1. How many joules of energy are given off when 5.0 g of water cool from 75.°C to 25.°C?
2. How many calories are given off in the above problem?
3. How many joules does it take to melt 35 g of ice at 0°C? (ΔHfus = 333 J/g)
4. How many calories are given off when 85 g of steam condense to liquid water? (ΔHvap =
539.4 cal/g)
5. How many joules of heat are necessary to raise the temperature of 25 g of water from
10.°C to 60.°C?
6. How many calories are given off when 50. g of water at 0°C freezes? (ΔHfus = 79.72
cal/g)
7. Calculate the change in heat if an empty aluminum can cools from 36°C to 22°C. The
can has a mass of 15.00 g and a specific heat of 0.897 J/g°C.
8. How much heat is required to change 54.0 g of H2O from liquid water at 46.°C to steam
at 110.°C.
ΔHvap = 2.26 kJ/g
Csteam = 2.042 J/g°C
9. Calculate the amount of energy required to raise 25.0 g of water from
-10.°C to 115.°C.
ΔHvap = 2.26 kJ/g
Csteam = 2.042 J/g°C
ΔHfus = 0.333 kJ/g
Cice = 2.08 J/g°C
10. Ethanol melts at -114°C and boils at 78°C.
ΔHvap = 0.837 kJ/g
Cliquid = 2.3 J/g°C
ΔHfus = 10.9 kJ/g
Csolid = 0.97 J/g°C
How much heat is required to convert 75.0 g of ethanol at -120°C to the vapor phase at
78°C?
11. Your air conditioner may use chlorofluorocarbon, CCl2F2, as the heat transfer fluid. Its
normal boiling point is -30.°C, and the latent heat of vaporization is 165 J/g. The gas and
liquid have specific heat capacities of 0.61 J/g°C and 0.97 J/g°C, respectively. How
much heat must be evolved when 10.0 g of CCl2F2 is cooled from +40.°C to -40.°C?
12. The heat required to melt 1.00 g of ice at 0°C is 333 J. If one ice cube has a mass of
62.0 g, and a tray contains 16 ice cubes, what quantity of energy is required to melt a tray
of ice cubes at 0°C?
13. Chloromethane, CH3Cl, is used as a topical anesthetic. The temperature at which CH3Cl
liquid turns to vapor is -24.09°C. What quantity of heat must be absorbed by the liquid to
convert 150. g of liquid to vapor at -24.09°C? The heat of vaporization of CH3Cl is 0.432
kJ/g.
14. Mercury, with a freezing point of -39°C, is the only metal that is liquid at room
temperature. How much heat energy (in joules) must be released by mercury if 1.00 mL
of the metal is cooled from room temperature (23.0°C) to -39.°C, and then crystallized to
a solid? The density of mercury is 13.6 g/cm3. Its specific heat is 0.140 J/g°C, and its
heat of fusion is 11.4 kJ/g.
15. How much heat energy (in joules) is required to raise the temperature of 1.00 lb of tin
from room temperature (25.0°C) to its melting point, 231.9°C, and then melt the tin at that
temperature? The specific heat of tin is 0.227 J/g°C, and the metal requires 59.2 J/g to
convert the solid to a liquid.
16. The heat of fusion of ice is 333 J/g. Suppose that three 45 g ice cubes at 0°C are
dropped into 500. mL of tea initially at 20.0°C. How much of the ice melts? Assume the
tea is weak enough that its specific heat is the same as pure water.
17. For many years drinking water has been cooled in hot climates by evaporating it from the
surfaces of canvas bags or porous clay pots. How many grams of water can be cooled
from 35°C to 22°C by the evaporation of 50 g of water? ΔHvap = 2.4 kJ/g.
18. As mentioned earlier, chlorofluorocarbon, CCl2F2 is widely used as a refrigerant, but is
now being replaced by compounds that are believed to be less harmful to the
environment. What mass of this substance must evaporate in order to freeze 100. g of
water initially at 18°C?
ΔHvap, chloro = 289 J/g. ΔHfus, water = 334 J/g