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
E._____(10) 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
Answers
I.__ ___(05) Effusion & Diffusion of Gases-Section 12.7 Answers
L._____(10) 5thMultiple Choice Chapter 12 5thAnswers
Chapter 12withAns 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:
http://www.fccj.us/chm2045/KotzMovies/chapter12/index.html
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 |
|
|
Units |
Value |
|
L atm/mol K |
0.08206 |
|
cal/mol K |
1.987 |
|
J/mol K |
8.314 |
|
m3 Pa/mol K |
8.314 |
|
L torr/mol K |
62.36 |
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.
Part E: Gas Law Problems-
Sections 12.2, 12.3, 12.5 Answers
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
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 Answers
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 Answers
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 Answers
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.