# Chapter 5 Outline

## 5.1 Breathing: Putting Pressure to Work

## 5.2 Pressure: The Result of Molecular Collisions

This section introduces the concept of pressure. We will be using it extensively in this chapter, so you must understand what it means. Fortunately it is something that you are familiar with from daily experience. Know what the units are and since we have not done much math for awhile, brush up on your unit conversions.

## 5.3 The SimpleGas Laws: Boyle's Law, Charles's Law, and Avogadro's Law

The gas laws were originally derived from experimental data. The laws described here fit the experiental data for different relationships between pressure, temperature, volume and the amount of material (moles). The relationships should be pretty straightforward and you should be able to derive the laws by thinking about how gas molecules behave and what causes pressure. The model of an ideal gas as a small ball bouncing around should be very useful for checking these problems. The laws are given in many different forms, but they are all related. You should be able to derive all the equations for this section from the combined gas law (V1P1/T1 = V2P2/T2)

## 5.4 The Ideal Gas Law

An ideal gas is one where the molecules occupy no volume and they do not interact with each other. Under many conditions the behavior of real gases is close to idea. As a result, the mathematics of their behavior is greatly simplified. The ideal gas law is a single equation that relates Pressure, Volume, Temperature, and moles. The gas law constant (R) is needed for this equation. Remember it. PV=nRT. Watch the units on R so that the units for the variables work out.

## 5.5 Applications of the Ideal Gas Law: Molar Volume, Density and Molar Mass of a Gas

Gas laws are used for many different practical applications. Here are a couple.

## 5.6 Mixtures of Gases and Partial Pressures

Dalton's law of partial pressures seems a bit strange at first. But remember if we assume an ideal gas it does not matter what the gas is, only the number of molecules. So we can talk about pressures of individual gases in a mixture, or we can add the pressures and get a total pressure. A partial pressure is just the part of the total pressure caused by a specific gas.

## 5.7 Gases in Chemical Reactions: Stoichiometry Revisited

This section is a stoichiometry refresher. Instead of using mass or concentration and volume to find the number of moles for a reaction, now you can use the gas laws.

## 5.8 Kinetic Molecular Theory: A Model for Gases

The Kinetic-Molecular theory carries understanding of gas behavior a bit further. It is simply related to the speed that gas molecules move. The speed is related to temperature and the mass of the molecule. Not all gas molecules move at the same speed, the range of speeds is called the boltzman distribution. This distribution is temperature dependent.

## 5.9 Mean Free Path, Diffusion , and Effusion of Gases

This section describes the mixing and motion of gases.

## 5.10 Real Gases: The Effects of Size and Intermolecular Forces

Recall the asusmptions made for the ideal gas behavior. Well, these assumptions are not always valid. There are several different mathematical relationships that account for the non-ideal behavior of real gases. This section introduces the Van der Walls equation, one of the more commonly used equations for non-ideal gases. The mathematics are more complex, but you should be able to solve for pressure given volume, temperature, and moles.

This page is maintained by

Scott Van Bramer

Department of Chemistry

Widener University

Chester, PA 19013
Please send any comments, corrections, or suggestions to
svanbram@science.widener.edu.

This page has been accessed
4342
times since 5/30/97.

Last Updated Friday, May 25, 2001 2:11:17 PM