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Chapter 5 Gases Kim Shih Ph.D.. Gases Pushing Gas molecules are constantly in motion As they move and strike a surface, they push on that surface push.

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Presentation on theme: "Chapter 5 Gases Kim Shih Ph.D.. Gases Pushing Gas molecules are constantly in motion As they move and strike a surface, they push on that surface push."— Presentation transcript:

1 Chapter 5 Gases Kim Shih Ph.D.

2 Gases Pushing Gas molecules are constantly in motion As they move and strike a surface, they push on that surface push = force If we could measure the total amount of force exerted by gas molecules hitting the entire surface at any one instant, we would know the pressure the gas is exerting pressure = force per unit area

3 The Effect of Gas Pressure Gas flows from an area of high pressure to an area of low pressure the bigger the difference in pressure, the stronger the flow of the gas If there is something in the gass path, the gas will try to push it along as the gas flows Differences in air pressure result in weather and wind patterns The higher in the atmosphere you climb, the lower the atmospheric pressure is around you

4 Pressure Imbalance in the Ear If there is a difference in pressure across the eardrum membrane, the membrane will be pushed out – what we commonly call a popped eardrum

5 The Pressure of a Gas Gas pressure is a result of the constant movement of the gas molecules and their collisions with the surfaces around them The pressure of a gas depends on several factors number of gas particles in a given volume volume of the container average speed of the gas particles

6 Gases and Gas Pressure

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8 Manometer for this sample, the gas has a larger pressure than the atmosphere, so

9 Manometer

10 Measuring Air Pressure gravity We measure air pressure with a barometer Column of mercury supported by air pressure Force of the air on the surface of the mercury counter balances the force of gravity on the column of mercury

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12 1.The height of the column increases because atmospheric pressure decreases with increasing altitude 2.The height of the column decreases because atmospheric pressure decreases with increasing altitude 3.The height of the column decreases because atmospheric pressure increases with increasing altitude 4.The height of the column increases because atmospheric pressure increases with increasing altitude 1.The height of the column increases because atmospheric pressure decreases with increasing altitude 2.The height of the column decreases because atmospheric pressure decreases with increasing altitude 3.The height of the column decreases because atmospheric pressure increases with increasing altitude 4.The height of the column increases because atmospheric pressure increases with increasing altitude Practice – What happens to the height of the column of mercury in a mercury barometer as you climb to the top of a mountain?

13 Common Units of Pressure

14 Brain Exercises

15 A high-performance bicycle tire has a pressure of 132 psi. What is the pressure in mmHg? Convert 45.5 psi into kPa

16 Manometers The pressure of a gas trapped in a container can be measured with an instrument called a manometer Manometers are U-shaped tubes, partially filled with a liquid, connected to the gas sample on one side and open to the air on the other A competition is established between the pressures of the atmosphere and the gas The difference in the liquid levels is a measure of the difference in pressure between the gas and the atmosphere

17 The Gas Laws Ideal Gas: A gas whose behavior follows the gas laws exactly. The physical properties of a gas can be defined by four variables: Ppressure Ttemperature Vvolume nnumber of moles

18 The Gas Laws ---- Boyles Law Boyles Law constant n and T PV = k Pressure of a gas is inversely proportional to its volume

19 Boyles Law P initial V initial = P final V final

20 Boyles Law: A Molecular View Pressure is caused by the molecules striking the sides of the container When you decrease the volume of the container with the same number of molecules in the container, more molecules will hit the wall at the same instant This results in increasing the pressure

21 Boyles Law and Diving Scuba tanks have a regulator so that the air from the tank is delivered at the same pressure as the water surrounding you. This allows you to take in air even when the outside pressure is large. Because water is more dense than air, for each 10 m you dive below the surface, the pressure on your lungs increases 1 atm at 20 m the total pressure is 3 atm If your tank contained air at 1 atm of pressure, you would not be able to inhale it into your lungs you can only generate enough force to overcome about 1.06 atm

22 If a diver holds her breath and rises to the surface quickly, the outside pressure drops to 1 atm According to Boyles law, what should happen to the volume of air in the lungs? Because the pressure is decreasing by a factor of 3, the volume will expand by a factor of 3, causing damage to internal organs. Always Exhale When Rising!! Boyles Law and Diving

23 The Gas Laws ---- Charles Law Charles Law V α T constant n and P = k T V Volume is directly proportional to temperature

24 = T final V final T initial V initial Charles Law

25 If the lines are extrapolated back to a volume of 0, they all show the same temperature, °C, called absolute zero If you plot volume vs. temperature for any gas at constant pressure, the points will all fall on a straight line

26 The pressure of gas inside and outside the balloon are the same At high temperatures, the gas molecules are moving faster, so they hit the sides of the balloon harder – causing the volume to become larger The pressure of gas inside and outside the balloon are the same At low temperatures, the gas molecules are not moving as fast, so they dont hit the sides of the balloon as hard – therefore the volume is small Charless Law – A Molecular View

27 The Gas Laws ---- Avogadros Law Avogadros Law constant T and P = k n V = n final V final n initial V initial V α n Volume directly proportional to the number of gas molecules

28 The Gas Laws =Avogadros Law: P initial V initial = P final V final Boyles Law: n final V final n initial V initial Charles Law: = T final V final T initial V initial Summary

29 The General Gas Law Avogadros Law: PV = nRT = kBoyles Law: Charles Law: = k P RT n V == k P nR T V = (n and P are constant) (P and T are constant) (n and T are constant) General Gas law: = n final T final P final V final n initial T initial P initial V initial

30 The Ideal Gas Law Standard Temperature and Pressure (STP) for Gases Ideal Gas Law: PV = nRT P = 1 atm T = 0 °C ( K) R is the gas constant and is the same for all gases. R = K mol L atm

31 The Ideal Gas Law What is the volume of 1 mol of gas at STP? = LV = P nRT = (1 atm) (1 mol) K mol L atm ( K)

32 Brain Exercises

33 A gas occupies 10.0 L at 44.1 psi and 57 °F. What volume will it occupy at standard conditions? Calculate the volume occupied by 637 g of SO 2 (MM 64.07) at 6.08 x 10 4 mmHg and –23 °C

34 Density of Gas PM=DRT PV=nRT PV=(Mass/M.W.)RT P x M.W. = (Mass/V) RT Density is directly proportional to molar mass

35 Density & Pressure Pressure is the result of the constant movement of the gas molecules and their collisions with the surfaces around them When more molecules are added, more molecules hit the container at any one instant, resulting in higher pressure also higher density

36 Calculate the density of a gas at 775 torr and 27 °C if moles weighs g Calculate the density of N 2 at 125°C and 755 mmHg

37 Molar Mass of a Gas One of the methods chemists use to determine the molar mass of an unknown substance is to heat a weighed sample until it becomes a gas, measure the temperature, pressure, and volume, and use the ideal gas law

38 Calculate the molar mass of a gas with mass g that has a volume of L at 55°C and 886 mmHg What is the molar mass of a gas if 12.0 g occupies 197 L at 380 torr and 127 °C?

39 Mixtures of Gases When gases are mixed together, their molecules behave independent of each other – all the gases in the mixture have the same volume all completely fill the container each gass volume = the volume of the container – all gases in the mixture are at the same temperature therefore they have the same average kinetic energy Therefore, in certain applications, the mixture can be thought of as one gas – even though air is a mixture, we can measure the pressure, volume, and temperature of air as if it were a pure substance – we can calculate the total moles of molecules in an air sample, knowing P, V, and T, even though they are different molecules

40 Partial Pressure and Daltons Law P total = P 1 + P 2 + … + P N Mole Fraction (X) = Daltons Law of Partial Pressures: The total pressure exerted by a mixture of gases in a container at constant V and T is equal to the sum of the pressures of each individual gas in the container. X i = P total PiPi X i = n total nini or Total moles in mixture Moles of component

41 Lake Nyos

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43 Brain Exercises

44 Find the partial pressure of neon in a mixture with total pressure 3.9 atm, volume 8.7 L, temperature 598 K, and 0.17 moles Xe Find the mole fractions and partial pressures in a 12.5 L tank with 24.2 g He and 4.32 g O 2 at 298 K

45 Collecting Gases Gases are often collected by having them displace water from a container The problem is that because water evaporates, there is also water vapor in the collected gas The partial pressure of the water vapor, called the vapor pressure, depends only on the temperature so you can use a table to find out the partial pressure of the water vapor in the gas you collect if you collect a gas sample with a total pressure of mmHg* at 25 °C, the partial pressure of the water vapor will be mmHg – so the partial pressure of the dry gas will be mmHg Table 5.4*

46 Collecting Gas by Water Displacement

47 Vapor Pressure of Water

48 1.02 L of O 2 collected over water at 293 K with a total pressure of mmHg. Find mass O moles of H 2 is collected over water in a 10.0 L container at 323 K. Find the total pressure.

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51 Stoichiometric Relationships with Gases 2Na(s) + 3N 2 (g)2NaN 3 (s) The reaction used in the deployment of automobile airbags is the high-temperature decomposition of sodium azide, NaN 3, to produce N 2 gas. How many liters of N 2 at 1.15 atm and 30.0 °C are produced by decomposition of 45.0 g NaN 3 ? P, V, T of Gas Amole Amole BP, V, T of Gas B

52 Kims Law If you dont know where to start, always start with mole number

53 Stoichiometric Relationships with Gases 2Na(s) + 3N 2 (g)2NaN 3 (s) 45.0 g NaN g NaN 3 1 mol NaN 3 2 mol NaN 3 3 mol N 2 xx Volume of N 2 produced: = 1.04 mol N 2 Moles of N 2 produced: = 22.5 LV = P nRT = (1.15 atm) (1.04 mol) K mol L atm (303.2 K)

54 Brain Exercises

55 How many liters of O STP can be made from the decomposition of g of PbO 2 ? 2 PbO 2 (s) 2 PbO(s) + O 2 (g) (PbO 2 = 239.2, O 2 = 32.00) What volume of H 2 is needed to make 35.7 g of CH 3 OH at 738 mmHg and 355 K? CO(g) + 2 H 2 (g) CH 3 OH(g) What volume of O 2 at atm and 313 K is generated by the thermolysis of 10.0 g of HgO? 2 HgO(s) 2 Hg(l) + O 2 (g) MW HgO = g/mol

56 The Kinetic-Molecular Theory of Gases 1.A gas consists of tiny particles, either atoms or molecules, moving about at random. 2.The volume of the particles themselves is negligible compared with the total volume of the gas; most of the volume of a gas is empty space. 3.The gas particles act independently of one another; there are no attractive or repulsive forces between particles. 4.Collisions of the gas particles, either with other particles or with the walls of a container, are elastic (constant temperature). 5.The average kinetic energy of the gas particles is proportional to the Kelvin temperature of the sample.

57 Kinetic Energy and Molecular Velocities Average kinetic energy of the gas molecules depends on the average mass and velocity – KE = ½mv 2 Gases in the same container have the same temperature, therefore they have the same average kinetic energy If they have different masses, the only way for them to have the same kinetic energy is to have different average velocities – lighter particles will have a faster average velocity than more massive particles

58 The Kinetic-Molecular Theory of Gases

59 Molecular Speed vs. Molar Mass To have the same average kinetic energy, heavier molecules must have a slower average speed average speed molar mass

60 Temperature and Molecular Velocities _ KE avg = ½N A mu 2 – N A is Avogadros number KE avg = 1.5RT – R is the gas constant in energy units, J/molK 1 J = 1 kgm 2 /s 2 Equating and solving we get – N A mass = molar mass in kg/mol As temperature increases, the average velocity increases

61 Molecular Velocities All the gas molecules in a sample can travel at different speeds However, the distribution of speeds follows a statistical pattern called a Boltzman distribution We talk about the average velocity of the molecules, but there are different ways to take this kind of average The method of choice for our average velocity is called the root-mean-square method, where the rms average velocity, u rms, is the square root of the average of the sum of the squares of all the molecule velocities

62 Boltzman Distribution

63 Calculate the velocity of O 2 at 25 °C T = = 298KMM of O 2 = 32g/mol Calculate the rms velocity of CH 4 (MM 16.04) at 25 °C

64 Mean Free Path Molecules in a gas travel in straight lines until they collide with another molecule or the container The average distance a molecule travels between collisions is called the mean free path Mean free path decreases as the pressure increases

65 Diffusion and Effusion The process of a collection of molecules spreading out from high concentration to low concentration is called diffusion The process by which a collection of molecules escapes through a small hole into a vacuum is called effusion The rates of diffusion and effusion of a gas are both related to its rms average velocity For gases at the same temperature, this means that the rate of gas movement is inversely proportional to the square root of its molar mass

66 Grahams Law: Diffusion and Effusion of Gases

67 Grahams Law of Effusion For two different gases at the same temperature, the ratio of their rates of effusion is given by the following equation: Thomas Graham (1805–1869)

68 Calculate the molar mass of a gas that effuses at a rate times N 2 Calculate the ratio of rate of effusion for oxygen to hydrogen

69 Ideal vs. Real Gases Real gases often do not behave like ideal gases at high pressure or low temperature Ideal gas laws assume 1.no attractions between gas molecules 2.gas molecules do not take up space based on the kinetic-molecular theory At low temperatures and high pressures these assumptions are not valid

70 Real Gas Behavior Because real molecules take up space, the molar volume of a real gas is larger than predicted by the ideal gas law at high pressures

71 The Behavior of Real Gases The volume of a real gas is larger than predicted by the ideal gas law.

72 Real Gas Behavior Because real molecules attract each other, the molar volume of a real gas is smaller than predicted by the ideal gas law at low temperatures

73 The Behavior of Real Gases Attractive forces between particles become more important at higher pressures.

74 van der Waals Equation Combining the equations to account for molecular volume and intermolecular attractions we get the following equation – used for real gases

75 PV/RT Plots

76 Structure of the Atmosphere The atmosphere shows several layers, each with its own characteristics The troposphere is the layer closest to the Earths surface Pollution added to the troposphere has a direct effect on human health and the materials we use because we come in contact with it The stratosphere is the next layer up(ozone layer) – less air mixingand weather in the stratosphere means that pollutants last longer before washing out The boundary between the troposphere and stratosphere is called the tropopause

77 Pollutant Gases, SO x SO 2 and SO 3, oxides of sulfur, come from coal combustion in power plants and metal refining – as well as volcanoes Lung and eye irritants Major contributors to acid rain 2 SO 2 + O H 2 O 2 H 2 SO 4 SO 3 + H 2 O H 2 SO 4

78 Pollutant Gases, NO x NO and NO 2, oxides of nitrogen, come from burning of fossil fuels in cars, trucks, and power plants – as well as lightning storms NO 2 causes the brown haze seen in some cities Lung and eye irritants Strong oxidizers Major contributors to acid rain 4 NO + 3 O H 2 O 4 HNO 3 4 NO 2 + O H 2 O 4 HNO 3

79 Stratospheric Ozone Ozone occurs naturally in the stratosphere Stratospheric ozone protects the surface of the earth from over-exposure to UV light from the Sun O 3 (g) + UV light O 2 (g) + O(g) Normally the reverse reaction occurs quickly, but the energy is not UV light O 2 (g) + O(g) O 3 (g)

80 Ozone Depletion Chlorofluorocarbons became popular as aerosol propellants and refrigerants in the 1960s CFCs pass through the tropopause into the stratosphere There, CFCs can be decomposed by UV light, releasing Cl atoms CF 2 Cl 2 + UV light CF 2 Cl + Cl Cl atoms catalyze O 3 decomposition and remove O atoms so that O 3 cannot be regenerated – NO 2 also catalyzes O 3 destruction Cl + O 3 ClO + O 2 O 3 + UV light O 2 + O ClO + O O 2 + Cl


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