Chapter 9 Gases Revised PowerPoint PDF

COLLEGE CHEMISTRY 2E Chapter #9 Gases PowerPoint Image Slideshow Chapter 9 GASES PowerPoint Image Slideshow  The Gas Laws The density of a gas decreases as its temperature increases.  History of Science...

COLLEGE CHEMISTRY 2E Chapter #9 Gases PowerPoint Image Slideshow Chapter 9 GASES PowerPoint Image Slideshow  The Gas Laws The density of a gas decreases as its temperature increases.  History of Science Gas Laws Air has been known to exist since ancient times. The Greeks considered air one of the four fundamental elements of nature along with earth, water, and fire. 18th century—Lavoisier, Cavendish, and Priestley: Determined that air is actually a mixture of gases. Gay-Lussac’s law Dalton announces his atomic theory Avagadro’s particle Number theory Boyle’s law Charles’s law 1650 1700 1750 1800 1850 Mogul empire in India Constitution of the United States signed U.S. Congress bans (1526-1707) importation of slaves United States Bill of Rights ratified Napoleon is emperor(1804- 12) Latin American countries gain independence (1791- 1824) Haiti declares independence Herron, Frank, Sarquis, Sarquis, Schrader, Kulka, Chemistry, Heath Publishing,1996, page 220  Scientists Evangelista Torricelli (1608-1647) – Published first scientific explanation of a vacuum. – Invented mercury barometer. Robert Boyle (1627- 1691) – Volume inversely related to pressure (temperature remains constant) Jacques Charles (1746 -1823) – Volume directly related to temperature (pressure remains constant) Joseph Gay-Lussac (1778-1850) – Pressure directly related to temperature (volume remains constant)  Gas Pressure Pressure is defined as the force exerted on a given area. Gas pressure is a result of the force exerted by gas molecules colliding with the surface of objects. English system units for pressure: pounds per square inch (psi) A mercury barometer is a device used to measure the atmospheric pressure. Gas: expands uniformly to fill the container in which it is placed.  Figure 9.2 The atmosphere above us exerts a large pressure on objects at the surface of the earth, roughly equal to the weight of a bowling ball pressing on an area the size of a human thumbnail.  Units of Pressure Due to the way gas pressure is measured, millimeters of mercury (mm Hg) is a commonly used unit of pressure. Probably the most common unit of gas pressure is the atmosphere (atm). – 1 atm is the “normal atmospheric pressure” – 1 atm = 760 mm Hg SI unit for pressure is the Pascal (Pa). – 1 Pa is defined as the pressure exerted by a 0.1 mm high film of water on the surface beneath it. – Kilopascal (kPa) A related unit of pressure is the bar. – 1 bar = 105 Pa Pressure conversions: – 1.013 bar = 1 atm = 760 mm Hg = 760 torr – = 14.7 psi = 101.3 kPa  Temperature Always use absolute temperature (Kelvin) when working with gases. ºF -459 32 212 ºC -273 0 100 K 0 273 373 K = ºC + 273 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  STP TABLE A STP Standard Temperature & Pressure 0°C 273 K - OR - 1 atm 101.325 kPa 760 mm Hg 760 Torr Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Pressure KEY UNITS AT SEA LEVEL 101.325 kPa (kilopascal) Sea level 1 atm 760 mm Hg 760 torr N 14.7 psi kPa  2 m Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Temperature/ Pressure Conversions Temperature Conversions K = °C + 273 Convert the following to Kelvin Convert the following to Celsius 1) 10o C ________ 1) 32o F ________ 2) 30o C ________ 2) 90o F ________ 3) 0o C ________ 3) 212o F ________ Convert the following to Kelvin Convert the following to Celsius 1) 0o C ________ 1) 112o K ________ 2) -50o C ________ 2) 200o K ________ 3) 90o C ________ 3) 273o K ________ 4) -20o C ________ 4) 350o K ________  Pressure Conversions 1 atm =101.3 kPa = 101,325 Pa = 760 mm Hg = 760 torr = 14.7 lb/in2 (psi) 1.The air pressure for a certain tire is 109 kPa. What is this pressure in atmospheres? 2.The air pressure inside a submarine is 0.62 atm. What would be the height of a column of mercury balanced by this pressure? 3. The weather news gives the atmospheric pressure as 1.07 atm. What is this atmospheric pressure in torr? 4. An experiment at Sandia National Labs in New Mexico is performed at an atmospheric pressure of 758.7 mm Hg. What is this pressure in atm? 5. A bag of potato chips is sealed in a factory near sea level. The atmospheric pressure at the factory is 761.3 mm Hg. The pressure inside the bag is the same. What is the pressure inside the bag of potato chips in Pa? 6. The same bag of potato chips from Problem 5 is shipped to a town in Colorado, where the atmospheric pressure is 99.82 kPa. What is the difference (in Pa) between the pressure in the bag and the atmospheric pressure of the town? 7. The pressure gauge on a compressed air tank reads 43.2 lb/in2. What is the pressure in atm? 8. The pressure in the tire of an automobile in 34.8 lb/in2. What is the pressure in kPa?  Answers 1. The air pressure for a certain tire is 109 kPa. What is this pressure in atmospheres? 1.08 atm 2.The air pressure inside a submarine is 0.62 atm. What would be the height of a column of mercury balanced by this pressure? 471 mm Hg 3. The weather news gives the atmospheric pressure as 1.07 atm. What is this atmospheric pressure in torr? 813 torr 4. An experiment at Sandia National Labs in New Mexico is performed at an atmospheric pressure of 758.7 mm Hg. What is this pressure in atm? 0.9983 atm 5. A bag of potato chips is sealed in a factory near sea level. The atmospheric pressure at the factory is 761.3 mm Hg. The pressure inside the bag is the same. What is the pressure inside the bag of potato chips in Pa? 101,498 Pa 6. The same bag of potato chips from Problem 5 is shipped to a town in Colorado, where the atmospheric pressure is 99.82 kPa. What is the difference (in Pa) between the pressure in the bag and the atmospheric pressure of the town? 1,678 Pa 7. The pressure gauge on a compressed air tank reads 43.2 lb/in 2. What is the pressure in atm? 2.94 atm 8. The pressure in the tire of an automobile in 34.8 lb/in2. What is the pressure in kPa? 240. kPa Kinetic Molecular Theory (KMT)  explains why gases behave as they do  deals w/“ideal” gas particles… 1. …are so small that they are assumed to have zero volume 2.…are in constant, straight-line motion 3.…experience elastic collisions in which no energy is lost 4.…have no attractive or repulsive forces toward each other 5.…have an average kinetic energy (KE) that is proportional to the absolute temp. of gas (i.e., Kelvin temp.) AS TEMP. , KE  Kinetic Molecular Theory Particles in an ideal gas… – have no volume. – have elastic collisions. – are in constant, random, straight-line motion. – don’t attract or repel each other. – have an avg. KE directly related to Kelvin temperature. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Real Gases Particles in a REAL gas… – have their own volume – Have slight attractive forces for each other ***Gas behavior is most ideal… – at low pressures – at high temperatures – in nonpolar atoms/molecules Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Characteristics of Gases Gases expand to fill any container. – random motion, no attraction Gases are fluids (like liquids). – no attraction Gases have very low densities. – no volume = lots of empty space Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Characteristics of Gases Gases can be compressed. – no volume = lots of empty space Gases undergo diffusion & effusion. – random motion Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem  Properties of Gases Gas properties can be modeled using math. Model depends on: V = volume of the gas (liters, L) T = temperature (Kelvin, K) P = pressure (atmospheres, atm) n = amount (moles, mol) Pressure - Temperature - Volume Relationship P V P TT V PP VV Boyle’s P1V1 = P2V2 (Indirect) Charles V1/T1 = V2/T2 (Direct) Gay-Lussac’s P1/T1 = P2/T2 (Direct)  Two formulas Combined Gas Law – P1V1 = P2V2 – T1 T2 Ideal Gas Law PV = nRT P = pressure R = gas law constant V = Volume T = Temperature n = # of moles  Practice Ideal Gas Law If I have 4 moles of a gas at a pressure of 5.6 atm and a volume of 12 liters what is the temperature? Sulfur hexafluoride is a gas used as a long-term tamponade (plug) for a retinal hole to repair a detached retina in the eye. If 2.50 g of this compound is introduced into an evacuated 500.0 mL container at 83℃ what pressure in atmospheres is developed. Iodine I2 is a solid at room temperature but converts to a gas when warmed. What is the temperature in a 73.3 mL bulb that contains 0.292 g of I2 vapor at a pressure of.462 atm?  A FEW THINGS ABOUT GAS LAWS: ANY GAS OCCUPIES THE SAME VOLUME. 1 MOLE OF ANY GAS IS EQUAL TO 22.4 LITERS. 1 MOLE OF PARTICLES OF ANY GAS IS 6.022 X 1023. IDEAL GAS LAW STATES: GASES ARE MOST IDEAL AT LOW PRESSURE AND HIGH TEMPERATURE. THEY ARE MOST REAL ……  Chapter Outline 9.1 Gas Pressure 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions 9.4 Effusion and Diffusion of Gases 9.5 The Kinetic-Molecular Theory 9.6 Non-Ideal Gas Behavior  Learning Objectives Gas Pressure – Define the property of pressure – Define and convert among the units of pressure measurements – Describe the operation of common tools for measuring gas pressure – Calculate pressure from manometer data Conversions  Measuring Atmospheric Pressure The barometer measures pressure in terms of the height of a column of liquid mercury. The atmosphere exerts a force on the liquid mercury causing it to rise. Commonly, the pressure of the atmosphere at sea level pushes the column of mercury ~760 mm high.  Atmospheric Pressure and Altitude Pressure varies with the height of the column of air above. We don’t notice atmospheric pressure but do notice changes in atmospheric pressure Measuring the Pressure of a Confined Gas—The Manometer The force of the gas in the flask is measured relative to the force of the atmosphere (atmospheric pressure). The distance and direction that the Hg travels is related to the pressure of the gas.  Volume of a Gas At constant temperature and pressure, a gas expands uniformly to fill the container in which it is contained.  Relating Pressure, Volume, Amount, Temperature: The Ideal Gas Law There are many different types of gases. – Elements: Ar, He, H2, N2, O2 – Compounds: CO2, CO, H2O, NH3 All gases have the same dependence on those four properties. – Volume – Amount – Temperature – Pressure  Pressure and Temperature The pressure of a gas is directly proportional to its temperature in K. At constant V and n This is Amonton’s Law or Gay-Lussac’s Law.  Volume and Temperature The volume of a gas is directly proportional to its temperature in K. At constant P and n This is Charles’s Law.  Figure 9.12 The volume and temperature are linearly related for 1 mole of methane gas at a constant pressure of 1 atm. If the temperature is in kelvin, volume and temperature are directly proportional. The line stops at 111 K because methane liquefies at this temperature; when extrapolated, it intersects the graph’s origin, representing a temperature of absolute zero.  Figure 9.13 When a gas occupies a smaller volume, it exerts a higher pressure; when it occupies a larger volume, it exerts a lower pressure (assuming the amount of gas and the temperature do not change). Since P and V are inversely proportional, a graph of vs. V is linear.  Volume and Pressure The volume of a gas is inversely proportional to its pressure. At constant T and n This is Boyle’s Law.  Figure 9.14 The relationship between pressure and volume is inversely proportional. (a) The graph of P vs. V is a parabola, whereas (b) the graph of vs. V is linear.  Figure 9.15 Breathing occurs because expanding and contracting lung volume creates small pressure differences between your lungs and your surroundings, causing air to be drawn into and forced out of your lungs.  Volume and Amount (Moles) The volume of a gas is directly proportional to its amount (moles of gas). At constant T and P This is Avogadro’s Law.  The Ideal Gas Law These four equations can be combined into a single law consisting of all four properties. The Ideal Gas Law: PV = nRT The four constants (k) were combined into a new constant, R, called the ideal gas constant.  The Ideal Gas Law PV = nRT P is the pressure in atm. V is the volume in L. n is the moles of gas. T is the temperature in K. L × atm R is the ideal gas constant. R = 0.08206 mol × K  The Ideal Gas Law PV = nRT The ideal gas law accurately describes the properties of an ideal gas. What is an ideal gas? What assumptions are being made?  Figure 9.16 Scuba divers use compressed air to breathe while underwater. (credit: modification of work by Mark Goodchild)  Figure 9.17 Scuba divers, whether at the Great Barrier Reef or in the Caribbean, must be aware of buoyancy, pressure equalization, and the amount of time they spend underwater, to avoid the risks associated with pressurized gases in the body. (credit: Kyle Taylor)  Standard Temperature and Pressure Sometimes it is important to define standard conditions. Standard temperature and pressure (STP) – 0 oC = 273 K – 1 atm Note: The IUPAC recommended value of standard pressure was changed to 1 bar in 1982, but many texts continue to use the 1 atm value. At STP, the volume of one mole of any gas can be calculated.  Modified Ideal Gas Law PV = nRT Modified version of the ideal gas law: – Moles kept constant – But other properties change  Figure 9.18 Regardless of its chemical identity, one mole of gas behaving ideally occupies a volume of ~22.4 L at STP.  Learning Objectives 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions – Use the ideal gas law to compute gas densities and molar masses – Perform stoichiometric calculations involving gaseous substances – State Dalton’s law of partial pressures and use it in calculations involving gaseous mixtures  Calculating the Density of a Gas Recall, the density of a substance mass d V We can rearrange the following equation grams PV = RT M Substitute into the density equation to obtain P d =M RT  Density of Gases P d =M RT Recall that density is an intensive property. – Does not depend on the amount of substance But, the density of a gas does depend on – Pressure  Calculating the Molar Mass of a Gas M Recall, the molar mass ( ) of a substance. M = grams = m mole n m n= M This expression can be rearranged. m PV = RT M And then substituted into the Ideal Gas  Figure 9.19 When the volatile liquid in the flask is heated past its boiling point, it becomes gas and drives air out of the flask. At tl ⟶ g, the flask is filled with volatile liquid gas at the same pressure as the atmosphere. If the flask is then cooled to room temperature, the gas condenses and the mass of the gas that filled the flask, and is now liquid, can be measured. (credit: modification of work by Mark Ott)  Dalton’s Law The ideal gas law also applies to mixtures of gases if the gases don’t react with each other. Dalton’s law of partial pressure: The total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the component gases. PT = PA + PB + PC …  Dalton’s Law of Partial Pressure PT = PA + PB + PC … For a mixture of gases A, B, C, … PT is the total pressure of the gas mixture PA is the partial pressure of gas A PB is the partial pressure of gas B PC is the partial pressure of gas C The pressure exerted by each individual gas in a mixture is called its partial pressure. – The partial pressure is equal to the pressure that the gas would exert if it were by itself.  Figure 9.20 If equal-volume cylinders containing gasses at pressures of 300 kPa, 450 kPa, and 600 kPa are all combined in the same-size cylinder, the total pressure of the gas mixture is 1350 kPa.  Partial Pressure and Mole Fraction We can express the partial pressure of a gas in a mixture in terms of its mole fraction. Consider a mixture of gases A and B. The mole fraction of gas A (ΧA) is the number of moles of A divided by the total number of moles of gas in the mixture.  Collecting a Gas Over Water A common way to collect and quantify a gas produced in a chemical reaction is to capture the gas in an inverted bottle that had been filled with water. The volume of water displaced can then be related back to the amount of gas produced. As the gas travels through the water, it picks up water vapor, H2O (g).  Figure 9.21 When a reaction produces a gas that is collected above water, the trapped gas is a mixture of the gas produced by the reaction and water vapor. If the collection flask is appropriately positioned to equalize the water levels both within and outside the flask, the pressure of the trapped gas mixture will equal the atmospheric pressure outside the flask (see the earlier discussion of manometers).  Wet Gases We refer to the gas collected over water as a wet gas mixture. Say, for example, that the gas being collected was H2. PT is the atmospheric pressure. PH is the partial pressure of the H2 (g) collected. 2 PH is the vapor pressure of water. 2O  Vapor Pressure Vapor pressure of water: The pressure exerted by water vapor in equilibrium with liquid water in a closed container. – Intensive property—does not depend on the amount of water – Temperature dependent  Figure 9.22 This graph shows the vapor pressure of water at sea level as a function of temperature.  Stoichiometry in Gaseous Reactions Gases may appear as reactants or products in a chemical reaction. The ideal gas law and a balanced chemical equation can be used to relate the P, V, T, n, and grams of a gas that takes part in a chemical reaction. Stoichiometric factors can be used to relate the moles of one substance to the moles of any other substance in a  Stoichiometry with Gas Volumes The volume ratio of any two gases in a reaction at constant temperature and pressure is the same as the stoichiometric ratio (mole ratio) in the balanced chemical equation. Recall that V = kn (constant T, P). The volume of a gas is directly proportional to the moles of gas.  Figure 9.23 One volume of N2 combines with three volumes of H2 to form two volumes of NH3.  Figure 9.24 Greenhouse gases trap enough of the sun’s energy to make the planet habitable— this is known as the greenhouse effect. Human activities are increasing greenhouse gas levels, warming the planet and causing more extreme weather events.  Figure 9.25 CO2 levels over the past 700,000 years were typically from 200–300 ppm, with a steep, unprecedented increase over the past 50 years.  Figure 9.26 Susan Solomon’s research focuses on climate change and has been instrumental in determining the cause of the ozone hole over Antarctica. (credit: National Oceanic and Atmospheric Administration)  Learning Objectives 9.4 Effusion and Diffusion of Gases – Define and explain effusion and diffusion – State Graham’s law and use it to compute relevant gas properties  Figure 9.27 (a) Two gases, H2 and O2, are initially separated. (b) When the stopcock is opened, they mix together. The lighter gas, H2, passes through the opening faster than O 2, so just after the stopcock is opened, more H2 molecules move to the O2 side than O2 molecules move to the H2 side. (c) After a short time, both the slower-moving O 2 molecules and the faster-moving H2 molecules have distributed themselves evenly on both sides of the vessel.  Learning Objectives 9.6 Non-Ideal Gas Behavior – Describe the physical factors that lead to deviations from ideal gas behavior – Explain how these factors are represented in the van der Waals equation – Define compressibility (Z) and describe how its variation with pressure reflects non-ideal behavior – Quantify non-ideal behavior by comparing computations of gas properties using the ideal gas law and the van der Waals equation