1. AP Chemistry Zumdahl Notes, 9 th ed. A Brief Collection of notes, Chapter 5 Feel free to open these files and annotate as you feel the need…this is for your success.

2. Pressure What is pressure? The molecules hitting the inside walls of whatever contains them; this is directly proportional to temperature, which is nothing more than a measure of the kinetic energy of the particles of whatever gas is being measured Barometer: device to measure barometric (atmospheric) pressure, invented by Torricelli, 760mmHg, at sea level Barometers are manometers, devices to measure differences in pressure Units 1 atm = 760 mmHg = 760 torr = 101.325 kPa (kilopascals) Yes, you will need to keep conversions in mind.

3. The gas laws of Boyle, Charles, Avogadro Boyle’s Law Pressure and volume are inversely proportional: PV = k, at constant temp. Reality check: a gas that strictly follows this is called an ideal gas This means that there are no intermolecular attractions, nor loss of energy during collisions of particles of said gas Second reality check: at pressures near 1 atm, our calculations work very well for real world gases Form of Boyle’s law that is used for calculations: P 1 V 1 = P 2 V 2 Rearrange as needed to solve for unknown

4. The gas laws of Boyle, Charles, Avogadro Charles’ Law Volume and temperature are directly proportional, at constant pressure Reality check: must use kelvin temperatures…need to always be positive (K=°C + 273) V = bT Second reality check: we have never actually reached absolute zero…very, very close but not quite there Form of Charles’ law that is used for calculations:V 1 /T 1 = V 2 /T 2 Rearrange as needed to solve for unknown

5. Gay-Lussac’s Law and Combined gas law…not in Zumdahl Another relationship: at constant volume, pressure is directly proportional to temperature in Kelvins P/T = C This becomes another part of combined gas law: P 1 V 1 /T 1 = P 2 V 2 /T 2 I wanted to include these, as you may run into them, even though the Zumdahls chose to leave them out of the 9 th edition

6. The gas laws of Boyle, Charles, Avogadro Avogadro’s Law The amount of a gas present is directly proportional to the volume it occupies, at constant temperature and pressure V/n = C Works best at low pressures…as pressure increases, intermolecular attractions, etc., start to affect real world gases…think liquification….

7. Ideal Gas Law If we combine Boyle’s, Charles’, and Avogadro’s, we can arrive at the mathematical formula: V = R (Tn/P), where R is a proportionality constant (universal gas constant, relating the units together) Rearranging, we get to the form we use: PV = nRT Normal reality check…we can solve for one unknown by algebraically rearranging this equation. This is one of those times where I remind you of the importance of labelling with units!

8. Gas Stoichiometry At STP, (standard temperature and pressure, O°C, 1 atm), one mole of an ideal gas occupies a volume of 22.42 L. Yes, you will use this a lot for these kinds of problems You will need to substitute this in some calculations which use one of the gas laws Another fun calculation/formula Molar mass = dRT/P, where d = density in g/L Yes, it is likely you will need to mess with this formula Yes, you can derive it if you need to as the units will guide you to data location

9. Dalton’s Law of Partial Pressures “For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.” Awesomely summarized…means that each one adds its part to the final pressure Formula: P total = P 1 + P 2 + P 3 + … Each gas, therefore, exerts what is referred to as its “partial pressure” Reality check: you can substitute “n 1 RT/V” as needed for each gas so as to calculate each gas’ contribution; this is dependent upon what information you are given to solve a problem with, as usual Mole fraction: the ratio of one component to the total (see page 210-211) X 1 = n 1 /n total

10. Kinetic Molecular Theory of Gases An attempt to explain the properties of ideal gases The particles are so small compared with the distances separating them that we consider the individual volumes to be zero The particles are in continuous motion and their collisions with the walls of their container create the pressure exerted by same particles The particles have no interactions with each other, as far as attraction/repulsion…consider similar to billiard balls The average KE of a collection is considered to be in direct proportion to the Kelvin temperature of the gas REALITY check: we know that real gases violate these, right? But…these do help us explain much of gas behavior…it allows us to develop the relationships we call the gas laws, which we have been discussing

11. Kinetic Molecular Theory of Gases, continued Important expression: KE avg = 3/2 RT This allows for the calculation of the energy of the gas at a given temperature Also allows for velocity calculations of gas particles…be thoughtful about units…need to be working in joules for energy

12. Effusion and Diffusion Effusion: movement into evacuated chamber Diffusion: movement into “occupied” chamber Graham’s Law of Effusion: Rate of effusion of gas 1 = (mass of gas 2) 1/2 Rate of effusion of gas 2 (mass of gas 1) ½ This shows the inversely proportional relationship between relative rates and the square roots of the masses of the gases involved Diffusion rates: yes, can be calculated similarly, but found experimentally to be less…showing a much more complicated analysis is necessary and beyond our needs to perform

13. Real gases Real gases behave most like ideal gases at low pressures and high temperatures…why?

14. Chemistry in the atmosphere Components of atmosphere: mostly nitrogen and oxygen Pollutants: sources of acid rain! Nitrogen oxides Sulfur oxides