1. What are the properties of gases?
Chapter 13 “Gases”
What are the properties of gases?

2. Chapt. 13 OBJECTIVES Describe gas behavior & explain how the
Kinetic-Molecular (K-M) Theory accounts for it.
Explain what ‘gas pressure’ means & describe how it is measured.
State the “Gas Laws” and use them to solve problems.
Understand the significance of the “Ideal Gas Law.”
Compare ideal & real gases.
Relate gas density to temperature and molar mass.

3. What are the Physical Properties of Gases?
Gases have mass.
It is easy to compress gases.
Gases fill their containers completely.
Different gases quickly move through each other. (This mixing is called diffusion.)
Gases exert pressure.
The pressure of a gas depends on its temperature.

4. How do we explain gas behavior?
Gas properties are explained by a model, called the kinetic-molecular model, that describes the behavior of the subatomic particles that make up a gas.
This is known as “The Kinetic Molecular Theory” or “K-M” Theory.

5. The Kinetic-Molecular (K-M) Theory
Gases are composed of very small particles (atoms or molecules) having mass.
Gas particles have low volume, but are spaced widely, with lots of empty space between them, making gases easy to compress.
There is no attraction or repulsion among the gas particles.
Gas particles move in constant, rapid, random motion, so a gas has no definite shape, but it fills its container completely.
The gas particles travel in short straight lines, with each particle moving independently.
When colliding with walls, other particles, etc., the gas particles rebound elastically, so the total kinetic energy (KE) is constant at a given temperature.
The average kinetic energy of the molecules of a gas is proportional to the temperature of the gas in kelvins.

6. 13-2 Measuring Gases
To completely describe a gas sample and to make predictions about its behavior under new conditions, we must deal with four variables:
Amount of gas (n, in moles)
n = mass/molar mass = m(g)/M(g/mol)
Volume of gas (V, in liters)
1 L = 1000mL = 1000 cm3
Temperature (T, in kelvins)
T(K) = T(°C) + 273
Pressure (P, in atm, torr, kPa, etc.)

7. How to Measure Pressure
Pressure is the force exerted by gas particles on the walls of the container.
Pressure = force/area (newtons/meter2 = pascals)
Atmospheric pressure results from the mass of the air being attracted by the Earth’s gravity.
Conversions:
1 atmosphere (atm) = kPa
1 atmosphere = 760 mm Hg (same as torr)
1 atmosphere = lb/in2
(See Fig , pg 426 for inter-conversions.)

8. Tools for Measuring Pressure
Barometer
Atmospheric pressure used to predict the weather
Pressure gauge
Tire pressure gauges or gas cylinder regulators
Manometer
Laboratory tools for working with gases in closed space
Examples: reading manometers
Do you remember STP?
STP = Standard Temperature & Pressure
STP = 273 K (0°C) & 1 atm (760 torr, etc.)
It is important for measuring gas properties.

9. 13-3 The Gas Laws (1) What happens when you put pressure on a gas?
Examples: Small balloon being squeezed, books piled on a small cylinder.
Robert Boyle (1600s) studied the effect of pressure (P) on the volume (V) of a gas.
Worksheet class activity
What is the relationship between V and P?
Boyle’s Law: The pressure and volume of a sample of gas are inversely proportional to each other at constant temperature.

10. Boyle’s Law

11. 13-3 The Gas Laws (2)
What happens when you change the temperature of a gas?
Examples: Hot air balloon inflating, automobile tires in very cold weather.
Jacques Charles (1700s) studied the effect of temperature (T) on the volume (V) of a gas.
Worksheet class activity
What is the relationship between V and T?
Charles’ Law: The volume of a sample of gas is directly proportional to the kelvin temperature at constant pressure.
Lab Activity: Predicting Absolute Zero

12. Charles’ Law

13. 13-3 The Gas Laws (3)
What happens to pressure of a fixed volume of gas if you change the temperature?
Example: “Empty” aerosol can thrown into an incinerator. (Read the warning on the label!)
Gay-Lussac (1700s) studied the effect of pressure (P) and temperature (T) on a fixed volume (V) of gas.
Worksheet class activity
What relationship do you see between P and T?
Gay-Lussac’s Law: The pressure and temperature of a fixed volume of gas are directly proportional to each other.

14. Gay-Lussac’s Law

15. 13-3 The Gas Laws (4)
How is the number of gas particles related to its volume under constant conditions?
Examples: Two balloons of different size.
Amedeo Avogadro (1800s) studied different gases to determine the relationship between the number of gas particles and the volume at a given pressure (P) and temperature (T).
Worksheet class activity
What is the relationship between V and n?
Avogadro’s Law: Equal volumes of gases at the same pressure and temperature contain an equal number of particles.

16. Avogadro’s Law Equal volumes of different gases have the same
number of particles under the same conditions.

17. 13-3 The Gas Laws (Combined)
These Gas Laws seem complicated, but fortunately we
can simplify things by combining the relationships into two
simple expressions, the one first being:
P1V1
P2V2 = T1 T2
This equation is used to solve “Combined Gas Law” problems, by simply
“plugging in” the numbers!
Just remember two things:
“Go Kelvin!” (Convert temperature to kelvins.)
Be sure all units are consistent for P and V.

18. 13-4 The “Ideal Gas Law”
The second equation that may be derived from the various gas laws is the “Ideal Gas Equation” PV = nRT
This describes the relationship among the four variables (P,V,n,T) of an ideal gas, where R is the gas constant.
R = atm-L/mol-K
R = Pa-m3/mol-K
R = J/mol-K
Problems may be solved by simple substitution, but be careful to use consistent units.

19. But what is an “Idea Gas”?
Ideal Gas: one that is described by the postulates of the Kinetic-Molecular Theory.
No such gases exist! The gases we encounter are “real” – not “ideal”!
However, most gases behave like ideal gases under ordinary conditions of temperature and pressure.
At low temperature and high pressure real gases behave in non-ideal ways. (Why is this so? See the next slide!)

20. Deviations from Ideal Gas Behavior Occur Because:
Kinetic-Molecular Theory makes two simplifying, but WRONG, assumptions:
Gas particles have no volume of their own.
Gas particles have no attraction for each other.
But, as pressure increases…
gas particles get closer together.
At very high pressure, the volume of the gas particles themselves become a significant part of the total volume, contrary to K-M Theory.
As temperature decreases…
gas particles slow down.
At very low temperature, attractive forces between gas particles become significant, contrary to K-M Theory.

21. Nevertheless…
K-M Theory is valid for studying gas behavior under ordinary conditions.
We are safe to use the Ideal Gas Equation (PV = nRT) to solve problems for real gases, except at low temperature and high pressure conditions where gases start to behave in a non-ideal manner.

22. 13-5 How Gases Work: Lift Lifting Power
This is the result of low density of a gas, but the density must be lower than that of the air.
Low molar mass gases (H2, He, NH3, CH4) may be used to gain lift.
Or the gas, usually air, may be heated.
Higher mass gases (CO2, Kr) are too dense, so they sink in air.
Can have disastrous effects.

23. 13-5 How Gases Work: Effusion
Related to ‘diffusion’ (one gas moving through another).
Effusion is the gas movement through a tiny hole one particle at a time.
Small, light gases have greater speeds than large, heavy gases at a given temperature, so they effuse faster.
Graham’s Law of Effusion relates effusion to mass.

24. Graham’s Law of Effusion
Gases will effuse at a rate that is inversely proportional
to the square root of the Molecular Masses, MMx, or:
rate1 √
MM2 =
rate2
MM1
See worksheet problems.