Gases: Kinetic Molecular Theory, Ideal and Real Gases

1. Chapter 10: Page 300-330 Chapter 11: Page 332-359 Chapter 10 & 11: Gases Chlorine gas was used as a weapon in WWI

2. Kinetic Molecular Theory • all matter is made up of particles (atoms) in random and constant motion. (colliding) • Gases have very low density • particles are spaced far apart. • Gases are compressible. • Extreme pressures-gases will compress until they become liquids (or solids, CO2). • Adding heat to a system • increases the temperature … • Temperature = measure of the average kinetic energy of the particles. • Increasing the pressure of a gas, • increases the density of the gas - the number of particles in a given space.

3. Ideal and Real Gases Ideal gasses • Ideal Gas: • Imaginary, perfect gas – makes calculations easier • Real Gas: • Gas that actually behaves in reality… • When compressed, real gases will form liquids, and even exhibit liquid-like behaviors when still in gas form. • Real gas molecules interact with each other -causing them to travel in non-linear paths and collide “inelastically.” • With real gases, the size of the gas molecules effects their behavior. Real Gasses 

4. Gases in our atmosphere • Nitrogen-78% • Oxygen-21% • Argon-<1% • Trace amounts of CO2, Ne, He, CH4, Kr, H2, O3, and others. • Some gases function as greenhouse gases, and work to hold heat on the earth’s surface. • Some gases function to block harmful UV radiation energy from the sun.

5. The Greenhouse Effect • The sun’s energy travels through space and warms the surface of the earth. • Some of the energy is reflected back into space. • Greenhouse Gases • trap heat that would leave the atmosphere. • H2O, CH4, and CO2 are common greenhouse gases. • “Global Warming” • Theory that increasing levels of Greenhouse gasses is causing the global average temps to increase.

6. The Ozone Layer (O3) Page 778 for more info • Ozone is • a corrosive poisonin the Troposphere(where we live) • frequently created and given off from free electrical ionization. • Ozone • Absorbs harmful Ultraviolet (UV) energy in the stratosphere, 11km (6 miles) above us. • Note: the Ozone layer is less than 1mm thick! • It is always moving, like a cloud, due to weather patterns and climate variations.

7. Pascal’s Principle and Pressure F = Force a = area • French physician, Blaise Pascal, showed that • fluids (including gasses) exert a uniform pressure on all the surfaces that they contact. • Exerting a force on the top surface of a gas, causes that force (pressure) to be exerted on all the walls of its container. • Pressure is due to the particles of a gas striking a surface. We can detect pressure from billions upon billions of gas molecules striking a surface at any point in time.  Which exerts a greater pressure? 

8. Pressure units… • The SI unit of pressure is the Pascal, Pa, equaling one newton per square meter. • Earth’s air pressure at sea level ~ 100,000 Pa. 100kPa • PSI (US) • Pound per square inch. Atmospheric pressure at sea level is about 14.5 PSI. • mmHg (EU, Asia) (AKA: Torr) • Millimeters of mercury. • Atmospheric pressure is 760 mmHg at sea level. • This has to due with the height of a column of liquid mercury raised in a barometer. • inHg • Inches of mercury. Used only in meteorology. • Atmospheric pressure is apx 30inHg.

9. And finally… • And, finally…the atmosphere, atm • the pressure exerted by the atmosphere at sea level, at 00C. (This creates STP…) • Standard Temperature and Pressure: • STP • usually used when referring to reactions with gases. STP is defined as: • 1 atm and 273.15 K • 101 kPa and 273.15 K • 760 mmHg and 273.15 K When doing work with gases, select the STP that matches the pressure you are using. (atm in this class)

10. Charles’ Law Simulation. constant volume • French chemist, Jacque Charles, showed that at constant pressure, • temperature and volume varied proportionally. That is… • V / T=k(k = some constant #) • We tend to write Charles’ Law as the volumes and temperatures under two conditions: c 1780’s

11. Boyle’s Law We’re leaving one law out… can you guess what it is? • A young, adventurous, British aristocrat named Robert Boyle found that • when temperature is kept constant, volume varies inversely proportional with pressure. That is: • P V = k (constant) • We tend to write Boyle’s Law as the volumes and pressures under two conditions: c 1660’s

12. Charles’ Law + Boyle’s Law + Avogadro’s Law=THE IDEAL GAS LAW • R is the “gas constant” and numerically depends upon the pressure units used. Pressure Volume (in Liters) Moles Constant Temperature (in Kelvin)

13. The Gas Constant • The Gas Constant is the numerical bridge between number of moles of a gas, its temperature, and volume or pressure. • R = 8.314 L٠kPa / mol٠K • R = 0.0821 L٠atm / mol٠K • Note that the first constant is in KILO Pascals. When given Pascals, you must first convert to kilopascals. • Our calculations will be done in Atm

14. Dalton’s Law of Partial Pressures • The total pressure in a system is the sum of the individual pressures exerted by each gas. • So, if gas A exerts a pressure of 2 units, and gas B exerts a pressure of 3 units, the total pressure of a system of equal parts of A and B, would be ? • Total= A + B……..2 + 3 = 5 units. • In our atmosphere, Oxygen is about 21%. If we have a sample of air at 1 atm, what is the pressure due to oxygen?

15. Graham’s Law of Gas Effusion • Effusion • motion of a gas through an opening in a container. • notDiffusion - dispersing from higher concentration to lower concentration. • Rates (speeds) of effusion are related to the molar mass of a gas. • The higher the molar mass, the slower the gas will effuse. • This is a property of real gases

16. Graham’s Law of Gas Effusion • At the same temperature… • The higher the molar mass, the slower the gas will effuse. • Graham’s Law of Effusion: Molar mass velocity Molar mass velocity Gas AvsGas B

17. Vapor Pressure Page 324 • All liquids exert a vapor pressure. • Vapor pressure = liquid’s molecules  gas phase. • Higher temperatures  greater molecular speed  greater vapor pressure. • Morevolatile liquids exert a greater vapor pressure than do less volatile liquids. • Can you explain why this is? • In lab: we collect gasses over water. There is a small amount of water vapor in our gas samples, due to water’s vapor pressure.

18. Phase Diagrams Example on page 381 • Phase diagrams • predict if a substance will be a solid, liquid or gas • depends upon the pressure and temperature of the substance. • Triple Point • point where solid, liquid, and gas all exist – for water, 00C. Notice, that as you increase pressure, the boiling point of water increases-this is why a pressure cooker works. What about Denver, the “mile-high city?” End of Gases lecture, Chapters 10,11, problems following

19. In-chapter problems: • Page 327, #5,7,8What is Pressure? • Page 327, #11-14Pressure Units • Page 327, #17-19Pressure Conversions • Page 330, #20-24eBoyle’s Law • Page 330, #25-27Charles’ Law • Page 330, #31-35oCombined Law • Page 331, #39,40Dalton’s Law of Partial Pressures • Page 357, #9-13oAvogadro’s Molar Gasses • Page 358, #17-20Ideal Gas Law • Page 358, #23-29oIdeal Gas Law and Stoichiometry • Page 359, #39-42 Graham’s Law of Gas Effusion End of Gases Unit, Chapters 10,11

20. CCSD Syllabus Objectives • 11.1: Kinetic Molecular Theory • 11.2: Physical Properties of Gasses • 11.3: STP • 11.4: Volume-Temp relationships • 11.5: Volume-Pressure relationships • 11.6: Density-Volume-Pressure-Temperature • 11.10: Ideal Gas Law • 11.11: Graham’s Law • 11.12: Ideal Gas vs Real Gas • 12.3: Evaporation, Condensation, Sublimation • 21.1: Environmental Chemistry