Gases are perhaps the best understood of the three states of matter.
Chapter five touches all of the important bases as far as the
study of gases is concerned. Students should strive to work as
many problems as possible in Chapter 5...and pay
particular attention to the units involved in all
variables and constants.
Chapter 5 Outline
5.1 An Overview of the Physical States of Matter
5.2 Measuring the Pressure of a Gas
5.3 The Gas Laws and Their Experimental Foundations
5.4 Further Applications of the Ideal Gas Law
5.5 The Ideal Gas Law and Reaction Stoichiometry
Breakdown, Discussion & Help With Sample Problems: What You Should Know
5.1 An Overview of the Physical States of Matter
Gases are different from liquids and solids in that they expand to fill the entire container in which they are confined. Gases, as opposed to liqujids and solids, are also highly compressible, thermally expandable, have low viscosities and low densities, and are highly miscible with other gases.
5.2 Measuring the Pressure of a Gas
Pressure is defined as force per unit area, barometers are devices used to measure atmospheric pressure, and manometers are devices used to measure the pressure of a gas involved in a lab experiment.
As far as pressure is concerned, chemists refer primarily to two sets of units.
The atmospheric pressure at sea level is defined as one atmosphere...and one atmosphere of pressure is defined as 760 torr.
5.3 The Gas Laws and Their Experimental Foundations
A gas sample can be described in terms of four variables: pressure (P), volume (V), temperature (T), and the number of moles of gas (n).
The essential point of this chapter can be gleaned from the following statement: any one of the above four variables can be determined by measuring the other three...assuming that the gas in question can be considered to be an ideal gas.
There are several fundamental laws discussed in this section...including--
*Boyle's Law: which states that, at constant temperature and # of moles, that the pressure and volume for a given ideal gas are inversely proportional to each other
*Charles' Law: which states that, at constant pressure and # of moles, that the volume and temperature of a given ideal gas are directly proportional to each other
*Avogadro's law: which states that, at constant temperature and pressure, that equal volumes of any ideal gas contain equal numbers of particles (or moles)
click here for a few words about SAMPLE PROBLEM 5.2 (page 181)
click here for a few words about SAMPLE PROBLEM 5.3 (page 184)
click here for a few words about SAMPLE PROBLEM 5.4 (page 185)
An important fact is discussed on page 186 of Silberberg's text: at standard temperature (273 K) and pressure (1 atm), one mole of an ideal gas has a volume of 22.4 L. The value of 22.4 L is therefore referred to as the standard molar volume.
Finally, the three laws listed above culminate in the Ideal Gas Law (eq 1), which relates the four variables listed at the beginning of this section [pressure (P), volume (V), temperature (T), and the number of moles of gas (n)]--
PV = nRT (1) where R is the universal gas constant. It is important to recognize that CHEM 200 students are not required to remember the exact value of R. R is a constant...and its numerical value depends on the units that are attached to it. CHEM 200 students should particular attention to Table 5.3 (p 187) and Sample Problem 5.5.
click here for a few words about SAMPLE PROBLEM 5.5 (page 188)
5.4 Further Applications of the Ideal Gas Law
This section includes numerous equations (all derived from eq 1 above) that enable determinations of various properties of gases, such as their density and molar mass.
click here for a few words about SAMPLE PROBLEMS 5.6 and 5.7 (pages 190 and 191, respectively)
There is yet one additional law discussed in this section: Dalton's Law. Dalton's Law of partial pressures states that in a mixture of unreacting gases, the total pressure is the sum of the partial pressures of the individual gases. Sample Problem 5.8 is concerned with Dalton's Law.click here for a few words about SAMPLE PROBLEM 5.8 (page 193)
5.5 The Ideal Gas Law and Reaction Stoichiometry
Section 5.5 includes yet another application of eq 1--determinations of reaction stoichiometry (for reactions that involve gases).
click here for a few words about SAMPLE PROBLEM 5.11 (page 197)
5.6 The Kinetic-Molecular Theory: A Model for Gas Behavior
This section includes a rather large amount of theoretical background that aims to explain the behavior of gases at a fundamental level.
Included in this section are equations that enable determination of the rms speed of a gaseous molecule (eq 5.12; p. 201) and the rate of effusion of a gas (Graham's Law).
Sample Problem 5.12 nicely summarizes important aspects of Section 5.6.
5.7 Real Gases: Deviations from Ideal Behavior
In a way, the essential points contained within Section 5.7 is a cruel trick on CHEM 200 students.
Specifically, Silberberg spends the better part of 35 pages describing how all gases can be treated as ideal gases...and now, in Section 5.7, he takes it all back.
Well, sort of. Actually, equation 1 (above) does still apply to gases...but it just needs to be tweaked up a bit with a couple of additional constants. The result is called the van der Waals eq (eq 5.14 on p. 210 of Silberberg), an equation that requires constants a and b...both of which are provided to CHEM 200 students on examinations.
click here for a few words about SAMPLE PROBLEM 5.13 (page 210-11)