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MCAT Gen. Chem Ch. 8

Gases Are The Least: Dense phase of matter
Gases Are Fluids, And Therefore: They conform to the shapes of their containers
Gases Are Easily: Compressible
Gas Systems Are Described By The Variables: T, temp, P press., V volume, n number of moles
1 atm = 760 mmHg = 760 torr = 101.325 kPa
Simple Mercury Barometer Measures incident (atmospheric) pressure. As pressure increases, more mercury is forced into the column, increasing its height. As pressure decreases, mercury flows out of the column under its own weight, and decreases its height
Standard Temperature And Pressure (STP) 273 K (0C) and 1 atm
Equimolar Amounts Of Two Gases Will Occupy The: Same volume at the same temperature and pressure. At STP, one mole of an ideal gas occupies 22.4 L.
Kinetic Molecular Theory Attempts to explain the behavior of gas particles. It makes a number of assumptions about the gas particles.
Assumptions Of The Kinetic Molecular Theory, Gas Particles: Have negligible volume, do not have intermolecular attractions or repulsions, undergo random collisions with each other and the walls of the container
Collisions Between Gas Particles Are: Elastic
Average Kinetic Energy Of Gas Particles Is: Directly proportional to temperature.
Diffusion Spreading out of particles from high to low concentration
Effusion Movement of gas from one compartment to another through a small opening under pressure
Real Gases Deviate From Idea Behavior Under: High pressure (low volume) and low temperature conditions.
At Moderately High Pressures, Low Volumes, Or Low Temperatures, Real Gases Will: Occupy less volume than predicted by the ideal gas law because the particles have intermolecular attractions.
At Extremely High Pressures, Low Volume, Or Low Temperatures, Real Gases Will Occupy more volume than predicted by the ideal gas law because the particles occupy physical space.
Eq. 8.1: Ideal Gas Law PV = nRT. R = ideal gas constant 8.21 x 10^-2 L*atm / mol*K
Eq. 8.2: Density Of A Gas D = m/v = PM / RT. M = Molar Mass. R = ideal gas constant 8.21 x 10^-2 L*atm / mol*K.
Eq. 8.3: Combined Gas Law P1V1/T1 = P2V2/T2
Eq. 8.4: Avogadro's Principle n/V = k or n1/V1 = n2/V2
Eq. 8.5: Boyle's Law PV = k or P1V1 = P2V2. k = a constant
Eq. 8.6: Charles's Law
Eq. 8.7: Guy-Lussac's Law P/T = k OR P1/T1 = P2/T2
Eq. 8.8: Dalton's Law (Total Pressure From Partial Pressures) PT = PA + PB + PC + ...
Eq. 8.9: Dalton' Law (Partial Pressure From Total Pressure) PA = XA*PT. XA = moles of gas A / total moles of gas
Eq. 8.10: Henry's Law [A] = kH * PA or [A]1/P1 = [A]2/P2 = kH. kH = Henry's constant that depends on the identity of the gas.
Eq. 8.11: Average Kinetic Energy Of A Gas KE = 1/2mv^2 = 3/2*ka*T. ka = Boltzmann constant, 1.38 x 10^-23 J/K.
Eq. 8.12: Root-Mean-Square Speed urms = Sq.root(3RT/M). R = ideal gas constant, 8.314 J/mol*K. T = temperature. M = molar mass.
Eq. 8.13: Graham's Law r1/r2 = Sq.root(M2/M1)
Eq. 8.14: Van der Waals Equation Of State (P + n^2*a / V^2) * (V-nb) = nRT. a = physical constant for attractive forces. b = physical constant for big particles.
Created by: SamB91