Chapter 11: Gases – Introductory Chemistry Study Notes

Gases: Properties and Behavior

Introduction to Gases

Gases are a fundamental state of matter characterized by their ability to expand and fill any container, low density, and high compressibility. Understanding gases is essential for explaining everyday phenomena such as breathing, weather, and air pollution.

Kinetic Molecular Theory

Basic Principles

The kinetic molecular theory provides a model for understanding the behavior of gases. It explains gas properties based on the motion and interactions of particles.

• Gas particles are in constant, straight-line motion.

• They do not attract or repel each other; collisions are elastic.

• There is a large amount of space between particles compared to their size.

• The average kinetic energy of gas particles is proportional to the temperature in kelvin.

Properties of Gases

• Compressibility: Gases can be compressed because of the empty space between particles.

• Shape and Volume: Gases assume the shape and volume of their container.

• Low Density: Gases have much lower densities than liquids and solids.

Pressure: The Result of Molecular Collisions

Definition and Effects

Pressure is the force per unit area resulting from collisions of gas particles with surfaces. It is a key property in understanding gas behavior and is involved in many everyday processes.

• Pressure increases with the number of gas particles in a given volume.

• Pressure changes with altitude, affecting phenomena such as ear pain during airplane travel.

Units of Pressure

Common Units

Pressure can be measured in several units:

• Atmosphere (atm): Average pressure at sea level.

• Pascals (Pa): SI unit, 1 Pa = 1 N/m2.

• Millimeters of mercury (mm Hg): Based on barometer measurements; 1 atm = 760 mm Hg.

• Torr: 1 mm Hg = 1 torr.

• Pounds per square inch (psi): Common in engineering.

Gas Laws

Boyle’s Law: Pressure and Volume

Boyle’s Law describes the inverse relationship between the pressure and volume of a gas at constant temperature and amount.

• Mathematical relationship: $P_1V_1 = P_2V_2$

• As volume decreases, pressure increases, and vice versa.

Charles’s Law: Volume and Temperature

Charles’s Law states that the volume of a gas is directly proportional to its temperature (in kelvin) at constant pressure and amount.

• Mathematical relationship: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• As temperature increases, volume increases.

• Absolute zero (0 K, -273°C) is the lowest possible temperature.

Combined Gas Law

The combined gas law relates pressure, volume, and temperature when the amount of gas is constant.

• Mathematical relationship: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Avogadro’s Law: Volume and Moles

Avogadro’s Law states that the volume of a gas is directly proportional to the number of moles at constant temperature and pressure.

• Mathematical relationship: $V_1/n_1 = V_2/n_2$

• As the number of moles increases, volume increases.

*Additional info: Avogadro’s Law is essential for understanding chemical reactions involving gases.*

Ideal Gas Law

The ideal gas law combines Boyle’s, Charles’s, and Avogadro’s laws into a single equation relating pressure, volume, temperature, and moles.

• Mathematical relationship: $PV = nRT$

• R: Ideal gas constant, 0.0821 L·atm/(mol·K)

• All quantities must be in the units specified by R.

*Additional info: The ideal gas law is used to calculate unknown properties of gases in various conditions.*

Partial Pressures and Gas Mixtures

Dalton’s Law of Partial Pressures

In a mixture of gases, each gas exerts its own pressure independently. The total pressure is the sum of the partial pressures of each component.

• Mathematical relationship: $P_{total} = P_a + P_b + P_c + \ldots$

• Partial pressure = Fractional composition × Total pressure

*Additional info: Dalton’s Law is important in physiology and environmental chemistry.*

Applications and Environmental Chemistry

Air Pollution

Air pollution is caused by various gaseous pollutants such as sulfur dioxide, carbon monoxide, ozone, and nitrogen dioxide. These gases have significant health and environmental impacts.

• Sulfur dioxide: Lung and eye irritant, precursor to acid rain.

• Carbon monoxide: Displaces oxygen in blood, causes sensory impairment.

• Ozone: Eye and lung irritant, damages lungs.

• Nitrogen dioxide: Causes haze, eye and lung irritant.

*Additional info: Legislation such as the Clean Air Act has reduced pollutant levels in many cities.*

Summary Table: Common Units of Pressure

Unit	Definition	Conversion
Atmosphere (atm)	Average pressure at sea level	1 atm = 101,325 Pa = 760 mm Hg
Pascals (Pa)	SI unit	1 Pa = 1 N/m2
Millimeters of mercury (mm Hg)	Barometer measurement	1 mm Hg = 1 torr
Pounds per square inch (psi)	Engineering unit	1 atm = 14.7 psi

Key Equations

• Pressure: $\text{Pressure} = \frac{\text{Force}}{\text{Area}}$

• Boyle’s Law: $P_1V_1 = P_2V_2$

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Avogadro’s Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$

• Ideal Gas Law: $PV = nRT$

• Dalton’s Law: $P_{total} = P_a + P_b + P_c + \ldots$

Learning Objectives

• Describe how kinetic molecular theory predicts the main properties of a gas.

• Identify and explain the relationship between pressure, force, and area.

• Convert among pressure units.

• Restate and apply Boyle’s, Charles’s, Avogadro’s, and the ideal gas laws.

• Restate and apply Dalton’s law of partial pressures.

• Apply stoichiometry to chemical reactions involving gases.