Chapter 11: Gases – Properties, Laws, and Applications

Gases: Properties and Behavior

Introduction to Gases

Gases are one of the fundamental states of matter, characterized by their ability to expand and fill any container. Understanding gases is essential for explaining phenomena such as breathing, weather, and chemical reactions involving gases.

How Straws Work: Atmospheric Pressure and Gas Behavior

Pressure Differences and Drinking with Straws

• Pressure Difference: Drinking from a straw works by creating a pressure difference between the inside and outside of the straw. Sucking air out of the straw lowers the internal pressure, allowing atmospheric pressure to push the liquid up.

• Atmospheric Pressure: The maximum height a liquid can be pushed up a straw by atmospheric pressure is about 10.3 meters (34 feet) for water.

• Force from Gas Molecules: Atmospheric pressure is caused by the force of gas molecules colliding with surfaces.

• Pressure at Sea Level: At sea level, atmospheric pressure averages 101,325 N/m2 (14.7 lb/in2).

Kinetic Molecular Theory: A Model for Gases

Postulates of the Kinetic Molecular Theory

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

• No Interactions: Gas particles do not attract or repel each other.

• Empty Space: There is a lot of space between gas particles compared to their size.

• Kinetic Energy and Temperature: The average kinetic energy of gas particles is proportional to the temperature in kelvin.

Properties of Gases Explained by Kinetic Molecular Theory

• Compressibility: Gases are compressible due to the large amount of 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 Constant Molecular Collisions

Definition and Effects of Pressure

• Pressure: The force per unit area resulting from collisions of gas particles with surfaces.

• Everyday Effects: Pressure allows us to drink from straws, inflate objects, and breathe.

• Atmospheric Variation: Pressure decreases with altitude, affecting the body (e.g., ear pain when climbing a mountain).

Factors Affecting Pressure

• Number of Particles: More gas particles in a given volume increase the pressure.

Units of Pressure

Common Units and Conversions

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

• Pascal (Pa): SI unit; 1 Pa = 1 N/m2.

• Millimeter of Mercury (mm Hg): Based on the height of a mercury column in a barometer; 1 atm = 760 mm Hg.

• Torr: 1 mm Hg = 1 torr.

• Other Units: Inches of mercury (in. Hg), pounds per square inch (psi).

Unit	Equivalent to 1 atm
Atmosphere (atm)	1 atm
Pascals (Pa)	101,325 Pa
Millimeters of mercury (mm Hg)	760 mm Hg
Torr	760 torr
Pounds per square inch (psi)	14.7 psi
Inches of mercury (in. Hg)	29.92 in. Hg

Pressure Unit Conversion Example

• To convert from atm to mm Hg: multiply by 760 mm Hg/atm.

Gas Laws

Boyle’s Law: Pressure and Volume

Boyle’s law states that the volume of a gas is inversely proportional to its pressure at constant temperature and amount of gas.

• Mathematical Form: $P_1V_1 = P_2V_2$

• Explanation: Decreasing the volume increases the pressure, and vice versa.

Applications: Scuba Diving and Boyle’s Law

• Scuba divers must ascend slowly to avoid lung overexpansion due to pressure changes.

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 of gas.

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

• Explanation: Heating a gas increases its volume if pressure is constant.

The Combined Gas Law

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

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

Avogadro’s Law: Volume and Moles

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

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

• Explanation: Adding more gas increases the volume if temperature and pressure are constant.

The Ideal Gas Law

The ideal gas law combines Boyle’s, Charles’s, and Avogadro’s laws into a single equation:

• Equation: $PV = nRT$

• R (Ideal Gas Constant): 0.0821 L·atm/(mol·K)

• Units: P in atm, V in L, n in mol, T in K

Mixtures of Gases and Partial Pressures

Dalton’s Law of Partial Pressures

• Each gas in a mixture exerts its own pressure, called partial pressure.

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

• Partial pressure is calculated as: Partial pressure = Fractional composition × Total pressure

Applications and Environmental Chemistry

Air Pollution

• Sulfur Dioxide (SO2): Emitted from electricity generation and metal refining; causes respiratory irritation and acid rain.

• Carbon Monoxide (CO): Produced by incomplete combustion; can displace oxygen in blood.

• Ozone (O3): Upper-atmosphere ozone protects from UV; ground-level ozone is a pollutant.

• Nitrogen Dioxide (NO2): Emitted by vehicles and power plants; causes haze and respiratory issues.

Summary Table: Key Gas Laws

Law	Equation	Variables Held Constant
Boyle’s Law	$P_1V_1 = P_2V_2$	n, T
Charles’s Law	$\frac{V_1}{T_1} = \frac{V_2}{T_2}$	n, P
Avogadro’s Law	$\frac{V_1}{n_1} = \frac{V_2}{n_2}$	P, T
Combined Gas Law	$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$	n
Ideal Gas Law	$PV = nRT$	None

Key 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 law.

• Apply Dalton’s law of partial pressures.

• Apply stoichiometry to chemical reactions involving gases.