1. Gases

2. Pressure and Temperature

3. Units of Pressure
Millimeters of Mercury (mm Hg) is the most common because mercury barometers are most often used. Average atmospheric pressure at sea level at 0°C is 760 mm Hg.
Torr is another name for pressure when a mercury barometer is used in honor of Torricelli for his invention of the barometer. 1 torr = 1 mm Hg

4. One atmosphere of pressure (1 atm) is defined as being exactly equivalent to 760 mm Hg.
One pascal (Pa) is defined as the pressure exerted by the force of one Newton (1 N) acting on an area of one square meter. We usually always express in kilopascals (kPa).
1 atm = kPa = 760 mm Hg = 760 torr

5. 738 mmHg
98.4 kPa

6. Convert the following pressures:
925 mmHg into atm
1.22 atm
b torr into kPa
106 kPa
c kPa into atm
2.00 atm
840 mmHg into torr
840 torr

7. Standard Temperature and Pressure
Because volumes of gases change so much when the temperature or pressure changes, scientists have agreed on standard conditions of exactly 1 atm pressure and 0˚C.
These are called standard temperature and pressure or STP.

8. When doing gas law calculations, temperature must be in Kelvin, not Celsius.
K = °C + 273
To convert from Celsius to Kelvin, add 273
To convert from Kelvin to Celsius, subtract 273

9. Convert the following temperatures.
48°C to K
321 K
25°C to K
298 K
275 K to °C
2°C
323 K to °C
50°C

10. Dalton’s Law of Partial Pressures
Partial Pressure is the pressure of each gas in a mixture of a gas.
Dalton’s Law of Partial Pressure states that the total pressure of a gas mixture is the sum of the partial pressures of the component gases.
PT = P1 + P2 + P3 + …
PT is the total pressure of the mixture.
P1, P2, P3, and so on, are the partial pressures of the component gases.

11. Example:
A container holds 3 gases: O2, CO2, and He. The partial pressure of the 3 gases are 2.00 atm, 3.00 atm, and 4.00 atm, respectively. What is the total pressure inside the container?
PT = 2.00 atm atm atm
PT = 9.00 atm

12. Example:
A container with 2 gases, He and Ar, has a total pressure of 4.00 atm. If the partial pressure of He is 2.30 atm, what is the partial pressure of Ar?
4.00 atm = 2.30 atm + PAr
PAr = = 1.70 atm

13. Section 2: The Gas Laws

14. The Combined Gas Law

15. Example: P1 V1 P2 V2 T1 T2

16. Continue Examples on the board or use the document camera.

17. Section 3: Gas Volumes and Ideal Gas Law

18. Standard Molar Volume of A Gas: the volume occupied by one mole of a gas at STP is 22.4 L.
22.4 L or 1 mol
1 mol L
Using this conversion, you can get from grams  moles  Liters or vice versa.
Remember: this only works at STP

19. Example
20.12 L
mol

20. How many grams of O2 are contained in 18.5 L at STP?
18.5 𝐿 𝑚𝑜𝑙 22.4 𝐿 𝑔 1 𝑚𝑜𝑙 = 26.4 g O2

21. Ideal Gas Law

22. Example R P V n T

23. Multiply moles by molar mass of Ne to get grams
How many grams of Ne are in a 2.50 L container at 273K and 2.00 atm?
2.00 atm • 2.50 L = n • • 273 K
n = mol
Multiply moles by molar mass of Ne to get grams
4.50 g Ne P V n R T

24. Section 4: Diffusion and Effusion

25. Graham’s Law of Effusion

26. Examples
Molar Mass of O2
Molar Mass of H2

27. Examples
Molar Mass of Ar
Molar Mass of He

28. Square both sides to get rid of square root.
A sample of H2 effuses 9 times faster than an unknown gas. Estimate the molar mass of the unknown gas.
𝑅𝑎𝑡𝑒 𝑜𝑓 𝐻 2 𝑅𝑎𝑡𝑒 𝑜𝑓 𝑈𝑛𝑘𝑛𝑜𝑤𝑛 = 𝑋 =9
Square both sides to get rid of square root.
𝑋 2.02 =81
X =