Chapter 11: Gases

Section 1: Gases and Pressure

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. Can also be expressed in kilopascals (kPa). 1 atm = x 10 5 Pa = kPa 1 atm = 760 mm Hg = 760 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.

738 mmHg 98.4 kPa

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. P T = P 1 + P 2 + P 3 + … P T is the total pressure of the mixture. P 1, P 2, P 3, and so on, are the partial pressures of the component gases.

Example: A container holds 3 gases: O 2, CO 2, 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? P T = 2.00 atm atm atm P T = 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 + P Ar P Ar = = 1.70 atm

Gases Collected by Water Displacement Gases produced in the lab are often collected over water, through a method known as water displacement. Water molecules mix with the gas molecules and the resulting water vapor experts its own pressure, known as vapor pressure. The pressure of the water vapor must be taken into account when determining the pressure of the gas. P atm = P gas + P H20 The vapor pressure of water is dependent on temperature (Table A-8 on page 859).

Example: Oxygen gas from the decomposition of potassium chlorate, KClO 3, was collected by water displacement. The barometric pressure and the temperature during the experiment were torr and 20.0°C. What was the partial pressure of the oxygen collected? P atm = torr and P H2O = 17.5 torr at 20.0°C P atm = P O2 + P H2O  P O2 = P atm – P H2O P O2 = torr – 17.5 torr = torr

Example: Some Hydrogen gas is collected over water at 20.0°C. The partial pressure of hydrogen is torr. What was the barometric pressure at the time the gas was collected? P H2 = torr and P H2O = 17.5 torr at 20.0°C P atm = P H2 + P H2O P atm = torr torr = torr

Section 2: The Gas Laws

Boyle’s Law: Pressure- Volume Relationship P 1 x V 1 = P 2 x V 2 What must stay constant for Boyle’s Law to work? Temperature Pressure can be in any unit, but both must be the same. Volume can be in any unit, but both must be the same.

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Example: A balloon filled with helium gas has a volume of 500 mL at a pressure of 1 atm. If the pressure decreases to 0.5 atm and the temperature remained the same, what volume does the gas now occupy? Trying to find the new volume V 2 Plug in the 3 known values and solve. 1 atm x 500 mL = 0.5 atm x V 2 V 2 = 1000 mL

Charles’s Law: Volume- Temperature Relationship

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Gay-Lussac’s Law: Pressure- Temperature Relationship

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The Combined Gas Law

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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 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 mmHg mol

Ideal Gas Law

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Section 4: Diffusion and Effusion

Graham’s Law of Effusion

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Examples