Ch 9 and Ch 5 terms

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

a 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 indvidual 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 (urms)

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

bomb calorimeter

device designed to measure the energy change for processes occurring under conditions of constant volume; commonly used for reactions involving solid and gaseous reactants or products

calorie (cal)

unit of heat or other energy; the amount of energy required to raise 1 gram of water by 1 degree Celsius; 1 cal is defined as 4.184 J

calorimeter

device used to measure the amount of heat absorbed or released in a chemical or physical process

calorimetry

process of measuring the amount of heat involved in a chemical or physical process

chemical thermodynamics

area of science that deals with the relationships between heat, work, and all forms of energy associated with chemical and physical processes

endothermic process

chemical reaction or physical change that absorbs heat

energy

capacity to supply heat or do work

enthalpy (H)

sum of a system's internal energy and the mathematical product of its pressure and volume

enthalpy change (ΔH)

heat released or absorbed by a system under constant pressure during a chemical or physical process

exothermic process

chemical reaction or physical change that releases heat

expansion work (pressure-volume work)

work done as a system expands or contracts against external pressure

first law of thermodynamics

internal energy of a system changes due to heat flow in or out of the system or work done on or by the system

heat (q)

transfer of thermal energy between two bodies

heat capacity (C)

extensive property of a body of matter that represents the quantity of heat required to increase its temperature by 1 degree Celsius (or 1 kelvin)

Hess's law

if a process can be represented as the sum of several steps, the enthalpy change of the process equals the sum of the enthalpy changes of the steps

hydrocarbon

compound composed only of hydrogen and carbon; the major component of fossil fuels

internal energy (U)

total of all possible kinds of energy present in a substance or substances

joule (J)

SI unit of energy; amount of energy used when a force of 1 newton moves an object 1 meter, 1 J = 1 kg m2/s2 and 4.184 J = 1 cal

kinetic energy

energy associated with an object's motion, equal to one-half the product of the object's mass and the square of its velocity, 12𝑚𝑣2 (where m = mass and v = velocity)

nutritional calorie (Calorie)

unit used for quantifying energy provided by digestion of foods, defined as 1000 cal or 1 kcal

potential energy

energy of a particle or system of particles derived from relative position, composition, or condition

specific heat capacity (c)

intensive property of a substance that represents the quantity of heat required to raise the temperature of 1 gram of the substance by 1 degree Celsius (or 1 kelvin)

standard enthalpy of combustion (Δ𝐻∘c)

heat released when one mole of a compound undergoes complete combustion under standard conditions

standard enthalpy of formation (Δ𝐻∘f)

enthalpy change of a chemical reaction in which 1 mole of a pure substance is formed from its elements in their most stable states under standard state conditions

standard state

set of physical conditions as accepted as common reference conditions for reporting thermodynamic properties; 1 bar of pressure, and solutions at 1 molar concentrations, usually at a temperature of 298.15 K

state function

property depending only on the state of a system, and not the path taken to reach that state

surroundings

all matter other than the system being studied

system

portion of matter undergoing a chemical or physical change being studied

temperature

intensive property of matter that is a quantitative measure of "hotness" and "coldness"

thermal energy

kinetic energy associated with the random motion of atoms and molecules

thermochemistry

study of measuring the amount of heat absorbed or released during a chemical reaction or a physical change

work (w)

energy transfer due to changes in external, macroscopic variables such as pressure and volume; or causing matter to move against an opposing force