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Chapter 11

Chapter 11. Gases. 10. 1 Kinetic Molecular Theory. State the kinetic-molecular theory of matter, and describe how it explains certain properties of matter. List the five assumptions of the kinetic-molecular theory of gases. Define the terms ideal gas and real gas.

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Chapter 11

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  1. Chapter 11 Gases

  2. 10. 1 Kinetic Molecular Theory State the kinetic-molecular theory of matter, and describe how it explains certain properties of matter. List the five assumptions of the kinetic-molecular theory of gases. Define the terms ideal gas and real gas. Describe each of the following characteristic properties of gases: expansion, density, fluidity, compressibility, diffusion, and effusion. Describe the conditions under which a real gas deviates from “ideal” behavior.

  3. What is the Kinetic Molecular Theory? • Break it down: • Kinetic: movement • Molecular: particles • Theory: tested ideas Tested ideas about the movement of particles! This theory is used to explain the energy and forces that cause the properties of solids, liquids, and gases.

  4. KMT of Gases • Ideal gas: hypothetical gas that satisfies all 5 ideas of KMT • pressure is not too high • temperature is not too low • Gases consist of large numbers of tiny particles that are far apart relative to their size. • Most of the volume is empty space • Collisions between gas particles and between particles and container walls are elastic collisions. • elastic collisionwhen there is no net loss of total kinetic energy

  5. KMT cont. • Gas particles are in continuous, rapid, random motion. • There are no forces of attraction between gas particles. • The temperature of a gas depends on the average kinetic energy of the particles of the gas. • The kinetic energy of any moving object is given by the following equation:

  6. Effusion • Because gases have motion, they can travel. • Effusion: process by which gas particles pass through a tiny opening What determines have fast a gas effuses? Mass • Gases at the same temperature have the same KE so… • Heavier gases  travel slower • Lighter gases  travel faster

  7. Gas Behavior KMT applies only to ideal gasses. Which parts are not true for real gases?

  8. 11.1 Gases and Pressures • Definepressure, give units of pressure, and describe how pressure is measured. • State the standard conditions of temperature and pressure and convert units of pressure. • Use Dalton’s law of partial pressures to calculate partial pressures and total pressures.

  9. 4 Variables of Gases • Pressure (P) Volume (V) • Temperature (T) Mols (n) • What causes pressure? •  collisions of the gas molecules with each other and with surfaces with which they come into contact. •  depends on volume (mL or L), temperature (oF, oC, K), and the number of molecules present (mol, mmol).

  10. Equation for Pressure Pressure (P): the force per unit area on a surface. Pressure = Force Area  More force on a given area, the greater the pressure.  smaller the area is on which a given force acts, the greater the pressure.

  11. Pressure Video

  12. Relationship Between Pressure, Force and Area

  13. Measuring Pressure • barometer: device used to measure atmospheric pressure • Pressure of atmosphere supports a column of Hg about 760 mm above surface of mercury in dish • Can change depending on weather & elevation

  14. Measuring Pressure

  15. Units for Measuring Pressure • mm Hg : millimeters of mercury • A pressure of 1 mm Hg is also called 1 torr in honor of Torricelli for his invention of the barometer. • atm : atmosphere of pressure • kPa : kiloPascal • Others… • psi : pounds per square inch • Bar • torr 1 atm = 101.3 kPa = 760 mmHg = 760 torr

  16. Review- Units of Pressure

  17. Pressure Conversions The average atmospheric pressure in Denver, Colorado is 0.830 atm. Express this pressure in: a. millimeters of mercury (mm Hg) and b. kilopascals (kPa) Given:atmospheric pressure = 0.830 atm Unknown: a. pressure in mm Hg b. pressure in kPa

  18. Pressure Conversions Answers A) B)

  19. STP • STP : Standard Temperature & Pressure • 1.0 atm (or any of units of equal value) • 0 oC • Used by scientists to compare volumes of gases

  20. Dalton’s Law of Partial Pressures • The pressure of each gas in a mixture is called the partial pressure. • John Dalton discovered that the pressure exerted by each gas in a mixture is independent of that exerted by other gases present. • Dalton’s law of partial pressures: the total pressure of a gas mixture is the sum of the partial pressures of each gas.

  21. Dalton’s Law of Partial Pressures • Dalton derived the following equation: PT = P1 + P2 + P3 + … Total Pressure = sum of pressures of each individual gas

  22. Dalton’s Law of Partial Pressures

  23. Gases Collected by Water Displacement • Water molecules at the liquid surface evaporate and mix with the gas molecules. Water vapor, like other gases, exerts a pressure known as vapor pressure. • Gases produced in the laboratory are often collected over water. The gas produced by the reaction displaces the water in the reaction bottle.

  24. Particle Model for a Gas Collected Over Water

  25. Gases Collected by Water Displacement (ctd) • Step 1: Raise bottle until water level inside matches the water level outside. (Ptot = Patm) • Step 2: Dalton’s Law of Partial Pressures states: Patm = Pgas + PH2O To get Patm, record atmospheric pressure. • Step 3: look up the value of PH2Oat the temperature of the experiment in a table, you can then calculate Pgas.

  26. Dalton’s Law of Partial Pressures Sample Problem KClO3decomposes and the oxygen gas was collected by water displacement. The barometric pressure and the temperature during the experiment were 731.0 torr and 20.0°C. respectively. What was the partial pressure of the oxygen collected? Given: PT = Patm= 731.0 torr PH2O = 17.5 torr(vapor pressure of water at 20.0°C, from table A-8 in your book) Patm = PO2+ PH2O Unknown:PO2in torr

  27. Dalton’s Law Sample Problem Solution • Solution: Patm = PO2+ PH2O PO2 = Patm- PH2O • substitute the given values of Patm and into the equation: PO2 =731.0 torr – 17.5 torr = 713.5 torr

  28. Mole Fractions (X) Mole fraction of a gas(XA) = mole fraction: ratio of the number of moles of one component of a mixture to the total number of moles Moles of gas A (nA) Total number of moles of a gas (ntot)

  29. Calculating Partial Pressure Partial pressures can be determined from mole fractions using the following equation: PA= XAPT

  30. 11.2 The Gas Laws • Use the kinetic-molecular theory to explain the relationships between gas volume, temperature and pressure. • Use Boyle’s law to calculate volume-pressure changes at constant temperature. • Use Charles’s law to calculate volume-temperature changes at constant pressure. • Use Gay-Lussac’s law to calculate pressure-temperature changes at constant volume. • Use the combined gas law to calculate volume-temperature-pressure changes.

  31. Boyle’s Law • If you increase the pressure on a gas in a flexible container, what happens to the volume? • If you decrease the pressure, what happens the volume? • Pressure and volume are ________ related. P1V1 = P2V2 Variables: pressure & volume Constant: temperature, amount of gas

  32. Boyle’s Law

  33. Boyle’s Law Video

  34. Boyle’s Law Problem A sample of oxygen gas has a volume of 150.0 mL when its pressure is 0.947 atm. What will the volume of the gas be at a pressure of 0.987 atm if the temperature remains constant? P1 = 0.947 atm P2 = 0.987 atm V1 = 150.0 mL V2 = ?

  35. Boyle’s Law Problem Solution

  36. Charles’ Law • If you increase the temperature of gas, what will happen to the volume? • If you decrease the temperature of a gas, what will happen to the volume? • Volume and temperature are ______ related. • Variables: volume & temperature • Constant: pressure & amount of gas

  37. Charles’ Law

  38. Charles’ Law Video

  39. Temperature in Charles Law • To Convert to Kelvin K = 273 + °C. • absolute zero: when all motion stops O K = -273 oC

  40. Charles’ Law Problem A sample of neon gas occupies a volume of 752 mL at 25°C. What volume will the gas occupy at 50°C if the pressure remains constant? Temperature must be in KELVIN!!! V1 = 752 mL V2 = ? T1 = 25°C T2 = 50°C

  41. Charles’ Law Sample Problem Solution

  42. Gay-Lussac’s Law • If you increase the temperature of a gas what will happen to the pressure? • If you decrease the temperature of gas what will happen to the pressure? • Pressure and temperature are _____ related. • Variables: pressure & temperature • Constant: volume & amount of gas

  43. Gay-Lussac’s Law

  44. GL Law Video

  45. Gay-Lussac’s Law Problem The gas in a container is at a pressure of 3.00 atm at 25°C. Directions on the container warn the user not to keep it in a place where the temperature exceeds 52°C. What would the gas pressure in the container be at 52°C? Temperature must also be in KELVIN!!! P1 = 3.00 atm P2 = ? T1 = 25°C T2 = 52°C

  46. Gay-Lussac’s Law Problem Solution P2 = P1T2 = (3.00 atm) (325 K) = 3.27 atm T1 298 K

  47. The Combined Gas Law Constant: amount of gas • combined gas law: used when pressure, temperature, and volume change within a system NOTE: P & V are directly related to T, while P is inversely related to V

  48. Combined Gas Law Problem A helium-filled balloon has a volume of 50.0 L at 25.0°C and 1.08 atm. What volume will it have at 0.855 atm and 10.0°C? Temperature must be in KELVIN!! P1 = 1.08 atm P2 = 0.855 atm V1 = 50.0 L V2 = ? T1 = 25.0°C T2 = 10.0°C

  49. Combined Gas Law Problem Solution

  50. End of Material for Quiz #1

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