Chapter 8: Gas Laws

Pressure

collision of gas molecules with wall of container

Temperature

related to average speed of gas molecules

A Closed-End Manometer

P = d × g × h

Boyle’s Law: Pressure and Volume Equation

P1V1 = P2V2

Charles’s Law Equation

V1/T1 = V2/T2

P versus V

obtain a curved line called a

hyperbola

V versus 1/P

we obtain a straight line

Gas Law of Boyle

Volume increases, pressure decreases

Gas Law of Charles

Temperature increases, volume increases

Charles’s Law Graph - By plotting V versus T (℃)

we obtain straight lines

Gas Law of Avogadro

Moles increase, volume increases

Gas Law of Avogadro Equation

V1/n1 = V2/n2

The Ideal Gas Law Equation

PV=nRT

Combined Gas Law Equation

P1V1/T1 = P2V2/T2

Molar Mass of a Gas Equation

dRT/P

All gases, regardless of their molecular mass, have the same

average kinetic energy at the same temperature.

Mole Fraction of a Gas

the ratio of the number of moles of a given

component in a mixture to the total number of moles in the

mixture.

❖ Represented by chi ( χ):

n1 n1

χ1 = –––– = –––––––––––––––

ntotal n1 + n2 + n3 + …

Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3 + …

Diffusion

refers to the process of particles (atoms, molecules, or ions)

moving from an area of higher concentration to an area of lower

concentration.

Effusion

is the process in which a

gas escapes from its container

through a tiny hole, or orifice into

a vacuum.

Real gases deviate from the ideal gas law

for two main reasons:

1.) Intermolecular forces of attraction cause the

measured pressure of a real gas to be less than

expected.

2.) When molecules are close together, the volume of

the molecules themselves becomes a significant

fraction of the total volume of a gas.

• Ideal Gas - No volume and no intermolecular attractions

• High Pressure and Low Temperatures cause there to be deviations

Postulates of the Kinetic Molecular Theory

1.) The particles are so small compared to the distances between

them that the volume of the individual particle can be

assumed to be negligible or zero.

2.) The particles are in constant motion.

• The collisions of the particles with the walls of the container

are the cause of the pressure exerted by the gas.

3.) The particles are assumed to exert no forces on each other.

• They are assumed neither to attract nor to repel each other.

4.) The average kinetic energy of a collection of gas particles is

directly proportional to the gas temperature.

• All gases, regardless of their molecular mass, have the same

average kinetic energy at the same temperature.

Ozone formation

• O₂ → O₃ via UV light

Ozone depletion

• : CFCs release Cl radicals

Greenhouse gases

• (CO₂, CH₄) trap heat

Environmental Impact

• Smog, acid rain, global warming

• Strategies: reduce emissions, catalytic converters, renewable energy

Real gases deviate from the ideal gas law: Reason #1- forces of attraction cause the

measured

Intermolecular forces of attraction cause the
measured
pressure of a real gas to be less than

expected.

Real gases deviate from the ideal gas law: Reason #2 - molecules are close together, the volume of

the molecules themselves

becomes a significant fraction of the total volume of a gas.

Van der Waals Equation of State for Gases

Ideal behavior of real

gases at high temperatures can be explained with this equation

Number of Moles INCREASE

• Average kinetic energy remains the same

• Average velocity remains the same

• Collision Frequency increases

Pressure increases

• Average kinetic energy remains the same

• Average velocity remains the same

• Collision Frequency increases

When Temp increases

• Average KE increases

• Average V increases

• Collision frequency increases

Volume Decreases

• KE remains the same

• Velocity remains the same

• Collision frequency increases