polyphosphates become a sticky colloidalprecipitate which must be filtered (e.g., with a five micron cartridge fil-ter or #20 filter sand at the bottom of a softener resin bed) or dissolvedwith acid (such as citric, acetic, or phosphoric acid in regenerationbrine) to be removed.

128

Water containing a high concentration of calcium that is in a con-dition to precipitate, treated with only 0.5 mg/L of polyphosphate, canbe kept from depositing CaCO3 scale. Polyphosphates can also holdiron and manganese in solution (sequestration) in a water environmentwhere they would otherwise precipitate, e.g., in the presence of oxy-gen or chlorine at a high pH.

Polyphosphates also inhibit scale build-up by distorting the usualcalcite crystal form when CaCO3 precipitates. The scale structure isweak and not as capable to build upon itself. Because polyphosphatesare strongly charged, they adsorb on silt particles and help to keepthem from settling because the individual particles repel one another.

Blended phosphates comprise various mixtures of chain polyphos-phates and orthophosphate chemicals. The blended formulations pro-vide both the corrosive-inhibiting properties of the orthophosphateion and the sequestering (complexing) ability of the polyphosphates.

Corrosion control phosphates typically utilize a blend of about 60percent poly- and 40 percent orthophosphate. Phosphate films are notpermanent. They will dissolve and wash away over a period of time.To combat this, it is necessary to feed phosphates into the water moreor less continuously.

Initially, phosphate inhibitors must be dosed at rates up to tentimes that required for normal maintenance needs. This is because thepassivating film must be built up in order to be effective. Once the filmis in place, doses of about 0.5 to one mg/L are effective to maintain auniform passivating film plus the amounts needed to accommodateiron, manganese, hardness, and other constituents in the water. Forbest results, generally feed corrosion control phosphate blends at a rateof about one to three mg/L and maintain pH at 7.4 to 7.8.

For sequestering, blends of about 80 percent poly- and 20 percentorthophosphates are applied. Feed rates are about one mg/L for eachmg/L of iron and manganese, plus one mg/L for each ten grains pergallon of total hardness. For sequestering magnesium hardness, potas-sium-based phosphate compounds are more effective than sodium-based phosphates.

Phosphate liquids and phosphate crystals are manufactured for water supply treatment and used via the following applicationmethods.

1. Injection by means of a chemical feed pump.2. Use of specially prepared compounds in a small tank or car-

tridge. Polyphosphates are combined with calcium and mag-nesium during their manufacture. When placed in a smallmixing tank or cartridge, these compounds dissolve slowly.If some of the water is directed through the tank, it will

129

absorb polyphosphates needed in the system. These com-pounds can be fed into water by use of a pot feeder or aphosphate cartridge filter.

The weakness of this second method is in its inability to producean even concentration of polyphosphates. To correct this problem, theuse of a mixing tank or equalization/retention tank to even out thehighs and lows of feed is advisable. Further, the general practice is tofeed a somewhat higher concentration than the required 5 to 20 ppmto compensate for incomplete mixing.

Other Chemical Feeds

Municipalities and industrial firms sometimes feed sodium silicate(Na2Si4O9), or water-glass, into partially softened water to reduce cor-

130

Solution Pot-Type Feeder

When solid polyphosphates are used for treatment of corrosion, it isconsidered advisable to use a dissolving tank like the one above.

rosive tendencies. Like polyphosphates, silicate forms a protective film.In some cases, caustic soda is also added to raise the pH as an aid in ar-resting corrosion. Sodium silicate combines with calcium to form ahard, dense calcium silicate film (CaSiO3). There are disadvantages tothe use of silicates. It is difficult to find a silicate that dissolves at theright rate, neither too slow nor too fast, and careful control is neces-sary in the feeding of sodium silicate to a water supply. Further, aschanges of water temperature occur, protective silica films do not re-main fixed to the container walls. Oxygen scavengers are another in-dustrial approach to the corrosion problem through the use ofchemicals. Sodium sulfite and sodium nitrite have long been used asoxygen scavengers.

Some corrosion experts hold that dissolved oxygen is the primarycause of corrosion. If so, one might well expect that oxygen scavengerscan control corrosion at all temperatures. And they have produced re-sults justifying their long use. Further, oxygen scavengers provide aninexpensive method for controlling corrosion. Unfortunately, they mayrender a water nonpotable. For this reason they can be used only forcertain industrial purposes.

Below are the reactions which occur in water when sodium sulfiteand sodium nitrate are used as oxygen scavengers:

2NaNO2 + O2 —> 2NaNO3

Sodium Nitrite Plus Free Oxygen Reacts to Produce Sodium Nitrate

2Na2SO3 + O2 —> 2Na2SO4

Sodium Sulfite Plus Free Oxygen Reacts to Produce Sodium Sulfate

Coatings

Coatings are a good basic defense against corrosive water. For this rea-son, the galvanizing of steel pipe is a well established practice. The prac-tice of coating steel pipe surfaces with zinc does two things: (1) itprevents contact between the water and the steel; and (2) it tends to“heal” breaks in the coating. Thus, if the zinc coating is damaged byeither mechanical or corrosive action, nearby zinc goes into solution,and then deposits as zinc oxide on the exposed steel to form a new pro-tective coating. In this case, “galvanic action” is used to protect thesteel. Other coatings are also widely used. Among these are “glass” lin-ings, paints, and other organic coatings. All have proved helpful—moreor less.

The major problem with coatings is the difficulty of applyingcontinuous cover. It is hard to coat a metal surface without leaving

131

“holidays” or bare spots. Unfortunately, such breaks in the coating mayserve to intensify corrosive action. Even a tiny pinhole may cause trou-ble. In a glass-lined container, for example, a pinhole may enable cor-rosion to spread like a cancer under the glass coating. Coatingtechniques have made remarkable strides in recent years. Glass liningis now so reliable, for example, that manufacturers guarantee waterheaters for ten or even 15 years. While these results are excellent, thepractice is generally limited to the interiors of tanks and large vessels.Except in rare instances, it is too costly to coat long lengths of smallpipe in the same manner.

Insulating Unions

Where dissimilar metals must be connected in a water system, the useof nonconductive fittings significantly reduces galvanic corrosion.These nonconductive fittings break the connection between the dis-similar metals. This sharply reduces the flow of current between themetals. The result is almost complete stoppage of electrochemical cor-rosion across the joint. Insulating bushings work well where the con-nections involve relatively large surface areas. The inlet and outlet ofwater heaters are good examples. The use of such bushings may not bepractical at every valve location in a system, however. This handicapis especially a factor in the case of small joints.

The Relation of Naturally Soft Water and Softened Water to Corrosion

It is unfortunate that softened water has been automatically classifiedas corrosive by some “experts.” No doubt, this concept is the result ofexperiences with many of the water supplies in the New England states.

In the New England area much of the water is naturally relativelysoft. In fact, it is very low in any total dissolved solids (TDS) contents.It is the paucity of TDS, not softness per se, that makes such waters ag-gressive and corrosive. This water also contains significant quantitiesof carbon dioxide. Thus, the water is both soft and corrosive.

The waters of the Southwest contrast sharply with those of the NewEngland area in terms of hardness. Instead of being two to three grainshard, the waters of the Southwest range from 200 to 300 grains per gal-lon of hardness in some areas.

These waters of the Southwest are corrosive primarily because oftheir high mineral content which increases electrical conductivity. In

132

turn, this leads to acceleration of galvanic corrosion. As can be seenfrom these two examples, hard water can be just as corrosive as softwater. But softness is neither cause nor cure for corrosion. And soften-ing water through the use of ion exchange equipment does not changethe factors which affect corrosion: total minerals, electrical conductiv-ity, oxygen content, acidity, and temperature. There is nothing aboutsodium ions versus calcium and magnesium ions in water that affectscorrosivity.

133

134

CO

RR

OSI

ON

CO

NT

RO

L F

EE

D R

AT

ES

Op

tion

Cau

stic

Sod

a A

shB

icar

bon

ate

Lim

ePh

osp

hat

eSi

licat

eD

osa

ge (

mg/

L)

5–30

10–4

05–

305–

201–

310

–30

TEC

HN

IQU

E

Lim

esto

ne

Ch

ip F

ilte

r

Soda

Ash

Fee

d

CO

RR

OSI

ON

CO

NT

RO

L T

EC

HN

IQU

ES

DES

CR

IPTI

ON

OF

THE

MET

HO

D

Wat

er p

asse

s th

rou

gh t

ank

con

tain

ing

lim

esto

ne

chip

s. T

he

acid

s co

mbi

ne

wit

h t

he

lim

esto

ne

carb

onat

es t

o p

ro-

duce

inof

fen

sive

bic

arbo

nat

es.

A s

olu

tion

of

sodi

um

car

bon

ate

is f

edin

to t

he

wat

er t

o n

eutr

aliz

e ac

idit

y.

USE

D T

OTR

EAT

Aci

d w

ater

Aci

d w

ater

AD

VAN

TAG

ES

Init

ial l

ow c

ost.

Lit

tle

serv

ice

or m

ain

ten

ance

.

No

har

dnes

s ad

ded

tow

ater

. Am

oun

t of

neu

tral

izin

g so

luti

onca

n b

e re

gula

ted

clos

ely.

DIS

AD

VAN

TAG

ES

Add

s h

ardn

ess

to w

ater

.Ir

on p

robl

em in

crea

ses

asac

id is

neu

tral

ized

.

Hig

her

fir

st c

ost

than

lim

esto

ne

filt

er. M

ayag

grav

ate

iron

pro

blem

s.

135

TEC

HN

IQU

E

Phos

ph

ate

Solu

tion

or

Cry

stal

Fee

d

Sili

cate

(e.

g.So

diu

mSi

lica

te)

Solu

tion

Fee

d

CO

RR

OSI

ON

CO

NT

RO

L T

EC

HN

IQU

ES

(Con

tin

ued

)

DES

CR

IPTI

ON

OF

THE

MET

HO

D

A b

len

ded

ph

osp

hat

e so

luti

on is

fed

or

cart

ridg

e cr

ysta

ls a

re d

isso

lved

into

th

ew

ater

to

form

a g

lass

-lik

e co

atin

g on

met

al s

urf

aces

an

d to

seq

ues

ter

dis-

solv

ed m

etal

ion

s in

th

e w

ater

. Ble

nds

som

etim

es c

onta

in s

ilic

ates

for

im-

pro

ved

corr

osio

n in

hib

itio

n w

ith

cop

-p

er a

nd

cop

per

all

oys

and

wit

h s

oft,

acid

ic w

ater

s. D

osag

es o

f 5

to 1

0 m

il-

ligr

ams

per

lite

r ar

e re

com

men

ded

wh

en t

reat

men

t is

init

iate

d. A

fter

on

eto

tw

o m

onth

s a

un

ifor

m p

rote

ctiv

efi

lm is

est

abli

shed

, an

d th

e do

sage

s ca

nbe

red

uce

d an

d m

ain

tain

ed a

t ap

pro

xi-

mat

ely

one

mg/

L.

Sodi

um

sil

icat

e (w

ater

gla

ss)

has

bee

nu

sed

for

over

50

year

s to

red

uce

cor

ro-

sivi

ty. A

s a

gen

eral

ru

le, s

tart

sod

ium

sili

cate

fee

d ra

tes

at 2

0–30

pp

m f

orab

out

two

mon

ths,

th

en m

ain

tain

fe

ed a

t 8–

15 p

pm

.

USE

D T

OTR

EAT

Cor

rosi

veh

ard

wat

er

Cor

rosi

vew

ater

wit

hlo

w h

ardn

ess,

low

alk

alin

ity,

and

pH

less

than

8.4

AD

VAN

TAG

ES

Goo

d co

ntr

ol f

or m

ost

met

als

use

d in

hou

seh

old

plu

mbi

ng.

Als

o h

elp

s p

reve

nt

diss

olve

d ir

on f

rom

pre

cip

itat

ing

to r

edw

ater

sta

ins.

Cor

rosi

on a

nd

red

wat

eror

blu

e/gr

een

sta

inco

mp

lain

ts e

ffec

tive

lyre

duce

d in

gal

van

ized

iron

, yel

low

bra

ss, a

nd

cop

per

plu

mbi

ng

inbo

th h

ot a

nd

cold

wat

er.

DIS

AD

VAN

TAG

ES

Poly

ph

osp

hat

es m

ayre

vert

in h

ot w

ater

an

ddr

op s

equ

este

red

iron

and

man

gan

ese.

May

inte

rfer

e w

ith

oth

er w

ater

trea

tmen

t p

roce

sses

su

chas

ion

exc

han

ge,

iron

/man

gan

ese

rem

oval

,an

d re

vers

e os

mos

is.

Sili

ca s

cale

can

be

obje

ctio

nab

le, e

.g.,

inbo

iler

fee

dwat

ers

and

spot

tin

g fr

om w

ash

wat

ers.

May

inte

rfer

ew

ith

oth

er w

ater

trea

tmen

t p

roce

sses

su

ch a

s io

n e

xch

ange

,ir

on/m

anga

nes

e re

mov

al,

and

reve

rse

osm

osis

.

136

TEC

HN

IQU

E

Oxy

gen

Scav

enge

rs

Coa

tin

gs

Insu

lati

ng

Un

ion

s

CO

RR

OSI

ON

CO

NT

RO

L T

EC

HN

IQU

ES

(Con

tin

ued

)

DES

CR

IPTI

ON

OF

THE

MET

HO

D

Sodi

um

sil

icat

e, s

odiu

m s

ulf

ite,

an

dce

rtai

n n

atu

ral o

rgan

ic c

omp

oun

ds a

rew

idel

y u

sed

for

indu

stri

al a

pp

lica

tion

s.R

edu

ctio

n o

f di

ssol

ved

oxyg

en c

once

n-

trat

ion

to

less

th

an 0

.5 m

g/L

wil

l usu

-al

ly p

rovi

de a

deq

uat

e co

rros

ion

con

trol

pro

tect

ion

. Sig

nif

ican

t co

rro-

sion

can

be

exp

ecte

d if

dis

solv

ed o

xy-

gen

is o

ver

thre

e m

illi

gram

s p

er li

ter

inw

ater

.

Var

iou

s m

etal

an

d n

onm

etal

coa

tin

gsar

e ap

pli

ed t

o su

rfac

es in

con

tact

wit

hw

ater

. Th

e co

atin

gs s

hie

ld t

he

met

alfr

om c

onta

ct w

ith

wat

er.

Non

met

alli

c bu

shin

gs u

sed

in t

he

con

-n

ecti

ons

betw

een

dis

sim

ilar

met

als

brea

k th

e fl

ow o

f cu

rren

t be

twee

n t

he

two

met

als

to p

reve

nt

elec

troc

hem

ical

corr

osio

n.

USE

D T

OTR

EAT

Dis

solv

edox

ygen

Var

iou

sco

atin

gs c

anbe

use

d fo

rsp

ecif

icco

rros

ion

pro

blem

s.

Con

nec

tion

to d

issi

mil

arm

etal

sex

pos

ed t

oco

ndu

ctiv

ew

ater

.

AD

VAN

TAG

ES

Act

ual

ly t

reat

s th

eso

urc

e of

th

e p

robl

emra

ther

th

an a

ttem

pti

ng

to p

rote

ct a

gain

stco

rros

ive

acti

on.

Prov

ides

pro

tect

ion

wh

en t

he

coat

ing

isti

ght

and

un

ifor

m. C

anbe

“ta

ilor

ed”

to a

ny

spec

ific

cor

rosi

onp

robl

em.

Bes

t so

luti

on t

op

robl

em w

hen

diss

imil

ar m

etal

s m

ust

be jo

ined

.

DIS

AD

VAN

TAG

ES

Gen

eral

ly c

ann

ot b

e u

sed

in w

ater

for

hu

man

con

sum

pti

on.

Bes

t su

ited

to

tan

ks,

vess

els,

etc

. Too

cos

tly

for

mos

t p

ipin

g (e

xcep

tga

lvan

ized

ste

el p

ipe)

. If

hol

e do

es d

evel

op,

corr

osio

n m

ay b

eex

trem

ely

seve

re.

Wor

ks b

est

wh

ere

cou

pli

ng

area

is la

rge.

Not

su

ited

to

man

y sm

all j

oin

ts.

137

Lesson 7Introduction

When total dissolved solids and many of the other impurities discussed inLesson 7 occur in high concentrations, such waters become objectionable.For the most part, serious concentrations of many of these contaminantsoccur in certain rather limited areas. While these particular problems maynot be as universal as those mentioned in earlier lessons, they can be ex-tremely troublesome to those who depend on such waters for daily needs.

Total Dissolved Solids, Hydrogen Sulfide,Fluorides, and Other Water Impurities

In general, water for drinking and cooking should be wholesome. Itshould be both potable and palatable. It must be bacteriologically andchemically safe for drinking and be good tasting. It should be clear, col-orless, and have no unpleasant taste or odor.

In our present-day world, we need at least three basic types of water ofsomewhat different quality, depending on the requirements of each use:

1 .Utility Water. Water which is suitable for use in sanitationand lawn sprinkling; adequate in quantity, bacteriologicallysafe, but not necessarily treated to the highest quality.

2. Softened Water. Water which is optimum for bathing,shampooing, personal grooming, laundering and dishwash-ing. Since many of these uses demand hot water, fully soft-ened water produces better results with minimum soap anddetergent usage, and, in addition, provides conservation ofenergy required for water heating.

3. Drinking Water. Water to be used for drinking and cook-ing must be of high quality. It must meet or exceed the bac-teriological and chemical requirements of both the EPAInterim Primary and Secondary Drinking Water Regulations.Since water used for drinking and cooking amounts to aboutone percent of the total water supplied by a community, thisamounts to two gallons per person per day of the 183 gallonsper person per day furnished by the community. The re-maining amount (over 181 gallons per person per day) is used for a variety of purposes such as sprinkling lawns andirrigation, flushing toilets, fighting fires, cleaning streets, aswell as commercial and industrial uses within the community.

4. Of course, many commercial establishments (laundries,beauty salons, car washes, etc.), industries (for manufactur-ing and specific processes), and institutions (hospitals, forexample for laboratory use, hemodialysis, etc.) will want toprovide extremely high quality water of different types forspecific applications at the point of use.

Today, more than ever before, water is what we make it—not onlyfor community water supplies, but also for individual water supplies.Point-of-use water treatment today is an extremely viable and readilyavailable means by which water of extremely high quality can be pro-vided. Moreover, since treatment takes place just before the water isused, point-of-use water treatment also provides distinct and uniqueadvantages in that only the amount of water needed for each specific

138

purpose or application is treated to the desired quality and also there isvirtually no opportunity for recontamination of the water from the dis-tribution system after treatment.

EPA Primary and Secondary Drinking WaterRegulations

Since it is recognized that our highest priority is safe drinking water, itis essential that the WQA Water Treatment Fundamentals Course providethe EPA Regulations which set maximum contaminant levels for waterimpurities. The following charts of Primary and Secondary MaximumContaminant Levels (MCLs) provides a summary of these regulations.Primary MCLs cover contaminants in drinking water that the USEPAhas determined may cause an adverse effect on the health of persons;they are federally enforceable for public water systems. Secondary MCLsare USEPA advisory levels pertaining to aesthetic qualities of water.

Primary Drinking Water RegulationsMaximum Contaminant Levels (MCLs)

Primary (Health-Related) Microbial & Turbidity Contaminants

Contaminants MCLG† MCL‡

Turbidity — 0.5 to 1 NTU in 95% of samples; maximum of 5 NTU under

certain circumstances

Coliform Bacteria zero zero in 95% of samples

Viruses zero 99.99% reduction or inactivation

Giardia lamblia andCryptosporidium Cysts zero 99.9% reduction or inactivation

Primary (Health-Related) Radionuclide Contaminants

Contaminants MCLG† MCL‡

Beta Particle and Photon Activity zero 4 mrem/year(formerly man-made radionuclides)

Gross Alpha Particle zero 15 pCi/L

Radium 226 & Radium 228 zero 5 pCi/L

Radon zero 300 pCi/L (P)*

Uranium zero (P) 0.03 mg/L

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Primary (Health-Related) Inorganic Contaminants

Contaminants MCLG†, mg/L MCL‡, mg/L

Antimony 0.006 0.006

Arsenic (total) zero 0.01

Asbestos 7MFL 7 million fibers perliter (MFL) (longerthan 10 microns)

Barium 2. 2.

Beryllium 0.004 0.004

Cadmium 0.005 0.005

Chromium (total) 0.1 0.1

Copper 1.3 1.3 (action level)

Cyanide 0.2 0.2

Fluoride 4 4

Lead zero 0.015 (action level)

Mercury (total) 0.002 0.002

Nickel 0.1 0.1

Nitrate plus Nitrite 10 10(as nitrogen)

Nitrite (as nitrogen) 1 1

Selenium (total) 0.05 0.05

Sulfate 500 (P) 500 (P)

Thallium .0005 0.002

Primary (Health-Related) Organic Contaminants

Contaminants MCLG†, mg/L MCL‡, mg/L

Acrylamide zero 0.0005 (action level)

Benzene zero 0.005

Benzo(a)pyrene (PAH) zero 0.0002

Bromate zero 0.010

Carbofuran 0.04 0.04

Carbon Tetrachloride zero 0.005

Chloramines 4 4

Chlordane zero 0.002

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Primary (Health-Related) Organic Contaminants continued

Contaminants MCLG†, mg/L MCL‡, mg/L

Chlorine 4 4

Chlorine Dioxide 0.3 0.8

Chlorite 0.8 1.0

Chrysene (PAH) zero (P) 0.0002 (P)

2,4-D 0.07 0.07

Dalapon 0.2 0.2

Di[2-ethylhexyl]adipate 0.4 0.4

Dibromochloropropane (DBCP) zero 0.0002

Dichlorobenzene (ortho-) 0.6 0.6

Dichlorobenzene (para-) 0.075 0.075

Dichloroethane (1,2-) zero 0.005

Dichloroethylene (1,1-) 0.007 0.007

Dichloroethylene (cis-1,2-) 0.07 0.07

Dichloroethylene (trans-1,2-) 0.1 0.1

Dichloromethane zero 0.005(methylene chloride)

Dichloropropane (1,2-) zero 0.005

Diethylhexyl Phthalate (PAE) zero 0.006

Dinoseb 0.007 0.007

Diquat 0.02 0.02

Endothall 0.1 0.1

Endrin 0.002 0.002

Epichlorohydrin zero 0.002 (action level)

Ethylbenzene 0.7 0.7

Ethylene Dibromide (EDB) zero 0.00005

Glyphosate 0.7 0.7

Halocetic Acids (HAA5) zero 0.060

Heptachlor zero 0.0004

Heptachlor Epoxide zero 0.0002

Hexachlorobenzene zero 0.001

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Primary (Health-Related) Organic Contaminants continued

Contaminants MCLG†, mg/L MCL‡, mg/L

Hexachlorocyclopentadiene 0.05 0.05

Lindane 0.0002 0.0002

Methoxychlor 0.04 0.04

Monochlorobenzene 0.1 0.1

Oxamyl (Vydate) 0.2 0.2

Pentachlorophenol zero 0.001

Picloram 0.5 0.5

Polychlorinated Byphenyls zero 0.0005(PCBs)

Simarzine 0.004 0.004

Styrene 0.1 0.1

2,3,7,8-TCDD (Dioxin) zero 3×10–8

Tetrachloroethylene zero 0.005

Toluene 1. 1.

Toxaphene zero 0.003

2,4,5-TP (Silvex) 0.05 0.05

Trichlorobenzene (1,2,4) 0.07 0.07

Trichloroethane (1,1,1-) 0.2 0.2

Trichloroethane (1,1,2-) 0.003 0.005

Trichloroethylene zero 0.005

Trihalomethanes (THMs) zero 0.080ChloroformBromodichloromethaneDibromochloromethaneBromoform

Vinyl Chloride zero 0.002

Xylenes (total) 10. 10.

*(P) = Proposed Standard

†MCLG = Maximum Contaminant Level Goal established at the level at which no known or an-ticipated adverse effects on the health of persons occur and which allows an adequate mar-gin of safety; expressed in milligrams per liter unless otherwise specified.

‡MCL = Maximum Contaminant Level established as close to the MCLG as feasible takinginto consideration costs and treatment techniques applicable at public water systems.

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Secondary Drinking Water Regulations MaximumContaminant Levels (MCLs)

Contaminants SMCL, mg/L†

Aluminum 0.05 to 0.2 depending oncase-by-case circumstances

Chloride 250

Color 15 color units

Copper 1.0

Corrosivity Noncorrosive

Fluoride 2.0

Foaming Agents (MBAS) 0.5(Methylene Blue Active Substances)

Iron 0.3

Manganese 0.05

Odor 3 threshold odor numbers

pH 6.5–8.5

Silver 0.1

Sulfate 250

Total Dissolved Solids (TDS) 500

Zinc 5

*(P) = Proposed Standard

†SMCL, mg/L = Secondary Maximum Contaminant Level expressed in milligrams per liter(unless otherwise specified).

Total Dissolved Solids

Ionic Contaminants and Other Contaminants inSolution and Suspended Contaminants

As we learned previously, many dissolved inorganic water contami-nants or impurities exist as ions in solution, the most common of theseions are:

Cations AnionsCalcium Ca++ Bicarbonate HCO3

–

Magnesium Mg++ Chloride Cl–

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Sodium Na+ Sulfate SO4– –

Iron Fe++ Nitrate NO3–

Manganese Mn++ Carbonate CO3– –

These electrically charged dissolved particles make ordinary natu-ral water a good conductor of electricity. Coversely, pure water has ahigh electrical resistance, and resistance is frequently used as a meas-ure of its purity.

Since only a few of these most common ionic water contaminantsare health related, most natural water supplies are safe to drink fromthe standpoint of dissolved inorganic chemical contaminants. How-ever, even though found more rarely—and in much smaller quantities—certain inorganic ions can be toxic. These contaminants are listed,along with their maximum allowable levels in the EPA regulations,which also includes maximum levels for radiological ionic contami-nants, maximum levels for water turbidity (cloudiness), and maximumlevels for coliform bacteria (which indicate the presence of human oranimal fecal contamination). Turbidity and bacteria are examples ofsuspended water contaminants.

In addition, water supplies can contain dissolved organic chemi-cal contaminants which are usually pollutants that enter water as aresult of man’s activities, such as insecticides, pesticides, and herbi-cides. These are usually chronically, rather than acutely, toxic to manand other life in extremely small amounts. The trihalomethanes aredissolved organic contaminants, such as chloroform, which are formedby the reaction of chlorine used to disinfect water, with humic and ful-vic acids from vegetation decay. Other organics can enter both surfaceand groundwater through chemical spills, such as trichlorethylene,tetrachlorethylene (TCEs), polychlorinated biphenyls (PCBs), dioxin,etc. Many of the organic contaminants are probably carcinogenic (can-cer-producing). The organics do not necessarily exist in water in theform of dissolved ions.

The Secondary Drinking Water Regulations control contaminantsin drinking water that primarily affect the aesthetic qualities of water.Several of these—chloride, sulfate, copper, iron, manganese, zinc, andtotal dissolved solids—are ionized contaminants.

Color and odor are contaminants which cause objectionable sen-sory responses to the water.

pH is a measure of the acid or alkaline strength of a water supplyand corrosivity refers to the ability of a water supply to disintegratepipes and containers.

The discovery and development of membrane water treatmentprocesses (ultrafiltration, electrodialysis, and reverse osmosis) began inthe 1920s, but these did not really become practical in water treatment

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for 40 years. Rapid technological developments of these processes havetaken place since 1960, and today they are invaluable adjuncts to bothwater and waste treatment technology—yet all these membrane tech-nologies, and more, are routinely performed throughout the humanbody such as in the kidneys; the wall of every living cell is a semi-permable membrane.

Reverse osmosis provides a very important unit process for use bythe present day point-of-use water treatment professional. Togetherwith filtration and ion exchange, which can be used in pretreatmentor post-treatment, reverse osmosis dramatically extends the expertiseand really rounds out the technological capabilities of the water treat-ment professional.

The table at the end of this lesson, “Drinking Water Contaminantsand Their Control,” demonstrates the ability and versatility of a reverseosmosis-activated carbon system at a water pressure of 60 psi.

First, such membranes constitute the finest particulate filter known,which is shown by the 100% removal of turbidity and the virtually100% removal of asbestos.

They show greater than 90% removal of total dissolved solids, withcorresponding substantial ability to remove many other cations andanions listed in the EPA Drinking Water Regulations from 30% removalfor silver and 40% removal for nitrate to 98% removal for sulfate.

Likewise, the ability of RO/carbon to remove dissolved organics isexceptional.

In a paper by Hanes, Bratina, and Brown entitled, “Lead Removal inHome Water Purifiers,” presented at the Annual Conference of theAmerican Water Works Association, the authors show that a reverseosmosis household unit operating at normal household water pressurereduced the lead level from an influent of 762 micrograms per liter toa mean lead concentration of 22 micrograms per liter (97% removal).Then, the activated carbon unit further reduced the 22 microgram perliter RO effluent to a mean lead level of 5 micrograms per liter (99%removal for both reverse osmosis and activated carbon).

Alkalinity

Alkalinity of water may be due to the presence of one or more of anumber of ions. These include hydroxides, carbonates, and bicarbon-ates. As discussed in Lesson 2, hydroxide ions are always present inwater, even if the concentration is extremely small. However, signifi-cant concentrations of hydroxides are unusual in natural water sup-plies, but may be present after certain types of treatment. Smallamounts of carbonates are found in natural water supplies in certain

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sections of the country, rarely exceeding 3 or 4 gpg. They may also befound in water after treatment, such as lime soda ash softening. Bicar-bonates are the most common sources of alkalinity. Almost all naturalsupplies have a measurable amount of this ion, ranging from 0 to about50 gpg.

Alkalinity. The alkalinity of water may be defined as its ca-pacity to neutralize acid. Alkali substances in water includehydroxides or bases. They can be detected by their acridtaste and by the fact that they cause litmus paper to turnblue.

Phosphates and silicates, which can also contribute to alkalinity,are rarely found in natural supplies in concentrations significant in thehome. Compounds containing these ions may be used in a variety ofwater treatment processes. Moderate concentrations of alkalinity aredesirable in most water supplies to balance the corrosive effects of acid-ity. However, excessive quantities cause a number of problems. Theseions are, of course, free in the water, but have their counterpart incations such as calcium, magnesium, and sodium or potassium.

You probably will not notice an alkaline condition due to bicar-bonate ions except when present in large amounts. In contrast, youshould readily detect alkalinity due even to fairly small amounts of car-bonate and hydroxide ions.

Strongly alkaline waters have an objectionable “soda” taste. TheEPA Secondary Drinking Water Regulations limit alkalinity only interms of total dissolved solids (500 ppm) and to some extent by thelimitation on pH.

Highly mineralized alkaline waters also cause excessive drying ofthe skin due to the fact that they tend to remove normal skin oils.

Troublesome amounts of alkalinity can be removed by reverse os-mosis, along with other total dissolved solids. Other methods of watertreatment remove total dissolved solids and alkalinity, but they aresomewhat less suitable for household use than reverse osmosis. Thesemethods are distillation and deionization (demineralization).

Several other methods of water treatment will remove alkalinity,but these methods are not satisfactory for household use. They include:

1. Lime softening (described in Lesson 4) removes hardness.At the same time, this process will precipitate an equivalentamount of alkalinity. Lime softening is usually restricted toindustrial and municipal installations.

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Lime softening. While reducing total alkalinity, lime soft-ening does convert HCO3

– to CO3––, a stronger alkalinity ion.

2. An anion resin regenerated with sodium chloride removessubstantially all the anions (carbonates, bicarbonates, andsulfates, as well as nitrates). It replaces these anions with achemically equivalent amount of chloride ions. The disad-vantage of this process is that in almost all cases a high chlo-ride ion concentration results. At the point of exhaustion,the resin has the tendency to unload high concentrations ofthe anions it carries including the nitrates. If deal Kalizersare allowed to operate to the point of exhaustion and suchanion dumping, the results are as undesirable as the originalalkalinity.

3. The feed of a mineral acid will neutralize the alkalinity of awater. Hydrochloric acid, sulfuric acid, or a combination ofthese can be used. This process converts the bicarbonatesand carbonates present into carbonic acid. At this point, itis advisable to provide some method to permit the resultingcarbon dioxide gas to escape into the atmosphere. The dis-advantages of this acid feed technique are obvious. There areneeds for precise control of the process and caution in han-dling the strong acid.

Free Carbon Dioxide

Almost all natural waters contain some carbon dioxide which they gainin several ways. Carbon dioxide gas (CO2) is present in the air to theextent of 0.03 percent by volume and 0.05 percent by weight. As rainfalls through the air, it absorbs some of this gas.

Free carbon dioxide. Refers to carbon dioxide gas dis-solved in water. The term is used to distinguish a solutionof the gas from the combined carbon dioxide present in bi-carbonate and carbonate ions.

On reaching the earth, the rainwater—now slightly acid—willabsorb additional amounts of carbon dioxide if it flows throughdecaying vegetation. At the same time, the carbon dioxide becomescarbonic acid. If the water now passes through limestone formations,

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its carbonic acid content will react with the limestone to form solublecalcium bicarbonate. In this process, the carbonic acid is partiallyneutralized.

Limestone. A sedimentary rock formation wholly, or in largedegree, composed of calcium carbonate. There are manyimportant varieties of limestone, such as chalk, travertine,and marble.

On the other hand, if water passes through rock formations, suchas granite, no such reaction occurs. The carbonic acid is not neutral-ized. It continues as carbonic acid until drawn to the surface where itcan then cause corrosion if not neutralized.

Granite. A type of rock that consists primarily of quartz, al-kali feldspar, and mica. The quartz and feldspar are alwayspresent in granite. Other minerals are sometimes present aswell. These are all silicates.

If nature or chemical agents do not neutralize carbonic acid, it willcause corrosion of both copper and galvanized plumbing systems (re-view Lesson 6). In those parts of the country where the problem isprevalent, it is serious for it can lead to serious damaging of plumbingequipment. Carbon dioxide, together with carbonic acid, is primarily aproblem in water containing relatively low concentrations of miner-als. In such water, there are not sufficient alkaline salts to buffer the ef-fect of the carbonic acid.

The simplest method for removal of carbonic acid is to pass thewater through a tank containing limestone chips. A neutralizing filterof this type affects the carbonic acid just as does the underground lime-stone formation. The limestone in the filter reacts with the carbonicacid to produce calcium bicarbonate. In the same way, lesser amountsof magnesium bicarbonate are formed. Note: Not all forms of limestoneare suitable for this purpose. Excessively soft material may break downto form a solid mass and block the filter. The best types are hard, stronggranules which retain their physical structure, even as they are dissolved.

Another type of material used in this neutralizing process is mag-nesium oxide. Although this procedure does add hardness and alkalinesalts to the water, it effectively neutralizes a considerable amount ofcarbonic acid at a relatively low cost.

Where high carbon dioxide concentrations are encountered, a so-lution of soda ash or sodium carbonate (Na2CO3) may be fed into the

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water. The carbonic acid and the sodium carbonate react directly toform sodium bicarbonate. This method of treatment offers the advan-tage of not adding hardness to the water. Also, it is especially effectivewhere it is necessary to remove carbonic acid from large volumes ofwater. This method, as we have seen, has the disadvantage of requir-ing more attention in the preparation and maintenance of proper feeds.

Where water is obtained from a private well, a small positive dis-placement pump can be used to feed the soda ash solution into thewater. Normally, such pumps are wired to act in conjunction with theoperation of the well pump. This permits the proportioning of the so-lution with a good degree of accuracy.

Where a private water system is not used to draw water to thehousehold lines, some other type of feeding device is necessary. How-ever, the design of such devices is limited only by the ingenuity ofpump manufacturers and installation personnel.

It is important to feed soda ash solutions into the water in advanceof some type of tank or mixing device. This is necessary to provide forreasonably consistent concentrations in the water to be treated.

Chloride and Sulfate

Almost all natural waters contain chloride and sulfate ions. Their con-centrations vary considerably according to the mineral content of theearth in any given area. In small amounts, they are not significant. Inlarge concentrations, they present problems. Usually chloride concen-trations are low. Sulfates can be more troublesome because they gen-erally occur in greater concentrations. Low to moderate concentrationsof both chloride and sulfate ions add palatability to water. In fact, theymay be considered desirable for this reason. Excessive concentrations ofeither, of course, can make water unpleasant to drink.

The EPA Secondary Drinking Water Regulations recommend a max-imum concentration of 250 mg/L for chloride ions and 250 mg/L forsulfate ions (expressed as Cl– and SO4

– –, not as CaCO3).Water containing calcium sulfate (gypsum) ions is likely to have a

characteristic taste…somewhat bitter and astringent. In fact, it has beencompared to the way dissolved gypsum might taste in water. When 30to 40 grains per gallon of calcium sulfate are dissolved in water, mostpeople can detect the taste.

If equal amounts of magnesium sulfate or sodium sulfate are dis-solved in water, the taste would not be noticeable. All sulphates possessdefinite laxative effects in concentrations above 30 grains per gallon.In this way, they can be troublesome especially to people not accus-tomed to such water. In addition to their laxative properties and pos-

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sible medicinal taste, sulfate water can mean extreme hardness, largeamounts of sodium salts, or acidity. Alone or together, these can posespecial problems in the conditioning of water.

Chlorides give water a salty taste. At what concentrations this be-comes noticeable again depends upon the individual. In large concen-trations, chlorides cause a brackish, briny taste that definitely isundesirable. Although chlorides are extremely soluble, they possessmarked stability. This enables them to resist change and to remainfairly constant in any given water unless the supply is altered by dilu-tion or by industrial or human wastes. Both chlorides and sulfates con-tribute to the total mineral content of water. As indicated above, thetotal concentration of minerals may have a variety of effects in thehome. High concentrations of either sulfate or chloride ions add to theelectrical conductivity of water.

Chlorides and sulfates can be substantially removed from water byreverse osmosis. Deionization (demineralization) or distillation will alsoremove chlorides and sulfates from water.

Fluoride

Fluorides in water can be detrimental or beneficial. It all depends onthe concentration. Surface water supplies are normally low in fluorides(less than 0.5 ppm). Some have no fluoride at all. Well waters may con-tain excessive amounts of fluoride above the recommended amount (1 mg/L) for drinking water.

Fluorides are important because they have a definite relation todental health. Research has shown that a concentration of 1 mg/L offluoride in drinking water reduces tooth decay. On the other hand,some children under nine years of age exposed to levels of fluoridegreater than about 2 mg/L may develop a condition known as “en-demic dental fluorosis.” Sometimes called “Colorado Brown Stain,” thiscondition appears as a dark brown mottling or spotting of the perma-nent teeth. In certain cases, the teeth become chalky white in appear-ance. Further, federal regulations require that fluoride not exceed aconcentration of 4 mg/L in drinking water. This is an enforceable max-imum contaminant level standard, and it has been established to pro-tect public health. Exposure to drinking water levels above 4 mg/L formany years may result in cases of crippling skeletal fluorosis, which isa serious bone disorder.

Research studies indicate that fluoride concentrations of 1 mg/Lare optimum. Authorities generally agree: (1) where concentrations aregreater than 4 mg/L, the excess fluorides must be removed from water;(2) where concentrations are less than 1 mg/L, fluorides may be added

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to drinking water for the prevention of dental caries. In today’s soci-ety, however, fluorides are ever more prevalent in many foods and bev-erages, and especially in toothpastes and mouthwashes. Cities arepresently required by some states to add fluorides to municipal watersupplies. Where the fluoride concentration is too great, it is necessaryto reduce the amount to acceptable limits.

Various methods have been suggested for reducing fluorides. Thesecan be classified broadly in three groups:

1. Reverse osmosis.2. Those involving treatment with chemicals, such as alu-

minum sulfate, magnesium or calcium phosphate, andothers.

3. Those involving percolation through a bed of material, suchas activated alumina, granular tricalcium phosphate, oranion exchange resins.

The first treatment method has obvious advantages. Methods inthe second category have distinct disadvantages. They require use ofelaborate treatment plants, careful control of chemical dosage and pH.In some cases, further treatment is necessary to restore the pH of thetreated water to normal.

Methods in the third category do not require such elaborate con-trol. Of these, the only widely used method of reducing fluoride con-tent involves the use of a tricalcium phosphate filter. Such a filterfunctions in much the same way as a carbon filter. As the water flowsthrough a tricalcium phosphate filter, the fluorides are absorbed.

Hydrogen Sulfide

Hydrogen sulfide is a gas present in some waters. There is never anydoubt as to when it is present due to its offensive “rotten egg” odor.This characteristic odor is sometimes apparent in concentrations as lowas 0.5 mg/L. Obnoxious as are the taste and odor of hydrogen sulfide,these are only two of the problems it presents. Hydrogen sulfide pro-motes corrosion due to its activity as a weak acid. Further, its presencein the air causes silver to tarnish in a matter of seconds. High concen-trations of hydrogen sulfide gas are both flammable and poisonous.While such concentrations are rare, their presence in drinking waterhas been known to cause nausea, illness, and in extreme cases, death.High concentrations of dissolved hydrogen sulfide can also foul thebed of an ion exchange softener. Its continued presence will lead tolower and lower capacity and may finally necessitate replacement ofthe resin bed. Generally, hydrogen sulfide occurs in concentrations of

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less than 10 ppm (milligrams per liter). Occasionally, the amount goesas high as 50 to 75 mg/L. Hydrogen sulfide is more common to wellwaters than to surface water supplies.

There are several methods for removing hydrogen sulfide fromwater. Most of them involve converting the gas into elemental sulfur.This insoluble yellow powder can then be removed by filtration. Low tomoderate concentrations of hydrogen sulfide can be eliminatedthrough use of an oxidizing filter of the same type as satisfactory foriron removal. Because the elemental sulfur precipitate tends to clog thefilter material, it is necessary to backwash adequately and sometimesto replace the filter bed material from time to time.

Chemical treatment is recommended for medium to high concen-trations of hydrogen sulfide. In such cases, solutions of householdbleach, ozone, or potassium permanganate serve as satisfactory oxi-dizing agents. When these oxidizing agents—such as household bleachand permanganate solution—are used, a small chemical feed pump willserve to feed the agent into the water. A ratio of 2 mg/L chlorine per 1mg/L H2S is suggested as a starting dosage. This level will normally pro-vide a high enough chlorine residue to insure complete oxidation ofthe sulfide to sulfur. The feeding rate of the chlorine solution may beadjusted from the original settings to provide the most efficient oper-ation. As in the case of iron, the chlorine solution should enter thewater upstream from the mixing or storage tank to provide sufficientcontact time. A contact time of at least 20 minutes should be allowedfor complete reaction. After this contact time, the water should passthrough a sediment and activated carbon filter to remove the now in-soluble sulfur and excess chlorine.

If potassium permanganate can be used as the oxidizing agent, aniron filter is recommended to remove the insoluble products from thewater. (Theoretically, 6.2 mg/L of pure KMnO4 are necessary to oxidize1 mg/L H2S.) However, a slight excess of permanganate, as shown by alight pink color, should be fed to keep the filter in a “regenerated” state.In this way, it acts as a reserve to protect against any unexpected in-crease in the hydrogen sulfide content of the water.

An activated carbon filter alone will remove trace amounts of hy-drogen sulfide. In this process the carbon simply absorbs the gas on itssurface areas. The use of an activated carbon filter can be economicalwhen small amounts of the gas are present. Periodic replacement of theactivated carbon is necessary. With moderate to high concentrations ofhydrogen sulfide this becomes impractical from an economic standpoint.

Some large users of water depend on aeration to remove hydrogensulfide from water. Although this is the simplest basic method, it’s notnormally used for household applications. It has the disadvantage ofhigh initial cost and incomplete removal of the gas. There has been

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some use of the ion exchange process for removal of hydrogen sulfide.The ion exchange material for this purpose is a strong base anion sub-stance which can be regenerated with salt or a mixture of salt andsodium bicarbonate. This technique has the advantage of simplicity inoperation. On the other hand, it offers relatively low flow rate and aneffluent water that has all chloride anions.

Nitrate (Nitrate Nitrogen)

Nitrate Nitrogen. The concentration of nitrates is com-monly expressed as NO3

–. The term “nitrate nitrogen” isused to refer to the nitrogen present which is combined inthe nitrate ion, or nitrate measured and expressed as theamount of nitrogen it contains. This nomenclature is usedto differentiate nitrate nitrogen from nitrogen in the formof ammonia (ammonia nitrogen), from nitrogen in the formof nitrite (nitrite nitrogen), etc. The concentrations are usu-ally expressed in milligrams per liter of nitrogen.

Many groundwaters contain small amounts of nitrate nitrogen.Concentrations range from 0.1 mg/L to 3 or 4 mg/L in most areas.Amounts as high as 100 mg/L have been found, however. Nitrates mayoccur in both shallow and deep well supplies, but they are most com-mon in water from shallow wells. Nitrate nitrogen can result from theseepage of water through soil containing nitrate-bearing minerals. Itmay also occur as the result of using certain fertilizers in the soil; how-ever, nitrates are one of the products of decomposition of animal andhuman wastes. Thus, the presence of nitrates in a water supply may in-dicate possible sewage pollution of the water.

Nitrate nitrogen has been much publicized in recent years in rela-tions to the problem of “blue babies.” In concentrations as low as 10to 20 mg/L, nitrate nitrogen has caused illness and even death amonginfants under six months of age. If such water is used for supplementalor for complete bottle feeding, it may affect the ability of the blood tocarry oxygen. This oxygen starvation is called methemoglobinemia, ormore commonly, the “blue baby” condition. This serious illness in in-fants is caused because nitrate is converted to nitrite in the higher pHconditions existing in the stomachs and intestinal tracts of infants undersix months of age. Nitrite interferes with the oxygen carrying capacityof a child’s or baby animal’s blood. This is an acute disease in that thesymptoms can develop rapidly. In most cases, health deteriorates rapidly

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over a period of days. Symptoms include shortness of breath and blue-ness of skin.

In the process of decomposition, raw sewage undergoes a chemi-cal change. Among the end products is nitrate nitrogen. When ni-trate nitrogen occurs, it is considered evidence of pollution eitherfrom septic tank fields, cesspools, or other sewage sources. Where agroundwater is known to contain little or no nitrate nitrogen natu-rally, the appearance of any significant increase is a probable indica-tion of pollution. Because of these factors, well waters containingnitrate nitrogen should be checked periodically by local or statehealth authorities.

The best method for treatment of large nitrate nitrogen concen-trations due to human or animal wastes is prevention. Wells should beproperly located and constructed in order to prevent sewage contami-nation. Nitrates can be removed through distillation, anion exchange,or reverse osmosis. Even though about 95% of ionic nitrates can be re-moved by reverse osmosis, nonionic forms of nitrogen are not rejectedand pass through the membrane. Nitrogen compounds are weakly ion-ized in solution, which may result in a total rejection of as low as 40%by reverse osmosis especially when feed pressures drop below 50 psi.Bottled water can also be a practical source of nitrate-free water for in-fants. In commercial and industrial water supplies, nitrates do not usu-ally present serious problems.

Oxygen

As rainwater falls through the atomosphere, it collects oxygen gas. Thisdissolved oxygen is not the same as the oxygen in the water molecule.Dissolved oxygen is present in all rainwaters and surface supplies dueto contact with the atmosphere. Just how much dissolved oxygen awater supply will contain depends on several factors:

1. Under high pressure, relatively large quantities of oxygendissolve in water. When the pressure is reduced, a propor-tionate weight of the gas escapes (Henry’s Law).

Henry’s Law. The English chemist, William Henry, formu-lated a law regarding the effect of pressure on a gas. Thelaw states: The weight of a gas that dissolves in any givenliquid is directly proportional to the pressure, providing thetemperature remains constant. If one gram of oxygen, forexample, dissolves in 100 cubic centimeters of water at at-

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mospheric pressure, two grams of oxygen will dissolveunder twice the normal atmospheric pressure, providingthere has been no change in temperature.

2. The amount of minerals in a water affects its ability to dis-solve oxygen. Distilled water can absorb more oxygen thanwell waters with higher mineral content. Obviously seawater, for this same reason, holds less dissolved oxygen thanfresh surface water.

Well waters usually contain smaller amounts of dissolved oxygenthan surface supplies. In deep wells, there may be a total absence ofthe gas. However, an article in Science Magazine, June 11, 1982, pages1227–30, states:

Contrary to the prevailing notion that oxygen-depleting reactions inthe soil zone and in the aquifer rapidly reduce the dissolved oxygencontent of recharge water to detection limits, two to eight milligramsper liter of dissolved oxygen is present in water from a variety of deep(100 to 1000 meters) aquifers in Nevada, Arizona, and the hot springsof the folded Appalachians and Arkansas. Most of the waters sam-pled are several thousand to more than 10,000 years old and someare 80 kilometers from their point of recharge.

Oxygen adds to the taste of water. For this reason, a small amountof it is desirable in drinking water. We are all familiar with the “flat”taste which water often possesses after it has been standing in an opencontainer for some time. The taste can be improved simply by shakingthe water in a partially filled bottle. This reintroduced oxygen into thewater will give it a more appealing taste. Despite this desirable feature,dissolved oxygen can be a source of serious trouble in a householdwater supply. The fact is that oxygen causes corrosion. In cold water,oxygen normally has little corrosive effect. In contrast, when the wateris heated, the oxygen can cause serious corrosion problems. Any dis-solved oxygen levels over 0.5 mg/L can be the cause of copper corro-sion problems, for example.

A number of chemicals are used in industry to remove oxygen froma water supply. Sodium sulfite (Na2SO3) is probably most widely usedfor this purpose. It reacts with oxygen at high temperatures to formsodium sulfate (Na2SO4), in this way reducing the oxygen. There are anumber of chemicals that react similarly with oxygen to effect its re-moval. The degree of success varies. For household purposes, treatmentis normally limited to the use of polyphosphates to coat the insides ofwater lines to protect the metal from contact with the oxygen.

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Silica

Many water supplies contain silica. This is not surprising since siliconis the second most abundant chemical element in the earth.

Silica. (silicon dioxide) A compound of silicon and oxygen(SiO2). It is a hard, glassy mineral substance which occurs ina variety of forms such as sand, quartz, sandstone, and gran-ite. It also is found in the skeletal parts of various animalsand plants.

Silicon. (Si) One of the nonmetallic elements in abundantsupply as part of various compounds in the crust of the earth.

The solid crust of the earth contains 80% to 90% silicates or othercompounds of silicon. Water passing through or over the earth dissolvessilica from sands, rocks, and minerals as one of the impurities it collects.

Silicates. Compounds which contain silicon and oxygen incombination with such metals as aluminum, calcium, mag-nesium, iron, potassium, sodium, and others. Silicates areclassed as salts. Silicates are widely distributed in such min-erals as asbestos, mica, talc, lava, etc.

The silica content of water ranges from a few parts per million insurface supplies to well over 100 ppm in certain well waters.

Colloids. Extremely small solid particles, 0.01 to 0.1 micronsin size, that are suspended in a solution such as water. Theweight of the individual particle is so low that a true colloidwill not settle out, even after standing for an indefinite pe-riod. Colloidal particles are thought to have a charge whichcauses the particles to repel each other and prevent theiraggregation into larger agglomerates. A colloid diffuses veryslowly or not at all through a membrane, and has little, ifany, effect upon the freezing point, boiling point, or vaporpressure of the solution.

In its colloidal form, silica consists of very fine particles in suspen-sion. These can usually be removed by coagulation and settling orfiltering.

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Sodium

Sodium salts are present to a greater or lesser degree in all natural wa-ters. Their concentrations vary from a few parts per million in somesurface supplies to several hundred grains per gallon in certain wellsupplies. Sodium is extremely soluble and increases its solubility as thetemperature of water rises. Because of this characteristic, sodium saltsdo not form scale when water is heated. Likewise, sodium salts do notproduce curd when combined with soap. In fact, ordinary soap is anorganic sodium compound. As such, it does not react with the sodiumin water.

Soap may be made from a fatty acid and a strong alkali:

C17H35COOH + NaOH —> C17H35COONa + H2O

Stearic Acid Plus Sodium Hydroxide Reacts to Produce Sodium Stirates Plus Water

From this, it is evident that soap is actually a salt formed from anacid and a base.

High concentrations of sodium, on the other hand, mean hightotal minerals and tend to increase the corrosive action of water. Inconcentrations over 30 to 40 grains per gallon, sodium salts may givewater an unpleasant taste. Further, sodium ions in large amounts ham-per the operation of ion exchange softeners in the removal of hard-ness. Where water contains appreciable amounts of both hardnessminerals and sodium, several grains of hardness may continue to ap-pear in the softened water. This occurs because of the regenerative ac-tion of the sodium ions on the ion exchange material.

Reverse osmosis, distillation, and deionization remove sodium fromwater.

Methane

Wells that contain methane are generally located in areas where gasand oil wells are common sights. Amounts run from 0.1 to 11.6 cubicfeet per 1,000 gallons. This is roughly equivalent to 0.8 to 87 millilitersof methane per liter of water. Methane is objectionable in drinkingwater because of the odor and flammability. When water containsmethane gas, it is advisable to aerate it prior to use for either industrialor household purposes. This is necessary to avoid the dangers of fire orexplosion. The aerator must be vented to the open air to permit thegas to escape into the atmosphere.

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Phenol

There is a growing trend by both government and private groups tocontrol pollution of water due to the discharge of industrial waste ma-terials. One of the offensive wastes is phenol. Phenol (C6H5OH) im-parts a medicinal taste and odor to water when the latter is chlorinated.This objectionable taste in a chlorinated water occurs in concentrationsas low as one part per billion due to the formation of chlorophenols,which may be removed by activated carbon filtration.

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DRINKING WATER CONTAMINANTS AND THEIR CONTROLWITH REVERSE OSMOSIS WATER TREATMENT

Nominal Rejection Performance for Reverse Osmosis Membranesat 60 psi Net Pressure and 77°F1

Inorganic CTA* TFC*Contaminant Rejection Rejection

Sodium 85–90% 90–98%Calcium 90–95% 93–99%Magnesium 90–95% 93–99%Potassium 85–90% 90–98%Iron2 90–95% 93–99%Manganese2 90–95% 93–99%Aluminum 90–95% 93–99%Copper 90–95% 93–99%Nickel 90–95% 93–99%Zinc 90–95% 93–99%Strontium 90–95% 93–99%Cadmium 90–95% 93–99%Silver 90–95% 93–99%Mercury 90–95% 93–99%Barium 90–95% 93–99%Chromium 90–95% 93–99%Lead 90–95% 93–99%Chloride 85–95% 90–98%Bicarbonate 85–90% 90–98%Nitrate3 40–50% 85–95%Fluoride 85–90% 90–98%Phosphate 90–95% 93–99%Chromate 85–90% 90–98%Cyanide 85–90% 90–98%Sulfate 90–95% 93–99%Boron 30–40% 55–80%Arsenic+3 60–70% 70–80%Arsenic+5 85–90% 93–99%Selenium 90–95% 93–99%Radioactivity 90–95% 93–99%

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Biological & Particulate Contaminants4

CTA* TFC*Contaminant Rejection Rejection

Bacteria > 99% > 99%Protozoa > 99% > 99%Ameobic Cysts > 99% > 99%Giardia > 99% > 99%Asbestos > 99% > 99%Sediment/Turbidity > 99% > 99%

Organic Contaminants

Organic molecules with amolecular weight > 300 > 90% > 99%Organic molecules with amolecular weight < 3005 0–90% 0–99%

*CTA—Celulosic Membrane*TFC—Thin Film Composite Membrane

1. This table of nominal rejection performance is for the two types of membranes used indrinking water systems operating at a net pressure (feed pressure less back pressure andosmotic pressure) of 60 psi and 77°F water temperature.

The actual performance of systems incorporating these membranes may be less due tochanges in feed pressure, temperature, water chemistry, contaminant level, net pressureon membrane, and individual membrane efficiency.

Countertop RO drinking water systems produce better overall rejection performance thanundercounter systems due to maximizing of net pressure on membrane.

2. While iron and manganese are effectively removed by the membrane, they also can eas-ily foul its surface with deposits even at low concentrations. Generally, iron and manganeseshould be removed by other water treatment methods prior to RO treatment.

3. Nitrate removal depends on factors such as pH, temperature, net pressure across mem-brane, and other contaminants present.

4. While reverse osmosis membranes theoretically remove virtually all known microorgan-isms, including virus, they cannot offer foolproof protection when incorporated into aconsumer drinking water system. Potential seal leaks and manufacturing imperfectionsmay allow some microorganisms to pass into the treated water.

Therefore, small home RO drinking water systems should never be used as a primarymeans of removing biological contamination to make a water supply fit for consumption.

5. The degree of rejection of organic molecules less than molecular weight (MW) 300 de-pends on the size and shape of the molecule. Activated carbon is always incorporatedalong with reverse osmosis to insure complete removal of these lower molecular weight or-ganic contaminants.

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Lesson 8Introduction

Water must have eye appeal and taste appeal before we will drink it withmuch relish. Instinctively, we draw back from the idea of drinking dirty,smelly, cloudy, or discolored water.

Actually, far more important to our well-being is whether or not awater is safe to drink. If it holds disease bacteria, regardless of its clarityand sparkle, we should avoid it.

Let’s consider now these two highly important aspects of water…potability and palatability.

Taste, Odor, and Turbidity

Regardless of any other factors, water piped into the home must bepotable. To be potable, it should be completely free of disease organ-isms. Water is the breeding ground for an almost unbelievably largevariety of organisms. Water does not produce these organisms. Itmerely is an ideal medium in which they can grow. These organismsgain entry into water through a variety of sources. They enter waterfrom natural sources, surface drainage, and sewage. Many of the or-ganisms in water are harmless. In fact, they are extremely beneficialto man. Others have a mild nuisance value. And still others are asource of disease.

In general, those organisms which are potential disease-producersare of primary concern. These are of five types: (1) bacteria, (2) proto-zoa, (3) worms, (4) viruses, and (5) fungi. The presence of certain or-ganisms of these various types can lead to such infectious diseases astyphoid fever, dysentery, cholera, jaundice, hepatitis, giardiasis, undu-lant fever and tularemia, as well as other diseases which spread throughdrinking unsafe water.

Tremendous strides have been made in the control of these diseaseswithin recent years. Much of the credit must go to sanitary engineersfor their careful, consistent control of public water supplies. As proof,outbreaks of typhoid fever in either this country or Canada are rare.Natural disasters can play havoc with water supplies, but under rou-tine conditions typhoid is no longer a serious threat. Paradoxically, thefreedom from typhoid and other similar waterborne diseases makesnecessary even greater vigilance today. For now, whole generationshave grown up without the opportunity to develop a natural immu-nity to such diseases. Thus, a failure in the protective system could re-sult in far more people succumbing to the disease than in the past.

As was previously indicated, many waterborne organisms are ex-tremely beneficial to man. Bacteria, protozoa, and fungi that purifypolluted water are essential to our well-being. Many of these organismsset into motion the chain reactions that result in purification.

We can classify living organisms in many ways and into manygroups. Modern taxonomy categorizes living organisms into five king-doms: Monera, Protista, Fungi, Plantae, and Animalia. Monera includesingle-celled bacteria and photosynthetic blue-green algae. They differfrom all other organisms in that their more primitive cell structureslack a nuclear membrane as well as other membrane-bound organelles.They are called prokaryotes.

All other organisms are eukaryotes, that is organisms with cells thathave distinct nuclei surrounded by nuclear membranes, as well as avariety of other well-defined membranous organelles. Organelles are

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specialized parts of cells (as mitochondria, chloroplasts, or endoplas-mic reticulum) performing functions analogous to those of organs inmany-celled plants and animals.

Members of the kingdom Protista are known as protists. They aresolitary, single-celled eukaryotes (but some species form loose aggre-gations of cells called colonies). Animal-like protists are the protozoa;they are generally larger than bacteria and are mobile. Plant-like pro-tists include several divisions of algae; these contain chlorophyll andcarry on photosynthesis.

The Fungi are a diverse group of eukaryotes that are plant-like butthat cannot carry on photosynthesis. They serve as decomposers, ab-sorbing nutrients from dead leaves or other organic matter in soil andwater. Fungi produce spores during the reproductive process. They con-sist of slime molds, such as the slimy masses found on decaying leavesand wood, and the true fungi, such as molds, yeasts, mildew, andmushrooms.

Plantae (plant) and Animalia (animal) kingdoms consist of the mul-ticellular and well-developed plants and animals we are all familiarwith. Plant cells contain photosynthetic pigments, such as chlorophyll,and plants carry out photosynthesis. Animal cells lack photosyntheticpigments, so animals must obtain nutrients by eating other organisms.

Forms of Lower Life in Water

Algae

These organisms are found throughout the world. Simple algaes exist inthe Monera and Protista kingdoms. Other algaes are plants. They con-stitute single-celled or simple multicellular photosynthetic organismsthat are important producers—produce their own food by using energyfrom sunlight to synthesize complex molecules from carbon dioxideand water—both in sea and fresh water. Algae range in size from mi-croscopic organisms to giant seaweeds several hundred feet in length.They contain chlorophyll and other pigments which give them a va-riety of colors. They manufacture their food by photosynthesis.

Photosynthesis. This is a process of nature in which greenplants use energy from the sun to manufacture carbo-hydrates out of water and carbon dioxide. Only those plantswhich possess chlorophyll (green pigment) are able to carryon the process of photosynthesis.

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Algae thrive well in stagnant surface waters, especially during thewarm weather. Algae give water fishy, grassy, earthy, musty, and othereven more objectionable odors. While algae-laden waters are replusiveto man, animals will drink them, and the presence of blue-green algaehas been known to cause the death of cattle drinking this water.

Diatoms

Diatoms are algal protists belonging to the plant-like (algae) portion ofthe Protista kingdom. Some exist as single cells, others are found asgroups or colonies. More than 15,000 forms of diatoms are known toexist. Diatoms have silica-impregnated cell walls. At times, they releaseessential oils which give water a fishy taste.

Fungi

Fungi have many varieties. Included among these are molds, mildews,mushrooms, yeast, rust, and smut. Fungi are not able to manufacturetheir own food; they contain no chlorophyll. They exist by feeding onliving things or on dead organic matter. Like the bacteria, the fungi areimportant decomposers that breakdown the wastes and the bodies ofdead organisms making their components available for reuse. De-pending on their individual characteristics, they are usually colorless,but may vary in this respect.

Molds

One important category of fungi is molds. This group of fungi feedsentirely on organic matter. They decompose carbohydrates, such assugars, starches, and fats, as well as proteins and other substances. Theythrive ideally in water that has a temperature range of approximately80°F to 100°F. The presence of molds is generally a strong indicator ofheavy pollution of water.

Bacteria

Bacteria are another important class of prokaryotes in the Monerakingdom. Bacteria cells range in size from less than 1 to 10 microns inlength and from 0.2 to 1 micron in width. Despite their small size, ithas been estimated that the total weight of all bacteria in the world ex-ceeds that of all other organisms combined. Bacteria, along with fungi,are an important component of the ecosystem because they decom-pose. If these decomposers did not exist, nutrients would becomelocked up in the dead bodies of plants and animals, and the supply of

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elements required by living systems would soon be exhausted. Amongthe higher organisms in this group are the iron, manganese, and sul-fur bacteria. These higher bacteria gain their energy from the oxida-tion of simple inorganic substances. (Review Lesson 5 for more details.)Lower forms of bacteria can be grouped as those that are helpful andthose that are harmful to man. Those harmful to man are mainly thedisease-producing organisms. Helpful organisms hasten the process ofdecomposing organic waste matter. And by feeding on waste materi-als, they aid in the purifying of water.

All bacteria are sensitive to the temperature and pH of a water.Some bacteria can tolerate acid water. But for the most part, they thrivebest in waters that have a pH between 6.5 and 7.5, that is, essentiallyneutral waters. As to temperature, most pathogenic or disease bacteriathrive best in water of body temperature. Beyond this, no hard and faststatements can be made.

Some bacteria are more resistant to heat than are others. Some aremore sensitive to cold. At low temperatures, for example, some bacte-ria may become dormant for long periods of time, but will still con-tinue to exist. Interestingly enough, the waste products of their owngrowth can hamper bacteria and may even prove toxic to them.

Worms

Worms belong to the animal kingdom. There are three types of worms(flatworms, roundworms [nematodes], and rotifers) found in water. Forthe most part, they dwell in the bed of material at the bottom of lakesand streams. There they do important work as scavengers. The rotifersare the only organisms in this category at or near the surface. They liveprimarily in stagnant fresh water. The eggs and larvae of various intes-tinal worms found in man and warm-blooded animals pollute thewater at times. They do not generally cause widespread infection forseveral reasons: they are relatively few in number and are so large theycan be filtered out of water with comparative ease. The typical size ofparasitic worms or helmiths, such as flukes, tapeworms, hookworms,ascris, pinworms, trichina worms, and filaria worms is 30–50 micronsin diameter.

Protozoa

A basic classification in the Protista kingdom is that group of micro-scopic animal-like protists known as protozoa. These one-celled or-ganisms live mainly in water either at or near the surface or at greatdepths in the oceans. Many live as parasites in the bodies of men andanimals. Like other organisms, protozoa can be classed as helpful or

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injurious. Sometimes drinking water becomes infested with certain pro-tozoa which are not disease-producing. When present, they give thewater a fishy taste and odor. Some protozoa are aerobic, that is, theyexist only where free oxygen is available. Some exist where no free oxy-gen is available. Others can be either aerobic or anaerobic.

Note: One important group of protozoa are those which com-monly form cysts. These protozoans have somewhat bladder-like sacs orvesicles which form a resistant protective wall when they find them-selves in unfavorable surroundings. On entering more favorable sur-roundings (such as the body of a warm-blooded animal), the cystabandons this wall and dwells in the blood stream of the animal. Oneof the most common of these cysts carries the waterborne diseaseamoebic dysentery, for which there is no universal cure at this time.The protists Giardia lamblia and cryptosporidium are a cause of acute gas-trointestinal illness (AGI), which is the most frequently diagnosedwaterborne disease in the United States. The symptoms of the AGI ill-ness, giardiasis, may include severe dehydration, weight loss, and fa-tigue. Giardiasis can persist for several months or longer. Giardiasis isusually associated with unfiltered surface water that has not been dis-infected sufficiently to kill or inactivate the protozoan cysts. Fortunately,these cysts, being typically 2 to 50 microns in diameter, are much largerthan bacteria and can be removed from water by fine filtration.

Nematodes

Nematodes belong to the worm family. They are commonly calledroundworms. Nematodes have long, cylindrical bodies which have nointernal segments. Interestingly enough, those nematodes, which arefound in the bodies of men and warm-blooded animals, are largeenough to be visible to the naked eye. Those living in fresh water andin the soil are microscopic. Nematodes can be a problem in drinkingwater because they impart objectionable tastes and odors to water. Theyare also under suspicion of being carriers of the type of disease-bearingbacteria found in the intestines of warm-blooded animals, though stud-ies show, however, this possibility is somewhat remote. Nematodes areapt to be found in municipal waters derived from surface supplies.

Viruses

The smallest of the infectious microorganisms is that group of parasiticforms known as viruses. Too small to be seen under a microscope,viruses are capable of causing disease in both plants and animals.Viruses can pass through porcelain filters that are capable of screening

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out bacteria. Viruses, such as those producing infectious hepatitis,poliomyelitis, meningitis, and gastroenteritis, can be waterborne.Drinking water contaminated with any of these viruses is hazardous.

Virus. A minute (0.004 to 0.1 micron in diameter) infectiousagent that is much smaller than bacteria. Viruses are gen-erally considered parasites that are incapable of growth ex-cept in the presence of living cells. They can be preservedindefinitely even when frozen or dried.

As you can see from even this brief summary, there is a tremendousvariety of living organisms in water. To understand and classify theircountless varieties requires an immense amount of knowledge andtime. Where these organisms are pathogenic or disease-producing, theymay make water unsafe to drink. For obvious reasons, even where thereis just a possibility that water contains pathogenic organisms, it mustbe considered contaminated. While there is a large and varied numberof pathogens, no single contaminated water supply is apt to containmore than a few varieties at a time. On one hand, this is fortunate. Butat the same time, it makes detection of pathogens extremely difficultin terms of a routine water analysis.

Note: Not only are speed and accuracy essential in testing sourcesof drinking water for purity, but frequency is also highly important.Municipal systems run tests on a sliding scale: the more inhabitantsthere are in the community, the more frequent the tests. A sanitary en-gineer for a community of 10,000 would be required to run a mini-mum of ten tests a month; an engineer for a city of 1,000,000 wouldrun at least 300 water sample tests a month.

Tests on private water systems are more seldom. In many cases,one sample is taken. If it shows the water is safe, no further tests maybe made on that well. Unfortunately, a serious limitation of coliformbacteria tests is that they indicate the condition of a given sample andno more. Once a test shows lack of contamination, there is no guar-antee water cannot become contaminated even within a short time.Proper location and construction of a well are important factors.Equally vital is regular disinfection of the water and frequent con-tamination tests.

Since both speed and accuracy are essential, laboratory scientistsneed a sure way to expedite detection of pathogens. They have adependable answer in a group of readily identified organisms thatindicate possible contamination. These indicator organisms are thefecal coliform bacteria.

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Fecal coliform bacteria. Fecal coliform bacteria, such asEscherichia coli (E. coli), grow in the intestines of human be-ings and other warm-blooded animals. Since they are dis-carged in astronomical numbers (approximately 400 billionper day) in human excrement, their presence in a water sam-ple is an indication of human sewage.

Actually the total number of coliform bacteria that mayenter a source of drinking water is reduced by several fac-tors: (1) they die in large numbers because they cannot gen-erally maintain themselves outside the warm-blooded bodyor in cleaner water; (2) they are removed in water purifica-tion processes; and (3) they are destroyed in sewage treat-ment operations.

Research has shown that the presence of fecal coliform bacteria in-dicates the entrance of human or animal wastes into water since col-iform bacteria naturally exist in the intestines of humans and certainanimals. Thus, the presence of these bacteria in water is accepted asproof that the water has been contaminated by human or animalwastes. Although such water may contain no pathogens, an infectedperson or animal, or a carrier of disease, could introduce pathogens atany moment, and immediate corrective action must be taken. The pres-ence of fecal coliform bacteria shows water is contaminated by humanwastes and is potentially contaminated with pathogens. In short, thesebacteria are a measure of guilt by association.

Conversely, the absence of coliform bacteria does not assure absenceof pathogens, but their presence is considered unlikely. Just how canwater be tested for the presence of coliform bacteria? These organismscause the fermentation of lactose (the crystalline sugar compound inmilk). When water containing coliform bacteria is placed in a lactose cul-ture, it will cause fermentation resulting in the formation of gas. Thisconfirms the suspicions. Or, coliform bacteria can be captured on a 0.45micron membrane filter and grow into colonies large enough to see andcount when incubated with a suitable agar growth medium or broth.

A recently developed coliform and E. coli analytical method makesuse of the fact that coliform bacteria react with o-nitrophenyl-B-d-galac-topyranoside (ONPG) to form a visible yellow color, and E. coli reactwith 4-methylumbelliferyl-B-d-glucuronide (MUG) to produce visiblefluorescence under long wave ultraviolet light. By incorporating thesechemicals along with growth broth in test tubes to be innoculated withthe water sample, both total coliform and E. coli can be easily and ac-curately detected within 24 hours.

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The EPA Primary Drinking Water Regulations specify that watershould contain an absence of coliform organisms or no more than onecoliform-positive sample result when 5 to 39 samples are analyzed ina month, and no more than 5 percent coliform-positive sample resultswhen 40 or more samples are analyzed each month.

Note: The standard of approximately one or 5 percent coliform-positive sample results is, of course, a standard of expediency. Witheven a single organism of this type in the water, there is always the pos-sibility of infection.

Recognizing the danger, what can be done to provide adequate pro-tection against contamination? When a water supply becomes con-taminated, correct the problem at once. This means going beyondtreatment alone, important as this may be. It is a basic rule of watersanitation to get to the source of the problem and eliminate it. If a well,for example, becomes badly contaminated, it is necessary to trace thecontamination to its source and, if possible, remedy the situation. Itmay even be necessary to seek out a new source of supply.

Note: Coliform bacteria were selected as a biological indicator ofcontamination or pollution because they satisfied the following re-quirements:

1. An organism serving as a reliable measure of contaminationmust indicate the potential presence of specific contami-nating organisms in either a natural water or one subjectedto treatment. Such an organism must react exactly as do thecontaminating organisms both in the natural water supplyand in a treated water.

2. The indicator organism must be present in greater numberthan is the contaminating organism. Unless this is true, thecontaminating organism itself would serve the same purposemore directly.

3. The indicator organism must be readily identifiable bymeans of relatively simple analytical tests.

4. It is also important to evaluate the quantity of indicator or-ganisms in the water since the degree of contamination isan important factor.

Treatment of a water supply is a safety factor, not a corrective meas-ure. Keep this in mind in the discussion that now follows.

There are a number of ways of purifying water. In evaluating themethods of treatment available, the following points regarding waterdisinfectants should be considered:

1. A disinfectant should be able to destroy all types ofpathogens and in whatever number present in the water.

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2. A disinfectant should destroy the pathogens within the timeavailable for disinfection.

3. A disinfectant should function properly regardless of anyfluctuations in the composition or condition of the water.

4. A disinfectant should function within the temperature rangeof the water.

5. A disinfectant should not cause the water to become toxicor unpalatable.

6. A disinfectant should be safe and easy to handle.7. A disinfectant should be such that it is easy to determine its

concentration in the water.8. A disinfectant should provide residual protection against re-

contamination.

Techniques such as filtration may remove infectious organismsfrom water. They are, however, no substitute for disinfection.

Water Disinfection Methods

The following are specific methods for disinfecting water.

Boiling Water

Place water in a container over heat. Bring it to the boiling point. Holdit at this temperature for at least one minute. This will disinfect thewater. Perhaps you have used this technique after a flood or when awater main has burst as an emergency aid. Boiling water is an effectivemethod of treatment because no important waterborne diseases arecaused by heat-resisting organisms.

Ultraviolet Light

The use of ultraviolet light is an attempt to imitate nature. As you re-call, sunlight destroys some bacteria in the natural purification of water.Exposing water to ultraviolet light destroys pathogens. To assure thor-ough treatment, the water must be free of turbidity and color. Other-wise, some bacteria will be protected from the germ-killing ultravioletrays. Since ultraviolet light adds nothing to the water, there is little pos-sibility of its creating taste or odor problems. Similarly, ultraviolet lighttreatment has no residual effect. Further, it must be closely checked toassure that sufficient ultraviolet energy is reaching the point of appli-cation at all times.

Advantages of ultraviolet light: automatic, no taste or odor,and low contact time.

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Disadvantages of ultraviolet light: low penetration power,shielding by turbidity, deposits may develop and coat the tube, no sim-ple operational test of disinfectant dosage, no longer lasting residualeffects, and ultraviolet tube gradually loses power.

Various Chemical Disinfectants

The most common method of disinfecting water for contamination isto use one of the various chemical agents available. Among these arechlorine, bromine, iodine, potassium permanganate, copper and silverions, alkalis, acids, and ozone. Let us review them briefly here.

Bromine

Bromine is an oxidizing agent that has been used quite successfully inthe disinfecting of swimming pool waters. It is rated as a good germi-cidal agent. Bromine is easy to feed into water and is not hazardous tostore. It apparently does not cause eye irritations among swimmers,nor are its odors troublesome.

Chlorine

One of the most widely used disinfecting agents to insure safe drink-ing water is chlorine. Chlorine in cylinders is used extensively by mu-nicipalities in water disinfection. However, in this form, chlorine gas(Cl2) is far too dangerous for any home purposes.

For use in the home, chlorine is readily available as sodiumhypochlorite (household bleach) which can be used both for launder-ing or disinfecting purposes. This product contains a 5.25 percentsolution of sodium hypochlorite which is equivalent to 5 percent avail-able chlorine.

Chlorine is also available as calcium hypochlorite, which is sold inthe form of dry granules. In this form, it is usually 70 percent availablechlorine. When calcium hypochlorite is used, this chlorinated limeshould be mixed thoroughly and allowed to settle, pumping only theclear solution. For a variety of reasons, not the least of which is con-venience, chlorine in the liquid form (sodium hypochlorite) is morepopular for household use. Chlorine is normally fed into water withthe aid of a chemical feed pump.

The first chlorine fed into the water is likely to be consumed in theoxidation of any iron, manganese, or hydrogen sulfide that may be pres-ent. Some of the chlorine is also neutralized by organic matter normallypresent in any supply, including bacteria, if present. When the “chlorine

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demand” due to these materials has been satisfied, what’s left over—thechlorine that has not been consumed—remains as a “chlorine residual.”

Chlorination terms. There are three basic terms used inthe chlorination process: chlorine demand, chlorine dosage,and chlorine residual.

Chlorine demand is the amount of chlorine which willbe reduced or consumed in the process of oxidizing impu-rities in the water.

Chlorine dosage is the amount of chlorine fed into thewater.

And chlorine residual is the amount of chlorine still re-maining in water after oxidation takes place.

For example, if a water has 2.0 ppm chlorine demand,and a chlorine dosage of 5.0 ppm is fed into the water, thechlorine residual would be 3.0 ppm.

The rate of feed is normally adjusted with a chemical feed pumpto provide a chlorine residual of 0.5–1.0 ppm after 20 minutes of con-tact time. This is enough to kill coliform bacteria, but may or may notkill all viruses or cysts which may be present. Such a chlorine residualnot only serves to overcome intermittent trace contamination from co-liform bacteria, but also provides for minor variations in the chlorinedemand of the water. The pathogens causing such diseases as typhoidfever, cholera and dysentery succumb most easily to chlorine treat-ment. Cyst-forming protozoa which cause amoebic dysentery, cry-tosporidiosis, and giardiasis are most resistant to chlorine.

Viruses are also susceptible, similar to bacteria, to inactivation bychlorine disinfects.

Iodine

For emergency purposes, iodine may be used for treatment of drinkingwater. Much work at present is being done to test the effect of iodine indestroying viruses, which are now considered among the pathogensresistant to treatment. Tests show that 20 minutes exposure to 8.0 ppmof iodine is adequate to render a potable water. As usual, the residualrequired varies inversely with contact time. Lower residuals requirelonger contact time, while higher residuals require shorter contact time.While such test results are encouraging, not enough is yet known aboutthe physiological effects of iodine-treated water on the human system.

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For this reason, its use must be considered only on an emergency orshort-term exposure basis.

People who suffer from hyperthyroidism (one percent of the fe-male population and 0.1 percent of the male population) may be atrisk from exposures to over 700 micrograms of iodine per day. Excessiodine and iodide levels from water treatment should, therefore, bescavenged from water that is to be utilized for long-term drinking watersupplies.

Silver

Silver in various forms has been used to inhibit the growth of mi-croorganisms. It is most frequently found combined with activatedcarbon in filters. When some bacteria species come into contact withthis silver, they are rendered inactive. There is disagreement amongthe experts as to the effectiveness of this process because silver ions inwater kill E.coli very well and probably also salmonella, shigella, andvibrio bacteria, but it has found lesser effect on viruses, cysts, and otherbacteria species. Silver does not produce offensive tastes or odors whenused in water treatment. Further, organic matter does not interfere withits effectiveness as is the case with free chlorine. Its high cost, inter-ferences by chlorides and sulfides, need for long periods of exposure,and incomplete bactericidal action have hindered its widespread ac-ceptance.

Copper

Copper ions are used quite frequently to destroy algae in surface wa-ters. But these ions are relatively ineffective in killing bacteria. Coppersulfate, for example, is also used to kill algae in reservoirs.

Alkalies and Acids

Disease-bearing organisms are strongly affected by the pH of a water.They will not survive when water is either highly acid or highly alka-line. Thus, treatment which sharply reduces or increases pH in relationto the normal range of 6.5 to 7.5 can be an effective means of reduc-ing microorganisms.

Other Agents

There are numerous other agents which have proved to be successfulin destroying pathogens. Many of these must still be subjected to pro-

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longed testing with regard to their physiological effect on man. Amongthese are certain surfactants which aid in destroying pathogens. Thecationic detergents readily kill pathogens. Anionic detergents are onlyweakly effective in destroying pathogens. Because of their objection-able flavor and possible toxic effects, however, surfactants have notbeen seriously considered for treating drinking water.

Chlorine dioxide has unusually good germ killing power. Becauseit is such a strong oxidizing agent, a larger residual of chlorine dioxidewould probably be needed than is the case with chlorine.

At present, chlorination in one form or another is regarded as themost effective disinfectant available for all general purposes. It has fullacceptance of health authorities. Still, there are certain factors whichaffect its ability to disinfect waters. These should always be kept inmind. They are:

1. “Free” chlorine residuals are more effective than “combined”or “chloramine” residuals. Disinfection, regardless of thetype of chlorine, becomes more effective with increasedresiduals.

Chloramine. The compound formed by feeding both chlo-rine and ammonia to the water. This treatment has beenused for controlling bacterial growth in long pipe lines andin the applications where its slower oxidizing action is ofparticular benefit.

2. A pH of 6.0 to 7.0 makes water a far more effective mediumfor chlorine as a disinfecting agent than to higher pH valuesof around 0 to 10.0.

3. The effectiveness of disinfection increases with the amountof contact time available.

4. The effectiveness of chlorine residuals increases with highertemperatures within the normal water temperature range.

5. All types of organisms do not react in the same way undervarious conditions to chlorination.

6. An increase in the chlorine demand of a water increases theamount of chlorine necessary to provide a satisfactory chlo-rine residual.

7. Chlorination of drinking water always results in the forma-tion of trihalomethanes and other chlorination productswhich cause adverse birth outcome and cancer risks. USEPAlimits trihalomethanes to 0.08 mg/L and haloacetic acids to0.06 mg/L in safe drinking water.

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In order to insure the destruction of pathogens, the process of chlo-rination must achieve certain control of at least one factor, and prefer-ably two, to compensate for fluctuations that occur. For this reason,some authorities on the subject stress the fact that the type and con-centration of the chlorine residual must be controlled to insure ade-quate disinfection. Only in this way, they claim, can chlorinationadequately take into account variations in temperature, pH, chlorinedemand, and types of organisms in the water. While possible to in-crease minimum contact times, it is difficult to do so. Five to ten min-utes is normally all the time available with the type of pressure systemsnormally used for small water supplies. Superchlorination is a meansto achieve greater disinfection with shorter contact times.

Briefly, what is this technique, and how does it operate?The success of superchlorination-dechlorination depends on put-

ting enough chlorine in the water to provide a residual of 3.0 to 5.0ppm. This is considerably greater than a chlorine residual of 0.1 to 0.5ppm usually found in municipal water supplies when drawn from thetap. A superchlorination-dechlorination system consists of two basic

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Typical layout of superchlorination-dechlorination equipment on a private watersystem. With dechlorinator on main waterline, this arrangement can be used whereiron and/or manganese are present in the water.

units. A chlorinator feeds chlorine into the raw water. This chlorinefeed is stepped up to provide the needed residual. A dechlorinator unitthen removes the excess chlorine from the water before it reaches thehousehold taps.

The chlorinator should be installed so that it feeds the chlorineinto the water before it reaches the retention tank. A general purposechemical feed pump (such as described in Lesson 5) will do the job.The size and the placement of the dechlorinator unit depends on thetype of treatment necessary. This will usually be an activated carbonfilter. If pathogen kill is all that is required, a small dechlorinator canbe installed at the kitchen sink. This unit then serves to remove chlo-rine from water used for drinking and cooking. Since many familiesalso drink water from bathroom taps, it may be necessary to installdechlorinators at these locations as well. The advantage in dechlori-nating only a part of the water is obvious. A smaller filter unit doesthe job. And since only a small portion of the total water is filteredunder such conditions, the unit lasts longer before either servicing orreplacement is necessary. Essentially, dechlorination is not needed toinsure a safe drinking water. Once the water is chlorinated, the healthhazard is gone. The chlorine residual is removed merely to make thewater palatable.

If the problem is compounded due to the presence of iron and/ormanganese, all the water must be filtered. Under such conditions, alarge central filter is necessary and should be placed on the main lineafter the pressure tank.

The prime advantage of the superchlorination-dechlorinationprocess is that it saturates water with enough chlorine to kill bacteria.

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Typical layout of superchlorination-dechlorination equipment on a private watersystem. With dechlorinator on main waterline, this arrangement can be used whereiron and/or manganese are present in the water.

Simple chlorination sometimes fails its objective because homeownersmay set the chlorine feed rate too low in order to avoid giving theirwater a chlorine taste.

We have discussed, at some length, various types of pathogens andmethods of destroying them in the process of making water potable—safe to drink. This is highly important, but it is not the whole story;for water must be palatable as well as potable.

Questions and Answers on Dechlorinating Filters

How long will a dechlorinating filter operate?With a bad iron problem, it may be necessary to service the filter fre-quently due to the clogging of the filter with precipitated iron. Whenneither iron or sulfur is present, a filter continues to operate for longperiods of time. How long depends on the amount of turbidity in thewater and the design of the unit. With reasonably clear water, a poudechlorinating filter should be good for 5,000 gallons of water, morethan the average household uses for cooking and drinking in a year.

How can you service a dechlorinator?Some units can be backwashed. Others are so designed that the home-owner can remove the filter from the container and hose it off. Stillothers are disposable units.

What happens when a dechlorinator ceases to operate?The homeowner quickly knows his unit is not working because he willtaste chlorine in the water, or the pressure drop will become excessive dueto plugging by turbidity. He should then backwash or replace the filter.

Palatability

What makes a water palatable? To be palatable, a water should be freeof detectable taste and odors.

What constitutes a detectable taste or odor? Undoubtedly, you havetasted waters which have had unpleasant tastes or odors. Natives in thearea may be surprised to note your reaction. For after drinking thewater for many years, they find nothing peculiar to either the taste orodor of the water. And then there are those waters which have tastesand odors so obnoxious (hydrogen sulfide water, for example) even thelong-time inhabitant can’t stomach them.

Turbidity, sediment, and color also play important roles in deter-mining whether a water is palatable.

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Odors and Tastes

Various odors and tastes may be present in water. They can be tracedto many conditions. Unfortunately, the causes of bad taste and odorproblems in water are so many, it is impossible to suggest a single treat-ment that would be universally effective in controlling these problems.

Tastes are generally classified in four groups—sour, salt, sweet, andbitter. Odors, on the other hand, possess many classifications. Thereare some 20 of them commonly used, all possessing rather picturesquenames. In fact, the names in many cases, are far more pleasant thenthe odors themselves. To name a few of them—nasturtium, cucumber,geranium, fishy, pigpen, earthy, grassy, and musty. Authorities furtherclassify these odors in terms of their intensity from very faint, faint,distinct and decided, to very strong,

All taste buds and olfactory organs are not necessarily of the sameacuteness, but generally you should not be aware of any tastes or odorsin water if there is to be pleasure in drinking it. If you are conscious ofa distinct odor, the water is in need of treatment.

In many cases, it is difficult to differentiate between tastes andodors. Both the taste buds and olfactory organs work so effectively to-gether it is hard to determine where one leaves off and the other be-gins. To illustrate: hydrogen sulfide gives water an “awful” taste, yetactually, it is the unpleasant odor of this gas that we detect rather thanan unpleasant taste. Unfortunately, there is little in the way of stan-dard measuring equipment for rating tastes and odors. Tastes and odorsin water can be traced to a number of factors. They include:

1. decaying organic matter;2. living organisms;3. iron, manganese, and the metallic products of corrosion;4. industrial waste pollution from substances such as phenol;5. chlorination;6. high mineral concentrations;7. dissolved gases.

In general, odors can be traced to living organisms, organic matterand gases in water. Likewise, tastes can be traced generally to the hightotal minerals in water. There are, however, some tastes due to variousalgae and industrial wastes. Now how can these objectionable tastesand odors be removed from water?

Some tastes and odors, especially those due to organic substances,can be removed from water simply by passing it through an activatedcarbon filter. Other tastes and odors may respond to oxidizing agentssuch as chlorine, ozone, and potassium permanganate. Where theseproblems are due to industrial wastes and certain other substances,

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some of the above types of treatment may completely fail. In somecases, for example, chlorination may actually intensify a taste or odorproblem. Potassium permanganate and ozone have been found to beextremely effective in removing many musty, fishy, grassy, and moldyodors. Two factors make this compound valuable—it is a strong oxi-dizing agent, and it does not form obnoxious compounds with organicmatter. However, a filter must be used to remove the manganese diox-ide formed when the permanganate is reduced.

In any case, you may have to try a number of methods in an at-tempt to rid a water of objectionable tastes and odors. If methods con-sidered here do not work, it may be more economical to seek out a newsource of drinking water.

Turbidity

Pick up a glass of water and hold it to the light. Can you see any finelydivided, insoluble particles suspended in the water? Or does the waterseem hazy? If so, the water is turbid.

Turbidity. Turbidity and suspended matter are not synony-mous terms, although most of us use the terms more or lessinterchangeably. Correctly speaking, suspended matter isthat material which can be removed from water through fil-tration or the coagulation-filtration process. Turbidity, onthe other hand, is a measure of the amount of light scat-tered and absorbed by water because of the suspendedmatter in the water.

There is also some danger of confusion regarding turbid-ity and color. Turbidity is the lack of clarity or brilliance in awater. Water may have a great deal of color—it may even bedark brown—and still be clear and without suspended matter.

When water has a large amount of such suspended particles, welose our zest for it. While it may be safe to drink, it seems offensive.The EPA Primary Drinking Water Regulations specify that turbidity ofa potable water be less than 1 unit in 95 percent of all samples and lessthan 5 units under special conditions. The suspended particles cloud-ing the water may be due to such inorganic substances as clay, rockflour, silt, calcium carbonate, silica, iron, manganese, sulfur, or indus-trial wastes. Again, the clouding may be caused by organic substancessuch as various microorganisms, finely divided vegetable or animalmatter, grease, fat, oil, and others.

While turbidity may be due to a single foreign substance in water,chances are it is probably due to a mixture of several or many sub-stances. These particles may range in size from fine colloidal materialsto coarse grains of sand that remain in suspension only as long as thewater is agitated. Those particles which quickly sink to the bottom areusually called sediment. There are, however, no hard and fast rules forclassifying such impurities. Water from a swiftly flowing river or streamcan contain a considerable amount of sediment. In contrast, watertaken from a lake or pond is usually much clearer. In these more quiet,nonflowing waters there is greater opportunity for settling action. Thus,all but very fine particles sink to the bottom. Least apt to contain sed-iment are wells and springs. Sediment is generally strained from thesewaters as they percolate through sand, gravel, and rock formations.

Turbidity varies tremendously even within these various groupings.Some rivers and streams have water that appears crystal clear with justtrace amounts of turbidity in them, especially at points near theirsources. These same moving waters may contain upwards of 30,000ppm of turbidity-causing suspended solids at other points in theircourse to the oceans. In fact, turbidity in amounts well over 60,000ppm of turbidity sediment has been registered.

Again, there are significant fluctuations in the amount of turbid-ity in a river at different times in a year. Samples taken at various timesfrom the same site show the Allegheny River carried from a high of 65ppm of turbidity-causing suspended material to a low of 2 ppm. Theaverage ran around 21 ppm. The Flint River at Albany, Georgia rangedfrom a high of 560 ppm to a low of 12 ppm. An even more striking ex-ample of these seasonal variations in turbidity is provided by theArkansas River. At Arkansas City, Kansas the high was 27,500 ppm ofturbidity-causing suspended solids, the low 14 ppm and the average2,300 ppm.

Heavy rainfalls, strong winds, and convection currents can greatlyincrease the turbid state of both lakes and rivers. Warm weather andincreases in the temperature can also add to the problem. For withwarmer weather, microorganisms and aquatic plants renew their ac-tivity in the water. As they grow and later decay, these plant and animalforms substantially add to the turbid state of a water. Also, they fre-quently cause an increase in odor and color problems.

Mechanical Filtration

Mechanical filtration will remove almost all forms of turbidity. Ofcourse, the smaller the turbid particles, the finer must be the filteropenings in order to strain them out. Under some circumstances, the

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openings have to be so small that they cause an excessive pressure dropas the water percolates through the filter, and the unit may be im-practical. In many cases, filters containing specially graded and sizedgravel and sand are effective in screening out turbid particles. Withsuch units a periodic backwashing to remove the filtered material is allthe maintenance necessary. (See drawing above.) Some filter manufac-turers also provide a “filter-aid” which is added onto the top of the fil-ter bed immediately after backwashing. The filter-aid traps fine dirtparticles, producing a more sparkling clear water and keeps dirt frompenetrating the filter bed, insuring better bed cleansing during back-washing. In some cases, cartridge filters are effective. These will filterwater used for cooking and drinking or as pretreatment sediment re-moval with other water treatment technologies. Generally, these car-tridge filters are just installed on the drinking water lines. There areseveral reasons for this: (1) they produce a significant pressure drop.This would be a handicap if these filters were installed on the mainwaterline. (2) their replacement cost is higher when filtering all thewater.

Municipal and industrial systems frequently make use of the co-agulation process to aid in the removal of turbidity. In this economi-cal process, a coagulating agent, such as aluminum sulfate, is fed into

Graded sand and gravel filters, such as the one shown in this schematic, areeffective in clearing up mild turbidity.

the water. After rapid mixing, the coagulating agent forms a “floc”generally in the form of a gelatinous precipitate. This floc gives the ap-pearance of a soft, gentle snowfall. A settling period is then needed toallow the floc to fall gently through the water. As the floc forms andsettles, it tends to collect or entrap the turbid particles and form theminto larger particles which sink to the bottom or can be captured by afilter. On large installations, settling basins provide the necessary timeand space for the process. After the settling period, the water flowsthrough a filter to remove the last traces of the coagulant and any re-maining turbid particles. Small coagulation-filtration systems are some-times utilized for household purposes when turbidity is particularlyoffensive. The difficulties with the use of this process for home pur-poses are in determining what type of coagulant will give best results,and creating flow that won’t break apart with the varying flow ratesand turbulences common with household water usages. Experience inthis area is essential because the chemical properties of the water mustbe considered in relation to the coagulant or combination of coagu-lants employed. Further, adequate mixing and coagulation times areessential. Coagulation-filtration also requires considerably more at-tention to maintenance than do the simple filtration processes men-tioned earlier.

MultiMedia Filters (Depth Filters)

Multimedia filters represent a significant improvement over single-media filters. This is due primarily to improved filter bed action basedon the innovative use and selection of filter media. Multimedia filtra-tion permits delivery of high quality filtered water at much faster flowrates, as compared to a conventional sand filter.

In a conventional sand filter, lighter and finer sand particles arefound at the top of the filter bed, and coarser, heavier sand particlesremain at the bottom after backwashing. Filtration takes place in thetop few inches of the filter bed.

The multimedia filter is radically different. The multimedia filterbed, in comparison to the sand filter bed, is upside down. Coarse, butlighter, particles backwash to the top, whereas finer, but heavier, par-ticles remain at the bottom of the bed. The innovation lies in the se-lection of suitable media. This configuration has many advantages. Theentire bed acts as a filter, rather than only the top few inches. Turbid-ity is trapped throughout the bed, enabling the filter to hold far moresolids filtered from the water before backwashing is necessary.

Typically, the filter bed is made up of three layers of filter media.The total bed depth is about 26 to 40 inches. In a three layer filter, the

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top layer is made up of large, lighter weight particles of anthracite coaland is from 15 to 18 inches in depth (particle size 1.0 to 1.5 milli-meters, density 1.35 to 1.75). The middle layer contains from 8 to 15inches of heavier and smaller particles of calcined aluminum silicateor sand (particle size 0.5 to 0.6 millimeters, density 2.65). The bottomlayer contains from 3 to 6 inches of heavier garnet (particle size 0.2 to0.3 millimeters, density 4.0 to 4.2). This semiprecious red silicate min-eral is 50 to 60 percent heavier than sand.

A multimedia filter is backwashed in the same manner as a sandfilter, using reverse or upward flow of water through the filter bed. Thevarious layers of media retain their stratification because each materialhas a different density.

In a four-media filter a fourth or top layer contains from 3 to 6inches of lighter and larger plastic pillows (particle size 2.0 to 4.0 mil-limeters, density 1.1 to 1.2). Their density is slightly above the densityof water which is 1.0.

Advantages

1. The multimedia filter can operate for much longer periodsof time (five or more times as long at the same filtration rate),

MultiMedia Filter

before backwashing is necessary because the bed can holdmore turbidity. Turbidity is trapped and held throughout theentire bed depth, rather than the top one or two inches.

2. Multimedia filtration is much better suited for use in a closedpressure tank since cracking of the bed, and subsequentbreakthrough of turbidity is greatly reduced and the needfor visual inspection is less necessary.

The use of pressure tanks, rather than open basins or fil-ters, is an obvious advantage for point-of-use filtration andcould also be of real importance in the filtration of smallcommunity water supplies.

More rapid filtration flow rates in multimedia filtrationallow the use of smaller diameter tanks with equal or betterresults.

3. A very high degree of clarity is achieved in the filtered waterbecause of the fact that the finer particles of garnet at thebottom trap finer turbidity particles.

4. Another important advantage is that the multimedia filtercan clarify water at a much higher flow rate than a single-media sand filter (5.5 to 8 gallons per minute, as comparedto 1.5 to 3 gallons per minute in a 12 inch diameter tank).This is 7 to 10 gpm per square foot of bed area, as comparedto 2 to 4 gpm per square foot of bed area. These and othervery important differences between a single media sand fil-ter and a multimedia depth filer are shown in the chartbelow.

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Central Systems

In small community water supply filtration the conventional massivesedimentation tank, which allows larger particles of turbidity to settle,is replaced by the centrifugal separator which does a similar job in 1 percent of the space.

Centrifugal separators have been used in mining and mineral re-covery for many years. Solid particles entering the separation chamberare acted on by high centrifugal forces which move the particles to theouter separator walls and then down to a collection device at the bot-tom. At the same time, the clarified water moves toward the center ofthe separation chamber and upward to the clear water outlet at the top.

Separators can remove up to 98 percent of all suspended particles,down to a particle size as small as three thousandths of an inch (74 mi-crometers). A human hair has a thickness of about 100 micrometers.

In the multimedia filter, the traditional feed of alum as a coagulantis reduced. At the same time, it is supplemented with a polymer (poly-electrolyte) which forms a stronger floc and is applicable over a broaderturbidity range.

Contact Clarification

A separate tank, called a contact clarifier, provides hydraulic contactflocculation and surface storage clarification. This replaces traditionalpaddle flocculation and four hours of quiescent clarification. The sandfilter which followed in the traditional system has been replaced by amore efficient multimedia filter. Thus, without process shortcuts,process time has been reduced from the traditional 41⁄2–6 hours toabout 10 minutes!

Typical results for multi-media systems include reduction from 200NTU to 0.42 NTU on a high turbidity water, and from 25 NTU to 0.15NTU on a low turbidity water.

Color

What color is water?Ordinarily, we think of water as being blue in color. When artists

paint bodies of water, they generally color them blue or blue-green.While water does reflect blue-green light, noticeable in great depths, itshould appear colorless as used in the home.

Ideally, water from the tap is not blue or blue-green. If such is thecase, there are certain foreign substances in the water. Infinitely smallmicroscopic particles add color to water. Colloidal suspensions and

noncolloidal organic acids as well as neutral salts also affect the colorof water. The color in water is primarily of vegetable origin and is ex-tracted from leaves and aquatic plants. Naturally water draining fromswamps has the most intense coloring. The bleaching action of sun-light plus the aging of water gradually dissipates this color, however.All surface waters possess some degree of color. Likewise, some shal-low wells, springs and an occasional deep well can contain notice-able coloring. In general, however, water from deep wells is practicallycolorless.

An arbitrary standard scale has been developed for measuring colorintensity in water samples. When a water is rated as having a color of5 units, it means: the color of this water is equal in intensity to thecolor of distilled water containing 5 milligrams of platinum as potas-sium chloroplatinate per liter. Highly colored water is objectionable formost process work in the industrial field because excessive color causesstains. And while color is not a factor of great concern in relation tohousehold applications, excessive color lacks appeal from an aestheticstandpoint in a potable water. USEPA Secondary Drinking Water Reg-ulations recommend that a potable water possess color of less than 15units. In general, color is reduced or removed from water through theuse of coagulation, settling, and filtration techniques. Aluminum sul-fate is the most widely used coagulant for this purpose. Superchlori-nation, activated carbon filters, ozone, and potassium permanganatehave been used with varying degrees of success in removing color.

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Lesson 9Introduction

By now, it is obvious that water plays host to an almost endless list ofcontaminants and impurities. When various combinations of these con-taminants are considered, the possibilities are virtually infinite.

Fortunately, there are certain rather well defined patterns into whichmost waters can be grouped for greater ease of treatment. In this lesson,we will consider analysis of a variety of waters with the purpose of set-ting up a sound basis of treatment in each case.

How to Interpret Water Analyses

Draw a glass of water from the tap. Examine it closely. Hold it up tothe light. Sniff it. And, finally, take a drink of it. On the basis of yourexamination, are you in a position to tell much about this water?

Such sensory tests cannot go far in analyzing water. The look,taste, and smell of a water often give no clues as to the types andamounts of some contaminants. Some contaminants, of course, areobvious. Your eyes will quickly tell you when there is a significantamount of sediment in the water. And your nose will not fool youwhen there is hydrogen sulfide present. Likewise, algae, industrialwastes, and certain other contaminants will also give water distinc-tively unpleasant tastes. Such contaminants as calcium, magnesium,fluorides, nitrate nitrogen, ferrous iron, and free oxygen provide in-dicators that they are present in water. Actually, it is probably onlythrough their after-effects that we know for sure that a certain rawwater contains these troublesome contaminants. We would have noway of knowing the quantities.

Test for Water Hardness

Take a small bottle or jar (about two ounce capacity) andput in it one ounce of water. Add liquid soap (use a tinctureof green soap Lilly No. 100 or U.S.P. XV, which can be ob-tained at any drug store), one drop at a time. After eachdrop, cap the bottle and shake. Continue to add soap untila good quantity of lasting suds has been formed. The num-ber of drops of soap required to form suds will indicate theapproximate hardness of the water in grains per gallon.(One drop equals one grain.) If more than one drop of soapis added, the water is considered hard. Soft water will formsuds after the addition of one drop of soap.

Even the presence of harmful bacteria may go completely unno-ticed. In fact, if a given family has built up an immunity to certain dis-ease-producing bacteria in the water, it is possible they may neversuffer any ill effects from drinking it. Unfortunately, guests in the homemay not enjoy such immunity. Thus, there is the distinct possibility oftheir becoming ill after drinking such a contaminated water. The bestway to determine the kinds and amounts of various contaminants inwater is by means of laboratory analysis. Such analysis may be eitherphysical, microbiological, or chemical. Where the homeowner main-tains his own water system, he is well advised to have periodic bacte-

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riological tests run. Such tests can assure him that the water is safe todrink, though admittedly a source could be contaminated at any time.

Municipal water systems make regular analyses of their water sup-plies, the larger the system, the more frequent the tests. The program ineach city conforms to EPA Drinking Water Regulations.

While there are no controls requiring periodic bacteriological testsover private water systems, wise owners of private water systems dohave such tests run on their drinking water at least once a year. A bac-teriological test of a water sample will show the possible presence ofdisease organisms. A report will normally indicate whether the wateris potable or not from a bacterial standpoint. Such tests may be run fora small fee. Generally, any state or city health department has thefacilities for running the tests.

The Water Quality Association recommends these basic tests for allprivate water supplies:

1. Coliform bacteria2. Lead and copper from water that has been standing in the

household plumbing system for more than a few hours, and3. Nitrates

At a minimum, private wells, for example, should be tested for coliformbacteria once each year and for nitrates every two to three years.

If nitrate is detected, the drinking water should be tested for nitrateevery year. Anytime that a private well owner notices a change in thequality of the water or after work is done on the well, he or she shouldhave the well tested again for coliform bacteria and nitrate. Before col-lecting a water sample for the testing, the homeowner should contacta certified water testing laboratory or a qualified local water treatmentcompany and water treatment specialist for assistance.

To determine whether to consider testing for other contaminants,consumers are advised to evaluate the circumstances of the water sup-ply. If the water has a troublesome taste, odor, appearance, corrosion,staining, or scaling characteristic, one or more of the following shouldbe checked.

1. Water Hardness2. Iron3. Manganese4. Hydrogen sulfide or sulfides5. Total dissolved solids6. pH7. Alkalinity8. Dissolved oxygen9. Turbidity

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10. Chloride, fluoride, and sulfate11. Chlorine and chloramines12. Voltage and amperage between ground and areas of the

plumbing system

Further assessment of the water supply is necessary to direct any addi-tional testing. This assessment should consider the type of troublesomecharacteristic noted in the water supply, concerns from previous waterquality reports, the proximity and vulnerability of the well or water sup-ply to sources of contamination, any unique or changing odors, tastes,or appearance to the water, and any particular susceptibilities of individ-uals that are exposed to and/or drinking the water. To determine whetherand what testing is appropriate, consumers should discuss the water as-sessment with their state or local health/environmental protectionagency, a certified water quality testing laboratory, and/or a qualified localwater treatment specialist.

Additional water analyses that may be appropriate include thosefor one or more of the following analyte categories. Private well own-ers, for example, who have no information on the well’s water qualityor who have a new well of unknown characteristics, may be advised toobtain an initial analytical scan for priority items in these health-related contaminant groups.

1. Heavy metal inorganic chemicals such as aluminum, arsenic,barium, cadmium, chromium, copper, lead, mercury, nickel,selenium, silver, sodium, and zinc

2. Lead and copper from standing water that represents variousparts of the household and water service plumbing systems

3. Volatile organic chemicals including benzene, vinyl chlo-ride, carbon tetrachloride, trichlorethylene, tetrachloroeth-ylene, total trihalomethanes (TTHMs), methyl tertiary-butylether (MTBE), and others

4. Pesticides, herbicides, PCBs, and especially the common her-bicide atrazine or the triazines

5. Radon6. Radiochemicals such as radium 226, radium 228, and

uranium7. Protozoan cysts such as Giardia and Cryptosporidium.

Regardless of where or how an analysis is performed, there are cer-tain precautions which should be observed in preparing a water samplefor shipment to a laboratory. Careless handling of a sample can pro-duce highly misleading results.

To achieve valid results, the following steps should be taken:

1. Secure a container (16 oz. size preferable) of a type that willnot permit contamination of the sample. Clean rubber-

192

stoppered, resistant glass bottles, or polyethylene containersare recommended. A plastic screw-capped bottle may beused. It is advisable to use a new bottle for each test. If this isnot possible, the bottle and cap should be washed in soapand water and thoroughly rinsed before use.

2. Before taking samples through metal lines and valves, firstallow sufficient flow of water to wash out the system, exceptfor lead and copper analyzes which may be originating fromplumbing and soldering materials.

3. Rinse the container thoroughly. Use the water that is to beanalyzed for this purpose.

4. Fill the bottle to a point just below the shoulder. Leave justa little air space.

5. Immediately after filling the container, jot down all perti-nent information regarding the water at the time the sampleis taken. Facts to include: source of the sample; physical ap-pearance of the water (clear, dirty, highly colored, etc.); odor,if any; taste as reported by the homeowner, if any.

6. Send the sample to the lab for testing.

Many certified laboratories will furnish you the clean sample bot-tles and sampling instructions for you to use and follow.

Is a “complete” water analysis necessary in every case? Definitelynot! In many cases, a “complete” analysis would be unwarranted.Where an unusual situation exists, it is perhaps wise to make a “com-plete” analysis of the water. Such tests, when completed, would giveinformation on such chemical contaminants as those listed in the EPAPrimary Drinking Water Regulations, and others such as alkalinity,sodium, sulfates, chlorides, tastes, odors, turbidity, and color in addi-tion to hardness, iron, pH, etc.

After a water analysis has been made, there is still need for correc-tive action.

Sound treatment is a matter of judgement based on know-how, ex-perience, and the particular requirements of the individual owner ineach case. Often a water analysis will reveal conditions which mightbe treated in one of several possible ways. In many cases, there maynot be any one best possible answer. Several solutions may be satisfac-tory depending on the factors involved.

For example, let’s now consider several water analyses. Each ofthem is a comparatively simple analysis. In two of the cases, at leastthree solutions are practical. Which is the most feasible? Which is themost economical?

Examine each of these analyses carefully now. Then consider thesuggested solutions. Note the reasons given for these solutions in eachcase. Perhaps other sound solutions come to mind.

193

Water Analysis No. 1

Date Collected . . . . . . . . . . . . . . . . . . . 8/14/83Source . . . . . . . . . . . . . . . . . . . . . . . . . . . . . WellDate Analyzed . . . . . . . . . . . . . . . . . . . 8/23/83Appearance when drawn . . . . . Clear, odorlesspH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7.5Iron. . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.2 ppmBicarbonate . . . . . . . . . . . . . . . . . . . . . 11.3 gpgSulfate . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5 gpgChloride . . . . . . . . . . . . . . . . . . . . . . . . . 1.6 gpgTotal Anions . . . . . . . . . . . . . . . . . . . . 16.4 gpgCalcium Hardness . . . . . . . . . . . . . . . . 10.5 gpgMagnesium Hardness . . . . . . . . . . . . . . 5.3 gpgTotal Hardness. . . . . . . . . . . . . . . . . . . 15.8 gpgCalcium Bicarbonate. . . . . . . . . . . . . . 10.5 gpgMagnesium Bicarbonate . . . . . . . . . . . . 0.8 gpgMagnesium Sulfate . . . . . . . . . . . . . . . . 3.5 gpgMagnesium Chloride. . . . . . . . . . . . . . . 1.0 gpgSodium Chloride . . . . . . . . . . . . . . . . . . 0.6 gpg

*All values are reported as CaCO3 equivalent except pH and iron.

Let’s see now what we can determine from the information re-vealed in this particular analysis.

Study of Water Analysis No. 1 shows that the water is very hard ac-cording to the Water Quality Association’s hard water classifications.While total hardness equals 15.8 gpg, the total cations equal 16.4 gpg.The sodium in the water (0.6 gpg) accounts for the difference betweentotal anions and the total hardness. The amount of iron, however, isnegligible and would produce no staining of washables in the laundry.

A sound solution for treating this water would be the installation ofan ion exchange softener of proper capacity on both the hot and coldwater lines. The outdoor sillcocks should be bypassed. Where possible,the toilets could also be bypassed as the amount of iron in the wateris not sufficient to produce any appreciable staining.

The capacity of a unit necessary for the installation would dependon the number of people in the household and to some extent on thenumber of water-using appliances. The type of model (time clock ordemand initiated regeneration, DIR) would be a matter of personalpreference on the part of the buyer.

A look at Water Analysis No. 1 shows that the iron is reported inparts per million while all the hardness minerals are listed in terms ofgrains per gallon. Why is this practice followed? Primarily for

194

convenience in reporting concentrations of minerals—some of whichare found in water in abundance and some of which are found in onlyscant, or even trace, amounts.

Actually, there are four basic units of measure used in water analy-sis work: parts per million (ppm) or milligrams per liter (mg/L); grainsper U.S. gallon (gpg); equivalents per million (epm); and grains per im-perial gallon (gpg imp).

In Lesson 4, we discussed how to make a conversion from parts permillion to grains per gallon. As you recall, simply divide the parts permillion by 17.1 to convert to grains per gallon. Thus…

= grains per gallon

At the end of this lesson are basic tables which will enable you toconvert milligrams per liter or parts per million, grains per U.S. gallon,parts per hundred thousand and grains per imperial gallon at will. Oneword of caution: the term parts per million means one part in a mil-lion parts. Correctly, 1 ppm could be translated as one ounce in a mil-lion ounces of water, or one pound in a million pounds of water. Onthe other hand, it would be incorrect to interpret 1 ppm to mean onepound in a million gallons of water without first converting the fig-ures. For obviously, pounds and gallons are not similar units of meas-ure. In the same way, 1 milligram in a liter of water, by definition oneliter of water weighs one million milligrams, is 1 milligram per liter or1 part per million.

The point-of-use water conditioning industry generally expresseshardness in terms of grains per gallon. This is done to avoid workingwith large figures in many instances. An exception would be trace sub-stances which are expressed in terms of milligrams per liter. Anotherpoint to remember…the minerals we are considering in these analysesare expressed in terms of hypothetical combinations in terms of grainsper gallon or milligrams per liter as calcium carbonate (CaCO3).

As you recall in Lesson 4, we stated in effect:

In order to make calculations for all the various hardness min-eral compounds, the concentrations of the various ions mustbe expressed in equivalent units to permit direct addition andsubtraction in analysis work. This is similar to the conversionof 1⁄3 and 1⁄4 to 4⁄12 and 3⁄12, respectively, to facilitate use of thesefractions in addition and subtraction.

If it is stated that a water contains minerals in the amount of 10grains per gallon as CaCO3, these minerals may consist of calcium or magnesium carbonates, bicarbonates, sulfates or chlorides, or a

milligrams per liter or parts per million�����17.1

195

combination of these compounds. But in every case, the combinedconcentration of these various hypothetical combinations is chemi-cally equivalent to 10 grains per gallon of calcium carbonate.

Undoubtedly, calcium carbonate serves as the standard because ithas a molecular weight of approximately 100 (100.089) and an equiv-alent weight of 50 (50.45). The concentration of the various mineralcompounds as calcium carbonate (CaCO3) in a water supply can read-ily be determined upon analysis of the water. The concentrations ofeach of the minerals is divided by the equivalent weight of the com-pound and then multiplied by the equivalent weight of CaCO3.

Thus, to determine the equivalent weight of any mineral com-pound in terms of calcium carbonate, follow this formula to give youthe concentration of that mineral as CaCO3:

concentration of mineral ×

If necessary, review portions of Lesson 4 for full details on hypo-thetical combinations. Now let us consider another water analysis.

Water Analysis No. 2

Date Collected . . . . . . . . . . . . . . . . . . . . . . . 7/12/83Source . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . WellDate Analyzed . . . . . . . . . . . . . . . . . . . . . . . 7/18/83Appearance when drawn . . . . . . . . . . Clear, yellow

color, no odor, “iron” tastepH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8.0Iron . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.6 ppmBicarbonate . . . . . . . . . . . . . . . . . . . . . . . . . . 3.9 gpgSulfate . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.5 gpgChloride . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1 gpgTotal Anions . . . . . . . . . . . . . . . . . . . . . . . . . 6.5 gpgCalcium Hardness . . . . . . . . . . . . . . . . . . . . . 3.4 gpgMagnesium Hardness . . . . . . . . . . . . . . . . . . 1.4 gpgTotal Hardness . . . . . . . . . . . . . . . . . . . . . . . . 4.8 gpgCalcium Bicarbonate. . . . . . . . . . . . . . . . . . . 3.4 gpgMagnesium Bicarbonate . . . . . . . . . . . . . . . . 0.5 gpgMagnesium Sulfate . . . . . . . . . . . . . . . . . . . . 0.9 gpgMagnesium Chloride . . . . . . . . . . . . . . . . . . . —Sodium Chloride . . . . . . . . . . . . . . . . . . . . . . 1.1 gpgSodium Sulfate . . . . . . . . . . . . . . . . . . . . . . . 0.6 gpgColor . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 50

*All values are reported as CaCO3 equivalent except pH and iron.

equivalent wt of CaCO3���equivalent wt of mineral

196

What does this analysis mean in terms of mineral content andsound corrective treatment? From this analysis, the water in this sam-ple should be classified as moderately hard.

This water contains 50 units of color. Corrective treatment is ad-visable when water contains more than the USEPA Secondary Drink-ing Water limit of 15 units of color.

The iron content exceeds the USEPA secondary limit of 0.3 ppm;it, therefore, presents the potential for a significant staining problem.

The sodium content of this water is 1.7 gpg. You get this figure bysubtracting total hardness from total anions.

This water with a pH of 8 is definitely alkaline. The alkalinity canbe traced to its bicarbonate content. For effective precipitation of ironand manganese, as discussed in Lesson 5, a minimum of 100 mil-ligrams per liter of alkalinity and a pH greater than 7.0 is necessary.

Total anions, of course, equals 6.5 gpg.There are several solutions possible for the treatment of this water.Solution No. 1. You could install a softener unit to remove both

hardness and iron. Some manufacturers recommend installation of sof-teners on such concentrations of iron; others do not. Removal of irondepends to a degree on the type of iron in the water that is beingtreated. Application of resin cleaners with the salt or in the brine tankcan automatically clean iron from the resin bed at each regenerationto keep the bed clean and prevent fouling so that both iron and hard-ness may be effectively removed.

If Solution No. 1 is selected, it is advisable to regenerate the unitmore frequently than would be the case if hardness alone were the prob-lem. Taking this precaution will minimize the danger of the iron foul-ing the exchanger bed. Install the unit on both the hot and cold waterlines and connect to the toilets as well to prevent iron staining. Thisshould prove to be a most satisfactory solution where families are small.

Solution No. 2. Install an iron filter on the hot and cold waterlines. Connect it to the piping to the toilets. Install a softener just on thewater (both hot and cold) going to the taps. Bypass the toilets. This is asomewhat more expensive method, but is a better way to handle theiron problem where there is a larger number of individuals in a family.

Solution No. 3. Install a chemical feed pump. Use an oxidizingagent such as chlorine in the form of household bleach, or perman-ganate, a mixing tank and a sand or iron removal filter for removal ofthe coagulated material. Both hot and cold water should be treated to-gether with the toilets. Also, connect a softener to the hot and coldwater lines, bypassing the toilet, if possible, for removal of hardness.

If the exterior of the home into which the equipment is to beplaced is white or some light color, it would be wise to filter the waterflowing to the sillcocks as well. There is a definite possibility that where

197

the water is not treated for correction of color, and iron, it may stainthe exterior of the house during the sprinkling process.

There is a strong indication of color in Water Analysis No. 2. If thecolor is due to some organic contamination, Solution No. 3 would bebest, though more costly.

Now consider one other water analysis which will present severaladditional problems.

Water Analysis No. 3

Date Collected . . . . . . . . . . . . . . . . . . . . . . . 8/15/83Source . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . WellDate Analyzed . . . . . . . . . . . . . . . . . . . . . . . 8/24/83Appearance when drawn . . . . . . . . Clear, colorless,

no odor, “alkali” tastepH . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8.2Iron. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 0.2 ppmBicarbonates . . . . . . . . . . . . . . . . . . . . . . . . 52.0 gpgChlorides . . . . . . . . . . . . . . . . . . . . . . . . . . . 15.6 gpgSulfates. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31.4 gpgTotal Anions . . . . . . . . . . . . . . . . . . . . . . . . 99.0 gpgCalcium . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15.0 gpgMagnesium . . . . . . . . . . . . . . . . . . . . . . . . . . 9.1 gpgTotal Hardness . . . . . . . . . . . . . . . . . . . . . . . 24.1 gpgCalcium Bicarbonate. . . . . . . . . . . . . . . . . . 15.0 gpgMagnesium Bicarbonate . . . . . . . . . . . . . . . . 9.1 gpgSodium Bicarbonate . . . . . . . . . . . . . . . . . . 27.9 gpgSodium Chloride . . . . . . . . . . . . . . . . . . . . . 15.6 gpgSodium Sulfate . . . . . . . . . . . . . . . . . . . . . . 31.4 gpg

*All values are reported as CaCO3 equivalent except pH and iron.

What does this analysis mean in terms of minerals? What type ofcorrective treatment would you recommend?

Study of Water Analysis No. 3 would indicate that the water is veryhard with 24.1 gpg of hardness minerals. The water is also definitelyalkaline with a pH of 8.2 due to its high bicarbonate content.

While there is iron in the water, it would not cause staining ofwhite fabrics and porcelain fixtures unless deposits were permitted tobuild up. Under these circumstances, 0.2 ppm of iron could producestaining in time.

The sodium sulfate content of this water might prove to be laxa-tive to some people.

198

With approximately 75 grains of sodium salts, it might be difficultto obtain “zero” soft water with an ion exchange softener because ofthe regenerative effect of the sodium in the raw water.

The suggested treatment for this water would be an ion exchangesoftener. Such a unit would provide soft water, though perhaps not“zero” soft water. It would also be able to handle the iron content quitesatisfactorily.

Treatment of this water supply with a reverse osmosis system willsubstantially reduce the total dissolved solids including the alkalinity,sodium, and laxative sulfates to provide an excellent high quality waterfor drinking and cooking purposes.

199

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200

201

Calcium, Ca

Calcium Bicarbonate,Ca(HCO3)2

Calcium Chloride, CaCl2Calcium Hydroxide Ca(OH)2

Calcium Nitrate, Ca(NO3)2

Calcium Sulfate, CaSO4

Magnesium, Mg

Magnesium Bicarbonate,Mg(HCO3)2

Magnesium Carbonate, MgCO3

Magnesium Chloride, MgCl2Magnesium Hydroxide,

Mg(OH)2

Magnesium Nitrate, Mg(NO3)2

Magnesium Sulfate, MgSO4

Alkalinity expressed asbicarbonate ion, HCO3

Sodium, Na

Sodium Bicarbonate, NaHCO3

Sodium Carbonate, Na2CO3

Sodium Chloride, NaCl

Sodium Hydroxide, NaOH

Sodium Nitrate, NaNO3

Sodium Sulfate, Na2SO4

Sulfuric Acid, H2SO4

Sulfate Ion, SO4

Chlorine, Cl

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

2.50

.617

.902

1.35

.610

.735

4.12

.684

1.19

1.05

1.72

.674

.831

.820

2.17

.596

.944

.856

1.25

.588

.705

1.02

1.04

1.41

CONVERSION TABLE

Conversion of Compounds to Calcium Carbonate Equivalents(Conversion to like units may also be necessary)

To convert To Multiply by

202

Calcium, Ca

Calcium Bicarbonate,Ca(HCO3)2

Calcium Chloride, CaCl2Calcium Hydroxide Ca(OH)2

Calcium Nitrate, Ca(NO3)2

Calcium Sulfate, CaSO4

Magnesium, Mg

Magnesium Bicarbonate,Mg(HCO3)2

Magnesium Carbonate, MgCO3

Magnesium Chloride, MgCl2Magnesium Hydroxide,

Mg(OH)2

Magnesium Nitrate, Mg(NO3)2

Magnesium Oxide, MgO

Magnesium Sulfate, MgSO4

Alkalinity expressed asbicarbonate ion, HCO3

Sodium, Na

Sodium Bicarbonate, NaHCO3

Sodium Carbonate, Na2CO3

Sodium Chloride, NaCl

Sodium Hydroxide, NaOH

Sodium Nitrate, NaNO3

Sodium Oxide, Na2O

Sodium Sulfate, Na2SO4

Sulfuric Acid, H2SO4

Sulfate Ion, SO4

Sulfur Trioxide, SO3

Hydrochloric Acid, HCl

Chlorine, Cl

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

Calcium Carbonate, CaCO3

.400

1.62

1.11

.740

1.64

1.36

.243

1.46

.843

.952

.583

1.48

.403

1.20

1.22

.460

1.68

1.06

1.17

.800

1.70

.620

1.42

.980

.960

.800

.729

.709

CONVERSION TABLE

Conversion from Calcium Carbonate Equivalent to Compound Concentration

(Conversion to like units may also be necessary)

To convert To Multiply by

203

Lesson 10Introduction

It is obvious now that water requires careful analysis in order to deter-mine what type of water treatment, if any, is necessary to provide the bestwater quality for the intended uses.

Once the problem is analyzed, the next step is to provide the propercorrective equipment for the water. Here, as in all phases of a water treat-ment program, maximum care is essential. For this reason, good instal-lation practices provide the basis for this lesson in the Water TreatmentFundamentals course.

Recommended Installation Procedures

There are three highly important factors to keep in mind when makingan installation of water conditioning equipment:

1. The individual responsible for installing equipment (soften-ers, chemical feeders, filters) should be fully trained for thejob. Further, he/she should know and appreciate the needfor extreme care in making the installation.

2. In making any water conditioning installation, he/sheshould comply with all state and local codes where perti-nent. This applies to pertinent plumbing, electrical, build-ing, and/or specialist codes.

3. The installer should be mindful of the manufacturer’s in-structions. Where these may be in conflict with existingbuilding codes, obviously, the codes take precedence.

Note: The Water Quality Association recommends: “watersofteners shall be installed only by or under the supervisionof reputable and responsible persons fully trained andexperienced with water softening equipment.”

204

Installation for a Taste and Odor Filter

All installations shall be in compliance with applicableplumbing, electrical, building, and qualified specialist codes,as well as with the manufacturer’s instructions and specifi-cations. Where in conflict, the codes take precedence.

In making any installation, one of the first decisions is whether ornot to treat all the water.

In general, it is a must to soften the water for both the hot and coldwater lines. Such an installation is illustrated below.

Both hot and cold soft water are needed for all personal groomingneeds (baths, showers, and shampoos), laundering, dishwashing, andgeneral homemaking tasks. When only the hot water supply is soft-ened, the end result will not provide soft water benefits whenever coldwater is used and hot and cold water are mixed.

Note in the installation below that the softener is placed betweenthe water meter (or pressure tank) and the water heater. Further notethat 3⁄4" pipe or larger is recommended for all connections to the sof-tener. If larger size piping is used in the plumbing, this should bematched. In this way, you avoid the possibility of the softener instal-lation causing an excessive pressure drop. In any event, local codesmust be observed.

205

Installation for Soft Water Throughout the Home

There are many materials that can be used for the piping: brass,copper, galvanized steel, iron, and plastic. Most commonly used arecopper and plastic or pvc piping. Rigid copper tubing is one of the bet-ter materials for installations. On the other hand, the use of soft coppertubing is not recommended.

Flexible plastic pipe is readily available and convenient to use. It isgaining in popularity where permitted under local plumbing codes. Aword of caution: where plastic piping is permitted, you should insiston material which has been validated by an ANSI-accredited certificationorganization. This organization has carried on extensive plastic pipetests, both toxicological and physical. Failure to use NSF tested materialscould lead to unsatisfactory results as far as the consumer is concerned.

Where plastic pipe is employed, it is essential that a ground strapbe connected to metal piping at both ends of the plastic pipe to avoidbreaking the continuity of existing electrical grounding.

In the diagram on page 205, cold soft water flows to all the toiletflush boxes. Is this essential? Soft water for the toilet flush boxes is de-sirable, especially where there is some possibility of iron staining. Also,it should be considered where scaling of the flush mechanism may bea problem.

Under other circumstances, some authorities feel that it is unnec-essary to use soft water for flushing the toilets. They point to the fact

206

Installation for Soft Water Throughout the Home Exceptfor Toilet Flush Boxes

(Note: Same arrangement would be used with other types of filters)

that where there is a leaky toilet a great deal of soft water could bewasted.

Even where soft water for the toilets may not be necessary, theymay have to be included due to the construction of the home. For ex-ample, when a home has been built with a slab-floor construction, allthe pipes are imbedded in the concrete. Under these conditions, it ispractically impossible to bypass the toilets. Again, for second floorbathrooms, it might be necessary to run an extra cold water line upthrough the walls in order to bypass the toilets. Here, the expense ofthe additional piping could well outweigh any savings possible fromnot softening water for the toilets. As you can see, a good deal of com-mon sense must be used in order to provide for a sound, yet practicalinstallation.

In making a soft water installation, there are several other consid-erations to keep in thought:

1. Bypass the outside sillcocks. It would be a waste to sprinklethe lawns with soft water. Also, it could be injurious to thelawn if the sodium content of the water is extremely high.

2. If the water softener is being relied on to remove iron fromthe iron supply, outside spiraling water should be softenedto prevent red and brown iron stains on driveways and sidesof homes. Some homeowners also may desire to have softwater for washing their cars.

3. Again, there are some people who may prefer hard water fordrinking. A hard water bypass of the cold water line to thekitchen sink is sometimes provided for this purpose.

Figure 1 is a selection guide for manual and semi-automatic house-hold water softeners. Guides for automatic and fully automatic unitscan be found in Figures 2 and 3. These charts can be used for ready ref-erence in sizing units properly according to the water usage, numberof people in the household, and grains per gallon of hardness.

Placement of Equipment

After the basic plumbing requirements are taken into account, nextconsider the location of the equipment. Ideally, water conditioningequipment should be placed so that it is readily accessible and yet outof the way. The best location is adjacent to the water meter or besidethe well pressure tank. Space requirements may be such that it will benecessary to place a softener in one area and the brine tank (in the caseof automatic equipment) somewhere else. With the smart looking en-closures available today, softeners can readily be installed in the

207

208

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kitchen or utility room. Under such conditions, it may well be advis-able to place the brine tank elsewhere. However, the manufacturer’slimitations on the length of the brine line should be checked. Too longa line may prevent proper regeneration. In this case, a salt-in-head sof-tener should be considered.

Essentially, a good rule of thumb to follow is this: place equipmentso as to supply soft water to all desired points of use in the simplest,neatest, least expensive, and most effective way. In making a placement,the location of the water heater, main supply line, sillcocks, drain andelectrical outlets must be taken into account.

Installation of Equipment

Once you have full agreement with the homeowner on all require-ments for installing the equipment, you are ready to outline the job tothe installation person. After all necessary bypass lines are installed,place the softener in position. Be sure that it is level (see diagram, “In-stallation for Soft Water Throughout the Home”). Take measurementsfor the connecting pipes that will run from the main line to the sof-tener inlet and outlet. Shut-off valves should be installed on both theinlet and outlet lines. Also, a bypass valve should be positioned on themain line in between the inlet and outlet connections. This permitsthe cutting out of the softener at any time should servicing of the unitbecome necessary. It is important to install an electrical continuityground strap connecting both ends of metal household plumbing oneach side of a water treatment equipment installation. This is especiallypertinent if the equipment consists of plastic valves or connections.Many household water lines can carry elecrical grounding currents; itis important for safety and also corrosion considerations to not disruptsuch electrical currents.

An important aspect of the installation is the setting up of properdrainage facilities. WQA recommends:

Regeneration wastes from permanently installed softeners shallbe discharged to the building waste system, subject to the fol-lowing precautions:a. The softener drain line shall not be connected directly to

the waste system, but shall be emptied into a laundry tray,floor drain, or properly trapped special outlet, preserving anair gap of at least two times the diameter of the drain line,but in no case less than one inch above the top of the re-ceptacle used.

b. Installations requiring rinsing of brine through water sup-ply lines shall not be acceptable.

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The purpose of the air break is to avoid the possibility of havingsewage water drawn from the drain onto the drinking water treatmentequipment. Unless a sufficient air break is maintained, the followingis possible. Suppose during the recharging of a softener a fire breaksout in the neighborhood. Under these circumstances the fire fightingequipment could create such a strong suction that water would flowbackward. Now if this suction were strong enough, some fluid fromthe sewer pipes might be drawn back through the drain line into thesoftener, through the house, and even out to the water lines. The dan-ger, of course, is that the water may be polluted with the possibilityof spreading a disease-producing condition into the drinking water.Obviously, public health standards make it essential to observe the air-gap rule in providing for drainage.

A floor drain, if away from the flow of traffic, is ideal for the dis-charge of the backwash and regenerating effluents. With such an in-stallation, rigid pipe should extend from the softener but stop aminimum of one inch above the drain. A drain or laundry sink willalso serve for the discharge of the backwash and regenerating effluents.The homeowner must be cautioned against having the sink stopper inat any time when the softener is in the process of being regeneratedday or night. For if this occurs, there is the distinct possibility of over-flowing the tub. The use of a long section of garden hose rather thanrigid piping for the drain pipe is not recommended. Do not restrict thedrain line, or use a long garden hose, as such restriction may preventproper backwash and regeneration. The larger the softener or filter, themore important this becomes.

Where softeners are installed in rural areas, it is acceptable to dis-charge the softener wastes into properly sized septic tank systems. Drywells also frequently can provide excellent drainage for the unit. In ef-fect, such installations use the earth as a filter. Naturally such installa-tions should follow local codes closely. Once proper drainage isprovided, the unit is usually ready for operation.

The installer should load the softener or filter, backwash, disinfectand place the unit in operation, and check the results. In general, it isnot necessary to regenerate a new softener. Especially important at thispoint in the installation of a softener is the matter of disinfecting it.Essentially this process calls for the addition of chlorine to the unit.Why is this essential?

The Water Quality Association Guidelines for Disinfection and Sani-tation of Water Treatment Equipment state the following:

“The materials of construction of the modern water softenerare inert toward bacterial nutrition and will not contaminatea water supply. However, conditions existing during shipping,

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storage, and installation indicate the advisability of disinfect-ing a softener after installation and before the softener is placedinto service. In addition, during normal use, a softener may be-come fouled with organic matter, or in some cases, with bac-teria from an untreated water supply. Every softener should bedisinfected after installation. Some will require periodic disin-fection during their normal life, and in some cases, disinfec-tion with every regeneration may be required.”

There are several ways in which to proceed in order to disinfect sof-teners, largely depending on the type of disinfectant employed. Com-monly used for this purpose are sodium hypochlorite or calciumhypochlorite.

Sodium hypochlorite (household bleach) is quite frequently thechoice for disinfecting a new unit. With polystyrene resins, the rec-ommended dosage is 1.2 fluid ounces per cubic foot of resin. Withnon-resinous exchange materials, 0.8 fluid ounces per cubic foot is sug-gested. For salt-in-head or brine tank units using downflow regenera-tion design: (1) begin by backwashing, draining to the point wherewater is about 1⁄2 inch above the exchange resin; (2) pour in therequired amount of sodium hypochlorite; (3) refill the unit from thebottom by backwashing; and (4) close the unit when full and proceedwith normal downflow regeneration. For upflow regeneration units,the disinfectant may be added to the brine well.

When calcium hypochlorite is used as the disinfectant, the dosageshould be 2 grams (approximately 0.1 oz.) per cubic foot. Calciumhypochlorite, which is 70 percent available chlorine, is available undersuch trade names as H.T.H. and Perchloron. These products, which areavailable in both tablet and granule forms, may be placed directly inthe unit without first dissolving them. With downflow regeneration,the steps are then: (1) backwash the unit, (2) pour in the requiredamount of disinfectant from the top of unit, (3) refill unit from bot-tom by backwashing, and (4) close the unit and proceed with normaldownflow regeneration. For upflow regeneration softeners, the disin-fectant may be added to the brine well.

Some softener manufacturers suggest the addition of dry chlorinecompounds to the salt when filling the unit as a means of providingfor continuous disinfection of the softener. Units are also available thatgenerate amounts of chlorine from chlorides in the brine as it is beingdrawn into the mineeral tank for media regeneration.

In making an installation, it’s sound practice at the outset to drainall the pipes. Shut off the water at the meter or tank. Open laundryfaucets, hot and cold. Proceed to open all other hot and cold waterfaucets in the house. Also trip the toilets. This lets air into the lines and

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permits the pipes to drain. It is advisable to reduce the water level inthe water heater at the same time. Some dealers drain and clean theheater while making their installation. This permits them to providetheir customers soft hot water right from the start. If you follow thispractice, care should be taken to turn off the main burner if it is a gasunit, or pull the switch if it is an electrical unit. Upon completion ofthe job, care should be taken to fill all pipes slowly and completely. Asair is cleared from the lines, and water flows steadily from the taps,they should be closed. Begin with the basement taps and then moveto the first and second floors as water flows to the higher levels.

Pressure Drop Considerations

Up to this point, we have not considered the effect of pressure drop ona water softener installation. This factor should not be overlooked forsatisfactory results. While we do not normally associate friction withliquids, friction actually can be a significant factor in a water system.For every foot of pipe in a system, there is a definite and measureableresistance to flow due to friction. Every change in the direction of flow,every plumbing fitting, every restriction no matter what its cause, pro-duces a loss in the energy developed by the water system pump,whether it be a private or municipal system. The pressure of water isexpressed in several different units. The common units used in our in-dustry are: pounds per square inch (psi); “head” in feet (of water); andsometimes as inches of mercury.

TABLE 1Conversion Factors

Feet of MercuryWater Inches PSI

Feet of water 1.0 0.833 0.433Inches of mercury 1.133 1.0 0.489Pounds per square inch 2.31 2.04 1.0

Example: To convert “Feet of water” to “Inches of mercury,” go tohorizontal line marked “Feet of water” and move across to verticalcolumn “Mercury Inches.” The figure there is 0.833. Multiply this bythe “Total Feet of Water.”

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TABLE 2Pressure Loss in Common Pipe Fittings, Expressed in

Equivalent Feet of Straight Pipe

1⁄2" 3⁄4" 1"

Standard elbow 1.7 2.1 2.6Straight thru tee 1.11 1.4 1.8Tee thru side outlet 3.3 4.2 5.3Globe valve, full open 18.6 23.1 29.4Angle valve, fun open 9.3 11.5 14.7Gate valve, full open 0.35 0.44 0.56Tank inlet or outlet 0.92 1.2 1.5Coupling or union 0.35 0.44 0.56

The concept of “head” of water may be illustrated in diagrams. Ifinstead of using conventional pressure gauges, it were practical to in-stall a series of long tubes (top open to the atmosphere) in a water line,the pressure on the water in the line would force water up the tubesuntil the weight of water in each tube just balanced the water pressure.The height of water in each tube, or “head,” is proportionate to thepressure in the water line at the entrance to the tube.

Under conditions where the water is not flowing through the line,the water in each of the tubes would be at the same level (Figure A).However, as soon as water flows in the line, the height of the water inthe tubes will change, as indicated (Figure B). This change in height ofwater is referred to as “head loss” or “pressure drop,” and increases rap-idly with increases in the flow rate. Note that the diameter of the tubeshas no effect on the height of water in the column.

It is obvious that it would be most inconvenient to install a tall trans-parent tube in a line at every point where a pressure reading might bedesired, particularly when 2.31 feet of tube is equal to 1 psi. Thus, a 20psi pressure would produce a water column 46.20 feet tall. In addition,it is more convenient to use the smaller numbers common with the units

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Figure A Figure BNo Flow Flow

of pounds per square inch. However, a great deal of hydraulic data is stillexpressed in feet of head, but it may be converted readily to other units.

Mercury is frequently used in laboratory manometers to take pres-sure readings for several reasons: (1) It is a primary standard which doesnot get out of adjustment readily. (2) It is approximately 13.5 times asheavy as water, making possible a much shorter tube. Thus, it providesexcellent accuracy within a reasonable height. The factors for the var-ious conversions of pressure units are given in Table 1.

When an installation of water conditioning equipment is beingplanned, it is wise to check the plumbing in the home carefully. De-pending upon the size of the pipe, the number of fittings, and theheight to which the water must be delivered, the proposed softener,for example, may be adequate or may be seriously undersized.

The pressure drop in a water system varies with the length of pipe,and this must always be measured. In addition, the pressure drop ef-fect of various common pipe fittings can be expressed in terms ofequivalent length of pipe. Table 2 gives such values for three sizes ofcommon steel pipe.

Table 3 gives the pressure losses through 100 feet sections of steelpipe. Tables 4 and 5 give the same data for two common types of cop-per tubing, and Table 6 gives the pressure losses through 100 feet sec-tions of plastic pipe.

How are these data used? The drawing below shows a simplifiedplumbing plan of a typical home in which the pressure tank, heater,and softener are in the basement together with the laundry facilities.On the first floor of the home are a single bath and kitchen.

Let us assume the following information: the pressure tank oper-ates on a 20–40 psi setting, all pipe is 3⁄4 inch, the proposed softenerhas a 6 psi pressure drop at 4.0 gpm, and that a flow of 4 gpm, half hotand half cold water, is desired at the kitchen faucet.

From the drawing and Table 2 the following list can be made:

Points A to B we have 4 gpm flow.

1 Tank outlet × 1.2 ft. = 1.2 ft.3 Elbows × 2.1 ft. = 6.3 ft.3 Side outlet tees × 4.2 ft. = 12.6 ft.2 Globe valves × 23.1 ft. = 46.2 ft.2 Unions × 0.44 ft. = 0.9 ftActual feet of pipe 20.0 ft.

Total 87.2 ft.

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From Table 3 we find that 4 gpm thru 3⁄4 inch pipe has a 3.0 psipressure drop per 100 feet of pipe. Therefore,

�31.000pfsti.� × 87.2 ft. = 2.6 psi

The water must also be lifted to a total of six feet from the bottomof the pressure tank to the height of “B.” From Table 1, one foot ofheight is equal to .433 psi. Thus,

6 ft. × 0.433 = 2.6 psi

It was stated that the pressure drop through the softener was 6.0psi at the flow rate of 4 gpm. This must also be added in. Thus,

Plumbing pressure drop 2.6 psiNet rise pressure loss 2.6 psiSoftener pressure drop 6.0 psiTotal drop, points A to B 11.2 psi

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The same procedure is used for the pressure drops from points B toC. However, as the flow is now in two paths, each must be calculatedseparately.

Hot Water—from Tables 1, 2, and 3:

4 Side outlet tees × 4.2 ft. = 16.8 ft.1 Globe valve × 23.1 ft. = 23.1 ft.2 Unions × 0.44 ft. = 0.9 ft.2 Tank entrances × 1.2 ft. = 2.4 ft.2 Elbows × 2.1 ft. = 4.2 ft.25 ft. of pipe 25.0 ft.

Total 72.4 ft.

at 2 gpm, �01.8040

pft

s.i

� × 72.4 ft. = 0.61 psi

Net rise, 5 ft. × 0.433 = 2.2 psiTotal 2.8 psi

Cold Water:

1 Straight thru tee × 1.4 ft. = 1.4 ft.3 Side outlet tees × 4.2 ft. = 12.6 ft.1 Elbow × 2.1 ft. = 2.1 ft.25 ft. of pipe 25.0 ft.

Total 41.1 ft.

at 2 gpm, �01.8040

pft

s.i

� × 41.1 ft. = 0.35 psi

Net rise, 5 ft. × 0.433 = 2.2 psiTotal 2.6 psi

In this case, the hot water system has a higher pressure drop thanthe cold water lines and, therefore, the hot water drop is used in fig-uring the total pressure drop.

Points A to B 11.2 psiPoints B to C—Hot Water 2.8 psi

Total 14.0 psi

As the minimum pressure at the pressure tank is 20 psi, 20.0–14.0psi or as low as 6.0 psi pressure would be available at the kitchen tap,under the conditions outlined in this problem.

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Note: It is an important rule that a minimum of ten pounds ofpressure should be available at any tap in order to provide an adequateflow of water. Although it could be argued that the six pounds wouldbe available only when the pressure tank was at its miminum, it is alsotrue that flow could exist at other taps in this home. Therefore, itwould be advisable, under these circumstances, to increase the size ofthe softener to minimize the pressure drop.

Summary

In place of a written summary, this lesson concludes with a series ofdiagrams which show proper placement of various types of waterconditioning equipment. These self-explanatory diagrams should bestudied.

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TABLE 3FRICTION LOSSES PER 100 FT. OF STEEL PIPE

C = 100

1⁄2 Inch 3⁄4 Inch 1 Inch

Flow .622" ID* .824" ID 1.049" ID

gpm Ft. psi Ft. psi Ft. psi

0.5 .582 .25

1.0 2.10 .91

1.5 4.44 1.92 1.13 .49

2.0 7.57 3.28 1.93 .84 .595 .26

2.5 11.4 4.94 2.91 1.26

3.0 16.0 6.93 4.08 1.77 1.26 .55

3.5 21.3 9.22 5.42 2.35

4.0 27.3 11.8 6.94 3.00 2.14 .93

4.5 33.9 14.7 8.63 3.74

5.0 41.2 17.8 10.5 4.54 3.24 1.40

6.0 57.8 25.0 14.7 6.36 4.54 1.97

7.0 76.8 33.3 19.6 8.48

8.0 98.3 42.6 25.0 10.8 7.73 3.35

9.0 122 52.8 31.1 13.5

10.0 149 64.5 37.8 16.4 11.7 5.06

11.0 45.1 19.5

12.0 53.0 22.9 16.4 7.10

13.0 61.5 26.6

14.0 70.5 30.5 21.8 9.44

16.0 90.2 39.2 27.9 12.1

18.0 112 48.5 34.7 15.0

20.0 136 58.9 42.1 18.2

*Inner Diameter

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TABLE 4FRICTION LOSSES PER 100 FT. OF TYPE L RIGID COPPER TUBING

C = 130

1" 5⁄8" 3⁄4" 1"

Flow .995" ID* .652" ID .745" ID 1.025" ID

gpm Ft. psi Ft. psi Ft. psi Ft. psi

0.5 .681 .29 .256 .11

1.0 2.46 1.06 .925 .40 .416 .18

1.5 5.26 2.28 1.96 .85

2.0 8.89 3.85 3.34 1.45 1.50 .65 .411 .18

2.5 13.4 5.80 5.03 2.18

3.0 18.8 8.14 7.09 3.07 3.18 1.38 .871 .38

3.5 25.1 10.9 9.45 4.09

4.0 32.0 13.8 12.0 5.20 5.40 2.34 1.48 .64

4.5 39.6 17.1 14.9 6.45

5.0 48.2 20.9 18.2 7.88 8.15 3.53 2.23 .97

6.0 67.7 29.3 25.5 11.0 11.5 4.98 3.13 1.36

7.0 89.8 38.9 33.8 14.6 15.2 6.58 4.15 1.80

8.0 116 50.2 43.5 18.8 19.5 8.44 5.35 2.32

9.0 143 61.9 53.9 23.3 24.2 10.5 6.63 2.87

10.0 174 75.3 65.6 28.4 29.4 12.7 8.08 3.50

11.0 78.1 33.8 35.1 15.2

12.0 91.7 39.7 41.2 17.8 11.3 4.89

13.0 107 46.3 47.9 20.7

14.0 54.9 23.8 15.0 6.49

16.0 70.3 30.5 19.2 8.31

18.0 87.5 37.9 23.9 10.3

20.0 29.0 12.6

*Inner Diameter

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TABLE 5FRICTION LOSSES PER 100 FT. OF TYPE K RIGID COPPER TUBING

C = 130

1⁄2" 5⁄8" 3⁄4" 1"

Flow .527" ID* .652" ID .745" ID .995" ID

gpm Ft. psi Ft. psi Ft. psi Ft. psi

0.5 .804 .35 .285 .12

1.0 2.90 1.26 1.03 .45 .535 .23

1.5 6.20 2.68 2.18 .94

2.0 10.5 4.54 3.72 1.61 1.93 .84 .475 .21

2.5 15.8 6.84 5.61 2.43

3.0 22.2 9.61 7.88 3.41 4.09 1.77 1.01 .44

3.5 29.5 12.8 10.5 4.54

4.0 37.7 16.3 13.4 5.80 6.95 3.02 1.71 .74

4.5 46.8 20.3 16.6 7.18

5.0 56.9 24.6 20.2 8.74 10.5 4.54 2.58 1.12

6.0 79.7 34.5 28.3 12.2 14.7 6.36 3.62 1.57

7.0 106 45.9 37.6 16.3 19.6 8.48 4.80 2.08

8.0 136 58.9 48.2 20.9 25.2 10.9 6.18 2.68

9.0 169 73.2 60.0 26.0 31.2 13.5 7.65 3.32

10.0 205 88.7 73.1 31.7 37.9 16.4 9.34 4.04

11.0 87.0 37.7 45.1 19.5

12.0 102 44.2 53.1 23.0 13.0 5.63

13.0 119 51.5 61.5 26.6

14.0 70.6 30.6 17.4 7.53

16.0 90.5 39.2 22.2 9.61

18.0 112.3 43.6 27.6 11.9

20.0 33.5 14.5

*Inner Diameter

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TABLE 6PLASTIC — *C = 140

PRESSURE DROP IN PSI PER 100 FEET

Nominal Size 1⁄2" 3⁄4" 1" 11⁄4" 11⁄2"

Actual ID .622" .824" 1.049" 1.380" 1.61"

Flow U.S. gpm.

1.0 0.49

2.0 1.77 0.45 0.14 — —

3.0 3.74 0.96 0.29 — —

4.0 6.41 1.62 0.50 0.13 0.062

5.0 9.61 2.45 0.76 — —

6.0 13.5 3.44 1.06 0.28 0.13

7.0 18.0 4.59 — — —

8.0 22.9 5.84 1.80 0.48 0.23

9.0 28.6 7.27 — — —

10.0 34.8 8.83 2.74 0.72 0.34

12.0 — 12.4 3.83 1.01 0.48

14.0 — 16.5 5.11 1.34 0.63

16.0 — 21.0 6.54 1.71 0.81

18.0 — 26.2 8.10 2.13 1.01

20.0 — 31.8 9.87 2.60 1.23

25.0 — 3.92 1.84

30.0 20.8 5.50 2.60

35.0 27.8 7.32 3.44

40.0 35.5 9.35 4.42

— — 6.67

*C = a constant accounting for surface roughness. The higher thenumerical value of C, the smoother the inside surface of the pipe. Forgood, clean, new plastic pipe, a value of 140 represents normalsmoothness of the inside pipe surface.

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224

Recommended Installation for Water Softener and Iron Filter

Illustration of Phosphate Feeder for Treatment of Entire HouseholdSupply for Iron Stabilization (small amounts) and Corrosion Control

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Illustration of Water Softener and Phosphate Feeder to Prevent Rusty Hot Water or to Control Corrosion in the Hot Water System

Installation for Chemical Feed Pump

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Illustration for Water Filters Removing Iron, Sand, or Sediment:No Softener Required

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Lesson 11Introduction

The installation of water conditioning equipment has but one purpose…toprovide a definite improvement in the quality of water.

Some consumers have certain rather vague ideas about what waterconditioning can do for them. Others are completely uninformed on thesubject. Few know specifically what to expect when they consider the pur-chase, for example, of a softener, a filter, a reverse osmosis unit, or achemical feeder.

This lesson points up some of the common questions about water con-ditioning and provides answers. Of course, these are only a few of themany questions which may be asked.

Because of the nature of this lesson, we have omitted the prelessonquestionnaire.

Common Questions and Answers in Regard toWater Conditioning

QUESTION: Will my family be deprived of minerals nec-essary to good health in drinking water softened by the ionexchange process?

ANSWER: No. The human body gains the minerals necessary togood health primarily through eating foods, not through drinkingwater. The body may absorb or use the minerals in water but, in mostcases, the amount would not be significant. In order for a person toobtain sufficient minerals from water, it would be necessary to drinkmany gallons daily. In general, neither a water with a high mineral con-tent, nor a fully softened water, could be considered a significant sourceof minerals. In contrast, one glass of milk provides the mineral equiv-alent of multiple gallons of ordinary well water. (Cow’s milk containsabout 8,000 milligrams per liter of dissolved minerals.)

Note: Certain trace elements, such as fluoride, iodine, etc.,may be obtained from water. However, these would not beremoved through common household water softening.

QUESTION: Is soft water safe for tropical fish?ANSWER: Yes, soft water is satisfactory for most tropical fish. Ac-

cording to several authorities, both fully soft water and municipallysoftened water would have no undesirable or toxic effect for use in anaquarium.

When making the change from hard to soft water, it is necessaryto make the substitution on a gradual stepwise basis of new water forold. This follows the basic pattern in regard to any change in the en-vironment for tropical fish. This applies to temperature, pH of water,food, as well as to hardness. Drastic changes, of course, would consti-tute a shock to the delicate systems of such fish and could result in fa-talities. Preferably replace about one-fourth of the aquarium water atweekly intervals with soft water. Eventually, the aquarium will have asupply consisting of essentially soft water, and the fish will suffer noill effects as a result of the change.

Note: While soft water is an improvement in that it reducesthe clouding and scaling of the glass panels of an aquarium,it does not, of itself, necessarily provide a suitable environ-ment for the breeding of tropical fish. Authorities indicatethat water of low dissolved solids and pH control may be

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more desirable for breeding, though this depends on thespecies. Since total dissolved solids content of a softenedwater is the same as that of the untreated raw water, a sup-ply with a lower dissolved solids content must be gained insome other way. Blending of softened water with reverseosmosis or distilled water may produce the conditions con-ducive to breeding.

QUESTION: Does soft water affect the operation of a hu-midifier?

ANSWER: Yes. Soft water provides for easier maintenance of ahumidifier. When hard water is evaporated, the mineral residue con-sists of a hard scale which normally requires some drastic treatment(such as chipping or acid) for its removal. When soft water is used, theresidue is commonly called soft and can usually be removed by flush-ing the unit with water or going over the surface with a brush.

A point to remember: Softening water does not reduce the totalamount of minerals present; ion exchange softening merely convertsthe calcium and magnesium minerals to sodium minerals. The humid-ifier most common in homes has an open pan, a small tube connectedto a water source, and float valve. When water evaporates, the float valveopens to permit make-up water to flow into the pan. Sooner or later thistype of unit fills with minerals deposited by the inflowing water.

Many humidifiers today automatically accomplish periodic flush-ing with fresh soft water to keep the mineral concentration down, andthe unit operating satisfactorily. Soft water minerals will flush or rinseaway much easier than will hard water mineral deposits.

Note: A modification of the pan-type humidifier uses wicksto increase the surface of water exposed to the air and thusincrease the evaporation rate. The wicks in such humidifiersare particularly susceptible to clogging due to scale. Whenthis occurs, the wicks must be replaced. When soft water isused, however, the mineral deposit can be redissolved bysoaking the wicks in fresh soft water.

There is also another type of humidifier that physicallysprays a fine mist of water into the air. If minerals are pres-ent in the water, they settle out of the air as a fine powder.Depending on the mineral concentration, the amount ofwater evaporated, and the use of the humidified air, a widevariety of problems may be encountered.

In homes, a few grains of minerals per gallon of watermay be tolerated, but higher concentrations may lead to

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large quantities of fine dust throughout the home. Again,the severity of the problem depends upon the amount ofwater evaporated. In industries, where the fine dust may actas an abrasive in machinery and equipment, the problemmay be much more severe. Thus, the best water for such hu-midifiers is water which is free of all dissolved minerals, suchas demineralized water.

QUESTION: Can an ion exchange softener effect the re-moval of radioactive contaminants from water?

ANSWER: Yes, a water softener will effectively remove from 80 to97 percent of the cationic radioactive substances such as radium foundin water. Water softening will not, however, remove radioactive radongas or anionic radioactive species (e.g., uranium).

Radioactive radium, barium, and strontium, for example, are muchpreferred by ion exchange water softening resins over the water hard-ness cations calcium and magnesium. Therefore, one can always beconfident that so long as a water softener is softening water, it is evenmore effectively removing these cationic radioactive isotopes.

Note: The cationic radioactive contaminants likely to bepresent in a water supply include:

ATOMIC WEIGHT

Barium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .137Cerium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .144Cesium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .137Radium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .226Radium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .228Strontium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 89Strontium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90Yttrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91Yttrium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 90

Also, the following anionic radioactive contaminants, espe-cially iodine, are likely to be present in a contaminated water:

Iodine . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 131Ruthenium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 103Ruthenium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106Uranium . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 238

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The most common radioactive constituent naturally pres-ent, especially in groundwater, is the radioactive gas:

Radon . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .222

The discharge of water softener regeneration wastes results in thesame course for the radioactive constituents in the sewer system as ifthe softener was not used to protect the drinking water. They usuallyend up in the sewage sludge and at no higher concentration thanwould occur without ion exchange softening.

Reverse osmosis and distillation will also substantially remove bothradioactive cations and anions with continuous concentrated wastedisposal.

Because some radioactive substances are anionic, while others arecationic, demineralization units employing strongly acidic action resinand a strong base anion resin can be utilized to remove both speciestypes.

Small “throw-away” demineralizers are available for emergencyneeds. Such units offer a very satisfactory solution to this problem, butshould be used following reverse osmosis.

Activated carbon and/or deaeration are effective means to reduceradon gas from a water supply.

QUESTION: Should softened water be used for wateringhouse plants or for sprinkling the garden or lawn?

ANSWER: Where the amount of hardness minerals in the wateris only moderate (less than 10 gpg), it is doubtful whether the sodiumconcentration would be sufficient to be a serious hazard to plants. Mosthouse plants require specific soil conditions for healthy growth. Manythrive best in slightly acid soils. If there is a high hardness concentra-tion in the water being softened, the necessarily higher sodium con-centration of the soft water may be harmful to plants.

For outside sprinkling purposes, the use of softened water, for econ-omy reasons, is not recommended unless necessary to prevent ironstains on buildings and concrete. Again, where the concentration ofhardness minerals is heavy, the sodium salts replacing them might re-tard growth and might be sufficient to kill the grass.

Note: Where rainfall is rare, sodium accumulation is apt tobe greatest. Heavy rain “ rinses “ the earth.

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QUESTION: Will soft water produce clearer ice cubes?ANSWER: Actually removing calcium and magnesium from the

water has little effect on the quality of ice prepared in the home. Hereagain the reason is that softening the water does not reduce the totalmineral concentration. To the extent that a softener removes sedimentiron and manganese, for example, from water, this would help to pro-duce at least cleaner ice. Filters, of course, can be helpful in removingiron, turbidity, tastes and odors from water used for icemaking. De-mineralized water such as from reverse osmosis, distillation, or deion-ization is most ideal for icemaking of all types.

The use of polyphosphates is an economical method of treatingwater used in typical restaurant icemaking units. The polyphosphateskeep the minerals in the water dispersed and, in this way, minimizethe cloudy appearance of ice cubes.

If fed in proper concentrations, polyphosphates also control scaleformations and corrosion in the ice cube machine. Approximately 5ppm are recommended for scale prevention and 10 ppm for both scaleprevention and corrosion control.

Note: Total minerals must be below 10 grains per gallon forfirst quality ice. Large commercial ice producers have foundthat water containing more than 20 gpg of minerals causesdifficulties in the freezing process. Further, water with sucha mineral content may make a brittle ice of poor quality.

Large commercial ice plants use such processes as reverseosmosis, lime softening, alkalinity reduction, filtrationand/or deaeration to produce the high quality of waterneeded for quickly freezing quality ice with a minimum oflabor and expense. Reverse osmosis filter units are availablein sizes small enough to be used in restaurants, homes, andother small installations, but the other processes are toolarge for these applications.

QUESTION: Does a water softener have any harmful ef-fect on a septic tank?

ANSWER: Several studies have been made to determine the exactnature of water softener recharge waste effluents and their effects onprivate sewage disposal systems. These studies evaluated three majorareas, all dealing with the effect of effluents developed during therecharge of household water softeners.

First, it was important to study the effect of dissolved salts in softenerrecharge effluents on biological action in septic tank systems. These stud-ies demonstrated that recharge effluent from water softeners had no dele-

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terious effect on the biological action in a septic tank and that therecharge waste effluents may actually stimulate biological action.

Second, it was felt important to assess the hydraulic effect of thevolume of water softener waste water. These studies demonstrated thatthe volume of recharge effluent from a water softener is less than thatof present day automatic clothes washers. The amount of waste effluentdeveloped by a typical household water softener during recharge isabout 50 gallons containing calcium, magnesium, and sodium chlo-rides. The frequency of recharge is dependent on water hardness, waterusage, and recharge salt dosage.

The last area of study concerned the effect of softener recharge ef-fluents on soil percolation in septic system drain fields. This portionof the study is important since much of the literature on irrigation con-tains references to the adverse effects of high sodium water on soilstructure and permeability, particularly in clay-type soils. The studyconcluded that there was an important difference between water sof-tener effluents and sodium effluents, which has an important bearingon soil percolation and permeability.

The important difference is that water softener effluents containsignificant amounts of calcium and magnesium and thus are not re-ally sodium effluents alone. Calcium and magnesium counteract theeffect of sodium and help maintain and sustain soil permeability, evenin susceptible clay-type soils. Thus, it appears that water softenerrecharge effluent brine will not affect biological digestion, hydraulicload, or leach field permeability in a septic tank system. However, ifthe leach field is composed of swelling clays, permeability will be re-duced regardless of the presence of water softener effluent. Moreover,calcium and magnesium contained in recharge effluents actually in-creased soil permeability.

Salts in the waste effluent from recharge of water softeners createdno hydraulic conductivity or percolation problems in a properly designedseptic tank seepage field. In fact, it was found that soil percolation wasincreased by water softener recharge effluents, as compared to soil re-ceiving household sewage effluents without the addition of effluentsfrom the recharge of water softeners. In other words, lower hydraulicconductivity (HC) might result if regeneration or recharge wastes fromwater softeners were not allowed to enter the septic tank seepage field.In this case, the beneficial effects of calcium and magnesium would belost. This would occur if the regeneration wastes were not discharged tothe septic system, but to a dry well, roadside ditch, or other point.

One study was conducted by soil scientists at the University of Wis-consin and dealt solely with anaerobic septic tank systems. The otherstudy, conducted by the National Sanitation Foundation, dealt solelywith aerobic septic tank systems. Conclusions reached in this study

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were as follows: Water softener regeneration wastes demonstrated no adverseeffects on home aerobic waste water treatment plant performance, even whenstressed by loading at a use rate simulating ten persons (twice the averageuse rate). There was no difference in performance between days in which theplant received regeneration wastes and days in which it did not.

QUESTION: Is the sodium in softened water harmful topeople on restrictive salt diets?

ANSWER: The amounts of sodium in softened water are minis-cule compared to other normal dietary sources of sodium. In fact, ionexchange softening of water with even 75 grains per gallon of totalwater hardness would add less sodium to the drinking water than is al-lowed in beverages meeting the U.S. Food and Drug Administrationregulations for “Low Sodium” labeling.

In establishing a salt-free diet for patients, physicians should notoverlook the fact that even hard water may contain appreciableamounts of sodium. To determine the amount, a complete analysis ofthe water is necessary.

How can the sodium content of a softened water be determined interms of milligrams of sodium?

1. First, determine the sodium content of the natural water.Multiply the water’s sodium content in grains per gallon ex-pressed as calcium carbonate, by 7.86. This will give you thesodium content of the water in milligrams per liter of water(gpg CaCO3 × 17.118 mg/L/gpg × 22.99 Na+/50.0436 CaCO3= 7.86).

2. Next, determine the additional sodium content of water asthe result of ion exchange softening. Here, multiply the totalhardness of the water in grains per gallon, expressed as cal-cium carbonate, by 7.86.

3. A simple addition of the results of both steps No. 1 and 2will give the sodium content of the softened water in mil-ligrams of sodium per liter. One to two liters (1 liter equals1.057 quarts) is commonly accepted as normal daily waterconsumption.

Actually, the amount of sodium present in softened water is smallwhen compared to the sodium present in foods.

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MilligramsThe Food Amount of SodiumMilk 2 cups 244Eggs 1 59Meat 4 ounces 76Bread 3 slices 342Cereal 1 cup 256Potatoes 1 medium 5Vegetables (peas) 3 oz. 402Vegetables (cauliflower) 3 servings 15Fruit (apple) 1 2Fruit (grapefruit) l⁄2 1

Total milligrams of sodium 1,405

It is important to note that about 2⁄3 of the daily water intake of theindividual is through food and only about 1⁄3 from water itself.

QUESTION: How can iron stains be removed from fabrics?ANSWER: Iron stains can be removed from most washable fabrics

with oxalic acid, provided that certain precautions are taken. Oxalicacid can be obtained at most drug stores.

Note: Oxalic acid is a poison, and everything it touches shouldbe thoroughly washed after use. In addition, it may be a skinirritant, particularly if the hands have small cuts or otherbreaks in the skin. Therefore, rubber gloves should be used.

A solution of the oxalic acid may be made up by simply dissolvingthe crystals in water. As the solubility of the acid is only about 10 per-cent, it cannot be made up too strong. The solution should be preparedin a plastic container.

A small amount of the solution should be tried on the inside of ahem or other inconspicuous location, as the acid may bleach certaindyes. If it is apparent that no bleaching has occurred, the entire gar-ment may be repeatedly dipped in the solution. The dipping shouldcontinue until the iron stains are gone. Allowing the garment to soakin the acid solution is not recommended.

After the stains have disappeared, the garment should be thor-oughly rinsed in several changes of fresh water. To be sure that any re-maining acid has been neutralized, the garment should be immediatelylaundered in the normal manner, with the regular amount of soap. Thealkalinity in the soap will eliminate the last traces of acidity.

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The acid solution may be poured down the drain, but it should befollowed with a thorough flushing with fresh water.

QUESTION: What general procedures should be used withhousehold dishwashers?

ANSWER: In any case, the instructions of the dishwasher manu-facturer should be followed. However, these may be supplemented withthe following general considerations.

Avoid the use of sudsing soaps or synthetic detergents as the sudswhich develop can “muffle” the spray action used in most dishwashersto provide scrubbing action. A number of washing compounds are onthe market which have been specifically formulated for dishwashers.These materials do not generally contain any soap—they depend uponthe presence of greases and oils on the dishes to produce a cleaning so-lution.

Even with the use of soft water, spotting of dishes is sometimes re-ported. In some cases, the cause may be traced to the use of too muchwashing compound. In other cases, the dishwasher is loaded too heav-ily, and improper cleaning occurs.

Another common cause of spotting is too rapid drying. Ideally, thedishes should first be allowed to drain completely in a humid atmos-phere so that the water with its minerals runs off, rather than evapo-rates. In some cases, the water temperature is too high, and the waterevaporates rather than drains. In other cases, the dishwasher lid opensat the end of the cycle and allows cool air to contact the utensils. This,too, results in rapid evaporation.

Depending upon the construction of the dishwasher, it may be pos-sible to adjust one or both of these factors to improve the results sig-nificantly.

QUESTION: Can softened water be used for a steam iron?ANSWER: Neither hard nor soft water should be used with a steam

iron. Distilled water, or water treated by reverse osmosis, is acceptablefor use over a period of time. Bear in mind that the softening of waterdoes not remove the minerals, but that soft water minerals can be moreeasily rinsed from the iron.

QUESTION: How much soap or detergent should be usedin softened water?

ANSWER: When first using softened water for household clean-ing chores, it is best to use as little soap as possible. If necessary, thehomeowner can gradually increase the quantity to produce the resultsdesired.

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The habit of using far too much soap is not easily broken. There-fore, anyone contacting the customer should stress the coffee measureas being adequate, rather than the cup or more that may have beennecessary prior to the installation of a water softener.

QUESTION: Can pink water stains be caused by micro-organisms?

ANSWER: Yes. Homeowners can experience an unusual problemof a pinkish substance on bathroom fixtures that is very persistent andappears in the shower, the sink, and especially along the water line oftoilet bowls.

This pink residue is less likely a problem associated with water qual-ity than with naturally occurring airborne bacteria. The bacteria pro-duces a pinkish film, and sometimes a dark gray film, on surfaces thatare regularly moist, including toilet bowls, showerheads, sink drains,and tiles. The problem also more commonly occurs in humid regionsof the country.

To determine the exact species of bacteria would require lengthyand costly laboratory testing, and for those reasons most homeownersare reluctant to have the tests performed. Although the exact speciesof bacteria is not known, most experts have concluded that this pinkstaining is most likely from the bacteria Serratia marcescens. Membersof the Serratia genus are essentially harmless organisms that produce acharacteristic red pigment. These bacteria thrive on moisture, dust, andphosphates and are widely distributed, having been found naturally insoil, food, and also in animals. The conditions for the survival of Ser-ratia marcescens are minimal, and the bacteria may even feed upon it-self in the absence of other nutrients.

Many times, the pinkish film appears during and after new con-struction or remodeling activities. The dirt and dust stirred up from thework probably contains Serratia bacteria. Once airborne, the bacteriaseek moist environments to proliferate. Some people have even notedthe pink residue in their pet’s water bowl. It causes no apparent harmand can be easily cleaned off. Others have indicated that their experi-ence with this nuisance occurs during a time of year that their win-dows are open for the majority of the day. These airborne bacteria cancome from any number of naturally occurring sources.

The best solution to keeping fixture, sink, and bathroom surfacesfree from this bacterial film is continual cleaning. Chlorine bleach canbe periodically stirred into the toilet tank and flushed into the bowlitself. As the tank refills, more bleach can be added. Three to five ta-blespoons of fresh bleach should be all that is necessary. A toilet cakethat contains a disinfectant can keep a residual in the water at all

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times. The porous walls of a toilet tank can harbor many opportunis-tic organisms.

Cleaning and flushing with chlorine will not necessarily eliminatethe problem, but will help to control these bacteria. Keep bathtubsand sinks wiped down and dry to avoid this problem. Using a cleaningsolution that contains chlorine will help curtail the onset of thebacteria.

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Lesson 12Introduction

The preceding lessons have attempted to cover clearly and briefly most ofthe fundamentals of water treatment. It has not been the purpose of thecourse to examine the subject exhaustively, but to present most of the factsnecessary to an understanding of point-of-use water treatment.

The most important points—points you should have at your finger-tips—are the points that are stressed in this 12th and final lesson ofWater Treatment Fundamentals.

The fact that certain points are omitted, however, does not indicatelack of significance. The purpose of this lesson is to stress the high points.

General Review

Pure Water

The term “pure water” is somewhat lacking in precise meaning. In anabsolute sense, there is no body of water that contains only H2O mol-ecules. To the general public, “pure water” undoubtedly means waterthat is both safe and enjoyable to drink…appealing to the senses oftaste, smell, and sight. To the bacteriologist, pure water signifies a ster-ile liquid which contains no living organisms. And to the chemist,water is pure when it contains no mineral, gaseous, or organic impu-rities. Obviously, water that measures up to the chemist’s or bacteriol-ogist’s definition of purity may well be found only in laboratories, andeven there only under ideal conditions.

Most available water contains contaminants in a wide range ofcombinations and amounts. The term “removal” in this text meanssubstantial reduction or a sufficient degree of removal of specific im-purities or contaminants.

Acceptable Levels of Various Contaminants

In its Primary and Secondary Drinking Water Regulations, the USEPAlists Maximum Contaminant Levels for water impurities. While theEPA regulations list specific limits of contaminants, the more generalrequirements for water may be considered as follows:

1. That it shall contain no disease-producing organisms.2. That it be colorless and free from turbidity.3. That it be reasonably soft.4. That it be good tasting, free from odors and preferably cool.5. That it be free from objectionable gases, such as hydrogen sul-

fide, and objectionable minerals, such as iron and manganese.6. That it be neither scale forming nor corrosive.

Water Impurities or Contaminants

Water is frequently described as “the universal solvent.” Because of itssolvent action, it picks up or dissolves small amounts of metals, rocksand minerals, gases, dust, and numerous other inorganic and organicsubstances with which it comes in contact. In addition, water is an ex-cellent environment for countless organisms, a few of which are harm-ful to man.

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How Water Collects Contaminants

Water follows a continuous cycle from sky to earth and back again.This never-ceasing process is known as the hydrologic cycle. When mil-lions of vapor particles combine in the atmosphere, they become dropsof moisture. As these increase in size, they finally become so heavythey fall to earth as rain, snow, sleet, or hail. In nature’s balanced op-erations, all this precipitation sooner or later is drawn back into the at-mosphere in the process of evaporation.

Actually about 70 percent of all precipitation returns quickly to theatmosphere. In fact, the sun causes some evaporation even as moisturefalls. It also draws water from the top soil, surface run-off, streams,lakes, oceans, and from leaves through the process of transpiration.Only about 30 percent of all precipitation seeps down into the groundor becomes part of oceans, rivers and lakes. As moisture falls to earth,it has a cleansing effect on both the atmosphere near the earth’s sur-face and on the earth itself. The result is that even rainfall contains sig-nificant amounts of matter, such as gases, dust, organic matter, andvarious minerals. Because the cleansing action of water continues afterit reaches the earth and percolates into it, water acquires furtheramounts of hardness minerals in addition to clay, silt, and decayed an-imal and vegetable matter.

Surface and Groundwater Supplies

Groundwater

Under most conditions, groundwater contains greater amounts of dis-solved minerals than do surface supplies. But as water percolatesthrough sand, rock, and clay formations, it loses much of the sus-pended matter, color, and bacterial contamination which it had col-lected on the earth’s surface. Thus, deep wells are likely to providewater that is clear, colorless, low in bacteria, and high in minerals.Springs also derive from groundwater and usually contain significantamounts of dissolved minerals.

Major considerations in the use of groundwater include:

1. The presence of hardness and other minerals in largeramounts, as a rule, than in the surface waters of the same lo-cality.

2. Iron and manganese in many well supplies.3. Hydrogen sulfide sometimes present.4. The cost of pumping well water usually greater than that for

pumping surface water.

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5. The mineral content of several wells may differ widely eventhough located close to each other.

6. There are uncertainties of supply.7. There is a lower possibility of bacterial contamination as

compared to surface waters.8. Generally more constant temperature and composition than

surface waters.

Surface Water

We classify lakes, rivers, reservoirs and ponds, for example, as surfacewater. These bodies receive water directly from precipitation and fromsurface run-off. They also derive a portion of their supplies from un-derground springs connected with groundwater sources.

While surface water, as a rule, has the advantage of lower mineralcontent, it also has certain disadvantages to be considered:

1. The presence of more contaminated matter and sewage-typepollution making water unfit for human consumption untilproperly treated.

2. Industrial and municipal pollution of many supplies.3. Surface run-off bringing mud and decayed vegetation into

the water.4. Possibility of animal and human wastes in the water.5. A potentially good environment for algae and bacteria.

Regardless of the mineral and organic makeup of a water source,both deep wells and large lakes make available water that is of more orless consistent quality from season to season. In contrast, many smallbodies of water, shallow wells, and springs often reflect seasonal andeven daily variations in their mineral content.

Let’s turn now to a consideration of the specific mineral contami-nation of water.

Variation in Water Quality from Different Sources

To understand these variations, it is necessary to know somethingabout the chemistry of water.

When suspended in the atmosphere, water vapor approximates dis-tilled vaporized water. It is free from impurities and remains thus as longas it stays aloft. When water vapor condenses sufficiently to fall to earth,it comes into contact with gases in the surrounding air—chiefly carbondioxide, nitrogen and oxygen, with the possibility of industrial fumesand waste gases. In falling, moisture absorbs amounts of the atmos-

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pheric gases because these are partially soluble in water. Atmosphericdust may also contain minute particles of silica, oxides of iron, miner-als, and other material together with pollen and some microorganisms.

If we chemically diagram the action of water as it dissolves someof the carbon dioxide in the air, it would look this way:

H2O + CO2 → H2CO3

Water Dissolves and Collects Carbon Dioxide to Produce Carbonic Acid

Normally when such water reaches the earth, it is slightly acid, cor-rosive, and relatively soft. Here it may pick up additional amounts ofcarbon dioxide from decaying vegetable matter. Equipped with thisbooster action, it acquires even greater potential for dissolving miner-als and other impurities on or below the surface. If water, now slightlyacid, has the opportunity to seep into the soil and pass through a lime-stone stratum, the acidity will be neutralized In this process, the waterbecomes increasingly mineralized.

Chemically this can be diagrammed:

H2CO3 + CaCO3 → Ca(HCO3)2

Carbonic Acid Reacts With Insoluble Calcium Carbonate to ProduceSoluble Calcium Bicarbonate

Limestone, a common rock formation, contains varying amountsof both calcium and magnesium carbonates. These are the unseenhardness minerals which plague so many water supplies. The basic re-action shown in the above diagram holds true for both minerals. Ironand manganese are found in water supplies less frequently. But againtheir basic chemical reaction in water is quite similar. Because it is asolvent, water also picks up readily soluble chlorides, sulfates and ni-trates of calcium and magnesium. Similarly, it dissolves the carbonate,bicarbonate, chloride, sulfate and nitrate compounds of sodium, as wellas quantities of silica. It is obvious that as these various minerals gointo solution in water, changes occur. Some of these changes are phys-ical in nature, others are chemical.

Physical Change

If you were to smash a rock into any number of pieces, you would pro-duce a physical change. The turning of water into either ice or vaporis a not so obvious example of physical change, for here there appearsto be a definite change in the properties of the substance.

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The point to remember: When a physical change takes place, the struc-ture of the molecules is not changed.

In the case of water, the only difference between the solid, liquid,and vapor forms is this: the molecules in ice essentially have no free-dom. They only vibrate within the crystal. The molecules in the liquidform are free to move within the limits of the container as limited bygravity. The molecules in steam are completely free to move within thecontainer. They are essentially unaffected by gravity.

Chemical Change

Chemical changes occur when new molecules are formed as the resultof a chemical reaction. For example, when sodium (a metal) combineswith chlorine (a gas), common salt results. The molecules of salt arenot like those of either sodium or chlorine. Hence a chemical changehas occurred.

Before going further, consider a few definitions with which weshould all be familiar…

Element

An element is a substance which cannot be chemically decomposedinto an even more basic substance. While only a few elements exist intheir elemental form in nature, chemists have discovered and isolatedmore than 100 of these substances out of which the universe is built.

Atom

The atom is the smallest possible unit of an element. It consists of acore made up of neutrons and protons, the protons each possess a pos-itive charge. Around this core are one or more electrons. These in turneach possess a negative charge. In the atom, the sum of the positivecharges equals the sum of the negative charges. In other words, thenumber of electrons equals the number of protons, and the atom is aneutral particle.

Molecule

A molecule may consist of one or more atoms. In some cases, such aselemental sulfur, the molecule of the element may be a single atom. In

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other cases, two identical atoms of an element form a molecule (Cl2).In the case of compounds, molecules consist of two or more differentatoms because the molecules in a compound are the result of the unionof the atoms of two or more different elements.

As a result of chemical processes, atoms or molecules lose or gainelectrons. When this occurs, they become ions.

Ion

An ion is an electrically charged atom or molecule. It possesses its elec-trical charge due to a loss or gain of one or more of the electrons thatsurround the nucleus or core of the atom.

Several ions may combine into a group having a charge that is thesum of the positive and negative charges of the individual ions in thegroup. Most atoms have a strong tendency to ionize (that is, to becomeions). Why? It appears that they seek to gain or lose electrons in orderto achieve the same type of atomic structure as that possessed by theinert gases. These inert gases, with the exception of helium, all possesseight electrons in their outermost orbit to give them remarkable sta-bility.

Valence

The valence of an ion is the charge due to unequal numbers of elec-trons and protons. The valence is a small whole number such as 1, 2, or3 and may be either positive or negative. This valence expresses thecombining power of the ion with other ions. To form a compound, thesum of the negative valences must equal the sum of the positive va-lences. Thus, two hydrogen ions (each with a single plus charge) com-bine with one oxygen ion (it has a minus 2 charge) to form H2O orwater.

As elements and compounds gain or lose electrons, oxidation-reduction occurs.

Oxidation and Reduction

In the broad sense, oxidation of an atom or ion refers to an increase inpositive valence or a decrease in negative valence. Reduction has ex-actly the opposite meaning—a decrease in positive valence or an in-crease in negative valence. The following chart illustrates what happensin the process:

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Oxidation Valence Reduction

When a substance loses an electron, it becomes oxidized. At thesame time, the substance that receives the electron is reduced. Here areexamples of oxidation and reduction that occur in water:

Oxidation:Fe++ – e → Fe+++

Ferrous Iron Less 1 Electron Reacts to Produce Ferric Iron

Reduction:Cl2 + 2e → 2Cl–

Chlorine Plus Two Electrons Reacts to Produce Two Chloride Ions

As we have seen, in the process of solution, the molecules of manyelements or compounds dissociate into two or more ions which aretheoretically free to move about as independent particles. The word“ion” is derived from the Greek language where it meant to wander.Remember, ions differ from molecules or atoms only in that they arenot neutral particles. Rather, they carry positive or negative charges.In water chemistry, if an ion holds a positive charge, it is called acation. If the charge is negative, the ion is referred to as an anion. Inwater, the total of all positive charges (the cations) equals the sum ofall the negative charges (the anions).

Listed here are a few of the common cations and anions which arefamiliar to the water conditioning industry:

CATIONS ANIONSHydrogen +1 Hydroxide –1Sodium +1 Chloride –1Calcium +2 Bicarbonate –1Magnesium +2 Carbonate –2Aluminum +3 Sulfate –2Ferrous Iron +2 Nitrate –1Ferric Iron +3 Phosphate –3Manganese +2, +4, +7

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(In

crea

se)

(Decrease)

++++++

––––––

+++++

+

–––

–––

Acids, Bases, and Salts

Acids can be defined as compounds that release hydrogen ions (H+) ina solution. In general, they have (1) a more or less sour taste; (2) theychange the color of indicators (i.e., litmus paper to red); and (3) theyreact with bases to form a salt and water.

However, in addition to their properties held in common, acidshave other properties which may vary widely. These specific propertiesof each acid are due to the anion present and to the undissociated mol-ecules. Thus, the molecules of various acids are capable of releasing dif-ferent amounts of free hydrogen ions in solution.

Examples of the release of hydrogen ions (H+) in a solution are:

Hydrochloric Acid:

HCl → H+ + Cl–

Acetic Acid:

HC2H3O2→← H+ + C2H3O2

–

Bases are substances containing (OH–) ions. Sodium hydroxide(NaOH) and ammonium hydroxide (NH4OH) are examples. These hy-droxides ionize as follows:

Sodium Hydroxide:

NaOH → Na+ + OH–

Ammonium Hydroxide:

NH4OH →← NH4+ + OH–

There are a great many bases. In general, their solutions (1) tastebitter rather than sour; (2) feel slippery; and (3) reverse the colorchanges produced by the acids in indicators (i.e., bases turn litmuspaper blue). For example, they turn litmus paper blue.

As in the case of acids, each base has individual properties. Theseare due in each compound to the cation present and the nonionizedmolecules of the base. As in the case of acids, bases demonstrate vary-ing degrees of ionization. Those that ionize only slightly are weak.

Strong acids and bases separate into their ions and stay sep-arated. This is indicated by a single arrow pointing in onedirection. Weak acids and bases are in a continuous processof ionization, but the free ions are also continuously

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recombining to form molecules. At any one time, only a por-tion of the acid or base is present as ions. This equilibriumprocess is indicated by arrows pointing in both directions.

Salts are formed from a metallic ion (positively charged) and a non-metallic ion (negatively charged). Salts are of three types. They are clas-sified as normal, acid, and basic. A salt is a compound produced by theunion of the cations of any base and the anions of any acid. Salts pos-sess no one ion characteristic, although they are nearly all strongly ion-ized. In regard to their solubility, they range between wide limits. Somedissolve less than one milligram per liter. Others dissolve in much lessthan their own weight of water.

No discussion of the chemistry of water conditioning is completewithout review of the term pH.

pH

pH indicates intensity of a given water in terms of alkalinity or acid-ity. Acid or alkaline strength is measured on a scale that goes from 0 to14. Seven, the midpoint, is the measure of a neutral water, neither acidnor alkaline. pH figures larger than 7 indicate alkaline solutions withthe intensity of alkalinity increasing as the number becomes larger. pHfigures lower than 7 indicate acid solutions with the intensity of acid-ity increasing as the numbers get smaller.

There are two important things to bear in mind when consideringpH:

1. Always it is an intensity measure, not one of quantity. Justas a thermometer will tell how cold a room is but not howmuch warm air is necessary to heat it.

2. It is an exponential function. pH 10 is ten times as alkalineas pH 9, and one hundred times as alkaline as pH 8. Simi-larly, a pH 2 is one hundred times as acid as pH 4, and onethousand times as acid as pH 5.

Now let’s look at specific problems that occur in water and how thewater treatment industry can solve them at the point of use.

Water Problems

The most prevalent and common problem is that of hardness.

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Hardness Reduction or Removal

The ions of a number of calcium and magnesium compounds are to befound in water:

Calcium carbonateCalcium bicarbonateCalcium sulfateCalcium chlorideMagnesium carbonateMagnesium bicarbonateMagnesium sulfateMagnesium chloride

Classification of Hardness

Water hardness can be measured as grains per gallon (gpg). Other unitsof hardness measurement are parts per million (ppm), or milligramsper liter (mg/L). Because by definition a liter of water weighs one mil-lion milligrams, a milligram per liter is the same as a part per million.Here is what an analysis in grains per gallon or parts per million meansto you, according to the U.S. Department of the Interior and WaterQuality Association:

milligrams perliter (mg/L)

grains per parts per milliongallon (gpg) (ppm)less than 1.0 less than 17.1 soft1 to 3.5 17.1 to 60 slightly hard3.5 to 7.0 60 to 120 moderately hard7.0 to 10.5 120 to 180 hard10.5 and over 180 and over very hard

Generally speaking, water of up to 60 mg/L or ppm of hardnessdoes not require softening. When water is only slightly or moderatelyhard, many people are not as readily aware of it. Yet, when a switch ismade to soft water, they become aware of an improvement in waterquality.

In areas where water is very hard, water utilities can partially softenthe entire supply. The resulting water typically contains 5 to 10 grainsper gallon, or more, of hardness—still hard to moderately hard. A water

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softening appliance produces water containing less than one grain pergallon of hardness. One of the unique advantages offered by point-of-use water softening is the opportunity for homemakers to have eitherhard or soft water for drinking. This choice is not available if the watersupply is softened municipally.

One rather confusing point is that calcium bicarbonate, magne-sium bicarbonate, and any of the various other hardness compoundsdo not exist as such in water. What happens? In water, the compoundsionize, that is, they become ions with positive or negative charges. Inwater, for example, calcium bicarbonate and magnesium bicarbonateexist as calcium ions, magnesium ions, and bicarbonate ions. Their ex-istence in water as calcium bicarbonate or magnesium bicarbonate ismerely as designations in terms of hypothetical combinations.

Hypothetical Combinations

Hypothetical combinations are a very useful method of describing thevarious dissolved minerals in water. While the minerals exist only asions in water, a laboratory technician making an analysis of the waterwould probably report it in terms of hypothetical combinations ascompounds. This would enable a chemist to prepare a water samplehaving the same chemical characteristics as shown in the analysis.

When hypothetical combinations are calculated, the ions are com-bined in their order of increasing solubility. As calcium compounds aregenerally less soluble than other compounds, calcium is usually firston the list of cations. Magnesium is second, and sodium or potassiumis last.

Similarly, the anions are calculated in the following order: hy-droxides, carbonates, bicarbonates, sulfates, chlorides, and nitrates.

Note: Traces of elements or compounds are not normally consid-ered in these calculations. Iron, for example, would not be included,unless present in extremely high concentration.

Since all the compounds listed in terms of hypothetical combina-tions in water have differing equivalent weights, or, in other words,unique weights per valence charge of chemical activity, they cannotbe totaled together. The concentration of ions of the various mineralsis expressed in terms of calcium carbonate (CaCO3) that is the amountof CaCO3 that would represent equal chemical values to each respec-tive mineral or constituent, to facilitate calculations, just as we convert1⁄3, 1⁄4, and 1⁄6 to 4⁄12, 3⁄12, and 2⁄12 respectively for ease in adding and sub-tracting. A table of equivalent weights for many compounds in solu-tion in water is found in Lesson 4.

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To express the equivalent weight of any mineral compound interms of calcium carbonate:

concentration of the mineral ×

Hardness Removal

Hardness can be removed from water by ion exchange softening,deionization, distillation, reverse osmosis, or lime-soda ash treatment.Let’s review each process briefly.

Ion Exchange Softeners

Softening water for home needs is done almost exclusively through theuse of ion exchange equipment.

Inside the softener unit is a bed of ion exchange material (usuallybead-like resin). The resin softening material is made up of permanentinsoluble anions to which sodium cations are chemically bonded.

Hard water enters the exchange column (the softener tank). As itflows through it, the hardness cations are drawn to the anions of theexchange material. In the process, the hardness minerals are adsorbedand a chemically equivalent number of sodium ions are released intothe water. In brief, harmless sodium ions have replaced the trouble-producing hardness ions.

Ion exchange occurs literally billions of times between the mate-rial in the exchange column and the minerals in the water as soften-ing proceeds.

After a vast number of hardness ions in the water have become af-fixed to the softening material through the attraction of positive andnegative charges, and most of the sodium ions have been released, theunit can no longer soften the water. It has become temporarily ex-hausted, though in actual practice a small number of sodium ions re-main on the softening material after the unit is exhausted.

If no new chemical reaction is set into operation at this point, theincoming hardness ions flow untouched through the unit. Rechargingor regeneration becomes necessary. A reverse ion exchange operation isnow put into motion. In this reverse process, it is necessary to bom-bard the exchange material with the original type of cations in con-centrated solution. In this way, the affinity of the exchanger for thehardness ions is overcome.

equivalent wt of CaCO3���equivalent wt of mineral

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What occurs in all examples of ion exchange is a “swap” or bal-anced exchange of ions. The hardness ions are not destroyed. Theyhave merely been replaced by a chemically equivalent amount ofsodium ions. When the unit is regenerated, these hardness ions arewashed down the drain and sodium ions from the salt brine becomeattached to the resin.

Manufacturers produce two basic types of ion exchange softeners:fully automatic (time clock controlled) and demand initiated (meteror sensor controlled) regeneration models.

Ion Exchange Softener Capacity

The number of gallons of soft water that can be anticipated betweenregenerations may be determined by dividing the rated capacity of thesoftener by the hardness of the water in grains per gallon. Thus:

= Gallons of soft water per regeneration

Deionization

The process used for removal of all minerals from water (primarily forindustrial application) is referred to as deionization. This process treatswater with two ion exchange beds in order to effect the removal of alldissolved salts.

As water flows through the first ion exchange bed, the process re-moves all positive ions and replaces them with hydrogen ions, insteadof with sodium ions as in softening. As the positive ions in the wateraffix themselves to the exchange material, the latter releases its hy-drogen ions on a chemically equivalent basis. Now because of the rel-atively high concentration of hydrogen ions in the partially treatedwater, the solution is extremely acid.

At this point, the deionization process is just half complete. Whilethe positive ions have been removed, the water now contains positivehydrogen ions, and the anions originally in the raw water. The partiallytreated water now flows through a second unit, which contains ananion exchange bed. This second exchange material normally consistsof replaceable hydroxyl anions and fixed insoluble cations. At thispoint, the negative ions in solution (the anions) are adsorbed onto theanion exchange material. Released in their place are hydroxyl anions.All that emerges from such a two unit system is mineral-free water. It

Rated softener capacity���Hardness of the water

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still contains the positive hydrogen ions released in the initial exchangeplus the negative hydroxyl ions released in the second exchange.

What has become of these two ions? Through the magic of chem-istry, they combine (positive to negative) to produce HOH or watermolecules (H2O) which are in no way different from the water inwhich they were produced.

Equipment for use in the deionization process may be of severaltypes. Available are both multiple bed and single bed units. Multiplebed units have pairs of tanks, one for the cation bed, the other for theanion bed. Single bed units incorporate both the cation and anion ex-changers mixed in a single tank. Thus, the resin beads (both cation andanion types), as discussed above, are mixed. This provides millions ofopportunities for cation and anion exchange in the same unit, pro-ducing extremely high quality deionized water. In this type of deion-izer, the two types of resin (cation and anion) must be separated beforeregeneration, and then remixed before being used to deionize water.The 40% less weight typical of anion exchange resins facilitates theirseparation from cation resins. Sometimes concentrated salt brine isused to achieve the separation because anion resins tend to float in saltbrine while cation resins will sink. As in the two separate bed system,the cation exchange bed is regenerated with acid to replace the hy-drogen ions which have been given up. At the same time, the anionexchange bed is regenerated with caustic soda to replace the hydroxylions which have been given up.

Distillation

Efforts to produce very low conductivity water through multiple dis-tillation have proved to be extremely complex and require elaborateequipment.

Lime-Soda Ash Treatment

The hardness of water can be reduced through lime-soda ash treat-ment. This particular method of removing hardness is sometimes usedby municipal water plants to reduce the amount of calcium and mag-nesium in a water supply. While it is quite effective in reducing hard-ness, it is not a complete removal treatment.

Often when a city has a raw water source that has 35 to 40 grainhard water, the local water system will use the lime-soda ash treatmentto reduce hardness to between 5 and 10 grains.

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Iron

Iron is a common water contaminant. Iron in water can cause stainson porcelain and fabrics in concentrations exceeding 0.3 mg/L (partsper million).

In its insoluble forms, iron can form sludge deposits in pressuretanks, pipe lines, water heaters, and other plumbing appliances andfixtures.

Iron may be present in water in several forms:

1. soluble ferrous ions;2. ferric ions, soluble only in very acid water;3. ferric hydroxide, insoluble in neutral or alkaline water;4. ferric oxide, which appears as particles of rust from pipes;5. in combination with organic compounds or as iron bacteria.

Frequently when iron is drawn from a well, it is present in waterin the soluble ferrous state. Upon exposure to the air, molecular oxy-gen begins to enter the water as carbon dioxide escapes. The oxygenthen oxidizes the ferrous ions (Fe++) changing them to ferric ions(Fe+++). At this point, the ferric ions combine with free hydroxyl ions(OH–) to form the insoluble gelatinous compound ferric hydroxide[Fe(OH)3].

The water is initially clear and colorless when pumped, but as theindividual molecules join together, characteristic rust color (oftencalled “red water” or “rusty water”) appears. And finally, a gelatinousprecipitate of ferric hydroxide settles to the bottom of the container.In this way, the soluble ferrous ions convert into the insoluble ferrichydroxide.

When iron is found in surface supplies, the water may well be ex-tremely acid, or the iron may be combined in various complex mole-cules which resist oxidation. In some surface waters, iron may bepresent in an organic (chelated) form. Such water usually contains agreat deal of color.

Iron bacteria frequently thrive in iron-bearing water. As they de-velop, these bacteria form reddish-brown growths which may clogpipes and reduce flow rates. A decaying mass of these iron bacteria cancause bad tastes and odors in a water supply, as well as severe discol-oration problems.

The chemistry of iron removal is not difficult once the cause hasbeen clearly determined. Corrective measures present difficulties insome instances only because it is not always easy to determine the spe-cific type of iron and because the operation of mechanical equipmentmay, perhaps, be unfamiliar.

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Iron Removal

Generally speaking, there are three basic methods for removing ironfrom water:

(1) Ion Exchange Water Softeners

The use of a water softener unit is considered a satisfactory way to re-move limited amounts of iron from water supplies. No hard and fastrules can be given on the amount of iron that can be treated. The an-swer in each case depends upon the design of the softener as well ason other variables.

(2) Oxidation and Filtration

(a) Iron Filters

For medium concentrations of iron, the use of an oxidizing filter canbe a most effective means of treatment. When used, such a filter shouldbe placed in the water line ahead of the softener. Oxidizing filters nor-mally contain a base material that has been coated with manganesedioxide. This may be a manganese-treated greensand, a manufacturedmanganese material, natural manganese-bearing ores, and similar ma-terials. These manganic oxide deposits convert the soluble ferrous ironin the water into ferric iron. As the ferric hydroxide forms, it is filteredfrom the water by the granular material in the tank.

(b) Feeding Oxidizing Agent (chlorine, ozone, peroxide, orpotassium permanganate plus a filter)

With high concentrations of iron, small pumps, eductors, or other de-vices may be used to feed chemical oxidizing agents such as householdbleach, ozone, hydrogen peroxide, or a permanganate solution in thewater. Like the manganese dioxide in an iron filter, these chemical ox-idizing agents convert ferrous iron to the ferric state. The contaminantcan then be removed by running the water through a simple sand filterwhen chlorine, ozone, or peroxide is used and through an iron filterwhen permanganate is used.

(c) Superchlorination-Dechlorination

The use of superchlorination-dechlorination as a means of removingiron from water has been gaining more and more acceptance. This vari-ation of the oxidation-filtration concept uses two basic devices: onefeeds chlorine into the water, the other removes any excess chlorine.

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The injection of the full strength chlorine bleach into the water con-verts the ferrous iron into ferric iron. This insoluble gel is then removedfrom the water by simple filtration. A dechlorinator unit (usually anactivated carbon filter), either on the main line or at one or more spe-cific outlets, is then installed to remove excess chlorine, especially fromdrinking water. These methods are appropriate for use in treating wa-ters that have an essentially neutral pH.

(3) Sequestration (Use of such materials as polyphosphates)

(a) Solution Feeders

There are, in addition, several other techniques for the control of sol-uble iron in water. These make use of polyphosphates to keep the ironin solution. Polyphosphates do not remove iron from water. Ratherthey stabilize and sequester the iron so that the water remains clearand does not produce iron stains. Solutions of the very soluble sodiumpolyphosphates may be fed into the water with various small chemi-cal feed pumps, eductor units, etc. In general, these units proportionthe feed to the flow of water. In this way, they provide economical useof the polyphosphate and accurate, consistent control of the concen-tration.

(b) Cartridge or Pot Feeders

A simple cartridge or pot-type feeder is a less expensive original unit.These feeders utilize polyphosphates which commonly have calciumor magnesium incorporated into them to provide a product which dis-solves slowly and evenly. The feeder is installed in the water line sothat all or a part of the water supply passes through the cartridge ortank. The water picks up some of the polyphosphate solution and car-ries it into the water line.

Manganese

Manganese is similar to iron and is often found in iron-bearing waters,but more rarely than iron. Chemically it can be considered closelyrelated to iron since it occurs in much the same forms. In low con-centrations, it causes objectionable stains. USEPA Drinking Water Reg-ulations place a limit of 0.05 mg/L (ppm) on the manganeseconcentration in water.

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Manganese Removal

The same methods as those outlined for iron removal apply equally tothe treatment of manganese problems. Light concentrations can be re-moved with a softener. Larger quantities call for chemical treatmentplus filtration.

Corrosion

Corrosion can be defined as the destructive disintegration—the “eat-ing away”—of metals due to electrochemical action. There are threebasic types of corrosion. They are due to:

1. oxygen and other dissolved gases2. acids3. galvanic action

A corrosive water becomes even more corrosive when certain fac-tors are altered, among these:

1. An increase in the temperature of water.2. An increase in the flow rate of water.3. Use of several different metals in a plumbing system.4. An increase in the dissolved solids content of the water.

It is sometimes difficult to pinpoint the specific cause of corrosionin any given instance, but it can have a very deleterious effect onplumbing.

Corrosion Control

If the problem is due to acid water conditions, install a limestone chipfilter to raise the pH. The limestone carbonate combines with the acidto produce inoffensive bicarbonates. In the case of soda ash feed, a so-lution of sodium carbonate neutralizes the acidity of the water. Whendissimilar metals must be joined together in a plumbing system, theuse of insulating unions can be used to break the flow of electric cur-rent between the two metals, thus preventing or at least significantlyretarding electrochemical corrosion. If trouble is due to corrosive hardwater, a polyphosphate solution feed or slow dissolving phosphatecompound provides possible solutions. In both cases, polyphosphatesform a glass-like coating on the inner metal surfaces of the plumbingsystem.

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Oxygen scavengers are widely used to treat corrosion due to dis-solved oxygen in water. These compounds (such as sodium sulfite) arewidely used for industrial treatment but are not used where water is in-tended for human consumption.

Various types of metal and nonmetal coatings can be used for con-trol of specific corrosion problems.

Alkalinity

Alkalinity can be due to hydroxides, carbonates, or bicarbonates inwater. Of the three, bicarbonates predominate in waters with pH be-tween 4.3 and 8.5. Bicarbonates, therefore, represent the most com-mon source of alkalinity. Alkalinity is not particularly noticeable whendue to bicarbonate ions except if present in large amounts. In contrast,alkalinity resulting from fairly small amounts of carbonate and hy-droxide ions is very noticeable.

Alkalinity Removal

Troublesome amounts of alkalinity can best be removed in the house-hold by reverse osmosis. Other methods include lime soda ash soften-ing, anion resin exchangers, and mineral acid feed which are moresuitable for industrial applications.

Free Carbon Dioxide

Almost all natural waters contain some carbon dioxide or carbonicacid. In waters containing relatively low concentrations of minerals,carbonic acid becomes a problem in that not sufficient alkaline saltsare present in the water to buffer its effect, and the water is corrosive.

Carbonic Acid Removal

The simplest method for removal of carbonic acid is to pass the waterthrough a tank containing limestone chips. The limestone in the filterreacts with the carbonic acid to produce calcium bicarbonate. In thesame way, lesser amounts of magnesium bicarbonate are formed.

A decarbonator is an effective device for remova] of excess carbonicacid. It consists of a packed tower into which water is fed from the topand allowed to trickle downward through a bed of plastic air spacers by

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gravity. A flow of air is directed upward through the tower which scrubsthe carbon dioxide out of the water. This removal can be illustrated asfollows:

H2CO3 → H2O + CO2

Chloride and Sulfate

Almost all natural waters contain chloride and sulfate ions. Low tomoderate concentrations of both chloride and sulfate ions add palata-bility to water and are desirable. Excessive concentrations can makewater unpleasant to drink. Both chlorides and sulfates contribute tothe total mineral content of water. The total concentration may have avariety of effects in the home, everything from extreme hardness toelectrochemical corrosion. Excessive sulfates may also have laxativeeffects.

Chloride and sulfate can be removed from household water sup-plies by reverse osmosis, deionization, or anion exchange, and by dis-tillation.

Fluoride

Fluoride in water can be detrimental or beneficial. It all depends on theconcentration. Research has shown that a concentration of about onemilligram per liter (ppm) of fluoride in drinking water reduces toothdecay.

When drinking water contains excessive fluoride above two ppm,it causes “endemic dental fluorsis.” Sometimes called “Colorado BrownStain,” it appears as a dark brown mottling or spotting of the teeth orcauses them to become chalky white. Above four milligrams of fluo-ride per liter crippling skeletal fluorosis, a serious bone disorder, andosteopetrosis or brittle “marble” bones, can occur.

Many cities now add fluorides in concentrations of one ppm towater to reduce tooth decay. Reverse osmosis and distillation are thesimplest household methods for reducing the fluoride content of water.

Hydrogen Sulfide

Hydrogen sulfide is a gas present in some waters. There is never anydoubt as to when it is present due to its offensive “rotten egg” odor ap-parent in concentrations as low as 0.5 ppm.

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There are several methods for removing hydrogen sulfide fromwater. Most of them involve converting the gas into elemental sulfur.This insoluble yellow powder can then be removed by filtration.

Low to moderate concentrations of hydrogen sulfide can be elim-inated through use of an oxidizing filter of the same type satisfactoryfor iron removal. Because the elemental sulfur precipitate tends to clogthe filter material, it is usually necessary to replace this material fromtime to time.

Chemical treatment is recommended for medium to high concen-trations of hydrogen sulfide. In such cases, solutions of householdbleach, ozone, peroxide, or potassium permanganate serve as satisfac-tory oxidizing agents.

An activated carbon filter alone will remove trace amounts of hy-drogen sulfide. In this process, the carbon simply adsorbs the gas onits surface. The use of an activated carbon filter can be economicalwhen small amounts of the gas are present.

Nitrate (Nitrate Nitrogen)

Many groundwaters contain small amounts of nitrate nitrogen.Concentrations range from 0.1 ppm to 3 or 4 ppm in most areas. Thepresence of nitrates in a water supply indicates possible pollution ofthe water with animal wastes. In concentrations as low as 10 to 20 mil-ligrams per liter, nitrate nitrogen has caused illness and even deathamong baby animals and infants under six months of age. Althoughthis problem is serious, public health officials are also concerned withnitrates as a strong indicator of water polution. Certainly, where agroundwater is known to contain little or no nitrate nitrogen naturally,the appearance of any significant increase is a probable indication ofpollution. Here, prevention of sewage contamination is the best pos-sible treatment.

Bottled water is a practical source of nitrate-free water for infants.The best method for removal of nitrates in the household is reverseosmosis, anion exchange, or deionization.

Oxygen

Dissolved oxygen, as a cause of corrosion, can be a source of serioustrouble in a household supply. While a number of chemicals effects itsremoval, polyphosphates or sodium silicates can be used for hometreatment needs.

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Silica

Silica exists in water in two forms—as both dissolved or reactive silicicacid and associated anions and undissolved or inert particulates andcolloids. In the colloidal form, it consists of very fine particles in sus-pension. These can usually be removed by coagulation and settling orfiltering. In the dissolved form, silica is slightly soluble and extremelydifficult to remove by either chemical or physical means. Fortunately,silica is not too significant in water for household uses.

Sodium

Sodium salts are present to a greater or lesser degree in all natural wa-ters. Their concentrations vary from a few parts per million in somesurface supplies to several hundred grains per gallon in certain wellsupplies, and up to 30,000 to 40,000 ppm in areas of the oceans.

Extremely soluble, sodium salts do not form scale when water isheated, nor do they produce curd when combined with soap. Highconcentrations tend to increase the corrosive action of water and maygive water an unpleasant taste. Sodium ions in large amounts alsohamper the operation of ion exchange softeners. Where water containsmuch hardness and sodium, several grains of hardness may remain inthe softened water.

Reverse osmosis, along with distillation or deionization, are effec-tive methods for removing sodium from household water supplies.Again, the use of sodium-free bottled water is a viable alternative.

Biological Contaminants

No discussion of water would be complete without a review of the im-portant palatability and potability factors. While both are necessary indrinking water, the factor of potability is as vital as life itself.

Potability

A water is potable, that is, safe to drink, when and only when it is freeof disease organisms, as well as toxic chemical contaminants.

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Biological Organisms

Most of the organisms in water are harmless. In fact, some are ex-tremely beneficial to man. Others are of little interest or have a mildnuisance value. Some, however, are the source of disease. The poten-tial disease-producers are found among five of the subgroups of livingorganisms in water. They include bacteria, protozoa, worms, viruses,and fungi. The presence of certain organisms of these various types canlead to such infectious diseases as typhoid fever, dysentery, giardiasis,cryptosporidiosis, cholera, jaundice, hepatitis, undulant fever, andtularemia.

There are many ways in which all waterborne organisms can beclassified. For our purposes, it is sufficient to state that they may bemembers of either the plant or animal kingdoms. Broad and easy clas-sifications to be sure, still there are some species that even men of sci-ence cannot readily classify.

Among the lower life forms found in water are algae, diatoms,fungi, molds, bacteria, worms, protozoa, nematodes, and viruses.Where there is even the possibility that water contains pathogenic or-ganisms, that supply must be considered contaminated. It is not safeto take risks with water where human life is involved.

While there is a large and varied number of pathogens, no singlecontaminated water supply is apt to contain more than a few of thecountless variety. On one hand, this is fortunate. At the same time, itmakes detection of pathogens extremely difficult in terms of a routinewater analysis.

Since both speed and accuracy are essential, laboratory technicianslook for coliform bacteria as an indicator of the fact that water is con-taminated. Fecal coliform bacteria grow in the intestines of human be-ings and other warm-blooded animals, thus their presence in a watersample is an indication of sewage contamination. Coliform bacteriaare used as a biological indicators of contamination or pollution be-cause they can satisfy the following requirements:

They serve as a reliable measure of contamination indicatingthe possible presence of specific contaminating organisms ineither a natural water or one subjected to treatment. Further,they react exactly as do possible contaminating organisms inwater.

Coliform bacteria can be expected to be present in a con-taminated water in greater number than are the contaminat-ing organisms.

They are readily identifiable as a result of relatively simpleanalytical tests.

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Bear in mind that coliform bacteria show that water is contami-nated with animal wastes and is possibly contaminated withpathogens. The coliform bacteria, except for rare strains such as0157:H7, are not disease organisms or pathogens in themselves. Onthe other hand, their absence is no absolute assurance that water doesnot contain pathogens.

The USEPA Primary Drinking Water Regulations indicate that watershould contain an absence of coliform organisms or, in multiplesamples, no more than one coliform-positive sample result when 5 to39 samples are analyzed in a month, and no more than five percentcoliform-positive sample results when 40 or more samples are analyzedeach month.

Note: The standard of approximately one or five percent coliform-positive sample results is, of course, a standard of expediency. With evena single organism of this type in the water, there is always the possibility,though remote, of infection. Recognizing the danger, what can bedone to provide adequate protection against contamination?

Above all else, when a water supply becomes contaminated, cor-rect the problem at once. It is a basic rule of water sanitation to get tothe source of the problem and eliminate it. If a well, for example, be-comes badly contaminated, it is necessary to trace the contaminationto its source and, if possible, remedy the situation. The best solutionsare protecting the source of water supply, finding a new source of sup-ply, or disinfection of the water.

There are a number of ways of disinfecting waters. Among theseare boiling, ultraviolet light, and a number of chemicals, the mostwidely used of which is chlorine. Each method has its advantages. Inevaluating them, the following points should be considered:

1. A disinfectant should be able to destroy all types ofpathogens and in whatever number present in the water.

2. A disinfectant should destroy the pathogens within the timeavailable for disinfection.

3. A disinfectant should function properly regardless of anyfluctuations in the composition or condition of the water.

4. A disinfectant should function within the temperature rangeof the water.

5. A disinfectant should not cause the water to become toxicor unpalatable.

6. A disinfectant should be safe and easy to handle.7. A disinfectant should be such that it is easy to determine its

concentration in the water.8. A disinfectant should provide residual protection against re-

contamination.

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At present, chlorination in one form or another is regarded as themost effective disinfectant available for all general purposes.

Chlorine is normally fed into water with the aid of a chemical feedpump. The first chlorine fed into the water is likely to be consumed inthe oxidation of any iron, manganese, or hydrogen sulfide present inany supply, including bacteria, if present.

When the “chlorine demand” due to these materials has been sat-isfied, what’s left over—the chlorine that has not been consumed—remains as a “chlorine residual.”

The rate of feed is normally adjusted with a chemical feed pumpto provide a chlorine residual of 0.5–1.0 ppm after 20 minutes of con-tact time. This is enough to kill coliform bacteria but may or may notkill all viruses or cysts.

Some authorities believe greater feed and shorter contact time is amore effective technique. In their opinion, a satisfactory chlorine resid-ual alone can provide adequate control. The technique that fills thisbill is superchlorination-dechlorination.

A superchlorination-dechlorination system consists of two basicunits. The chlorinator feeds a stepped-up chlorine dosage into thewater to provide a residual of 3.0 to 5.0 ppm. The dechlorinator thenremoves the excess chlorine from the water before it reaches the house-hold taps.

In evaluating the use of chlorine, bear these points in mind:

1. “Free” chlorine residuals are more effective than “combined”or “chloramine” residuals.

2. A pH of 6.0 to 7.0 makes water a far more effective mediumfor chlorine as a disinfecting agent than do higher pH val-ues above pH 7.5 or 8.

3. The effectiveness of chlorine residuals increases with highertemperatures within the normal water temperature range.

4. All types of organisms do not react in the same way undervarious conditions to chlorination.

5. An increase in the chlorine demand of a water increases theamount of chlorine necessary to provide a satisfactory chlo-rine residual.

Palatability

To be palatable, a water must be free of detectable tastes and odors. Itmust also be free of turbidity as well as strong color. Tastes and odorscan be traced to one or more of the following: decaying organic matter;living organisms; iron or manganese; the metallic products of corro-

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sion, industrial waste pollution, chlorination; and high mineral con-centrations.

Taste and Odor Removal

Some tastes and odors, especially those due to organic substances, canbe removed from water simply by passing it through an activated car-bon filter. Other tastes and odors may respond to oxidizing agents suchas chlorine, ozone, and potassium permanganate.

In any case, you may have to try a number of methods in an at-tempt to rid a water of objectionable tastes and odors. If methods out-lined in this course do not work, it may be more economical to seekout a new source of drinking water.

Turbidity

Noticeable amounts of suspended solids also can affect our zest forwater. For although turbid water may be safe to drink, it is apt to bedistasteful to most individuals. The USEPA Interim Primary DrinkingWater Regulations recommend that the amount of turbidity in waterbe less than 0.5 Turbidity Unit (T.U.), or 5 T.U. under special condi-tions.

Precisely speaking, the terms “turbidity” and “suspended matter”are not synonomous. Suspended matter is that material which can beremoved from water through filtration or the coagulation-filtrationprocess. Turbidity is a measure of the light absorbed by water becauseof its content of suspended matter.

Turbidity Removal

Mechanical filtration will remove most forms of turbidity. Of course,the smaller the turbid particles, the finer must be the filter openings.In many cases, filters containing specially graded and sized gravel andsand are effective. Periodic backwashing to remove all filtered matteris all the maintenance necessary.

Some filter manufacturers also provide a “filter-aid” which is addedonto the top of the filter bed immediately after backwashing. The filteraid traps fine dirt particles producing a more sparkling clear water andkeeps dirt from penetrating the filter bed. Further, in many cases, car-tridge filters are effective on drinking water lines.

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Thus far in this lesson, we have reviewed the wide range of con-taminants to be found in water. Now, how do we determine the kindsand amounts of various contaminants in water?

Determining the Kinds and Amounts of Water Contaminants

The best way to do this is through laboratory analysis of water. Suchanalyses may include either physical, microbiological, and/or chemi-cal tests, depending on the purpose.

Chemical analyses made by laboratories show iron in milligramsper liter and hardness minerals in grains per gallon or milligrams perliter.

Grains Per Gallon and Milligrams Per Liter (Parts Per Million) in Water Analysis

Use of both grains per gallon and milligrams per liter is a practice fol-lowed primarily for convenience in reporting concentrations of min-erals, some of which are abundant in water and some of which arefound only in trace quantities.

Actually, there are four units of measure commonly used in wateranalysis work: milligrams per liter (mg/L) or parts per million (ppm);grains per U.S. gallon (gpg); equivalents per million (epm); and grainsper imperial gallon (gpg imp).

To convert from milligrams per liter or parts per million to grainsper gallon, divide the former by 17.1. Similarly, multiply grains per gal-lon by 17.1 to arrive at milligrams per liter or parts per million.

In making all conversions, it is always necessary to keep to a com-mon unit of measure. Thus, 1 ppm must be translated as one ounce ina million ounces of water…not gallons, quarts, or pounds.

Further, the various hardness minerals, as we have seen, are gener-ally expressed in terms of hypothetical combinations as grains per gal-lon of calcium carbonate. The hardness conversion tables for variousminerals can be found in Lesson 9.

After the analysis of water and sales of proper corrective equipmenthave been made, the final step is installation of the product. Here, too,scientific care is necessary to insure ideal water at the installation site.To achieve such results, installers should place equipment so as to sup-ply soft water to all desired points of use in the simplest, neatest, leastexpensive and most effective way. In making placements, the location

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of the water heater, main supply line, sillcocks, drain, and electricaloutlet must be taken into account.

Further, it is essential to install a unit of the proper capacity. Failureto size equipment properly means unsatisfactory operation of the unit.Here the amounts of hardness and the number of persons using the in-stallation are paramount factors. In case of any doubt, refer to the Se-lection Guides in Lesson 10. It is also highly important to determinethe amount of pressure loss which will occur as a result of the instal-lation of a softener. If the pressure loss is too excessive, it may be nec-essary to install a second unit in parallel. Procedures for determiningpressure drop losses are discussed below.

The following points are important to the installation of a unit:

1. Shut-off valves should be installed on both the inlet and out-let lines, and an electrical ground strap installed across anybreak in metallic water lines made for the installation oftreatment equipment.

2. A bypass valve should be positioned on the main line in be-tween the inlet and outlet connections.

3. Softener drainage should be discharged into the householdwaste system.

4. The drain line should not be connected directly to the watersystem, as an air-breadth of at least one inch or two timesthe diameter of the drain line between the line and the out-let must be maintained.

5. At the time of installation, a softener should be disinfectedas a precaution against any possible contamination duringthe shipping, storage, or installation of the unit. Disinfec-tants commonly used for this purpose are sodium hypochlo-rite and calcium hypochorite.

Pressure Loss Through Water ConditioningEquipment

There is a definite and measurable resistance to flow due to friction forevery foot of pipe in a system. Every change in the direction of flow,every plumbing fitting, every restriction, produces some loss in the en-ergy developed by a water system. While some loss due to friction isinevitable, this must be kept to a minimum. Depending upon the sizeof the pipe, the number of fittings, the height to which the water mustbe delivered, a softener may be adequate or may be seriously under-sized. Pressure drop in the system varies with the length of pipe andthe number and type of fittings.

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These must be taken into account to determine just how great thehead loss may be. The tables in Lesson 10 will provide figures for eval-uating just how great the head loss or pressure drop will be throughthe entire system.

Summary

In studying these twelve lessons, it should be quite obvious by nowthat water conditioning is a large and extremely complex field. Aknowledge of chemistry, hydraulics, mathematics, bacteriology, andmuch more is necessary in gaining a well-rounded background.

It is our sincere hope that you will continue to review this mate-rial, making more and more of it your own as time goes on and thatyou will put the knowledge gained as a result to practical use at everyopportunity.

Also bear in mind that the course can and should be used as a ref-erence should any point occur for which you want any specific infor-mation. We hope you will find that Water Treatment Fundamentalscontinues to contribute not only to your own success, but to the well-being of the general public. For all of us, the aim should be to providewater of optimum quality for each use.

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Lesson 1Water and Water Quality

The first lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course covers the water cycle. It covers the sky-to-earth-back-to-sky cycle which water pursues.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Choose the word or phrase to make each of the following themost accurate statement possible.

1. Water is __________ compressible.

(a) very slightly

(b) moderately

(c) extremely

2. By weight, water consists of __________ hydrogen.

(a) 31.852%

(b) 11.188%

(c) 22.140%

By weight, water consists of __________ oxygen.

(d) 77.860%

(e) 88.812%

(f) 68.148%

3. Compared to the Great Lakes, sea water contains __________dissolved metallic substances.

(a) more

(b) less

(c) approximately the same amount of

4. Connate waters are sometimes referred to as “fossil waters.”They are troublesome because they generally have largequantities of __________.

(a) salt

(b) hydrogen sulfide

(c) toxic chemicals

5. Pollution of water is neutralized in nature through __________.

(a) sunlight

(b) bacteria

(c) algae

(d) oxidation

(e) all of the above

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Complete the following statements with the best possible word.

6. Many scientists regard water as the universal __________.

7. The treatment of water for the removal of impurities in whichtwo ion exchange resins are used is known as __________.

8. Water dissolves and collects carbon dioxide as it falls throughthe air. In doing this, the carbon dioxide and water react toproduce __________.

9. Limestone, a common rock formation, contains both calciumand magnesium __________.

10. Approximately __________ percent of all moisture that falls toearth is drawn back in the process of evaporation before it seepsinto soil or goes far in the process of run-off.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

11. Generally there are quite appreciable differencesin the amounts of dissolved hardness mineralsto be found in water drawn from a deep well and a reservoir. _____ _____

12. When water is rated “potable,” this means it is safe to drink. _____ _____

13. A heavy rainfall will wash all impurities out of the air giving it a fresh clean smell. _____ _____

14. The moisture in the atmosphere providesprotection against extremes of both heat and cold. _____ _____

15. As water falls either as rain or snow, it picks upsome hardness minerals even before it reaches the ground. _____ _____

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TRUE FALSE16. Normally when water reaches the earth, it is

slightly acid, corrosive, and relatively soft. _____ _____

17. Although springs are popularly believed toprovide sparkling clear water, some of them contain a good deal of turbidity at times. _____ _____

18. It is not generally necessary to chlorinatecistern water if it is to be used for drinking purposes. _____ _____

19. Surface waters, such as lakes and rivers, get agreat deal of their water from underground springs. _____ _____

20. Snowfall high in the mountains above thetimberline contains a great deal of hardness due to its exposure to rock formations. _____ _____

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Lesson 2Introduction to Water Chemistry

The second lesson in the Water Quality Association Water TreatmentFundamentals correspondence course covers the study of the activities ofatoms, molecules, ions, etc.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Complete the following statements with the word or wordswhich you feel will make the most accurate statement.

1. An electrically charged atom or group of atoms is called an__________.

2. If you were to take a large lump of coal and break it into smallerpieces, the resulting change would be called a “__________change.” If you set fire to some of these pieces, and they turn toashes, this is called a “__________ change.”

3. The molecular weight of ordinary water is __________.

4. The protons and neutrons in an atom are to be found in the__________.

5. One prevalent theory is that all elemental atoms seek to achievethe __________ type of structure.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

6. A molecule is a substance which cannot bedecomposed chemically into a simpler substance. _____ _____

7. Scientists claim there are more than 25 basicparticles, such as the electron, proton, and neutron in the universe. _____ _____

8. An electron can be said to possess an electrical charge of –1. _____ _____

9. The number of protons in the nucleus of anatom equal the number of neutrons outside the nucleus. _____ _____

10. Paper, egg yolk, nylon, air, turbid water, and gasoline are all mixtures. _____ _____

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TRUE FALSE11. The number of electrons surrounding the

nucleus of an atom have little effect on its weight. _____ _____

12. The evaporation of moth balls is an example of chemical change. _____ _____

13. The valence of an element is expressed as asmall whole number and is designated as plus or minus. _____ _____

Choose the answer which makes each of the following themost accurate statement possible.

14. The weight of an electron is extremely light. Actually its specificweight is

(a) .000,000,000,000,000,000,000,000,009 gram

(b) .000,000,000,000,001,837 gram

(c) .000,000,000,009 gram

15. Some of the following statements are true of compounds, butnot of mixtures. Check those which are true of compounds.

(a) have definite freezing and boiling points

(b) can be separated into constituents by purely physicalmeans

(c) are heterogenous, that is, relative amounts of theconstituents may vary within the sample.

(d) usually retain the properties of their constituents

(e) have compositions that may be varied over a limitedrange

16. Atoms with the same atomic number but with different atomicweight due to varying numbers of neutrons in their nuclei arecalled __________.

(a) electrons

(b) ions

(c) isotopes

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17. The molecular weight of gaseous nitrogen (N2) is __________.

(a) 7

(b) 14

(c) 21

(d) 28

(e) none of these

18. A mixture is a substance with __________.

(a) a definite freezing point

(b) no exact ratio of ingredients

19. Many ions are electrically charged __________.

(a) groups of atoms

(b) protons

(c) electrons

20. __________ determine chemical properties.

(a) Protons

(b) Neutrons

(c) Electrons

__________ simply add mass.

(d) Protons

(e) Neutrons

(f) Electrons

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Lesson 3Water Chemistry

The third lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course moves into the specific subject of waterchemistry.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

1. Place the following true or false statements about water underthe correct headings below:• Water is an excellent oxidizing agent.• Water is an extremely stable compound.• Water does not ionize readily.• Water contains 88.81% oxygen.• The chemical formula for water is H2O.• When water ionizes, it forms hydrogen and oxygen ions.• At a temperature of 3600°F, only about 22% of water

molecules dissociate or separate.

TRUE STATEMENTS ABOUT WATER:

INACCURATE STATEMENTS ABOUT WATER:

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

2. A negatively charged ion in solution is known as an anion. _____ _____

3. Hydrogen has a valence of +2. _____ _____

4. The bicarbonate ion consists of 1 hydrogen atom, 3 carbon atoms, and 1 oxygen atom. _____ _____

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TRUE FALSE5. When sodium, with a valence of +1, combines

with chlorine, which has a variable valence, thechloride ion in the compound must have a valence of –1. _____ _____

6. Acids can be defined as compounds that release oxygen ions (O–) in a solution. _____ _____

7. Bases are substances which can release hydroxide (OH–) ions. _____ _____

8. Salts consist of metallic cations and nonmetallic anions. _____ _____

9. Carbon dioxide is a compound present in mostwater supplies in concentrations ranging from almost zero to about 50 parts per million. _____ _____

10. Iron, present in water exceeding as little as 0.3ppm, can cause brown stains on plumbing fixtures and laundry. _____ _____

11. Choose correct statements (one or more) that apply in regard toacids.

(a) Acids can be defined as compounds that releasehydrogen ions (H+) in a solution.

(b) Acids have a bitter rather than a sour taste.

(c) Molecules of various acids are capable of releasingdifferent amounts of free hydrogen ions in solution.

(d) Acids will turn litmus paper blue.

(e) Acids react with bases to form a salt and water.

continued on next page

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12. Choose statements (one or more) that are correct in regard toions.

(a) In solution many molecules dissociate into two ormore ions which theoretically are free to move andwander around.

(b) Ions are the same as atoms or molecules except thatthey carry positive electrical charges.

(c) The terms “cation” and “anion” were developedabout 40 years ago by early pioneers in the waterconditioning industry.

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Lesson 4Water Hardness

The fourth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course covers the problem of hard water andvarious methods of treating it.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Choose the answer (one or more) for each of the followingwhich you feel will make the most accurate statement.

1. Water that contains __________ of hardness minerals isconsidered slightly hard.

(a) 0–2 gpg

(b) 5–7.5 gpg

(c) 1–3.5 gpg

(d) 7.0–10.5 gpg

2. Soft water has the following characteristics:

(a) Soft water saves on soap.

(b) Soft water is ideal for watering household plants.

(c) Soft water prevents the clogging of hot water pipes.

(d) Soft water prevents unsightly rings around thebathtub.

3. The hardness minerals (calcium and magnesium) are __________in the earth.

(a) rather common

(b) scarce

They are found only in __________.

(c) elemental form

(d) various compounds

4. Hypothetically, the specific magnesium compounds commonlyfound in natural waters include __________.

(a) magnesium hydroxide

(b) magnesium sulfate

(c) magnesium chloride

5. Common calcium compounds found in the earth include__________.

(a) chalk

(b) marble

(c) limestone

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6. On a sliding scale of solubility, the most soluble hardnesscompound is __________.

(a) calcium carbonate

(b) magnesium chloride

(c) calcium sulfate

7. If water contains 7 gpg of magnesium bicarbonate asMg(HCO3)2, this concentration is equivalent to __________grains per gallon of calcium carbonate.

(a) 23.5

(b) 18.0

(c) 2.3

(d) 4.8

(e) none of these

8. Hardness minerals in the softening process are __________ bysodium ions.

(a) destroyed

(b) neutralized

(c) replaced

9. A fully automatic water softener contains a __________.

(a) water flow meter

(b) time clock

(c) hardness sensor

10. Circle the correct answers. A softener with a rated capacityof [50,000 / 10,000 / 7,000] grains operating on a 15 grains pergallon water will soften approximately [3,333 / 2,000 / 1,000]gallons of water before regeneration will be necessary.

continued on next page

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11. Below are a number of statements regarding the completeremoval of minerals from water through two ion exchangematerials. Indicate the proper sequence of these statements bynumbering from 1 through 9. Note that a few steps in thisprocess occur almost simultaneously.

(a) An anion exchange material with replaceablehydroxyl ions.

(b) Hydrogen ions replacing metallic ions on achemically equivalent basis.

(c) Water containing hydrogen and hydroxyl ions.

(d) Nonmetallic anions affixing themselves to the anionresin.

(e) A cation exchange material containing replaceablehydrogen ions.

(f) Hydrogen and hydroxyl ions combining to producewater molecules.

(g) Metallic cations affixing themselves to the cationresin.

(h) Partially treated water containing hydrogen ions andanions originally present in the solution.

(i) Hydroxyl anions replacing the nonmetallic anions inthe water.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

12. In order to convert grains per gallon into parts per million, divide grains per gallon by 17.1. _____ _____

13. Hardness minerals can be removed from water through the lime soda-ash treatment. _____ _____

14. The lime soda-ash treatment of hard water isused exclusively for industrial treatment of boiler waters. _____ _____

15. A chemist may report the results of a water analysis in terms of hypothetical combinations. _____ _____

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TRUE FALSE16. When water containing carbon dioxide comes

into contact with magnesium carbonate, thislatter compound is transformed into highly soluble magnesium bicarbonate. _____ _____

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Lesson 5Iron and Manganese

The fifth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course deals with iron and manganese prob-lems and possible methods of treating these conditions.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Below are 24 statements regarding iron and manganese, aswell as the treatment of these problems. Check whether thefollowing statements are True (completely accurate) or False(incorrect in any degree) in the proper column.

TRUE FALSE1. The Drinking Water Regulations of the U.S.

Environmental Protection Agency recommend a maximum of 0.1 ppm of iron in potable water. _____ _____

2. Iron may be found in water as soluble ferrousions and as suspended ferric hydroxide as well as in other compounds. _____ _____

3. Iron bacteria grow in stagnant waters where no light is available because they are aerobic. _____ _____

4. Manganese-bearing water produces dark brownor black stains on fabrics when these are washed in such water. _____ _____

5. Manganese is generally more readily removedfrom water than iron because it oxidizes more rapidly. _____ _____

6. Because iron problems are due to a variety ofconditions, no snap judgements should bemade as to the proper treatment in any given case. _____ _____

7. A water softener will remove small amounts of ferrous iron. _____ _____

8. Oxidizing filters are an inexpensive method ofremoving iron from water. When the rustaccumulation on these filters becomesexcessive, the unit can be cleansed bybackwashing. Some necessitate regeneration of the oxidizing media. _____ _____

9. When chemical oxidizing agents such ashousehold bleach are injected into iron-bearingwater, they convert the ferrous iron into ferric iron. _____ _____

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TRUE FALSE10. Polyphosphates are sometimes added to water

containing ferrous iron to prevent iron staining. _____ _____

11. Eductors are sometimes used to draw solutions of oxidizing agents into a water supply. _____ _____

12. Iron-bearing waters with a low pH requiredifferent treatment for removal of the iron than waters which are neutral or alkaline. _____ _____

13. Organic (chelated) iron, although occurring insome shallow wells, and even more rarely indeep wells, is found primarily in certain surface waters. _____ _____

14. Solutions of potassium permanganate will servewell as agents to oxidize high concentrations of ferrous iron. _____ _____

15. Telltale evidence of iron bacteria’s presence in ahousehold water supply is sometimes providedby the appearance of slimy red fungus-like growths in the flush tanks of water closets. _____ _____

16. The homeowner must take into considerationthe concentration of iron in a water supply incalculating the frequency of regeneration of his ion exchange softener. _____ _____

17. Ferrous iron can give water a somewhatastringent and disagreeable taste. This can ruinthe flavor of tea and coffee made with iron water. _____ _____

18. Underground water supplies which are iron-bearing hold the iron in its insoluble ferric state as long as the water remains below ground. _____ _____

19. Tests show that potassium permanganate is oneof the most effective ways to oxidize chelated iron. _____ _____

continued on next page

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TRUE FALSE20. Although some individuals popularly refer to

iron bacteria as “crenothrix,” this is actually only one form of the bacteria. _____ _____

21. When a water is acid and contains largeamounts of iron, the use of soda ash to raise thepH and chlorine to precipitate the iron is auseful technique. Later the precipitated iron can be removed with the use of a simple sand filter. _____ _____

22. Polyphosphates can not be used for control of iron in municipal water supplies. _____ _____

23. The presence of iron in surface water supplies is rather uncommon. _____ _____

24. A low pH and a high carbonate content arecharacteristic of natural waters that contain ferrous iron. _____ _____

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Lesson 6Corrosion

The sixth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course discusses the causes of corrosion andvarious methods of treatment available.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Complete the following statements with the word or wordswhich you feel will make the most accurate statement.

1. Major causes of corrosive water are __________ __________,__________, and __________.

2. Substances which ionize in solution and, as a result, make watercapable of carrying electric current are known as __________.

3. Neutralizing filters are used primarily to correct ____________________ conditions. They have several advantages. List twoof them below:

4. Some experts state that the number one and permanent causeof corrosion problems is due to __________.

5. The use of water softening equipment is neither the __________nor __________ of corrosion problems.

Check whether the following statements are True (com-pletely accurate) or False (inaccurate in any detail) in theproper column.

TRUE FALSE6. Iron, one of the common metals found in the

earth’s crust, is also the most susceptible to corrosion. _____ _____

7. “Rust” is the common name for the corrosion product of iron. _____ _____

8. When oxygen combines with another substance, the process is called oxidation. _____ _____

9. Oxidation occurs only when oxygen combines with metals. _____ _____

10. Acid waters can always be traced to thepresence of carbon dioxide or dissolved gases in water. _____ _____

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TRUE FALSE11. Acid waters cause deep pitting of the walls of

metal containers. _____ _____

12. Insulating unions at bimetallic points reduce galvanic corrosion almost completely. _____ _____

13. Polyphosphates can be effective in reducing the corrosive action of water on metal waterpipe. _____ _____

14. Soda ash treatment does not increase the hardness of a water supply. _____ _____

15. Polyphosphate films are permanent as long aswater is kept under 120°F. Higher temperatures tend to dissolve these thin protective films. _____ _____

Choose the answer which makes each of the following themost accurate statement possible.

16. A measurable electric current can __________ be produced as theresult of two dissimilar metal plates being immersed in water.

(a) always

(b) never

(c) sometimes

17. Oxygen scavengers are an inexpensive method of treatingcorrosion due to free oxygen in industrial water supplies. Oneof the common oxygen scavengers is __________.

(a) ferrous hydroxide

(b) potassium sulfate

(c) sodium sulfite

18. One common cause of corrosion is __________.

(a) iron bacteria

(b) extremely hard water

(c) free oxygen

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19. Soda ash treatment has certain disadvantages which are listedbelow. One of the statements, however, is inaccurate. Check thestatement that is inaccurate.

(a) Soda ash treatment has a relatively high initial cost.

(b) It may be necessary to install a sand filter if iron ispresent.

(c) Soda ash is used mainly for industrial applications.

(d) Soda ash treatment adds hardness to the water.

20. When the temperature of water is raised from 60°F to 140°F, therate of corrosion of steel may increase some __________ timesthe normal rate.

(a) 6 to 8

(b) 3 to 4

(c) 8 to 10

(d) 3 to 5

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Lesson 7Total Dissolved Solids, Hydrogen Sulfide,Fluorides, and Other Water Impurities

The seventh lesson in the Water Quality Association Water TreatmentFundamentals correspondence course covers a number of miscellaneousproblems that render many water supplies difficult to use. Among the con-taminants to be discussed are hydrogen sulfide, chlorides, and sulfates,nitrate nitrogen, and others.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Complete the following statements with the word or wordswhich you feel will make the most accurate statement.

1. Carbon dioxide in solution reacts with water to form ____________________. This resulting substance is a __________ [ weak /strong ] acid.

2. Authorities recommend a fluoride concentration at __________ppm to reduce tooth decay. Higher amounts in water canproduce two bad effects which are: __________ and __________.

3. The USEPA Drinking Water Regulations recommend amaximum of __________ ppm for sulfates in drinking water.

4. Several methods for control of the hydrogen sulfide probleminvolve conversion of the hydrogen sulfide gas into ____________________.

5. Silica may be present in water in both __________ and__________ forms.

Choose the answer in each of the following that makes themost accurate statement.

6. Analyses show that rainwater __________ contains free oxygen.

(a) often

(b) seldom

(c) always

(d) never

7. Large concentrations of fluorides in natural waters are anadverse health danger for which of these reasons:

(a) They produce dark brown mottling on teeth.

(b) They cause bone disorders.

(c) They are highly inflammable.

(d) They cause heavy decay of teeth among adults.

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8. A significant increase in the nitrate nitrogen concentration inwater is an indication of __________ contamination.

(a) possible

(b) probable

(c) certain

9. Which of the following are possible methods of treatment ofalkaline waters? (select one or more)

(a) ion exchange softener

(b) hydrochloric acid feed

(c) demineralization

(d) solution pot feed with polyphosphates

(e) chloride anion exchange

10. Hydrogen sulfide in solution makes a water unsatisfactory forsome of the following reasons. (select one or more)

(a) unpleasant odor

(b) corrosiveness

(c) inflammable even in low concentrations

(d) an indicator of coliform bacteria

(e) tarnishes silver

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

11. Strongly alkaline waters have an objectionable soda taste. _____ _____

12. Carbonic acid slows down the corrosive actionof water in both copper and galvanized plumbing equipment. _____ _____

13. Low to moderate concentrations of chlorides and sulfates may add palatability to water. _____ _____

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297

TRUE FALSE14. The use of aluminum sulfate, magnesium, or

calcium phosphate for removal of excessiveamounts of fluorides sometimes seriously affects the pH of water. _____ _____

15. The “rotten egg” odor of hydrogen sulfide canbe readily detected in water containing as little as 0.5 ppm of this gas. _____ _____

16. Well waters generally contain less hydrogen sulfide than do surface waters. _____ _____

17. Nitrate nitrogen does not usually present any problem in water for industrial purposes. _____ _____

18. In large amounts sodium salts in the waterlower the efficiency of ion exchange softener in the removal of hardness. _____ _____

19. Methane gas is frequently referred to as marsh gas. _____ _____

20. Activated carbon filters are recommended foruse in removing large concentrations of hydrogen sulfide in solution. _____ _____

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Lesson 8Taste, Odor, and Turbidity

The eighth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course discusses pathogens, coliform bacteria,turbidity, tastes, odors, and color.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Complete the following statements with the word or wordswhich you feel will make the most accurate statement.

1. There are a number of diseases due to waterborne infection.Among diseases due to infectious organisms in water are__________ __________ and __________.

2. Laboratory analysts determine if a water is contaminated bytesting for the presence of __________ __________.

3. When using ultraviolet light for water purification, it isnecessary that the water be free of __________ and __________.

4. If you saw the words nasturtium, pigpen, geranium, cucumber,and grassy used in connection with some water samples, youwould know that references are being made to __________.

5. If chlorination fails to provide complete pathogen kill, thefailure may be due to the homeowner’s effort to set the chlorinefeed rate very low in order to avoid having a ____________________ in the water.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE

6. Algae are among the most primitive oforganism forms on earth. They live as parasites,taking food from other living creatures found primarily in water. _____ _____

7. Aerobic bacteria thrive only when there is a supply of free oxygen in water. _____ _____

8. All bacteria present in water are classified as pathogens. _____ _____

9. Viruses may produce disease in plants as well as animals. _____ _____

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TRUE FALSE10. Chlorine demand is the amount of the total

chlorine fed into water that is neutralized in theoxidation of metallic impurities before a pathogen kill can be effected. _____ _____

11. Boiling water for one minute is effective indestroying pathogens because no heat-resisting bacteria are disease producers. _____ _____

12. Tastes and odors in water are due frequently to living organisms or industrial wastes. _____ _____

13. Rotifers are a type of worm found in the mudand slime at the bottom of rapidly moving fresh waters. _____ _____

14. Reverse osmosis with no bypass through sealleaks or other imperfections will remove virusesfrom water. _____ _____

15. A water that is safe to drink is sometimes referred to as a potable water. _____ _____

Choose the answer which makes each of the following themost accurate statement possible.

16. Disease bacteria are sensitive to the pH of a water. They live best in

(a) alkaline waters with a pH between 7.5 and 9.0.

(b) acid waters with a pH of 2.0 to 3.0.

(c) neutral waters with a pH of 6.5 to 7.5.

17. The discharge of coliform bacteria in a person’s waste materialin a single day may run as high as

(a) 50.

(b) 350,000.

(c) 400,000,000,000.

(d) 2,000,000.

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18. Some detergents have proved to be effective in destroyingpathogens in water. Those producing the best results are the

(a) neutral detergents.

(b) cationic detergents.

(c) anionic detergents.

19. Silver has proved an effective way to reduce the numbers oftotal bacteria because

(a) silver inactivates some microorganisms on contact.

(b) it is easy to maintain a proper amount of silver inwater.

(c) silver has a residual effect that may last for days.

20. Removal of excess chlorine is provided in the treatment of water to

(a) effect a pathogen kill.

(b) remove certain metallic ions from water.

(c) make the water palatable.

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Lesson 9How to Interpret Water Analyses

The ninth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course discusses how to evaluate and use awater analysis. Also covered are steps necessary to correct problems indi-cated by the water sample.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Complete the following statements with the word or wordswhich you feel will make the most accurate statement.

1. Sometimes families have contaminated water supplies whichthey drink without ill effects. They may be able to drink thewater because of __________ they have built up to a givendisease over a period of time.

2. There are three types of water analyses possible with any givenwater sample. They are microbiological, __________, and__________.

3. USEPA Drinking Water Regulations specify for public watersystems that the maximum contaminate level (MCL) ofcoliform bacteria acceptable in safe drinking water is __________in __________ percent of all samples collected and analyzed.

4. In reporting the condition of a water at the time a test sample istaken, both odor and __________, as reported by thehomeowner, should be noted.

5. One grain per gallon is the equivalent of __________ parts permillion.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE6. Because iron conditions may be due to a

number of causes, it is not sound practice to treat all iron problems in the same way. _____ _____

7. Carbon dioxide can be removed by causingwater to pass through a neutralizing filter containing limestone chips. _____ _____

8. Even low to moderate concentrations ofchlorides and sulfates make a water unpalatable. _____ _____

9. In large quantities, the sodium salts in the rawwater can lower the efficiency of an ion exchange softener in the removal of hardness. _____ _____

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TRUE FALSE10. Either glass bottles or polyethylene containers

may be used to hold water samples to be sent to a laboratory. _____ _____

11. All owners of private wells are required tosubmit water samples for bacterial study to their state boards of health once each year. _____ _____

12. On water analysis reports some mineralcontaminants are reported in ppm; others are reported in gpg. _____ _____

13. If a water analysis indicates color in terms of 60units, this would mean that the water wouldhave a definite color even when contained in a small glass. _____ _____

14. If a water analysis indicated that the watercontained 3.5 ppm of iron, the water wouldcause definite staining problems on fixtures and washable items. _____ _____

15. The term “ten parts per million of manganese”in a water would mean that in ten milliongallons of water there would be a total of ten pounds of manganese. _____ _____

Choose the answer which makes each of the following themost accurate statement possible. (Note: There is only oneright answer in each of the questions below.)

16. The recommended maximum concentration of chlorides indrinking waters is 250 ppm according to the USEPA DrinkingWater Regulations. What is this chloride content in grains pergallon expressed as calcium carbonate?

(a) 33.2 grains per gallon

(b) 17.1 grains per gallon

(c) 20.6 grains per gallon

(d) 7.9 grains per gallon

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17. If you were to express the sodium chloride content of a naturalwater in terms of calcium carbonate (CaCO3), you shouldmultiply the NaCl content by which of the following factors:

(a) 1.72

(b) 0.856

(c) 0.612

18. A water analysis shows 19.2 gpg of calcium chloride as CaCl2.To convert this to a corresponding amount of CaCO3, it wouldbe necessary to multiply by .902. The answer would be:

(a) 173 parts per million

(b) 17.3 grains per gallon

(c) 1.73 grains per gallon

19. If a water analysis showed 22.3 grains per gallon of sodiumbicarbonate expressed as calcium carbonate, this would be theequivalent of how many grains per gallon expressed as sodiumbicarbonate?

(a) 30.9 grains per gallon

(b) 374 parts per million

(c) 37.5 grains per gallon

20. For effective iron and/or manganese precipitation, the watermust provide at least __________ milligrams per liter ofalkalinity is recommended.

(a) 100

(b) 10

(c) zero

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Lesson 10Recommended Installation Procedures

The tenth lesson in the Water Quality Association Water Treatment Fun-damentals correspondence course discusses points to consider in regard tothe installation of water conditioning equipment.

Remember, the purpose of this prelesson assignment is to help you de-termine how much you know about the subject in advance of your studyof the material. Please answer all questions carefully. Determine whereyou need to give special attention to the material to come.

Prelesson Questionnaire

Below are statements regarding the installation of waterconditioning equipment. Check whether the following state-ments are True or False in the proper column.

TRUE FALSE

1. If a manufacturer’s installation instructions everare at variance with area building codes, the installer should follow the local regulations. _____ _____

2. It is generally advisable to soften water flowing into both the hot and cold water lines. _____ _____

3. The friction caused by water flowing throughpipes is not significant due to the fact that water is a liquid. _____ _____

4. Plumbing codes specify water flow rate/pressure drop considerations for softener installations. _____ _____

5. Although installations of softeners frequentlyinclude water for the toilets, this is done for reasons of convenience, never because of need. _____ _____

6. Rinsing of brine through the household piping is not an acceptable practice. _____ _____

7. Sodium hypochlorite can be effectively used for disinfecting softeners. _____ _____

8. In making a softener installation, the use of 3⁄4" pipe is considered best for all conditions. _____ _____

9. Oxidizing filters can be used to remove iron from water inexpensively. _____ _____

10. Water softener backwash effluents will not harm septic tank waste disposal systems. _____ _____

11. Both the size of the pipe and its length fromwater meter to taps should be taken intoconsideration with any softener or filterinstallation as a practical aid in keeping pressure drops to a minimum. _____ _____

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TRUE FALSE12. Disinfection of a softener at the time of

installation is done as protection againstpossible contamination during the process of shipping and installation of the equipment. _____ _____

13. A 1⁄2" globe valve, when fully open, produces apressure loss equivalent to that occuring when water flows through 18.6 feet of straight pipe. _____ _____

14. A pressure drop is twice as great in 100 feet of 1⁄2" pipe as it would be in 100 feet of 1" pipe. _____ _____

15. Upon completion of a softener installation,open taps on upper floors should be closed first.Last to be closed should be those on the basement or on the lowest level. _____ _____

Choose the answer which makes each of the following themost accurate statement possible.

16. A softener with a rated capacity of 50,000 grains installed in ahome having 45 grain water will soften approximately howmany gallons before regeneration of the unit is necessary?

(a) 7,500 gallons

(b) 2,500 gallons

(c) 1,100 gallons

(d) 850 gallons

17. A DIR softener is one that

(a) requires only initiation of the regeneration cyclemanually.

(b) has a timer which initiates every step in theregeneration cycle.

(c) requires initiation of the brining cycle manually.

(d) has a water or hardness sensor which initiatesregeneration.

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18. A flow rate of 6 gpm would produce how much of the pressuredrop under the following conditions? The water travelshorizontally through 80 feet of 3⁄4" steel pipe plus two globevalves, five elbows and one gate valve.

(a) 12.25 psi

(b) 7.48 psi

(c) 8.71 psi

19. If a fully automatic softener were installed to soften a 30 grainwater for a family of six using an estimated 50 gallons each perday, the unit would have to have a minimum daily capacity ofapproximately

(a) 7,000 grains capacity.

(b) 9,000 grains capacity.

(c) 15,000 grains capacity.

20. In installing a pipe from the softener to the drain for dischargeof backwash brining effluents, there should be a gap betweenthe pipe and drain of not less than

(a) 2 inches.

(b) 31⁄2 inches.

(c) 11⁄2 inches.

(d) 1 inch.

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Lesson 1Water and Water Quality

Questionnaire

Fill in your answers to the following questions.

1. Which of the following statements are correct? Place checksbeside them.

(a) Sea water contains about 35,000 mg/L of dissolvedsolids.

(b) About 40 percent of all precipitation evaporates beforeseeping into the soil or becoming surface runoff.

(c) Water evaporates even while it falls to earth asprecipitation.

(d) Surface waters generally have higher mineral contentthan ground waters.

2. Which of the following are effective aids in nature’s efforts topurify water? Check those which apply.

(a) Sunlight

(b) Algae

(c) Oxidation

(d) Chlorides

(e) Erosion

(f) Sulfates

(g) Bacteria

(h) Viruses

(i) Hydroxides

Fill in the blanks in the following:

3. The amount of free carbon dioxide in rainwater generally variesfrom __________ to __________ mg/L.

4. Normally when water reaches the earth, it is ____________________, __________ and relatively soft water.

5. There are a number of factors which help to determine whathappens to water after it falls to earth. Among these are surfacetopography and the degree of the soil’s saturation at the time ofa rainfall. Name two other general factors which also affect thedisposition of rainfall. ___________________ ___________________

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6. Water and carbon dioxide react to produce a weak acid knownas __________ __________.

7. Deep well waters generally show __________ in compositionover a period of time.

(a) no change

(b) slight change

(c) important change

8. H2CO3 + CaCO3 —> Ca(HCO3)2This is the chemical reaction which shows how carbonic acidcombines with insoluble calcium carbonate to produce soluble__________ __________.

9. In nature’s balanced operations, evaporation equals __________.

10. Spring water usually contains rather large amounts of dissolved_____________________.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE11. Snowfall high in the mountains above the

timberline contains a great deal of hardness due to its exposure to rock formations. _____ _____

12. Although springs are popularly believed toprovide sparkling clear water, some of them contain a good deal of turbidity at times. _____ _____

13. A heavy rainfall will wash all impurities out of the air giving it a fresh clean smell. _____ _____

14. As water falls either as rain or snow, it picks upsome hardness minerals even before it reaches the ground. _____ _____

15. Water has its maximum density at a temperature of 39.2°F. _____ _____

continued on next page

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TRUE FALSE16. Water from which the minerals have been

removed by the ion exchange process is known as “distilled water.” _____ _____

17. Water decreases in temperature at about 5°F for each 64 feet of additional depth to a well. _____ _____

18. The United States Environmental ProtectionAgency provides practical standards for water in terms of its potability. _____ _____

19. When water percolates into the ground, it loses some of the impurities it has absorbed in the air. _____ _____

20. Generally, there are quite appreciable differencesin the amounts of hardness minerals to befound between water drawn from a lake and a deep well. _____ _____

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315

Lesson 2Introduction to Water Chemistry

Questionnaire

Fill in your answers to the following questions.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE1. An element is a substance which cannot be

broken into a simpler substance without disrupting the atom. _____ _____

2. Compounds are heterogeneous because they combine the atoms of two or more elements. _____ _____

3. Hydrogen has an atomic weight of 16. Theatomic weights of all other elements are expressed in relation to that of hydrogen. _____ _____

4. While the electron, proton, and neutron are thefundamental particles of matter, there are otherssuch as mesons, positrons, and antiprotons. _____ _____

5. The valence of an element is expressed as asmall whole number with a plus or minus designation. _____ _____

6. An ion is an electrically charged atom or group of atoms. _____ _____

7. Pure water has a molecular weight of 18 asshown by substituting the atomic weights in the formula H2O: (2 × 1) + 16 = 18. _____ _____

8. Isotopes are atoms that are related, that is, theyhave the same atomic number but have different atomic weights. _____ _____

9. The gram, the basic unit of mass of the metricsystem, is meant to be the weight of distilledwater at 4°C in a cube whose edge is one hundredth of a meter. _____ _____

10. The electrons surrounding the atom nucleusbehave in much the same way as do the planets in their travels around the sun. _____ _____

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Complete the following statements with the words that bestsupply the correct meaning.

11. In the magnesium chloride compound, the magnesium atom,in order to achieve stability, must throw off ____________________. At the same time, the chlorine atom is capable ofreceiving only a single magnesium electron, thus requiring twochlorine atoms to make a balanced compound.

12. Water and carbon dioxide react to produce a weak acid knownas __________ acid.

13. The chlorine atom has __________ electrons in its outer shell.The carbon atom has __________.

Choose the word or words which make the following themost accurate statements.

14. The following numbers represent the electrons in the outershells of the atoms of five different elements. On the basis ofthis information alone, indicate (underline) which element youwould expect to be the least active chemically and indicate(circle) the element that would be the most active.

(a) 1 (b) 3 (c) 5 (d) 6 (e) 8

15. Which of the following statements (one or more) is true of theatoms of all the alkaline earths?

(a) Same number of orbits containing electrons

(b) Same number of electrons in the outermost orbit

(c) Same total number of extranuclear electrons

(d) Same net positive charge on the nucleus

(e) Same number of electrons in all orbits or shells lowerthan the outermost one

continued on next page

317

16. Given the correct chemical formula—HCl, NaCl, MgCl2, andAlCl3, one would conclude (one or more) that

(a) chlorine’s valence may be either 1, 2, or 3.

(b) hydrogen and sodium have the same valence.

(c) the valence of aluminum is six times that ofmagnesium.

(d) chlorine has a negative valence.

(e) aluminum has a valence three times that of sodium.

17. When water turns to steam at 212°F, a __________ changeoccurs.

(a) chemical

(b) physical

18. The maximum number of electrons in the “L” shell of an atomis __________.

(a) 2

(b) 8

(c) 18

(d) some other number.

Circle the correct answers.

19. A compound is [ homogeneous / heterogeneous ] and has a(n) [ inconsistent / definite and unvarying ] composition.

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319

Lesson 3Water Chemistry

Questionnaire

Fill in your answers to the following questions.

Complete the following statements with words that bestsupply the correct meaning.

1. Iron, present in water in excess of as little as __________ mg/L,can cause staining of plumbing fixtures.

2. Water is an extremely __________ compound; it contains__________% of oxygen and __________% of hydrogen.

3. The protons and neutrons of an atom are to be found in the__________.

4. Atoms that are related, that is, have the same atomic numberbut have different atomic weights, are called __________.

5. The union of two atoms, one of hydrogen, one of oxygen, hasthe formula (OH–). It performs as though it were a single atom.It is called __________.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE6. Fluorides can be detrimental or beneficial in

water. They may cause a dark brown stain to form on teeth. _____ _____

7. Chlorides seldom are found in natural water supplies. _____ _____

8. pH indicates the degree of intensity of the acidity or alkalinity of water. _____ _____

9. Sodium can be removed from water onlythrough such processes as distillation, deionization, or reverse osmosis. _____ _____

10. Manganese oxidizes much faster than iron. Forthis reason, it is generally easier to remove from water. _____ _____

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TRUE FALSE11. Alkalinity is caused by the presence of

bicarbonates, carbonates, and hydroxides in water. _____ _____

12. Salts are substances containing metallic cations and metallic anions. _____ _____

13. A pH of 2 is 100 times as acid as a pH of 4. _____ _____

14. A bicarbonate ion consists of 3 hydrogen atoms, 1 carbon atom and 1 oxygen atom. _____ _____

15. Surface waters generally contain more carbon dioxide than do well waters. _____ _____

Choose the answer or answers which make the most accu-rate statement for the following:

16. In a broad sense, the term oxidation of an atom or ion refers to

(a) decrease in positive valence.

(b) increase in negative valence.

(c) increase in positive valence.

17. Check the statements which are correct in regard to ions.

(a) In solution many molecules dissociate into two ormore ions which theoretically are free to move orwander around.

(b) An ion differs from an atom or molecule in that anion carries an electric charge.

(c) The terms “cation” and “anion” were developedabout 40 years ago by early pioneers in the waterconditioning industry.

18. A pH of __________ indicates a neutral water with acidity andalkalinity in balance.

(a) 7

(b) 3

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321

19. In the ionization of water, a hydrogen ion separates from themolecule. This hydrogen ion now consists of __________.

(a) one neutron

(b) one proton

(c) two electrons

20. Magnesium sulfate (MgSO4) found in water is __________.

(a) highly insoluble

(b) variable in solubility

(c) extremely soluble

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323

Lesson 4Water Hardness

Questionnaire

Fill in your answers to the following questions.

Choose the word or words which make the most accuratestatement of the following:

1. The hardness minerals found in water are __________ by sodiumions in the ion exchange process.

(a) neutralized

(b) replaced

(c) destroyed

2. The lime-soda ash treatment is not effective for home softeningof water for some of the following reasons:

(a) Difficulty in feeding the lime and soda ash into rawwater

(b) Failure to remove calcium hardness

(c) Produces a bitter taste in the water

(d) Close control necessary during the settling andfiltering processes

3. A water softener with a rated capacity of 20,000 grains shouldbe able to soften __________ gallons of a 20 grain hard waterbefore regeneration is necessary.

(a) 15,000

(b) 6,666

(c) 2,000

(d) 1,000

(e) 400

4. The equivalent weight in grains per gallon of the varioushardness compounds is expressed in terms of __________.

(a) sodium sulfate

(b) magnesium bicarbonate

(c) calcium carbonate

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5. If hard water contains a solution of calcium bicarbonate, it willcontain __________ after it has been ion exchanged softened.

(a) sodium sulfate or potassium sulfate

(b) sodium bicarbonate or potassium sulfate

(c) sodium chloride or potassium sulfate

6. The following are common impurities occurring in water.Match the impurities with the most suitable method ofremoving or counteracting that impurity.

Bacteria FiltrationCa(HCO3)2 DistillationMud ChlorinationNaCl Ion exchange softening

Check whether the following statements are True or False inthe proper column.

TRUE FALSE7. The amount of hardness for water in the

United States ranges from 1 to 350 grains. _____ _____

8. Water does not ionize readily, and it is extremely stable. _____ _____

9. The process of removing all minerals from water is known as cation exchange. _____ _____

10. A fully automatic softener is equipped with atiming device that initiates every step in the regeneration of the unit. _____ _____

11. In all compounds, the sum of the positivecharges equals the number of the negative charges. _____ _____

12. When water contains carbon dioxide, itconverts the insoluble forms of calcium and magnesium into highly soluble bicarbonates. _____ _____

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325

TRUE FALSE13. A molecule is a substance which cannot be

changed chemically into a simpler substance. _____ _____

14. Excessive alkalinity gives water a “soda” taste. It has a drying effect on the skin. _____ _____

15. In the demineralization process, hydroxyl ionsare used in the cation exchanger and hydrogen ions in the anion exchanger. _____ _____

16. No natural water supply is completely free of hardness. _____ _____

17. Rainwater is as soft as zero soft water. _____ _____

Fill in the missing word or phrase in each of the following:

18. When hypothetical combinations are calculated, the ions arecombined in their order of __________ solubility.

19. Calcium and magnesium ions react with soap to produceinsoluble soap __________.

20. What type of ions replace calcium ions in the water softeningexchange process? ___________________________

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327

Lesson 5Iron and Manganese

Questionnaire

Fill in your answers to the following questions.

Choose the answer or answers which make the most accu-rate statements in the sentences below. In some cases, theremay be more than one correct answer.

1. Iron water is a household troublemaker. It can do some of thefollowing things:

(a) Stain plumbing fixtures

(b) Turn coffee a putty gray color

(c) Foul the bed of a water softener

(d) Give water a metallic taste

(e) Discolor washable fabrics

2. Organic (chelated) iron is especially hard to remove from water.If a water has an essentially neutral pH, the most effectivemeans of removing this type of iron is considered to be:

(a) ion exchange softener.

(b) oxidation and filtration, using a powerful oxidantfeed.

(c) chlorination.

(d) sequestration using polyphosphates.

3. A neutral hard water with an iron content may contain the ionsof various hardness and iron compounds, including some of thefollowing:

(a) calcium bicarbonate

(b) magnesium sulfate

(c) ferric hydroxide

(d) magnesium chloride

(e) ferric sulfate

(f) calcium chloride

(g) ferrous hydroxide

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4. The presence of iron bacteria can be indicated in several ways.Choose the methods which apply.

(a) Appearance of slimy red fungus growths in flushtanks

(b) Microscopically

(c) Sudden discharge of extremely turbid water

(d) Sharp increase in the acidity of water

5. Below is a group of statements regarding the treatment of iron-bearing acid water. Corrective equipment and chemicals includea chemical feed pump, a sand filter, chlorine, and soda ash.Place these statements in proper sequence by matching theletters to the numbers.

(a) pH of water rises (1) _____

(b) Iron free water flows to household taps (2) _____

(c) Solutions flow into retention tank (3) _____

(d) Feed pump injects chlorine into water (4) _____

(e) Water flows from pressure tank to sand filter (5) _____

(f) Iron precipitates (6) _____

(g) In retention tank, soda ash and chlorine is in contact with water for a minimum of 20 minutes (7) _____

(h) Feed pump injects soda ash into water (8) _____

(i) Iron-bearing acid water enters the house (9) _____

(j) Insoluble iron precipitate is removed from the water (10) _____

Complete the following statements with the word or wordsthat supply the correct meaning.

6. Inlet and outlet of a retention tank should be so arranged as toassure __________ in the tank.

7. Polyphosphates may be fed into water through use of a smallpositive displacement pump. A __________ ____________________ can also be used for polyphosphate feed.

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329

8. A minimum of __________ ppm of alkalinity and __________ pHis necessary in the water to enable complete precipitation ofiron.

9. Polyphosphates do not remove iron from the water. They__________ and __________ the iron so that the water remainsclear and does not produce __________ __________.

10. Oxidizing filters act to convert ferrous iron into __________ iron.

11. Because dissolved manganese __________ more slowly than iron,it is more difficult to remove from water.

12. Not all iron problems are due to the presence of ferrous andferric compounds as well as iron bacteria. Some iron problemsare due to __________.

13. The USEPA recommends a __________ mg/L limit for iron and a__________ mg/L limit for manganese in its Drinking WaterRegulations.

14. Iron in its ferrous state is __________ and soluble. As ferrichydroxide, it is rust-colored and __________.

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331

Lesson 6Corrosion

Questionnaire

Fill in your answers to the following questions.

Choose the answer or answers which make the most accu-rate statements in the sentences below.

1. One common oxygen scavenger frequently used in industrialefforts for corrosion is:

(a) potassium sulfate.

(b) sodium sulfite.

(c) ferrous hydroxide.

This chemical is not used for corrosion control in the homebecause these scavengers __________ at high temperature.

(d) are too expensive

(e) are not dependable

(f) make water impotable

2. Hardness minerals are __________ when an ion exchangesoftener is used.

(a) replaced

(b) neutralized

(c) destroyed

3. When dissimilar metals are in contact in a solution capable ofcarrying an electric current, certain actions occur. (Check anystatement below that is correct.)

(a) An electric current flows between the two metals.

(b) Convection currents are set up in the water.

(c) A deep etching of both the metals occurs.

(d) One of the metals gradually dissolves.

(e) Actually nothing happens if one of the dissimilarmetals is copper.

332

4. Coating steel surfaces with zinc does the following. (Check thestatements that are correct.)

(a) The coating prevents contact between the corrosivewater and metal steel.

(b) The coating prevents pinhole corrosion fromspreading.

(c) The coating causes zinc oxide to form. This provides aprotective coating when the metal surface is damagedby either mechanical or corrosive action.

(d) The coating can be tailored to any specific corrosionproblem.

Complete the following statements with the proper word orwords:

5. If a water is corrosive and iron-bearing, a neutralizing filter may__________ the iron. As a result, the iron gel formed in the filterbed could well __________ the the unit.

6. In a broad sense, oxidation refers to __________ of positivevalence or __________ in negative valence.

7. The most chemically active metal known is __________; the leastactive is ___________.

8. The use of coatings has become quite popular in combatingcorrosion. These coatings prevent contact between __________and __________. Zinc coatings tend to __________ wheredamaged by either mechanical or corrosive action.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE9. Neutralizing filters are used primarily to correct

corrosive hard water. _____ _____

10. Polyphosphates reduce corrosion of all metals except gold and iron. _____ _____

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333

TRUE FALSE11. Acid waters may cause the deep etching of all

metals except copper and brass. _____ _____

12. Breaks in galvanized coatings tend to “heal” themselves. _____ _____

13. The only drawback to soda ash feed is that itrequires the use of a sand filter following the feeder to remove any iron present in the water. _____ _____

14. Insulating unions cannot be satisfactorily used where joints between metals are too small. _____ _____

15. The presence of electrolytes in water enables it to carry electricity. _____ _____

16. Millions of dollars are spent annually inpreventing iron rusting; this is because iron is the most corrosive of all metals. _____ _____

17. Distilled water carries about the same amount of electric current as soft water. _____ _____

18. Hydrogen gas is a byproduct of corrosion. It may protect against further corrosion. _____ _____

19. Where corrosion is due to conductivity, anincrease in dissolved mineral content increases the rate of corrosion. _____ _____

20. The use of ion exchange softeners does not cause or correct corrosion problems. _____ _____

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335

Lesson 7Total Dissolved Solids, Hydrogen Sulfide,Fluorides, and Other Water Impurities

Questionnaire

Fill in your answers to the following questions.

Choose the answer or answers which make the most accu-rate statements in the sentences below. Note that in severalinstances, more than one answer is correct.

1. The recommended maximum for sulfates is 250 mg/L (USEPADrinking Water Regulations). What is this sulfate content ingrains per gallon expressed as calcium carbonate?

(a) 12.1 grains per gallon

(b) 15.2 grains per gallon

(c) 39.6 grains per gallon

(d) 27.0 grains per gallon

(e) None of the above

2. Which of the following are accurate statements regardingcarbon dioxide?

(a) Carbon dioxide reacts with water to form carbonicacid.

(b) The formula for carbon dioxide gas is C2O3.

(c) Carbon dioxide from decaying vegetation is a sourceof water acidity.

(d) Five percent of the air by volume consists of carbondioxide gas.

(e) Limestone reacts with carbon dioxide in water toform soluble bicarbonates.

3. Atoms which have the same atomic number but differentatomic weights are known as:

(a) riffles.

(b) ions.

(c) isotopes.

(d) molecules.

(e) mesons.

336

4. Hydrogen sulfide gas in solution makes a water unsatisfactoryfor which of the following reasons?

(a) Nonflammable in high concentrations

(b) Tarnishes silverware

(c) Can cause nausea in high concentrations

(d) Obnoxious odor

(e) Promotes corrosion of plumbing systems

5. Under normal conditions, water from which of the belowmentioned sources is likely to contain the least amounts of freeoxygen?

(a) rivers

(b) lakes

(c) deep wells

(d) oceans

(e) RO water

Check whether the following statements are True or False inthe proper column.

TRUE FALSE6. Strongly alkaline waters have an objectionable

“soda” taste. _____ _____

7. Alkalinity may be removed by ion exchange. _____ _____

8. Carbonic acid can be removed from water by passing water through a bed of limestone chips. _____ _____

9. Dissolved sulfates in water can have stronglaxative effects in amounts over 500milligrams per liter. _____ _____

10. Originally the term oxidation signified achemical reaction involving the addition of oxygen to a compound. _____ _____

11. Well waters generally contain more hydrogen sulfide than do surface waters. _____ _____

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TRUE FALSE12. Even in minute quantities (10–20 mg/L),

nitrate-nitrogen is harmful to young babieswhen water containing this contaminant is used for bottle feedings. _____ _____

13. Oxygen starvation caused by nitrate-nitrogen causes a condition called cyanosis. _____ _____

14. Sodium salts produce hard water curd when combined with soap. _____ _____

15. Potassium permanganate is an excellent oxidizing agent for ferrous iron. _____ _____

Complete the following statements with the pertinent wordor words.

16. Soda ash and carbon dioxide react to form _____________________________.

17. List three indications of the presence of alkalinity in water:

(a)

(b)

(c)

18. Silica is found in water in both ____________________ and____________________ forms.

19. Low to moderate concentrations of chlorides and sulfates mayadd what desirable quality to water? ____________________

20. A valence is equal to the number of electrons that an elementloses or gains in the formation of ____________________.

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Lesson 8Taste, Odor, and Turbidity

Questionnaire

Fill in your answers to the following questions.

Choose the one most accurate answer in each of the following:

1. The terms nasturtium, cucumber, grassy, violets, and vegetablein relation to water refer to

(a) color.

(b) taste.

(c) odor.

2. The discharge of coliform bacteria from an individual’s wastematerial may run as high as

(a) 400,000,002,000.

(b) 400,000,000,000.

(c) 400,003,500,000.

3. Nematodes are members of the worm family. Nematodes rangein size from

(a) microscopic forms in water to those visible to thenaked eye in warm-blooded animals.

(b) submicroscopic forms in warm-blooded animals tomicroscopic units in water.

(c) forms visible to the naked eye in water to microscopicforms in warm-blooded animals.

4. Algae are a primitive life form. They obtain their food by

(a) taking food parasitically from other living creaturesin water.

(b) feeding on dead fish.

(c) manufacturing their own food throughphotosynthesis.

5. Chlorine demand is a term common to the chlorinationprocess. It means:

(a) the amount of chlorine fed into the water.

(b) the chlorine still remaining in water after oxidationoccurs.

(c) the amount of chlorine necessary to oxidizeimpurities.

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6. Pathogenic bacteria thrive best in water that has a temperature

(a) anywhere from 0° to about 180°F.

(b) somewhere between 32° and 78°F.

(c) akin to that of warm-blooded animals.

7. Iodine should be used for purification of drinking water only onan emergency basis because

(a) excessive iodine needed for treatment gives water amusty taste.

(b) not enough is presently known about thephysiological effects of iodine-treated water on thehuman system.

(c) sufficient contact time is not normally possible.

8. One of the most effective ways to remove organic (chelated)iron from water with essentially neutral pH is

(a) an ion exchange softener.

(b) oxidation and filtration using a powerful oxidant feed.

(c) sequestration using polyphosphates.

9. Bacteria are sensitive to the pH of a water. They live best in

(a) neutral waters with a pH between 6.5 and 7.5.

(b) acid waters with a pH between 3.5 and 4.5.

(c) alkaline waters with a pH between 8.0 and 9.0.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE10. Photosynthesis is the process whereby plants

manufacture carbohydrates out of water and chlorophyll. _____ _____

11. Algae and fungi belong to the same basic category of plant forms. _____ _____

12. Nematodes are considered objectionablebecause they impart objectionable tastes and odors to water. _____ _____

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TRUE FALSE13. Some protozoa are classed as anaerobic. This

means that they can exist where no free oxygen is available. _____ _____

14. Viruses can survive indefinitely under such conditions as freezing or drying. _____ _____

15. The use of ultraviolet light for disinfection ofwater is not effective where the temperature of the water is increased above 78°F. _____ _____

16. Chlorine is readily available for domestic needs in the form of calcium hypochlorite. _____ _____

17. Superchlorination requires the addition ofmuch more chlorine than usually found in municipal water at the tap. _____ _____

18. Tastes and odors in water can be traced to anumber of factors. Among these are decayingorganic matter, living organisms, and industrial wastes. _____ _____

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Lesson 9How to Interpret Water Analyses

Questionnaire

Fill in your answers to the following questions.

Read over the following water analysis. Then check whetherthe statements are True or False in the proper column.

WATER ANALYSIS

Date Collected 5/3/01Source WellDate Analyzed 5/10/01Appearance When Drawn Clear, odorless, slight iron tastepH 5.7Iron 0.2 mg/LBicarbonate 0.4 gpgSulfate 0.5 gpgChloride 0.6 gpgTotal Anions 1.5 gpgCalcium Hardness 0.4 gpgMagnesium Hardness 0.2 gpgTotal Hardness 0.6 gpgCa(HCO3)2 0.4 gpgMg(HCO3)2 — gpgMgSO4 0.2 gpgNaCl 0.6 gpgNa2SO4 0.3 gpg

TRUE FALSE1. The amount of iron in this sample would cause

some staining of porcelain fixtures in the home. _____ _____

2. It would be correct to say that this is a relatively soft, acid water. _____ _____

3. A neutralizing filter could be used to raise the pH of the water. _____ _____

4. The big advantage to using a neutralizing filteron this water is that it would not increase thehardness of the water as would be the case with some other possible methods of treatment. _____ _____

5. Total hardness of this water expressed in milligrams per liter is 12.6. _____ _____

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TRUE FALSE6. The use of a solution feeder pumping soda ash

is another possible method of raising the pH of this water. _____ _____

7. This soda ash treatment followed by filtrationwill also be effective in removing the iron from the water. _____ _____

8. The use of a polyphosphate feeder wouldprovide an economical method of treating boththe iron and corrosion problems in a plumbing system using galvanized pipe. _____ _____

9. This analysis shows that water contains amountsof sodium (sodium chloride—0.6 gpg; sodiumsulfate—0.3 gpg). This amount of sodium couldlower the operating efficiency of an ion exchange softener installed for treatment of the water. _____ _____

10. Low concentrations of chlorides and sulfates(less than 5 grains) are desirable in a drinking water. _____ _____

11. Manganese-bearing water produces dark brown or black stains on fabrics washed in this water. _____ _____

12. Highly alkaline waters have an objectionable“soda” taste. They also have a tendency towarddrying the skin. While alkalinity can be treatedin several ways, no method is too satisfactory for whole-house use. _____ _____

13. Potassium permanganate alone providesexcellent control for any amount of ferrous iron. _____ _____

14. A water containing 5 gpg of sodium sulfate may have a strong laxative effect on many people. _____ _____

15. Demineralization is the only satisfactory way of removing sodium salts from water. _____ _____

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TRUE FALSE16. The recommended maximum concentration for

sulfates in drinking water according to USEPADrinking Water Regulations is 250 mg/L. Ingrains per gallon expressed as calcium carbonate this would be 20.6 gpg. _____ _____

17. This water is likely to be highly corrosive. _____ _____

18. In order to determine comparable chemicalvalue amounts of calcium chloride andmagnesium hydroxide both would be normally expressed in terms of calcium bicarbonate. _____ _____

19. This water would not be safe to drink. _____ _____

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347

Lesson 10Recommended Installation Procedures

Questionnaire

Complete the following:

1. Study the above diagram carefully. Can you give four reasonswhy this would not be rated as a satisfactory installation?

(a)

(b)

(c)

(d)

2. A variety of materials are used for piping. Two that arecommonly employed for water softener installations include:

(a)

(b)

3. It is good practice to bypass outside sillcocks when making awater softening installation for two reasons:

(a)

(b)

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Fill in your answers to the following questions.

4. Polyphosphates can be fed into water in two ways with the aid of:

(a)

(b)

Choose the most accurate answer for each of the following:

5. If water were to rise from the basement softener to the showerin a second floor bathroom, a distance of 18 feet, there wouldbe a pressure drop due to the increase in height alone of

(a) 23.12 psi.

(b) 7.79 psi.

(c) 11.44 psi.

6. A flow rate of 4 gpm would produce how much of a pressuredrop under the following conditions? The water travels through64 feet of 3⁄4" steel pipe plus two globe valves, seven elbows, oneside outlet tee, two unions and will travel up a height of 11 feet.

(a) 7.84 psi

(b) 8.66 psi

(c) 3.90 psi

Check whether the following statements are True or False inthe proper column.

TRUE FALSE7. With many modern softeners, the brine tank

may be located some distance from the softener tank. _____ _____

8. Not only the pressure drop rating of thesoftener but total length of pipe, fittings, andheight of the water head must be figured inorder to determine whether there will be adequate pressure at upper floor taps. _____ _____

9. Installation of water conditioning equipmentdoes not require the services of fully trained installers. _____ _____

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349

TRUE FALSE10. The rinsing of brine through the household

pipes is not acceptable practice. _____ _____

11. When water flows through a standard elbow,there is a pressure loss equivalent to the flow through 2.1 feet of 3⁄4" pipe. _____ _____

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351

Lesson 11Common Questions and Answers on

Water Conditioning

Questionnaire

Fill in your answers to the following questions.

Choose the one most correct answer in each of the following:

1. The use of an ion exchange softener for a humidifier is stronglyrecommended because

(a) it reduces the total mineral content of the water.

(b) it prevents formation of hard water scale.

(c) it eliminates the need to periodically flush the unit.

2. Water treatment for improving the quality of ice cubes is

(a) an iron removal filter.

(b) a water softener.

(c) a chlorinator.

(d) a reverse osmosis or distillation system.

3. The human body gains most of the minerals necessary to goodhealth from

(a) food.

(b) water.

(c) other sources.

4. The flow of softened water (regardless of the hardness of theraw water) to outside sillcocks is not recommended because

(a) any sodium content in the water would be harmful tothe grass.

(b) there would be considerable waste of soft water.

(c) the lack of the calcium and magnesium ions in thewater could retard healthy growth of grass andflowers.

5. The use of soft water in a tropical fish acquarium is

(a) not acceptable.

(b) especially ideal.

(c) normally satisfactory.

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6. A septic tank with a 300 gallon liquid capacity would containapproximately how many pounds of water?

(a) 2,000 lbs.

(b) 2,500 lbs.

(c) 3,000 lbs.

7. Based upon a salt consumption of 3,000 grains of hardnessexchange per pound of salt used (.333 lb/Kgr.) a softener thathas removed 30,000 grains of hardness during the service cyclewould release how many calcium carbonate equivalent poundsof calcium, magnesium, and sodium salts in the regenerationprocess?

(a) approximately 8.6 lbs.

(b) approximately 10 lbs.

(c) approximately 12 lbs.

Check whether the following statements are True or False inthe proper column.

TRUE FALSE8. The use of a polyphosphate feeder is one

economical way for restaurants to minimize clouding of water in the making of ice cubes. _____ _____

9. In making a change from hard to fully softwater in an aquarium, it is a good practice to replace only one-fourth of the water at a time. _____ _____

10. Total mineral content of water should not exceed 10 gpg for production of top quality ice. _____ _____

11. It is important to remember that the softeningof any supply does not reduce total minerals from the water. _____ _____

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Prelesson Questionnaire Answers

Prelesson Questionnaire, Lesson 1

1. a; 2. b, e; 3. a; 4. a; 5. e; 6. solvent; 7. deionization or demineraliza-tion; 8. carbonic acid; 9. carbonate; 10. 70; 11. T; 12. T; 13. F; 14. T; 15. T; 16. T; 17. T; 18. F; 19. T; 20. F

Prelesson Questionnaire, Lesson 2

1. ion; 2. physical, chemical; 3. 18 grams per mole; 4. nucleus; 5. inert;6. F; 7. T; 8. T; 9. F; 10. T; 11. T; 12. F; 13. T; 14. a; 15. a; 16. c; 17. d; 18 b. 19. a; 20. c, e

Prelesson Questionnaire, Lesson 3

1. True statements about water: Water is an extremely stable com-pound. Water does not ionize readily. Water contains 88.81% oxygen.The chemical formula for water is H2O. Inaccurate statements aboutwater: Water is an excellent oxidizing agent. When water ionizes, itforms hydrogen and oxygen ions. At a temperature of 3600°F, onlyabout 22% of water molecules dissociate or separate.; 2. T; 3. F; 4. F; 5. T; 6. F; 7. T; 8. T; 9. T; 10. T; 11. a, c, e; 12. a

Prelesson Questionnaire, Lesson 4

1. c; 2. a, c, d; 3. a, d; 4. b, c; 5. a, b, c; 6. b; 7. d; 8. c; 9. b; 10. 50,000,3,333; 11. 5, 3, 8, 6, 1, 9, 2, 4, 7; 12. F; 13. T; 14. F; 15. T; 16. T

Prelesson Questionnaire, Lesson 5

1. F; 2. T; 3. F; 4. T; 5. F; 6. T; 7. T; 8. T; 9. T; 10. T; 11. T; 12. T; 13. T; 14. T; 15. T; 16. T; 17. T; 18. F; 19. T; 20. T; 21. T; 22. F; 23. T; 24. T

Prelesson Questionnaire, Lesson 6

1. dissolved oxygen content; pH or dissolved CO2 or carbonic acid con-tent; electrical conductivity or total dissolved solids content; temper-ature; aggressiveness caused by absence of dissolved substances; 2. electrolytes; 3. low pH, low cost, little maintenance; 4. dissolvedgases such as oxygen, carbon dioxide, and hydrogen sulfide; 5. cause,cure; 6. F; 7. T; 8. T; 9. F; 10. F; 11. F; 12. T; 13. T; 14. T; 15. F; 16. T; 17. c; 18. c; 19. d; 20. b

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Prelesson Questionnaire, Lesson 7

1. carbonic acid, weak; 2. one, staining or molting of teeth, osteopor-rosis, or brittleness of bones; 3. 250; 4. elemental sulfur; 5. reactive (dis-solved), inert (nonreactive), undissolved, or colloidal; 6. c; 7. b; 8. b; 9.b, c, e; 10. a, b, e; 11. T; 12. F; 13. T; 14. T; 15. T; 16. F; 17. T; 18. T; 19.T; 20. F

Prelesson Questionnaire, Lesson 8

1. giardiasis, cryptosporidiosis, amoebic dysentery, typhoid fever,cholera, gastrointestinal illness; 2. coliform bacteria or fecal coliformbacteria or E. coli; 3. turbidity, hardness, iron, color; 4. odor; 5. chlo-rine taste; 6. F; 7. T; 8. F; 9. T; 10. F; 11. T; 12. T; 13. F; 14. T; 15. T; 16.c; 17. c; 18. b; 19. a; 20. c

Prelesson Questionnaire, Lesson 9

1. immunity; 2. chemical, physical; 3. zero, 95; 4. appearance; 5. �171.1�;

6. T; 7. T; 8. F; 9. T; 10. T; 11. F; 12. T; 13. T; 14. T; 15. F; 16. c; 17. b; 18.b; 19. c; 20. a

Prelesson Questionnaire, Lesson 10

1. T; 2. T; 3. F; 4. T; 5. F; 6. T; 7. T; 8. F; 9. T; 10. T; 11. T; 12. T; 13. T; 14. F; 15. F; 16. c; 17. d; 18. c; 19. b; 20. d

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Questionnaire Answers

Questionnaire, Lesson 1

1. a, c; 2. a, b, c, g; 3. 2, 6; 4. slightly acid, corrosive; 5. surface vegeta-tion, porosity of soil, atmospheric condition; 6. carbonic acid; 7. b; 8. calcium bicarbonate; 9. precipitation; 10. solids or mineral matter;11. F; 12. T; 13. F; 14. T; 15. T; 16. F; 17. F; 18. T; 19. T; 20. T

Questionnaire, Lesson 2

1. T; 2. F; 3. F; 4. T; 5. T; 6. T; 7. T; 8. T; 9. T; 10. T; 11. two electrons; 12. carbonic; 13. 7, 4; 14. e (least active), a (most active); 15. b; 16. b, d,e; 17. b; 18. b; 19. homogeneous; definite and unvarying

Questionnaire, Lesson 3

1. 0.3; 2. stable, 88.81, 11.19; 3. nucleus; 4. isotopes; 5. hydroxide orhydroxyl ion; 6. T; 7. F; 8. T; 9. T; 10. F; 11. T; 12. F; 13. T; 14. F; 15. T;16. c; 17. a, b; 18. a; 19. b; 20. c

Questionnaire, Lesson 4

1. b; 2. a, d; 3. d; 4. c; 5. b; 6. bacteria-chlorination, Ca(HCO3)2-ionexchange softening, mud-filtration, NaCl distillation; 7. T; 8. T; 9. F;10. T; 11. T; 12. T; 13. F; 14. T; 15. F; 16. T; 17. F; 18. increasing; 19. curd; 20. sodium or potassium

Questionnaire, Lesson 5

1. a, c, d, e; 2. b; 3. a, b, c, d, e, f, g; 4. a, b, c; 5. i, h, a, d, c, g, f, e, j, b;6. mixing and contact time; 7. cartridge or pot-type feeder; 8. 100, 7.0;9. stabilize and sequester, iron stains; 10. ferric; 11. oxidizes; 12. cor-rosion; 13. 0.3, 0.05; 14. colorless, insoluble

Questionnaire, Lesson 6

1. b, f; 2. a; 3. a, d; 4. a, b, c; 5. precipitate, clog; 6. increase, decrease; 7. potassium, gold; 8. water, metal, heal; 9. F; 10. F; 11. F; 12. T; 13. F;14. T; 15. T; 16. F; 17. F; 18. T; 19. T; 20. T

357

Questionnaire, Lesson 7

1. b; 2. a, c, e; 3. c; 4. b, c, d, e; 5. c; 6. T; 7. T; 8. T; 9. T; 10. T; 11. T; 12. T; 13. T; 14. F; 15. T; 16. sodium bicarbonate; 17. soda taste, dryingof skin, increased pH, turns red litmus paper blue; 18. dissolved (reac-tive), inert (nonreactive), undissolved, or colloidal; 19. palatability; 20. ions or compounds

Questionnaire, Lesson 8

1. c; 2. b; 3. a; 4. c; 5. c; 6. c; 7. b; 8. b; 9. a; 10. F; 11. F; 12. T; 13. T; 14. T; 15. F; 16. T; 17. T; 18. T

Questionnaire, Lesson 9

1. F; 2. T; 3. T; 4. F; 5. F; 6. T; 7. F; 8. T; 9. F; 10. T; 11. T; 12. T; 13. F; 14. F; 15. F; 16. F; 17. T; 18. F; 19. F

Questionnaire, Lesson 10

1. hot side only water softening, i.e., cold water not softened, whichwill diminish softening benefits; direct connection to sewer drain (noair gap); no bypass arrangement; no inlet and outlet shut-off valves;no insulating unions and electrical strap around the softener valve in-stalled; 2. rigid copper tubing; plastic pipe; 3. expensive to use softenedwater for large volume lawn watering; sodium in softened water maybuild up in lawn soil and reduce soil structure; 4. chemical solutionfeeder (e.g., with a positive displacement pump), and cartridge or pot-type feeder for use with crystalline granules; 5. b; 6. b; 7. T; 8. T; 9. F;10. T; 11. T

Questionnaire, Lesson 11

1. b; 2. d; 3. a; 4. b; 5. c; 6. b; 7. a; 8. T; 9. T; 10. T; 11. T

358

359

Index

360

361

362

363

364

365

366