kinetic-molecular theory - explains behavior of gases
- gases made up of large number of molecules in constant, random motion
- gas molecules are extremely tiny
- attractive/repulsive forces between gas molecules don’t really do anything
- energy transferred between molecules during collisions, but average kinetic energy stays the same (elastic collisions)
- individual molecules have different speeds
- average kinetic energy directly proportional to absolute temperature
- constant temperature >> constant average kinetic energy
- rms speed (u) - speed of molecule w/ average kinetic energy
- average kinetic energy = 1/2 mu2
- u = (3RT/molar mass)1/2
effusion - escape of gas molecules through a tiny opening
- Graham’s law - compares rates of effusion under identical conditions
- shows that lighter gas effuses more rapidly
- r1/r2 = (M2/M1)1/2
- diffusion - spread of substance throughout a space or 2nd substance
- molecular collisions >> motion of gas molecules constantly changes >> diffusion is slow
- mean free path - average distance a molecule moves between collisions
behavior of real gases - different than behavior of ideal gases
- ideal gas equation assumes that gas molecules take up no space, have no intermolecular forces
- less deviation w/ higher temperature, lower pressure
- gas volumes usually slightly greater than predicted by ideal-gas equation
- gas pressure usually slightly lesser than predicted by ideal-gas equation
van der Waals equation - takes into account the gas volume and attractive forces
- (P + n2a/V2)(V - nb) = nRT
- constants a, b different for each gas
Find the pressure in atm that O2 exerts at 70.6° C if 1.850 moles occupies 16.5 L.
- Given:
- a = 1.36 (L2 atom/ mol2)
- b = 0.0318 (L/mol)
- R = 0.08206 (L-atm/mol-K)
- (P + n2a/V2)(V - nb) = nRT
- (P + (1.850)2(1.36)/(16.5)2)(16.5 - (1.850)(0.0318)) = (1.850)(0.08206)(343.6)
- P = [(1.850)(0.08206)(343.6)] / [(16.5 - (1.850)(0.0318))] - (1.850)2(1.36)/(16.5)2
- P = 3.16 atm
Subject:
Subject X2:
