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four variables that define the state of a gaseous sample

pressure (P), volume (V), temperature (T), and number of moles (n)


gas pressure units

atmospheres (atm) or millimeters of mercury (mmHg), which are equivalent to torr; SI unit is the pascal (Pa); mathematical relationship between all of these: 1 atm = 760 mmHg = 760 torr = 101.325 kPa


standard temp and pressure (STP)

conditions at 273 K (0 degrees C) and 1 atm; generally used for gas law calculations


standard state conditions

298 K, 1 atm, 1 M concentrations; used when measuring standard enthalpy, entropy, free energy changes and electrochemical cell voltage


ideal gas

represents a hypothetical gas with molecules that have no intermolecular forces and occupy no volume


ideal gas law

PV = nRT; where P = pressure, V = volume, n = number of moles, T = temp, R = ideal gas constant (8.21 x 10^-2 L x atm/mol x K or 8.314 J/K x mol which is derived when SI units of Pa and cubic meters (for volume) are substituted into the ideal gas law


density (fancy p)

ratio of the mass per unit volume of a substance; usually expressed for gases in units of grams per liter; density derived from ideal gas law: PV = nRT; n = m/Molar mass; PV = (m/molar mass) RT and density = m/V = P(molar mass)/RT


how much space does a mole of an ideal gas at STP occupy?

22.4 L


combined gas Law

can be used to relate changes in temp, volume, and pressure of gas: Psub1Vsub1/Tsub1 = Psub2Vsub2/Tsub2; where the subscripts 1 and 2 refer to the two states of the gas; this equation assumes number of moles stays constant


change in volume

Vsub2 = Vsub1[Psub1/Psub2][Tsub1/Tsub2]


change in volume used to find density

density = m/Vsub2


molar mass from density

molar mass = (density at STP)(22.4 L/mol)


Avogadro's principle

states that all gases at a constant temp and pressure occupy volumes that are directly proportional to the number of moles of gas present; equal amounts of all gases at the same temp and pressure will occupy equal volumes: n/V = k or nsub1/Vsub1 = nsub2/Vsub2; k = a constant; nsub1 and nsub2 = number of moles of gas 1 and gas 2; Vsub1 and Vsub2 = volumes of gases 1 and 2


Boyle's Law

for a given gaseous sample held at constant temp (isothermal), the volume of the gas is inversely proportional to its pressure: PV = k or Psub1Vsub1 = Psub2Vsub2; k = a constant; subscripts 1 and 2 = different sets of pressure and volume condidtions; in terms of gas law, n and T are constant here... As pressure increases volume decreases


Charles's Law

states that, at constant pressure, volume of a gas is proportional to its absolute temp in kelvins: V/T = k or Vsub1/Tsub1 = Vsub2/Tsub2; k = proportionality constant; subscripts 1 and 2 = two different sets of temp and volume conditions; n and P (with reference to ideal gas law) are held constant... as temp increases, volume increase


Gay-Lussac's law

relates pressure to temp: P/T = k or Psub1/Tsub1 = Psub2/Tsub2; subscripts 1 and 2 = two different sets of temp and pressure conditions; n and V (in terms of ideal law) are constant; as temp increases, pressure increases


combined gas law

combination of many of the preceding laws; relates pressure and volume in numerator, and variations in temp to both volume and pressure


Dalton's Law of partial pressures

if two or more gases that do not chemically interact are found in the same container each one will act independently of the other; law states total pressure of a gaseous mixture is equal to the sum of the partial pressure of the individual components


partial pressure

pressure exerted by each gas in a container where multiple gases that do not chemically interact are found


Dalton's Law equation

PsubT = PsubA + PsubB + PsubC + etc..; where PsubT = total pressure in container; PsubA through C = partial pressures of respective gases


determining partial pressure

PsubA = (XsubA)(PsubT); where XsubA = moles of gas A/total moles of gas and PsubT = total pressure of container


vapor pressure

pressure exerted by evaporated particles above the surface of a liquid


Henry's Law

[A] = ksubH x PsubA or [A]sub1/Psub1 = [A]sub2/Psub2 = KsubH; [A] = concentration of A in solution; ksubH = Henry's constant; PsubA = partial pressure of A


kinetic molecular theory

used to explain the behavior of gases; assumptions are: 1. Gases are made up of particles with volumes that are negligible compared to volume of container; 2. gas atoms or molecules exhibit no intermolecular attractions or repulsions; 3. Gas particles are in continuous, random motion, undergoing collisions with other particles and the container walls; 4. Collisions bt any two gas particles or particles and container are elastic meaning conserving momentum and kinetic energy; 5. avg kinetic energy of gas particles is proportional to the absolute temp of the gas (in kelvins), and it is the same for all gases at a given temp


kinetic energy of a gas particle

KE is proportional to the absolute temp of gas: KE = 1/2mv^2 = 3/2ksubB(T); where ksubB = Boltzmann constant (1.38 x 10^-23 J/K)


Boltzmann constant

1.38 x 10^-23 J/K, serves as a bridge between macroscopic and microscopic behaviors of gases (as a bridge between behavior of the gas as a whole and the individual gas molecules)


root-mean-square speed (u sub rms)

a way to define the average speed of gases, determined by the average kinetic energy per particle and then calculating the speed which corresponds to it; given by the equation: u sub rms = square root of (3RT/M); R = ideal gas constant; T = temp, M = molar mass


Maxwell-Boltzmann distribution curve

shows the distribution of gas particle speeds at a given temp


Graham's Law

under isothermal and isobaric conditions, rates at which two gases diffuse are inversely proportional to the square roots of their molar masses: r sub 1/r sub 2 = square root of (Msub2/Msub1); r1 and r2 are diffusion rates of gases 1 and 2; M1 and M2 are molar masses of gases 1 and 2; noted that a gas with 4 times the molar mass of another gas will travel half as fast as the lighter gas



the flow of gas particles under pressure from one compartment to another through a small opening; for two gases at same temp, rates of effusion are proportional to avg speeds