1. Gases: Chapter 10 10.1 – Characteristics of Gases Physical properties of gases are all similar. Composed mainly of nonmetallic elements with simple formulas and low molar masses. Unlike liquids and solids, gases  expand to fill their containers.  are highly compressible.  have extremely low densities. Two or more gases form a homogeneous mixture.

2. 10.2 -- Pressure Pressure = force per unit area (P = F/A) SI Unit = N/m 2 = 1 pascal (Pa). kPa is more common Atmospheric pressure (the force exerted by the atmosphere on a given surface area) -- measured with a barometer: 1 atm = 760mmHg = 760 torr = 101.3kPa Standard pressure = 1atm

3. Manometer – measures the difference in pressure between atmospheric pressure the pressure of a gas in a container. P gas = P atm + h Note: h is positive if P gas > P atm h is negative if P gas < P atm

4. 10.3 – the Gas Laws Boyle’s Law: At a constant temp., the pressure and volume of a gas are inversely proportional. PV = a constant, or P 1 x V 1 = P 2 x V 2

5. Charles’s Law: The volume of a fixed amount of gas at constant pressure is directly proportional to its Kelvin temperature V 1 /T 1 = V 2 /T 2 (note: temp must be in Kelvin)

6. Avogadro’s Law: At a fixed T and P, the volume of a gas is directly proportional to the number of moles of gas. At STP (1atm and 0°C), 1 mole of a gas = 22.4L (molar volume of a gas at STP)

7. 10.4 – The Ideal Gas Law The Ideal Gas Law (an ideal gas is a hypothetical gas that follows the ideal gas law under all pressure, volume, and temperature conditions) PV=nRT R=.0821 L atm/mol K (see other values on page 408) The Combined Gas Law

8. 10.5 – Applications of the Ideal Gas Law Determination of Gas Densities: using the ideal gas law, set the volume equal to 1L and solve for moles (that gives you moles/L) – then convert to g/L using the molar mass. Or use the following equation: d = MP/RT, where M is the molar mass of the gas (note: this equation is not on the equation sheet, so you would have to memorize it)

9. Determination of the molar mass of a gas: calculate moles with the ideal gas equation, then divide mass/moles Or use the following equation: M = mRT/PV, where M is the molar mass. (note: this equation is also not on the equation sheet)

10. Gases in Reaction Stoichiometry Law of Combining Volumes: If all the gases in the reaction are at the same conditions, then the mole ratio is also the volume ratio. For other problems, use the ideal gas law to relate the moles of a gas to other gas units, such as volume and pressure (you can’t use the 22.4L/mole conversion if the conditions aren’t STP)

12. 10.6 – Gas Mixtures and Partial Pressures If two gases that don’t react are combined in a container, they act as if they are alone in the container. Dalton’s Law of Partial Pressures: P total = P 1 + P 2 + P 3 + … Mole fraction = moles of one component / total # of moles in mixture. For gases in a mixture, the mole fraction = the pressure fraction (P gas /P total ) So, P A = P tot x X A

13. Collecting a gas over water: -there’s a mixture of the desired insoluble gas and water vapor P total = P bar = P gas + P H2O

14. 10.7 – The Kinetic-Molecular Theory of Gases 1)Gases consist of large numbers of molecules that are in continuous, random motion. 2)The combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained. 3)Attractive and repulsive forces between gas molecules are negligible. 4)Energy can be transferred between molecules during collisions, but the average kinetic energy of the molecules does not change with time, as long as the temperature of the gas remains constant. 5)The average kinetic energy of the molecules is proportional to the absolute (kelvin) temperature

15. The Maxwell-Boltzmann distribution describes the distribution of particle speeds in an ideal gas. Notice that: 1. As the molar mass of the gas increases, the speed of the molecules decreases and the range gets smaller. 2. As the temp. increases, the speed increases and the range gets larger.

16. How Fast Do Gas Molecules Move? Temperature is related to the average kinetic energy of the gas molecules. Individual molecules can have different speeds of motion. The figure shows three different speeds:  u mp is the most probable speed (most molecules are this fast).  u av is the average speed of the molecules.  u rms, the root-mean-square speed, is the one associated with the average kinetic energy.  Molar mass

17. 10.8 – Molecular Effusion and Diffusion Effusion is the escape of gas molecules through a tiny hole into an evacuated space. Diffusion is the spread of one substance throughout a space or a second substance.

18. Graham’s Law of Effusion Graham’s Law relates the molar mass of two gases to their rate of speed of travel. The “lighter” gas always has a faster rate of speed.

19. 10.9 – Real Gases: Deviations from Ideal Behavior The assumptions made in the kinetic-molecular model (negligible volume of gas molecules themselves, no attractive forces between gas molecules, etc.) break down at high pressure and/or low temperature.

20. Real Gases  In the real world, the behavior of gases only conforms to the ideal-gas equation at relatively high temperature and low pressure. This is because real gases do have volumes and do attract one another.  At high pressures, the volumes of the gas molecules are not negligible, so the gas volume would be slightly greater than that predicted by the Ideal Gas Law.  At low temps, the attractive forces are not negligible, which lessens the forces with which gas molecules hit the wall of their container, giving a lower pressure than would be predicted by the Ideal Gas Law