BackChapter 11: Gases – Properties, Laws, and Applications
Study Guide - Smart Notes
Tailored notes based on your materials, expanded with key definitions, examples, and context.
Gases and Atmospheric Pressure
How Straws Work and the Role of Atmospheric Pressure
Straws function based on the principles of atmospheric pressure. When you drink from a straw, you reduce the number of air molecules inside the straw, creating a pressure difference between the inside and outside. The higher external atmospheric pressure pushes the liquid up the straw and into your mouth. However, even with perfect materials and seals, a straw longer than about 10.3 meters (34 feet) would not work because atmospheric pressure can only support a column of water up to this height.
Pressure is the force exerted per unit area by gas molecules as they collide with surfaces.
At sea level, atmospheric pressure averages 101,325 N/m2 (14.7 psi).
Straws work because sucking creates a pressure difference; the external pressure pushes the liquid up until the weight of the liquid column balances atmospheric pressure.
Kinetic Molecular Theory: A Model for Gases
Postulates of the Kinetic Molecular Theory
The kinetic molecular theory provides a model for understanding the behavior of gases. It is based on several key assumptions:
Gases are collections of particles in constant, straight-line motion.
Gas particles do not attract or repel each other; their collisions are elastic.
There is a large amount of empty space between gas particles compared to their size.
The average kinetic energy of gas particles is proportional to the temperature in kelvins.
Properties of Gases Explained by Kinetic Molecular Theory
Compressibility: Gases are compressible because of the large amount of empty space between particles.
Shape and Volume: Gases assume the shape and volume of their container due to negligible intermolecular attractions and constant motion.
Low Density: Gases have much lower densities than liquids and solids because of the empty space between particles.
Pressure: The Result of Constant Molecular Collisions
Definition and Effects of Pressure
Pressure is the result of constant collisions between gas molecules and the surfaces around them. It is essential for many everyday phenomena, such as drinking through a straw, inflating objects, and breathing. Atmospheric pressure decreases with altitude, which can cause discomfort (e.g., ear pain) due to pressure imbalances.
Pressure (P) is defined as force per unit area:
Pressure depends on the number of gas particles in a given volume; more particles mean higher pressure.
Units of Pressure
Pressure can be measured in several units:
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 torr = 1 mm Hg.
Pounds per square inch (psi): 1 atm = 14.7 psi.
Unit | Average Air Pressure at Sea Level |
|---|---|
pascal (Pa) | 101,325 Pa |
atmosphere (atm) | 1 atm |
millimeter 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 |
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:
As volume decreases, pressure increases, and vice versa.
Mathematically:
Charles’s Law: Volume and Temperature
Charles’s law states that the volume of a gas is directly proportional to its temperature (in kelvins) at constant pressure and amount:
As temperature increases, volume increases.
Mathematically:
Absolute zero (0 K or -273°C) is the theoretical temperature at which a gas would have zero volume.
The Combined Gas Law
The combined gas law relates pressure, volume, and temperature for a fixed amount of gas:
Mathematically:
Useful when two or more variables change simultaneously.
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:
As the number of moles increases, volume increases.
Mathematically:
The Ideal Gas Law
The ideal gas law combines Boyle’s, Charles’s, and Avogadro’s laws into a single equation:
R (ideal gas constant) =
All quantities must be in the correct units: P (atm), V (L), n (mol), T (K).
Partial Pressures and Gas Mixtures
Dalton’s Law of Partial Pressures
In a mixture of gases, each component exerts a pressure independently of the others. The total pressure is the sum of the partial pressures of each component:
Dalton’s law:
Partial pressure = fractional composition × total pressure
Physiological Effects of Partial Pressures
Our bodies are adapted to breathe oxygen at a partial pressure of 0.21 atm.
Low partial pressures of oxygen (e.g., at high altitudes) can cause hypoxia (oxygen starvation).
High partial pressures of oxygen (e.g., deep diving) can cause oxygen toxicity.
High partial pressures of nitrogen can cause nitrogen narcosis.
Gases in Chemical Reactions and Molar Volume
Stoichiometry with Gases
In reactions involving gases, the ideal gas law can be used to relate the volume, pressure, temperature, and amount of gas. At standard temperature and pressure (STP: 0°C, 1 atm), one mole of any ideal gas occupies 22.4 L.
Use the ideal gas law to convert between moles and volume for gases in reactions.
At STP: 1 mol gas = 22.4 L
Chemistry in the Environment: Air Pollution
Major Gaseous Air Pollutants
Sulfur dioxide (SO2): Emitted from electricity generation and metal refining; causes respiratory irritation and acid rain.
Carbon monoxide (CO): Produced by incomplete combustion; displaces oxygen in blood, causing health risks.
Ozone (O3): Upper-atmosphere ozone protects from UV light; ground-level ozone is a pollutant and lung irritant.
Nitrogen dioxide (NO2): Emitted by vehicles and power plants; causes haze, respiratory irritation, and acid rain.
Legislation such as the Clean Air Act has significantly reduced pollutant levels in the U.S. over recent decades.
Chapter 11 in Review
Kinetic molecular theory explains the properties of gases.
Pressure results from molecular collisions.
Gas laws (Boyle’s, Charles’s, Avogadro’s, and the ideal gas law) describe relationships among pressure, volume, temperature, and amount.
Partial pressures are important in mixtures and physiological contexts.
Air pollution is a significant environmental and health concern.
