1. Unit 6 Gases, Phase Changes and Introduction to Thermochemistry
Part I: Gases
Characteristics of Gases
Pressure
Kinetic-Molecular Theory
The Gas Laws
Partial Pressures
Effusion and Diffusion
Real Gases

2. Properties of Gases Three phases of matter solid liquid
Definite shape and volume
Definite volume, shape of container
Shape and volume of container

3. Properties of Gases
A gas is a collection of molecules that are very far apart on average.
In air, gas molecules occupy only 0.1% of the total volume.
In liquids, molecules occupy ~ 70% of the total space.

4. Properties of Gases Gases are highly compressible.
Volume decreases when pressure is applied.
Gases form homogeneous mixtures with each other regardless of the identities or relative proportions of the different gases.
Water and gasoline = heterogeneous mixture.
Water vapor and gasoline vapor = homogeneous mixture.

5. Properties of Gases
Properties of gases vary depending on their composition.
Air: ~ 78% N2 and ~ 21% O2
CO2: colorless, odorless
CO: colorless, odorless, highly toxic
NO2: toxic, red-brown, irritant
N2O: colorless, sweet odor (laughing gas)

6. Pressure Four quantities are commonly needed to describe a gas:
amount of gas (n)
Temperature (T)
Volume (V)
Pressure (P)

7. Pressure Gases exert pressure on the objects in their surroundings.
Pressure is caused by collisions between the gas molecules and objects with which they are in contact.
Pressure: the force exerted on a unit area
P = F A

8. Pressure
Atmospheric pressure: the pressure exerted by gas molecules in the air on all objects exposed to the atmosphere
Atmospheric pressure varies with altitude.
Altitude
(ft above sea level)
Atmospheric Pressure
in. Hg
Torr
psi
29.92
760
14.7
5000
24.9
632.5
12.23
10,000
20.58
522.7
10.1

9. Pressure
Why does atmospheric pressure decrease with increasing altitude?
Gravity decreases
Density of gas decreases
Fewer gas molecules
Fewer collisions
Lower pressure

10. Must know units and abbreviations!!
Pressure
Many different units used to report pressure.
millimeters of Hg (mm Hg)
inches of Hg (in. Hg)
pounds per square inch (psi)
atmosphere (atm)
torr (torr)
pascal (Pa) = SI base unit
kilopascal (kPa)
Must know units and abbreviations!!

11. Pressure Relationships between different pressure units:
1 atm = 760 mm Hg
= 760 torr
= in. Hg
= 14.7 psi
= x 105 Pa
Must be able to interconvert between units.
Memorize the ones in red…I’ll give you the others.
You must know that 1 kPa = 1000 Pa

12. Pressure
Example: The measured pressure inside the eye of a hurricane was 669 torr. What was the pressure in atm?

13. Pressure
Example: On a nice sunny day in Chicago the barometric pressure was in. Hg. What was the pressure in Pa?

14. Pressure
Example: On Titan, the largest moon of Saturn, the atmospheric pressure is Pa. What is the pressure in atm?

15. Kinetic Molecular Theory
The behavior of gases can be described and explained using kinetic molecular theory.
the “theory of moving molecules”
You must know the basic ideas that are part of kinetic molecular theory.

16. Kinetic Molecular Theory
Gases consist of large numbers of molecules that are in continuous, random motion.
The combined volume of all the molecules of the gas is negligible compared to the total volume in which the gas is contained.
i.e. the molecules are very far apart on average

17. Kinetic Molecular Theory
Attractive and repulsive forces between gas molecules are negligible.
Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change as long as the temperature remains constant.
Collisions are perfectly elastic.

18. Kinetic Molecular Theory
The average kinetic energy of the molecules is proportional to the absolute temperature.
At any given temperature all molecules of a gas have the same average kinetic energy.
As T (in K) increases,
KE increases.

19. Gas Laws
Four variables are needed to define the physical condition or state of any gas:
Temperature (T)
Pressure (P)
Volume (V)
Amount of gas (moles: n)
Equations relating these variables are known as the gas laws.

20. Gas Laws
Consider a fixed amount of gas that is confined to a container with a certain volume. P
At a specific temperature, the gas sample will exert a certain pressure on the container.

21. Gas Laws What will happen to the pressure if the volume is decreased?
decreases

22. Gas Laws
As the volume of a fixed quantity of gas decreases, the pressure increases because:
gas molecules are more tightly packed together
i.e. denser
more collisions between gas molecules and the container
greater pressure

23. Gas Laws
Boyle’s Law:
The volume of a fixed quantity of gas maintained at constant temperature is inversely proportional to the pressure.
Mathematically,
V = k x or PV = k or P1V1 = P2V2 P
at constant temperature and quantity of gas

24. Gas Laws
As liquid nitrogen (-196oC) is poured over a balloon, the volume of the balloon decreases.

25. At constant pressure and quantity of gas Remember: T must be in Kelvin
Gas Laws
Charles’ Law:
The volume of a fixed amount of gas maintained at constant pressure is directly proportional to its absolute temperature.
V = k x T or V = k or V1 = V2
T T1 T2
At constant pressure and quantity of gas
Remember: T must be in Kelvin

26. Gas Laws
On a molecular level, as the temperature of a gas maintained at constant pressure decreases,
KE decreases
fewer collisions between gas molecules and the environment (i.e. container)
volume decreases in order to maintain constant pressure

27. Gas Laws
What happens when you “blow up” a balloon?

28. Gas Laws
What happens when you “blow up” a balloon?

29. Gas Laws
What happens when you “blow up” a balloon?

30. Gas Laws
What happens when you “blow up” a balloon?

31. Gas Laws What happens when you “blow up” a balloon?
the number of moles of gas (n) increases
and
the volume of the gas (balloon) increases

32. At constant temperature and pressure
Gas Laws
Avogadro’s Law:
The volume of a gas maintained at constant temperature and constant pressure is directly proportional to the number of moles of the gas.
Mathematically,
V = constant x n
At constant temperature and pressure

33. Gas Laws
At any given temperature and pressure, as the amount of gas increases,
the number of gas molecules increases
the number of collisions between gas molecules and the environment (container) increases
the volume must increase in order to maintain constant pressure

34. Gas Laws
In a chemical reaction, we use the coefficients to tell us how many moles or molecules are used or produced in a chemical reaction.
N2 (g) + 3 H2 (g)  2 NH3 (g)
1 mole of nitrogen reacts with 3 moles of hydrogen to produce 2 moles of ammonia

35. Gas Laws
Since the volume of a gas is directly proportional to the number of moles of gas at constant temperature and pressure, we can also use the coefficients to represent the volume of a gas involved in a reaction. (Avogadro’s Hypothesis)
N2 (g) + 3 H2 (g)  2 NH3 (g)
1 liter of nitrogen reacts with 3 liters of hydrogen to produce 2 liters of ammonia

36. Gas Laws
Boyle’s Law, Charles’ Law, and Avogadro’s Law can be combined to make a more general gas law:
Ideal Gas Law:
PV = nRT
where P = pressure
V = volume
n = moles
T = temperature (K)
R = gas constant

37. Gas Laws
The value of the gas constant (R) depends on the units of P, V, n, and T.
T must always be in Kelvin
n is usually in moles
If P (atm) and V (L),
then R = atm.L
mol.K
If P (torr) and V (L),
then R = L.torr
I will give you these on the test.