Gas Laws

Why Gases Behave As They Do

Factors Affecting Gases• Pressure

Atmospheremm Hg, Torrpsi (pounds per square inch; #/in2)Paschal (N/ m2)

• VolumeLiter

• Temperature (Kelvin)• Amount

Mole

Pressure Equivalencies

• 1 Atmosphere =• 760 mm Hg =• 760 torr =• 101.3 kPa =

1 paschal = 1 Newton/ meter2

• 14.7 psi

Ideal Gas Law

• Applies to kinetic-molecular theory gases– Real gases at high temperatures & low pressures

• P V = n R TP = Pressure (SI unit is Paschal)V = Volume in Litersn = # molesR = gas constant (adjusts to units of pressures)T = Temperature in Kelvin {273.15 + oC}

Variations of PV = nRT

• Boyle’s LawP1 V1 = P2 V2

• Charles’ LawV1/ T1 = V2/ T2 or V1 T2 = V2 T1

• Gay-Lussac’s LawP1/ T1 = P2/ T2 or P1 T2 = P2 T1

• Combined Gas Law P1 V1/ T1 = P2 V2/ T2 or P1 V1 T2 = P2 V2 T1

Other Values of Importance• Dalton’s Law of Partial Pressures

Total pressure of a gas = S partial pressures• Graham’s Law of Effusion

½ MAnA2 = ½ MBnB

2 {M = molar mass; = nvelocity}

The lighter gas must effuse faster.Rate of effusion of A/ rate of B = √MB / √MA

• Standard Temperature & Pressure (STP)1 atmosphere; 0o C = 273 K

• Gas Molar Constant = 22.41410 Liters/ mole