chem: gases

STP

standard temp and pressure; 0°C (273 K), 1 atm

1 atm \=

101.3 kPa, 760 mmHg, 760 Torrs

1 dm^3

1 L

1 cm^3 or 1 cc

1 mL

1000 mL

1 L

1 m^3

1000 L

ideal gases have

very small mass, little to no volume, no attraction between molecules

boyle's law

P1V1 \= P2V2; inversely related

charles' law

V1/T1 \= V2/T2; directly proportional

avogadro's law

V1/n1 \= V2/n2; directly proportional

guy-lussac's law

P1/T1 \= P2/T2; directly proportional

combined gas law

P1V1/n1T1 \= P2V2/n2T2

law

way of generalizing behavior, uses math

theory

explains why, no math involved

4 assumptions of kinetic molecular theory

all particles are in constant and random motion, all collisions are perfectly elastic, volume of particles is negligible, average kinetic energy is directly proportional to the kelvin temp

elastic

collisions do not lose energy

dalton's law

P1 + P2 + P3... \= Ptotal

molar mass of a gas

MM \= DRT/P; density is in g/L

diffusion

mixing of gases; moving from a higher concentration to a lower concentration

effusion

passage of gas through a tiny opening into an evacuated chamber

graham's law

rate a/rate b \= sqrt of molar mass b/molar mass a; gas a \= lighter, heavier, gas b \= heavier, slower

real gases

behave ideally at ordinary temperatures and pressures; why? molecules are very close, molecular interactions

van der waal's equation

accounts for the behavior of real gases at low temp/high pressure; (P+(n^2a/v^2))(v-nb) \= nRT

van der waal's equation "a"

accounts for intermolecular attraction

van der waal's equation "b"

accounts for volume of gas molecules