Gases

Kinetic Molecular Theory of Gases

A gas consists of small particles (atoms/molecules) that move randomly with rapid velocities

Further Information

They move faster when heated.

The attractive forces between particles of a gas

can be neglected

Do you think this is accurate?

Why would this be important for calculations?

The actual volume occupied by a gas molecule is

extremely small compared to the volume

that gas occupies.

Is this true in the real world?

Why would this be helpful with calculations?

The average kinetic energy of a gas molecule is

proportional to Kelvin temperature

What is kinetic energy? Why Kelvin temperature and not Celsius or Fahrenheit? What does proportional mean?

Gas particles are in constant motion, moving rapidly in

straight paths. *

Is this true?

What do we know about their motion?

Why would the real situation make the calculations more difficult?

Ideal Gases

An imaginary gas that perfectly fits all the assumptions of the kinetic molecular theory (KMT).

Expansion

Gases do not have definite shape or volume.

The expand to any container they are enclosed in.

A gas in a 1 L container is then put into a 2 L container. How much volume does it have now?

Fluidity

In an ideal gas, the gas particles glide past each other.

This feature allows gases to be referred to as fluids just like liquids.

Low Density

Density of a gas substance is only about 1/1000 of the same substance in liquid or solid state.

Why is this true?

Compressibility

This is a crowding effect of gases when the volume is decreased

Diffusion

Spontaneous (does not require energy) mixing of particles of two substances caused by their random motion

Properties of a Gas

Units of Measure

Pressure

Pressure is not the same as force.

Pressure is a force over an area.

Example: psi = Pounds per in2

Measuring Pressure

A barometer measures atmospheric pressure.

Units of Pressure

kPa, atm, mm of Hg, torr

Helpful Conversions

1 atm = 760 mm Hg

1 atm = 760 torr

1 mm Hg = 1 torr

1 atm = 101.325 kPa

Volume

L, mL or cm3

Helpful conversions 1000 mL = 1 L 1 mL = 1 cm3

Temperature

0C or K

Helpful conversions:0C = K – 273

K = 0C + 273

moles

Number of moles = n

If you are given grams, how would you convert to moles?

Standard Temperature and Pressure (STP)

Standard Temperature is 00C or 273 K

Standard Pressure is 101.3 kPa or 1 atm

Boyle’s Law:

Pressure and volume are inversely proportional

P1V1 = P2V2

Boyle’s Law

http://www.grc.nasa.gov/WWW/K-12/airplane/boyle.html

Charles’ Law:

Temperature and Volume are directly proportional

V1 / T1 = V2 / T2

Gay-Lussac’s Law:

Pressure and Temperature are directly proportional

P1/T1 = P2/T2

Combined Gas Law

P1V1 = P2V2

T1 T2

If you remember this law, hold constant the other variables not used and you have all the gas laws we’ve used so far.

Molar Volume

1 mole = 22.4 L of a gas at STP

Dalton’s Law of Partial Pressure

Be sure all units of pressure are the same.

If not, convert all units to the same unit of measure

The total pressure is equal to the sum of the partial pressures

Ideal Gas LawPV=nRT

P = Pressure (atm)

V = volume (L)

n = number of moles

R = 0.0821 atm x L / moles x K

T = temperature (K)

You must use these units for the R constant to be correct.

Name the Law!

You will be given a series of laws and asked to name the law or you will be

given the name and be asked to come up with the formula!

PV =nRT

Ideal Gas Law

V1 / T1 = V2 / T2

Charles’ Law

Boyle’s Law

P1V1 = P2V2

Combined Gas Law

P1V1 = P2V2

T1 T2

Gay-Lussac’s Law:

P1/T1 = P2/T2

Avagadro’s Law:

V1 / n1 = V2/n2