2. Thermal Energy
A. Temperature & Heat
1. Temperature is the measure of the
average kinetic energy of the
particles in a substance.
3. 2. SI unit for temp. is the Kelvin
(but you will see Celsius used)
a. K = C + 273 (10C = 283K)
b. C = K – 273 (10K = -263C)
3. Thermal Energy –
the total of all the
kinetic and potential
energy of all the
particles in a
substance.
4. Quantity of Heat
• It is measured in joules
• Common usage is a heat unit called the
calorie (the energy needed to raise the
temperature of 1 gram of water by 1o
C)
• Chemists use the kilocalorie (1000
calories=4200 Joules)
• Nutritionists call it a food Calorie
5. 4. Thermal energy relationships
a. As temperature increases, so
does thermal energy (because the
kinetic energy of the particles
increased).
b. Even if the temperature doesn’t
change, the thermal energy in a
more massive substance is higher
(because it is a total measure of
energy).
6. 5. Heat
a. The flow of
thermal energy
from one object to
another.
b. Heat always
flows from warmer
to cooler objects.
Ice gets warmer
while hand gets
cooler
Cup gets cooler while
hand gets warmer
7. Heat Transfer
• Heat flows from hot to cold.
– If you hold something cold, heat flows from hand
to object.
– If you hold something hot, heat flows from object
to hand
• Conduction- transfer of thermal energy
through matter by the direct contact of
particles
– Occurs because particles are in constant motion
– KE transferred as particles collide
9. Conduction• Heating of metal pan-
– Particles in handle of pan move slowly
– Fast moving particles from the bottom bump into
slower particles and speed them up
– Occurs until all particles move the same speed
• Conduction works best in solids- especially
metals- because particles are close together
10. Conduction and Convection
• Metals- good conductors-because electrons move
easily & transfer KE to nearby particles
• Fluid- any materials that flows
• Convection- transfer of energy in a fluid by the
movement of heated particles
• Convection currents transfer heat from warmer to
cooler parts of a fluid.
• Convection vs. Conduction-
– Conduction involves collisions and transfers of energy.
– Convection involves movement of the energetic particles
from one location to another
11.
12. Convection
• Convection- results in changes in density
– As particles move faster, they get farther apart
– Fluid expands as temperature increases
– Larger volume = smaller density
– Decreasing density results in the rise of the warmer
fluid
• Lava Lamp-
– Cool oil = dense = sits on the bottom
– Warmer oil = less dense than alcohol & rises
– As it rises, it loses energy through conduction
• Causes decrease in density = sinking
13. • When oil is
cool
Oil isOil is
warm, so itwarm, so it
risesrises
Oil starts to lose
heat by
conduction and
falls
14. Convection Currents
• Currents in which warm portions of the fluid
move through the substance- convection
• The warm portions transfer energy to the cool
section through conduction
15. Heat Transfer on Earth
• At equator- earth experiences the most heat from the
sun.
– Result: evaporation of water and large accumulations of
clouds.
– As the water vapor rises, it cools and condenses, forming
rain
• After the rain = dry air
– Dry air causes moisture to evaporate, drying out the ground
– causes desert
• Convection currents create deserts and rain forests
over different regions of Earth
16. Radiation
• Transfer of heat to the earth – occurs through
radiation
• Radiation- the transfer of energy by
electromagnetic waves. The waves travel
through space even without matter
17. Controlling the Flow of Heat
• To control the flow of heat: Use clothing,
blankets, layers of fat, fur, etc.
• Insulator- material that does not allow heat to
flow through easily
• Gases – like air- are good insulators because:
– Gas particles are very far apart & can’t transmit E
through conduction.
– If the gas is also held in place, particles can’t move
around and warm up the rest of the gas
18. Insulation
• Insulation is made of fluffy materials
containing pockets of trapped air – prevents
heat loss
• Thermos- vacuum layer between 2 layers of
glass
– Vacuum contains few particles so conduction &
convection don’t occur.
• Thermos- coated in aluminum
– Reflects electromagnetic waves that would either
heat the substance or allow the substance to cool
19.
20. 6. Specific Heat
a. Some things heat up or cool
down faster than others.
Land heats up and cools down faster than water
21. b. Specific heat is the amount of
thermal energy required to raise the
temperature of 1 kg of a substance
one degree (C or K).
1) C water = 4184 J / kg C
2) C sand = 664 J / kg C
This is why land heats up quickly
during the day and cools quickly at
night and why water takes longer.
22. Specific Heat
• The higher the specific heat, the more
energy is required to cause a change in
temperature.
• Substances with higher specific heats must
lose more thermal energy to lower their
temperature than do substances with a low
specific heat.
• Water is slower to heat but is also slower to
lose heat
23. Why does water have such a
high specific heat?
Water molecules form strong bonds
with each other; therefore it takes
more heat energy to break them.
Metals have weak bonds and do not
need as much energy to break them.
water metal
24. c. A calorimeter is
used to help measure
the specific heat of a
substance.
First, mass and
temperature of
water are measured
Then heated
sample is put
inside and heat
flows into water
T is measured
for water to help
get its heat gain
This gives the
heat lost by the
substance
27. Expansion of Water
• This is why lakes and ponds and rivers
freeze with the ice on top
• If they didn’t, no aquatic life would be
possible
28. Thermostat
•A thermostat is a device
that controls the
temperature
•The switch of a thermostat
is a bimetallic strip
29. Bimetallic Strip
• Two different metals that are
bound together
• They expand at different rates
when heated
• Used as a switch in a
thermostat
30. Copyright 1999,
PRENTICE HALL
Chapter 11 30
Phase ChangesPhase Changes
• Sublimation: solid → gas.
• Vaporization: liquid → gas.
• Melting or fusion: solid → liquid.
• Deposition: gas → solid.
• Condensation: gas → liquid.
• Freezing: liquid → solid.
Energy Changes AccompanyingEnergy Changes Accompanying
Phase ChangesPhase Changes
• Energy change of the system for the
above processes are:
31. Copyright 1999,
PRENTICE HALL
Chapter 11 31
Phase ChangesPhase Changes
Energy Changes AccompanyingEnergy Changes Accompanying
Phase ChangesPhase Changes
– Sublimation: (endothermic).
– Vaporization: (endothermic).
– Melting or Fusion: (endothermic).
– Deposition: (exothermic).
– Condensation: (exothermic).
– Freezing: (exothermic).
• Generally heat of fusion (enthalpy of fusion) is
less than heat of vaporization:
– it takes more energy to completely separate
molecules, than partially separate them.
32. Copyright 1999,
PRENTICE HALL
Chapter 11 32
Phase ChangesPhase Changes
Energy Changes AccompanyingEnergy Changes Accompanying
Phase ChangesPhase Changes
• All phase changes are possible under
the right conditions (e.g. water sublimes
when snow disappears without forming
puddles).
• The sequence heat solid → melt → heat
liquid → boil → heat gas is endothermic.
• The sequence cool gas → condense →
cool liquid → freeze → cool solid is
exothermic.
33. Copyright 1999,
PRENTICE HALL
Chapter 11 33
Phase ChangesPhase Changes
Energy Changes AccompanyingEnergy Changes Accompanying
Phase ChangesPhase Changes
34. Copyright 1999,
PRENTICE HALL
Chapter 11 34
Phase ChangesPhase Changes
Heating CurvesHeating Curves
• Plot of temperature change versus heat
added is a heating curve.
• During a phase change, adding heat
causes no temperature change.
–These points are used to
calculate ∆Hfus and ∆Hvap.
• Supercooling: When a liquid is cooled
below its melting point and it still
remains a liquid.
• Achieved by keeping the temperature
low and increasing kinetic energy to
37. Copyright 1999,
PRENTICE HALL
Chapter 11 37
Phase ChangesPhase Changes
Critical Temperature and PressureCritical Temperature and Pressure
• Gases liquefied by increasing pressure
at some temperature.
• Critical temperature: the minimum
temperature for liquefaction of a gas
using pressure.
• Critical pressure: pressure required for
liquefaction.
39. Phase Diagrams
Copyright 1999,
PRENTICE HALL
Chapter 11 39
A phase diagrams show what
phases exist at equilibrium and what
phase transformations we can
expect when we change one of the
parameters of the system (T, P,
composition).
40. Copyright 1999,
PRENTICE HALL
Chapter10 40
Phase DiagramsPhase Diagrams
• Phase diagram: plot of pressure vs.
Temperature summarizing all equilibrium
between phases.
• Given a temperature and pressure,
phase diagrams tell us which phase will
exist.
• Features of a phase diagram:
– Triple point: temperature and pressure at which all three
phases are in equilibrium.
– Vapor-pressure curve: generally as pressure increases,
temperature increases.
– Critical point: critical temperature and pressure for the
gas.
– Melting point curve: as pressure increases, the solid phase
is favored if the solid is more dense than the liquid.
– Normal melting point: melting point at 1 atm.
42. Copyright 1999,
PRENTICE HALL
Chapter 11 42
Phase DiagramsPhase Diagrams
The Phase Diagrams of HThe Phase Diagrams of H22 O and COO and CO22
• Water:
– The melting point curve slopes to the left
because ice is less dense than water.
– Triple point occurs at 0.0098°C and 4.58
mmHg.
– Normal melting (freezing) point is 0°C.
– Normal boiling point is 100°C.
– Critical point is 374°C and 218 atm.
43. Copyright 1999,
PRENTICE HALL
Chapter 11 43
Phase DiagramsPhase Diagrams
The Phase Diagrams of HThe Phase Diagrams of H22 O and COO and CO22
• Carbon Dioxide:
– Triple point
occurs at -56.4°C
and 5.11 atm.
– Normal
sublimation point
is -78.5°C. (At 1
atm CO2 sublimes
it does not melt.)
– Critical point
occurs at 31.1°C
and 73 atm.
