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Chapter 12: Gases and Their Properties

Chapter 12: Gases and Their Properties. Properties of Gases. Gases form homogeneous mixtures Gases are compressible All gases have low densities air 0.0013 g/mL water 1.00 g/mL iron 7.9 g/mL Gases expand to fill their containers uniformly A gas exerts a pressure.

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Chapter 12: Gases and Their Properties

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  1. Chapter 12: Gases and Their Properties

  2. Properties of Gases • Gases form homogeneous mixtures • Gases are compressible • All gases have low densities • air 0.0013 g/mL • water 1.00 g/mL • iron 7.9 g/mL • Gases expand to fill their containers uniformly • A gas exerts a pressure

  3. Kinetic Molecular Theory • Gases consist of molecular particles moving in straight lines at any given instant. • Molecules collide with each other and the container walls without any net loss of energy. • Gas molecules behave independently -- attractive/repulsive forces between them are negligible. • Gas molecules are widely spaced, the actual volume of molecules is negligible compared to the space they occupy. • The average kinetic energy of the gas particles is proportional to the temperature.

  4. Kinetic Graph

  5. Pressure, Volume, and Temperature Relationships • Pressure, volume, and temperature units • Boyle's law • Charles' law • Gay-Lussac's law • Avogadro's law

  6. A. Pressure, Volume, and Temperature Relationships force unit area SI: 1 pascal (Pa) = 1 N/m2 common: 1 atm = 760 mm Hg = 760 torr = 14.7 lb/in2 (psi) = 101.325 kPa = 1.013 bar • Pressure = • -measured with barometer • or manometer Volume: mL, L Temperature: K

  7. Pressure • Defined as force per unit area

  8. Atmospheric pressure

  9. Barometer

  10. Manometers

  11. Pressure Units • 101.325 kPa • = 760 mmHg • =760 torr • =1 atm • =30 in Hg • =14.7 psi

  12. constant P B. Boyle’s law Robert Boyle, 1662: for a sample of gas at constant T, V 1/P V = or PV = constant  P1V1 = P2V2 (at constant T)

  13. Boyles Law

  14. V T V1 T1 V2 T2 C. Charles’ law Jacques Charles, 1787: for a sample of gas at constant P, V T (K) V = constant x T or = constant  = (at constant P)

  15. Charles Law

  16. P T P1 T1 P2 T2 D. Gay-Lussac’s law Joseph Gay-Lussac, ~1800: for a sample of gas at constant V, P T (K) P = constant x T or = constant  = (at constant V)

  17. V n V1 n1 V2 n2 E. Avogadro’s law Amadeus Avogadro, 1811: at constant T and P, V n V = constant x n or = constant  = (at constant T & P)

  18. PV nT The Ideal gas laws • Since PV = constant, P/T = constant, V/T = constant, and V/n = constant •  = constant, R (universal gas constant) • or PV = nRT

  19. PV nT (1.00 atm)(22.4 L) (1.00 mol)(273 K) STP and molar volume STP: 0ºC (273 K) and 1.00 atm (760 torr) molar volume = average volume occupied by one mole of gas at STP = 22.4 L/mol R = = = 0.0821 L·atm/mol·K = 6.24 x 104 mL·torr/mol·K

  20. m n PV RT mRT PV dRT P m V Gas densities and molar mass • mw = • since n =  mw = • and since d =  mw = • e.g., If one finds that a 0.108-g sample of gas occupies a volume of • 238 mL at 25ºC and 525 torr, what is the molecular weight of the gas?

  21. Dalton’s law of partial pressures1. Dalton’s law partial pressure, p = pressure exerted by each gas in a mixture of gases e.g., If 6.00 g of O2 and 9.00 g of CH4 are placed in a 15.0-L container at 0º, what is the partial pressure of each gas and the total pressure in the container.

  22. Dalton’s law of partial pressures2. mole fraction PT = p1 + p2 + p3…  p1 = X1PT e.g., Air is 78 mol % N2 and 22 mol % O2. What is the partial pressure of each gas if the atmospheric pressure is 713 torr?

  23. Dalton’s law of partial pressures3. collecting gases over water vapor pressure = pressure exerted by evaporation of a liquid (or sublimation of a solid) pvap vacuum equilibrium at some T liquid or solid • pvap increases with T • e.g., H2O T pvap • 0ºC 4.6 torr • 25ºC 23.8 torr • 100ºC 760 torr (boils!)

  24. Dalton’s law of partial pressures3. collecting gases over water • Mixture of gas and water vapor. • PT = pgas + pvap(H2O)

  25. Mole Fraction

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