General Chemistry: Chapter 6

when temperature increases, how does that impact kinetic energy and collisions?

KE increases and collisions increase

area for pressure using area and force

pressure = force / area

Boyle’s Law

P1V1 = P2V2 (inverse relationship)

Avogadro’s Law

V1/n1 = V2/n2 (direct relationship)

Charles’ Law

V1/T1 = V2/T2 (inverse relationship)

Gay-Lussac’s Law

P1/T1 = P2/T2 (direct relationship)

5 points of the Kinetic-Molecular Theory of Gases

1. particles in gas = constant & random motion

2. combined volume of particles are negligible

3. particles exert no forces on one another

4. gas molecules collisions are completely elastic (no IMF)

5. all gases have the same average KE at given temperature

gases behave most ideally at (high/low) pressures and (low/high) temperatures

LOW pressure (gases far apart = decreased interactions = decreased IMF)

HIGH temperatures (gases move faster = increased KE = decreased IMF)

ideal gas law

PV = nRT (R = 0.0821)

values for STP

273 K, 1 atm, 1 mol of gas = 22.4 L

Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3…

P1 = X1 x Ptotal (X1 = mole fraction)

gas density equation

density = mass / volume = (pressure x Molar mass) / (R x temperature)

diffusion vs effusion

diffusion: high concentration to low concentration

effusion: confined gas escaping through a hole (gas > liquid > solid)

effusion law

r1/r2 = sqrt(M2/M1)

r = rate of effusion

M = molar mass