Kinetic Theory of Matter

1. Kinetic Theory of Matter Why Johnny can’t sit still (Johnny is a gas particle)

2. Kinetic model of gases • Ideal gas particles are point masses • Particles travel in a straight line until they run into something – around 100 -1000 m/s • Collisions with walls of container cause pressure • Diffusion – dispersion of a gas by random motion – heavier gases diffuse more slowly

3. Kinetic model of gases • Collisions are perfectly elastic – no other interactions between gas particles – like air hockey pucks • Temperature is related to the average kinetic energy of the gas molecules – higher temp = faster speed

4. Kinetic model of gases Plot of speed vs. # molecules

5. Kinetic model of gases • Brownian motion – Random motion of suspended particles in liquid or gas • Due to collisions between particles and atoms of gas or liquid • Used by Einstein to prove atomic theory of matter

6. Brownian motion

7. Brownian motion animation

8. Properties of gases • Gases can flow • Gases take the shape of the container • Gases have no definite volume • Gases and liquids are fluids (anything that can flow)

9. Kinetic model of liquids • Particles are much closer together than gases • Interparticle interactions are significant • Particles slide past each other like magnetized marbles • Flow • Take shape of container • Have a definite volume

10. Kinetic model of liquids • Particles cannot move in a straight line • Particles vibrate along random paths • Higher temp means more vibration and faster speed

11. Kinetic model of solids • Particles vibrate in place • Higher temp means faster/wider vibrations • Crystalline solids – regular arrangement of particles (salt, diamond) • Amorphous solid – random arrangement (wax, rubber, glass)

12. Liquid crystals • Substances that lose organization in only one dimension as they melt • Used in electronic displays because their characteristics change with electric charge

13. Plasmas • Most like gases • Composed of ions and subatomic particles at high energy – candle flame, fluorescent lights

14. Kinetic energy and temperature • Temperature scales • Celsius – based on melting point (0ºC) and boiling point (100ºC) of water • Kelvin – based on absolute zero (temperature at which all atomic movement ceases)

15. Kinetic energy and temperature • Kelvins are the same size as ºC • Absolute zero is the same as –273ºC • K=C+273 • Find the Kelvin equivalent of room temperature (25ºC) K = 25 + 273 = 298K (no “º”)

16. Kinetic energy and temperature • Kelvins are directly proportional to kinetic energy • Molecules at 400K have twice as much energy as molecules at 200K • Degrees Celsius are not directly proportional to kinetic energy

17. Mass and energy • Kinetic energy depends on mass and speed • At the same temperature, heavier molecules move more slowly • Heavier molecules diffuse more slowly than light ones

18. Mass and energy • Consider the following gases He at 300K Rnat 300K H2 at 100K Br2 at 100K • In which gas are the molecules moving the fastest? • In which gas are particles moving the slowest?

19. Specific heat capacity • Heat it takes to raise the temperature of one gram of stuff 1ºC • Unit is J/gºC; symbol is CP • Metals have low heat capacity • Water has a very high heat capacity (4.184J/gºC, or 1cal/gºC)

20. Specific heat capacity • q = mCPT • Find the heat necessary to raise the temperature of a 5g slug of lead from 22-100ºC. CP for lead = 0.13J/gºC • H = mCPT = 5(0.13)(100-22) = 50.7J

21. Changing state • Gas – liquid • Evaporation – some molecules of a liquid have enough energy to escape – happens at RT • Boiling point – temperature at which the vapor pressure of a liquid equals the atmospheric pressure

22. Liquid state to gas state • Vapor pressure – pressure exerted by molecules trying to leave the surface of a liquid – increases with increasing temperature • Boiling point depends on: • Molar mass - higher MM, higher BP • Polarity – high polarity, high BP • Atmospheric pressure – high AP, high BP

23. Liquid state to gas state • Heat of vaporization – heat necessary to vaporize one gram of a liquid at its boiling point • Hv = 2260 J/g for water • J = Joule • 1 calorie is the heat necessary to raise the temperature of 1g of water 1ºC. 1 cal = 4.184 J

24. Liquid state to gas state • Heat transfer – when a liquid boils or evaporates, heat goes from surroundings to the liquid (sweating) • When a gas condenses, heat is transferred from the gas to the surroundings (steam burns)

25. Liquid state to gas state • Heat = mHv • Find the heat necessary to boil 230g water. • Heat = 230gx2260J/g = 519,800 Joules

26. Solid state to liquid state • Melting – molecules get enough energy to acquire linear motion • Freezing – molecules slow down enough so they get trapped in place • Heat of fusion – heat released when one gram of a substance freezes – Hf = 334J/g for water

27. Solid state to liquid state • Math is the same as for boiling • Find the heat released when 10.0g water freezes to form ice. • q = Hfxm = 10.0gx334J/g = 3340J • Heat transfer happens without temperature changes during phase change

28. Heating curves

29. Sublimation • Solid – gas – sublimation – happens when pressure is low • Dry ice and iodine sublime readily at standard atmospheric pressure • Below freezing, ice will sublime slowly • Many substances can be made to sublime under a vacuum

30. Sublimation • Sublimation involves heat transfer from the surroundings to the substance • Opposite process is deposition (heat goes from substance to surroundings)