Kinetic Theory &

Boyles Law

Kinetic Theory of Gases

• All matter consists of tiny particles in constant motion

Kinetic Energy – energy an object has due

to it’s motion.

Assumptions

• Particles are small, hard spheres with insignificant volume

• Particle motion is rapid, constant & random– Move in straight path until collision occurs

• All collisions btwn particles are ELASTIC– Kinetic energy is transferred with no loss (so

total kinetic energy stays constant)

Gas Pressure

• Due to simultaneous collisions of billions of particles of gas on a object

• UNITS:– Pascals (Pa) (SI)– Standard atmosphere (atm)– Millimeters of mercury

(mmHg)

1atm = 760mm Hg = 101.3 kPa

Atmospheric Pressure

• Results from collisions of atoms & molecules in the air with objects

• As you increase altitude the atm pressure decreases– b/c density decreases

Gases & Temperature• Heated particles store

energy

• Causes particles to move faster!– Causes kinetic energy

to change, so we use the avg. kinetic energy

• UNITS: Kelvin– Directly proportional to

avg KE of particles

Gas Property

• Compressibility– Measure of how much the

volume of matter decreases under pressure

– Think of squeezing gases into a smaller container!

– Easier with gases b/c of the space btwn particles

Factors Affecting Gas Pressure

1. Amount of Gas (mol) - More particles, more collisions, more pressure

2. Volume (L) - More volume, less pressure

- What happens if you compress a gas?

3. Temperature (K)

- increase temp, increase collisions, increase pressure

Boyle’s Law

Pressure & Volume

Boyle’s Law

• If temperature & mass is constant, as pressure increases the volume decreases.

• P1V1 = P2V2

Example Problem

Nitrous oxide (N2O) is used as an anesthetic. The pressure on 2.5L of N2O changes from 105 kPa to 40.5 kPa. If the temperature does not change, what will the new volume be?

V1= P1=

V2 = P2 =