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Unit 5: Under Pressure

Topics/ Daily Outline: Day A Day B Day Content: CW: HW: TEXT:

1 3/16 3/13 Newton’s Laws at the Atomic Level 1, 2 -- --

2 3/18 3/17 Behavior of Gases 3 1 13.1, 14.1 3 3/20 3/19 Gas Law Demos 4 -- 14.2

4 3/24 3/23 Combined Gas Law, Ideal Gas Law 5, 6 2 14.2, 14.3 5 3/26 3/25 Dalton’s Law, Graham’s Law 7, 8 3 14.4

6 3/30 3/27 Quarterly Assessment -- 4 -- 7 4/1 3/31 Weighing by Reaction 9 -- 14.3

8 4/3 4/2 Flex Day/ Review -- -- --

9 4/6 4/7 Unit Test -- -- 14.3 Important Due Dates: Homework Assignments: Assignment A Day B Day xxxxxxxx HW 1: Gas Law Lab Design

Poster Project 3/30 3/31 HW 2: Gas Law Calculations Weighing by Reaction 4/16 4/15 HW 3: Review for Quarterly Assessment

Threaded Ions Project 4/29 4/30 HW 4: Review for Unit Test

For tutorials and additional resources:

http://www.leffellabs.com If you are absent, please use this sheet to determine what you missed and collect the materials from the make-up work bins up front. Get help from a friend, the links above, or the instructor.

Name: Class:

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Drills

Date: Outcome:

Drill: Date: Outcome:

Drill:

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Date: Outcome:

Drill:

Date: Outcome:

Drill:

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CW 1: Newton’s Laws at the Atomic Level

Directions 1. Read ALL of the directions so you don’t waste your time. 2. Go to http://mw.concord.org/modeler/. Click on “Download MW.” Choose “No thanks.

Just give me the download.” 3. Next to the “Featured Activities” header, click “Browse all activities…” 4. Click “Newton’s Laws” under the Physics section, in the atoms and molecules row. 5. Read the statement or question at the top of each page. Consider this information as

you complete each page, making sure to READ the words on the page, including directions for how to run each molecular model.

6. As you READ, define/ explain the key phrases below. 7. Answer the questions on the separate hand out as you work through the activity. Plan

on handing this in at the end of class.

Key Phrases

• Newton’s Laws (Page 1)

• Brownian Motion (Page 2)

• Velocity Vector (Page 3)

• Mass Spectroscopy (Page 5)

• Vibrational Frequency (Page 8)

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CW 2: Kinetic Molecular Theory

Assumptions of the Kinetic Molecular Theory of Gases The experimental observation about the behavior of gases can be explained with a simple theoretical model known as the kinetic molecular theory. This theory is based on the following postulates, or assumptions.

1. Gases are composed of many particles that behave like hard, spheres in a state of constant, random motion.

2. Gas particles move in a straight line until they collide with another particle or the walls of the container.

3. Gas particles are much smaller than the distance between them. In other words, most of the volume of a gas is empty space.

4. There are no attractive forces between gas particles or between gas particles and the walls of their container.

5. Collisions between gas particles or with the walls of the container are elastic, meaning no energy lost when it collides with another particle or with the walls of the container.

6. The average kinetic energy of a collection of gas particles depends on the temperature of the gas and nothing else.

A gas that follows all these assumptions is called an ideal gas. We use these assumptions because they vastly simplify the behavior of gases, even at the cost of losing some accuracy.

Modeling Kinetic Molecular Theory for Gases 1. You will work in groups to compare the behavior of gas particles to marbles. 2. In particular, focus on how the marbles are similar to and different from gas particles. 3. Each group will explore one assumption, then share their findings with the rest of class. 4. Use the butcher paper to form and summarize your ideas. 5. Fill out the table on the next page once you have been approved to share your findings.

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Assumption How are gas particles like marbles? How are gas particles unlike marbles? 1. Gases are composed of many

particles that behave like hard, spheres in a state of constant, random motion.

2. Gas particles move in a straight line until they collide with another particle or the walls of the container.

3. Gas particles are much smaller than the distance between them. In other words, most of the volume of a gas is empty space.

4. There are no attractive forces between gas particles or between gas particles and the walls of their container.

5. Collisions between gas particles or with the walls of the container are elastic, meaning no energy lost when it collides with another particle or with the walls of the container.

6. The average kinetic energy of a collection of gas particles depends on the temperature of the gas and nothing else.

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CW 3: Behavior of Gases

Learning Goals

• Design experiments to measure the relationship between pressure, volume, and temperature, and moles of an ideal gas.

• Describe these relationships graphically (using graphs) and quantitatively (using numbers).

Predictions

1. For each of the relationships below, sketch what you think the graphs would look like. Include an explanation for your sketch.

Explain your prediction:




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Experiment

2. Go to http://phet.colorado.edu/en/simulation/gas-properties and click on the picture. 3. Familiarize yourself with the controls. 4. For each of the predictions you made, write a short description of the experiment you

need to complete using the simulation to determine if your prediction was correct. 5. Summarize your findings by explaining the relationship between variables.

Variables Constant Explain your Experiment Relationship Between Variables

V and P T

V and T P

T and P V

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Gas Law Notes

Gas Law Variables Constants Direct/ Inverse?

Explanation Equation Graph

Boyle’s Law

Charles’ Law

Gay-Lussac’s Law

Avogadro’s Law

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6. A gas begins at a pressure of 1 atm and a volume of 2 L. What will the pressure of the gas be if its volume expands to 4 L?

7. A gas begins at a temperature of 300 K and a volume of 2 L. What will the volume of the gas be if its temperature drops to 150 K?

8. A gas begins at a temperature of 300 K and a pressure of 10 atm. What will the temperature of the gas be if its pressure drops to 5 atm?

9. A 1.5 mole sample of H2 is contained in a volume of 2.0 L. What will the volume be if an additional 1.5 moles of gas are added to the container?

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Using the gas laws, explain each of the following scenarios. Include the graph that can be used as evidence for your answer.

10. Explain why bicycle tires seem flatter in the winter than in summer.

11. Explain why a can of soda pop explodes if left in the hot sun.

12. A rigid container (Nalgene bottle) is filled with a gas is placed in ice. What will happen to the pressure of the gas? What do you think will happen to the volume?

13. An infected tooth forms an abscess (area of infected tissue) that fills with gas. The abscess puts pressure on the nerve of the tooth, causing a toothache. While waiting to see a dentist, the person with the toothache tried to relieve the pain by treating the infected area with moist heat. Will this treatment help? Why or why not?

Journal Write 1 On a cold winter morning, Ms. Leffel leaves her house to go to work. When she gets in her car, the TPMS (tire pressure monitoring system) indicates that the tire pressure is low. Should she call Ms. Johnson for a ride? What gas law can help her figure this out? Explain.

http://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=0ahUKEwil4JGNwdHRAhVG6iYKHf2GBjQQjRwIBw&url=http://www.vectors4all.net/vectors/japanese-map-symbol-library-clip-art&bvm=bv.144224172,d.ZGg&psig=AFQjCNEmIeIwFQ2xrxhmNxMJapY4dV4ZsQ&ust=1485028662556102

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CW 4: Gas Law Demos

Vapor Pressure Vapor pressure is a measure of the force exerted by a gas above a liquid. The vapor pressure of a liquid can be measured using a U-shaped device called a manometer. One end of the manometer is connected to a liquid sample. The other end is open to the atmosphere. If the vapor pressure of the liquid is greater than the pressure of the atmosphere at the open end, the mercury in the U is pushed up.

1. Which image shows the pressure inside the flask as being equal to the pressure of the

atmosphere? How do you know?

2. Comparing the two ethanol samples, which one shows a greater vapor pressure? How do you know?

3. If atmospheric pressure is equal to 761.0 mmHg, what is the vapor pressure of a. The ethanol at 0°C?

b. The ethanol at 20°C?

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Pressure is measured in several units. The most common in chemistry are mmHg (millimeters of mercury), kPa (kilopascals), and atm (atmospheres). You can convert between these units using the conversion factors below.

101.3 𝑘𝑃𝑎 = 1 𝑎𝑡𝑚 = 760 𝑚𝑚𝐻𝑔

4. Convert the vapor pressure of the ethanol at 0°C from mmHg to kPa.

5. Convert the vapor pressure of the ethanol at 20°C from mmHg to atm. For a liquid to become a vapor, the molecules must have enough kinetic energy to escape the attractive intermolecular forces between liquid particles. If the vapor pressure of the liquid is equal to the pressure of the atmosphere, the liquid boils.

6. What is the relationship between the vapor pressure, the temperature of the liquid, and the strength of the intermolecular forces?

7. Consider the image below. Why does water boil at 100°C at sea level, but at 70°C on top of Mount Everest?

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Boiling in a Vacuum

8. Watch the boiling in a vacuum demonstration. Explain how the water boiled despite being below its boiling point. Use the following words in your explanation: vapor pressure, intermolecular forces, atmospheric pressure.

9. The label below is from a box of rice. Explain why a longer cooking time is given for high altitude preparation.

Crush the Can

10. Watch the crush the can demonstration. Draw in the particles on the cans below, and then explain why the can was crushed. Use the following words in your explanation: liquid water, gaseous water, vacuum, atmospheric pressure.

Before heating After Heating As soon as turned upside

down in cold water

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Student in a Plastic Bag

11. Watch the student in a plastic bag demonstration. Explain why the student found it difficult to move. Use the following words in your explanation: vacuum, atmospheric pressure.

Balloon in the Erlenmeyer

12. Watch the balloon in the bottle demonstration. Explain why the balloon was forced into the bottle. Use the following words in your explanation: temperature, expanding air, partial vacuum, atmospheric pressure.

13. To remove the balloon, a straw is inserted alongside the balloon, and air is blown in. How does this make the balloon come out?

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CW 5: The Combined Gas Law

The Combined Gas Law combines Boyle’s, Charles’, and Gay-Lussac’s laws into one expression:

𝑃1𝑉1

𝑇1=

𝑃2𝑉2

𝑇2

This only holds true so long as the amount of gas in moles is constant.

1. How can the combined gas law be used to derive the others, saving you from memorizing so many equations?

2. A sample of nitrogen gas has a pressure of 6.58 kPa at 539 K. If the volume does not

change, what will the pressure (in atm) be at 211 K? (Answer: 0.0255 atm)

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3. A sample of 21 liters of a gas held at a pressure of 78 atm and a temperature of 900. K, what will the volume of the gas be if the pressure decreases to 45 atm and the temperature decreases to 750 K?

(Answer: 30 L)

4. A balloon has an internal pressure of 1.05 atm and a volume of 5.0 L. If the temperature where the balloon is released is 20.°C, what will happen to the volume when the balloon rises to an altitude where the pressure is 0.65 atm and the temperature is –15°C?

(Answer: 7.1 L)

Journal Write 2 Use a sticky note to write your own combined gas law problem. Then trade your sticky note with a partner and solve. Tape the sticky note to your journal.


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CW 6: The Ideal Gas Law

The ideal gas law relates several variables that affect gases by combining them into one expression:

𝑃𝑉 = 𝑛𝑅𝑇 R is the gas law constant, which has many possible values. The correct R value is selected by considering the units that need to cancel out to find the answer in the desired units.

Common Gas Law Constant Values

𝑅 = 8.31 𝐿 ∙ 𝑘𝑃𝑎

𝑚𝑜𝑙 ∙ 𝐾

𝑅 = 0.0821 𝐿 ∙ 𝑎𝑡𝑚

𝑚𝑜𝑙 ∙ 𝐾 𝑅 =

62.364 𝐿 ∙ 𝑚𝑚𝐻𝑔


For each of the following:

• Identify the knowns and unknowns.

• Algebraically rearrange the ideal gas law to solve for the unknown.

• Plug in the knowns and solve.

1. If I have an unknown quantity of gas at a pressure of 1.2 atm, a volume of 31 liters, and a temperature of 87°C, how many moles of gas do I have? (Answer: 1.3 mol)

Knowns

Unknowns

Algebraic Rearrangement Solve

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2. If I contain 3 moles of gas in a container with a volume of 60 liters and at a temperature of 400 K, what is the pressure inside the container in mmHg? (Answer: 1000 mmHg)

Knowns

Unknowns

Algebraic Rearrangement

Solve

3. If 4.00 moles of a gas are at a pressure of 560.0 kPa and a volume of 12.0 liters, what is

the temperature in K? (Answer: 202 K) Knowns

Unknowns

Algebraic Rearrangement

Solve

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CW 7: Dalton’s Law of Partial Pressures

1. The container below contains H2O (g), H2 (g), and O2 (g). According to Dalton’s Law of Partial Pressures, the total pressure of a mixture of gases is equal to the sum of the partial pressures of the gases.

𝑃𝑡𝑜𝑡𝑎𝑙 = 𝑃1 + 𝑃2 + 𝑃3 + ⋯

a. If the total pressure is 100 kPa, and the partial pressure of O2 and H2 are 50 kPa and 25 kPa respectively, what is the partial pressure of H2O?

b. What is the total pressure of the gases if O2 exerts 100 kPa, H2 exerts 50 kPa, and H2O exerts 200 kPa?

c. Based on the last question, which gas do you think is more common in the mixture, assuming the mixture is at a constant temperature? Why?

2. Explain the relationship between gas pressure and collisions.

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3. Heliox is a mixture of helium and oxygen gas.

a. Count the number of gas molecules in each picture, then determine the ratios of

each component to the mixture.

b. How does the ratio of helium atoms and oxygen molecules in Heliox compare to the ratio of the partial pressures?

4. Consider the gas samples below.

cccccccc

a. Which would be easier to compress and why?

b. Why are gases so much easier to compress than solids or liquids?

A B

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CW 8: Graham’s Law of Diffusion

Learning Goals

• Define effusion and explain the relationship between molar mass and effusion.

• Calculate the velocity ratio for two given gases.

Directions

1. Go to http://intro.chem.okstate.edu/Project112801/GrahamsLawEntry.html (Or Google: Graham's Law Simulation)

2. Change the volume to be about 50 mL. 3. Explain the experimental setup, the vacuum flask.

4. Effusion is the process of a confined gas escaping through a tiny hole in its container. Compare the molar masses of the gases in the table below. Make a prediction about how this will affect the time to effuse.

5. Complete the table below. Gas Molar Mass Volume Time

Hydrogen: H2

Oxygen: O2

Xenon: Xe

Krypton: Kr

Gas X ---

Gas Y ---

Gas Z ---

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6. What is the relationship between molar mass and the time for the gas to effuse?

7. Rank gas X, Y, and Z from lightest molar mass to heaviest molar mass. Use a table of elements to suggest possible elements these may be.

Demonstration

A glass tube is stuffed at both ends with cotton. On one side, HCl is added to the cotton. On the other side, NH3 is added to the cotton. When these substances react, they form a cloudy white solid.

𝑁𝐻3 (𝑔, 𝑐𝑜𝑙𝑜𝑟𝑙𝑒𝑠𝑠) + 𝐻𝐶𝑙 (𝑔, 𝑐𝑜𝑙𝑜𝑟𝑙𝑒𝑠𝑠) → 𝑁𝐻4𝐶𝑙 (𝑠, 𝑐𝑙𝑜𝑢𝑑𝑦 𝑤ℎ𝑖𝑡𝑒)

8. Make a prediction about where the white gas will form: closer to the HCl end, or closer

to the NH3 end?

9. Find the velocity ratio (𝑋𝑁𝐻3:𝐻𝐶𝑙) of NH3 to HCl. This will tell you which gas moves faster.

𝑋𝑁𝐻3:𝐻𝐶𝑙 = √𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐻𝐶𝑙

𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑁𝐻3

Journal Write 3

Observe the demonstration. Was your prediction correct? Justify your answer with the velocity ratio you calculated.


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CW 9: Weighing by Reaction

Experimental Background Magnesium and hydrochloric acid react to form hydrogen gas.

𝑀𝑔 (𝑠) + 2 𝐻𝐶𝑙 (𝑎𝑞) → 𝑀𝑔𝐶𝑙2 (𝑎𝑞) + 𝐻2 (𝑔) Given a sample of gas, the ideal gas law may be used to determine how many moles of the gas are present.

𝑛 =𝑃𝑉

𝑅𝑇

Then, using reaction stoichiometry, you can determine the mass of magnesium that must have reacted to form that number of moles of hydrogen gas.

Pre-Lab Calculations Consider the experimental set up and the data given below. Atmospheric pressure 755.16 mmHg

Volume of H2 gas collected 36.86 mL

Temperature of water bath 23.1 °C Water height difference 132.1 mmH2O

1. Which is exerting more gas pressure, the atmosphere

or the gas in the tube? Explain. There are two corrections we must make to the given pressure.

2. First, correct for the difference in pressure inside and outside the gas collection tube. a. Convert the height difference of the water in the tube (in mmH2O) into mmHg.

(HINT: 13.595 mmH2O = 1 mmHg)

b. Should this value be added to or subtracted from the given pressure? Explain.

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3. Second, we must correct for the fact that the collected gas is “wet” with water vapor. a. What is the partial pressure of water at the given temperature? (HINT: Page 34)

b. Should this value be added to or subtracted from the given pressure? Explain.

4. Based on your answers to questions 2 and 3, what is the pressure of the H2 gas?

5. Use the corrected value for pressure to solve for moles of H2 gas in the ideal gas law.

6. Use the balanced equation to determine the mass of magnesium to produce the moles of H2 (g) you found in the last question.

2 𝐻𝐶𝑙 (𝑎𝑞) + 𝑀𝑔 (𝑠) → 𝐻2 (𝑔) + 𝑀𝑔𝐶𝑙2 (𝑎𝑞)

7. If the actual mass of the magnesium was 0.0452 g, what is the percent error?

% 𝐸𝑟𝑟𝑜𝑟 = |(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)

𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒| × 100%

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Materials

• Goggles

• Apron

• 2 M HCl

• Copper wire

• Thermometer

• 1000 mL beaker

• Magnesium sample

• Metric ruler

• Eudiometer (gas collection tube)

• 1-hole stopper (#00)

• Ring stand

• Utility clamp

Procedure

1. Obtain and wear goggles. 2. Obtain a piece of magnesium ribbon. Record its identification number. 3. Obtain a copper wire cage #00 stopper. Carefully thread the magnesium ribbon into the

cage so that it won’t escape/ float away 4. Clamp the eudiometer into place, open end up, in the burette clamp. Carefully pour

about 25 mL of 2M HCl into the tube. 5. Slowly fill the rest of the eudiometer with distilled water from a squirt bottle. Try to

avoid mixing the acid with the distilled water as much as possible (let the water slowly trickle down the side of the tube).

6. Carefully lower the rubber stopper into the eudiometer, so that the copper “cage” is inserted into the water. Some water spillage is normal and should be expected. The copper cage with the magnesium should be above the 50 ml mark on the tube.

7. Fill a 1000 ml beaker about halfway full of tap water. Place this below the burette clamp.

8. Place your finger over the hole in the rubber stopper and invert the gas measuring tube, ensuring that no air goes into the tube. Lower the stoppered end of the tube into the beaker of water. Clamp the tube in place so that the stoppered end is a few centimeters above the bottom of the beaker.

9. Let the apparatus stand for about five minutes longer than it takes for the magnesium to completely react. Then, tap the sides of the eudiometer to dislodge any gas bubbles that may have become attached to the sides of the tube. Record the volume of gas in the tube.

10. Use a metric ruler to measure and record the height difference between the top water level in the eudiometer and the top water level in the beaker (in cm).

11. Record the water bath temperature. 12. Ask your teacher for the atmospheric pressure for the day (in mmHg). 13. Use your sample identification number to look up the actual mass of the magnesium. 14. Clean up: dump all water, rinse the eudiometer, and return all materials to the bin

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5: Under Pressure Name: Class: Weighing by Reaction Lab Report Rubric Partner(s): Date: To complete this lab, you will complete a lab report, using this sheet to guide you. Your final lab report should be typed and submitted in class on the due date. The following sections should be completed, in order, as they appear below. COVER SHEET: A cover sheet with:

• Title of the lab

• Your name

• Your partners’ names

• Due date of the lab report

• Class period

PURPOSE: What are we trying to determine/ do in this experiment? PROCEDURE: A step by step procedure for setting up the experiment and collecting data over the course of the experiment. Directions should be numbered; and read something like a recicpe. Underline any materials you will need once you have written the procedure. DATA: Organize ALL data into a neat data table. This means you will need in depth observations. Things to observe:

• Measurements needed for calculations

• Observations

• Possible sources of error

• Your Mg sample number CONCLUSION: Answer the following questions using complete sentences. You should use at least 5 sentences for each question. For calculations, show all work, including units. This is worth the most points; therefore you should put A LOT of effort into this part.

1. You will have to make two corrections to the atmospheric pressure. a. First, convert the ‘cm of water’ height difference to ‘mm of Hg’. If the water level

was higher in the tube, pressure inside was lower than atmospheric pressure, so you have to subtract your correction factor. If the water level was lower in the tube, pressure inside was higher than atmospheric, so you have to add your correction.

b. Second, you collected the hydrogen gas as it bubbled through water, meaning the hydrogen gas is ‘wet’. You must subtract out the partial pressure of the water vapor. Use the chart and the temperature of the water bath to find this value. Subtract it from the answer you got in the previous step.

2. Explain, in your own words, why the corrections from question 1 are required. 3. Use your corrected pressure to solve for the moles of hydrogen gas using the ideal gas

law. You measured the volume of the gas, and the temperature of the gas is the same as the water bath temperature.

4. Using the moles of hydrogen, calculate the mass of the magnesium sample. CON’T NEXT PAGE!

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CONCLUSION (CON’T.) 5. Using the actual value for the mass of magnesium, calculate the percent error in your

value.

% 𝐸𝑟𝑟𝑜𝑟 = |(𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑉𝑎𝑙𝑢𝑒)

𝐴𝑐𝑐𝑒𝑝𝑡𝑒𝑑 𝑉𝑎𝑙𝑢𝑒| × 100%

6. Considering your experimental yield, discuss a specific source of error that affected your calculated mass of the magnesium sample (and how it affected it).

Grading Rubric

Cover Page

• Title of the lab

• Your name

• Your partners’ names

• Due date of the lab report

• Class period

/2.5

Purpose

• Clear statement about what we set out to do with this lab

/2.5

Procedure

• Step by step, repeatable, clear

• Materials are underlined in procedure

• Complete, no steps skipped or assumed

/5

Data

• Table is neat, organized, readable, complete

• Reflects student understanding of lab concepts and practices

• Quality of observations

/10

Conclusion

• Question 1: 5 points






/30

TOTAL:

/50

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Poster Project 2020 (Individual)

Introduction and Theme of the Project: According to the Environmental Protection Agency, 267.8 million tons of trash were generated in the United States in 2017. About half of it, 139.6 million tons, went into landfills. With large population growth, managing and reducing waste will continue to pose challenges to our society. Resources are finite, and recycling materials often requires less energy and environmental damage than making them from scratch. How can we curb our waste production and what role does chemistry play in reducing waste? Directions:

1. Create a poster on the topic “Chemistry of Reducing, Reusing, and Recycling”. 2. The poster must be two-dimensional and entirely original. No computer printouts of

images. 3. The display should be on one side of a sheet of cardboard or poster board whose

dimensions are within two centimeters of 56 cm x 71 cm. 4. The poster will be graded on originality, effectiveness, how memorable it is, how

attention getting it is, and the appropriateness of the message. 5. Submit posters with your name and class period on the back. 6. Due Date: 3/30 (A) and 3/31 (B)

Hints:

• Be creative, present ideas in a unique way – a bedtime story, comic strip, board game

• Formulate/ brainstorm ideas as questions, create a poster to answer the question

• Fill and color all spaces. If you have neat handwriting, handwrite in pencil first then go over it with a sharpie. If you cannot handwrite, you may type, but take extra care to make it look good and cohesive on the poster.

Rubric: 20% Originality: concept and presentation

20% Effectiveness: conveys meaning well 20% How memorable: stands out in a good way

20% Attention grabbing: attractive, makes you want to read it

20% Appropriateness of the message: relates chemistry to the theme in a meaningful way *Note: to get a perfect score in any category, you must impress your instructor.

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Threaded Ions Project 2020 (Individual)

Introduction and Theme of the Project: A supersaturated solution refers to a solution that contains more of the dissolved

material than could be dissolved by the solvent under normal circumstances. Typically, supersaturated solutions are prepared by adding extra solute to a saturated solution at an elevated temperature or pressure. Bottling soda under high pressures creates supersaturated CO2 solutions, and rock candy forms from supersaturated sugar solutions.

During this project, your goal is to use principles of solutions and solubility to grow a single crystal of potassium aluminum sulfate (KAl(SO4)2 ∙ 12H2O), sodium aluminum sulfate (NaAl(SO4)2 ∙ 12H2O), or ammonium aluminum sulfate (NH4Al(SO4)2 ∙ 12H2O). You may purchase your own (these compounds are used in pickling, alum in the spice section) or obtain some from your instructor. Hints for growing crystals may researched online. Directions:

1. Grow a single crystal from one of the given compounds. NOTE: We want a single crystal, not a cluster of crystals.

2. The crystal must be grown on a piece of “Coats and Clark Ming Teal thread, #5760.” You can buy the thread at a craft store or get a piece from your instructor. NO OTHER THREAD WILL BE ACCEPTED.

3. The thread with the single crystal grown on it should be taped to a index card with your name and class period on the back and placed in a clear plastic sandwich bag.

4. Due Date: 4/29 (A) and 4/30 (B) Resources and Hints:

• How to Grow Crystals: https://bit.ly/1ICkFAT

• There are many procedures available online. Provide citations if you use any of these.

• Make multiple crystals at a time.

• Do not dump solutions – reuse them to make the alum go further

• You can get some thread and alum from your teacher.

• Start early, as there are several variables you will want to experiment with to get the best crystal.

Rubric:

/40% Perfection of the Crystal: Few flaws/ cracks. Sharp, square edges.

/40% Pattern of Growth: Even sided growth, highly symmetric crystal /20% Size of Crystal: Determined by mass

Note: the other categories are worth more than this one. Do not gain mass at the loss of even growth and neat edges.

NO CRYSTALS WILL BE ACCEPTED IF THEY ARE ON THE WRONG COLOR STRING!

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Reference Materials

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Partial Pressure of Water Vapor for Select Temperatures Temperature

(°C) Pressure (mmHg)

Temperature (°C)

Pressure (mmHg)

15 12.8 26 25.2

16 13.6 27 26.7

17 14.5 28 28.3 18 15.5 29 30.0

19 16.5 30 31.8 20 17.5 31 33.7

21 18.7 32 35.7

22 19.8 33 37.7 23 21.1 34 39.9

24 22.4 35 42.2 25 23.8 36 44.6

Polyatomic Ions Ion Name Ion Name

H3O+ hydronium CrO42– chromate

Hg22+ dimercury(I) Cr2O7

2– dichromate

NH4+ ammonium MnO4

– permanganate

C2H3O2–

CH3COO– acetate

NO2– nitrite

NO3– nitrate

C2O42– oxalate O2

2– peroxide CO3

2– carbonate OH– hydroxide

HCO3–

hydrogen (bi)carbonate

CN– cyanide

PO43– Phosphate SCN– thiocyanate

ClO– hypochlorite SO32– sulfite

ClO2– chlorite SO4

2– sulfate

ClO3– chlorate

HSO4

– hydrogen sulfate

ClO4– perchlorate S2O3

2– thiosulfate

Common Gas Law Constant Values

𝑅 = 8.31 𝐿 ∙ 𝑘𝑃𝑎


𝑅 =

0.0821 𝐿 ∙ 𝑎𝑡𝑚

𝑚𝑜𝑙 ∙ 𝐾 𝑅 =

62.364 𝐿 ∙ 𝑚𝑚𝐻𝑔


Pressure Conversion Factors

1 atm = 760 mmHg = 101.3 kPa

Temperature Conversions

𝐾 = ℃ + 273.15

℃ =5

9(℉ − 32)

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