States of Matter

1. States of Matter Chapter 13

2. What You Need to Master • How to use the kinetic-molecular theory to explain the physical properties of gasses, liquids, and solids • Compare types of intermolecular forces • Explain how kinetic energy and intermolecular forces combine to determine the state of a substance • Describe the role of energy in phase changes

3. Gases Chapter 13.1

4. Kinetic-Molecular Theory • Jan Baptista Van Helmont (Flemish) used the word Chaos (without order) to describe products of reactions that had no fixed shape or volume. • By 18th century, scientists knew how to collect gasses by displacing water • They could now measure properties of gasses • 1860 Ludwig Boltzmann and James Maxwell proposed models to explain the properties of gasses • Kinetic-molecular theory

5. Kinetic Molecular Theory Assumptions • Particles are so small compared to distances between the particles that the volume of the particles can be assumed to be negligible (zero). • Not totally true • The particles are in constant motion. Collisions of the particles with the walls of the container are the cause of pressure exerted by the gas. • Basically true • The particles are assumed to exert no forces on each other; they are assumed to neither attract nor repel each other. • Not totally true • The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temperature of the gas. • True

6. Kinetic-Molecular Theory • Kinetic – Greek word meaning “to move” • any object in motion has “kinetic energy” • Kinetic-Molecular Theory describes the behavior of gasses in terms of particles in motion. • Model makes several assumptions about size, motion, and energy of particles

7. Kinetic-Molecular Theory • Assumptions: • Particle Size : particles are so small and so far apart that there no significant forces of attraction • Particle motion: • particles are in constant, random motion. • Particles move in a straight line until they collide with another particle or the container wall • Particle collisions are elastic (no energy is lost, but may be transferred)

8. Kinetic-Molecular Theory • Assumptions: • Particle Energy: 2 factors, mass & velocity • KE = ½ mv2 • All particles do NOT have the same energy • Temperature is a measure of the average kinetic energy of the particles

9. How K-M Theory Explains Gas Behavior • Low Density • Density is mass/volume • Cl2 density at 20°C is 2.95 x 10-3 g/mL • Au density is 19.3 g/mL • Because Chlorine is a gas, K-M theory says there must be a lot of empty space between molecules. • Thus, fewer molecules in the same volume

10. How K-M Theory Explains Gas Behavior • Compression and Expansion • Because there is a lot of empty space between molecules/atoms, you can compress or squeeze together the atoms/molecules • Removes some of the empty space • Allowing the atoms to return to normal is expansion

11. How K-M Theory Explains Gas Behavior • Diffusion and Effusion • Diffusion – describes the movement of one material through another • Effusion – gas escaping through a tiny opening • Like puncturing a balloon or tire • Graham’s Law of Diffusion • Rate of effusion ͒1/Ã molar mass • ͒means ‘is proportional to’ • Larger particles diffuse slower than smaller particles

12. How K-M Theory Explains Gas Behavior • Rate Relationships: • Ratea • Rateb Molar massa Molar massb à =

13. Gas Pressure • Pressure = force/area • An elephant has less pressure on its foot than a human. Why? • Gasses exert pressure when they collide with walls of the container. • Small mass means small pressure • But there are about 1022 particles in one liter • Pressure can be substantial • Earth’s air pressure varies – gets lower with altitude

14. Measuring Air Pressure • Evangelista Torricelli (Italian – before 1650) first demonstrated air pressure • Designed equipment to measure air pressure using mercury in a tube. • Now called a barometer

15. Measuring Air Pressure • A ‘manometer’ measures air pressure in a u-shaped tube, in a closed container.

16. Units of Pressure • SI unit is ‘Pascal’ • 1 Pascal = 1 newton/m2 • PSI (pounds per square inch) is still used • Barometers and manometers still report millimeters of mercury (mm Hg) • At sea level, air pressure is 760 mm Hg at 0°C • Air pressure is frequently reported as 1 atmosphere (atm) = 760 mm Hg or 760 torr or 101.3 kilopascals

17. Units of Pressure • The units 1 atm, 760 mm Hg, 760 torr, and 101.3 kilopascals are considered to be defined units • They have as many significant digits as necessary

18. Dalton’s Law of Partial Pressure • Dalton studied mixtures of gasses: • Found that each gas exerts pressure independently. • Law of Partial Pressures: • The total pressure is equal to the sum of pressures by each gas in the mixture • Partial pressure depends on the moles of gas, size of container and temperature. • Does NOT depend on the identity of the gas

19. Dalton’s Law of Partial Pressure • At a given temperature and pressure, • PTOTAL = PA + PB + PC + PD + …. PN • Dalton’s Law can be used for: • Determine amount of gas produced by a reaction • Partial pressures of gasses at same temperature are related to concentrations. • Look up values in reference tables • At 20° C, partial pressure of water is 2.3 kPa