Chapter 5

Imagine a book that is falling from a shelf. At a particular moment during its fall, the book has a kinetic energy of \(13 \mathrm{~J}\) and a potential energy with respect to the floor of \(72 \mathrm{~J}\). How does the book's kinetic energy and its potential energy change as it continues to fall? What is its total kinetic energy at the instant just before it strikes the floor? [Section 5.1]

Imagine that you are climbing a mountain. (a) Is the distance you travel to the top a state function? Why or why not? (b) Is the change in elevation between your base camp and the peak a state function? Why or why not? [Section 5.2]

Which will release more heat as it cools from \(50^{\circ} \mathrm{C}\) to \(25^{\circ} \mathrm{C}, 1 \mathrm{~kg}\) of water or \(1 \mathrm{~kg}\) of aluminum? How do you know? [Section 5.5]

Does \(\Delta H_{\mathrm{rxn}}\) for the reaction represented by the following equation equal the standard enthalpy of formation for \(\mathrm{CH}_{3} \mathrm{OH}(l) ?\) Why or why not? [Section 5.7] $$ \mathrm{C}(\text { graphite })+4 \mathrm{H}(\mathrm{g})+\mathrm{O}(\mathrm{g}) \longrightarrow \mathrm{CH}_{3} \mathrm{OH}(l) $$

In what two ways can an object possess energy? How do these two ways differ from one another?

Suppose you toss a tennis ball upward. (a) Does the kinetic energy of the ball increase or decrease as it moves higher? (b) What happens to the potential energy of the ball as it moves higher? (c) If the same amount of energy were imparted to a ball the same size as a tennis ball, but of twice the mass, how high would it go in comparison to the tennis ball? Explain your answers.

(a) Calculate the kinetic energy in joules of a 45-g golf ball moving at \(61 \mathrm{~m} / \mathrm{s}\). (b) Convert this energy to calories. (c) What happens to this energy when the ball lands in a sand trap?

A watt is a measure of power (the rate of energy change) equal to \(1 \mathrm{~J} / \mathrm{s}\). (a) Calculate the number of joules in a kilowatt- hour. (b) An adult person radiates heat to the surroundings at about the same rate as a 100 -watt electric incandescent lightbulb. What is the total amount of energy in kcal radiated to the surroundings by an adult in 24 hours?

(a) What is meant by the term system in thermodynamics? (b) What is a closed system?

In a thermodynamic study a scientist focuses on the properties of a solution in an apparatus as illustrated. A solution is continuously flowing into the apparatus at the top and out at the bottom, such that the amount of solution in the apparatus is constant with time. (a) Is the solution in the apparatus a closed system, open system, or isolated system? Explain your choice. (b) If it is not a closed system, what could be done to make it a closed system?

(a) What is work? (b) How do we determine the amount of work done, given the force associated with the work?

(a) What is heat? (b) Under what conditions is heat transferred from one object to another?

Identify the force present, and explain whether work is done when (a) a positively charged particle moves in a circle at a fixed distance from a negatively charged particle; (b) an iron nail is pulled off a magnet.

(a) State the first law of thermodynamics. (b) What is meant by the internal energy of a system? (c) By what means can the internal energy of a closed system increase?

(a) Write an equation that expresses the first law of thermodynamics in terms of heat and work. (b) Under what conditions will the quantities \(q\) and \(w\) be negative numbers?

Calculate \(\Delta E\), and determine whether the process is endothermic or exothermic for the following cases: (a) A system absorbs \(105 \mathrm{~kJ}\) of heat from its surroundings while doing \(29 \mathrm{~kJ}\) of work on the surroundings; (b) \(q=1.50 \mathrm{~kJ}\) and \(w=-657 \mathrm{~J} ;\) (c) the system releases \(57.5 \mathrm{~kJ}\) of heat while doing \(22.5 \mathrm{~kJ}\) of work on the surroundings.

(a) What is meant by the term state function? (b) Give an example of a quantity that is a state function and one that is not. (c) Is work a state function? Why or why not?

Indicate which of the following is independent of the path by which a change occurs: \((a)\) the change in potential energy when a book is transferred from table to shelf, (b) the heat evolved when a cube of sugar is oxidized to \(\mathrm{CO}_{2}(\mathrm{~g})\) and \(\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\), (c) the work accomplished in burning a gallon of gasoline.

(a) Why is the change in enthalpy usually easier to measure than the change in internal energy? (b) For a given process at constant pressure, \(\Delta H\) is negative. Is the process endothermic or exothermic?

(a) Under what condition will the enthalpy change of a process equal the amount of heat transferred into or out of the system? (b) During a constant- pressure process the system absorbs heat from the surroundings. Does the enthalpy of the system increase or decrease during the process?

You are given \(\Delta H\) for a process that occurs at constant pressure. What additional information do you need to determine \(\Delta E\) for the process?

Suppose that the gas-phase reaction \(2 \mathrm{NO}(g)+\mathrm{O}_{2}(g)\) \(\longrightarrow 2 \mathrm{NO}_{2}(g)\) were carried out in a constant-volume container at constant temperature. Would the measured heat change represent \(\Delta H\) or \(\Delta E ?\) If there is a difference, which quantity is larger for this reaction? Explain.

The complete combustion of acetic acid, \(\mathrm{CH}_{3} \mathrm{COOH}(l)\), to form \(\mathrm{H}_{2} \mathrm{O}(l)\) and \(\mathrm{CO}_{2}(g)\) at constant pressure releases \(871.7 \mathrm{~kJ}\) of heat per mole of \(\mathrm{CH}_{3} \mathrm{COOH}\). (a) Write a balanced thermochemical equation for this reaction. (b) Draw an enthalpy diagram for the reaction.

The decomposition of zinc carbonate, \(\mathrm{ZnCO}_{3}(\mathrm{~s})\), into zinc oxide, \(\mathrm{ZnO}(\mathrm{s})\), and \(\mathrm{CO}_{2}(g)\) at constant pressure requires the addition of \(71.5 \mathrm{~kJ}\) of heat per mole of \(\mathrm{ZnCO}_{3}\) (a) Write a balanced thermochemical equation for the reaction. (b) Draw an enthalpy diagram for the reaction.

Consider the following reaction, which occurs at room temperature and pressure: $$ 2 \mathrm{Cl}(g) \longrightarrow \mathrm{Cl}_{2}(g) \quad \Delta H=-243.4 \mathrm{~kJ} $$ Which has the higher enthalpy under these conditions, \(2 \mathrm{Cl}(g)\) or \(\mathrm{Cl}_{2}(g) ?\)

Without referring to tables, predict which of the following has the higher enthalpy in each case: (a) \(1 \mathrm{~mol} \mathrm{CO}_{2}(\mathrm{~s})\) or \(1 \mathrm{~mol} \mathrm{CO}_{2}(g)\) at the same temperature, (b) \(2 \mathrm{~mol}\) of hydrogen atoms or \(1 \mathrm{~mol}\) of \(\mathrm{H}_{2}\) (c) \(1 \mathrm{~mol} \mathrm{H}_{2}(\mathrm{~g})\) and \(0.5 \mathrm{~mol}\) \(\mathrm{O}_{2}(g)\) at \(25^{\circ} \mathrm{C}\) or \(1 \mathrm{~mol} \mathrm{H}_{2} \mathrm{O}(g)\) at \(25^{\circ} \mathrm{C}\), (d) \(1 \mathrm{~mol} \mathrm{~N}_{\mathbf{2}}(g)\) at \(100^{\circ} \mathrm{C}\) or \(1 \mathrm{~mol} \mathrm{~N}_{2}(g)\) at \(300^{\circ} \mathrm{C}\).

Consider the following reaction: $$ 2 \mathrm{Mg}(s)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{MgO}(s) \quad \Delta H=-1204 \mathrm{~kJ} $$ (a) Is this reaction exothermic or endothermic? (b) Calculate the amount of heat transferred when \(2.4 \mathrm{~g}\) of \(\mathrm{Mg}(\mathrm{s})\) reacts at constant pressure. (c) How many grams of \(\mathrm{MgO}\) are produced during an enthalpy change of \(-96.0 \mathrm{~kJ} ?\) (d) How many kilojoules of heat are absorbed when \(7.50 \mathrm{~g}\) of \(\mathrm{MgO}(s)\) is decomposed into \(\mathrm{Mg}(\mathrm{s})\) and \(\mathrm{O}_{2}(g)\) at constant pressure?

Consider the following reaction: $$ \mathrm{CH}_{3} \mathrm{OH}(g) \longrightarrow \mathrm{CO}(g)+2 \mathrm{H}_{2}(g) \quad \Delta H=+90.7 \mathrm{~kJ} $$ (a) Is heat absorbed or released in the course of this reaction? (b) Calculate the amount of heat transferred when \(45.0 \mathrm{~g}\) of \(\mathrm{CH}_{3} \mathrm{OH}(g)\) is decomposed by this reaction at constant pressure. (c) For a given sample of \(\mathrm{CH}_{3} \mathrm{OH}\), the enthalpy change on reaction is \(25.8 \mathrm{~kJ}\). How many grams of hydrogen gas are produced? What is the value of \(\Delta H\) for the reverse of the previous reaction? (d) How many kilojoules of heat are released when \(50.9 \mathrm{~g}\) of \(\mathrm{CO}(g)\) reacts completely with \(\mathrm{H}_{2}(g)\) to form \(\mathrm{CH}_{3} \mathrm{OH}(g)\) at constant pressure?

When solutions containing silver ions and chloride ions are mixed, silver chloride precipitates: $$ \mathrm{Ag}^{+}(a q)+\mathrm{Cl}^{-}(a q)-\longrightarrow \mathrm{AgCl}(s) \quad \Delta H=-65.5 \mathrm{~kJ} $$ (a) Calculate \(\Delta H\) for production of \(0.200 \mathrm{~mol}\) of \(\mathrm{AgCl}\) by this reaction. (b) Calculate \(\Delta H\) for the production of \(2.50 \mathrm{~g}\) of AgCl. (c) Calculate \(\Delta H\) when \(0.150 \mathrm{mmol}\) of AgCl dissolves in water.

At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat \(\mathrm{KClO}_{3}\) : \(2 \mathrm{KClO}_{3}(s)-\longrightarrow 2 \mathrm{KCl}(s)+3 \mathrm{O}_{2}(g) \quad \Delta H=-89.4 \mathrm{~kJ}\) For this reaction, calculate \(\Delta H\) for the formation of (a) \(0.632 \mathrm{~mol}\) of \(\mathrm{O}_{2}\) and (b) \(8.57 \mathrm{~g}\) of \(\mathrm{KCl}\). (c) The decomposition of \(\mathrm{KClO}_{3}\) proceeds spontaneously when it is heated. Do you think that the reverse reaction, the formation of \(\mathrm{KClO}_{3}\) from \(\mathrm{KCl}\) and \(\mathrm{O}_{2}\), is likely to be feasible under ordinary conditions? Explain your answer.

Consider the combustion of liquid methanol, \(\mathrm{CH}_{3} \mathrm{OH}(l)\) : $$ \begin{array}{r} \mathrm{CH}_{3} \mathrm{OH}(l)+\frac{3}{2} \mathrm{O}_{2}(g)-\longrightarrow \mathrm{CO}_{2}(g)+2 \mathrm{H}_{2} \mathrm{O}(l) \\ \Delta H=-726.5 \mathrm{~kJ} \end{array} $$ (a) What is the enthalpy change for the reverse reaction? (b) Balance the forward reaction with whole-number coefficients. What is \(\Delta H\) for the reaction represented by this equation? (c) Which is more likely to be thermodynamically favored, the forward reaction or the reverse reaction? (d) If the reaction were written to produce \(\mathrm{H}_{2} \mathrm{O}(g)\) instead of \(\mathrm{H}_{2} \mathrm{O}(l)\), would you expect the magnitude of \(\Delta H\) to increase, decrease, or stay the same? Explain.

Consider the decomposition of liquid benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}(l)\), to gaseous acetylene, \(\mathrm{C}_{2} \mathrm{H}_{2}(g)\) : $$ \mathrm{C}_{6} \mathrm{H}_{6}(l) \longrightarrow 3 \mathrm{C}_{2} \mathrm{H}_{2}(g) \quad \Delta H=+630 \mathrm{~kJ} $$ (a) What is the enthalpy change for the reverse reaction? (b) What is \(\Delta H\) for the formation of 1 mol of acetylene? (c) Which is more likely to be thermodynamically favored, the forward reaction or the reverse reaction? (d) If \(C_{6} \mathrm{H}_{6}(g)\) were consumed instead of \(\mathrm{C}_{6} \mathrm{H}_{6}(l)\), would you expect the magnitude of \(\Delta H\) to increase, decrease, or stay the same? Explain.

Two solid objects, \(\mathrm{A}\) and \(\mathrm{B}\), are placed in boiling water and allowed to come to temperature there. Each is then lifted out and placed in separate beakers containing \(1000 \mathrm{~g}\) water at \(10.0^{\circ} \mathrm{C}\). Object \(\mathrm{A}\) increases the water temperature by \(3.50^{\circ} \mathrm{C} ; \mathrm{B}\) increases the water temperature by \(2.60{ }^{\circ} \mathrm{C}\). (a) Which object has the larger heat capacity? (b) What can you say about the specific heats of \(\mathrm{A}\) and \(\mathrm{B}\) ?

(a) What is the specific heat of liquid water? (b) What is the molar heat capacity of liquid water? (c) What is the heat capacity of \(185 \mathrm{~g}\) of liquid water? (d) How many \(\mathrm{kJ}\) of heat are needed to raise the temperature of \(10.00 \mathrm{~kg}\) of liquid water from \(24.6^{\circ} \mathrm{C}\) to \(46.2^{\circ} \mathrm{C}\) ?

The specific heat of iron metal is \(0.450 \mathrm{~J} / \mathrm{g}-\mathrm{K}\). How many \(J\) of heat are necessary to raise the temperature of a 1.05-kg block of iron from \(25.0^{\circ} \mathrm{C}\) to \(88.5^{\circ} \mathrm{C}\) ?

The specific heat of ethylene glycol is \(2.42 \mathrm{~J} / \mathrm{g}-\mathrm{K} .\) How many J of heat are needed to raise the temperature of \(62.0 \mathrm{~g}\) of ethylene glycol from \(13.1^{\circ} \mathrm{C}\) to \(40.5^{\circ} \mathrm{C}\) ?

When a 9.55-g sample of solid sodium hydroxide dissolves in \(100.0 \mathrm{~g}\) of water in a coffee-cup calorimeter (Figure 5.17), the temperature rises from \(23.6^{\circ} \mathrm{C}\) to \(47.4^{\circ} \mathrm{C}\). Calculate \(\Delta H\) (in \(\mathrm{kJ} / \mathrm{mol} \mathrm{NaOH}\) ) for the solution process $$ \mathrm{NaOH}(s) \stackrel{-\cdots}{\mathrm{Na}}^{+}(a q)+\mathrm{OH}^{-}(a q) $$ Assume that the specific heat of the solution is the same as that of pure water.

(a) When a 3.88-g sample of solid ammonium nitrate dissolves in \(60.0 \mathrm{~g}\) of water in a coffee-cup calorimeter (Figure 5.17), the temperature drops from \(23.0^{\circ} \mathrm{C}\) to \(18.4^{\circ} \mathrm{C}\). Calculate \(\Delta H\left(\right.\) in \(\mathrm{kJ} / \mathrm{mol} \mathrm{NH}_{4} \mathrm{NO}_{3}\) ) for the solu- tion process $$ \mathrm{NH}_{4} \mathrm{NO}_{3}(s) \rightarrow \mathrm{NH}_{4}{\underline{\phantom{xx}}}^{+}(a q)+\mathrm{NO}_{3}^{-}(a q) $$ Assume that the specific heat of the solution is the same as that of pure water. (b) Is this process endothermic or exothermic?

A \(2.200-g\) sample of quinone \(\left(\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{O}_{2}\right)\) is burned in a bomb calorimeter whose total heat capacity is \(7.854 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\). The temperature of the calorimeter increases from \(23.44^{\circ} \mathrm{C}\) to \(30.57^{\circ} \mathrm{C}\). What is the heat of combustion per gram of quinone? Per mole of quinone?

A \(1.800-g\) sample of phenol \(\left(C_{6} H_{5} O H\right)\) was burned in a bomb calorimeter whose total heat capacity is \(11.66 \mathrm{~kJ} /{ }^{\circ} \mathrm{C}\). The temperature of the calorimeter plus contents increased from \(21.36^{\circ} \mathrm{C}\) to \(26.37^{\circ} \mathrm{C}\). (a) Write a balanced chemical equation for the bomb calorimeter reaction. (b) What is the heat of combustion per gram of phenol? Per mole of phenol?

Under constant-volume conditions the heat of combustion of glucose \(\left(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}\right)\) is \(15.57 \mathrm{~kJ} / \mathrm{g}\). A \(2.500-\mathrm{g}\) sample of glucose is burned in a bomb calorimeter. The temperature of the calorimeter increased from \(20.55^{\circ} \mathrm{C}\) to \(23.25^{\circ} \mathrm{C}\). (a) What is the total heat capacity of the calorimeter? (b) If the size of the glucose sample had been exactly twice as large, what would the temperature change of the calorimeter have been?

Under constant-volume conditions the heat of combustion of benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)\) is \(26.38 \mathrm{~kJ} / \mathrm{g}\). A \(1.640\) \(g\) sample of benzoic acid is burned in a bomb calorimeter. The temperature of the calorimeter increases from \(22.25^{\circ} \mathrm{C}\) to \(27.20^{\circ} \mathrm{C}\). (a) What is the total heat capacity of the calorimeter? (b) A 1.320-g sample of a new organic substance is combusted in the same calorimeter. The temperature of the calorimeter increases from \(22.14^{\circ} \mathrm{C}\) to \(26.82^{\circ} \mathrm{C}\). What is the heat of combustion per gram of the new substance? (c) Suppose that in changing samples, a portion of the water in the calorimeter were lost. In what way, if any, would this change the heat capacity of the calorimeter?

What is the connection between Hess's law and the fact that \(H\) is a state function?

Consider the following hypothetical reactions: $$ \begin{array}{ll} \mathrm{A} \rightarrow \mathrm{B} & \Delta H=+30 \mathrm{~kJ} \\ \mathrm{~B} \longrightarrow \mathrm{C} & \Delta H=+60 \mathrm{~kJ} \end{array} $$ (a) Use Hess's law to calculate the enthalpy change for the reaction \(\mathrm{A}-\cdots \mathrm{C}\) (b) Construct an enthalpy diagram for substances \(\mathrm{A}, \mathrm{B}\), and \(\mathrm{C}\), and show how Hess's law applies.

From the enthalpies of reaction \(2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(g) \quad \Delta H=-483.6 \mathrm{~kJ}\) \(3 \mathrm{O}_{2}(g) \stackrel{-\cdots}{\longrightarrow} 2 \mathrm{O}_{3}(g) \quad \Delta H=+284.6 \mathrm{~kJ}\) calculate the heat of the reaction $$ 3 \mathrm{H}_{2}(g)+\mathrm{O}_{3}(g) \longrightarrow 3 \mathrm{H}_{2} \mathrm{O}(g) $$

Given the data $$ \begin{aligned} \mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \cdots-\rightarrow 2 \mathrm{NO}(g) & \Delta H=+180.7 \mathrm{~kJ} \\ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g)-\cdots 2 \mathrm{NO}_{2}(g) & \Delta H=-113.1 \mathrm{~kJ} \\ 2 \mathrm{~N}_{2} \mathrm{O}(g)-\cdots 2 \mathrm{~N}_{2}(g)+\mathrm{O}_{2}(g) & \Delta H=-163.2 \mathrm{~kJ} \end{aligned} $$ use Hess's law to calculate \(\Delta H\) for the reaction $$ \mathrm{N}_{2} \mathrm{O}(g)+\mathrm{NO}_{2}(g)-\cdots 3 \mathrm{NO}(g) $$

(a) What is meant by the term standard conditions, with reference to enthalpy changes? (b) What is meant by the term enthalpy of formation? (c) What is meant by the term standard enthalpy of formation?

(a) Why are tables of standard enthalpies of formation so useful? (b) What is the value of the standard enthalpy of formation of an element in its most stable form? (c) Write the chemical equation for the reaction whose enthalpy change is the standard enthalpy of formation of glucose, \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(s), \Delta H_{f}^{\circ}\left[\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}\right]\).

Using values from Appendix \(C\), calculate the standard enthalpy change for each of the following reactions: (a) \(2 \mathrm{SO}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{SO}_{3}(g)\) (b) \(\mathrm{Mg}(\mathrm{OH})_{2}(s) \longrightarrow \mathrm{MgO}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) (c) \(\mathrm{N}_{2} \mathrm{O}_{4}(g)+4 \mathrm{H}_{2}(g) \longrightarrow \mathrm{N}_{2}(g)+4 \mathrm{H}_{2} \mathrm{O}(g)\) (d) \(\mathrm{SiCl}_{4}(l)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow \mathrm{SiO}_{2}(s)+4 \mathrm{HCl}(g)\)

Using values from Appendix \(C\), calculate the value of \(\Delta H^{\circ}\) for each of the following reactions: (a) \(4 \mathrm{HBr}(\mathrm{g})+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{Br}_{2}(l)\) (b) \(2 \mathrm{Na}(\mathrm{OH})(s)+\mathrm{SO}_{3}(g) \longrightarrow \mathrm{Na}_{2} \mathrm{SO}_{4}(s)+\mathrm{H}_{2} \mathrm{O}(g)\) (c) \(\mathrm{CH}_{4}(g)+4 \mathrm{Cl}_{2}(g) \longrightarrow \mathrm{CCl}_{4}(l)+4 \mathrm{HCl}(g)\) (d) \(\mathrm{Fe}_{2} \mathrm{O}_{3}(s)+6 \mathrm{HCl}(g) \longrightarrow 2 \mathrm{FeCl}_{3}(s)+3 \mathrm{H}_{2} \mathrm{O}(g)\)