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Kinetic Theory

Kinetic Theory . Behavior of Gases. Properties of a gas. Kinetic Theory of Gases Gas is mostly empty space and the particles are in constant random motion. The distance between the particles is so large that it is assumed to be insignificant. This makes gases compressible.

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Kinetic Theory

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  1. Kinetic Theory Behavior of Gases

  2. Properties of a gas • Kinetic Theory of Gases • Gas is mostly empty space and the particles are in constant random motion. • The distance between the particles is so large that it is assumed to be insignificant. • This makes gases compressible. • The assumption that there exist no attractive forces between particles.

  3. Kinetic Energy and Temperature. 1. When there is an increase in temperature, there is an increase in kinetic energy and when there is an increase in kinetic energy, there is an increase in pressure. 2. Absolute zero (0 K or -273.15°C) is the temperature at which all motion of particles theoretically cease. 3. The Kelvin scale reflects the relationship between temperature and average kinetic energy and is directly proportional to the average kinetic energy of the particle of a substance.

  4. Variable That Describe a Gas • P = Pressure in atmospheres,mm Hg, or kilopascals (kPa) • V = Volume in liters • T = Temperature in Kelvin • n = Moles

  5. Factors Affecting Gas Pressure • Amount of Gas- low pressure = fewer particles ; high pressure = more particles • Volume - decrease the volume = increase in pressure • Temperature- increase temperature = increase pressure

  6. Avogardro’s Principle • Equal volumes of gases at the same temperature and pressure contain the same number of particles. • Gas is mostly empty space and can be compressed. • 1mole=6.02 x 1023 particles =22.4dm3 • Standard temperature = 0ºC= 273 K • Standard pressure = 1 atm = 101.3 kPa = 760 mm Hg

  7. 1. Pressure is the force per unit area. 2. Pressure is also the force of collisions and the number of collision with the wall of the container. 3. To measure the thinness of air, a manometer is used. There are two type of manometers: a. open-armed manometers determines the difference between air pressure and the gas being studied. b. closed-armed manometers are independent of air pressure. GasPressure

  8. Measuring Gas Pressure • The open-tube manometerpressure of the gas is given by h (the difference in mercury levels) in units of torr or mmHg. Atmospheric pressure pushes on the mercury from one direction, and the gas in the container pushes from the other direction. In a manometer, since the gas in the bulb is pushing more than the atmospheric pressure, you add the atmospheric pressure to the height difference: • Pgas > PatmGas pressure = atmospheric pressure +h (height of the mercury) • Pgas < Patm Gas pressure = atmospheric pressure -h (height of the mercury) • You could be asked about a closed-tube manometer. Closed-tube manometers look similar to regular manometers except that the end that’s open to the atmospheric pressure in a regular manometer is sealed and contains a vacuum. In these systems, the difference in mercury levels (in mmHg) is equal to the pressure in torr.

  9. Measuring Gas pressure

  10. Barometers A barometer is a scientific instrument used in meteorology to measure atmospheric pressure. Barometer measure the air pressure changes . Air pressure and differences in pressure are among the most important weather makers. The centers of storms are areas of relatively low air pressure, compared to pressures around the storm. High air pressure generally brings good weather. Keeping track of how the pressure is changing is important for forecasting the weather. Differences in air pressure between places cause the winds to blow - air moves from high toward low pressure.

  11. Vaporization of Gas 1. The conversion of a liquid to a gas or vapor is called vaporization. 2. When the conversion occurs at the surface of a liquid that is not boiling, it is called evaporation. 3. Only molecules with the a certain minimum kinetic energy can break free of the surface. 4. Evaporation is also a cooling process because the remaining molecules have a lower average kinetic energy then the molecules that escaped and the temperature decreases. 5. The vapor particles collide with the wall and produce a vapor pressure. 6. In a closed container, a liquid and it vapor will reach equilibrium at a specific pressure for any particular temperature. 7. As time passes, some of the particles return to the liquid and an equilibrium is established. 8. If the temperature increase, the vapor pressure will increase causing a shift in equilibrium.

  12. Vapor Pressure • Here's a REAL IMPORTANT point about vapor pressure: the vapor pressure depends ONLY on the temperature!! • That is true because as the temperature goes up, there are more and more molecules with the right combination of energy and direction to break free of the liquid's surface. • Finally, you might think that, if you were to put more liquid in, the vapor pressure would go up. That's NOT correct. Keep in mind that there are two opposing processes at work: (1) molecules leaving the liquid's surface and (2) molecules returning to the liquid at its surface. • Process #1 depends only on temperature and process #2 depends only on how many molecules are in the vapor phase. Adding more liquid affects neither process.

  13. Vapor Pressure • Here is a question about vapor pressure: • If a liquid is sealed in a container, kept at constant temperature how does its vapor pressure change over time? Advice: reword the below, if you happen to get the above question (or something like it) as part of your homework. • Answer: let's assume there was zero vapor at the instant of sealing the liquid in the container. The pressure in the sealed container would increase up to a certain value, which depends on two things: (1) the identity of the substance and (2) the temperature. • Once the pressure got to that certain value, it would remain there and NOT increase any more. (Keep in mind that this is all happening at constant temperature.) • Also, remember that molecules will continue to move from the liquid into the gas state. However, the rate of that will equal the rate of molecules returning from the gas state to the liquid state. • Lastly, remember that for the above to work, there must be some liquid remaining in the sealed container. Vapor pressure is the behavior of the gas while still in contact with its liquid (or solid).

  14. Volatile and Nonvolatile • Notice the word "nonvolatile" in the title. The volatility of a substance refers to the readiness with which it vaporizes. Generally speaking, substances with a boiling point below 100 °C are considered volatile and all others are called nonvolatile. Ethyl alcohol and pentane are examples of volatile substances; sugar and sodium chloride are considered nonvolatile.

  15. Vapor Evaporation and Condensation Because of hydrogen bonding, water absorbs a large amount of heat as it evaporates or vaporizes. Hydrogen bonding is an extensive network of hydrogen bond attracted to the non hydrogen ion in a compound and tightly holds the molecules in liquid together. Those bonds must be broken before a liquid can changes from a liquid state to the vapor state. 4. This process is called evaporation. 5. This occurs naturally without the addition of heat energy. 6.The reverse of evaporation is condensation.

  16. Conversions • Celsius to Kelvin • K = °C + 273.15 °C = K – 273.15 • Atmospheres to millimeters Mercury • 1 atm = 760 mm Hg • 1 torr = 760 mm Hg • Atmospheres to Kilopascals • 1 atm = 101.3 kPa • Kilopascals to millimeters Mercury • 1 kPa = 7.501 mm Hg (760 mm Hg/ 101.3 kPa)

  17. Gas Laws • Combined - P1V1 = P2V2 T1 T2 • Boyles- P1V1 = P2V2 • Charles - V1 = V2 T 1 T2 • Guy Lussac - P1 = P2 T1 T2 • Dalton’s -PTotal= P1 +P2 + P3 +P4….. • Graham’s -

  18. Boyle’s Law • Boyle’s law • For a given mass of a gas at a constant temperature… • The volume of a gas varies inversely with the pressure. • P1V1=P2V2

  19. Charles’ Law • Charles’ law • For a given mass of a gas at a constant pressure… • The volume of a gas varies directly proportional with the temperature. • V1/T1=V2 /T2

  20. Guy Lussac’s Law • Guy Lussac’s law • For a given mass of a gas at a constant volume… • The pressure of a gas varies directly proportional with the temperature. • P1/ T1=P2/T2

  21. Dalton’s Law of Partial Pressure • The individual pressure of each gas in a mixture is called partial pressure of a gas. • John Dalton suggested that the total pressure of a mixture of gases is equal to the sum of the partial pressure of the individual gases in the mixture.(constant volume and temperature. • P total = P1 + P2 + P3 +…..

  22. Gas Pressure practice • http://www.chm.davidson.edu/vce/gaslaws/pressure.html • http://www.chemteam.info/GasLaw/KMT-Gas-Laws.html • http://www.kentchemistry.com/links/GasLaws/Pressure.htm

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