1. Chapter 10 & 11: Gases Chapter 10: Page 300-330 Chapter 11:
Chlorine gas was used as a weapon in WWI

2. Kinetic Molecular Theory
all matter is made up of particles (atoms) in random and constant motion. (colliding)
Gases have very low density
particles are spaced far apart.
Gases are compressible.
Extreme pressures-gases will compress until they become liquids (or solids, CO2).
Adding heat to a system
increases the temperature …
Temperature = measure of the average kinetic energy of the particles.
Increasing the pressure of a gas,
increases the density of the gas - the number of particles in a given space.

3. “Ideal” Gases
Ideal gasses
Make a two-column table
When we talk about gases, we refer to hypothetical “ideal” gases, for simplicity of calculations.
In reality, “real” gases behave different than ideal gases. Some differences:
When compressed, real gases will form liquids, and even exhibit liquid-like behaviors when still in gas form.
Real gas molecules interact with each other -causing them to travel in non-linear paths and collide “inelastically.”
With real gases, the size of the gas molecules effects their behavior.
Real Gasses 

4. Gases in our atmosphere
Nitrogen-78%
Oxygen-21%
Argon-<1%
Trace amounts of CO2, Ne, He, CH4, Kr, H2, O3, and others.
Some gases function as greenhouse gases, and work to hold heat on the earth’s surface.
Some gases function to block harmful UV radiation energy from the sun.

5. The Greenhouse Effect
The sun’s energy travels through space and warms the surface of the earth.
Some of the energy is reflected back into space.
Greenhouse Gases
trap heat that would leave the atmosphere.
H2O, CH4, and CO2 are common greenhouse gases.
“Global Warming”
Theory that increasing levels of Greenhouse gasses is causing the global average temps to increase.

6. The Ozone Layer (O3) Ozone is Ozone
Page 778 for more info
Ozone is
a corrosive poison in the Troposphere (where we live)
frequently created and given off from free electrical ionization.
Ozone
Absorbs harmful Ultraviolet (UV) energy in the stratosphere, 11km (6 miles) above us.
Note: the Ozone layer is less than 1mm thick!
It is always moving, like a cloud, due to weather patterns and climate variations.

7. Pascal’s Principle and Pressure
F = Force
a = area
French physician, Blaise Pascal, showed that
fluids (including gasses) exert a uniform pressure on all the surfaces that they contact.
Exerting a force on the top surface of a gas, causes that force (pressure) to be exerted on all the walls of its container.
Pressure is due to the particles of a gas striking a surface. We can detect pressure from billions upon billions of gas molecules striking a surface at any point in time.
 Which exerts a greater pressure? 

8. Pressure units…
The SI unit of pressure is the Pascal, Pa, equaling one newton per square meter.
Earth’s air pressure at sea level ~ 100,000 Pa. 100kPa
PSI (US)
Pound per square inch. Atmospheric pressure at sea level is about 14.5 PSI.
mmHg (EU, Asia) (AKA: Torr)
Millimeters of mercury.
Atmospheric pressure is 760 mmHg at sea level.
This has to due with the height of a column of liquid mercury raised in a barometer.
inHg
Inches of mercury. Used only in meteorology.
Atmospheric pressure is apx 30inHg.

9. Standard Temperature and Pressure:
And finally…
And, finally…the atmosphere, atm
the pressure exerted by the atmosphere at sea level, at 00C. (This creates STP…)
Standard Temperature and Pressure:
STP
usually used when referring to reactions with gases. STP is defined as:
1 atm and K
100 kPa and K
760 mmHg and K
When doing work with gases, select the STP that matches the pressure you are using. (atm in this class)

10. Charles’ Law
Simulation. constant volume
French chemist, Jacque Charles, showed that at constant pressure,
temperature and volume varied proportionally. That is…
V / T=k (k = some constant #)
We tend to write Charles’ Law as the volumes and temperatures under two conditions:
c 1780’s

11. Boyle’s Law P V = k (constant)
We’re leaving one law out… can you guess what it is?
Boyle’s Law
A young, adventurous, British aristocrat named Robert Boyle found that
when temperature is kept constant, volume varies inversely proportional with pressure. That is:
P V = k (constant)
We tend to write Boyle’s Law as the volumes and pressures under two conditions:
c 1660’s

12. Charles’ Law + Boyle’s Law + Avogadro’s Law = THE IDEAL GAS LAW
R is the “gas constant” and numerically depends upon the pressure units used.
Pressure
Volume
(in Liters)
Moles
Constant
Temperature
(in Kelvin)

13. The Gas Constant
The Gas Constant is the numerical bridge between number of moles of a gas, its temperature, and volume or pressure.
R = L٠kPa / mol٠K
R = L٠atm / mol٠K
Note that the first constant is in KILO Pascals. When given Pascals, you must first convert to kilopascals.
Our calculations will be done in Atm

14. Dalton’s Law of Partial Pressures
The total pressure in a system is the sum of the individual pressures exerted by each gas.
So, if gas A exerts a pressure of 2 units, and gas B exerts a pressure of 3 units, the total pressure of a system of equal parts of A and B, would be ?
Total = A + B …… = 5 units.
In our atmosphere, Oxygen is about 21%. If we have a sample of air at 1 atm, what is the pressure due to oxygen?

15. Graham’s Law of Gas Effusion
motion of a gas through a opening in a container.
Do not confuse Effusion with Diffusion - the tendency to disperse from an area of higher gas density to an area of lower gas density.
Rates of effusion are related to the molar mass of a gas.
The higher the molar mass, the slower the gas will effuse.
This is a property of real gases

16. Graham’s Law of Gas Effusion
We calculate effusion of a gas as a function of another gas’s know effusion rate.
The higher the molar mass, the slower the gas will effuse.
Graham’s Law of Effusion:
Molar mass
velocity
Molar mass
velocity
Gas A vs Gas B

17. Vapor Pressure All liquids exert a vapor pressure.
Page 324
All liquids exert a vapor pressure.
Vapor pressure is due to the liquid’s molecules entering the gas phase and leaving the liquid phase.
At higher temperatures, liquids will exert a greater vapor pressure.
More volatile liquids exert a greater vapor pressure than do less volatile liquids.
Can you explain why this is?
In lab: we collect gasses over water There is a small amount of water vapor in our gas samples, due to water’s vapor pressure.

18. Phase Diagrams
Example on page 381
Phase diagrams allow us to predict if a substance will be a solid, liquid or gas, depending upon the pressure and temperature of the substance.
Triple Point
point where solid, liquid, and gas all exist – for water, 00C.
Notice, that as you increase pressure, the boiling point of water increases-this is why a pressure cooker works.
What about Denver, the “mile-high city?”
End of Gases lecture, Chapters 10,11, problems following

19. In-chapter problems: Page 327, #5,7,8 What is Pressure?
Page 327, #11-14 Pressure Units
Page 327, #17-19 Pressure Conversions
Page 330, #20-24e Boyle’s Law
Page 330, #25-27 Charles’ Law
Page 330, #31-35o Combined Law
Page 331, #39,40 Dalton’s Law of Partial Pressures
Page 357, #9-13o Avogadro’s Molar Gasses
Page 358, #17-20 Ideal Gas Law
Page 358, #23-29o Ideal Gas Law and Stoichiometry
Page 359, #39-42 Graham’s Law of Gas Effusion
End of Gases Unit, Chapters 10,11

20. CCSD Syllabus Objectives
11.1: Kinetic Molecular Theory
11.2: Physical Properties of Gasses
11.3: STP
11.4: Volume-Temp relationships
11.5: Volume-Pressure relationships
11.6: Density-Volume-Pressure-Temperature
11.10: Ideal Gas Law
11.11: Graham’s Law
11.12: Ideal Gas vs Real Gas
12.3: Evaporation, Condensation, Sublimation
21.1: Environmental Chemistry