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Chapter 14“The Behavior of Gases”
Pre-AP Chemistry
Charles Page High School
Stephen L. Cotton
Section 14.1The Properties of Gases
OBJECTIVES:Explain why gases are easier to compress than solids or liquids are.
Section 14.1The Properties of Gases
OBJECTIVES:Describe the three factors that affect gas pressure.
CompressibilityGases can expand to fill its
container, unlike solids or liquidsThe reverse is also true:
They are easily compressed, or squeezed into a smaller volume
Compressibility is a measure of how much the volume of matter decreases under pressure
Compressibility This is the idea behind placing “air
bags” in automobilesIn an accident, the air compresses
more than the steering wheel or dash when you strike it
The impact forces the gas particles closer together, because there is a lot of empty space between them
Compressibility At room temperature, the distance
between particles is about 10x the diameter of the particleFig. 14.2, page 414
This empty space makes gases good insulators (example: windows, coats)
How does the volume of the particles in a gas compare to the overall volume of the gas?
Variables that describe a Gas The four variables and their common
units:
1. pressure (P) in kilopascals
2. volume (V) in Liters
3. temperature (T) in Kelvin
4. amount (n) in moles
• The amount of gas, volume, and temperature are factors that affect gas pressure.
1. Amount of GasWhen we inflate a balloon, we are
adding gas molecules. Increasing the number of gas
particles increases the number of collisionsthus, the pressure increases
If temperature is constant, then doubling the number of particles doubles the pressure
Pressure and the number of molecules are directly related
More molecules means more collisions, and…
Fewer molecules means fewer collisions.
Gases naturally move from areas of high pressure to low pressure, because there is empty space to move into – a spray can is example.
Common use? A practical application is Aerosol
(spray) cansgas moves from higher pressure to
lower pressurea propellant forces the product outwhipped cream, hair spray, paint
Fig. 14.5, page 416 Is the can really ever “empty”?
2. Volume of Gas In a smaller container, the
molecules have less room to move.
The particles hit the sides of the container more often.
As volume decreases, pressure increases. (think of a syringe)Thus, volume and pressure are
inversely related to each other
3. Temperature of Gas Raising the temperature of a gas increases
the pressure, if the volume is held constant. (Temp. and Pres. are directly related)The molecules hit the walls harder, and
more frequently! Fig. 14.7, page 417 Should you throw an aerosol can into a
fire? What could happen? When should your automobile tire pressure
be checked?
Section 14.2The Gas Laws
OBJECTIVES:OBJECTIVES:
Describe the relationships among the temperature, pressure, and volume of a gas.
Section 14.2The Gas Laws
OBJECTIVES:OBJECTIVES:
Use the combined gas law to solve problems.
The Gas Laws are mathematicalThe gas laws will describe HOW
gases behave.Gas behavior can be predicted by
the theory.The amount of change can be
calculated with mathematical equations.
You need to know both of these: the theory, and the math
Robert Boyle(1627-1691)
• Boyle was born into an aristocratic Irish family
• Became interested in medicine and the new science of Galileo and studied chemistry.
• A founder and an influential fellow of the Royal Society of London
• Wrote extensively on science, philosophy, and theology.
#1. Boyle’s Law - 1662
Pressure x Volume = a constant
Equation: P1V1 = P2V2 (T = constant)
Gas pressure is inversely proportional to the volume, when temperature is held constant.
Graph of Boyle’s Law – page 418
Boyle’s Law says the pressure is inverse to the volume.
Note that when the volume goes up, the pressure goes down
- Page 419
Jacques Charles (1746-1823)• French Physicist• Part of a scientific
balloon flight on Dec. 1, 1783 – was one of three passengers in the second balloon ascension that carried humans
• This is how his interest in gases started
• It was a hydrogen filled balloon – good thing they were careful!
#2. Charles’s Law - 1787The volume of a fixed mass of gas is directly proportional to the Kelvin temperature, when pressure is held constant.
This extrapolates to zero volume at a temperature of zero Kelvin.
VT
VT
P1
1
2
2 ( constant)
Converting Celsius to Kelvin•Gas law problems involving temperature will always require that the temperature be in Kelvin. (Remember that no degree sign is shown with the kelvin scale.)
•Reason? There will never be a zero volume, since we have never reached absolute zero.
Kelvin = C + 273 °C = Kelvin - 273and
- Page 421
Joseph Louis Gay-Lussac (1778 – 1850) French chemist and physicist Known for his studies on the physical properties of gases. In 1804 he made balloon ascensions to study magnetic forces and to observe the composition and temperature of the air at different altitudes.
#3. Gay-Lussac’s Law - 1802•The pressure and Kelvin temperature of a gas are directly proportional, provided that the volume remains constant.
2
2
1
1
T
P
T
P
•How does a pressure cooker affect the time needed to cook food? (Note page 422)
•Sample Problem 14.3, page 423
#4. The Combined Gas LawThe combined gas law expresses the relationship between pressure, volume and temperature of a fixed amount of gas.
2
22
1
11
T
VP
T
VP
Sample Problem 14.4, page 424
The combined gas law contains all the other gas laws!
If the temperature remains constant...
P1 V1
T1
x=
P2 V2
T2
x
Boyle’s Law
The combined gas law contains all the other gas laws!
If the pressure remains constant...
P1 V1
T1
x=
P2 V2
T2
x
Charles’s Law
The combined gas law contains all the other gas laws!
If the volume remains constant...
P1 V1
T1
x=
P2 V2
T2
x
Gay-Lussac’s Law
Section 14.3Ideal Gases
OBJECTIVES:OBJECTIVES:
Compute the value of an unknown using the ideal gas law.
Section 14.3Ideal Gases
OBJECTIVES:OBJECTIVES:
Compare and contrast real an ideal gases.
5. The Ideal Gas Law #1 Equation: P x V = n x R x T Pressure times Volume equals the
number of moles (n) times the Ideal Gas Constant (R) times the Temperature in Kelvin.
R = 8.31 (L x kPa) / (mol x K) The other units must match the value of
the constant, in order to cancel out. The value of R could change, if other
units of measurement are used for the other values (namely pressure changes)
We now have a new way to count moles (the amount of matter), by measuring T, P, and V. We aren’t restricted to only STP conditions:
P x V R x T
The Ideal Gas Law
n =
Ideal Gases We are going to assume the gases
behave “ideally”- in other words, they obey the Gas Laws under all conditions of temperature and pressure
An ideal gas does not really exist, but it makes the math easier and is a close approximation.
Particles have no volume? Wrong! No attractive forces? Wrong!
Ideal GasesThere are no gases for which this
is true (acting “ideal”); however,Real gases behave this way at
a) high temperature, and b) low pressure.Because at these conditions, a gas will stay a gas!
Sample Problem 14.5, page 427
#6. Ideal Gas Law 2 P x V = m x R x T
M Allows LOTS of calculations, and some
new items are: m = mass, in grams M = molar mass, in g/mol
Molar mass = m R T P V
Density Density is mass divided by volume
m
V
so,
m M P
V R T
D =
D = =
Ideal Gases don’t exist, because:
1. Molecules do take up space
2. There are attractive forces between particles
- otherwise there would be no liquids formed
Real Gases behave like Ideal Gases...
When the molecules are far apart.
The molecules do not take up as big a percentage of the space We can ignore the particle
volume. This is at low pressure
Real Gases behave like Ideal Gases…
When molecules are moving fastThis is at high temperature
Collisions are harder and faster.Molecules are not next to each
other very long.Attractive forces can’t play a role.
Section 14.4Gases: Mixtures and Movements
OBJECTIVES:
Relate the total pressure of a mixture of gases to the partial pressures of the component gases.
Section 14.4Gases: Mixtures and Movements
OBJECTIVES:
Explain how the molar mass of a gas affects the rate at which the gas diffuses and effuses.
#7 Dalton’s Law of Partial Pressures
For a mixture of gases in a container,
PTotal = P1 + P2 + P3 + . . .
•P1 represents the “partial pressure”, or the contribution by that gas.•Dalton’s Law is particularly useful in calculating the pressure of gases collected over water.
Collecting a gas over water – one of the experiments in Chapter 14 involves this.
Connected to gas generator
If the first three containers are all put into the fourth, we can find the pressure in that container by adding up the pressure in the first 3:
2 atm + 1 atm + 3 atm = 6 atm
Sample Problem 14.6, page 434
1 2 3 4
Diffusion is:
Effusion: Gas escaping through a tiny hole in a container.
Both of these depend on the molar mass of the particle, which determines the speed.
Molecules moving from areas of high concentration to low concentration.Example: perfume molecules spreading across the room.
•Diffusion: describes the mixing of gases. The rate of diffusion is the rate of gas mixing.
•Molecules move from areas of high concentration to low concentration.
•Fig. 14.18, p. 435
Effusion: a gas escapes through a tiny hole in its container
-Think of a nail in your car tire…
Diffusion and effusion are explained by the next gas law: Graham’s
8. Graham’s Law
The rate of effusion and diffusion is inversely proportional to the square root of the molar mass of the molecules.
Derived from: Kinetic energy = 1/2 mv2
m = the molar mass, and v = the velocity.
RateA MassB
RateB MassA
=
Sample: compare rates of effusion of Helium with Nitrogen – done on p. 436
With effusion and diffusion, the type of particle is important: Gases of lower molar mass diffuse and
effuse faster than gases of higher molar mass.
Helium effuses and diffuses faster than nitrogen – thus, helium escapes from a balloon quicker than many other gases!
Graham’s Law
