Phase Changes

Kinetic= motion

All matter consists of small particles

The molecules are in constant, random, rapid motion

All collisions are elastic (no net loss of energy)

Kinetic Theory

As temperature increases, the molecules velocity increases, increasing the pressure on the container.

The average distance traveled between collisions.

An oxygen molecules will collide with other molecules 4.5 billion times per second!

Mean Free Path

The temperature of a substance is a measure of the average kinetic energy of its particles.

Absolute zero – the temperature at which molecular motion stops (-273.15oC)

05

1015202530354045

100 200 300 400

temperature(K)

Kin

eti

c energ

y

Caused by gas molecules colliding with the sides of a container

Force per unit area

Units:

Gas Pressure

2

1

m

Newton

1 atm = 101.3 kPa = 760 mmHg = 760 torr

How many mmHg is in one kPa?

kPa

mmHgkPa

3.101

7601 7.5 mmHg

Open Manometer – atmosphere exerts pressure on one side and gas sample exerts pressure on the other side◦ Add if gas pressure is greater◦ Subtract if air pressure is greater

Closed Manometer (barometer) – vacuum on one side, gas pressure on the other side◦ No addition or subtraction necessary

Measuring Pressure

States of Matter

State Shape Volume Compressible?

Flows?

Solid Definite Indefinite

No No

Liquid Indefinite

Definite No Yes

Gas Indefinite

Definite Yes Yes

The pressure produced when vapor particles above a liquid collide with container walls; a dynamic equilibrium exists between the liquid and the vapor.

Vapor pressure increases with temperature. A substance with weak intermolecular forces

has a high vapor pressure and low boiling points. (these are volatile – alcohols, ethers)

A substance with strong intermolecular forces has a low vapor pressure and high boiling point (these are nonvolatile – water, molasses, glycerol)

Vapor Pressure

Critical Point – above this temperature, no amount of pressure can liquefy it.

Triple Point – all three phases are in equilibrium

Phase Diagrams