The Nature of Gases

I. The Kinetic Theory and a Model for GasesA. Assumptions of the Kinetic Theory

1. Gases consist of large numbers of tiny particles that are far apart• Most of the volume of a gas is empty space• Explains why gases are easily compressed• Accounts for low density of gases

2. Gas particles undergo elastic collisions, no net loss of kinetic energy

3. Gas particles are in constant, random motion – possess kinetic energy

4. There are no forces of attraction or repulsion between gas particles

5. Average kinetic energy of gas particles depends on the temperature

II. Kinetic Energy and TemperatureA. Average kinetic energy of gas particles depends

on temperatureKE = ½mv2

B. All gases at the same temperature have the same average K.E.• At the same temperature lighter gas particles have

higher average speed than do heavier particles

C. Kelvin temperature scale – directly proportional to the average kinetic energy of a substance

↑Temperature ↑speed and ↑K.E. or ↓Temperature ↓speed and ↓K.E.

Effect of Temperature

III. Kinetic Theory and the Nature of GasesA. Expansion

1. Gases do not have a definite shape or volume2. Gases take the shape of their container3. Gases evenly distribute themselves in a

container

B. Fluidity – gas particles flow past one another

C. Low Density1. A substance in the gaseous state has 1/1000th

the density of the same substance in the liquid or solid state

D. Compressibility1. Gases are mostly empty space2. Gases can be compressed, decreasing the

distance between particles, and decreasing the volume occupied by the gas

E. Diffusion – spontaneous mixing of particles caused by random motion

1. Rate of diffusion depends on three properties of the gas particles: speed, diameters, and attractive forces between them

E. Effusion – process by which gas particles under pressure pass through a tiny opening

1. Rates of effusion are directly proportional to velocities of particles

2. Molecules of low mass effuse faster than molecules of high mass

IV. Gas Pressure A. Pressure and Force

1. Pressure = force area

2. Newton (N) – SI unit of force3. When a force is applied to a smaller and

smaller area, the pressure increases 4. Gas molecules exert pressure on any surface

with which they collide

Pressure & Volume of a Gas

B. Measuring Pressure1. Barometer – device used to measure

atmospheric pressure• Evangelista Torricelli – invented 1st

barometer in early 1600’s• Experiments show that at sea level, 0°C,

average pressure of the atmosphere can support 760 mm column of Hg

2. Manometer – used to measure pressure of enclosed gas sample

How a Barometer Works

A. Units of Pressure1. Pascal (Pa) – SI unit of pressure • 1 Pa = 1 N/m2

2. mmHg – millimeter of mercury3. Atmospheres – one standard atmosphere (atm) 4. 1 atm = 760 mmHg = 101.3 kPa = 1.01325 x 105

pascal (Pa)B. STP – standard temperature and pressure, 0° C and

1 atm