1. Chapter The
14 Behavior of
Gases

2. Section 14.1
The Properties of Gases
 OBJECTIVES:
Explain why gases are
easier to compress than
solids or liquids are.

3. Section 14.1
The Properties of Gases
 OBJECTIVES:
Describe the three factors
that affect gas pressure.

4. Compressibility
 Gases can expand to fill its
container, unlike solids or liquids
 The 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

5. Compressibility
 Thisis the idea behind placing “air
bags” in automobiles
 In 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

6. Compressibility
 Atroom temperature, the distance
between particles is about 10x the
diameter of the particle
 Fig. 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?

7. 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.

8. 1. Amount of Gas
 When we inflate a balloon, we are
adding gas molecules.
 Increasing the number of gas
particles increases the number of
collisions
thus, the pressure increases
 If temperature is constant, then
doubling the number of particles
doubles the pressure

9. 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.

10. Common use?
A practical application is Aerosol
(spray) cans
 gas moves from higher pressure to
lower pressure
 a propellant forces the product out
 whipped cream, hair spray, paint
 Fig. 14.5, page 416
 Is the can really ever “empty”?

11. 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

12. 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?

14. Section 14.2
The Gas Laws
 OBJECTIVES:
Describe the relationships
among the temperature,
pressure, and volume of a
gas.

15. Section 14.2
The Gas Laws
 OBJECTIVES:
Use the combined gas
law to solve problems.

16. The Gas Laws are mathematical
 Thegas 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

17. Robert Boyle • Boyle was born into an
aristocratic Irish family
(1627-1691)
• 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.

18. #1. Boyle’s Law - 1662
Gas pressure is inversely proportional to the
volume, when temperature is held constant.
Pressure x Volume = a constant
Equation: P1V1 = P2V2 (T = constant)

19. 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

20. - Page 419

21. 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!

22. #2. Charles’s Law - 1787
The 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.
V1 V2
= ( P = constant)
T1 T2

23. 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 and °C = Kelvin - 273

24. - Page 421

25. 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.

26. #3. Gay-Lussac’s Law - 1802
•The pressure and Kelvin temperature of
a gas are directly proportional, provided
that the volume remains constant.
P P2
1
=
T1 T2
•How does a pressure cooker affect the time
needed to cook food? (Note page 422)
•Sample Problem 14.3, page 423

27. #4. The Combined Gas Law
The combined gas law expresses the
relationship between pressure, volume
and temperature of a fixed amount of
gas.
PV1 P2V2
1
=
T1 T2
Sample Problem 14.4, page 424

28.  The combined gas law contains
all the other gas laws!
 If the temperature remains
constant...
P1 x V1 P2 x V2
=
T1 T2
Boyle’s Law

29.  The combined gas law contains
all the other gas laws!
 If the pressure remains constant...
P1 x V1 P2 x V2
=
T1 T2
Charles’s Law

30. The combined gas law contains
all the other gas laws!
If the volume remains constant...
P1 x V1 P2 x V2
=
T1 T2
Gay-Lussac’s Law

31. Section 14.3
Ideal Gases
 OBJECTIVES:
Compute the value of an
unknown using the ideal
gas law.

32. Section 14.3
Ideal Gases
 OBJECTIVES:
Compare and contrast real
an ideal gases.

33. 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

34. The Ideal Gas Law
 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:
PxV
n=
RxT

35. 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!

36. Ideal Gases
 There 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

37. #6. Ideal Gas Law 2
Equation: P x V = mxRxT
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
PV

38. Density
 Density is mass divided by volume
m
D=
V
so,
m MP
D= =
V RT

40. 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

41. 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

42. Real Gases behave like Ideal Gases…
 When molecules are moving fast
This 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.

43. Section 14.4
Gases: Mixtures and Movements
 OBJECTIVES:
Relate the total pressure
of a mixture of gases to
the partial pressures of the
component gases.

44. Section 14.4
Gases: Mixtures and Movements
 OBJECTIVES:
Explain how the molar
mass of a gas affects the
rate at which the gas
diffuses and effuses.

45. #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.

46. Connected
to gas
generator
Collecting a gas over water

47.  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
1 2 3 4

48. Diffusion is:
Molecules moving from areas of high
concentration to low concentration.
Example: perfume molecules spreading
across the room.
 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.

49. •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.

50. 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

51. 8. Graham’s Law
RateA √ MassB
=
RateB √ MassA
 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.

52. Graham’s Law
 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