1. Chapter 2 p.63-67
Behavior of Gases

2. Behavior of Gases
The behavior of gases refers to the way gases react to different conditions.
The 4 conditions we will now look at are:
1) Compressibility
2) Expansion
3) Diffusion
4) Pressure

3. Gases can expand to fill their container, unlike solids or liquids
Compressibility
Gases can expand to fill their container, unlike solids or liquids
The reverse is also true:
They are easily compressed, or squeezed into a smaller volume.
For example, the compressed air used by divers contains 6-18 litres of compressed air at about 200 times the normal atmospheric pressure.

4. Expansion
Gases do not have a definite shape or volume, so they can expand to fit the volume available to them.

5. Diffusion
Gas particles constantly collide with each other but these collisions are random.
When a gas is introduced into a container, it will diffuse (mix) through out the container even if there are other gases present. This can be a slow process.

6. Pressure
Gases exert pressure on the objects they come in contact with.
To calculate the pressure use the following formula:
P = F A
P= pressure in pascals (Pa), F is force in Newtons(N), A is the area (m2)

7. Real Gases
and
Ideal Gases

8. 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!

9. Ideal Gases don’t exist, because:
Molecules do take up space
There are attractive forces between particles

10. 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.

11. 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

12. 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.

13. 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.

14. 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.

15. 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.

16. 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.
Thus, volume and pressure are inversely related to each other

17. 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!
Should you throw an aerosol can into a fire? What could happen?

18. Robert Boyle (1627-1691) do not copy
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.

19. Boyle’s Law
Gas pressure is inversely proportional to the volume, when temperature is held constant.
Equation: P1V1 = P2V2 (T = constant)

20. Graph of Boyle’s Law
Boyle’s Law says the pressure is inverse to the volume.
Note that when the volume goes up, the pressure goes down

21. Practice Question P1V1 = P2V2 (250)*(20.0) = (400) * V2 12.5 ml = V2
A chemistry collects 20.0 ml of a gas at a pressure of 250 kPa. What will be the volume of the sample of gas if the pressure in increased to 400kPa.
P1V1 = P2V2
(250)*(20.0) = (400) * V2
12.5 ml = V2

22. Jacques Charles ( )
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!
DO NOT COPY

23. Charles’s Law
The volume of a fixed mass of gas is directly proportional to the Kelvin temperature, when pressure is held constant.

24. How Volume Varies With Temperature
If we place a balloon in liquid nitrogen it shrinks:
So, gases shrink if cooled.
Conversely, if we heat a gas it expands (as in a hot air balloon).

25. Temperature vs. Volume Graph5 10 15 20 25 30
Volume (mL)
Temperature (C)
100
– 273

26. The Kelvin Temperature Scale
If a volume vs. temperature graph is plotted for gases, most lines can be interpolated so that when volume is 0 the temperature is -273 C.
Naturally, gases don’t really reach a 0 volume, but the spaces between molecules approach 0.
At this point all molecular movement stops.
–273C is known as “absolute zero” (no EK)
Lord Kelvin suggested that a reasonable
temperature scale should start at a true zero value.

27. 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.)
Kelvin = C + 273

28. Practice Question
A sample of gas occupies 3.5 L at 300 K. What volume will it occupy at 200 K?
V1 = 3.5 L, T1 = 300K, V2 = ?, T2 = 200K
Using Charles’ law: V1/T1 = V2/T2
3.5 L / 300 K = V2 / 200 K
V2 = (3.5 L/300 K) x (200 K) = 2.3 L

29. Joseph Louis Gay-Lussac (1778 – 1850) do not copy
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.

30. Gay-Lussac’s Law
The pressure and Kelvin temperature of a gas are directly proportional, provided that the volume remains constant.
How does a pressure cooker affect the time needed to cook food?

31. The Combined Gas Law

32. These are all subsets of a more encompassing law: the combined gas law
Combining the gas laws
So far we have seen these gas laws:
Robert Boyle
Jacques Charles
Joseph Louis Gay-Lussac V1 T1 = V2 T2 P1 T1 = P2 T2
P1V1 =
P2V2
These are all subsets of a more encompassing law: the combined gas law

33. The Ideal Gas Law R = 8.31 (L x kPa) / (mol x K)
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)

34. How many grams of Cl2(g) can be stored in a 10
How many grams of Cl2(g) can be stored in a 10.0 L container at 1000 kPa and 30°C?
PV = nRT
P= 1000 kPa, V= 10.0 L, T= 303 K
(8.31 kPa•L/K•mol)(303 K)
(1000 kPa)(10.0 L)
= n = 3.97 mol
3.97 mol x 70.9 g/mol = 282 g
At 150°C and 100 kPa, 1.00 L of a compound has a mass of g. Calculate molar mass.
PV = nRT
P= 100 kPa, V= 1.00 L, T= 423 K
(8.31 kPa•L/K•mol)(423 K)
(100 kPa)(1.00 L)
= n = mol
g/mol = g / mol = 88.1 g/mol

35. Dalton's Law of Partial Pressures

36. Dalton’s Law of Partial Pressures
For a mixture of gases in a container,
PTotal = P1 + P2 + P
P1 represents the “partial pressure” of gas 1...
Dalton’s Law is particularly useful in calculating the pressure of gases collected over water.

37. Connected to gas generator
Collecting a gas over water – one of the experiments in this chapter involves this.

38. 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

39. Diffusion: describes the mixing of gases
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.

40. 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

41. 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.

42. Practice Question #1
A molecule of CH4 diffuses a distance of m/sec from a point source. An unknown gas diffuses m/sec under identical conditions. You are told the gas could be O2, N2, CO2, or Ne.
Identify the unknown gas. Show all your work.

43. Graham’s Law (Question 2)
Compare rates of effusion of Helium with Nitrogen.
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!