1. GASES
Unit 1
Chapter 1

2. Motion of particles: In solids the particles:
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Motion of particles:
In solids the particles:
are moving relatively slowly.
have low kinetic energy
In liquids the particles:
molecules move faster.
have higher kinetic energy.
In gases, the particles:
move fastest,
have high kinetic energy.

4. Kinetic Theory Model of States
Liquid
Particles vibrate, rotate, tumble and “flow”, but cohesion (molecular attraction) keeps them close together.
Gas
Particles move freely through container. The wide spacing means molecular attraction is negligible.
Solid
Particles vibrate but don’t “flow”. Strong molecular attractions keep them in place. 4

5. Particles can have 3 types of motion:
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Particles can have 3 types of motion:
Vibrational kinetic energy (vibrating)
Rotational kinetic energy (tumbling)
Translational kinetic energy (flying around) 5

6. Kinetic Theory of s, l & g. When it is cold, molecules move slowly
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Kinetic Theory of s, l & g.
When it is cold, molecules move slowly
In solids, they move so slowly that they are held in place (only vibrational energy)
In liquids they move a bit faster, they can tumble and flow, but they don’t escape from the intermolecular attraction with other molecules (mostly rotational energy, some vibration & translation)
In gases they move so quickly that can overcome the intermolecular attractions and leave the container (more translational energy, with a little bit of rotation & vibration). 6

7. Plasma, the “Fourth State”
When strongly heated, or exposed to high voltage or radiation, gas atoms may lose some of their electrons. As they capture new electrons, the atoms emit light—they glow. This glowing, gas-like substance is called “plasma” 7

8. Properties of gases: Gases: can be atoms or molecules
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Properties of gases:
Gases:
can be atoms or molecules
I Have No Bright Or Clever Friends
I2 H2 N2 Br2 O2 Cl2 F2 are all diatomic gases,
have mass,
no definite volume or shape,
are compressible & can expand.
Therefore properties of gasses can only be compared under specific conditions.
Gas particles are very spaced out!

9. Find the properties of some common gases: Finish for homework!
Read pages from your textbook.
For N2, O2, CO2, radon & methane gas:
List their:
Abundance
General use (Do we breath it? Do plants use it etc)
Technological applications
For O3 how is it a useful and harmful gas?

10. Fun Gases (of no real importance)
Nitrous Oxide (N2O)
AKA: Laughing gas, Happy gas, Nitro, NOS
Uses
anaesthetic in dentist offices, this sweet-smelling gas reduces pain sensitivity and causes euphoric sensations.
It is an excellent oxidizer, reigniting a glowing splint much like oxygen would.
It is used in racing where it is injected into the carburetor to temporarily increase an engine’s horsepower.
Sulfur Hexafluoride
One of the densest gases in common use. Fun with Sulfur hexafluoride 10

11. Match the gas with the problem it causes
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Match the gas with the problem it causes
Gas Problem
Carbon Dioxide Ozone layer depletion
CFCs Climate Change
Methane Toxic poisoning
Carbon monoxide Noxious smell
Sulfur dioxide Acid Rain

12. Last class: We talked about the motion of molecules:
Vibrational
Rotational
Translational
Gasses have mostly translational motion.
As you increase temperature their motion also increases.
Increase in translational movement = increase in velocity

13. Pressure in gases: A force applied over a unit of area.
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Pressure in gases:
A force applied over a unit of area.
For a gas, pressure results from gas molecules colliding with the wall of its container.
Measured in Pascals (Pa) or kiloPascals (kPa)

14. If container is not strong enough walls can rupture
Adding a gas:
Adds more gas molecules
More collisions
Increased net pressure
Ex. Double # of molecules =
double pressure
If container is not strong enough walls can rupture

15. If container is not strong enough walls can collapse
Removed a Gas:
Removes gas molecules
Less collisions
Pressure decreases
If container is not strong enough walls can collapse

16. Change Size of Container:
Decrease container size
Decreases space for molecules to move
Increases collisions
Increases pressure

17. Change Size of Container:
Increase container size
Increases space for molecules to move
Decreases collisions
Decreases pressure

18. Heating a Gas: Gas molecules absorb heat Molecules move more rapidly
Increase collisions
Increase pressure

19. Cooling a Gas: Gas molecules release heat Molecules move more slowly
Decrease collisions
Decrease pressure

20. Write this!
Gases exert a pressure as they collide with the walls of containers. The total pressure is dependent on magnitude & quantity of collisions.
Concentration:
Add more gas  Pressure
Remove gas  Pressure
Container size:
decrease  Pressure
increase  Pressure
Temperature:
Increase Temp  Pressure
Decrease Temp  Pressure

21. Kinetic Molecular Theory (K.M.T): A gas is composed of particles
Gas particles move rapidly & are in constant random motion
All collisions are perfectly elastic
Kinetic energy is proportional to temperature
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22. Kinetic energy & temperature:
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As temperature increases molecules move faster & have a greater KE.
Not all molecules are moving at the same speed.
The KE of moving objects is expressed by:
This shows that the KE of molecules is dependent on both their mass & velocity.

23. The mean & mode can help establish “average” molecules
The range of kinetic energies can be represented as a “bell curve.” Maxwell’s Velocity Distribution Curve.
The mean & mode can help establish “average” molecules
Most molecules
“Average”
molecules
Increasing # molecules
“Slow”
molecules
“Fast”
Molecules
mode
mean
Average
kinetic energy
Increasing kinetic energy

24. Distribution of Particles Around Average Kinetic Energies.
Conclusion:
As temperature increases.
Curve broadens.
Average KE increases.
Average kinetic energy of molecules
Average kinetic energy of warmer molecules
Average kinetic energy of colder molecules
Number of molecules
Slower
than
average molecules
Faster
than
average molecules
Kinetic Energy of molecules
(proportional to velocity of molecules)

25. Two different gases at the same temperature will:
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Two different gases at the same temperature will:
Have the same AVERAGE EK= ½ mv2
Lighter molecules will move faster.
Heavier molecules will move slower.
Fun Fact
The average speed of oxygen molecules at 20°C is 1656km/h.
At that speed an oxygen molecule could travel from Montreal to Vancouver in three hours…If it travelled in a straight line.

26. Observing gases
As scientists observed gases, they saw mathematical relationships that very closely, but not perfectly, described the behaviour of many gases.
They have developed theories & mathematical laws that describe a hypothetical gas, called “ideal gas.”
It is an approximation that helps us model and predict the behavior of real gases.

27. Kinetic Theory for ideal gases.
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The particles of a gas are infinitely small.
Explains effusion & compressibility.
The particles of a gas are in constant motion, and move in straight lines.
Until they run into another particle or wall.
Explains diffusion.
The particles do not attract or repel each other.
Explains why gases expand to fill a space.
The average kinetic energy of particles is proportional to the absolute temperature.
Explains observed changes in pressure.
Fun fact 
Each air molecule has about ten billion collisions per second 27

28. Textbook questions
Do Page 62 # 3,4,5,8,9,10,11 To be done for next class

29. And in my spare time I invented dialysis, which has saved the lives of thousands of kidney patients
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Thomas Graham ( )
Graham studies the speed of diffusion & effusion.
Diffusion is when gas molecules spread throughout a container until they are evenly distributed
Effusion is when gas molecules pass through tiny opening in container.
He derived his law on Ek = ½ mv2
m = mass (kg) v = velocity (m/s)

30. The equation shows the ratio of Gas 1’s speed to Gas 2’s speed.
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Graham’s Law
Rate of diffusion of a gas is inversely related to the square root of its molar mass.
The equation shows the ratio of Gas 1’s speed to Gas 2’s speed.
v = velocity M = molar mass
(Leave a space for one more variation!)
Same as

31. Relative rate means find the ratio “v1/v2”!
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Ex. 1
Determine the relative rate of diffusion for krypton and bromine.
Relative rate means find the ratio “v1/v2”!
Ans: Kr diffuses times faster than Br2.

32. Put the gas with the unknown speed as “Gas 1”.
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Ex. 2
A molecule of oxygen gas has an average speed of 12.3 m/s at a given temp and pressure. What is the average speed of hydrogen molecules at the same conditions?
Put the gas with the unknown speed as “Gas 1”.
C. Johannesson

33. Square both sides to get rid of the square root sign.
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Ex. 3
An unknown gas diffuses 4.0 times faster than O2. Find its molar mass.
The ratio “v1/v2” is 4.0.
Square both sides to get rid of the square root sign.
C. Johannesson

34. The equation shows the ratio of Gas 1’s speed to Gas 2’s speed.
Graham’s Law
Rate of diffusion of a gas is inversely related to the square root of its molar mass.
The equation shows the ratio of Gas 1’s speed to Gas 2’s speed.
v = velocity M = molar mass
Same as

35. Graham’s Law Version #2, Effusion Time
Sometimes it’s easier to measure the time it takes for a gas to effuse completely, rather than the speed.
Graham’s law can be changed for this, but the relationship between time and molar mass is direct as the square root:
Add the equation!
Careful the relationship is not inverted!

36. Kinetic Theory Trivia
The average speed of oxygen molecules at 20°C is 1656km/h.
At that speed an oxygen molecule could travel from Montreal to Vancouver in three hours…If it travelled in a straight line.
Each air molecule has about 1010 (ten billion) collisions per second
10 billion collisions every second means they bounce around a lot!
The number of oxygen molecules in a classroom is about:
that’s more than there are stars in the universe!
The average distance air molecules travel between collisions is about 60nm.
m is about the width of a virus. 36

37. Videos Kinetic Molecular Basketball Average Kinetic Energies
Average Kinetic Energies
Thermo-chemistry lecture on kinetics 37