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The Gaseous State

The Gaseous State. Chapter 10. Objectives. Understand the definition of pressure. Use the definition to predict and measure pressure experimentally Describe experiments that show relationships between pressure, temperature, volume, and moles of a gas sample

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The Gaseous State

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  1. The Gaseous State Chapter 10

  2. Objectives • Understand the definition of pressure. Use the definition to predict and measure pressure experimentally • Describe experiments that show relationships between pressure, temperature, volume, and moles of a gas sample • Use empirical gas laws to predict how change in one of the properties of a gas will affect the remaining properties. • Use empirical gas laws to estimate gas densities and molecular masses. • Use volume-to-mole relationships obtained using the empirical gas laws to solve stoichiometry problems involving gases.

  3. Objectives 6. Understand the concept of partial pressure in mixtures of gases. 7. Use the ideal kinetic-molecular model to explain the empirical gas laws. 8. List deficiencies in the ideal gas mode3el that will cause real gases to deviate from behaviors predicted by the empirical gas laws. Explain how the model can be modified to account for these deficiencies.

  4. Definition of Gas Gas: large collection of particles moving at random throughout a volume that is primarily empty space. Have relatively large amount of energy. Gas pressure: due to collisions of randomly moving particles with the walls of the container. Force/unit area

  5. Definition of Gases • STP: 0°C, and 1 atmosphere pressure • Elements that exist as gases at STP: hydrogen, nitrogen, oxygen, fluorine, chlorine and Noble Gases • Ionic compounds are all solids • Molecular compounds - depends on the intermolecular forces. Most are liquids and solids. Some are gaseous • http://www.chemistry.ohio-state.edu/betha/nealGasLaw/fr1.1.html

  6. Properties of Gases • Assume the volume and shape of their container • Compressible • Mix evenly and completely when confined to the same container • Lower densities than liquids and solids • Allotropes: O2 ↔O3

  7. Kinetic Molecular Theory of Gases • Tiny particles in continuous motion ( the hotter the gas, the faster the molecules are moving) with negligible volume compared to volume of container. • Molecules are far apart from each other • Do not attract or repel each other (?). • All collisions are elastic (gas does not lose energy when left alone). • The energy is proportional to Kelvin temperature. At a given temperature all gases have the same average KE.

  8. Properties of Gases

  9. Properties of Gases

  10. Atmospheric Pressure Intensive or Extensive Property?

  11. Pressure • Pressure is due to collisions between gas molecules and the walls of the container. Magnitude determined by: force of collisions and frequency. • Pressure: force per unit area: P =F/A • Standard temperature: 0ºC = 273.15 K • Standard pressure: 1 atm in US; 1 bar elsewhere

  12. Pressure

  13. Pressure: Examples • How much pressure does an elephant with a mass of 2000 kg and total footprint area of 5000 cm2 exert on the ground? • Estimate the total footprint area of a tyrannosaur weighing 16 000 kg. Assume it exerts the same pressure on its feet that the elephant does.

  14. Pressure • Measuring pressure: • Strategy: • Relate pressure to fluid column heights • You can’t draw water higher than 34 feet by suction alone. Why? • Hypothesis: atmospheric pressure supports the fluid column • Develop the equation

  15. Measuring Pressure

  16. Pressure: Barometer Barometer measures atmospheric pressure as a mercury column height.

  17. Pressure: Open-Manometer Manometer measures gas pressure as a difference in mercury column heights. Two types: closed manometer open manometer

  18. Measuring Gas Pressure Closed-manometer: the arm not connected to the gas sample is closed to the atmosphere and is under vacuum. Explain how you can read the gas pressure in the bulb.

  19. Pressure: Examples 3. Calculate the difference in pressure between the top and the bottom of a vessel exactly 76 cm deep filled at 25 ºC with a) water; b) mercury (d = 13.6 g/cm3) (7.43 x 103 Pa;100.9 x 103 Pa) 4. How high a column of air would be necessary to cause the barometer to read 76 cm of mercury, if the atmosphere were of uniform density 1.2 kg/m3? dHg = 13.53 kg/m3 (8.6 km) 5. A Canadian weather report gives the atmospheric pressure as 100.2 kPa. What is the pressure in atmospheres? Torr? Mm Hg?

  20. The Gas Laws: State of Gas

  21. The Gas Laws: State of Gas • Any equation that relates P, V, T, and n for a material is called an equation of state. • Experiment shows PV = nRT is an approximate equation of state for gases. • R is the gas law constant • Determined by measuring P, V, T, n and computing R = PV/nT • Value depends on units chosen for P, V, T • Notice: 1 Joule = 1 N m = 1(Pa) (m3)

  22. The Gas Laws Gas lawsdeal with theMACROSCOPIC view of gases and we try to explain the macroscopic properties by examining the microscopic behaviors (many molecule behaviors) http://www.phy.ntnu.edu.tw/ntnujava/viewtopic.php?t=42 http://jersey.uoregon.edu/vlab/Piston/index.html

  23. Prentice Hall Simulations of Gas Laws • http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/chapter5/deluxe.html • http://cwx.prenhall.com/bookbind/pubbooks/hillchem3/chapter5/deluxe.html

  24. Boyle’s Law: Experiment Relate volume to pressure when everything else is constant. Experiment: trapped air bubble at 298 K Graphs?

  25. Boyle’s Law: Experiment Relate volume to pressure when everything else is constant. Experiment: trapped air bubble at 298 K Graphs?

  26. Boyle’s Law: Volume/Pressure Relationship At constant n, and T, the volume of a gas decreases proportionately as its pressure increases. If the pressure is doubled, the volume is halved.

  27. Boyle’s Law: Volume/Pressure Relationship What happens to the volume of the gas as the pressure increases? Mathematical Relationship?

  28. Plot of Boyle’s Law V versus P V versus 1/P Type of Graphs?

  29. Boyle’s Law Boyle’s Law – the volume of a fixed amount of gas at constant temperature and constant number of moles is inversely proportional to the gas pressure. MOLECULAR VIEW

  30. Boyle’s Law Boyle’s Law – the volume of a fixed amount of gas at constant temperature and constant number of moles is inversely proportional to the gas pressure. MOLECULAR VIEW: Confining molecules to a smaller space increases the number (frequency) of collisions, and so increases the pressure

  31. Charles' Law (V/T Relationships) Relate volume to temperature, everything else is constant. Experiment: He bubble trapped at 1 atm.

  32. Charles' Law (V/T Relationships) Relate volume to temperature, everything else is constant. Experiment: He bubble trapped at 1 atm.

  33. Charles’ Law: Volume/Temperature Relationships At constant n and P, the volume of a gas increases proportionately as its absolute temperature increases, If the absolute temperature doubles, the volume is doubled. K = ºC + 273

  34. Charles’ Law A plot of V versus T for a gas sample. What type of graph? Equation?

  35. Charles' Law Kinetic Interpretation of Charles's Law? Why higher pressure? Equation? Frequency and force of collision…

  36. Charles’ Law The volume of the gas is directly proportional to its Kelvin temperature, when everything else is constant. MOLECULAR VIEW

  37. Charles’ Law The volume of the gas is directly proportional to its Kelvin temperature, when everything else is constant. MOLECULAR VIEW Raising temperature increases the number of collisions and force of collisions (KE increases) with container wall. If the walls are flexible, they will be pushed back and the gas expands.

  38. Charles’ Law Assume that you have a sample of gas at 350 K in a sealed container, as represented in (a). Which of the drawings (b) – (d) represents the gas after the temperature is lowered from 350 K to 150 K

  39. Gay Lussac’s Law The pressure of the gas is directly proportional to its Kelvin temperature, when everything else is constant. Molecular View;

  40. Gay Lussac’s Law The pressure of the gas is directly proportional to its Kelvin temperature, when everything else is constant. Molecular View; • Raising the temperature increases the number of collisions and the kinetic energy of the molecules. More collisions with greater energy (force) means higher pressure.

  41. Combined Gas laws

  42. Avogadro’s Law: Relates n to Volume

  43. Volume of Real Gases at STP

  44. Avogadro’s Law: Relates n to Volume At constant T and P, the volume of a gas is directly proportional to moles of gas. Molar volume is almost independent of the type of gas. Samples of two gases with the same V, P, T contain the same number of molecules. MOLECULAR VIEW

  45. Avogadro’s Law: Relates n to Volume At constant T and P, the volume of a gas is directly proportional to moles of gas. Molar volume is almost independent of the type of gas. Samples of two gases with the same V, P, T contain the same number of molecules (moles). MOLECULAR VIEW Type of gas does not influence distance between molecules too much.

  46. Avogadro’s Law: Example 6 Show the approximate level of the movable piston in drawings (a) and (b) after the indicated changes have been made to the initial gas sample.

  47. Avogadro’s Law: Answer to Example 6

  48. Example 7 Show the approximate level of the movable piston in drawings (a), (b), and (c ) after the indicated changes.

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