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| absolute zero | temperature at which the volume of a gas would be zero according to Charles’s law.( -273.15 °C)
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| 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
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| atmosphere (atm) | unit of pressure; 1 atm = 101,325 Pa
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| Avogadro’s law | volume of a gas at constant temperature and pressure is proportional to the number of gas molecules
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| bar | (bar or b) unit of pressure; 1 bar = 100,000 Pa
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| barometer | device used to measure atmospheric pressure
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| 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
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| 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
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| compressibility factor (Z) | ratio of the experimentally measured molar volume for a gas to its molar volume as computed from the ideal gas equation
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| 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.
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| 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)
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| effusion | transfer of gaseous atoms or molecules from a container to a vacuum through very small openings
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| Graham’s law of effusion | rates of diffusion and effusion of gases are inversely proportional to the square roots of their molecular masses
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| hydrostatic pressure | pressure exerted by a fluid due to gravity
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| ideal gas | hypothetical gas whose physical properties are perfectly described by the gas laws
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| ideal gas constant (R) | constant derived from the ideal gas equation R = 0.08226 L atm mol–1 K–1 or 8.314 L kPa mol–1 K–1
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| ideal gas law | relation between the pressure, volume, amount, and temperature of a gas under conditions derived by combination of the simple gas laws
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| kinetic molecular theory | theory based on simple principles and assumptions that effectively explains ideal gas behavior
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| manometer | device used to measure the pressure of a gas trapped in a container
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| mean free path | average distance a molecule travels between collisions
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| 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
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| partial pressure | pressure exerted by an individual gas in a mixture
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| pascal (Pa) | SI unit of pressure; 1 Pa = 1 N/m2
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| pounds per square inch (psi) | unit of pressure common in the US
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| pressure | force exerted per unit area
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| rate of diffusion | amount of gas diffusing through a given area over a given time
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| root mean square velocity (urms) | measure of average velocity for a group of particles calculated as the square root of the average squared velocity
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| standard conditions of temperature and pressure (STP) | 273.15 K (0 °C) and 1 atm (101.325 kPa)
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| standard molar volume | volume of 1 mole of gas at STP, approximately 22.4 L for gases behaving ideally
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| torr | unit of pressure; 1 torr = (1/760) atm
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| van der Waals equation | modified version of the ideal gas equation containing additional terms to account for non- ideal gas behavior
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| vapor pressure of water | pressure exerted by water vapor in equilibrium with liquid water in a closed container at a specific temperature
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