Module 6 Gases & Their Properties Assignment Outline (Chapter 12)

The test outline for Module 6 of Exam #4 covers all of Chapter 12 for the Kotz 6e text. Most should be review from CHM 1025C.  Below is a Part by Part test outline with links to sample tests and answers plus text reference sections to study for that objective:

Module Six: The Gaseous State  (Chapter 12)
A._____(05) Kinetic Molecular Theory-Section 1.5, 12.1, 12.6 Answer
B._____(05) Discussion Real vs Ideal Gas Equation-Sect 12.9 Answer bc
C._____(05) Standard Conditions/Molar Volume-Sect 12.3 Answer bc
D._____(05) Gas Laws/Vocabulary-Sections 12.2, 12.5 Answers
Gas Law Problems-Sections 12.2, 12.3, 12.5 Answers
F. _____(05)
Volume-Volume Stoichiometry Problem-Section 12.4
Answers fg
G._____(05) Mass-Volume Stoichiometry Problem-Section 12.4 Answers fg
H._____(05) Gas Densities/Molecular Mass Determination-Sect 12.3

I.__ ___(05)
Effusion & Diffusion of Gases-Section 12.7 Answers
L._____(10) 5th
Multiple Choice Chapter 12   5thAnswers  
Chapter 12
withAns  Old Final
______(60) Total
= ______% 

Don't forget to view the short movies clips as you study chapter 12 which are provided by our text at: 

To begin your study of Module 6, please read Chapter 12.

 Part A: Kinetic Molecular Theory-Section 1.5, 12.1, 12.6

The gas properties and laws discussed in Chapter 12 are based on the Kinetic Molecular theory.  The CHM 1025C texts list five or six basic assumptions. You may write those assumpions as you previously learned. The CHM 2045 text combines the statements into three basic assumpions which you may write for the answer for Part A. The Kotz 6e on page 567 lists:


(a) Gases consist of particles (molecules or atoms), whose separation is much greater than the size of the particles themselves.

(b) The particles of a gas are in continual, random, and rapid motion. As they move, they collide with one another and with the walls of their container, but they do so without energy loss.

(c) The average kinetic energy of gas particles is proportional to the gas temperature. All gases, regardless of molar mass, have the same average kinetic energy at the same temperature.

From the CHM 1025C texts:

1. Gases are composed of molecules*[1].  The distance between the molecules is very-very great compared to the size of the molecules themselves, and the total volume of the molecules is only a very-very small fraction of the entire space occupied by the gas.  Therefore, considering volume, we are primary considering empty space.  (This assumption explains why gases are highly compressed and have very low densities.)

2. No attractive forces exist between molecules in a gas.  (This is what keeps a gas from spontaneously becoming a liquid.)

3. The molecules of a gas are in a state of constant, rapid motion, colliding with each other and with the walls of the container in a perfectly random manner.  (This assumption explains why different gases normally mix completely.  The collisions between molecules and the walls of the container account for the pressure exerted by the gas.)

4. All of these molecular collisions are perfectly elastic. As a result, the system as a whole experiences no loss of kinetic energy, the energy derived from the motion of a particle.

5. The average kinetic energy per molecule of a gas is proportional to the absolute temperature, and the average kinetic energy per molecule is the same at a given temperature and pressure for all gases. 
[1]When we think of molecules of elemental gases, we usually think of the diatomic gases such as nitrogen, oxygen, hydrogen, etc. The Nobel gases exist as monoatomic gases such as Helium, Neon, etc.



Part B: Discussion Real vs Ideal Gas Equation-Sect 12.9

If you have an understanding of the Kinetic Molecular Theory above then when you read section 12.9 you apply the KMT to gases in non ideal behavior. At STP gases behave ideally. But under extreme condition which cause oevrcrowding, the KMT breaks down such that the Ideal gas Equation: PV=nRT has to be re-written to the Real Gas Equation. This leads to the following discussion questions:

(a) In the Real Gas Equation:   (P + an2/V2) (V - nb) = n RT a pressure correction factor was added. Why? (What assumptions of the kinetic theory breakdown under extreme conditions of temperature and pressure?)

(b) Also a volume correction factor was subtracted. Why? (What assumptions of the KMT breakdown under extreme conditions?)

From the Answers posted:

In Section 12.9 on page 576 the answer to the first question is found in the third paragraph!

 Another assumption of the kinetic molecular theory is that collisions between the molecules are elastic-that is, that the atoms or molecules of the gas never stick to one another by some type of intermolecular force.  This is not true at extreme conditions of overcrowding. When a molecule is about to strike the wall of its container, other molecules in the vicinity exert a slight attraction for the molecule and pull it away from the wall. As a result of the intermolecular forces, molecules strike the wall with less force than they would in the absence of intermolecular attractive forces. Therefore, in a real gas, the observed pressure is less than the predicted pressure by the ideal gas law and a pressure correction factor is added to account for this pressure loss.

Also a volume correction factor was subtracted. Why? (What assumptions of the KMT breakdown under extreme conditions?)

In Section 12.9 on page 576 the answer is in the second paragraph!

 The kinetic molecular theory and the ideal gas law are concerned with the volume available to the molecules to move about, not the volume of the molecules themselves. It is clear the volume occupied by the gas molecules is NOT negligible at high pressures (or extreme low temperatures. The available volume is less than the volume of the container. The volume the molecules occupy must be subtracted from the volume of the container to obtain the volume of free space the molecule can move.

A good multiple choice question is: under what conditions does ideal gas behavior break down?

Part C: Standard Conditions/Molar Volume-Sect 12.3

In section 12.1 The properties of gases  are discussed. This includes the introction to the concept of Gas pressure. From you previous chemistry you should already know some of the values of standard temperature and pressure. These are also listed on pages 548-550.

State standard conditions (STP) in three units of pressure (the last is your choice) and oC and K temperatures:

 _760__mmHg or _760__torre=  __1___atm = _29.9 in_  = _14.7 psi_= 101 kPa

 __0_oC   =    _273__K

 From CHM 1025C you should know the value of the gram molar volume constant to three significant figures. On page 558 an ideal gas occupies 22.414 L at STP. Therefore you would put 22.4 in any of the following blanks: 

What are the values for the Molar Gas Volume Constant for the following gases:

 1 moleCO2 =__22.4__L CO2@STP         1 moleH2 =__22.4___L H2@STP

 1 moleN2 =__22.4___L N2@STP           1 moleO2 =__22.4___L O2@STP

Calculate the value of R in the Ideal Gas Equation at STP:

If you substitute the values of the Molar Gas Constant into the ideal gas equation (PV=nRT) you can calculate the value of the constant R:

PV  = nRT   (you must enter Kelvin temperatures-not Celsius)

(1 atm) (22.4 L) = (1 mole) R (273 K)

R = 0.08206 L atm/mol K



 R can include energy units such as Joules or calories:

Values for the gas constant R



L atm/mol K


cal/mol K


J/mol K


m3 Pa/mol K


L torr/mol K


We usually use the first value: 0.08206 L atm/mol K in the calculation in Module 6.

Part D: Gas Laws/Vocabulary-Sections 12.2, 12.5

For Part D you simply write a statement of the gas laws covered in chapter 12. In section 12.2 Boyle's , Charles, Avogadro's Hypothesis, and the general gas Law. Section 12.3 introduces the Ideal gas law.  Dalton's law and Partial pressures are introduced in Section 12.5. Graham's Law of Diffusion is covered in Section 12.7

Part E: Gas Law Problems-
               Sections 12.2, 12.3, 12.5

Part E asks you to do simple gas law calculations. Problem #1 is Boyle's law. In Section  12.2 Example 12.2 on page 552 is a sample of the first problem. Work Exercise 12.2 on Page 552 as well as Problems #5-#6 on page 580.

Problem #2 is a Charles Law calculation. In Section 12.2 Example 12.3 p 554 is a Charles Law Application. Work Exercise 12.3 on page 554 as well as #7 & #8 on page 580.

Problem #3 is a Dalton's law of Partial Pressure application. Section 12.5 covers Dalton's Law. Example 12.11 on page 566 is a sample of a Dalton's law application. Work Exercise 12.10 on page 566 as well as  Problems #37-#40 page 582 for more samples of the third gas law application problem in Part E.


Problem #4 is a combined gas law application or an ideal gas equation calculation. Example 12.4 is a General Gas Law appplication in section 12.2 page 555.  Work Exercise 12.4 for practice on page 555. In Section 12.3 Example 12.6 is an ideal Gas law Calculation. Work Exercise 12.6 on page 558 as an addition application sample. There are several problems at the end of the chapter: General gas law Problems #9-14 page 580 and Ideal Gas Laws #27-30 page 581.

Part F: Volume-Volume Stoichiometry Problem-
Section 12.4

Volume-Volume stoichiometry applies directly Avogadro's Hypothesis. Example 12.5 in section 12.2 show that in an all gas phase reaction (all reactants and products are gases) you can apply the coefficients that balance an equation as a direct volume-volume ratio (similar to the Mole Ratio). Work Exercise 12.5 page 557 as a sample. Problems #15 & #16 on page 581 are addition exercises to work.


Part G: Mass-Volume Stoichiometry Problem-
Section 12.4

Mass-Volume Stoichiometric problems are covered in Section 12.4. You road map is Figure 12.10 on page 562 to solve these problem types. Study Examples 12.9 and Example 12.10 on page 562-3 as examples of Part G problems.  At the end of the chapter work Problems #31-#40 on page 582.

Part H: Gas Densities/Molecular Mass Determination-Sect 12.3 Answers

Gas densities are discussed in section 12.3 on page 559. Study Examples 12.7  page 559 and 12.8 page 560 for Part H problem types. Then work Exercise 12.7 on page 560 and Exercise 12.8 on page 561 for addition examples of Part H. At the end of the chapter work Problems #23-#26 on page 581.


Part I: Effusion & Diffusion of Gases-
Section 12.7

In Section 12.7 you will work Diffusion and Effusion problems. Note the formula 12.10 on page 572. Example 12.13 on page 573 is a sample Part I problem. Work exercise 12.12 on page 573. Please note rate of effusion is amount escaped a containing in a given amount of time. When ever times are given students usually miss the problems because the times are placed in the denominator of the rates, inverting the formula.  There are four additional problems at the end of the chapter: #47-50 Page 583 for you to work.