11.4-11.5
1. Define phase change/change of states
2. Draw a diagram that represents the change in energy for phase changes.
3. Define heat of fusion
4. Define fusion
5. Define sublimation, deposition, melting, freezing, evaporation, and condensation.
6. Define heat of vaporization
7. Define vapor pressure
8. Define heat of freezing and heat of deposition
9. Will the temperature of a system change during a phase change? (Example ice changing to water)
10. Draw a heat curve and label the parts.
11. Calculate the enthalpy change upon converting 1.00 mol of ice at -25 °C to steam at 125 °C under a constant
pressure of 1 atm. The specific heats of ice, liquid water, and steam are 2.03, 4.18, and 1.84 J/g-K,
respectively. For H2O, ΔHfus = 6.01 kJ/mol and ΔHvap = 40.67 kJ/mol.
12. Define critical pressure and critical temperature.
13. What is the relationship between intermolecular and critical temp and pressure?
14. Define vapor pressure
15. Draw and label the vapor pressure over a liquid
16. Define dynamic equilibrium
17. What is the vapor pressure of a liquid?
18. Define volatile.
19. Define boiling point.
20. Define normal boiling point
21. Define and describe the Clausius-Clapeyron Equation.
22. Use Figure 11.25 to estimate the boiling point of diethyl ether under an external pressure of 0.80 atm.
23. Name the phase transition in each of the following situations and indicate whether it is exothermic or
endothermic: (a) When ice is heated, it turns to water. (b) Wet clothes dry on a warm summer day. (c) Frost
appears on a window on a cold winter day. (d) Droplets of water appear on a cold glass of beer.
24. Explain why any substance’s heat of fusion is generally lower than its heat of vaporization.
25. 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 to 20 °C by the evaporation of
60 g of water? (The heat of vaporization of water in this temperature range is 2.4 kJ/g. The specific heat of
water is 4.18 J/gK.)
26. Ethanol (C2H5OH) melts at -114 °C and boils at 78 °C. The enthalpy of fusion of ethanol is 5.02 kJ/mol, and
its enthalpy of vaporization is 38.56 kJ/mol. The specific heats of solid and liquid ethanol are 0.97 and 2.3
J/gK, respectively. (a) How much heat is required to convert 42.0 g of ethanol at 35 °C to the vapor phase at
78 °C? (b) How much heat is required to convert the same amount of ethanol at -155 °C to the vapor phase
at 78 °C?
27. (a) What is the significance of the critical pressure of a substance? (b) What happens to the critical
temperature of a series of compounds as the force of attraction between molecules increases? (c) Which of
the substances listed in Table 11.6 can be liquefied at the temperature of liquid nitrogen (-196 °C)?
28. Which of the following affects the vapor pressure of a liquid? (a) Volume of the liquid, (b) surface area, (c)
intermolecular attractive forces, (d) temperature, (e) density of the liquid.
29. (a) Place the following substances in order of increasing volatility: CH4, CBr4, CH2Cl2, CH3Cl, CHBr3, and
CH2Br2. (b) How do the boiling points vary through this series? (c) Explain your answer to part (b) in terms
of intermolecular forces.
30. (a) Two pans of water are on different burners of a stove. One pan of water is boiling vigorously, while the
other is boiling gently. What can be said about the temperature of the water in the two pans? (b) A large
container of water and a small one are at the same temperature. What can be said about the relative vapor
pressures of the water in the two containers?
31. Using the vapor-pressure curves in Figure 11.25, (a) estimate the boiling point of ethanol at an external
pressure of 200 torr, (b) estimate the external pressure at which ethanol will boil at 60 °C, (c) estimate the
boiling point of diethyl ether at 400 torr, (d) estimate the external pressure at which diethyl ether will boil at
40 °C.