1. AP Chemistry 2014-2015 CH 5 GASES REVIEW

2.  1 atm = 760.00 mmHg = 760.00 torr = 101.325 kPa = 1.013 x 10 3 Pa = standard pressure (sea level)  STP = 273 K, 1 atm  R = 0.08206 Latm/molK or 8.314510 J/molK GENERAL THINGS TO KNOW

3. P 1 V 1 /T 1 = P 2 V 2 /T 2  memorize it (“peas and vegetables on the table”) COMBINED GAS LAW  USE TO GET BOYLE’S, GA-LUSSAC’S, CHARLES’S LAWS

4.  PV = nRT  R = 0.08206 Latm/molK, though it has other values and units as well; useful only at low pressures and high temperatures 5.3 THE IDEAL GAS LAW

5.  “Molecular Mass kitty cat”—all good cats put dirt (dRT) over their pee (P). Ew, but it works.  Remember that the densities of gases are reported in g/L not g/mL. THE DENSITY OF GASES

6.  P total = P 1 + P 2 + …. + P n  You can also use this concept when collecting a gas over water; total pressure = atmospheric pressure  P total = P gas + P water DALTON’S LAW OF PARTIAL PRESSURES

7.  Assumptions  All particles are in constant, random motion  All collisions between particles are perfectly elastic  The volume of the particles in a gas is negligible  The average kinetic energy of the molecules in a gas it is its Kelvin temperature  These assumptions ignore intermolecular forces.  IMFs are stronger for larger/polar particles, weaker for smaller/nonpolar particles THE KINETIC MOLECULAR THEORY OF GASES

8.  Follows a rough bell curve  At any temperature, some particle will have zero (or near-zero) velocity  As temperature increases, the curve shifts to the right and flattens  Average kinetic energy is only tied to temperature; velocity is tied to both temperature and molar mass DISTRIBUTION OF MOLECULAR SPEEDS

9.  U rms = root mean speed, m/s  T = temperature, Kelvin  MM = mass of a mole of gas particles in kg (weird, I know—respect the math though)  Use the “energy R” or 8.314510 J/molK for this equation since kinetic energy is involved. ROOT MEAN SPEED

10.  Make sure to keep track of which gas is “gas 1” and which gas is “gas 2” to keep from messing up  Low molar mass = faster effusion, high molar mass = slower effusion COMPARING RATES OF EFFUSION FOR GASES WITH DIFFERENT MOLAR MASSES

11.  Real gases behave most ideally  At high temperature  At low pressure  When they have weaker IMF’s (smaller, nonpolar molecules)  The van der Waal’s equation has terms that correct for  Volume of gas particles (term “b”)  IMF’s between gas particles (term “a”) IDEAL VS. REAL GASES