Complexometry: Determination Hardness of Water

I. Experiment Title : Complexometry and Its

Application

II. Experiment Date : December 19th, 2014 At 07.00

III. End of the Experiment : December 19th, 2014 At 11.30

IV. Experiment Purpose : 1. Making and

determining(standardization) of Na-

EDTA solution

2. Determine the hardness of water Of

PDAM water

V. Basic Theory

A. Complexometry

Complexometric titration (sometimes chelatometry) is

a form of volumetric analysis in which the formation of a

colored complex is used to indicate the end point of a

titration. Complexometric titrations are particularly

useful for the determination of a mixture of different

metal ions in solution. An indicator capable of producing

an unambiguous color change is usually used to detect the

end-point of the titration.

A complex is a molecule or

ion for-med by the reaction of

two or more ions or molecules

capable of independent exis-

tence. The most important

complexation reactions from an

analytical point of view are

those between a metal ion in

solution and a complexingGroup IV |Complexometry and Its Application 1

http://en.wikipedia.org/wiki/Complexometric_indicator

http://en.wikipedia.org/wiki/Volumetric_analysis

agent. A metal atom can usu-

ally form a bond with one or

more donor atoms which have at

least one unshared pair

of electrons. The number of donor atoms which bond with

a given atom depends on the number of electron pairs that

the metal ion can accept, in other words, the

coordination number of the metal ion. Complexing agents,

or ligands, which can provide more than one pair of

electrons (multidentate ligands), are also called chelating

agents.

Complexometric titrations are particularly useful for

determination of a mixture of different metal ions in

solution. An indicator with a marked color change is

usually used to detect the end-point of the titration.

In a complexometric titration, a solution containing

the free metal ion of interest is titrated with a

solution of chelating agent until all of the metal ions

are completely complexed. The endpoint is usually

measured with an indicator ligand that forms a colored

complex with the free metal ion.

Chelating Agent: EDTA

The most important chelating agent in analytical

chemistry is ethelyenedia-minetetraacetic acid

(EDTA). The tetrabasic form of this acid forms

complexes with virtually all metal ions. EDTA is a

hexadentate ligand; each of the acid oxy-gens and each

of the amine nitrogens can donate one electron pair.

Group IV |Complexometry and Its Application 2

The metal ion is usually held in a one-to-one complex

with EDTA. The complexes have four or five membered

rings, contributing significantly to their stability.

Unfortunately, EDTA cannot be easily used as a primary

standard. It is available in pure form, but must be

dried at 80°C for several days to obtain the precise

composition of the dihydrate. In any case,

standardization of EDTA titrant against a solution of

the metal ion to be determined helps to eliminate any

errors in endpoint selection.

It is important to realize that the electron pairs

of the carboxylic acid groups of EDTA are only

available to the metal ion when the acid is

dissociated. This means that the effectiveness of the

complexing agent is strongly affected by pH. At low pH

EDTA will be in the acid form and will not be an

effective complexing agent. Additionally, many metal

ions form complexes with hydroxide ions. Hydroxide

ions compete with the chelating agent for coordination

sites in the me-tal ion. Therefore, the effectiveness

of the complexing agent will also be reduced at high


pKa-2 =

0.0

pKa-1 =

1.5

pKa1 =

2.0

pH. For a given chelating agent and metal ion, there

will be an optimum pH for the titration which will

depend on the pKa values for the chelating agent and

the formation constants for the metal-hydroxide

complexes.

Complex Titration with EDTA

EDTA, ethylenediaminetetraacetic acid, has four

carboxyl groups and two amine groups that can act as

electron pair donors, or Lewis bases. The ability of

EDTA to potentially donate its six lone pairs of

electrons for the formation of co-ordinate covalent

bonds to metal cations makes EDTA a hexadentate

ligand. However, in practice EDTA is usually only

partially ionized, and thus forms fe-wer than six

coordinate covalent bonds with metal cations.

Disodium EDTA is commonly used to standardize

aqueous solutions of transition metal cations.

Disodium EDTA (often written as Na2H2Y) only forms four

coordinate covalent bonds to metal cations at pH

values ≤ 12. In this pH range, the amine groups remain

protonated and thus unable to donate electrons to the

formation of coordinate covalent bonds. Note that the

shorthand form Na4-xHxY can be used to represent any

species of EDTA, with x designating the number of

acidic protons bonded to the EDTA molecule.


EDTA forms an octahedral complex with most 2+

metal cations, M2+, in aqueous solution. The main

reason that EDTA is used so extensively in the

standardization of metal cation solutions is that the

formation constant for most metal cation-EDTA

complexes is very high, meaning that the equilibrium

for the reaction:

M2+ + H4Y → MH2Y + 2H+

lies far to the right. Carrying out the reaction in a

basic buffer solution removes H+ as it is formed, which

also favors the formation of the EDTA-metal cation

complex reaction product. For most purposes it can be

considered that the for-mation of the metal cation-

EDTA complex goes to completion, and this is chiefly

why EDTA is used in titrations / standardizations of

this type.

The equilibrium involved in EDTA titration:

1.The stability of complex formed: The greater the

stability constant for complex formed, larger the

charge in free metal concentration (pM) at

equivalent point and more clear would be the end

point.

2.The number of steps involved in complex formation:

Fewer the number of steps required in the formation

of complex, greater would be the break in titration

curve at equivalent point and clear would be the end

point.


3.Effect of pH: During a complexometric titration, the

pH must be constant by use of a buffer solution.

Control of pH is important since the H+ ion plays an

important role in chelation. Most ligands are basic

and bind to H+ ions throughout a wide range of pH.

Some of these H+ ions are frequently displaced from

the ligands (chelating agents) by the metal during

chelate formation.

Indicators

To carry out metal cation titrations using EDTA,

it is almost always neces-sary to use a complexometric

indicator to determine when the end point has been

reached. Common indicators are organic dyes such as

Fast Sulphon Black, Erio-chrome Black T, Eriochrome

Red B, Patton Reeder, or Murexide. Color change shows

that the indicator has been displaced (usually by

EDTA) from the metal cations in solution when the end

point has been reached. Thus, the free indicator

(rather than the metal complex) serves as the endpoint

indicator.

EBT indicators

The structure of Eriochrome Black T are shown

below


OH

N

NO2

-

O3SN

OH

Chelate metal formed with Eriochrome Black T

molecule with the disappe-arance of hydrogen ions

from phenolate –OH and formation bonding bet-ween

metal ion and oxygen atoms and also azo-group.

This indicator forming red wine stable complex 1:1

with some cations, such as Mg2+, Ca2+, Zn2+, and Ni2+.

Most of the EDTA titration occur in buffer pH 8-

10, a range where the dominant form of EBT is HIn2-

blue colored.

Calmagnite

Most metallochromic indicators also are weak

acids. One consequence of this is that the

conditional formation constant for the metal–

indicator com-plex depends on the titrand’s pH.

This provides some control over an indi-cator’s

titration error because we can adjust the strength

of a metal–indica-tor complex by adjusted the pH

at which we carry out the titration. Unfor-

tunately, because the indicator is a weak acid,

the color of the uncomplexed indicator also

changes with pH.

Figure beside, for example,

shows the color of the

indicator calmagite as a

function of pH and pMg,Group IV |Complexometry and Its Application 7

(Underwood, 1998)solution at pH 10. It turns red when Ca2+ ions

http://en.wikipedia.org/wiki/Buffered_solution

where H2In–, HIn2–, and

In3– are different forms of

the uncomplexed indicator,

and MgIn– is the Mg2+–

calmagite complex. Because

the color of calmagite’s

metal–indicator com-

plex is red, its use as a metallochromic indicator

has a practical pH range of approximately 8.5–11

where the uncomplexed indicator, HIn2–, has a blue

color.

Some other indicator known can be used for some

cations:

Application of EDTA Titration

EDTA titration mostly succeeds in every cations.

There are some procedure which used in this kind of

analytic gravimetric:


Direct Titration

It is the simplest and the most convenient method

used in chelometry. In this method, the standard

chelon solution is added to the metal ion solution

until the end point is detected. This method is

analogous to simple acid-base titrations. E.g.-

calcium gluconate injection, calcium lactate

tablets and compound sodium lactate injection for

the assay of calcium chloride (CaCl2.6H2O).

Limitations -slow complexation reaction -

Interference due to presence of other ions

Back Titration

In this method, excess of a standard EDTA solution

is added to the metal solution, which is to be

analyzed, and the excess is back titrated with a

standard solution of a second metal ion. E.g. -

Determination of Mn. This metal cannot be directly

titrated with EDTA because of precipitation of

Mn(OH)2. An excess of known volume of EDTA is

added to an acidic solution of Mn salt and then

ammonia buffer is used to adjust the pH to 10 and

the excess EDTA remaining after chelation, is back

titrated with a standard Zn solution kept in

burette using Eriochrome blackT as indicator. This

method is analogous to back titration method in

acidimetry. e.g.- ZnO

Replacement Titration


In this method the metal, which is to be analyzed,

displaces quantitatively the metal from the

complex. When direct or back titrations do not

give sharp end points, the metal may be determined

by the displacement of an equivalent amount of Mg

or Zn from a less stable EDTA complex.

Mn+2 + Mg EDTA-2 → Mg+2 + Mn EDTA-

2

Mn displaces Mg from Mn EDTA solution. The freed

Mg metal is then di-rectly titrated with a

standard EDTA solution. In this method, excess

quan-tity of Mg EDTA chelate is added to Mn

solution. Mn quantitatively displa-ces Mg from Mg

EDTA chelate. This displacement takes place

because Mn forms a more stable complex with EDTA.

By this method Ca, Pb, Hg may be determined using

Eriochrome blackT indicator.

Indirect Titration

This is also known as Alkalimetric titration. It

is used for the determination of ions such as

anions, which do not react with EDTA chelate.

Protons from disodium EDTA are displaced by a

heavy metal and titrated with sodium alkali.

Mn+ + H2X-2 → MX (n-4) +

2H+

Some important elements which could be determined

by complexometric titration are as follows: i) Group IV |Complexometry and Its Application 10

Direct Titration : Analysis of Cu, Mn, Ca, Ba, Br,

Zn, Cd, Hg, Al, Thallium, Sn, Pb, Bi, Vanadium,

Cr, Mo, Gallium, Fe, Co, Ni, and Pd.

B. Application of Complexometry: Determination the hardness

of water

Complexometric titration is an efficient method for

determining the level of hardness of water. Caused by

accumulation of mineral ions, pH of water is increased.

The Kf during the titration of hard water is reduced

because of the reduced amount of EDTA added.

Softening of hard water is done by altering the pH of

the water reducing the concentration of the metal ions

present.Could be performed in two phases: Basic pH for

ions with high Kf e.g. Ca2+ and Mg2+

Hard water is water that has high mineral content (in

contrast with "soft water"). Hard water is formed when

water percolates through deposits of calcium and magne-

sium-containing minerals such as limestone, chalk and

dolomite. Group IV |Complexometry and Its Application 11

Figure: End point for the titration of hardness with EDTA using EBT as anindicator; the indicator is: (a) red WINEprior to the end point due to thepresence of the Ca2+–indicator complex; (b) grey-blue at the titration’s end

Hard drinking water is generally not harmful to one's

health, but can pose seri-ous problems in industrial

settings, where water hardness is monitored to avoid

costly breakdowns in boilers, cooling towers, and other

equipment that handles water. In do-mestic settings, hard

water is often indicated by a lack of suds formation when

soap is agitated in water, and by the formation of lime

scale in kettles and water heaters. Wherever water

hardness is a concern, water softening is commonly used

to reduce hard water's adverse effects.

Minerals that cause hard water have a wide impact on

households. Hard water interferes with almost every

cleaning task from laundering and dishwashing to bathing

and personal grooming. Clothes laundered in hard water

may look dingy and feel harsh and scratchy. Dishes and

glasses washed in hard water may become spotted as they

dry. Hard water may cause a film on glass shower doors,

shower walls, and bath-tubs. Hair washed in hard water

may feel sticky and look dull.

Calcium

Calcium occurs in water naturally. Seawater

contains approximately 400 ppm calcium. One of the

main reasons for the abundance of calcium in water is

its natural occurrence in the earth's crust. Calcium

is also a constituent of coral. Rivers generally

contain 1-2 ppm calcium, but in lime areas rivers may

contains calcium concentrations as high as 100 ppm.

Examples of calcium concentrations in water organisms:


seaweed luctuca 800-6500 ppm (moist mass), oysters

appro-ximately 1500 ppm (dry mass).

In a watery solution calcium is mainly present as

Ca2+ (aq), but it may also occur as CaOH+ (aq) or

Ca(OH)2 (aq), or as CaSO4 in seawater. Calcium is an

important determinant of water harness, and it also

functions as a pH stabilizer, because of its buffering

qualities. Calcium also gives water a better taste.

Ca (s) + 2H2O (g) -> Ca(OH)2 (aq) + H2 (g)

This reaction forms calcium hydroxide that

dissolves in water as a soda, and hydrogen gas.

Calcium is a dietary mineral that is present in

the human body in amounts of about 1.2 kg. No other

element is more abundant in the body. Calcium phos-

phate is a supporting substance, and it causes bone

and tooth growth, together with vitamin D. Calcium is

also present in muscle tissue and in the blood. It is

required for cell membrane development and cell

division, and it is partially res-ponsible for muscle

contractions and blood clotting. Calcium regulates

membrane activity, it assists nerve impulse transfer

and hormone release, stabilizes the pH of the body,

and is an essential part of conception. In order to

stimulate these body functions a daily intake of about

1000 mg of calcium is recommended for adults. This may

be achieved by consuming dairy, grains and green

vegetables.


Calcium carbonate works as a stomach acid remedy

and may be applied to resolve digestive failure.

Calcium lactate may aid the body during periods of

calcium deficiency, and calcium chloride is a

diuretic. Hard water may assist in strengthening bones

and teeth because of its high calcium concentration.

It may also decrease the risk of heart conditions.

Drinking water hardness must be above 8.4 odH. Calcium

carbonate has a positive effect on lead water pipes,

because it forms a protective lead(II)carbonate

coating. This prevents lead from dissolving in

drinking water, and thereby prevents it from entering

the human body.

Calcium and magnesium ions present as sulfates,

chlorides, carbonates, and bicarbonates cause water to

be hard. Water chemists measure water impurities in

parts per million (ppm), but water hardness is often

expressed in grains of hardness per gallon of water

(gpg). The two systems can be converted

mathematically. Table 1 gives common classifications

for hard water with values listed in both parts per

million and grains per gallon. One grain of hardness

is the amount of calcium and magnesium equal in weight

to a kernel of wheat.

Table 1. Hard Water Classifications.


VI. Tools and Materials

A. Tools B. Materials1. Ipi vitamin bottle

2. Colored bottle

3. Measuring flask 100 mL

4. Burette

5. Erlenmeyer 250 mL

6. Pipettes

7. Volumetric pipette

8. Pikno mass

9. Measuring glass

1. Distilled water

2. Seawater

3. AgNO3

4. NaCl p.a

5. K2CrO4

6. K2CrO4 5%

VII. Flow Chart

A. Making of Na-EDTA ±0,01 M Solution

B. Determining (standardization) of Na-EDTA ±0,01 M with CaCl2 as standard solution


± 4 g of Na-EDTA

± 0,1 g of MgCl2.H2O

weighed weighed

Entered to volumetric flask 400 mlDilute with distilled waterMoved into bottle and dilute until 1 L

Na-EDTA ± 0,01 M solution

± 0,0811 g of CaCO3 p.a

Weighed accurately Moved to volumetric flask 100 ml using waterAdded HCl 6 M drop by drop until the gas stopped (the bubble disappear)Dilute with water until boundary lineShaked well

CaCl2 ± 0,01 M solution

C. Application of Complexometry: Hardness water of PDAM water


Weighed accurately Moved to volumetric flask 100 ml using waterAdded HCl 6 M drop by drop until the gas stopped (the bubble disappear)Dilute with water until boundary lineShaked well

CaCl2 ± 0,01 M solution


25 ml of CaCl2 ±0,01 M solutionEntered to volumetric

flask 300 mlAdded 5 ml of buffer solution (pH:10)Added 3 drops of EBT indicator

Entered to burette

Titrated

Color changes from red wine to

blueRead and write down the volume of Na-EDTACalculate the concentration of Na-EDTARepeated 3 timesAverage concentration of

Na-EDTA


10 ml of water sample

(PDAM)Entered to volumetric flask 300 mlAdded 5 ml of buffer solution (pH:10)Added 3 drops of EBT indicator

Entered to burette

Titrated


blueRead and write down the volume of Na-EDTACalculate the Hardness water of sampleRepeated 3 timesTotal hardness water of CaCO3

salt/liter


Total hardness water of CaCO3 salt/liter

VIII. Result Of The Experiment

Experiment ProcedureResult of The Experiment Assumption/

ReactionConclusion

Before AfterA. Making of Na-EDTA ±0,01 M Solution

Na-EDTA : colorless solution

Distilled water: colorless solution


The average

concentratio

n of Na-EDTA

is 0,0123 M

and total

hardness of

water of

PDAM water

from

Ketintang

Wiyata

region is189

ppm

B. Determining (standardization) of Na-EDTA ±0,01 M with CaCl2 as standard solution CaCO3 p.a:

white powder


CaCO3 p.a+HCl 6M ⟶ colorless solution

+ distilled water⟶ colorless solution


± 4 g of Na-EDTA

± 0,1 g of MgCl2.H2O

weighed weighed

Entered to volumetric flask 400 mlDilute with distilled waterMoved into bottle and dilute until 1 L


± 0,0811 g of CaCO3 p.a

Weighed accurately Moved to volumetric flask 100 ml using waterAdded HCl 6 M drop by drop until the gas stopped (the bubble disappear)Dilute with water until boundary lineShaked wellCaCl2 ± 0,01 M

solution



Buffere solution: colorless solution

EBT: red wine solution

CaCl2: colorless solution

CaCl2+ buffer solution ⟶ colorless solution

+3 drops of EBT ⟶ red wine solution

+ Na-EDTA⟶ bluesolution

V Na-EDTAV1= 16,3 mlV2= 16,5 mlV3= 16,4 ml

Ca2++Y4- ⟶ CaY2--


C.Applicaton of Complexometry: Hardness water of PDAM water


Buffere solution: colorless solution

EBT: red wine solution

Sample:

Sample + buffer solution ⟶ colorless solution

+3 drops of EBT ⟶ red wine solution

+ Na-EDTA⟶ bluesolution

V Na-EDTA

Ca2++Y4- ⟶ CaY2--



10 ml of water sample

(PDAM)Entered to volumetric flask 300 mlAdded 5 ml of buffer solution (pH:10)Added 3 drops of EBT indicator

Entered to burette

Titrated


blueRead and write down the volume of Na-EDTACalculate the Hardness water of sampleRepeated 3 times

Total hardness water of CaCO3 salt/liter

colorless

solution

V1= 4 mlV2= 3,9 mlV3= 3,8 ml



10 drops of KMNO4

Solution B

+ H2O until

10 drops of H2C2O4

Solution A

+ H2O until

IX. Analysis and Explanation:

Complexometric is a form of volumetric analysis in which

the formation of a co-lored complex is used to indicate the

end point of a titration. Complexometric titrations are

particularly useful for the determination of a mixture of

different metal ions in solution.

An indicator capable of producing an unambiguous color

change is usually used to detect the end-point of the

titration. In this experiment we will detect Ca2+ in

formation of Complex with Na-EDTA. This experiment we use

CaCl2 as a standard solution, because it not primary

standard so we must make standard solution from this matter.

A. Determining (standardization) of Na-EDTA ±0,01 M with CaCl2 as standard solution

We use CaCO3 pure analyte to make standard solution

CaCl2. Weight accurately 0.0811 gram of CaCO3 p.a (white

powder). Moved to volumetric flask 100 ml using water but

not until boundary line, after that add HCl 6M (colourless

solution) drop to drop until the gases is stop.

CaCO3(s) + 2HCl(aq) →CaCl2(aq) + H2O(l) + CO2(g)

Before added HCl the solution is turbid but after added

HCl the solution becomes colorless.The function of adding

HCl is to making loosen carbon dioxide, because from this

reaction carbon dioxide which can be seen in the bubble

form, calcium chloride, and water will be formed. If the

bubble has completely gone, it means the CaCl2 has been

formed and it can be used to standardization as the analyte.


http://en.wikipedia.org/wiki/Complexometric_indicator

http://en.wikipedia.org/wiki/Volumetric_analysis

This Titration used 25 ml of CaCl2 solution entered to

Erlenmeyer. Added 5 ml of buffer solution pH=10, the

solution still colorless. The function of the buffer

solution pH= 10 addition because in pH 8-10 is a range where

dominant form of EBT is in form blue HIn2- and this titration

of CaCl2 and Na-EDTA is in a basic condition during

titration. Buffer solution will defend the pH because this

titration needs the constant pH to make titration process

stable. After that, added 3 drops of EBT indicators. The

solution changes color from colorless to red wine. It is

suitable with the theory because EBT has red wine in color

(shown in picture below). The reason why using this

indicator is because this indicator forming red wine stable

complex 1:1 with some cations such as: Mg2+, Ca2+,

Zn2+, and Ni2+. Since this

titration using CaCO3 as the

analyte which contain ion of

Ca2+. So, EBT is suitable to be

used in this experiment

The reaction is:

Ca2+ + In3- ⇄ CaIn-

Metal ion indicator

indicator-metal (colorless) (blue)

complex(wine-red)


Based on the theory: EBT is blue in a buffered solution at pH 10. It turns red when Ca2+ ions

Based on the expe-riment: the addition of EBT changes the solu-tion to red wine solu-tion


After that, titrated the CaCl2

solution with Na-EDTA solution

until the color changes to blue.

It suitable with the theory

where EBT in a buffered solution

pH=10 is blue (shown in the

figure beside).The using of EDTA

as a titrant because EDTA

forming stable complex, soluble

in water forming complex 1:1

with metal ions.

The reaction is:

Ca2+ (aq) + EDTA4- (aq) → Ca(EDTA)2- (aq)

The experiment is repeated of three times. The volume

of Na-EDTA which require to turns the color from red wine

to blue are 16.3 ml, 16.5 ml, and 16.4 ml. after that,

calculate the concentration of Na-EDTA using equality :

Molek Na-EDTA=Molek CaCO3.

The concentration of Na-EDTA in the first trial is 0.0122

M, in the second trial is 0.0124 M and the third is 0.0123

M. So, the average concentration of Na-EDTA is 0,01M



Based on the expe-riment: indicator EBT is blue in a buffered solution at pH 10.




B. Application of Complexometry: Hardness water of PDAM water

This application is determine the hardness of water.

Here we are using sample of PDAM water which we took from

Ketintang Wiyata region. The PDAM water is used as the

analyte and EDTA as a titer. First, enter 10 ml of sample

to Erlenmeyer. Next, adding 2 ml of buffer solution pH=10.

The function of the buffer solution pH= 10 addition

because in pH 8-10 is a range where dominant form of EBT

is in form blue HIn2- and this titration of CaCl2 and Na-

EDTA is in a basic condition during titration. Buffer

solution will defend the pH because this titration needs

the constant pH to make titration process stable. After

that add 3 drops of EBT. The solution changes color from

colorless to red wine. It is suitable with the theory

because EBT has red wine in color (shown in picture

below).

The reason why using this

indicator is because this

indicator forming red wine

stable complex 1:1 with some

cations such as: Mg2+, Ca2+, Zn2+,

and Ni2+. Since this titration

has purpose to determine the

level of Ca2+ in the sample, so

EBT is suitable to be used in

this experiment

The reaction is:

Ca2+ + In3- ⇄ CaIn- Group IV |Complexometry and Its Application 26


Based on the expe-riment: the addition of EBT changes the solu-tion to red wine solu-tion


Metal ion indicator

indicator-metal (colorless) (blue)

complex(wine-red)

After that, titrated the sample

which has been added by EBT

indicator with Na-EDTA solution

until the color changes to blue.

It suitable with the theory

where EBT in a buffered solution

pH=10 is blue (shown in the

figure beside).The using of EDTA

as a titrant because EDTA

forming stable complex, soluble

in water forming complex 1:1

with metal ions.

The reaction is:

Ca2+ (aq) + EDTA4- (aq) → Ca(EDTA)2- (aq)

Initially, indicator is added to a solution containing

metal ions. The indicator reacts with the metal ions in

solution to form complex ions, and the solution takes on

color of indicator-metal complex ion. Prior to equivalence

point, added EDTA reacts with free metal ions in solution to

form complex ions. The solution color does not change during

this part of the titration. As the equivalence point is

approached, added EDTA displaces metal ions from the

indicator-metal complex ion.



Based on the expe-riment: indicator EBT is blue in a buffered solution at pH 10.




The experiment is repeated of three times. The volume

of Na-EDTA which require to turns the color from red wine

to blue are 4 ml, 3,9 ml, and 3,8 ml. After that,

calculate the concentration of sample using equation:

Molek of sample = Molek Na-EDTA

From the standardization experiment, we know that the

concentration of Na-EDTA is 0,0123 M. From data above, the

concentration of sample in the first trial is 0.0049 M, in

the second trial is 0.0047 M and the third is 0.0046 M.

So, the average concentration of sample is 0,0047 M

The next step is determining the hardness of water for

each concentration from the calculation we done before.

But, first we have to calculate the mass of Ca2+ in the

sample using equation:

Msample=massofCa

ArCa× 1000100

×dilutionfactor

In the first trial with concentration of sample is 0,0049

M we get mass of Ca2+ in the sample is 0,00196 g. After

that calculate the level of Ca2+ in the sample using

equation:

Level of Ca2+ in water sample ¿ MassofCavsample

And we get the level of Ca2+ is 196 ppm. For the second

trial, the mass of Ca2+ in the sample is 0,00188 g and the

level of Ca2+ in the sample is 188 ppm. Meanwhile, for the


third trial the mass of Ca2+ in the sample is 0,00184 g and

the level of Ca2+ in the sample is 184 ppm. So the average

level Ca2+ in the sample is 189 ppm.

XI.Conclusion

From the experiment, we can conclude that :

1. The concentration of Na-EDTA standard solution is 0,0123

m

2. The level of Ca2+ in seawater is 189 ppm

X. Question and Answer( standardization)

1. NaC10H16N2O8 and C20H12N3O7SNa

2. Calcium concentrations are measured in units of ppm as

calcium, ppm as CaCO3, moles per liter, or any other

convenient concentration unit. Table 1 indicates some of

the concentration units.

TABLE 1: Concentration Unit Conversion Factors

3. We'll take a mixture of ammonia and ammonium chloride

solutions as typical. Ammonia is a weak base, and the

position of this equilibrium will be well to the left:


Adding ammonium chloride to this adds lots of extra

ammonium ions. According to Le Chatelier's Principle,

that will tip the position of the equilibrium even

further to the left.

The solution will therefore contain these important

things:

lots of unreacted ammonia;

lots of ammonium ions from the ammonium chloride;

Enough hydroxide ions to make the solution alkaline.

Other things (like water and chloride ions) which are

present aren't important to the argument.

Adding an acid to this buffer solution

There are two processes which can remove the hydrogen

ions that you are adding.

Removal by reacting with ammonia

The most likely basic substance which a hydrogen ion is

going to collide wih is an ammonia molecule. They will

react to form ammonium ions.

Most, but not all, of the hydrogen ions will be removed.

The ammonium ion is weakly acidic, and so some of the

hydrogen ions will be released again.

Removal of the hydrogen ions by reacting with hydroxide ions

Remember that there are some hydroxide ions present from

the reaction between the ammonia and the water.


Hydrogen ions can combine with these hydroxide ions to

make water. As soon as this happens, the equilibrium tips

to replace the hydroxide ions. This keeps on happening

until most of the hydrogen ions are removed.

Again, because you have equilibria involved, not all of

the hydrogen ions are removed - just most of them.

Adding an alkali to this buffer solution

The hydroxide ions from the alkali are removed by a

simple reaction with ammonium ions.

Because the ammonia formed is a weak base, it can

react with the water - and so the reaction is slightly

reversible. That means that, again, most (but not all) of

the the hydroxide ions are removed from the solution.

Question and Answer(application)


1. At a very high pH values, hydroxide ions can penetrate

layers of metal coordination, and complexes such as

Cu(OH)Y3- may arise. Clear that in line with the decrease

in pH, the equilibrium moves away, the pH value below by

EDTA titration of copper is not feasible. Solution of the

metal ions to be titrated with EDTA buffered to pH

constant despite the release of H3O+ when the complex is

formed. At high pH metal ions tend to hydrolyze and even

precipitate as hydroxide.

2. ppm of CaCO3


M1 x V1 = Me xVe

M1 x 100 ml = 0,01016 M x 15,28 ml

M1 = 0,00155 M

Determination hardness of water of sample (PDAM)

Msample=massofCa

ArCa× 1000100

×dilutionfactor

0,00155=massofCa40

× 1000100

×1

Mass of Ca2+ = 0,0621 g

Level of Ca2+ in water sample ¿ MassofCavsample=0,0621100

=¿

0,000621 g/mL

= 621 mg/L

= 621 ppm

ppm MgCO3



M1 x V1 = Me xVe

M1 x 100 ml = 0,01016 M x 10,43 ml

M1 = 0,00106M


Msample=massofMg

ArMg× 1000100

×dilutionfactor

0,00106=massofMg24

× 1000100

×1

Mass of Mg2+ = 0,0254 g

Level of Mg2+ in water sample ¿ MassofMgvsample=0,0254

100=¿

0,000254 g/mL

= 254 mg/L

= 254 ppm

XI. Referencess

Day, R.A, Jr & Underwood, A.L. 2002. Analisis Kimia Kuantitatif:Edisi keenam.. Translated Iis Sopyan. Jakarta: Erlangga

Larsen, Delmar PhD .2014.9C Complexation Titration.(online),(http://chemwiki.ucdavis.edu/Analytical_Chemistry/Analytical_Chemistry_2.0/09_Titrimetric_Methods/9C_Complexation_Titrations, accessed on December 19th, 2014)

Noname.2013.Complexometry, (online),(http://www.en.wikipedia.org/wiki/complexometry , accessed on December 19th, 2014)

Noname.2013.EDTA, (online),(http://www.en.wikipedia.org/wiki/EDTA , accessed onDecember 19th, 2014)


http://www.en.wikipedia.org/wiki/EDTA

http://www.en.wikipedia.org/wiki/complexometry

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Analytical_Chemistry_2.0/09_Titrimetric_Methods/9C_Complexation_Titrations



Svehla, G.(1985). Vogel: Buku Teks Analisis Oraganik Kualitatif Makrodan Semimikro.(first edition). Translated Setiono, Land Handayana, P.A Jakarta: Kaliman Media Pusaka.

Tim Kimia Dasar. 2014. Panduan Praktikum Kimia Analitik I Dasar-Dasar Kimia Analitik. Surabaya: Jurusan Kimia FMIPA Unesa

Attachment 1

A. Determining (standardization) of Na-EDTA ±0,01 M with CaCl2 as standard solution



HCl 6 MColorless solution

Entered CaCO3

(white powder)into volumetric

Dilute withdistilled wateruntil boundary

Addition of HCl 6MProducing bubble

CaCO3 + distilledwater

White precipitate

Took 25 ml of CaCl2

and moved toErlenmeyer



Addition of buffersolution pH=10

Colorless solution

Addition of EBTIndicator

Red wine solution

Titration processwith Na-EDTA

After reaching the endpoint V1= 16,3 ml V2=16,4 ml V3=16,5 ml

Blue Solution


Addition of buffersolution pH=10

Colorless solution

10 ml of Sample(PDAM)

Colorless solution

Addition of EBTIndicator

Red wine solution

10 ml sample + buffer solution pH=10+ EBTindicator

Red wine solution

Attachment 2

A. Determining (standardization) of Na-EDTA ±0,01 M with CaCl2 asstandard solution

Reaction : Ca2++Y4- ⟶ CaY2—


Titration process with Na-EDTA

After reaching the endpoint V1= 4 ml V2= 3,9 ml V3= 3,8 ml

Blue Solution

After reaching theendpoint

Blue Solution

Mass of CaCO3 : 0,0811 g

Mr of CaCO3 : 100

Determining the concentration of Ca2+

MCa2+¿=

massofCaCO3

MrCaCO3× 1000100

¿

MCa2+¿=

0,0811g100

×1000100

¿

MCa2+¿=0,00811M ¿

V Na-EDTA : V1= 16,3 ml

V2= 16,5 ml

V3= 16,4 ml

Determining the concentration of Na-EDTA

Molek Na-EDTA= Molek CaCO3

M1 x V1 = M2 xV2

M1 x 16,3 ml = 0,0811 M x 25 ml

M1 = 0,0124 M


M1 x V1 = M2 xV2

M1 x 16,4 ml = 0,0811 M x 25 ml

M1 = 0,0123 M


M1 x V1 = M2 xV2

M1 x 16,5 ml = 0,0811 M x 25 ml

M1 = 0,0122 M

Average concentration of Na-EDTA = 0.0124+0,0123+0,01223

= 0,0123 M



Reaction : Ca2++Y4- ⟶ CaY2—

M Na-EDTA : 0,0123 M

V Na-EDTA : V1= 4 ml

V2= 3,9 ml

V3= 3,8 ml

Determination the concentration of sample (PDAM)


M1 x V1 = Me xVe

M1 x 10 ml = 0,0123 M x 4 ml

M1 = 0,0049 M


M2 x V2 = Me xVe

M2 x 10 ml = 0,0123 M x 3,9 ml

M1 = 0,0047 M


M3 x V3 = Me xVe

M3 x 10 ml = 0,0123 M x 3,8 ml

M3 = 0,0046 M

Average concentration of sample = 0.0049+0,0047+0,00463

= 0,0047 M


Msample=massofCa

ArCa× 1000100

×dilutionfactor

0,0049=massofCa40

×1000100

×1

Mass of Ca2+ = 0,00196 g



10=¿

0,000196 g/mL

= 196 mg/L

= 196 ppm

Msample=massofCa

ArCa× 1000100

×dilutionfactor

0,0047=massofCa40

×1000100

×1

Mass of Ca2+ = 0,00188 g


=¿

0,000188 g/mL

= 188 mg/L

= 188 ppm

Msample=massofCaArCa

× 1000100

×dilutionfactor

0,0046=massofCa

40×1000100

×1

Mass of Ca2+ = 0,00184 g


10=¿

0,000184 g/mL

= 184 mg/L

= 184 ppm

Average level of Ca2+ in sample (PDAM) = 196+188+1843


=189 ppm


Complexometry: Determination Hardness of Water

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