New 25 Gas Variables-S - Weeblymatthew-austin.weebly.com/uploads/1/5/7/0/15700450/... · 2019. 1. 15. · Chemistry II: The Behavior of Gases Practice Problems I. Gas Laws 1) The

Gas Variables 1

Gas VariablesHow are the variables that describe a gas related?

Why?Imagine buying a balloon bouquet at a party store. How will the helium gas in the bouquet behave if you carry it outside on a hot summer day? How will it behave if you carry it outside during a snowstorm? What happens if the balloons are made of latex, which can stretch? What happens if the balloons are made of Mylar®, which cannot stretch? What if you add just a small amount of gas to each balloon? What if you add a lot of gas? In this activity, you will explore four variables that quantify gases—pressure (P), volume (V), temperature (T), and moles (n) of gas. These four variables can be related mathematically so that predictions about gas behavior can be made.

Model 1 – Gases in a Nonflexible ContainerExperiment A (Adding more gas)

A1 A2 A3Volume = 1 unit Volume = 1 unit Volume = 1 unit

External pressure = 1 atm External pressure = 1 atm External pressure = 1 atmInternal pressure = 1 atm Internal pressure = 2 atm Internal pressure = 3 atm

Temperature = 200 K Temperature = 200 K Temperature = 200 K

Experiment B (Heating the gas)

B1 B2 B3Volume = 1 unit Volume = 1 unit Volume = 1 unit



*Note: Volume in this model is recorded in units rather than liters because 4 molecules of gas at the conditions given would oc-cupy a very small space (~1 × 10–22 μL). The particles shown here are much larger compared to the space between them than actual gas particles.

2 POGIL™ Activities for High School Chemistry

1. In Model 1, what does a dot represent?

2. Name two materials that the containers in Model 1 could be made from that would ensure that they were “nonflexible?”

3. In Model 1, the length of the arrows represents the average kinetic energy of the molecules in that sample. Which gas variable (P

internal, V, T or n) is most closely related to the length of the

arrows in Model 1?

4. Complete the following table for the two experiments in Model 1.

Experiment A Experiment B

Independent Variable

Dependent Variable

Controlled Variable(s)

5. Of the variables that were controlled in both Experiment A and Experiment B in Model 1, one requires a nonflexible container. Name this variable, and explain why a nonflexible container is necessary. In your answer, consider the external and internal pressure data given in Model 1.

Read This!Pressure is caused by molecules hitting the sides of a container or other objects. The pressure changes when the molecules change how often or how hard they hit. A nonflexible container is needed if the gas sample is going to have an internal pressure that is different from the external pressure. If a flexible con-tainer is used, the internal pressure and external pressure will always be the same because they are both pushing on the sides of the container equally. If either the internal or external pressure changes, the flex-ible container walls will adjust in size until the pressures are equal again.

Gas Variables 3

6. Name the two factors related to molecular movement that infl uence the pressure of a gas.

7. Provide a molecular-level explanation for the increase in pressure observed among the fl asks of Experiment A.

8. Provide a molecular-level explanation for the increase in pressure observed among the fl asks of Experiment B.

9. Predict what would happen to the volume and internal pressure in Experiment A of Model 1 if a fl exible container were used.

10. Predict what would happen to the volume and internal pressure in Experiment B of Model 1 if a fl exible container were used.

11. For each experiment in Model 1, determine the relationship between the independent and depen-dent variables, and write an algebraic expression for the relationship using variables that relate to the experiment (P

internal, V, T or n). Use k as a proportionality constant in each equation.

Experiment A Experiment B

Direct or Inverse Proportion?

Algebraic Expression


Model 2 – Gases in a Flexible ContainerExperiment C(Adding more gas)

C1 C2 C3Volume = 1 unit Volume = 2 units Volume = 3 units



Experiment D(Heating the gas)

D1 D2 D3Volume = 1 unit Volume = 2 units Volume = 3 units



Experiment E(Reducing theexternal pressureon the gas)

E1 E2 E3Volume = 1 unit Volume = 2 units Volume = 3 units

External pressure = 1 atm External pressure = 0.50 atm External pressure = 0.33 atmInternal pressure = 1 atm Internal pressure = 0.50 atm Internal pressure = 0.33 atm


12. Consider the gas samples in Model 2.

a. Name two materials that the containers in Model 2 could be made from that would ensure that they were “flexible”?

b. What is always true for the external and internal pressures of a gas in a flexible container?

Gas Variables 5

13. Complete the following table for the three experiments in Model 2.

Experiment C Experiment D Experiment E

Independent Variable

Dependent Variable

Controlled Variable(s)

14. Provide a molecular level explanation for the increase in volume among the balloons in Experi-ment C. (How often and/or how hard are the molecules hitting the sides of the container?)

15. Provide a molecular level explanation for the increase in volume among the balloons in Experiment D.

16. Provide a molecular level explanation for the increase in volume among the balloons in Experiment E.

17. Compare Experiment A of Model 1 with Experiment C of Model 2. How are these two experiments similar and how are they different in terms of variables?

18. Compare Experiment B of Model 1 with Experiment D of Model 2. How are these two experiments similar and how are they different in terms of variables?

19. If Experiment E of Model 2 were done in a nonflexible container, would there be any change to the internal pressure of the flask when the external pressure was reduced? Explain.


20. For each experiment in Model 2, determine the relationship between the independent and dependent variables, and write an algebraic expression for the relationship using variables that relate to those in the experiment (P

internal, V, T or n). Use k as a proportionality constant

in each equation.

Constant Pressure

Experiment C Experiment D Experiment E

Direct or Inverse Proportion?

Algebraic Expression

21. The three samples of identical gas molecules below all have the same internal pressure. Rank the samples from lowest temperature to highest temperature, and add arrows of appropriate size to illustrate the average kinetic energy of the molecules in the samples.

Gas Variables 7

Extension Questions22. Draw a sample of gas that is colder than all three of the samples in Question 21. Explain why

you are sure that it is colder.

23. Four of the relationships you investigated in Models 1 and 2 are named after scientists who discovered the relationships. Use the Internet or your textbook to match each of the scientists below with the appropriate law. Write the algebraic expression that describes the law in the box below each name.

Robert Boyle Jacques Charles Guillaume Amontons Amedeo Avogadro

Read This!Chemists combine all of the relationships seen in Models 1 and 2 into one law—the Ideal Gas Law. It is one equation that describes gas behavior and the relationship among all four variables, P, V, T, and n. In the Ideal Gas Law the proportionality constant is represented by the letter R (rather than the generic k).

24. Circle the algebraic equation below that best combines all of the relationships you identified among P, V, T, and n in this activity.

P = RTnV PT = RnV PV = nRT PTV = Rn

Chemistry II: The Behavior of Gases

Practice Problems

I. Gas Laws 1) The volume of a gas at 155.0 kPa changes from 22.0 L to 10.0 L. What is the new

pressure if the temperature remains constant? (341 kPa)

2) Exactly 10.0 L of O2 at -25ºC is heated to 100.0ºC. What is the new volume if the pressure is kept constant? (15.04 L)

3) A sample of water vapor occupies a volume of 10.5 L at 200ºC and 100.0 kPa. What volume will the water vapor occupy when it is cooled to 27ºC if the pressure remains

constant? (6.66 L)

4) A gas at a pressure of 6.2 atm and a temperature of 25ºC occupies a volume of 5.2 L. When the gas is heated to 100.0ºC the volume increases to 7.00 L. What is the new

pressure? (5.76 atm)

5) A sample of O2 with an initial temperature of 50.0ºC and a volume of 105 L is cooled to -25ºC. The new pressure is 783 mmHg and the new volume is 55.0 L. What was

the initial pressure of the sample? (534.18 mmHg)

II. Ideal Gases

6) A sample of argon gas is at a pressure of 1.24 x 104 kPa and a temperature of 24ºC in a rigid 25-L tank. How many moles of argon does this tank contain? (125.6 mol)

7) A 35.0-L tank contains 7.00 mol of compressed air. If the pressure inside the tank is 4.5 atm, what is the temperature of the compressed gas? (274 K)

8) How many grams of helium does a 25.0-L balloon contain at 790 mmHg and 24ºC? (4.27 g He)

9) Calculate the volume that 2.25 mol of O2(g) will occupy at STP. (50.4 L)

Chemistry II: The Behavior of Gases

Practice Problems

10) Calcium carbonate, CaCO3, also known as limestone, can be heated to produce calcium oxide (lime), an industrial chemical with a wide variety of uses. The

balanced equation for the reaction is CaCO3(s) CaO(s) + CO2(g). How many

grams of calcium carbonate must be decomposed to produce 5.00 L of carbon dioxide

gas when reaction is carried out at 745 mmHg and 300°C? (10.41 g CaCO3)

11) Tungsten, W, a metal used in light-bulb filaments, is produced industrially by the reaction of tungsten oxide with hydrogen (WO3(s) + 3H2(g) W(s) + 3H2O(l)).

How many liters of hydrogen gas at 35°C and 0.980 atm are needed to react

completely with 875 g of tungsten oxide? (292.1 L H2)

12) Using the equation given in #11, how many grams of tungsten is produced when 40 g of WO3 reacts with 750 mL of hydrogen gas at a pressure of 0.75 atm and 400°C?

(0.625 g W)

IV. Partial Pressure and Diffusion

13) A gaseous mixture consisting of nitrogen, argon, and oxygen in a 3.5-L vessel at 25ºC. Determine the number of moles of oxygen if the total pressure is 98.5 kPa and

the partial pressures of nitrogen and argon are 22.0 kPa and 50.0 kPa, repectively.

(0.0375 mol O2)

14) Compare the diffusion rates of O2 (molar mass, 32.0 g/mol) and N2 (molar mass, 28.0 g/mol). (rate O2 = 0.177, rate N2 = 0.189, or ratio O2:N2 = 0.94:1)

Chemistry II - The Behavior of Gases

Review Assignment Name_________________________________

Name_______________________________

Match each example below with the appropriate gas property or properties it illustrates.

Letters will be used more than once. (not a matching, write all letters that apply for each

number)

_____1. the fragrance of perfume spreads a. compressibility

through the room

_____2. smog forms over Atlanta during b. diffuses through other gases

summer days

_____3. a cylinder of oxygen used c. exerts pressure

in a hospital

_____4. a balloon is inflated with helium d. fills container

Match the variables used to describe gases to the correct unit.

_____5. psi

_____6. kPa a. pressure

_____7. oC b. temperature

_____8. mL c. volume

_____9. K

____10. mmHg

____11. atmospheres (atm)

____12. L

____13. oF

Complete the following statements by writing “decreases,” “increases,” or “remains the

same” on the line provided.

As a gas is compressed in a cylinder (example--air in a syringe)

1. its mass ______________________________.

2. the number of gas molecules ____________________________.

3. its pressure ___________________________.

4. its volume __________________________.

5. the distance between gas molecules ________________________.

6. its density _________________________.

7. its temperature _________________________.



Address each of the following.

8. Boyle’s Law states that the pressure of a gas is inversely proportional to its volume. Give the correct equation that represents this statement and explain two examples of

this law from the world around us.

9. Charles’ Law states that the volume of a gas is directly proportional to its temperature. Give the correct equation that represents this statement and explain

two examples of this law from the world around us.

10. Gay-Lussac’s Law states that the pressure of a gas is directly proportional to its temperature. Give the correct equation that represents this statement and explain

two examples of this law from the world around us.

11. If I initially have a gas with a pressure of 84 kPa and a temperature of 35°C and I heat it an additional 230°C, what will the new pressure be? Assume the volume of

the container is constant. (146.7 kPa)



1. My car has an internal volume of 2600 liters. If the sun heats my car from a temperature of 20°C to a temperature of 55°C, what will the pressure inside my car

be? Assume the pressure was initially 760 mmHg. Given these values, how many

moles of gas are in my car? (850.8 mmHg, 108.15 mol)

2. A balloon filled with air has an internal pressure of 1.25 atm, a volume of 2.50 L, and a temperature of 30°C. If I take the balloon to the bottom of the ocean where

the pressure is 95 atmospheres and the temperature is 3°C, what will the new

volume of the balloon be? How many moles of gas does the balloon hold? (0.03 L,

0.126 mol)

3. A gas that has a volume of 28 liters, a temperature of 45°C, and an unknown pressure has its volume increased to 34 liters and its temperature decreased to 35°C.

If I measure the pressure after the change to be 2.0 atm, what was the original

pressure of the gas? (2.51 atm)

4. Some students believe that teachers are full of hot air. If I inhale 2.2 L of gas at a temperature of 18°C and it heats to a temperature of 38°C in my lungs, what is the

new volume of the gas? (2.35 L)

5. Calculate the diffusion rate of laughing gas, N2O. (rate = 0.15)

6. Blast furnaces give off many unpleasant and unhealthy gases. If the total air pressure is 0.99 atm, the partial pressure of carbon dioxide is 0.05 atm, and the

partial pressure of hydrogen sulfide is 0.02 atm, what is the partial pressure of the

remaining air? If the air in this case contains 22% oxygen, what is the partial

pressure of oxygen near a blast furnace? (Pair = 0.92 atm, PO2 = 0.2024 atm)



7. What volume of oxygen gas in liters can be collected at 0.987 atm pressure and 25.0°C when 30.6 g of KClO3 decomposes by heating, according to the following

balanced equation: 2KClO3(s) 2KCl(s) + 3O2(g) (9.28 L)

8. If 8.50 L of chlorine gas are used in the following reaction which occurs at 0.95 atm and 22°C, how many grams of I2 would be produced? The unbalanced reaction is

KI(aq) + Cl2(g) KCl(aq) + I2(s). (84.5 g I2)

9. How many liters of water vapor can be made from 55 grams of oxygen gas and 5.8 g of hydrogen gas at a pressure of 12.4 atm and a temperature of 85°C? (6.80 L H2O)

10. What mass of iron (III) chloride can be produced if 0.54 g of iron reacts with 450 mL of chlorine gas at 27°C and 105 kPa according the following balanced equation?

2Fe(s) + 3Cl2(g) 2FeCl3(s) (1.57 g FeCl3)


Antacids and Balloons Mini-Lab

Name_______________________________ Lab Station #__________________

Title: Carbon Dioxide from Antacid Tablets

Date:

Partners:

Objective: To measure the amount of carbon dioxide gas given off when antacid

tablets dissolve in water.

Materials:

- Tape measure

- Clock or watch

- 3 rubber balloons

- 10-mL graduated cylinder

- Water

- 6 antacid tablets

Procedure:

1. Break six antacid tablets into small pieces. Put the pieces of one tablet into the first balloon. Put the pieces from two tablets into a second balloon. Put

the pieces from three tablets into a third balloon.

2. Use the graduated cylinder to pour 5 mL of tap water into each balloon. Immediately tie off each balloon.

3. Shake the balloons to mix the contents. Allow the contents to warm to room temperature.

4. Carefully measure and note the circumference of each balloon several times during the next 20 minutes. In the data table below, record only the maximum

circumference the balloon obtained.

5. Use the maximum circumference of each balloon to calculate the volume of each balloon. Assume that the balloons are spherical. Show one sample

calculation here. (Circumference = 2r and volume of a sphere = 4r3/3)

Data Table:

Balloon Maximum Circumference Volume

1

2

3


Antacids and Balloons Mini-Lab

Conclusions:

1. Describe the relationship between the number of tablets used and the volume of the balloon.

2. Using the gas pressure sensor/thermometer to find the temperature and pressure in the classroom, calculate the number of moles and the mass of

carbon dioxide in each balloon at maximum inflation.

3. If a typical antacid tablet contains 2.0 g of sodium hydrogen carbonate, perform a percent yield calculation to show how your carbon dioxide results

compare with the theoretical values for each balloon (you have to calculate

these). The sodium hydrogen carbonate to carbon dioxide mole:mole ratio in

this reaction is 2:1.

Name _____________________________________ Date _________________ Period ____

1

Determination of the Molar Mass of Gases and Volatile Liquids

Introduction: The molar masses of compounds are used daily in the chemistry profession. The

molar mass is defined as the mass, in grams, of 1 mole of any element or compound. How is

molar mass determined and how is molar mass of an unknown found? In this experiment, the

molar masses of sample gases are determined directly and the molar masses of several volatile

liquids will be calculated based on the measurements of their vapor density.

Concepts: molar mass, ideal gas law, buoyancy

Background:

The ideal gas law relates the four measureable properties of a gas (P, V, n, T). In this experiment,

the ideal gas law will be used to determine the molar mass of gases and volatile liquids.

PV = nRT (Equation 1)

The number of moles (n) of any pure substance is equal to its mass divided by its molar mass.

n = mass/molar mass (Equation 2)

Substituting for n in Equation 1 and then rearranging produces the equation for the molar mass of

a gas.

molar mass (g/mol) = [mass (g) * RT (K)] / [P (atm) * V (L)] (Equation 3)

In part 1, the mass of several unknown gases (X) are measured and compared to the mass of the

same volume of oxygen. If two identical volumes of different gases are at the same temperature

(T) and pressure (P), then their mass ratio must be equal to their molar mass ratio (Equation 4).

[mass of X / mass of O2] = [molar mass of X / molar mass of O2] (Equation 4)

Rearranging:

molar mass of X = [mass of X / mass of O2] * molar mass of O2 (Equation 5)

When measuring the mass of a gas, the effect of buoyancy must be taking into account. Air like

water exerts a positive or upward buoyant force on all objects. This force is compensated for in

balances when massing liquids and solids when massing gases however, this force is not

compensated for and is real. The apparent mass of a gas is less than the true mass of the gas.

True mass of gas = apparent mass of gas + mass of air displaced (Equation 6)

In part 1, the mass of each gas will be measured in a 60-ml gas syringe that has been evacuated

of air and set to a fixed volume. The syringe is first massed with no gas in the fixed volume, then

with a sample gas in the fixed volume. The effect of buoyancy is thus eliminated and the true

mass of each gas, including air, can be determined directly.

In part 2, the molar masses of several volatile liquids with boiling points well below the boiling

point of water are determined. A small sample of the liquid is placed in a tared 15-ml plastic

pipet and the pipet is then heated in boiling water to vaporize the liquid. The air and excess vapor

escape, leaving the pipet filled only with the volatile liquid vapor at atmospheric pressure and at

the temperature of boiling water. The pipet is then removed and cooled to condense the vapor.

Name _____________________________________ Date _________________ Period ____

2

Once cooled, the pipet is weighed. By massing the same pipet filled with deionized water, the

volume of the pipet is calculated. The molar mass of the volatile liquid is then determined from

Equation 3 using the mass of the condensed vapor, the volume of the pipet, the atmospheric

pressure, and the temperature of the boiling water.

Experiment Overview: The purpose of this to determine the molar masses of various gases and

volatile liquids. In part 1, the gases are massed with a special gas syringe and their molar masses

are determined by comparisons to data from oxygen measurements. In part 2, liquids are

volatilized and condensed in a fixed volume. The condensed vapor is massed and the liquid’s

molar mass is calculated from the experimental data.

Pre-Lab Questions:

1. The following data represent a determination of the “molar mass” of a sample of air by

comparison with the mass of oxygen as the reference gas. Assuming the air is 79% nitrogen,

20% oxygen, and 1% argon, and that these gases act as ideal gases, calculate both the

experimental and theoretical “molar mass” of air. See equation 4 in the Background section.

Mass of evacuated syringe 40.687 g

Air O2

Mass of syringe and gas 40.741 g 40.747 g

Mass of gas 0.054 g 0.060

Mass of gas/Mass of oxygen 1.00

Experimental molar mass ---

Theoretical molar mass 32.00

(Hint: For determining the theoretical molar mass of air, assume the percentages represent the

mole fraction of each gas in the solution of air.)

Name _____________________________________ Date _________________ Period ____

3

2. A determination of the molar mass of methyl alcohol (CH3OH) yielded the following data.

Temperature of boiling water bath

Barometric pressure

Temperature of room temp. water bath

Density of water at room temp.

99.5 ˚C

738 mm Hg

24.0 ˚C

0.9973 g/mL

Trial 1

Mass of empty pipet 1.557 g

Mass of pipet and condensed methyl alcohol 1.571 g

Mass of pipet and water 16.001 g

Mass of condensed methyl alcohol

Mass of water in filled pipet

Volume of pipet

Molar mass of methyl alcohol (experimental)

Molar mass of methyl alcohol (theoretical)

Using the data, fill the rest of the table. Calculate the molar mass of methyl alcohol using

Equation 3 and compare this value to the actual molar mass of methyl alcohol. The volume of the

pipet is equal to the volume of water in the pipet. Use the relationship of mass and density to

determine this volume. Once the volume of the pipet is determined, equation 3 in the

Background section can be used to calculate the molar mass of methyl alcohol.

Materials Part 2

Acetone, CH3COCH3, 2mL

Ethyl alcohol, CH3CH2OH, 2mL

Isopropyl alcohol, (CH3)2CHOH, 2mL

Beakers, 400-mL, 2

Boiling stones

Hot plate

Plastic tubing pipet holder

Test tube clamp

Beral-type pipets, super jumbo, narrow stem, 15-mL, 4

Balance, milligram (0.001-g precision)

Barometer

Pliers

Permanent marker

Thermometer

Ring stand

Scissors

Safety Precautions

Acetone, ethyl alcohol, and isopropyl alcohol are all flammable liquids and are fire risks.

Acetone and isopropyl alcohol are slightly toxic by ingestion and inhalation. Ethyl alcohol is

made poisonous by the addition of denaturant – it cannot be made non-poisonous. If ammonia or

chlorine gases are used in Part 1, release these gases in an efficient working fume hood. Wear

chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Exercise care

when working with the hot water bath. Wash hands thoroughly with soap and water before

leaving the laboratory.

Name _____________________________________ Date _________________ Period ____

4

Procedure

Part 2: Molar Masses of Volatile Liquids

1. Place a 400-mL beaker on the hot plate and add about 300 mL of water to the beaker, along with several boiling stones. Turn on the hot plate to boil the water.

2. Obtain three 15-mL jumbo Beral-type pipets. With pliers, pull the thin stems of each so that

a very fine “capillary” tip is formed where the

stem has been pulled (Figure 1).

3. Cut the pipet as shown in Figure 1 so that the capillary tip is less than 1-cm long.

4. Label the pipets #1, #2, and #3 with a permanent marker. 5. Mass each pipet to the nearest 0.001 g and record this mass in the Part 2 Data Table. 6. Draw 2-3 mL of the ethyl alcohol from the

labeled bottle in the hood into each of the

previously prepared and labeled pipets.

7. Insert the tips of the pipets containing the ethyl alcohol into the short piece of plastic tubing,

then secure the tubing with a test tube clamp

(Figure 2).

8. Lower the pipets into the boiling water bath. Make sure the entire bulb of each pipet is

below the water line.

9. Heat for at least 5 minutes. 10. Carefully remove the pipets from the water. Inspect each pipet. If any liquid remains in a

pipet bulb, heat the entire assembly for another minute.

11. Cool the pipets by lowering the pipet assembly into a bath of room temperature water in a 400-mL beaker.

12. Record the temperature of the boiling water bath and the barometric pressure of the room in the Part 2 Data Table.

13. Dry the pipets with paper towels and mass each pipet, which now contains only the condensed vapor, to the nearest 0.001 g. Record these values in the Part 2 Data Table.

14. Fill a 250-mL beaker with room temperature deionized water. 15. Fill pipet #1 with the deionized water, then expel the water into the sink to flush the

remaining ethyl alcohol from the pipet. Repeat this process several times.

16. To determine the volume of the pipet #1: Fill the pipet completely with deionized water, dry the outside, and mass the pipet and water. Record the mass in the Part 2 Data Table.

17. Repeat step 16 for pipets #2 and #3. 18. Repeat steps 2 through 13 for acetone and isopropyl alcohol.

Name _____________________________________ Date _________________ Period ____

5

Data Tables

Part 2

Temperature of boiling water bath

Barometric pressure

Temperature of room temp. water bath

Density of water at room temperature

__________ ˚C

__________ mm Hg

__________ ˚C

__________ g/mol

Jumbo Pipets

Jumbo Pipet #1 Jumbo Pipet #2 Jumbo Pipet #3

Mass of empty pipet

Mass of pipet and water

Mass of water in filled pipet

Volume of pipet

Volatile Liquids

Trial 1 Trial 2 Trial 3

Ethyl Alcohol

Mass of pipet and condensed ethyl alcohol

Mass of condensed ethyl alcohol

Molar mass of ethyl alcohol

Acetone

Mass of pipet and condensed acetone

Mass of condensed acetone

Molar mass of acetone

Isopropyl Alcohol

Mass of pipet and condensed isopropyl alcohol

Mass of condensed isopropyl alcohol

Molar mass of isopropyl alcohol

Name _____________________________________ Date _________________ Period ____

6

Post-Lab Questions and Calculations

Part 2

1. Determine the mass of condensed, volatile vapor for each pipet trial and for each unknown in Part 2. Enter these values in the Part 2 Data Table.

2. Use the CRC Handbook of Chemistry and Physics to determine the density of water at the temperature of the room temperature water bath used in this experiment. Enter this

density value in the Part 2 Data Table. Use this value and the mass of water in each filled

pipet to calculate the volume of each pipet.

3. Determine the mass of the condensed volatile liquid for each run. Enter these values in the Part 2 Data Table.

4. Calculate the molar mass of the liquid used in each run and the average of the three runs for each volatile liquid.

5. Volatile liquids with lower boiling points often give better results then those with higher boiling points. Suggest a reason for this.

6. What effect would vapor condensation in the neck of the jumbo pipets have on the reported molar mass? How large an error might this introduce?

7. Some liquids have enough attractions between molecules to form dimers. (Dimers are molecules formed from the combination of the identical molecules, A + A A2.) What

effect would this have on the experimental molar mass?

Chemistry II Name_______________________________

Gases TEST

Part 1--Multiple Choice

1. ______As the temperature of a fixed volume of gas increases, the pressure will a. vary inversely. b. decrease. c. be unchanged. d. increase.

2. ______Reducing the volume of a contained gas by one-third, while holding the temperature constant, causes pressure to

a. be decreased by two-thirds. b. be increased by two-thirds. c. be decreased by one-third. d. be increased by one-third.

3. ______A breathing mixture used by deep-sea divers contains helium, oxygen, and carbon dioxide. What is the partial pressure of oxygen when the total pressure is

101.3 kPa and PHe = 84.0 kPa and PCO2 = 0.10 kPa?

a. 10.3 kPa b. 17.2 kPa c. 34.4 kPa d. 185.4 kPa

4. ______Increasing the volume of a given amount of gas at constant temperature causes the pressure to decrease because

a. the molecules are striking a larger area with the same force. b. there are fewer molecules. c. the molecules are moving more slowly. d. there are more molecules.

5. ______When a container is filled with 3.00 mol of H2, 2.00 mol of O2, and 1.00 mol of N2, the pressure in the container is 465 kPa. The partial pressure of O2 is

a. 78 kPa. b. 116 kPa. c. 155 kPa. d. 212 kPa.

6. ______The volume of a gas is doubled while the temperature is held constant. The pressure of the gas

a. remains unchanged. b. is reduced by one-half. c. is doubled. d. depends on the kind of gas.

7. ______Which of the following would double the pressure on a contained gas at constant temperature?

a. doubling the volume of the container b. halving the number of particles in the container c. doubling the number of particles in the container d. none of these

8. ______A gas occupies a volume of 2.50 L at a pressure of 350.0 kPa. If the temperature remains constant, what volume would the gas occupy at 1750 kPa?

a. 5.00 L b. 0.500 L c. 12.5 L d. 140 L

9. ______A gas occupies 40.0 mL at –123°C. What volume does is it occupy at 27°C, assuming pressure is constant?

a. 182 mL b. 8.80 mL c. 80.0 mL d. 20.0 mL

10. ______A gas occupies a volume of 0.2 L at 25 kPa. What volume will the gas occupy at 2.5 kPa?

a. 4 L b. 20 L c. 2 L d. 0.02 L

11. ______If the temperature of a gas in a closed container increases, a. the pressure of the gas decreases. b. the average kinetic energy of the gas molecules in the container decreases. c. the gas molecules collide with the walls of the container less frequently. d. the pressure of the gas increases.

12. ______If a balloon containing 1000 L of gas at 50°C and 101.3 kPa rises to an altitude where the pressure is 27.5 kPa and the temperature is 10°C, its volume there

is

a. 271.5 L. b. 3227.5 L. c. 309.8 L. d. 18,418 L.

13. ______The temperature of 6.24 L of a gas is increased from 25.0°C to 55.0°C at constant pressure. The new volume of the gas is

a. 13.7 L. b. 5.67 L. c. 6.87 L. d. 2.84 L.

14. ______A temperature of –25°C is equivalent to a. 248 K. b. 25 K. c. –25 K. d. 298 K.

15. ______A sample of chlorine gas has a pressure of 7.25 kPa at 20.0°C. What will its pressure be at 60.0°C if its volume remains constant?


16. ______If a sample of oxygen occupies a volume of 2.15 L at a pressure of 58.0 kPa and a temperature of 25°C, what volume would this sample occupy at 101.3 kPa and

0°C?

a. 1.35 L b. 1.13 L c. 44.8 L d. 22.4 L

17. ______The volume that would be occupied by 5.00 mol of O2 at STP is a. 0.411 L. b. 22.4 L. c. 41.6 L. d. 112 L.

18. ______How many moles of H2 would be contained in 4.0 L of the gas at 2.0 atm and 127°C?

a. 89.6 mol b. 6.38 mol c. 0.24 mol d. 0.77 mol

19. ______A sample of H2 is collected over water such that the combined H2-water vapor sample is held at a pressure of 1 atm. What is the partial pressure of the H2 if that of

the water vapor is 2.5 kPa?


20. ______Among the gases listed, which would have the highest diffusion rate? a. NH3 b. CH4 c. SO2 d. NO2

Part 2--Problem Solving Show your work to receive credit.

1. A sample of hydrogen occupies a volume of 1.20 L at a pressure of 425 kPa. If the temperature of the gas is kept constant, what would be the new volume of the gas at

615 kPa?

2. If a sample of oxygen gas occupies a volume of 3.50 L at 57C and 80.0 kPa, what volume would the gas occupy at STP?

3. How many grams of CO2 would be contained in 8.0 L at 805 mmHg and 27C?

4. What volume is occupied by 5.6 g of N2 at STP?

5. How many molecules of N2 would be contained in a 10.0 L balloon at 0.65 atm and

–73C? (Hint: Use Avogadro’s Number to change from moles to mc)

6. The separation of uranium-235 from uranium-238 has been carried out using diffusion. Calculate the rate of diffusion of the following UF6 compounds: UF6

containing uranium-235 = 349 g/mol, UF6 containing uranium-238 = 352 g/mol.

7. Reacting citric acid (H3C6H5O7) with baking soda (NaHCO3) liberates carbon dioxide gas according to the following balanced equation: H3C6H5O7(aq) + 3NaHCO3(s)

3CO2(g) + 3H2O(l) + Na3C6H5O7(aq). If this reaction is carried out in lab where the

pressure is 742 mmHg and the temperature is 21°C, how many grams of baking soda

are needed to produce 3.78 L of CO2?

8. If the reaction described in #11 was conducted outside on a hot summer day when the temperature was 30°C, what volume of CO2 would be produced if 5 g of baking soda

were used? Assume the pressure is 0.98 atm.

9. The reaction of red phosphorus and chlorine gas is shown by the following equation: P4(s) + 6Cl2(g) 4PCl3(g). If 1.6 g of P4 and 2.00 L of Cl2 at 896 mmHg and 200°C

are used, how many grams of PCl3 would be produced?

25 Gas Variables-SUnit 7 Practice Problems (answers)Unit 7 Review Assignment (answers)Unit 7 Minilab AntacidLab Chem 2Unit 7. Test Chem II Test Version

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