Lab: Safety Section 1410 Section 1411 Chem 9 Lab Orientation Tue ...

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Lab: Safety Section 1410 Section 1411 Chem 9 Lab Orientation Tue, Feb 16 Thur, Feb 18 Home Lab Safety p. 2 Laboratory PPE and equipment p. 3-4

Lab: Properties of Water and Physical States Tue, Feb 23 Thur, Feb 25 Procedure p. 5-7 Report p. 8-9

Lab: Measurements Tue, Mar 2 Thur, Mar 4 Procedure p. 10-15 Report p. 16-19

Lab: Paper Chromatography of Colored Inks Tue, Mar 9 Thur, Mar 11 Procedure p. 20-23 Report p. 24-25

Lab: Flame Tests and EM Calculations Tue, Mar 16 (both sections) Procedure p. 26-28 Report p. 29-31

Lab: Detection and Absorption of Ultraviolet Light Tue, Mar 23 Thur, Mar 25 Procedure p. 32-34 Report p. 35-37

Lab: Lewis Structures and Molecule Shapes Tue, Mar 30 Thur, Apr 1 Procedure p. 38-41 Report p. 42-45

Lab: Combustion Reactions Tue, Apr 6 Thur, Apr 8 Procedure p. 46-49 Report p. 50-52

Lab: Single Replacement Reactions Tue, Apr 20 Thur, Apr 22 Procedure p. 53-55 Report p. 56-59

Lab: Double Replacement Reactions Tue, Apr 27 Thur, Apr 29 Procedure p. 60-64 Report p. 65-70

Lab: Conductivity of Aqueous Solutions Tue, May 4 Thur, May 6 Procedure p. 71-74 Report p. 75-77

Lab: Acids, Bases and pH Tue, May 11 Thur, May 13 Procedure p. 78-81 Report p. 82-85

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Lab: Synthetic Polymers and Plastics Tue, May 18 Thur, May 20 Procedure p. 86-89 Report p. 90-92

Coronavirus and the Science of Soap Tue, May 25 Thur, May 27 Procedure p. 93 Report p 94-96

Everyday Chemistry in Your Home Tue, Jun 1 Thur, Jun 3 Assignment p. 97

EVERYDAY CHEMISTRY (CHEM 9) HOME LABORATORY SAFETY RULES

The following is a list of rules that is designed to ensure your safety at home during our chemistry labs. Failure to follow these safety rules may result in loss of points from the related assignments. It can also result in the removal from class on the day of the incident and one subsequent class session. In addition, violation of the rules may result in a referral to the Office of Student Judicial Affairs for disciplinary action. 1. The lab equipment has been provided solely for use in the specified Chem 9 lab experiments, as

described in lab procedures. Make sure not to use included tools for purposes other than intended use.

2. Conduct only experiments as directed in the lab procedures. When not being used during a lab activity, all equipment used must be stored in such a manner as to be in an inconspicuous place and away from nonstudents.

3. Reactants to be used in the kitchen-safe experiments are standard, non-toxic household compounds such as vinegar and aluminum foil. Maintain safe handling of those compounds in your working area.

4. Safe handling of lab equipment glassware and your own kitchen equipment should be adhered to. Be careful when handling glassware and use caution for cleaning up broken glassware.

5. Color changing UV beads are small parts and could be a choking hazard for children. 6. Eating, drinking or gum chewing is not permitted during lab. 7. Appropriate clothing, including shirt, shoes, etc. are required. During lab experiments, protective

eyewear (i.e. goggles and/or eyeglasses) should be worn. Students must provide clothing and eyewear.

8. Keep all lab areas clear of extra books, clothing and other personal items to allow for a clear working space and emergency evacuation.

9. Photography, videos and recordings can only be taken with instructor consent. They may not be given, sold or published in print or online without the written consent of the instructor, and can only be used for the explicit purposes of the class. It will be considered a serious disciplinary offense for a student to take or possess pictures or recordings of the biological specimens and used otherwise.

10. If you should cut, puncture, or wound yourself with any instruments, immediately notify your instructor, give yourself first aid, and seek medical attention as necessary.

11. Experiment should only be conducted during established class time so that they can be observed by your instructor.

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Chemistry Lab Safety Personal Protective Equipment (PPE)

safety goggles lab coat

gloves

On your own: closed-toe shoes, no loose clothing, no jewelry, hair tied back In the lab: No food or drinks in lab No papers or backpacks on lab bench hazardous waste disposal emergency eye wash

emergency fire blanket emergency exit

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Chemistry Lab Equipment – Chem 9

Pipetter (with tip) Electronic balance Bunsen burner Wash bottle

Graduated cylinders Graduated cylinder Erlenmeyer flasks Beakers

Test tubes (in rack) Plastic test tubes Dropper pipet Chromatography strips

Thermometers Metric ruler UV beads

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Properties of Water and Physical States Objectives The objectives of this lab are to: a) observe capillary action and surface tension in liquid water b) understand particulate properties of liquid water c) understand particulate properties of gases by observing air pressure d) understand how heating and cooling affect a gas Background Water is found almost everywhere on Earth: on the surface, in underground reservoirs, in the atmosphere, and is present in animals and plants. Water is the only common substance that can exist in three physical states, simultaneously, as a solid, liquid or gas, on Earth. Water is vital for living organisms, including our lives. Humans depend on clean, fresh drinking water to survive. Consider how water is part of your daily routine, for drinking, cooking, sanitation. We see water all around us on Earth. What may not be as apparent is that water has several unusual and unique properties. Chemistry is the study of matter and the changes it undergoes. Chemists use three viewpoints to understand and describe matter. One is the macroscopic view, using observations and measurements, to understand physical and chemical properties. We can also describe matter using chemical symbols and chemical formulas; this is known as the symbolic representation. Liquid water, water vapor (gas), and solid ice all have the same chemical formula, H2O. We can also describe matter in terms of what its particles “look” like, which is the particulate view. These are models, because we cannot actually see atoms and molecules; they are submicroscopic particles. Chemists study how matter behaves by studying how its particles (atoms and molecules) behave. Macroscopic, particulate, and symbolic representations of the water molecule are shown in Figure 1. Water is essential for all living organisms. What is not apparent are the properties that water molecules have at the particulate level, that lead to the properties that we see at the macroscopic level. In fact, if water did not have these properties, life as we know it could not exist. The fact that water can exist in three physical states (solid, liquid, gas) on Earth is both unique and unusual. Solid ice, liquid water, and water vapor all play a critical role in our daily lives, but are often taken for granted. We breathe in water vapor every day. We drink liquid water every day. Our planet has solid ice at both poles, and supports ecosystems on those ice forms.

Figure 1: Three viewpoints of water

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Water molecules in the liquid state are strongly attracted to other water molecules. This is the property known as cohesion. Surface tension and water droplets are a result of cohesiveness of liquid water. Water molecules are also attracted to other substances, which is adhesion. Cohesion and adhesion together allow for the property of capillary action, in which water can travel through a porous material, even against gravity. Most water vapor is in Earth’s atmosphere. We cannot see water vapor, but our bodies take it in every time we inhale. The air we breathe is composed of gases that are also invisible to us. These gas particles exert pressure on its surroundings. At sea level, the gases in the atmosphere exert almost 15 pounds of pressure per square inch of surface area. Your body has adapted to atmospheric pressure, so it does not affect you. Procedure

Safety Perform your experiments in an area clear of any obstructions. Small objects can be a danger to young children or pets. Clean spilled water immediately. Personal protective equipment (PPE) optional: safety goggles Materials and Equipment two identical-sized cups (plastic or glass), plastic cup, paper towel, tissue, index card or playing card (should cover the mouth of the plastic cup), shallow bowl, small paper clip Experimental Procedure Part A: Capillary Action of Liquid Water 1. Place the two identical cups on a surface that will be

undisturbed for at least 24 hours. Set the cups about 2 inches apart.

2. Add water to one cup so that it is about 3/4 full. The other cup should be empty.

3. Use a half-sheet of paper towel and roll or fold to about 2 inches width. Place one end of the paper towel strip into the water and the other end into the empty cup. Take a photo to document the start of the experiment.

4. Let them sit for one hour. Take a photo to document the progress of the experiment. 5. Return the next day (or more than 24 hours) and take a photo to document the result.

Figure 2: macroscopic and particulate views of gas, liquid, and solid

Figure 3: Experimental set-up

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Part B: Surface Tension of Liquid Water 1. Place the shallow bowl on a stable surface. Add water so that it almost fills

the bowl. 2. Allow the water to settle for a couple of minutes. 3. Bend one end of the paper clip upwards, see Figure 3. 4. Using the bent end, carefully ease the paper clip into the water. If the paper

clip does not “float”, try again, but more gently, without disrupting the surface.

5. Once the paper clip is “floating”, take a photo to document the experiment. 6. Now disturb the surface of the water. Take another photo to document the result. Part C: Air Pressure 1. Do these steps over a sink, or a large bowl. 2. Add water to the cup so that it is about halfway full. 3. Cover the cup with the card, making sure that the card completely covers the mouth of the container. 4. Keep your hand on the card and turn the cup upside down (over the sink or bowl). 5. Slowly take your hand away.

Part D: Air Pressure with Changing Temperatures – Instructor Demo

Materials and Equipment two clear plastic cups, small plastic bottle (should fit into clear plastic cups), soap solution in a cup, warm or hot water (about 50 °C), cold water Experimental Procedure 1. Pour hot water into an empty cup until it is

about ½-full. 2. Turn the bottle over and dip the opening of

the bottle into the soap solution so that a film forms.

3. While holding the bottle, slowly push the bottom of the bottle into the hot water.

4. Pour cold water into another cup until it is about ½-full.

5. If there is still a bubble on the bottle, slowly push the bottom of the bottle into the cold water. If a bubble is not still on the bottle, make another film by re-dipping into the soap solution.

Figure 4: paper clip with bent end

Figure 5: Experimental set-up for instructor demo

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Name: _______________________________ Chem 9 Section: _____________ Experiment Date: ______________________

Properties of Water and Physical States Questions Draw circles to represent the molecules in a solid, liquid, and gas. Include motion lines/arrows if there is molecular motion. Solid Liquid Gas Do the molecules of a liquid have strong or weak attractions? Explain. Are the molecules of a liquid randomly-arranged or orderly-arranged? Explain. What is cohesion (of water)? Which experiment demonstrated the property of cohesion? What was unusual, and how does cohesion play a role in the observed results? What is adhesion (of water)? Which experiment demonstrated the property of adhesion?

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What was unusual, and how does adhesion play a role in the observed results? Do the molecules of a gas have strong or weak attractions? Explain. Are the molecules of a gas randomly or orderly arranged? Explain. When the molecules of a gas hit each other, do they normally stick together or bounce off? The particulate view of the experiment “Air Pressure with Changing Temperatures” is modeled at right. What caused the bubble to form when the bottle was placed in hot water? Include a description of the molecular motion of the gas molecules inside the bubble and the force on the bubble from the outside air (air pressure). Why did the bubble get smaller when the bottle was placed in cold water? Include a description of the molecular motion of the gas molecules inside the bubble and the force on the bubble from the outside air (air pressure). Remember to attach photos to your lab report submission. There should be five photos from the at-home experiments.

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Measurements

Objectives The objectives of this laboratory are to:

1. Use standard laboratory measurement devices to measure length, volume and mass amounts. 2. Use these measurements to calculate areas and volumes 3. Determine the density of aluminum (applying the technique of water displacement).

Background Chemistry is the study of matter. Our understanding of chemical processes thus depends on our ability to acquire accurate information about matter. Often, this information is quantitative, in the form of measurements. In this lab, you will be introduced to some common measuring devices, and learn how to use them to obtain correct measurements, each with correct precision. All measuring devices are subject to error, making it impossible to obtain exact measurements. Students will record all the digits of the measurement using the markings that we know exactly and one further digit that we estimate and call uncertain. The uncertain digit is our best estimate using the smallest unit of measurement given and estimating between two of these values. These digits are collectively referred to as significant figures. Note, the electronic balance is designed to register these values and the student should only record the value displayed. When making measurements, it is important to be as accurate and precise as possible. Accuracy is a measure of how close an experimental measurement is to the true, accepted value. Precision refers to how close repeated measurements (using the same device) are to each other. A metric ruler will be used to measure length in centimeters (cm).

Measuring Length

Here the “ruler” markings are every 0.1-centimeter. The correct reading is 1.65 cm. The first 2 digits 1.6 are known exactly. The last digit 1.65 is uncertain. You may have instead estimated it as 1.66 cm.

The measuring devices used in this lab may have different scale graduations than the ones shown. To find the precision, find the smallest unit on your measuring device, and add a decimal place (the uncertain digit). In general, the more decimal places provided by a device, the more precise the measurement will be.

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Measurements obtained in lab will often be used in subsequent calculations to obtain other values of interest. Thus, it is important to consider the number of significant figures that should be recorded for such calculated values.

• If multiplying or dividing measured values, the result should be reported with the lowest number of significant figures used in the calculation.

• If adding or subtracting measured values, the result should be reported with the lowest number of decimal places used in the calculation.

Significant Figures in Calculated Values 1) A student runs 18.752 meters in 54.2 seconds. Calculate his velocity (or speed). velocity = distance/time = 18.752 m / 54.2 s = 0.345978 m/s from calculator = 0.346 m/s for 3 significant figures 2) The mass of a glass is measured to be 12.456 grams. If 10.33 grams of water are added to

this glass, what is the total combined mass? total mass = 12.456 g + 10.33 g = 22.786 g from calculator = 22.79 g for 2 decimal places

In this lab, students will also determine the density of aluminum. Volume is the amount of space occupied by matter. An extensive property is one that is dependent on the amount of matter present. Volume is an extensive property. The volume of a liquid can be directly measured with specialized glassware, typically in units of milliliters (mL) or liters (L). In this lab, a beaker and a graduated cylinder will be used to measure liquid volumes, and their precision will be compared. Note that when measuring liquid volumes, it is important to read the graduated scale from the lowest point of the curved surface of the liquid, known as the liquid meniscus.

Measuring the Volume of a Liquid

Here, the graduated cylinder markings are every 1-milliliter. When read from the lowest point of the meniscus, the correct volume reading is 30.0 mL. The first 2 digits 30. are known exactly. The last digit 30.0 is uncertain. Even though it is a zero, it is significant and must be recorded.

For regularly shaped solids, such as a cube, sphere, cylinder, or cone, the volume (V) can be calculated from its measured dimensions (length, width, height, and diameter) by using an appropriate equation.

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For irregularly shaped solids, the volume can be indirectly determined via the volume of water (or any other liquid) that the solid displaces when it is immersed in the water (Archimedes’ Principle). The units for solid volumes are typically cubic centimeters (cm3) or cubic meters (m3). Note that 1 mL = 1 cm3.

Measuring the Volume of an Irregularly Shaped Solid

Volume water displaced = Final volume – Initial volume or V = Vf – Vi Volume water displaced = Volume of solid

Density is defined as the mass per unit volume of a substance. Density is a physical property of matter. Physical properties can be measured without changing the chemical identity of the substance. Since pure substances have unique density values, measuring the density of a substance can help identify that substance. Density is determined by dividing the mass of a substance by its volume.

Density is commonly expressed in units of g/cm3 for solids, g/mL for liquids, and g/L for gases.

Equations, Constants and Conversion Factors: Arectangle = lw Acircle = πr2 r = ½ diameter Vcube = lwh Vcylinder = πr2h ∆V = Vfinal - Vinitial D = m/V 1 mi = 1.609 km 1 gal = 3.785 L 1 lb = 454 g 1 in = 2.54 cm 1 qt = 0.946 L 1 mL = 1 cm3 π = 3.14

Procedure Safety All lab procedures should be conducted during lab session, under guidance of instructor; instructor to provide supervision and assistance. Use lab equipment as directed in procedure, no other purposes. Use caution when handling glassware. Return all equipment to storage, safely away from children and pets, when not in use. Personal protective equipment (PPE) optional: safety goggles

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Materials and Equipment Lab equipment: 100-mL graduated cylinder, 150-mL beaker Also needed: metric ruler, measuring cup, three-dimensional cube (e.g. cereal box), three-dimensional cylinder (e.g. soda can), tap water Experimental Procedure Part A: Measuring the Dimensions of Regular Geometric Shapes 1. Determine the precision of the metric ruler before you begin measurement.

2. Measure the dimensions of the two geometric shapes on the “shape sheet”, page 6 of this lab

procedure. Best results if measured from a printed page, or a screen set to 100% view. Measure length (l) and width (w) of the rectangle, and the diameter of the circle.

3. Use the equations from page 3 of lab procedure to calculate the area of the rectangle (Arectangle),

and the radius (r) and area of the circle (Acircle). Part B: Measuring the Dimensions of Regular Geometric Solids 1. Determine the precision of the metric ruler before you begin measurement.

2. For the cube/box, measure the dimensions of its length (l), width (w), and height (h). For the

cylinder/can, measure the dimensions of its circular base (diameter) its height (h).

3. Use the equations from page 3 of lab procedure to calculate the volume of the cube (Acube), and the radius (r) and volume of the cylinder (Vcylinder).

Part C: Measuring the Volumes of Liquids 1. Determine the precision of the beaker and the graduated cylinder before you begin measurement.

2. Measure ¼ cup tap water and pour into the beaker.

3. Determine the measurement value of the liquid (dependent on the precision of the beaker).

Remember to read the volume at the bottom of the meniscus.

4. Carefully pour the water from the beaker into the 100-mL graduated cylinder.

5. Determine the measurement value of the liquid (dependent on the precision of the graduated cylinder). Remember to read the volume at the bottom of the meniscus.

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Part D: Determining the Density of Aluminum Experimental procedure provided as instructor demo video (~ 1min) 1. Using the electronic balance in the weigh room to determine the mass of your 100-mL beaker. 2. Obtain 20-25 dry aluminum pellets from the front bench. Transfer pellets to the beaker weighed

in the previous step, and measure the mass of the beaker and pellets together. 3. Add 30-35 mL of distilled water into your 100-mL graduated cylinder. Precisely measure this

volume. 4. Carefully add all of the aluminum pellets to the water/cylinder, making sure not to lose any water

to splashing. Also make sure that the pellets are all completely immersed in the water. Measure the new volume of the water plus the pellets.

5. When finished, pour the water from the graduated cylinder into the sink. Retrieve the aluminum

pellets and return them to the front bench; there is a container for wet aluminum pellets. Do not put into the original container which is for dry aluminum pellets.

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Shape sheet Rectangle Circle

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Name:

Chem 9 Section: Experiment Date:

Lab Report: Measurements

Part A: Measuring the Dimensions of Regular Geometric Shapes Experimental Data

Shape Dimensions Precision Measurement # Significant Figures

Rectangle Length

Width

Circle Diameter

Data Analysis Indicate the precision for each measuring instrument below.

Perform the conversions indicated. Show your work and report your answers in scientific notation.

Convert the measured rectangle length to nm

Convert the measured circle diameter to Mm

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Use your measurements to calculate the area of each shape, in cm2 Equations and constants for areas of shapes are in Lab Procedure. Show your work and report your answers to the correct number of significant figures.

Area of rectangle

Area of circle Part B: Measuring the Dimensions of Regular Geometric Solids Experimental Data

Shape Dimensions Precision Measurement # Significant Figures

Cube

Length

Width

Height

Cylinder Diameter

Height

Data Analysis Use your measurements to calculate the volume of each shape, in cm3 Equations and constants for volumes of shapes are in Lab Procedure. Show your work and report your answers to the correct number of significant figures.

Volume of cube

Volume of cylinder

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Part C: Measuring the Volumes of Liquids Experimental Data

Measuring Device Precision Volume Measurement # Significant Figures Beaker

Graduated Cylinder Data Analysis Perform the conversions indicated below. Show your work and report your answers in scientific notation.

Convert the volume of the water measured in the graduated cylinder to hL

Convert the volume of the water measured in the graduated cylinder to gallons, gal

Part D: Determining the Density of Aluminum Experimental Data

Mass of empty beaker

Mass of beaker and pellets

Initial volume of water in cylinder

Final volume of water and pellets

Volume of pellets

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Data Analysis

1) Use the measurement values of “Mass of empty beaker” and “Mass of beaker with pellets” to

calculate the mass of pellets (by difference). Show your work. 2) Circle one: When performing the above calculation, significant figures / decimal places (circle one)

are the primary consideration. 3) Use the measurement values of “Initial volume of water in cylinder” and “Final volume of water in

cylinder” to calculate the volume of pellets (by difference). Show your work. 4) Use your values for mass and volume of to calculate the density of aluminum, in g/cm3. Show your

work and report your answer to the correct number of significant figures. 5) Indicate the appropriate units for the following measurements:

Length measured with a laboratory metric ruler

Mass measured with the electronic balance

Volume of a liquid (measured with a graduated cylinder or beaker)

Volume of a solid (calculated using equation)

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Paper Chromatography of Colored Inks


a) Use paper chromatography to identify whether certain colored inks are pure substances or mixtures.

b) Obtain a paper chromatogram c) Make accurate measurements using a metric ruler d) Identify components of inks by calculating Rf values

Background Chromatography is a method of physically separating mixtures into its individual components. It is a common laboratory technique used to identify unknown components in mixtures. A mixture is a combination of substances that can be separated because they are not chemically bonded (as to a compound, which has elements chemically bonded together). There are several types of chromatography; all types employ a mobile phase or eluent (it can be liquid or gas), which is forced through a stationary phase (a solid or semi-solid). Mixtures are separated because some components will be more attracted to the stationary phase (and stick to it) while some components will be more attracted to the mobile phase (and travel with it). By eye, we cannot know if an ink color is a single component color or a mixture of colors. Some inks are homogeneous mixtures; we cannot see the component colors. In paper chromatography, a mixture is dissolved and pulled across a piece of paper. The mixture separates because its components travel across the paper at different rates, based on their attraction to the paper or solubility in the solvent. The mobile phase takes advantage of differing solubility or polarity of the components in order to separate them. This process is called elution. The ink samples are first applied to the chromatography paper, on a line drawn near the bottom of the paper (starting line). The starting line and ink samples must be above the level of the mobile phase when the paper is placed inside the beaker. If the starting line is below the liquid level, the inks will wash out into the mobile phase rather than elute up the stationary phase. Another line is drawn further up, closer to the top edge of the paper. This is the finish line. When the mobile phase reaches that line, any inks that are mixtures should be clearly separated. When the solvent front reaches the finish line, the paper should be removed immediately from contact with the mobile phase. There are a few difficulties commonly encountered in the elution process. One problem is that spots tend the spread out as they elute and can bleed into each other as they proceed up the paper. Another problem is an uneven solvent front. This can happen if the container is moved, even slightly nudged. If the mobile phase sloshes inside, the elution trails may travel diagonally.

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The figure below shows a typical chromatogram. The distance travelled by a single component is the distance moved by solute, from the starting line = D. The distance the solvent travels is the solvent front, from the starting line to the finish line = F. These are used to calculate the retention factor, Rf, for each spot or component in a mixture.

A component with a given solubility travels along with the mobile phase at one rate, regardless of what other components are present in the sample. If the red part of purple ink travels at the same rate as pure red ink, and both stop in the same place, the two should be the same red ink. The two red spots should have the same retention factor, Rf values.

Comparing the Rf values allows for the confirmation of a component in multiple samples because unique components have unique Rf values.

Procedure

Safety All lab procedures should be conducted during lab session, under guidance of instructor; instructor to provide supervision and assistance. Use lab equipment as directed in procedure, no other purposes. Use caution when handling glassware. Return all equipment to storage, safely away from children and pets, when not in use.

Personal protective equipment (PPE) optional: safety goggles

Materials and Equipment Lab equipment: chromatography paper strips, 150-mL beaker

Also needed: metric ruler, various felt-tip colored ink pens (not gel ink), pencil, binder clip or tape, clear cup (glass or plastic), tap water

D

F

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Part A: Preparation of Chromatography Paper

1. Use 6 chromatography paper strips, one at a time. 2. Use pencil for measurements and marks on the paper prior to elution. Use the metric

ruler to measure centimeters, cm. Be sure to determine the precision of the ruler before using it for measurements.

3. Using the pencil, draw two straight lines across each paper strip. • 1.50 cm above the bottom edge. This is the starting line. • 8.50 cm above the bottom edge. This is the finish line.

4. Use a binder clip or tape to curl the top of the paper strip around a pencil. Set the

pencil across the top of a beaker or cup. Adjust the pencil to ensure that the paper strip hangs down without touching the bottom of the cup.

5. Remove the pencil with strip from the cup. On the starting line, apply a dot of ink color.

You can use any colors, but try to vary the colors as much as possible, and try to make sure one of the colors you test is a black ink. At the top of each strip, write the name of the ink spot color in pencil, e.g. “black”. You will use this to identify the chromatogram after elution.

6. Use the beaker to measure 30-mL tap water (or measure 2 tablespoons).

7. Place the cup where you intend to do the experiment and then pour the water into the

cup and let it settle.

Part B: Acquisition of Chromatogram 1. Gently place the pencil/paper strip into the cup. The water level should be below the marker

spot, or else the ink will just run off into the water. Ensure that the paper strip hangs down without touching the sides or bottom of the cup.

2. As the water (mobile phase) travels up the paper, it may begin to move some of the ink. It will take about 10 minutes for the solvent front to reach the finish line.

3. When the solvent front reaches the finish line, remove the pencil/paper strip from the

cup. Remove binder clip or tape and flatten out the strip. Allow to dry.

4. Repeat this process for all 6 chromatography paper strips.

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Part C: Interpretation of Chromatogram 1. After your chromatogram has dried completely, you may continue to use pencil, or you may

use any pen to write on it. Some ink samples may no longer be on the starting line and have travelled up the paper, becoming one or more separate color spots between the starting and finish lines. Other ink samples may not have moved at all off the starting line.

2. Draw a circle around each individual color spot, including any on the starting line. Some color spots may be spread out; include the entire spot around all edges.

3. Determine the approximate middle of each color spot and draw a plus sign (+) in the center

of each color spot.

4. Use the metric ruler to measure the distance from the starting line to each plus sign, for each color spot. Record this distance for each spot in your lab report. Though the ruler precision is 0.01 cm, round the values to the nearest 0.1 cm. These are the D values, in cm.

5. Measure the distance between the starting line and the finish line or, the farthest up that

the solvent front reached. This will most likely be greater than the original 8.50 cm because the water continues to travel along the filter paper as the chromatogram dries. Record this distance. This is the F values, in cm.

6. Calculate the retention factor (Rf) for each spot and record the values in your lab report.

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Name:

Chem 9, Section: Experiment Date:

Paper Chromatography of Colored Inks Part C: Interpretation of Chromatogram Experimental Data Record D values for each eluted spot, and F value for each strip. Although the metric ruler precision is 0.01 cm, round your measured values to 0.1 cm Draw an X through any unused boxes.

Ink Color Distance(s) Traveled by Eluted Spot(s), D Distance Traveled by solvent front, F value

Calculate and record the Rf value for each eluted spot, using the equation

Ink Color Rf Value for Eluted Spot

Show calculations for Rf below for one of your ink colors:

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Data Analysis

1. Record ink colors in the appropriate column; for mixtures, list the colors of the components. Mixtures Single-Component Colors

2. Which component travelled farthest? What does this indicate about that component’s

attraction to the mobile phase? To the stationary phase? 3. If a component has an Rf value of zero, what does that indicate about its component

attraction to the mobile phase? To the stationary phase?

4. What is F value for the spot? 5. What is D value for the spot? 6. What is the Rf value for the spot?

7. Air is a homogeneous mixture. Nitrogen gas is one component, comprising 78 % of the total. Convert the % value (parts per hundred) to parts per million (ppm).

8. The standard/limit for carbon monoxide is 35 ppm in one hour. Convert this value to %

(parts per hundred) Remember to include a photo of your chromatography strips when you submit your lab.

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Flame Tests and EM Calculations Objectives The objectives of this lab are to: a) Perform flame tests of metal cations in order to observe their characteristic colors, b) Perform calculations to determine the frequency and energy of the emitted photons. c) Relate these results to the types of electronic transitions occurring in these elements. Background Electromagnetic (EM) radiation is energy in the form of waves. Waves are characterized by their wavelength (λ) and frequency (υ). Wavelength is defined as the distance between successive crests (or troughs) on a wave and is measured in meters. Frequency is defined as the number of waves that pass a given point every second and is measured in 1/seconds (1/s or s-1), or Hertz (Hz).

All electromagnetic waves travel at the speed of light (c), or 3.0 x 108 m/s. The relationship between the wavelength, frequency and speed of a wave is given by the equation:

c = λν

Electromagnetic radiation also occurs as discreet “packets” called photons. The energy (E) of a photon (in Joules, J) is given by the equation:

E = hν Here, h is Planck’s constant, which has a value of 6.63 x 10-34 J•s. Visible light is the most familiar example of electromagnetic radiation. Differences in the wavelengths of visible light are manifested as different colors, shown in the color spectrum below (colors can be seen in the PDF document online). Other examples of electromagnetic radiation include X-rays, ultraviolet light, infrared light, microwaves and radio waves.

λ: 400 nm 500 nm 600 nm 700 nm

UV ← Violet Red → IR

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So, how does electromagnetic radiation relate to flame tests? Well, when an atom or ion absorbs energy, its electrons can make transitions from lower energy levels to higher energy levels. The energy absorbed could be in the form of heat (as in flame tests), or electrical energy, or electromagnetic radiation. However, when electrons subsequently return from higher energy levels to lower energy levels, energy is released predominantly in the form of electromagnetic radiation. The spacing between energy levels in an atom determines the sizes of the transitions that occur, and thus the energy and wavelengths of the collection of photons emitted.

Larger transition – higher energy photon released (shorter wavelength) Small transition – lower energy photon released (longer wavelength)

1 2 3 4 5 6 If emitted photons are in the visible region of the spectrum, they may be perceived as lines of different colors (note that photons outside the visible spectrum may also be emitted but cannot be seen). The result is called a line emission spectrum and can serve as a ‘fingerprint’ of the element to which the atoms belong. For example, the line spectra shown below for the elements helium and carbon are clearly quite different.

Helium

Carbon Unfortunately, techniques more sophisticated than those used in this lab are required to obtain such line spectra. To the naked eye, when an element is vaporized in a flame (or an electrical discharge) the emission spectrum will appear to be just one color. For example, helium gas when excited by an electrical discharge emits light that appears an orange-yellow color. This one color results from a combination of all lines of the emission spectrum, in proportion to their intensities. As many elements will still produce distinctive colors under such conditions, simple flame tests can be used to identify these elements. In fact, flame tests were used to identify elements long before the invention of modern techniques, such as emission spectroscopy.

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Procedure Safety Exercise appropriate caution when using the Bunsen burner. Personal protective equipment (PPE) required: safety goggles Materials and Equipment wooden boiling sticks, wash bottle with distilled water, Bunsen burner, solutions: LiCl (aq), NaCl (aq), KCl (aq), CuCl2 (aq), BaCl2 (aq), CaCl2 (aq). gas discharge tubes for hydrogen, helium, mercury, hand-held spectroscope Experimental Procedure Flame Tests of Metal Cations Instructor will dip a wooden boiling stick into one of the solutions supplied, and then hold it in the Bunsen burner flame. Students will record the dominant flame color observed. Analysis: For each metal cation flame test performed, determine the wavelength corresponding to the observed flame color from the table below.

Dominant Color Approximate Wavelength (in nm)* *Wavelength values here are given for the mid-range of the color indicated. Red 701

Red-Orange 622

Orange 609

Orange-Yellow 597

Yellow 587

Yellow-Green 577

Green 535

Green-Blue 492

Blue 474

Blue-Violet 455

Violet 423

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Name: Chem 9 Section:

Experiment Date:

Flame Tests and EM Calculations Experimental Data and Observations: Flame Tests of Metal Cations

Aqueous solution Chemical name Dominant Color Wavelength

(nm) BaCl2

CaCl2

CuCl2

KCl

LiCl

NaCl

Data Analysis: EM Calculations Using the dominant color and corresponding wavelength value in nm, calculate the wavelength, and frequency and energy for BaCl2. Show all work, including equations used in your calculations, and report answers in scientific notation. • Wavelength (in m):

• Frequency (in s-1): • Energy (in J):

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Using the wavelengths recorded on page 1, calculate the corresponding wavelengths, frequencies and photon energies for each compound tested. Record the values, in scientific notation, in the table below. (You do not need to show the work on these, only the calculated answers.)

Solution Wavelength (m) Frequency (s-1) Energy (J)

BaCl2

CaCl2

CuCl2

KCl

LiCl

NaCl

Questions 1) Calculate the wavelength when frequency is 4.3 x 103 Hz. Put your answer into scientific notation.

2) Calculate the energy when frequency is 900 M Hz. Put your answer into scientific notation.

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3) A high-band 5G cellular network uses a frequency of 25 GHz. Calculate its wavelength. Put your answer into scientific notation.

4) For the same radio station, calculate the energy. Put your answer into scientific notation.

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Detection and Absorption of Ultraviolet Light

Objectives The objectives of this laboratory are:

a) Understand what ultraviolet (UV) light is b) Understand the different types of UV light c) Test the abilities of various materials to absorb UV light

Background Sunlight contains 3 forms of energy. Visible light has wavelengths of 750 to 400 nm. Ultraviolet (UV) light has shorter wavelengths (see table below), cannot be seen, and has higher energy. Infrared (IR) radiation can be felt as warmth, and is the major source of heat for Earth. Though UV is a fraction of sunlight, it can be damaging to living organisms. All of these are forms of energy in the electromagnetic (EM) spectrum. Just as visible light can be broken into its components, so can UV light: UV-A, UV-B, UV-C and vacuum-UV. UV-A has lowest energy and is least damaging; UV-A is also called “black light.” UV-B and UV-C have higher energies and can cause break bonds of molecules, causing changes in DNA and ultimately, skin cancers.

UV light type wavelength relative energy comments UV-A 320 – 400 nm lowest energy reaches Earth in greatest amount UV-B 280 – 320 nm higher energy than UV-A,

but less than UV-C most is absorbed by ozone

UV-C 200 – 280 nm highest energy absorbed by ozone and oxygen The majority of UV-B is absorbed by ozone in the stratosphere. Though UV-C is most damaging, it is totally absorbed by oxygen and ozone. Depletion of the ozone layer has allowed more UV light to reach us, resulting in more cases of skin cancers. Consequently, we have become aware of the need to protect ourselves from UV light. What protects us from UV light? One strategy would be to avoid exposure to any type of sunlight. Since we cannot avoid sunlight while outdoors, we can physically or chemically block the sun. A wide variety of commercial sunscreens are available with sun protection factor (SPF) ranging from SPF 2 to SPF 100. These lotions contain organic molecules that absorb UV light.

Reprinted with permission from h2odistributors.com

i i i i h2 di t ib t

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Procedure

Safety

All lab procedures should be conducted during lab session, under guidance of instructor; instructor to provide supervision and assistance. Use lab equipment as directed in procedure, no other purposes. No materials used in this experiment should be ingested. Return all equipment to storage, safely away from children and pets, when not in use.

Personal protective equipment (PPE) optional: safety goggles or eyewear

Materials and Equipment

Lab equipment used by instructor: UV-sensitive beads, sunscreens, cheap and UV sunglasses, clear and opaque plastic, plexiglass plate, foil, cloth, paper, small plastic bags, laboratory UV light Needed for at-home experiments: UV beads, plastic bag, sunscreen, sunglasses Part A: Detecting UV light 1. Place 5 UV-sensitive beads in two small plastic bags. These beads will turn color in the

presence of UV light. The higher the intensity of UV light the stronger the color change. Label the two bags: “Control” and “Experiment”. The Control bag will only be exposed to indoor light.

3. Expose Experiment bag to the laboratory UV light. Record observations (color change or no

change) after 10 seconds exposure. 4. Take the Experiment bag outside, in an area of direct sunlight. Record observations after 10

seconds of exposure to sunlight. Find a shaded area nearby. Record what you see after 10 seconds of exposure to shade.

Part B: Absorption of UV Light by lotions 1. Place 5 UV-sensitive beads into two more small bags. You should now have a total of four

bags. Label with the SPF numbers. Note the color of the beads under indoor lighting.

2. Add one dab of each lotion onto one side of each bag. Spread the lotion and allow to dry.

3. Expose all bags to UV light. Record observations after 10 seconds exposure to UV light.

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Part C: Absorption of Outdoor UV light by various materials 1. Place 5 UV-sensitive beads into two small plastic bags. Label the two bags: “Control” and

Experiment”. 2. Place on a tray: Control and Experiment bags, and one of each of the materials to be tested

(clear plastic, opaque plastic, paper, cheap sunglasses, UV sunglasses, glass plate, foil and cloth). Take the tray outside to an area of direct sunlight.

3. Both Control and Experiment bags will begin exposed to the direct sunlight. Use each type

of material to block the sunlight to the Experiment bag by holding directly above it. 4. Record observations of the Experiment bag when shielded by each sample. The Experiment

bags begin colored due to UV exposure, so your observations will still include color changes, but to less colored or whitening.

Part D: Home experiments using sunscreen and other materials 1. Gather materials from your home to test for UV. Include at least one sunscreen and one

piece of paper. You may choose to test all the same materials (including different SPF sunscreens and sunblock) as shown in the instructor demo.

2. Prepare 2-5 clear plastic bags, depending on how many materials you would like to test. In each bag, add 5-10 UV-sensitive beads.

3. Label one bag “Control”, all other bags will be Experiment”. 4. For any of the lotions (sunscreen, sunblock), add one dab of each lotion onto one side of

each bag. Spread the lotion and allow to dry. 5. Take Control and Experiment bags, and all materials to be tested outside to an area of direct

sunlight. Test one lotion or material at a time, placing Control bag and Experimental bag alongside each other.

6. Use each type of material to block the sunlight to the Experiment bag by holding directly

above it. Record observations of the Experiment bag when shielded by each sample. The Experiment bags begin colored due to UV exposure, so your observations will still include color changes, but to less colored or whitening.

35

Name:

Chem 9, Section: Experiment Date:

Detection and Absorption of Ultraviolet Light Experimental Data and Observations Part A: Detecting UV Light

Condition Control beads Experiment beads

indoor UV lab light

outdoor sunlight

shade

Part B: Absorption of UV Light by lotions

Material tested Beads before UV Beads after UV

spray

SPF ___

SPF ___

SPF ___

Part C: Absorption of UV Light by various materials

Material tested Control beads Experiment beads

clear plastic

opaque plastic

cloth

foil

plexiglass plate

paper

UV sunglasses

cheap sunglasses

36

Part D: Home Experiments using sunscreen and other materials

Material tested Control beads Experiment beads

sunscreen

paper

Questions 1) Compare the energies of UV light to IR and visible light. Explain why UV light is

potentially more dangerous than IR or visible light. 2) Considering results from the various materials tested in Part C, if you were going to

construct a shelter to protect from UV light, which material(s) would be best? Which would be worst (and provide no protection)?

3) What is the purpose of an experimental control? Give one specific example of a control

and its experiment used in this lab.

4) Most sunscreen lotions claim to protect against UV-A and UV-B. Why don’t they mention UV-C light?

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5) From your data in Part B, did higher SPF values provide more protection from UV light? What do higher SPF values represent?

6) If the UV Index for the day is 10, what is the exposure category and advisory? Taking into account your skin pigmentation, how might this affect your plans for the day?

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Lewis Structures and Molecule Shapes

Objectives The objectives of this laboratory are to: a) Practice drawing Lewis Structures for various covalently bonded molecules and polyatomic ions. b) Use model kits to construct these molecules and polyatomic ions in order to explore their structure and

shapes. c) Practice predicting molecular shapes(using VSEPR theory. Background Nonmetal atoms join in covalent bonding, resulting in the formation of either covalent molecules or polyatomic ions. A covalent bond is formed when nonmetal atoms share their valence electrons, which they do in order to achieve a filled valence shell, like their nearest noble gas (Family 8A). This means that most bonded nonmetal atoms will acquire a total of eight valence electrons via the sharing process – often referred to as the octet rule. A notable exception is hydrogen, which only needs to acquire two electrons to be like its nearest noble gas, helium. Hydrogen thus has the duet rule, needing only two electrons to achieve stability. Lewis Structures A Lewis Structure is a representation of a covalent molecule (or polyatomic ion) where all the valence electrons are shown distributed around the bonded atoms as either shared electron pairs (covalent bonds, or bonding pairs) or unshared electron pairs (lone pairs). Total the number of valence electrons that each atom contributes. • Use the family/group number for each atom. • For ions, add electrons if negative or subtract electrons if positive.

Choose a central atom and connect outer atoms using bonds. The octet rule must be obeyed for all elements except hydrogen (follows a “duet” rule). • Each bond in contains two electrons. • Starting with the outer atoms, place electrons around every atom until it has a total of 8 electrons. • No atom should have more than 8 electrons. No atom should have less than 8 electrons (except

hydrogen). • If too many electrons were used, convert single bonds to double or triple bonds. Erase two pairs of

electrons from two atoms, then connect those two atoms with the additional bond.

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A shared pair of electrons is represented as a short line (single bond). Sometimes atoms can share two pairs of electrons, represented by two short lines (double bond). Atoms can share three pairs of electrons, represented by three short lines (triple bond). Pairs of electrons are used to represent lone pair electrons.

Examples:

VSEPR Theory

The shape of a molecule depends on the distribution of atoms in space around the central atom, and their bond angles. Bonding electrons and lone pair electrons repel one another. The bonds will be arranged around a central atom as far apart as possible in order to minimize repulsions. Lone pairs take up slightly more space, but do not appear in the 3D molecule shape This is known as Valence Shell Electron Pair Repulsion theory, or VSEPR theory.

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VSEPR Molecule Shapes

Shape Name 3D Model Sketch Description

Linear

Two atoms bonded. May be single, double or triple bond.

Linear

Two bonds on the central atom. No lone pairs. Outer atoms are arranged opposite to each other. Bond angles are exactly 180°

Trigonal Planar

Three bonds on the central atom. No lone pairs. Central and outer atoms all lie in the same plane (flat). Bond angles are exactly 120°

Tetrahedral

Four bonds on the central atom. No lone pairs. Four outer atoms are evenly arranged in 3D around the central atom as if at the corners of a regular tetrahedron. Bond angles are exactly 109.5°

Bent

Two bonds and one lone pair on the central atom. Overall V-shaped. Bond angles are slightly less than 120°

Bent

Two bonds and two lone pairs on the central atom. Overall V-shaped. Bond angles are slightly less than 109.5°

Trigonal Pyramid

Three bonds and one lone pair on the central atom. Central atom slightly above the three outer atoms, like a tripod. Bond angles are slightly less than 109.5°

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Resonance When an equivalent Lewis structure can be drawn, usually due to swapping the position of a double bond with a single bond, these are called resonance structures.

The model kit simulator is used to construct 3D molecule shapes. Each atom is represented by a color: C = black, H = white, O = red, N = blue, S = yellow, Br/Cl/F = green A “bond” to the central atom can consist of a single, double or triple bond when considering 3D

molecule shapes. Ammonium, (NH4)+

Total Valence Electrons:

3-D Model Sketch

Lewis Structure

VSEPR shape name: Resonance: Yes No

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Lewis Structures and Molecule Shapes 1. H2O

Total Valence Electrons: 3-D Model Sketch

Lewis Structure


2. N2


Lewis Structure


3. CH4


Lewis Structure


4. O2


Lewis Structure


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5. CO2 Total Valence Electrons: 3-D Model Sketch

Lewis Structure


6. NH3


Lewis Structure


7. SO2


Lewis Structure


8. O3


Lewis Structure


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9. carbonate ion (CO3-2)


Lewis Structure


10. CO


Lewis Structure


11. nitrate ion (NO3-)


Lewis Structure


12. CF2Cl2 (CFC = chlorofluorocarbon)


Lewis Structure


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Questions

1) When molecules absorb IR energy, there are different results. Atoms may dissociate, a molecule may rotate, or bonds may stretch and bend. Carbon dioxide has 4 vibrations, shown below. What is the effect of each vibration on the molecule: rotation, stretching, bending, or breaking?

2) Three different modes of vibration of a water molecule are shown. What is the effect of each vibration on the molecule: rotation, stretching, bending, or breaking?

3) The molecule shape of each atmospheric gas contributes to its ability to absorb IR energy, making it a greenhouse gas. Water vapor and carbon dioxide are greenhouse gases. Using information from the two questions above, explain why water and carbon dioxide are greenhouse gases.

4) Nitrogen gas (N2) and oxygen gas (O2) compose the majority of atmospheric gases. Using the molecules shapes, explain why nitrogen and oxygen are not greenhouse gases.

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Combustion Reactions


a) Generate and collect oxygen gas via the decomposition of hydrogen peroxide reaction b) Understand how the properties of oxygen gas allows for collection by downward

displacement of water c) Investigate combustion reactions by burning different substances in oxygen gas d) Compare differences of burning different substances in air versus oxygen gas e) Write balanced chemical equations from word equations

Background Oxygen is one of the most abundant elements on this planet. Earth’s atmosphere contains 21% oxygen gas. Oxygen is also found in compounds in the earth’s crust, such as water (89%), and in mineral oxides. The human body contains 65% oxygen atoms, by mass. Oxygen occurs naturally as a diatomic molecule, O2; these molecules are a gas at most temperatures. Oxygen gas exhibits many unique physical and chemical properties. For example, oxygen is colorless and odorless, with a density greater than that of air, and a very low solubility in water. Among the unique chemical properties of oxygen are its ability to support respiration in plants and animals, and its ability to support combustion reactions. In this lab, oxygen gas will be generated as a product of the decomposition of hydrogen peroxide. A catalyst is used to speed up the rate of the decomposition reaction. A catalyst is either not consumed as a reactant in the reaction or it is regenerated as a product of the reaction. Without it, the reaction would proceed too slowly. The catalyst used in this reaction is yeast.

Decomposition of Hydrogen Peroxide: 2 H2O2 (aq catalyst 2 H2O (l) + O2 (g) hydrogen peroxide water oxygen

The oxygen gas produced will be collected in bottles by a method known as the downward displacement of water. Once collected, several tests will be performed in order to investigate the role of oxygen in a variety of combustion reactions. A combustion reaction is commonly referred to as “burning”. During a combustion reaction, oxygen reacts with the substance being burned. Note that since our atmosphere is roughly 21% oxygen, many substances readily burn in air. Both oxygen and the substance being burned (the reactants) are consumed during the combustion reaction, while new substances (the products) and energy are generated. Since energy is produced, this is an exothermic process.

Combustion Reactions: Substance being burned + Oxygen Products + Energy

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The actual products of a combustion reaction depend on what substance is burned and how much oxygen is present. In general, however, when a pure element burns in oxygen the product is called an oxide. An oxide is a compound containing both the element and oxygen chemically combined together. Some examples of element combustion are shown below. Several such reactions will be performed using the oxygen gas collected in this lab.

Combustion of an Element: Element + Oxygen Oxide of Element + Energy

C (s) + O2 (g) CO2 (g) + Energy carbon oxygen carbon dioxide

2 Hg (l) + O2 (g) 2 HgO (s) + Energy mercury oxygen mercury(II) oxide

Procedure

Safety

Exercise caution and/or use gloves when using the hydrogen peroxide (H2O2); it can cause chemical burns and skin irritation. If it comes into contact with your skin, immediately rinse with water for a minimum of fifteen minutes and notify your instructor. Also, do not look directly at the burning magnesium. In addition to being very bright, it emits harmful UV radiation that could cause damage to the retina of the eye. Personal protective equipment (PPE) required: lab coat, safety goggles, closed-toe shoes


Materials: hydrogen peroxide solution, active yeast, wooden splint, candle, sulfur, steel wool, magnesium, aluminum pellets, 6M hydrochloric acid Equipment: 250-mL Erlenmeyer flask, five wide-mouth bottles, four plexiglass plates, pneumatic trough, “stopper + thistle tube + tubing” apparatus, utility clamp, ring stand, deflagration spoon, crucible tongs, small beaker. Part A: Generating and Collecting Oxygen Gas

1) Obtain the following equipment: • 250-mL Erlenmeyer flask, small/50-mL beaker, utility clamp (your locker) • “two-hole stopper + thistle tube + glass tubing + rubber tubing” apparatus (front) • four plexiglass plates (front) • ring stand (back corner of room) • plastic or metal trough with tray (under sink) • five wide mouth glass bottles (under sink)

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2) Fill the plastic/metal trough with tap water to just above the level of the tray. Bring the filled trough to the bench where you will be working.

3) Bring four of the glass bottles to a sink. (The fifth bottle will be used later). Fill each bottle

to the brim or slightly overflowing with tap water. Then gently slide a plexiglass plate over the mouth of each bottle. Make sure that there are no air bubbles in the water. Bring the bottles to your bench.

4) While holding the plexiglass plate, invert bottle into the water in the trough. Remove the

plexiglass plate. Repeat this for all four bottles. Place the plexiglass plates aside; they will be used later.

5) In the balance room, weigh ~ 1 g active yeast into the small beaker. Add the yeast to the

Erlenmeyer flask. Use the same small beaker to measure ~25-mL of distilled water. Add the distilled water to the Erlenmeyer flask.

6) Assemble as shown in the figure at right. Make sure that • the Erlenmeyer flask is connected to the ring stand with a

utility clamp • the Erlenmeyer flask is allowed to “swish” by loosening one

of the connectors • the end of the thistle tube is completely covered with water

at the bottom of the flask 7) Once you have setup your apparatus with four bottles

submerged, Erlenmeyer reaction flask complete, with small beaker nearby, call the professor to check the setup.

8) Obtain ~30-mL of hydrogen peroxide (H2O2) in the small

beaker. Carefully add ~5-mL of H2O2 through the thistle tube. Swish the Erlenmeyer flask. Oxygen gas should be seen as bubbles generated in the flask. If at any time the rate of the reaction in the Erlenmeyer flask appears to slow down, either swish the flask and/or add another 5-mL portion of H2O2.

9) The oxygen produced will fill the inverted bottle by displacing the water. When the first

bottle is completely filled with gas, slide a plexiglass plate under the brim and invert the bottle back to upright. Place the bottle on the bench with the mouth up and do not remove the plexiglass plate.

10) Place the second bottle on the metal rack in its place and allow it to fill in a like

manner. Repeat this for the third and fourth bottles.

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Part B: Combustion Reactions Test 1: Combustion of wood

Light a wooden split, and then blow it out. While it is still glowing, place it into the empty wide-mouth bottle (air-filled only). Record your observations. Now re-light the same wooden split, and again blow it out. Quickly insert the splint into Bottle #1 (oxygen-filled) while it is still glowing. Record your observations.

Test 2: Combustion of candle wax

Place a tealight candle on a plexiglass plate and light it. Lower the empty bottle over the candle. Measure and record the number of seconds that the candle continues to burn. Then re-light the candle and lower Bottle #2 over it. Again, measure and record the number of seconds that the candle continues to burn.

Test 3: Combustion of iron in steel

Pour about 30-mL of tap water into Bottle #4 and replace the plexiglass plate quickly. Take a loose (frayed out) 2 or 3 centimeter piece of steel wool and hold it in a Bunsen burner flame for a very brief instant with your crucible tongs (it will glow red). Then immediately lower the steel wool into Bottle #4. Record your observations. Repeat with the empty bottle and record your observations.

Test 4: Combustion of sulfur Take the empty bottle and Bottle #3 to the hood your instructor directs you to. Place a small

lump of sulfur in a deflagrating spoon. Light the Bunsen burner in the hood, and heat the sulfur in the spoon. The sulfur will first melt, then burn with an almost invisible blue flame. Insert the spoon with the burning sulfur in the empty bottle and record your observations. Then insert it in Bottle #3, and again record your observations. When finished, extinguish the burning sulfur in a beaker of water.

Test 5: Combustion of magnesium Hold a 1-inch piece of magnesium metal in a Bunsen burner flame with your crucible tongs

until it ignites (in air). Record your observations, remembering not to look directly at the burning magnesium!

Test 6: Combustion of hydrogen

To a medium test tube, add aluminum pellets, followed by about 3-mL of hydrochloric acid. Bubbles should begin to appear as hydrogen gas is produced. Place the test tube in the plastic test tube rack. After 30-60 seconds have elapsed, light a wooden splint. Do not blow it out. Hold the burning splint to the mouth of test tube (where the hydrogen gas is being evolved) and record your observations.

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Name: Chem 9, Section: Experiment Date:

Combustion Reactions Experimental Data

Test 1 Observations

Glowing splint in empty bottle

Glowing splint in Bottle #1

Test 2 Observations

Burning candle in empty bottle

Burning candle in Bottle #2

Test 3 Observations

Glowing steel in empty bottle

Glowing steel in Bottle #3

Test 4 Observations

Burning sulfur in empty bottle

Burning sulfur in Bottle #4

Test 5 Observations

Burning magnesium in air

Test 6 Observations

Burning hydrogen in air

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Data Analysis and Questions 1) Write the balanced chemical equation for the reaction used to generate oxygen gas.

2) What is the catalyst used in this reaction?

What is the purpose of this catalyst?

3) In addition to oxygen, what other substance is produced by this reaction? What happens to this

substance? 4) Why did you perform each test in two separate bottles (air-filled and oxygen-filled)? 5) Are the combustion reactions of oxygen exothermic or endothermic? Support your answer with

experimental evidence from the tests you performed. 6) Consider your results for the first four tests you performed. In which bottles, air-filled or oxygen-

filled, did the combustion reactions occur more vigorously? Why? 7) Consider your Test 2 results. Although the candle burns for a longer period of time in one bottle, it

eventually goes out in both the empty bottle and Bottle #2. Why does it go out? 8) The combustion of sulfur (Test 4) was performed in the oxygen-filled bottle, but the combustion of

hydrogen (Test 6) and combustion of magnesium (Test 5) were not. Why did Tests 5 and 6 not require the oxygen-filled bottles?

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9) When an element burns in oxygen gas, the product is called an oxide. a. The wood in the splint consists mostly of carbon. The combustion of carbon produces carbon

dioxide, CO2. Write the balanced chemical equation for the combustion of carbon.

b. The combustion of sulfur produces sulfur dioxide, SO2. Write the balanced chemical equation for the

combustion of sulfur. c. The combustion of hydrogen produces water, H2O. Write the balanced chemical equation for the

combustion of hydrogen. d. Steel wool consists mostly of iron. The combustion of iron produces iron(III) oxide, Fe2O3. Write

the balanced chemical equation for the combustion of iron. e. Aluminum forms a +3 ion when it combines in an ionic compound. The combustion of aluminum

produces aluminum oxide, Al2O3. Write the balanced chemical equation for the combustion of aluminum.

f. Methane gas (CH4) is burned in natural gas-powered power plants to generate electricity. The

combustion products are water and carbon dioxide. Write the balanced chemical equation for the combustion of methane.

10) The combustion reaction of coal (a hydrocarbon) uses oxygen as a reactant and has two products. a. Of these products, one is a gas that is produced in large amounts. Give its chemical formula and

name. b. Sulfur dioxide gas, SO2, is a minor product of coal combustion. Since coal is a hydrocarbon,

composed mostly of carbon and hydrogen atoms, how is SO2 produced from the reaction? c. Where does the SO2 product end up and why is it an issue?

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Single Replacement Reactions Objectives a) Perform and observe the results of a variety of single replacement reactions, b) Become familiar with some of the observable signs of these reactions, c) Predict and identify the products formed in each of these reactions, d) Write balanced chemical equations for each single replacement reaction. e) Understand how voltaic cells generate electrical energy. Background In Part A of this lab we will examine single replacement reactions. This is one type of oxidation-reduction reaction, or redox reaction, because it occurs via a transfer of electrons. Single replacement reactions have the general form: A + BC B + AC Here, A is an element and BC is usually an aqueous ionic compound or an acid (consisting of B+ and C- aqueous ions). Element A replaces element B in the compound BC; this results in the formation of a new element B and a new ionic compound or acid, AC. If the new element B is a metal, it will appear as a metallic deposit. If it is a gas, it will appear as bubbles. An activity series of elements (next page) is often used to determine if A will displace B in a single replacement reaction. The activity series is provided on the following page. As a rule, if A has a higher activity that B, a single replacement reaction will occur. However, if A has lower activity than B, a single replacement reaction will not occur. Example 1: magnesium + aluminum chloride Since Mg is more active than Al, a single replacement reaction will occur. The predicted products are aluminum metal and aqueous magnesium chloride Balanced Equation: 3 Mg + 2 AlCl3 2 Al + 3 MgCl2

A B C B A C Example 2: aluminum + magnesium chloride Since Al is not more reactive than Mg, a replacement will not occur. Balanced Equation: Al + MgCl2 No Reaction A B C During a chemical reaction both the form and composition of matter are changed. Some of the observable signs that a chemical reaction has occurred include: metallic deposit, bubbles, temperature change, color change, precipitate formation. Note that there are other observable signs for chemical reactions, but these are most likely to be seen in this lab.

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When a substance loses electrons, it is being oxidized (or has undergone oxidation). Oxidation is a loss of electrons. When a substance gains electrons, it is been reduced (or has undergone reduction). Reduction is a gain of electrons. Remember these concepts with the mnemonic “OIL RIG”: oxidation is loss, reduction is gain Example 1 revisited: magnesium + aluminum chloride Balanced Equation: 3 Mg + 2 AlCl3 2 Al + 3 MgCl2

In the above example, magnesium is oxidized because it has lost electrons. Mg Mg+2 + 2 e-

In addition, the aluminum atom in aluminum chloride is reduced because it has gained electrons. Al+3 + 3 e- Al Electricity can be described as a flow of electrons through a wire. This form of energy is caused by the motion of electrons. A device that creates electrical current from redox reactions is called an electrochemical cell, voltaic cell, galvanic cell or battery. Batteries serve as a source of energy for flashlights, radios, as well as car motors. In Part B we will build an electrochemical cell. It will consist of two metals (called electrodes) connected by a salt bridge between individual half-cells. A salt bridge, which contains a strong electrolyte, will join the two reactions and complete the circuit. It allows for the mixing of the two solutions. The metal strip where oxidation occurs is the anode and is labeled with a negative (-) sign. The metal strip where reduction occurs is called the cathode and is labeled with a (+) sign. These symbols can be seen on everyday batteries. Electrons flow from anode to the cathode.

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Single Replacement Reaction – Example Given two reactants: zinc and sulfuric acid 1) Use activity series to compare two metals; higher metal

will replace lower metal

2) If a reaction will occur, predict ion formed 3) Write ions and use crossover rule to write chemical

formula for new ionic compound 4) Write balanced chemical equation 5) Write chemical names for products; remember stock

system/Roman number for metals with more than one charge copper (Cu +1 or +2) chromium (Cr +2 or +3) iron (Fe +2 or +3) lead (Pb+2 or +4) tin (Sn +2 or +4)

6) Observations – metallic deposit, gas bubbles or no change (no reaction) Reactants’ chemical formulas: Compare two metals, underline more reactive

metal:

If reaction will occur, predict ion formed. If no products, explain why:

Products’ chemical formulas

Products’chemical names:

Balanced Chemical Equation:

Acid names Acid chemical formulas Hydrochloric acid HCl Sulfuric acid H2SO4 Nitric acid HNO3 Acetic acid HC2H3O2

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Procedure

Safety, Materials and Equipment (seen in demo videos) Use caution with acids. These substances can burn your skin copper, zinc, magnesium, HCl, HC2H3O2, AgNO3, Pb(NO3)2, Zn(NO3)2 medium test tubes, fruit, battery kit Personal Protective Equipment (PPE): safety goggles, lab coat, closed-toe shoes

Part A: Single Replacement Reactions

For each reaction, first predict products, and write a complete, balanced equation.

Every box should be filled in except the “Observations” box. When all equations have been predicted, have them checked and signed off.

Watch the video for each reaction and record your observations for each. 1. Zinc metal and hydrochloric acid 2. Copper metal and aqueous silver nitrate 3. Copper metal and aqueous zinc nitrate 4. Zinc metal and aqueous lead (II) nitrate 5. Magnesium metal and acetic acid

Part B: Conductivities of Citrus Fruit and Salt Bridge Battery Setup: A battery kit with 2 beakers, voltmeter, 2 alligator clips, copper and zinc metal strips has been set up to test the electrical conductivity of various samples. Connect the black electrode to the zinc strip (grey) using alligator clip, and the red electrode to the copper strip (brown) using another alligator clip. Insert zinc and copper strips directly into sample. Make sure that there is a closed circuit, and no metal parts contact with any part of the circuit. A schematic of the electrochemical cell/salt bridge battery is shown at right. Experimental data: Record the voltage measurement for each on the lab report form. Citrus fruit Electrochemical cell/salt bridge battery

https://www.youtube.com/watch?v=A4XITC225uk

https://www.youtube.com/watch?v=A4XITC225uk

https://youtu.be/1NKI0gxbQZA

https://youtu.be/L8X9_UL-E3A

https://www.youtube.com/watch?v=wlB8O5zbJug

https://www.youtube.com/watch?v=eGrR201SUdI

https://youtu.be/h5mxHgbFZZ0

https://youtu.be/sz_o0ZYD7no

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Name: Chem 9, Section: Experiment Date:

Single Replacement Reactions Part A: Single Replacement Reactions 1. Zinc metal and aqueous hydrochloric acid

Reactants’ chemical formulas: Compare two metals, underline more reactive metal:





Observations:

2. Copper metal and aqueous silver nitrate






Observations:

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3. Copper metal and aqueous zinc nitrate






Observations:

4. Zinc metal and aqueous lead (II) nitrate






Observations:

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5. Magnesium metal and aqueous acetic acid






Observations:

Part B: Conductivities of Citrus Fruit and Salt Bridge Battery - Voltage Measurements 1. citrus measured voltage = __________ V 2. salt bridge battery measured voltage = __________ V 3. salt bridge battery with alligator clips switched measured voltage = __________ V

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Questions

1. Use the metal activity series to determine the ion formed when solid silver reacts in a single replacement. Solid copper? Solid tin?

2. What is the purpose of the multimeter in the electrochemical cell/salt bridge battery? 3. What is the purpose of the string in the electrochemical cell/salt bridge battery? 4. What was the resulting voltage when the alligator clips connections to the zinc and copper strips

were switched? Why? 5. What do you predict would happen to the voltage measurement if the salt bridge is removed from the

salt bridge battery? Explain. 6. The figure shows a galvanic cell. As this

cell operates, a reddish coating of copper metal begins to appear on the surface of the copper cathode.

Label the following in the figure:

oxidation half-reaction reduction half-reaction anode cathode salt bridge voltmeter

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Double Replacement Reactions

Objectives The objectives of this lab are to: a) Predict products and write balanced chemical equations for double replacement reactions b) Perform and observe the results of a variety of double replacement reactions c) Become familiar with observable signs of precipitation, neutralization and gas-forming reactions d) Identify the products formed in each of these reactions, using solubility rules for physical states

Background During a chemical reaction both the form and composition of matter are changed. Old substances are converted to new substances, which have unique physical and chemical properties of their own. Some observable signs that a chemical reaction has occurred include:

• metallic deposit appears • bubbles appear • temperature change occurs • color change occurs • precipitate (cloudy, tiny particles) appears

Note that there are other observable signs for chemical reactions, but these are most likely to be seen in this lab. Double Replacement Reactions

All double replacement reactions have the general form: AB + CD AD + CB Reactions that can be classified as double replacement include precipitation reactions, neutralization reactions and gas-forming reactions. Precipitation Reactions Here AB and CD are usually aqueous ionic compounds, consisting of aqueous ions: A+ and B-, C+ and D-. When a double replacement reaction occurs, the cations and anions switch partners, resulting in the formation of two new ionic compounds AD and CB, one of which is in the solid state. This solid product is an insoluble ionic compound called a precipitate. To determine whether a product ionic compound will be soluble or insoluble, consult the Solubility Rules. Note that if both of the predicted products are soluble, a precipitation reaction will not occur.

Example: aqueous lead (II) nitrate + aqueous sodium chloride

The predicted products are lead (II) chloride (insoluble) and sodium nitrate (soluble). Since one of the predicted products is insoluble, a precipitation reaction is will occur.

Equation: Pb(NO3)2 (aq) + 2 NaCl (aq) 2 NaNO3 (aq) + PbCl2 (pp

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Neutralization Reactions

Here AB is an acid (consisting of H+ and X- aqueous ions) and BC is a base (consisting of M+ and OH- ions). When a double replacement reaction occurs, the cations and anions switch partners, resulting in the formation of water and a new ionic compound (or salt), which is usually soluble. Neutralization reactions are exothermic, and are generally accompanied by a noticeable release of heat.

Example: sulfuric acid + aqueous lithium hydroxide The predicted products are water and lithium sulfate.

Reaction Equation: H2SO4 (aq) + 2 LiOH (aq) Li2SO4 (aq) + 2 H2O (l)

Gas-Forming Reactions In these reactions one of the products (AD or CB) after the double replacement is in the gaseous state, such as hydrogen sulfide (H2S) or ammonia (NH3). One of the products could also be carbonic acid (H2CO3) or sulfurous acid (H2SO3). Both carbonic acid and sulfurous acid are unstable and will decompose to form carbon dioxide and sulfur dioxide gases, respectively: Carbonic acid H2CO3 (aq) H2O (l) + CO2 (g) Sulfurous Acid H2SO3 (aq) H2O (l) + SO2 (g)

Example: nitric acid + aqueous sodium bicarbonate

The predicted products are carbonic acid and sodium nitrate. However carbonic acid decomposes to carbon dioxide and water:

Reaction Equation: HNO3 (aq) + NaHCO3 (aq) NaNO3 (aq) + H2CO3 (aq) decomposes

Final Equation: HNO3 (aq) + NaHCO3 (aq) NaNO3 (aq) + H2O (l) + CO2 (g)

Writing Equations for Reactions – example Aqueous nitric acid (HNO3) and aqueous sodium hydroxide Reactants’ ions with charge values: Reactants’ chemical formulas, with physical states: Products’ ions with charge values: Products’ chemical formulas, with physical states: Products’ chemical names: Balanced Chemical Equation:

Reaction Type (if “no reaction”, why not?):

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Review: Chemical Formulas and Names of Ionic Compounds Ionic compounds are formed when positive cations and negative anions are attracted to each other via strong electrostatic forces. This attraction is called an ionic bond. The basic rules for writing the chemical formulas of ionic compounds: 1. Determine the formulas and charges on the cation and anion involved in the compound. 2. Combine the ions in a ratio that results in the formation of a neutral ionic compound. The

total charge of all the positive cations must equal the total charge of all the negative anions in the compound. The numbers of each element present in the compound become subscripts in the chemical formula.

Example: Write the formula for iron (III) chloride

First identify the cation and the anion in this compound. cation = iron (III) = Fe 3+ anion = chloride = Cl - For a neutral compound, one Fe 3+ is needed for every 3 Cl - The formula of the compound is FeCl3

Example: Write the formula for magnesium phosphate.

First identify the cation and anion in this compound. cation = magnesium = Mg 2+ anion = phosphate = PO4 3- For a neutral compound, three Mg 2+ ions needed for every two PO4 3- The formula of the compound is Mg3(PO4)2

SOLUBILITY RULES for IONIC COMPOUNDS 1. Alkali metal compounds, acetates, nitrates, and ammonium compounds are all soluble. 2. Hydroxides of alkali metals and NH4+, Ca 2+, Sr 2+, and Ba 2+ are soluble. All others are

insoluble. 3. All halides (chlorides etc.) are soluble except for those containing Ag +and Pb +2. 4. Most sulfates are soluble, except for BaSO4, SrSO4, Ag2SO4, PbSO4, and CaSO4. 5. Most phosphates, carbonates, chromates and sulfides are insoluble (except those of the

alkali metals and ammonium). 6. In addition, all acids are soluble.

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Common Polyatomic Ions

(OH) - hydroxide (NH4) + ammonium (NO3) - nitrate (NO2) - nitrite (ClO) - hypochlorite (CO3) 2- carbonate (HCO3) - bicarbonate or hydrogen carbonate (C2H3O2) - acetate (SO4) 2- sulfate (SO3) 2- sulfite (PO4) 3- phosphate Common Metals that form more than one cation

chromium Cr 2+ chromium (II) Cr 3+ chromium (III)

copper Cu + copper (I) or cuprous Cu 2+ copper (II) or cupric

iron Fe 2+ iron (II) or ferrous Fe 3+ iron (III) or ferric

lead Pb 2+ lead (II) or plumbous Pb 4+ lead (IV) or plumbic

tin Sn 2+ tin (II) or stannous Sn 4+ tin (IV) or stannic

zinc only forms 2+ ion (Zn 2+) silver only forms 1+ ion (Ag +)

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Procedure

Safety Be especially cautious when using the 6M HCl, 3M H2SO4 and 6M NaOH as they can burn your skin. Also be aware that skin discoloration will result from contact with AgNO3. If you feel any tingling sensations or see any color changes on your skin, flush with water immediately for a minimum of 15 minutes. Inform your instructor of any chemical contact as soon as possible. Personal Protective Equipment (PPE) required: safety goggles, lab coat, closed-toe shoes

Materials and Equipment Solids: solid sodium bicarbonate Solutions: 6 M sodium hydroxide, 3 M sulfuric acid, 6 M hydrochloric acid; all other solutions are 0.1 M and include silver nitrate, sodium chloride, iron (III) chloride, ammonium hydroxide, sodium carbonate, cobalt (II) nitrate, sodium phosphate, copper (II) sulfate, potassium nitrate, barium chloride Equipment: 8 test tubes, test tube rack or large beaker Instructions for Experiments For each reaction, first predict products, and write a complete, balanced equation. Every box should be filled in except the “Observations” box. When all equations have been predicted, have them checked and signed off. Watch the video for each reaction and record your observations for each. 1. Barium chloride and sodium sulfate 2. Potassium chloride and silver nitrate 3. Sodium phosphate and copper (II) sulfate 4. Nickel (II) nitrate and sodium hydroxide 5. Ammonium hydroxide and iron (III) chloride 6. Hydrochloric acid and sodium hydroxide and additional temperature measurements 7. Sodium carbonate and cobalt (II) chloride 8. Sodium nitrate and potassium chloride

At-Home Experiment

For reaction 9 (aqueous acetic acid and solid sodium bicarbonate), you can do this at home. Carefully pour into it about 20 mL vinegar into the 150-mL beaker. Note the appearance of the vinegar before the reaction. Carefully add one teaspoon of baking soda. Include any temperature change in your observations.

http://youtu.be/xLJa3sidewk?hd=1

http://youtu.be/RSoBtMKSPVs?hd=1

http://youtu.be/MZdtHq_OnzE?hd=1

http://youtu.be/HZAti9A27gs?hd=1

https://youtu.be/E9s1FVsGimg

https://youtu.be/1XyjlQF6bVs

https://youtu.be/Y6uwoFjSvIM?t=94

https://www.youtube.com/watch?v=lEdAkBNNqSY

https://youtu.be/XIylL1OLyCQ

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Double Replacement Reactions 1. Aqueous barium chloride and aqueous sodium sulfate Reactants’ ions with charge values: Reactants’ chemical formulas, with physical states: Products’ ions with charge values: Products’ chemical formulas, with physical states: Products’chemical names: Balanced Chemical Equation: Reaction Type (if “no reaction”, why not?): Observations: What compound (chemical formula) caused the

chemical change observed?

2. Aqueous potassium chloride and aqueous silver nitrate Reactants’ ions with charge values: Reactants’ chemical formulas, with physical states: Products’ ions with charge values: Products’ chemical formulas, with physical states: Products’chemical names: Balanced Chemical Equation: Reaction Type (if “no reaction”, why not?): Observations: What compound (chemical formula) caused the


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3. Aqueous sodium phosphate and aqueous copper (II) sulfate

Reactants’ ions with charge values: Reactants’ chemical formulas, with physical states: Products’ ions with charge values: Products’ chemical formulas, with physical states: Products’chemical names: Balanced Chemical Equation: Reaction Type (if “no reaction”, why not?): Observations: What compound (chemical formula) caused the


4. Aqueous nickel (II) nitrate and aqueous sodium hydroxide



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5. Aqueous ammonium hydroxide and aqueous iron (III) chloride



6. Aqueous hydrochloric acid and aqueous sodium hydroxide Reactants’ ions with charge values: Reactants’ chemical formulas, with physical states: Products’ ions with charge values: Products’ chemical formulas, with physical states: Products’chemical names: Balanced Chemical Equation: Reaction Type (if “no reaction”, why not?): Observations: What compound (chemical formula) caused the


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7. Aqueous sodium carbonate and aqueous cobalt (II) chloride



8. Aqueous sodium nitrate and aqueous potassium chloride



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9. Aqueous acetic acid and solid sodium bicarbonate



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Describe each of the following compounds as ionic or covalent. Use IUPAC nomenclature rules to write the chemical formula or the chemical name. Use solubility rules for ionic compounds to determine if aqueous or solid

Chemical Name Description ionic or covalent compound

Chemical Formula If ionic compound, aq or s/ppt

1. Aluminum sulfide

2. Calcium carbonate

3. Tin(IV) nitrate

4. Silver sulfate

5. Ethane

6. Carbon tetrachloride

7. CH4

8. P2O5

9. NO2

10. Zn(NO2)2

11. PbCl2

12. AlPO4

13. CuOH

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Conductivity of Aqueous Solutions Objectives The objectives of this laboratory are: a) To observe electrical conductivity of substances in various aqueous solutions b) To determine of the solution is a strong or weak electrolyte c) To write complete ionic and net ionic equations. Background Electrical conductivity is based on the flow of electrons. Metals are good conductors of electricity because they allow electrons to flow through the entire piece of material. Thus, electrons flow like a “sea of electrons” through metals. In comparison, distilled water is a very poor conductor of electricity since very little electricity flows through water. Highly ionized substances are strong electrolytes. Strong acids and soluble salts are strong electrolytes because they completely ionize (dissociate or separate) in solution. The ions carry the electric charge through the solution thus creating an electric current. The current, if sufficient enough, will light one or both LEDs on a conductivity meter.

Strong Electrolytes: Strong Acids, Strong Bases, and Soluble Salts

Hydrochloric acid HCl (aq)

Nitric acid HNO3 (aq)

Sulfuric acid H2SO4 (aq)

Potassium hydroxide KOH (aq)

Calcium hydroxide Ca(OH)2 (aq)

Sodium chloride NaCl (aq)

Potassium carbonate K2CO3 (aq)

Copper(II) sulfate CuSO4 (aq)

Molecular view of a strong electrolyte in aqueous solution:

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Slightly ionized substances are weak electrolytes. Weak acids and bases would be categorized as weak electrolytes because they do not completely dissociate in solution.

Weak Electrolytes: Weak Acids, Weak Bases, Slightly Soluble Salts

Hydrofluoric acid HF (aq)

Carbonic acid H2CO3 (aq)

Acetic acid HC2H3O2 (aq)

Ammonium hydroxide NH4OH (aq)

Magnesium hydroxide Mg(OH)2 (aq)

Silver chloride AgCl (s)

Calcium carbonate CaCO3 (s)

Molecular view of a weak electrolyte in aqueous solution:

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Substances that do not conduct an electric current are called non-electrolytes. Non-electrolytes do not ionize; they do not contain moveable ions. The LEDs on a conductivity meter will not light because there are no ions to carry the electric current.

Non-Electrolytes

Molecular views of non-electrolytes: polar compound in aqueous solution insoluble ionic compound distilled water or solid ionic compound

Distilled water H2O (l)

Methanol CH3OH (aq)

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Procedure Part A - Conductivity Apparatus If the solution lights the bulb, it contains electrolytes. This flow of charged particles completes the circuit, causing the bulb to light.

Part B - Conductivity of Solids (and their Solutions) As you watch this video, record the experimental data for Part B in the lab report table:

• physical state of the sample = solid, liquid, or aqueous solution • conductivity = high, low, or none • electrolyte status = strong, weak, or non-electrolyte • ionization = completely ionized, partially ionized, or not ionized

Part C - Conductivity of Aqueous Solutions

As you watch this video, report your experimental data for Part C in the lab report table: • physical state of the sample = solid, liquid, or aqueous solution • conductivity = high, low, or none • electrolyte status = strong, weak, or non-electrolyte • ionization = completely ionized, partially ionized, or not ionized

https://www.youtube.com/watch?v=4FLFy3mrjD4&feature=youtu.be

https://www.youtube.com/watch?v=4WillWjxRWw

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Name: Chem 9, Section:

Conductivity of Aqueous Solutions Part B - Conductivity of Solids (and their Solutions) As you watch this video, record the experimental data for Part B in the table below:

• physical state of the sample = solid, liquid, or aqueous solution • conductivity = high, low, or none • electrolyte status = strong, weak, or non-electrolyte • ionization = completely ionized, partially ionized, or not ionized

Original sample Physical state Conductivity Electrolyte status Ionization distilled water, H2O (l)

aluminum, Al

sucrose, C12H22O11

sodium chloride, NaCl

silicon dioxide, SiO2

Sample after water added Physical state Conductivity Electrolyte status Ionization aluminum, Al

sucrose, C12H22O11

sodium chloride, NaCl

silicon dioxide, SiO2

Original sample Physical state Conductivity Electrolyte status Ionization polyethylene, (CH2)n

ethanol, C2H6O

copper, Cu

calcium chloride, CaCl2

copper (II) sulfate, CuSO4

Sample after water added Physical state Conductivity Electrolyte status Ionization polyethylene, (CH2)n

ethanol, C2H6O

copper, Cu

calcium chloride, CaCl2

copper (II) sulfate, CuSO4

https://www.youtube.com/watch?v=4FLFy3mrjD4&feature=youtu.be

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Part C - Conductivity of Aqueous Solutions

As you watch this video, report your experimental data for Part C in the table below: • physical state of the sample = solid, liquid, or aqueous solution • conductivity = high, low, or none • electrolyte status = strong, weak, or non-electrolyte • ionization = completely ionized, partially ionized, or not ionized

Sample Physical state Conductivity Electrolyte status Ionization

tap water

distilled water

hydrochloric acid (HCl)

sodium hydroxide (NaOH)

sugar

vinegar

ethanol (EtOH)

barium sulfate (BaSO4)

https://www.youtube.com/watch?v=4WillWjxRWw

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Questions 1. In Part B, which solid samples conducted electricity? Why?

2. In Part B, after addition of water, which samples conducted electricity? Explain why the solution was able to conduct electricity.

3. From Part C, explain the difference in the conductivities of tap water compared to distilled water. 4. How can an electrolyte experiment be used to distinguish between a strong acid (like HCl) and a

weak acid (like vinegar)?

5. Write each compound as it exists in aqueous solution e.g. NaCl (aq) Na+ (aq) + Cl- (aq)

CuSO4 (aq) – soluble salt HNO3 (aq) –strong acid HC2H3O2 (aq) – weak acid Mg(OH)2 (aq) – weak base

6. For the chemical reaction HBr (aq) + NaOH (aq) NaBr (aq) + H2O (l) Write the complete ionic equation: Write the net ionic equation:

7. For the chemical reaction H2SO4 (aq) + 2 KOH (aq) 2 H2O (l) + K2SO4 (aq)

Write the complete ionic equation: Write the net ionic equation:

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Acids, Bases and pH Objectives

a) Set up and show how to use a pH indicator b) Determine the pH of common solutions c) Understand pH differences of acids and bases d) Learn to use a laboratory pH meter e) Understand relationship between pH and H+ ion concentration

Background A pH value is a number, usually between 0 and 14, that represents the acidity or basicity of a solution. The pH is always written with a lowercase “p” and an uppercase “H”, which stands for “power of hydrogen.” pH values are related to hydrogen ion (H+) concentrations. The mathematical relationship between pH and H+ is described by the equation pH = - log [H+] There is an inverse relationship between pH and H+ ion concentration (in brackets, expressed in units of molarity, M). As the H+ concentration decreases, the pH value increases, and vice versa. When the pH value is a whole number (e.g. pH 7), the number is equal to the negative exponent of the H+ ion concentration.

pH value = X [H+] = 10-X M The mathematical relationship between H+ and (OH)- concentrations is described by the equation

[H+][OH-] = 10-14 The pH values of everyday chemicals typically range from pH 0 to pH 14. Values between 0 and 7 indicate an acidic solution. Values between 7 and 14 indicate a basic solution. A pH of exactly 7 indicates that a solution is neutral, neither acidic or basic. Pure water is usually pH 7. The pH scale is shown below.

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14

acidic neutral basic

The lower the pH value, the more acidic the solution; the higher the pH value, the more basic the solution. Basic solutions are also called alkaline solutions. It should be noted that the pH scale does extend beyond 0 and 14. Strong laboratory acids typically have pH values less than 0 (negative pH values) and strong laboratory bases typically have pH values greater than 14. Thus, they are considerably more dangerous. The concept of pH is widely used in all areas of science including agriculture, biology, engineering and medicine. Many commercial products use pH as an advertisement tool, such as shampoo and water; more recently, food and drink of certain pH has been touted as more healthful.

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A pH indicator is a substance that, when a small amount of it is added to a solution of unknown pH, will change its color. This is a way to determine pH of a solution visually. The indicator used in this lab will be obtained from a natural source, red cabbage. Cabbage indicator yields a particular color depending on the pH of the solution. pH indicators are a good way to easily and quickly show the approximate pH by color when compared to a standard. An everyday example where a pH indicator is used is for testing a water sample from a swimming pool. While pH indicators are useful for qualitative purposes, when an exact quantitative value is needed, a pH meter is used. A laboratory pH meter typically has a special probe capped with a membrane that is sensitive to H+ ion concentrations. The meter reading shows an exact pH value of the solution probed. pH meters are used to measure pH values of water samples, such as determining acidity of rainwater samples. Rain water is contains dissolved carbon dioxide that produces a weakly acidic solution. Rain naturally has a pH between 5 and 6. The pH of rain in parts of the U.S. is less than pH 5, which is harmful to aquatic life and human health. This is acid rain. Living organisms are very sensitive to the effects of acids and bases in their environment. An excess of H+ or OH- can interfere with the functioning of biological molecules, especially proteins. Thus, in order to maintain homeostasis and survive, organisms must maintain a stable internal pH.

A buffer is a solution whose pH resists change on addition of small amounts of either an acid or a base. To be a good buffer, a solution should have a component that acts as a base (takes H+

out of solution) and a component that acts as an acid (puts more H+ into solution when there is an excess of OH-).

The buffering capacity of a solution is tested by adding small amounts of acid (for example, HCl) and base (for example, NaOH) and checking the pH after each addition. If the pH changes only slightly, the solution is a good buffer. Eventually its buffering capacity will be exhausted, however, and the pH will change dramatically.

Procedure

Safety All lab procedures should be conducted during lab session, under guidance of instructor; instructor to provide supervision and assistance. Use lab equipment as directed in procedure, no other purposes. Use caution when handling glassware. Return all equipment to storage, safely away from children and pets, when not in use. Personal protective equipment (PPE): safety goggles or eyewear

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Materials and Equipment Lab equipment used by instructor in demo video: beaker, dropper pipet, large test tubes, test tube rack, stirring rod, wash bottle with distilled water, laboratory pH meter, 0.1 M acetic acid, 0.1 M NaC2H3O2, 0.1 M acetic acid (HC2H3O2), 0.1 M hydrochloric acid (HCl), 0.1 M sodium hydroxide (NaOH), pH standards Experimental Procedure Part A: Preparing pH indicator and pH standards 1. Tear a few leaves of red cabbage into small pieces and place the leaves into a 500-mL

beaker. Add about 500-mL of distilled water to this beaker. 2. Gently boil the mixture on heating plate until it appears dark purple in color (5-10 min). Turn

off the heat and allow to cool (5 min). 3. Add cabbage indicator solution to pH standard solutions, labeled 1-12. Record the colors of

the pH standards. Part B: Qualitative Analysis for pH Values of Everyday Chemicals 1. Obtain a test tube rack with 9 large test tubes. Pour about 3-mL of each solution into each

test tube. 3. Using a dropper pipet, add two full squirts of cabbage indicator solution. 4. Record the resulting color of the sample after mixed with the cabbage indicator. Compare

this color with pH standards to determine the pH of the sample. The color may be between the pH standard colors (e.g. green-blue instead of green or blue

alone). For these, record the pH to 0.5 values (e.g. pH = 9.5 instead of 9 or 10). Part C: Quantitative Analysis for pH Values of Everyday Chemicals

1. Prepare the pH meter for measurements: remove probe from the storage bottle and thoroughly rinse the lower section of the probe with distilled water/wash bottle.

2. Use the same nine test tubes containing samples from Part B. Insert the pH probe directly

into each test tube.

3. Allow pH meter to equilibrate; the pH values will change then fluctuate until only 0.01-0.05 pH differences are noticed. Record the pH value (to 0.01 pH) shown on the pH meter screen.

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At-Home Experiments for Parts A & B

Kitchen and lab materials needed for at-home experiment: blender, 150-mL beaker or cup, large test tubes or smaller cups, red cabbage, distilled water Also clear or non-colored aqueous solutions to test. Examples: baking soda, sparking/soda water, vinegar, tap water, bleach, shampoo, liquid soap, lemon juice.

1. Tear purple cabbage leaves and place into blender. Add about 1 cup of water. Puree until the water is purple, similar to the cabbage juice indicator as demonstrated in the video in Part A.

2. Pour some of each of the solutions to be tested into the test tubes or small cups. It is

important that the tested chemical is clear so that you can observe the color change when the indicator is added.

3. Pour some of the cabbage indicator into each test tube.

4. Record the resulting color of the sample after mixed with the cabbage indicator. Compare

this color with pH standards colors: The color may be between the pH standard colors (e.g. green-blue instead of green or blue

alone). For these, record the pH to 0.5 values (e.g. pH = 9.5 instead of 9 or 10).

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Name: Chem 9, Section: Experiment Date ____________________

Acids, Bases and pH

Experimental Data Part A: Color of Red Cabbage Indicator with pH standards

pH standard Color with Cabbage Indicator

1

2

3

4

5

6

7

8

9

10

11

12

Parts B and C: pH of Everyday Chemicals

Chemical Color with Indicator

Qualitative pH (to 0.5)

Acidic, Basic, or Neutral?

Quantitative pH (to 0.01)

ammonia cleaner

bleach

bottled water

shampoo

soap

soda

vinegar

0.1 M HCl

0.1 M NaOH

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Questions 1) What is an acidic solution? What is a basic solution?

2) What is a buffer solution? Why is a buffer important in blood?

3) What is a pH indicator? What are common uses of pH indicators?

4) Write the mathematical equation that relates OH- and H+ ion concentration:

Circle correct choice:

Acids have (high OR low) pH, and (high OR low) H+ ion concentration.

Bases have (high OR low) pH, and (high OR low) H+ ion concentration. 5) When the H+ ion concentration is expressed in brackets [H+], what are the units of the for H+

ion concentration?

6) Does a solution with pH 4 have equal, less or more H+ ions than of a solution with a pH 6?

Calculate the [H+] and [OH-] for both solutions, include units in your answer:

For pH 4, [H+] = _______ [OH-] = ________

For pH 6, [H+] = _______ [OH-] = ________

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7) The two methods of determining pH values (pH indicator versus pH meter) should show similar pH values for those solutions. What was different?

8) Explain why rain is naturally acidic, but not all rain is classified as acid rain.

9) Here are examples of what an individual might do to reduce acid rain. For each, explain the

source of production of acid rain.

a. air-drying your hair instead of using a blow dryer

b. biking instead of driving to school c. completely turning off your cell phone between uses

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At-home experiments: Results

pH of Everyday Chemicals

Chemical Color with Indicator

Qualitative pH (to 0.5)

Acidic, Basic, or Neutral?

Did any of these results surprise you?

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Synthetic Polymers and Plastics

Objectives The objectives of this laboratory are to: a) Understand properties and uses of synthetic thermoplastics b) Compare the physical properties of “Big Six” plastics c) Identify everyday plastics by their physical properties d) Draw basic structures of polymers when given monomer structure

Background The word polymer means “many units”. A polymer can be made up of many repeating units, which are small monomer molecules that have been covalently bonded. The figure at right (from Chemistry in Context) shows a single monomer, and a polymer made of identical monomers linked together. A polymer can contain hundreds of monomers, totaling thousands of atoms. Examples of naturally-occurring polymers are silk, cotton, wood, cotton, starch, natural rubber, skin, hair and DNA. In the early 1900s, chemists began to replicate natural polymers, and create synthetic polymers, beginning with nylon which mimics silk in its strength and flexibility. Currently, more than 60,000 synthetic polymers are manufactured for industrial and commercial purposes. Roughly 75% of these are these common plastics, categorized as “The Big Six”.

No. Name Abbreviation Uses 1 polyethylene terephthalate PET lightweight, clear bottles and containers; soda and

water bottles 2 high-density polyethylene HDPE stronger, opaque bottles and containers, buckets,

crates, furniture 3 polyvinyl chloride PVC rigid form: pipes, credit cards; durable plastics

softened form: tubing, electrician’s tape

4 low-density polyethylene LDPE shopping/grocery bags, plastic wrap, bubble wrap 5 polypropylene PP yogurt containers, food containers 6 polystyrene PS expandable form= styrofoam: insulated containers,

packing peanuts crystal form: CD cases, glass replacement

These six polymers are thermoplastic: they can be melted and reshaped, or recycled. The numbers are used to ease identification of the plastics, so that they can be separated for recycling.

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The Big Six thermoplastics have these general attributes: • recyclable • insoluble in water • resistant to most chemicals • lightweight yet strong • can be shaped • can be colored with pigments • usually made from petroleum • used to make items that have no alternatives from other materials

The most common of the Big Six plastics is high-density polyethylene (HDPE). It is composed of repeating units of the monomer ethylene

— (H2C-CH2)n— Monomers are linked together in an addition polymerization reaction. Each new monomer adds to one end with a covalent bond; the total number of monomers in the polymer is represented by the subscript, n. The resulting polymer is a chain of monomers linked together. The figure at right depicts one part of the polymer chain. How many monomers are present? In this experiment, you will be qualitatively analyzing plastic polymers for physical characteristics of opacity, flexibility, durability, and breakability. You’ll also analyze the density of each plastic by checking whether pellet samples float or sink in three liquids of different densities. In this lab, you will make a polymer bouncy ball using a chemical reaction between borax and glue. Glue contains the polymer polyvinyl acetate, which cross-links to itself when reacted with borax (diagram below). After cross-linking, the glue is no longer fluid, but more solid. Adding cornstarch helps to bind the molecules together so that they hold their shape.

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Procedure

Safety

materials used in this experiment should be ingested. Return all equipment to storage, safely away from children and pets, when not in use. Personal protective equipment (PPE) optional: safety goggles


Samples of Big Six plastics (marked with recycling symbols), Lab equipment and materials used in demonstration video: beaker, glue, borax, cornstarch, water, ruler Optional at-home experiment: cup, spoon, white glue, liquid laundry detergent, cornstarch, water

Experimental Procedure

Part A: Physical Characteristics of Plastic Polymers Find three different “Big Six” plastics by looking for the 1- 6 numbers/recycling symbol. You can use three, as long as they are different. Upload a photo of the three samples you used, clearly showing the recycling symbols. Use these samples to analyze physical characteristics of each type of plastic. In each group, you may find a wide variation in properties.

• Clear – yes, no, or both • Opaque – yes, no, or both • Flexibility – can it be bent? yes, no or both • Durability – hard, soft or both • Breakability – can it be cracked? yes, no or both • Recyclability – yes or no

Part B: Density Tests of Big Six Plastics

Density values for the plastics commonly found in a single-stream recycling blue bin

Plastic Approximate density (g/cm3) polypropylene (PP) 0.90 low density polyethylene (LDPE) 0.91 high density polyethylene (HDPE) 0.95 polystyrene (PS) 1.05 polyvinylchloride (PVC) 1.30 polyethylene terephthalate (PET) 1.38

When dropped into a liquid, a plastic will float or sink depending on the density of the liquid.

• If the density of the plastic is less than the liquid, the plastic will float. • If the density of the plastic is more than the liquid, the plastic will sink.

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Part C: Polymer Bouncy Balls Three different combinations of glue, water, borax and cornstarch will be used to make polymer balls. For each ball, record its composition and bounce height (in cm).

• Polymer Bouncy Ball #1: glue + water + borax • Polymer Bouncy Ball #2: glue + water + cornstarch + borax • Polymer Bouncy Ball #3: glue + cornstarch + borax

Optional at-home Experiment for Part C: Polymer Bouncy Ball Measure two level spoonfuls of glue into a cup. Add 5 mL distilled water (measure with graduated cylinder). Use a plastic spoon to mix thoroughly. Add one level spoonful of liquid laundry detergent. Use the original plastic spoon from the glue to stir throughoughly. Once the mixture becomes impossible to stir, take it out of the cup and mold the ball with your hands. The ball will start out sticky and messy, but will solidify as it is kneaded.

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Synthetic Polymers and Plastics Part A: Physical Characteristics Remember to include a photo of the three samples you used, clearing showing the recycling symbols.

Plastic number

Short Name

Clear

Opaque

Flexibility

Durability

Breakability

Recyclable

Part B: Density Tests Use the densities of the plastics and the liquids to predict if the samples will float or sink. Indicate in each box: float or sink

Plastic sample

Approximate density (g/cm3)

Methanol D = 0.79 g/cm3

1:1 ethanol-water D = 0.94 g/cm3

Water D = 1.00 g/cm3

10% NaCl solution D = 1.08 g/cm3

PP

LDPE

HDPE

PS

PVC

PET

Relative Plastic Densities:

(lowest) _______ _______ _______ _______ _______ _______ (highest) Part C: Polymer Bouncy Balls

Polymer Ball composition Approximate bounce height (in)

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Questions

1. List the name of the city that you currently live in:___________________________

What items may be placed in blue bins where you live? (look this up, list all)

2. What happens to garbage/trash that ends up in black bins?

3. Which of the Big Six plastics would be the best material for each of the following examples? Use short names to identify each plastic (e.g. HDPE).

a take-out container for food?

a durable container for laundry detergent?

a container for hot or cold food/drink? a clear wrap for food? 4. How should each be properly disposed (blue, black, green, e-waste, hazardous, special, other):

newspaper

empty cereal box

yogurt container

mylar chips bag

pizza box with grease stains

dead alkaline battery

iPhone

plastic shopping/grocery bag

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5. An unknown plastic floats in a water but sinks in 1:1 ethanol /water solution. What is the range of possible density values this plastic may have? Suggest the composition of this plastic.

6. The figure depicts a polymer of polyvinylchloride (PVC).

Circle the original monomers and determine how many monomers are present.

What arrangement is this polymer? (head-to-tail, head-to-head-tail-to-tail, or random)

7. Polypropylene (PP) is composed of the propylene monomer. The monomer structure is shown at right.

Draw a polypropylene polymer composed of five monomers arranged in a head-to-head, tail-to-tail pattern.

8. For the polymer bouncy balls, what is the name of the monomer?

What is the role of each of the following in the formation of the solid?

glue

borax

cornstarch 9. Which ball bounced the highest? Based on your data in the table, which compound was

most likely responsible for this?

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Coronavirus and the Science of Soap

Background Soap making has remained unchanged over the centuries. The ancient Roman tradition called for mixing rain water, potash and animal tallow. Making soap was a long and arduous process. First, the fat had to be rendered (melted and filtered). Then, potash solution was added. Since water and oil do not mix, this mixture had to be continuously stirred and heated sufficiently to keep the fat melted. Slowly, a chemical reaction called saponification would take place between the fat and the hydroxide which resulted in a liquid soap. When the fat and water no longer separated, the mixture was allowed to cool. At this point salt, such as sodium chloride, was added to separate the soap from the excess water. The soap came to the top, was skimmed off, and placed in wooden molds to cure. It was aged many months to allow the reaction to run to completion. All soap is made from fats and oils, mixed with alkaline solutions. There are many kinds of fats and oils, both animal and vegetable. Fats are usually solid at room temperature, but many oils are liquid at room temperature. Liquid cooking oils originate from corn, peanuts, olives, soybeans, and many other plants. For making soap, all different types of fats and oils can be used – anything from lard to exotic tropical plant oils.

Saponification Reaction: Fat + Lye → Soap + Glycerol

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Coronavirus and the Science of Soap What is the Coronavirus? – Los Angeles Times 1. What are other names/titles for coronavirus? 2. Which virus is COVID closely related to? 3. What is the main difference between these two viruses? 4. How does the virus enter your body? 5. What cells in the body does it affect? 6. How does the virus regenerate in lung cells and then shed? The Science of Soap – Los Angeles Times

1. Will antibacterial soap kill COVID? Why not?

2. Why does the fatty, greasy layer of COVID easily adhere to skin?

3. How does soap remove virus from skin?

4. When soap cracks the COVID envelope, what happens to the virus?

https://youtu.be/2NDjXGXhngM

https://youtu.be/Xmje8wAbSlg

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Saponification Reaction: Making Soap 1. The products of the saponification reaction are glycerol and a crude soap. The chemical

formula of the soap is CH3(CH2)14COO- Na+. Draw the line-angle structure of the soap molecule.

2. On the above structure, circle the portion of the molecule that is water-soluble.

Why is this portion water-soluble? 3. On the above structure, box the portion of the molecule that is fat-soluble.

Why is this portion fat-soluble? 4. On the above structure, add interactions of the soap molecule to water molecules:

positive sodium ion to negative dipole of H2O negative carboxylic acid dipoles to positive dipoles of H2O

5. Olive oil contains oleic fatty acids. Oleic acid is an 18-carbon monounsaturated fatty acid.

The double bond is at carbon 9, and it is a cis double bond, resulting in an overall bent shape. Draw the line angle structure of oleic acid. (remember to add carboxylic acid group) What is its overall shape?

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6. From its overall shape, what is the physical state of oleic acid at room temperature? Explain.

7. Coconut oil contains lauric fatty acids. Lauric acid is a 12-carbon saturated fatty acid. Draw

the line angle structure of lauric acid. (remember to add carboxylic acid group) What is its overall shape?

8. From its overall shape, what is the physical state of lauric acid at room temperature?

Explain.

The Science of Soap – experiment in hand washing 1. Under UV light, germs on the hand are concentrated in what two areas? 2. Why are our hands and skin naturally oily?

3. Why does washing hands using only water not remove all germs/virus?

4. Which end of the soap molecule bonds to the germs/virus?

5. Which end of the soap molecule bonds to water? With the power of soap, we can each kill this virus, one nanoparticle at a time!

https://youtu.be/LSIe5guzF9w

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Name ______________________________________ Lab Section _____________

Everyday Chemistry in Your Home

1) Take a photo of a nutrition facts label from any processed food or drink. Use the calories from fat, carbohydrates and protein to calculate total calories.

Show your calculations and upload the photo with the assignment.

2) Take a photo of the label for a lotion or shampoo listing of the ingredients. Identify at least one functional group by the name ending. Add each to the table below and upload the photo with the assignment. (e.g. glycol distearate = -ate ending is an ester) Name name ending functional group

3) Look up the air quality in your neighborhood, using your ZIP code at the Purpleair website.

Report the AQI value: ___________ and AQI color: ______________

Which air pollutant is used to determine this AQI value?

Use Table 2.6 from Chemistry in Context or this AirNow AQI Levels table to the full description of the air quality condition for that AQI value.

Does this AQI value affect your plans for the day, in terms of outdoor air and activities? Explain.

4) In Chem 9, the overall goal is to connect chemistry and sustainability to everyday occurrences in your life and in the world. Has Chem 9 has changed how you view either chemistry or sustainability?

https://www2.purpleair.com/

https://www.airnow.gov/aqi/aqi-basics/

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