1. Gas Density: Summary The molar concentrations and densities of gases increase as they are compressed (less volume, right?), but decrease as they are heated (volume increases, right?). The density of a gas depends on its molar mass.

2. The Stoichiometry of Reacting Gases Many reactions occur in the gas phase and we can use the ideal gas law to determine the volume of gas produced or consumed in a chemical reaction –How much oxygen will it take to saturate the hemoglobin molecules in a red blood cell?

3. Steps to working with stoichiometry in the gas phase: 1.Balance the chemical equation 2.Calculate the number of moles of reactant consumed 3.Use the stoichiometric coefficients from the chemical reaction to relate the # moles of product made to the # of moles of reactant consumed.

4. Mixtures of Gases Most gases we encounter and use every day are actually mixtures –The atmosphere of the earth –The breath we exhale If the gases in a mixture do not react with each other, we may consider the mixture to be a single, pure gas for the sake of computation

5. Mixtures and Partial Pressures Dalton came up with the law that allows us to calculate the pressure of a mixture and the contribution of the individual gases that comprise it How did he arrive at this conclusion? He determined that if he combined the gases, the pressure of the mixture would be the sum of the Partial Pressures of the individual gases. And it is.

6. Daltons Law of Partial Pressures The total pressure of a mixture of gases is the sum of the partial pressures of its components

7. Mole Fractions The best way to explain/understand the relationship between total pressure and partial pressures is to look at the mole fractions of each gas in a mixture For a mixture of gases with components A, B and C, the mole fraction (x A ) is:

8. Mole Fractions We know that x A + x B + x C = 1 Each gas exerts a pressure that is the mole fraction of the gas times the total pressure in the vessel P A = x A P

9. Molecular Motion We have derived the gas laws and worked with an equation of state: The Ideal Gas Law Now, we need to look at the motion of the gas molecules themselves

10. Diffusion and Effusion Diffusion Effusion

11. Grahams Law of Effusion Scottish chemist Thomas Graham studied the effusion of gases He found that at constant temperature, the rate of effusion of a gas is inversely proportional to the square root of its Molar Mass. Does this make sense to you?

12. Grahams Law of Effusion The average speed of the molecules will be inversely proportional to the molar mass –Bigger molecules move slower at constant temperature than smaller ones. –Think of it in terms of energy and getting the molecules to move

13. Grahams Law of Effusion We can use this relationship to identify compounds/molecules with unknown molar mass by comparing the Rate of Effusion of an unknown gas to that of another with known Molar Mass Remember: rate is in units of m/s and time is just seconds, so we flip the relationship

14. Effusion Rate and Temperature If you raise the temperature, what happens to the Effusion rate? … This shows us that the average speed of the gas molecules is directly dependent on temperature This tells us that The average speed of the molecules in a gas is proportional to the square root of the temperature This is BIG!!!

15. Average Speed and Temperature We can combine the observations on the relationship between the average speed of a gas to the temperature and its molar mass The average speed of the molecules in a gas is directly proportional to the square root of the temperature and inversely proportional to the square root of the molar mass

16. Kinetic Molecular Theory All of this leads to the theory describing the behaviour of gas molecules: KMT 1.A gas consists of a collection of molecules in continuous, random motion 2.Gas molecules are infinitesimally small points 3.The molecules move in a straight line until they collide 4.The molecules do not influence each other until they collide (No attractive forces b/w molecules)

17. Kinetic Molecular Theory Root mean square Why use v rms ? E k =1/2(mv rms 2 )