1. Gas Laws Why Gases Behave As They Do

2. Factors Affecting Gases Pressure  Atmosphere  mm Hg, Torr  psi (pounds per square inch; #/in 2 )  Paschal (N/ m 2 ) Volume  Liter Temperature (Kelvin) Amount  Mole

3. Pressure Equivalencies 1 Atmosphere = 760 mm Hg = 760 torr = 101.3 kPa =  1 paschal = 1 Newton/ meter 2 14.7 psi

4. Ideal Gas Law Applies to kinetic-molecular theory gases – Real gases at high temperatures & low pressures P V = n R T  P = Pressure (SI unit is Paschal)  V = Volume in Liters  n = # moles  R = gas constant (adjusts to units of pressures)  T = Temperature in Kelvin {273.15 + o C}

5. Variations of PV = nRT Boyle’s Law  P 1 V 1 = P 2 V 2 Charles’ Law  V 1 / T 1 = V 2 / T 2 or V 1 T 2 = V 2 T 1 Gay-Lussac’s Law  P 1 / T 1 = P 2 / T 2 or P 1 T 2 = P 2 T 1 Combined Gas Law  P 1 V 1 / T 1 = P 2 V 2 / T 2 or P 1 V 1 T 2 = P 2 V 2 T 1

6. Other Values of Importance Dalton’s Law of Partial Pressures  Total pressure of a gas =  partial pressures Graham’s Law of Effusion  ½ M A A 2 = ½ M B B 2 {M = molar mass;  velocity}  The lighter gas must effuse faster.  Rate of effusion of A/ rate of B = √M B / √M A Standard Temperature & Pressure (STP)  1 atmosphere; 0 o C = 273 K Gas Molar Constant = 22.41410 Liters/ mole