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Chapter 5 Gases

Chapter 5 Gases. Introduction to Gases. Gases have been known to exist since ancient times The Greeks considered gases one of the four fundamental elements of nature 18 th Century

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Chapter 5 Gases

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  1. Chapter 5Gases

  2. Introduction to Gases • Gases have been known to exist since ancient times • The Greeks considered gases one of the four fundamental elements of nature • 18th Century • Lavoisier, Cavendish and Priestley: Air is primarily nitrogen and oxygen, with trace components of argon, carbon dioxide and water vapor

  3. 5.1 Measurements of Gases • Key variables in the measurement of gases: • Volume • Amount • Temperature • Pressure

  4. Volume, Amount and Temperature • A gas expands uniformly to fill the container in which it is placed • The volume of the container is the volume of the gas • Volume is commonly measured in L, mL, or cm3 • The temperature of a gas must be indicated on the Kelvin scale • Recall that K = °C + 273.15 • Amount of a gas is measured in moles

  5. Pressure • Pressure is force per unit area • In the English system, pounds per square inch or psi • Standard atmospheric pressure is 14.7 psi

  6. Measuring Pressure – The Barometer

  7. The Barometer • The barometer measures atmospheric pressure in terms of the height of a column of liquid mercury • The atmosphere exerts a force on a pool of mercury, causing it to rise • Standard atmospheric pressure is a column of mercury 760 mm high • Mercury is used because it is liquid at room temperature and has a very high density

  8. The Manometer

  9. Gas Pressure Measurement • The manometer measures gas pressure by differential • The height of the column of liquid is proportional to the pressure of the gas in the container • Gas pressure can be more or less than atmospheric pressure

  10. Other Units of Pressure 1.013 bar = 1 atm = 760 mmHg = 14.7 psi = 101.3 kPa = 29.92 in Hg = 760 torr

  11. Example 5.1

  12. 5.2 The Ideal Gas Law Basic Gas Laws: Charles’ Law Boyle’s Law Gay-Lussac’s Law Relationship between moles and pressure Combination Law

  13. Charles’ Law -Volume and Temperature are directly proportional -Volume units must match and temperature units must be in Kelvin V1 = V2 T1 T2

  14. Boyle’s Law -Volume and pressure are inversely proportional -Volume and pressure units must match P1V1 =P2V2

  15. Gay-Lussac’s Law -Pressure and Temperature are directly proportional -Pressure units must match and temperature units must be in Kelvin P1 = P2 T1 T2

  16. Relationship between moles and pressure -Moles and pressure are directly proportional -pressure units must match and n is moles P1 = P2 n1 n2

  17. Combination Law -Looking at the basic gas laws they can easily be summed together to form one law V1P1 = V2P2 n1T1 n2T2

  18. Figure 5.3

  19. The Ideal Gas Law

  20. Relationships between the variables: • Volume vs. temperature = direct • Volume vs. moles = direct • Volume vs. pressure = inverse • Pressure vs. temperature = direct • Pattern: • Variables are a product = inverse • Variables are a quotient = direct

  21. Temperature Effects – Charles’s Law

  22. The Ideal Gas Law PV = nRT • R is the gas constant • Units of R:

  23. Table 5.1 – Units of R

  24. Standard Temperature and Pressure • STP • 1 atm P • 273 K • At STP, the molar volume of a gas can be calculated as follows:

  25. 5.3 Gas Law Calculations • Final and initial problems (Basic laws; 1’s and 2’s) • Single-state problems (PV = nRT) • Density problem (D=MMxP/RT)

  26. Example 5.2

  27. Example 5.3

  28. Molar Mass and Density • Density = mass/volume • Recall that the molar mass has units of grams per mole (MM = m/n) • Now, look at the ideal gas law: Substitute m/MM for n PV = mRT MM Substitute Density for m/V P=mRT P=DRT D=P MM V MM MM RT

  29. Density of Gases • Density is an intensive property • Does not depend on the amount of substance • Density of a gas does depend on: • Pressure • Temperature • Molar mass

  30. Balloons

  31. Example 5.4

  32. 5.4 Stoichiometry in Gaseous Reactions • Gases can participate as reactants or products in any chemical reaction • All substances in a balanced chemical equation are related by mole ratios

  33. Example 5.5

  34. Example 5.6

  35. 5.5 Gas Mixtures: Partial Pressures • The ideal gas law applies to all gases, so it applies to mixtures of gases as well • A new term is needed for a mixture of gases • Partial pressure, the part of the total pressure due to each gas in the mixture • Sum of the partial pressures is the total pressure

  36. Dalton’s Law of Partial Pressures • The total pressure of a gas mixture is the sum of the partial pressures of the gases in the mixture PT = P1 + P2 + P3 …

  37. Vapor Pressure • The vapor pressure of a substance is the pressure of the gaseous form of that substance • Vapor pressure is an intensive property • Vapor pressure depends on temperature

  38. Collecting a Gas Over Water • When a gas is collected over water, the total pressure is the pressure of the gas plus the vapor pressure of water

  39. Wet Gases • P H2O is the vapor pressure of water • P H2O is dependent on temperature • Consider H2 gas collected over water:

  40. Example 5.8

  41. 5.6 Kinetic Theory of Gases The kinetic-molecular model 1. Gases are mostly empty space. The total volume of the molecules is small 2. Gas molecules are in constant, random motion 3. Collisions of gas particles are elastic 4. Gas pressure is caused by collisions of molecules with the walls of the container

  42. Figure 5.7 – The Kinetic Molecular Model

  43. Results from Kinetic Energy of Translational Motion • At a given temperature, all molecules of all gases have the same average kinetic energy • The average kinetic energy of a gas particle is directly proportional to the Kelvin temperature

  44. Effusion of Gases • Diffusion • Gases move through space from a region of high concentration to a region of low concentration • You can smell an apple pie baking as the particles responsible for the odor diffuse through the room • Effusion • Gas particles will escape through a small hole in a container • Air will slowly leak out of a tire or balloon through pores in the rubber

  45. Graham’s Law of Effusion • The rate at which gas B escapes divided by the rate at which gas A escapes is equal to the square root of the ratio of the molar mass of gas A to gas B

  46. Effusion of Gases

  47. Example 5.11

  48. Real Gases • Recall that the molar volume of a gas at STP is 22.4 L (Avogadro’s Hypothesis) • There are deviations from this volume that depend on the gas being studied • The molar volume of a real gas is less than that calculated by the ideal gas law

  49. Table 5.2

  50. Liquefaction of Gases • All gases can be liquefied • Lowering the temperature • Increasing the pressure • When a gas is liquefied, the attractive forces between gas particles becomes significant • The closer a gas is to the liquid state, the more it will deviate from the ideal gas law

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