1. Section 10.2 - Thermal Energy Transfer
Chapter 10

2. Specific Heat Capacity
the amount of heat required to raise the temperature of 1g of a substance by 1°C
formula:
Q = mcΔT
Q = amount of heat in joules (J)
m = mass (g)
c = specific heat capacity (J/g°C)
ΔT = change in temperature (°C)

3. Examples
How much energy is released when 35g of water cools from 25°C to 10°C?
If 2.09x103J of energy raised the temperature of moist air from 23°C to 37°C, what is the mass of the air?
**Table 101 on page 375**

4. Specific Heat Capacities

5. Significant Digits All numbers between 1 and 9 are significant
Zeros between numbers 1 and 9 are significant
Zeros to the right of the decimal point are significant but zeros to the left of the decimal point are not significant.
In scientific notation, all digits before the power of ten are significant.

6. Examples:
415 cm
0.46 g/cm3
2.089 L km
5.000 x 104 s

7. Assignment
How much heat is absorbed by 1000 kg of seawater in a large hole on the beach as the seawater’s temperature rises from 20°C to 25°C?
What is the specific heat capacity of a substance if it requires 2334 J of energy to change the temperature of 40g of the substance by 15°C?

8. Conduction, Convection and Radiation
mechanisms that help transfer energy between objects
CONDUCTION : the transfer of heat by contact; molecules run into each other transferring energy from hot to cold molecules
CONVECTION : the transfer of heat by air (or liquid) movement; hot air rises; cold air sinks – causes “convection currents”
RADIATION : transfer of heat by absorbing infrared radiation coming off the Earth: absorbed by the greenhouse gases: water vapour, carbon dioxide, methane, etc.

9. Conduction – energy transfer by contact
Convection – energy transfer by air movement
Radiation – transfer of heat by absorbing infrared energy coming off the earth

11. Phase Changes
water and other substances can change states with no change in temperature
i.e. water going from a liquid to a solid and vice versa happens at 0C
Graph of a phase change......
- a physical change from one state to another

13. Heat of Fusion (Hfus)
amount of energy needed to melt one mole of a substance
Q = nHfus
Q = amount of energy (J)
n = number of moles (mol)
Hfus = heat of fusion (J/mol)
SOLID TO LIQUID or LIQUID TO SOLID

14. Latent Heats of Fusion and Vaporization

15. Examples
How much thermal energy is required to melt mol of ice at 0°C to liquid water at 0°C?
A container holding exactly 360.4g of ice at 0°C melts to form liquid water at 0°C. How much energy must be added to convert the ice to water?

16. Heat of Vaporization (Hvap)
amount of energy required to convert one mole of a substance from a liquid to a gas
Q = nHvap
Q = amount of energy (kJ)
n = number of moles (mol)
Hvap = heat of vaporization (kJ/mol)
LIQUID TO GAS or GAS TO LIQUID
Table 10.2 page 381

17. Examples
How much energy is required to boil 48.8mol of liquid water at 100°C to steam at 100°C?
A beaker contains exactly g of water at 100°C. How much energy is required to convert the water to water vapour at 100°C?

18. Water Influences Climate
water has a low albedo and does absorb about 93% of solar radiation
temperature of large bodies of water stays relatively constant in comparison to air and land

19. Reasons Water has a high specific heat capacity
large amounts of energy are needed to change the temperature
have a moderating effect on the air of nearby land communities
Water has a high heat of vaporization
needs a lot of energy input in order to turn into a gas

20. Water has a high heat of fusion
water must lose a large amount of energy to turn into a solid
when water cools below 4°C it begins to expand allowing ice to float on the surface of water

21. Hydrologic Cycle