Gas Laws

Kinetic Theory (Gases) 1. Gas particles do not attract or repel each other 2. Gas particles are much smaller than the distances between them 3. Gas particles are in constant, random motion 4. No kinetic energy is lost when gas particles collide with each other or the walls of their container 5. All gases have the same average kinetic energy at a given temperature

Properties of Gases Based on ideal gases and kinetic theory FLUID – capable of flowing No definite shape or volume * expand to fill container * takes the shape of the container Transmit and exert pressure in all directions

Ideal Gases describe the behavior of gases under “ideal” condition according to kinetic theory includes a 4th variable/factor affecting gases

Formula: PV = nRT P = pressure in atmospheres V = volume in liters n = # of moles R = gas constant = 0.0821 L atm/mol K T = temperature in Kelvin (change ° C to K  + 273)

EX: A sample of 11.4 moles of dry ice (solid CO2) becomes gas at room temperature (25 °C). Calculate the volume of CO2 if pressure is 9.62 atm.

EX: Average lung capacity of humans is 4000 mL EX: Average lung capacity of humans is 4000 mL. Assuming you breathe pure oxygen, how many moles of O2 gas can your lungs hold at 37°C (body temperature) and 827 mm pressure?

Molar Volume Volume of gas in a mole 22.4 L/mol (conversion) At STP

STP Standard temperature and pressure Standard temperature: 0°C (aka 273 K) Standard pressure: 1 atm