Chapter 9: Gases

absolute zero

temperature at which the volume of a gas would be zero according to Charles’s law.

Amontons’s law

(also, Gay-Lussac’s law) pressure of a given number of moles of gas is directly proportional to its kelvin temperature when the volume is held constant

atmosphere (atm)

unit of pressure; 1 atm = 101,325 Pa

Avogadro’s law

volume of a gas at constant temperature and pressure is proportional to the number of gas molecules

bar

(bar or b) unit of pressure; 1 bar = 100,000 Pa

barometer

device used to measure atmospheric pressure

Boyle’s law

volume of a given number of moles of gas held at constant temperature is inversely proportional to the pressure under which it is measured

Charles’s law

volume of a given number of moles of gas is directly proportional to its kelvin temperature when the pressure is held constant

compressibility factor (Z)

ratio of the experimentally measured molar volume for a gas to its molar volume as computed from the ideal gas equation

Dalton’s law of partial pressures

total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the component gases

diffusion

movement of an atom or molecule from a region of relatively high concentration to one of relatively low concentration (discussed in this chapter with regard to gaseous species, but applicable to species in any phase)

effusion

transfer of gaseous atoms or molecules from a container to a vacuum through very small openings

Graham’s law of effusion

rates of diffusion and effusion of gases are inversely proportional to the square roots of their molecular masses

hydrostatic pressure

pressure exerted by a fluid due to gravity

ideal gas

hypothetical gas whose physical properties are perfectly described by the gas laws

ideal gas constant (R)

constant derived from the ideal gas equation R = 0.08206 L atm mol^–1 K^–1 or 8.314 L kPa mol^–1 K^–1

ideal gas law

relation between the pressure, volume, amount, and temperature of a gas under conditions derived by combination of the simple gas laws

kinetic molecular theory

theory based on simple principles and assumptions that effectively explains ideal gas behavior

manometer

device used to measure the pressure of a gas trapped in a container

mean free path

average distance a molecule travels between collisions

mole fraction (X)

concentration unit defined as the ratio of the molar amount of a mixture component to the total number of moles of all mixture components

partial pressure

pressure exerted by an individual gas in a mixture

pascal (Pa)

SI unit of pressure; 1 Pa = 1 N/m²

pounds per square inch (psi)

unit of pressure common in the US

pressure

force exerted per unit area

rate of diffusion

amount of gas diffusing through a given area over a given time

root mean square speed (u_rms)

measure of average speed for a group of particles calculated as the square root of the average squared speed

standard conditions of temperature and pressure (STP)

273.15 K (0 °C) and 1 atm (101.325 kPa)

standard molar volume

volume of 1 mole of gas at STP, approximately 22.4 L for gases behaving ideally

torr

unit of pressure; 1 torr=1/760atm

van der Waals equation

modified version of the ideal gas equation containing additional terms to account for non-ideal gas behavior

vapor pressure of water

pressure exerted by water vapor in equilibrium with liquid water in a closed container at a specific temperature