Gases

Pressure and Temperature

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

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 = 101.325 kPa = 760 mm Hg = 760 torr

  738 mmHg 98.4 kPa

Convert the following pressures: 925 mmHg into atm 1.22 atm b. 792 torr into kPa 106 kPa c. 203 kPa into atm 2.00 atm 840 mmHg into torr 840 torr

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.

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

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

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.

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 + 3.00 atm + 4.00 atm PT = 9.00 atm

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 = 4.00-2.30 = 1.70 atm

Section 2: The Gas Laws

The Combined Gas Law  

Example:   P1 V1 P2 V2 T1 T2

Continue Examples on the board or use the document camera.

Section 3: Gas Volumes and Ideal Gas Law

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 22.4 L Using this conversion, you can get from grams  moles  Liters or vice versa. Remember: this only works at STP

Example   20.12 L 0.0987 mol

How many grams of O2 are contained in 18.5 L at STP? 18.5 𝐿 1 1 𝑚𝑜𝑙 22.4 𝐿 32.00 𝑔 1 𝑚𝑜𝑙 = 26.4 g O2

Ideal Gas Law  

Example   R P V n T

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 • 0.0821 • 273 K n = 0.223 mol Multiply moles by molar mass of Ne to get grams 4.50 g Ne P V n R T

Section 4: Diffusion and Effusion

Graham’s Law of Effusion  

Examples   Molar Mass of O2 Molar Mass of H2

Examples   Molar Mass of Ar Molar Mass of He

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 2.02 =9 Square both sides to get rid of square root. 𝑋 2.02 =81 X = 163.62