Kinetic Molecular Theory and Gas Laws

Day 1

Kinetic-Molecular Theory – explains how particles in matter

behave

1.All matter is composed of small particles that are far apart. Gas is

mostly empty space.

2. Particles are in constant, random motion.

3.Particles collide with each other and walls of their containers:

collisions create pressure

4. Collisions are elastic = no KE lost

5. No attractive/repulsive forces between particles. Molecules move in straight lines.

PHASE SOLID LIQUID GASKinetic Energy Low Medium High

Shape Rigid Fluid Fluid – fills container; diffusion

Forces Strong Medium Weak

Volume Definite Definite Indefinite

Compressibility Incompressible Incompressible COMPRESSIBLE!“squish!”

Density High (particles close together)

Medium Low (spread out)

EXIT QUESTIONS:1) Each of these flasks contains the same number of molecules. Which container has the highest pressure? Explain your answer.

EXIT QUESTIONS:2) Which of the following changes to a system

will NOT result in an increase in pressure? Explain why you chose your answer.

a) Increasing the volume of container

b) adding more gas molecules

c) Decreasing the volume of the container

d) Raising the temperature

Kinetic Molecular Theory and Gas Laws

Day 2

Factors affecting gases1.Volume – amount of space an object occupies

• Measured in milliliters (mL) or Liters (L)

• 1000 mL = 1 L

We already have heard that 1 mol = 22.4 L @ STP The more moles wehave the bigger the

balloon will need to be!

Example 1

How many moles of nitrogen gas are in 89.6 L at STP?

89.6 L N2 1 mol N2

22.4 L N2

= 4.00 mol N2

Example 2:What volume does 76 grams of fluorine (F2) occupy at STP (normal conditions)?

76 g F2

37.996 g F2

1 mole F2

1 mole F2

22.4 L F2

= 44.8 L F2

**45 L

2. Temperature Average kinetic

energy of particles (how fast they go)

Measured in Kelvin K = oC + 273

Ex: Convert 17oC to

Kelvin:

17oC + 273 = 290 K

*C= 5/9 (*F-32)*F= 9/5 (*C) + 32K= *C + 273*C= K- 273

3. Pressure

Force exerted by a gas per unit area on a surface. Example: Pounds/in2 or psi

Results from the simultaneous collisions of billions of gas particles with the walls of the vessel containing the gas.

Standard pressure: 760 mm Hg = 1 atmosphere = 101.3 kPa

= 29.92 in. Hg = 14.7 psi = 760 torr

Measuring Atmospheric PressureMeasured with a

barometer. A barometer uses a

column of mercury that rises to an average height of 760 mmHg at sea level.

1 atmosphere (1 atm)

Standard Temperature and Pressure (STP)The conditions of standard

temperature and pressure are = 1.0 atm pressure and = 273 K (or 0C).

@STP 1 mole of gas = 22.4 L of gas

Example 1

The atmospheric pressure in Denver, CO is 0.830 atm on average. Express this pressure in mm Hg.

0.83 atm mm Hg

1 atm = 760 mm Hg

0.83atm

1 atm

760 mm Hg

= 630.8 mm Hg**631 mmHg

Example 2

Convert a pressure of 175 kPa to atmospheres.

175 kPa atm101.3 kPa = 1 atm

175 kPa

101.3 kPa

1 atm

1.72 atm

Gas Law Foldable

Fold the left and right to the middle. Cut along solid lines (but only to the crack!)

Dalton’s Law of Partial Pressure

The pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases.

PTotal = P1 + P2 + P3….

Example

A balloon is filled with air (O2,

CO2, & N2) at a pressure of

1.3 atm.

If PO2 = 0.4 atm and PCO2 =

0.3 atm, what is the partial pressure of the nitrogen gas?

PTotal = P1 + P2 + P3….

Ptotal = PO2 + PCO2 + PN2

1.3 atm = 0.4 atm + 0.3 atm + PN2

PN2 = 0.6 atm