Kinetic Theory of Matter

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# Kinetic Theory of Matter - PowerPoint PPT Presentation

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

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### Kinetic Theory of Matter

Why Johnny can’t sit still (Johnny is a gas particle)

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
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
Kinetic model of gases

Plot of speed vs. # molecules

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
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)
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
Kinetic model of liquids
• Particles cannot move in a straight line
• Particles vibrate along random paths
• Higher temp means more vibration and faster speed
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)
Liquid crystals
• Substances that lose organization in only one dimension as they melt
• Used in electronic displays because their characteristics change with electric charge
Plasmas
• Most like gases
• Composed of ions and subatomic particles at high energy – candle flame, fluorescent lights
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)
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”)

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
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
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?
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)
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
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
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
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
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)
Liquid state to gas state
• Heat = mHv
• Find the heat necessary to boil 230g water.
• Heat = 230gx2260J/g

= 519,800 Joules

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