gaseous state of any substance that normally exists as a liquid or a solid
a measure of the force exerted on a unit area. In chemistry, pressure if often expressed in units of atmospheres (atm) or torr: 760 torr = 1 atm; in SI units, pressure is expressed in pascals (Pa)
the SI unit of pressure: 1 Pa = 1 N/(m^2)
a unit of pressure equal to 100 kPa
standard atmospheric pressure
defined as 760 torr or, in SI units, 101.325 kPa
atmosphere (atm)
a unit of pressure equal to 760 torr; 1 atm = 101.325 kPa
a unit of pressure (1 torr= 1 mm Hg)
Boyle’s law
a law stating that at constant temperature, the product of the volume and pressure of a given amount of gas is constant
Charles’s law
a law stating that at constant pressure, the volume of a given quantity of gas is proportional to absolute temperature (K)
Avogadro’s hypothesis
a statement that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules
Avogadro’s law
a statement that the volume of a gas maintained at constant temperature and pressure is directly proportional to the number of moles of the gas
ideal-gas equation
an equation of state for gases that embodies Boyle’s law, Charles’s law, and Avogadro’s hypothesis in the form PV=nRT
ideal gas
a hypothetical gas whose pressure, volume, and temperature behavior is completely described by the ideal-gas equation
gas constant (R)
the constant of proportionality in the ideal-gas equation
standard temperature and pressure (STP)
defined as 0* C and 1 atm pressure; frequently used as reference conditions for a gas
partial pressures
the pressure exerted by a particular gas in a mixture
Dalton’s law of partial pressures
a law stating that the total pressure of a mixture of gases is the sum of the pressures that each gas would exert if it were present alone
mole fraction
the ration of the number of moles of one component of a mixture to the total moles of all components; abbreviated X, with a subscript to identify the component
kinetic-molecular theory of gases
a set of assumptions about the nature of gases. These assumptions, we translated into mathematical form, yield the ideal-gas equation
root-mean-square (rms) speed
the square root of the average of the squared speeds of the gas molecules in a gas sample
the escape of a gas through and orifice or a hole
Graham’s law
a law stating that the rate of effusion of a gas is inversely proportion to the square root of its molecular weight
the spreading of one substance through a space occupied by one or more other substances
mean free path
the average distance traveled by a gas molecule between collisions
van der Waals equation
an equation of state for nonideal gases that is based on adding corrections of the ideal-gas equation. The correction terms account for intermolecular forces of attraction and for the volumes occupied by the gas molecules themselves

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