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AP Chemistry by Satellite Lectureguide
Student Edition
Gas Laws
Chapter 10

Objectives

Following your study of this chapter, you should be able to

1. supply the standard state phase for first 89 elements in the periodic table.
2. supply the color of all of the elements with gas phase standard states.
3. list the major defining attributes of solid, liquid and gas phases.
4. define the term atmospheric pressure and describe how a mercury barometer works.
5. state the units for measuring pressure and carry out conversions between different pressure units.
6. state the mathematical relationship between pressure and the density of the liquid in a barometer.
7. draw a manometer and describe its use to measure the pressure of a sample of gas.
8. state the mathematical relationship for Boyle's Law and carry out calculations using the relationship.
9. state the mathematical relationship for Charles' Law and carry out calculations using the relationship.
10. describe two independent experiments that yield a value for absolute zero.
11. write the ideal gas equation and define each variable.
12. use the ideal gas equation to solve for any of its variables or a ratio of variables.
13. extract Boyle's Law, Charles' Law and Avogadro's Law from the ideal gas equation.
14. use the ideal gas law to calculate the amount of product formed, or reactant consumed, in a chemical reaction, given the initial amount of one reactant (or product) and assuming the other reactant is in excess.
15. display the relationship between Dalton's Law of partial pressures and the ideal gas equation and solve problems using Dalton's Law of partial pressures.
16. state the postulates of the kinetic-molecular theory.
17. explain the relationship between the temperature of a sample of gas and the kinetic energy of the particles of the gas.
18. relate the macroscopic behavior of an ideal gas to the kinetic-molecular model of a gas.
19. explain the difference between effusion and diffusion at the molecular level.
20. calculate relative effusion or diffusion rates of gases.

1a. Of the first 89 elements, list those which are gases in their standard state. Mark the gaseous elements on the periodic table.

b) Of the first 89 elements list those which are liquids in their standard state:

1c. What phase are all remaining elements?

2. List all of the elements that are gases and identify their color.

3. Distinguish between the gas, liquid and solid phase by listing the unique properties of each that are not shared by the others.

4a. Define standard atmospheric pressure.

b) How is atmospheric pressure measured?

4c. Label the picture of the device used to measure atmospheric pressure and describe how it works.

5. Given a pressure of 1 atm, express the quantity in mmHg, Pascals (Pa) and kilopascals (kPa).

6a. Beginning with the equation F = ma, derive an equation relating atmospheric pressure (P) to the density (r) of the liquid, the force of gravity (g) and the height of the liquid in a barometer.

7. Sketch a picture of a manometer. Describe how it is constructed and how it is used to measure the pressure of a sample of gas. .

8. Use the data table displayed in Figure I to record the data during the class experiment. When the lecture is over, use the graph paper provided below to plot the data. Label the 'x' and 'y' axis of the graph and clearly identify all data points. Select the range of values for each axis to allow all data points to be displayed.

8. (Continued)

Plot #1

Volume vs. Pressure Graph

Plot #2

Volume vs. 1/Pressure Graph

9. Use the data table displayed in Figure II to record the data during the class experiment. When the lecture is over, use the graph paper provided below to plot the data. Label the 'x' and 'y' axis of the graph and clearly identify all data points. Select the range of values for each axis to allow all data points to be displayed and the determination of the x-intercept (the point on the x-axis where y = 0).

9. (Continued)

Plot #1

Using the graph paper below, extrapolate the data to determine the value of absolute zero. How does this value compare to value obtained in the Charles' Law experiment? (Note: Remember the minimum value of the temperature (x-axis) should be near -290 ºC.)

10. Describe the two independent experiments that yield a value for absolute zero.

11a. Write the ideal gas equation and define each of the variables.

13. Beginning with the mathematical form of the ideal gas law, rearrange the equation and show how Boyle's law, Charles' law and Avogadro's law can be obtained.

Problem Set #15
AP Chemistry by Satellite

ALL work must be shown in all problems for full credit.

PS15.1. The pressure of the mercury in the gas (vapor) phase above the liquid mercury in a barometer is 2.0 x 10^-3 mmHg. Calculate the pressure in units of atmospheres.

PS15.2. A sample of SF₆ occupies a container of variable volume. If the the sample originally occupies a volume of 50.0 mL at 755 mmHg, calculate the pressure which must be exerted to lower the volume to 25.0 mL.

PS15.3. If the pressure of a sample of an ideal gas, initially at 1.25 atm, is tripled, by what factor will the volume of the gas change?

PS15.4. A 545 mL sample of nitrogen gas initially at -220 ºC is heated to 100 ºC. Calculate the new volume, assuming the pressure does not change.

PS15.5. Calculate the volume of 0.390 moles of an ideal gas at 750 mmHg and 23 ºC.

PS15.6. Calculate the density of SF₆ at 1.00 atm and 0.00 ºC.

PS15.7. Calculate the volume of a sample of helium at -33.0 ºC and 1.23 atm if it occupies a volume of 2.34 L at 54.5 ºC and 1026 mmHg.

PS15.8. A 0.751 mol sample of an ideal gas occupies a 10.0 liter flask at 27.0 ºC and 1.85 atm. If 0.257 mol of the gas are removed from the container, calculate the new pressure. (Assume the temperature remains constant.)

PS15.9. Which of the following contains the largest number of particles?

a) 5.00 g of He at 1.00 atm and 0 ºC

b) 225 g of Au

c) 34.5 L of an ideal gas at -5.0 ºC and 2000 mmHg

PS15.10. Which of the following samples has the greater mass?

a) 278 L of Ar at 25 ºC and 300 mmHg

b) 225 mL of CH₄ at 300 ºC and 5.34 atm

c) 34.5 L of Cl₂

15a. State Dalton's Law of partial pressures and explain how it is related to the ideal gas law equation.

16. State the principle postulates of the kinetic-molecular theory.

17. Explain in words the relationship between temperature of a sample of gas and the kinetic energy of the particles of the gas.

18. In each of the following experiments, explain the macroscopic behavior of gases using kinetic-molecular theory.

19a. Describe the difference between effusion and diffusion at the molecular level.

20a. Write the mathematical equation that describes the relative effusion or diffusion rates for ideal gases.

21a. List the conditions under which real gases deviate most from ideal behavior.

b) Discuss the postulates of the kinetic-molecular theory that are not true for real gases. Explain how real gases behave in terms of these postulates.

c) Write the van der Waals equation for real gases and define the variables.

Problem Set #16
AP Chemistry by Satellite

ALL work must be shown in all problems for full credit.

PS16.1. The first laboratory preparation of O₂, performed by Joseph Priestly, required the heating of mercury (II) oxide,

What volume of O₂ at 22.5 ºC and 753 mmHg will be produced when 15.3 g of HgO are completely decomposed?

PS16.2. Hydrogen, H₂, can be prepared by passing steam through a hollow air tube which has been heated to high temperature. The reaction is,

Calculate the volume of H₂ formed at 0.98 atm and a temperature of 450 ºC when 98.3 g of H₂O are passed through an iron tube.

PS16.3. Calculate the volume of SO₂(g) produced when 2.00 liters of O₂ react with excess sulfur at constant temperature and pressure.

PS16.4. A gaseous mixture contains 3.00 g of N₂, 0.430 moles of Ar and 2.15 x 10²³

molecules of CH₄. If the total pressure of the mixture is 3.00 atm, calculate the partial pressure of each component.

PS16.5. Experimentally oxygen gas is frequently collected over water. When a sample of oxygen, produced from a chemical reaction, is collected over water at 24 ºC, the total pressure was 765 mmHg. Calculate the pressure due only to oxygen.

PS16.6. A sample of oxygen collected over water at 29 ºC exerts a total pressure of 759 mmHg. If the volume of the container is 125 mL calculate the mass of oxygen present.

PS16.7. Can the speed of a molecule of a gas be doubled if the temperature of the gas sample is held constant, yes or no? Briefly explain your answer.

PS16.8. Explain, in terms of the kinetic molecular model, why increasing the volume of a sample of an ideal gas decreases the pressure of the gas at constant temperature.

PS16.9. Rank the following gases in order of rate of diffusion. Explain your order.

Ne, CH₄, SF₆, CO₂

PS16.10. Describe the two factors responsible for the deviation of real gases from ideal behavior as assumed in the ideal gas equation.

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