Chapter 13: GASES Chapter 13: GASES
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(2) The Nature of Gases • Air was the first gas to be studied about 300 years ago. • Representative particles of a gas may consist of one, two, or more atoms.. He. CH4. O2. He. He. O2. O2. CH4.
(3) Behavior/Physical Properties of Gases •Gases have mass: the mass is more spread out than in solids and liquids, so they have very low density.. • Easy to compress: if you squeeze a gas, you can reduce the volume. • Gases fill their container completely: volume of container = volume of gas.
(4) More Properties • Diffusion: gases can move through each other, and like to spread out as much as possible. examples:. • Pressure: gas particles push on other objects. That’s pressure! • Pressure vs. Temperature.
(5) Pressure/Temperature Relationship • The higher the temperature, the higher the pressure. • The lower the temperature, the lower the pressure..
(6) Kinetic Molecular Theory • Describes the behavior of the molecules in a gas. • The word kinetic comes from the Greek word kinetos, which means movable..
(7) Assumptions 1. Gases consist of small particles that have mass. 2. There is lots of space between the particles, 3. 4. 5. 6.. while the particles take up almost no space. The particles in a gas are in constant, rapid, random motion. All collisions are elastic. This means that they don’t ever slow down. They don’t lose any energy. The kinetic energy of the particles depends on their temperature. The gas particles are not attracted to each other, but they also don’t repel each other..
(8) Section 13-2 Measuring Gases • There are four variables that you have to consider when dealing with gases:. Amount of gas Volume Temperature Pressure.
(9) 1. Amount of Gas n (# of moles). =. mass (g) molar mass (g/mole). Example: How many moles of H2 gas do you have if you have 60.0 g of H2 gas? n (# of moles). = 60.0 g H2 = 30.0 moles 2.0 g/mole.
(10) 2. Volume (V) • A gas will completely fill up any container in which you put it. • Ex. If you put O2 gas into a 10 Liter tank, you have 10 Liters of O2 gas. • Ex. If you put H2 gas into a 25 Liter tank, you have 25 Liters of H2 gas..
(11) 3. Temperature (T). •For gas problems, the temperature always has to be in Kelvins. K = °C + 273.
(12) 4. Pressure (P) •Gas molecules exert pressure on anything they bump into, which basically means they exert pressure on everything..
(13) Atmospheric Pressure •definition: the pressure exerted by the air in the atmosphere •Gravity is the force of attraction between any two pieces of matter, or between any two objects that have mass. • The air in the atmosphere has mass, and so does the earth, so they are attracted to each other..
(14) • Pressure is measured in force per unit area, N/m2, which also equals 1 pascal. • There are several different units that we can use to measure pressure: 1 atm 101, 325 Pa 760 mm Hg 760 torr 14.70 lb/in2 101.325 kPa.
(15) • Because there are different units for. measuring air pressure, you can use the numbers from the list as conversion factors. • Example: If the pressure in a tire is 115 kPa, what is the pressure in lb/in2? 115 kPa x 14.70 lb/in2 = 16.68 lb/in2 101.325 kPa These numbers both come from the list on the previous slide..
(16) • A barometer is a device used to measure the atmospheric pressure. Sample Problem #1: The column of mercury in a barometer is 745 mm above the mercury in the dish. What is the atmospheric pressure in pascals? 745 mmHg × 101,325 Pa 760 mmHg = 99,300 Pa.
(17) Enclosed Gases • A manometer is an instrument that can be used to measure the pressure of a gas in a closed container..
(18) Sample Problem #2: You have a closed manometer like one shown in the diagram. The mercury in the open end of the tube is 27 mm higher than in the end attached to the gas. The atmospheric pressure is measured to be 755 mm Hg. What is the pressure of the gas in the container in atm?.
(19) Sample Problem #2, cont.. 1. Which is pushing harder on the mercury, the gas or the atmosphere? THE GAS 2. What is the difference in height? 27 mm Hg 3. If the gas is 27 mm Hg stronger than the atmosphere (755 mm Hg), what is the gas pressure? 755 mmHg + 27 mmHg = 782 mmHg 4. Convert to atm. 782 mmHg ×. 1 atm 760 mmHg. = 1.03 atm.
(20) Standard Temperature & Pressure • Temperature and pressure strongly influence the behavior of gases. • In order to compare different gases, they have to be at the same temperature and the same pressure. • Science came up with a standard temperature (0°C) and a standard pressure (1 atm) for dealing with gases..
(21) Chapter 13-3 The Gas Laws Pressure and Volume (Boyle’s Law) Temperature and Volume (Charles’ Law) Temperature and Pressure (Gay-Lussac’s Law).
(22) Boyle's Law The pressure of a gas is inversely related to the volume when T does not change Then the PV product remains constant. P1V1. =. P2V2. P1V1=. 8.0 atm x 2.0 L. = 16 atm L. P2V2=. 4.0 atm x 4.0 L. = 16 atm L.
(23) Boyle’s Law.
(24) Boyle’s Law Illustrated.
(25) Example: PV Problem Freon-12, CCl2F2, is used in refrigeration systems. What is the new volume (L) of a 1.6 L sample of Freon gas initially at 50 mm Hg after its pressure is changed to 200 mm Hg at constant T?.
(26) Find New Volume (V2) Solve for V2: V2 =. P1V1 = P2V2. V1 x P1 /P2. V2 = 1.6 L x 50 mm Hg = 200 mm Hg 0.4 L.
(27) Learning Check GL1 A sample of nitrogen gas is 6.4 L at a pressure of 0.70 atm. What will the new volume be if the pressure is changed to 1.40 atm? (T constant) Explain.. 1) 3.2 L 2) 6.4 L 3) 12.8 L.
(28) Solution GL1 A sample of nitrogen gas is 6.4 L at a pressure of 0.70 atm. What will the new volume be if the pressure is changed to 1.40 atm? (T constant) 6.4 L x 0.70 atm. =. 3.2 L (1). 1.40 atm Volume must decrease to cause an increase in the pressure.
(29) Charles’ Law: V and T At constant pressure, the volume of a gas is directly related to its absolute (K) temperature V1 = V2 T1 T2. Jacques Charles (17461823). Isolated boron and studied gases. Balloonist..
(30) Charles’s Law.
(31) Charles’ Law. V = 125 mL T = 273 K. V T. = 250 mL = 546 K. Observe the V and T of the balloons. How does volume change with temperature?.
(32) Learning Check GL3 Use Charles’ Law to complete the statements below: 1. If final T is higher than initial T, final V is (greater, or less) than the initial V. 2. If final V is less than initial V, final T is (higher, or lower) than the initial T..
(33) Solution GL3 V1 T1. = V2 T2. 1. If final T is higher than initial T, final V is (greater) than the initial V. 2. If final V is less than initial V, final T is (lower) than the initial T..
(34) Learning Check GL4 Example: VT Problem A sample of oxygen gas has a volume of 420 mL at a temperature of 18°C. What temperature (in °C) is needed to change the volume to 640 mL?.
(35) Solution GL4 T2 = 291 K x 640 mL = 420 mL 443 K = 443 K - 273 K. = 170°C.
(36) Avogadro’s Law When a gas is at constant T and P, the V is directly proportional to the number of moles (n) of gas (equal volumes of gases at the same T and P contain an = # of particles) V1 = V2 n1 n2. initial. • Or V = k x n. final n = # moles k = constant V = volume.
(37) Volume and Number of Moles.
(38) Daltons’ Law of Partial Pressures Partial Pressure Pressure each gas in a mixture would exert if it were the only gas in the container Dalton's Law of Partial Pressures The total pressure exerted by a gas mixture is the sum of the partial pressures of the gases in that mixture. PT = P1 +. P2 + P3 + ......
(39) Partial Pressures The total pressure of a gas mixture depends on the total number of gas particles, not on the types of particles. 1 mole H2 P = 1.00 atm. 0.5 mole O2 + 0.3 mole He + 0.2 mole Ar P = 1.00 atm.
(40) Gases in the Air The % of gases in air. Partial pressure (STP). 78.08% N2 593.4 mmHg 20.95% O2 159.2 mmHg 0.94% Ar 0.03% CO2. 7.1 mmHg 0.2 mmHg. PAIR = PN + PO + PAr + PCO = 760 mmHg 2. •. 2. Total Pressure 760. 2. mm Hg.
(41) Learning Check C6 Example: Partial Pressure A.If the atmospheric pressure today is 745. mm Hg, what is the partial pressure (mm Hg) of O2 in the air? 745 mm Hg x 20.95% O2 745 mm Hg x .2095 = 156 mm Hg.
(42) 13-4: IDEAL GAS LAW. P V = n R T Brings together gas properties. Can be derived from experiment and theory. BE SURE YOU KNOW THIS EQUATION!.
(43) Ideal Gas Equation Volume Pressure. PV=nRT Number of moles. Universal Gas Constant can be R = 0.0821 atm L / mol K or R = 8.314 kPa L / mol K. Universal Gas Constant. Temperature.
(44) Example: PV = nRT How much N2 is required to fill a small room with a volume of 960 cubic feet (27,000 L) to 745 mm Hg at 25 oC? Solution 1. Get all data into proper units V = 27,000 L T = 25 oC + 273 = 298 K P = 745 mm Hg (1 atm/760 mm Hg) = 0.98 atm And we always know R, 0.0821 L atm / mol K.
(45) 2. Now plug in those values and solve for the unknown. PV = nRT RT RT. n=. (0.98 atm)(2.7 x 104 L) (0.0821 L x atm/K x mol)(298 K) n = 1.1 x 103 mol (or about 30 kg of gas).
(46) Deviations from Ideal Gas Law • Real molecules have volume. The ideal gas consumes the entire amount of available volume. It does not account for the volume of the molecules themselves. • There are intermolecular forces. An ideal gas assumes there are no attractions between molecules. Attractions slow down the molecules and reduce the amount of collisions. – Otherwise a gas could not condense to become a liquid..
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