A coffee-cup calorimeter of the type shown in Figure 5.17 contains 150.0 g of water at

Imagine a book that is falling from a shelf. At a particular moment during its fall, the book has a kinetic energy of 13 J and a potential energy with respect to the floor of 72 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]

Consider the accompanying energy diagram. (a) Does this diagram represent an increase or decrease in the internal energy of the system? (b) What sign is given to ll.E for this process? (c) If there is no work associated with the process, is it exothermic or endothermic? [Section 5.2]

The contents of the closed box in each of the following illustrations represent a system, and the arrows show the changes to the system during some process. The lengths of the arrows represent the relative magnitudes of q and w (a) Which of these processes is endothermic? (b) For which of these processes, if any, is ll.E < 0? (c) For which process, if any, is there a net gain in internal energy? [Section 5.2]

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]

In the cylinder diagrammed below, a chemical process occurs at constant temperature and pressure. (a) Is the sign of w indicated by this change positive or negative? (b) If the process is endothermic, does the internal energy of the system within the cylinder increase or decrease during the change and is ll.E positive or negative? [Sections 5.2 and 5.3]

Imagine a container placed in a tub of water, as depicted in the accompanying diagram. (a) If the contents of the container are the system and 350 K 290 K heat is able to flow through the container walls, what qualitative changes will occur in the temperatures of the system and in its surroundings? What is the sign of q associated with each change? From the system's perspective, is the process exothermic or endothermic? (b) If neither the volume nor the pressure of the system changes during the process, how is the change in internal energy related to the change in enthalpy? [Sections 5.2 and 5.3]

Which will release more heat as it cools from 50 oc to 25 C, 1 kg of water or 1 kg of aluminum? How do you know? [Section 5.5]

A gas-phase reaction was run in an apparatus designed to maintain a constant pressure. (a) Write a balanced chemical equation for the reaction depicted, and predict whether w is positive, negative, or zero. (b) Using data from Appendix C, determine !l.H for the formation of one mole of the product. Why is this enthalpy change called the enthalpy of formation of the involved product? [Sections 5.3 and 5.7]

Consider the two diagrams below. (a) Based on (1), write an equation showing how !l.HA is related to !l.H8 and !l.Hc. How do both diagram (i) and your equation relate to the fact that enthalpy is a state function? (b) Based on (ii), write an equation relating !l.Hz to the other enthalpy changes in the diagram. (c) How do these diagrams relate to Hess's law? [Section 5.6]

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 m/s. (b) Convert this energy to calories. (c) What happens to this energy when the ball lands in a sand trap?

(a) What is the kinetic energy in joules of an 850-lb motorcycle moving at 66 mph? (b) By what factor will the kinetic energy change if the speed of the motorcycle is decreased to 33 mph? (c) Where does the kinetic energy of the motorcycle go when the rider brakes to a stop?

The use of the British thermal unit (Btu) is common in much engineering work. A Btu is the amount of heat required to raise the temperature of 1 lb of water by 1 F. Calculate the number of joules in a Btu.

A watt is a measure of power (the rate of energy change) equal to 1 Jjs. (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 proper- In ties of a solution in an apparatus as illustrated. A solution is continuously flowing into the apparatus at the top Out 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?

Identify the force present, and explain whether work is being performed in the following cases: (a) You lift a pencil off the top of a desk. (b) A spring is compressed to half its normal length.

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 intem111 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 11, and determine whether the process is endothermic or exothermic for the following cases: (a) A system absorbs 105 kJ of heat from its surroundings while doing 29 kJ of work on the surroundings; (b) q = 1.50 kJ and w = -657 J; (c) the system releases 57.5 kJ of heat while doing 22.5 kJ of work on the surroundings.

For the following processes, calculate the change in internal energy of the system and determine whether the process is endothermic or exothermic: (a) A balloon is heated by adding 850 J of heat. It expands, doing 382 J of work on the atmosphere. (b) A 50-g sample of water is cooled from 30 C to 15 C, thereby losing approximately 3140 J of heat. (c) A chemical reaction releases 6.47 kJ of heat and does no work on the surroundings.

A gas is confined to a cylinder fitted with a piston and an electrical heater, as shown in the accompanying illustration. Suppose that current is supplied to the heater so that 100 J of energy is added. Consider two different situations. In case (1) the piston is allowed to move as the energy is added. In case (2) the piston is fixed so that it cannot move. (a) In which case does the gas have the higher temperature after addition of the electrical energy? Explain. (b) What can you say about the values of q and w in each of these cases? (c) What can you say about the relative values of flE for the system (the gas in the cylinder) in the two cases?

Consider a system consisting of two oppositely charged spheres hanging by strings and separated by a distance r1, as shown in the accompanying illustration. Suppose they are separated to a larger distance r2, by moving them apart along a track. (a) What change, if any, has occurred in the potential energy of the system? (b) What effect, if any, does this process have on the value of !J.E? (c) What can you say about q and w for this process?

(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?

(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, !J.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 !J.H for a process that occurs at constant pressure. What additional information do you need to determine !J.E for the process?

Suppose that the gas-phase reaction 2 NO(g) + 02(g) ----> 2 N02(g) were carried out in a constant-volume container at constant temperature. Would the measured heat change represent !J.H or !J.E? lf there is a difference, which quantity is larger for this reaction? Explain.

A gas is confined to a cylinder under constant atmospheric pressure, as illustrated in Figure 5.3. When the gas undergoes a particular chemical reaction, it releases 79 kJ of heat to its surroundings and does 18 kJ of P-V work on its surroundings. What are the values of flH and !J.E for this process?

A gas is confined to a cylinder under constant atmospheric pressure, as illustrated in Figure 5.3. When 378 J of heat is added to the gas, it expands and does 56 J of work on the surroundings. What are the values of t.H and D. for this process?

The complete combustion of acetic acid, CH3COOH(I), to form H20(1) and C02{g) at constant pressure releases 871.7 kJ of heat per mole of CH3COOH. (a) Write a balanced thermochemical equation for this reaction. (b) Draw an enthalpy diagram for the reaction.

The decomposition of zinc carbonate, ZnC03(s), into zinc oxide, ZnO(s), and C02{g) at constant pressure requires the addition of 71.5 kJ of heat per mole of ZnC03. (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: t.H = -243.4 kJ Which has the higher enthalpy under these conditions, 2 Cl{g) or Clz(g)?

Consider the following reaction: 2 Mg(s) + 02{g) -> 2 MgO(s) t.H = -1204 kJ (a) Is this reaction exothermic or endothermic? (b) Calculate the amount of heat transferred when 2.4 g of Mg(s) reacts at constant pressure. (c) How many grams of MgO are produced during an enthalpy change of -96.0 kJ? (d) How many kilojoules of heat are absorbed when 7.50 g of MgO(s) is decomposed into Mg(s) and 02{g) at constant pressure?

Consider the following reaction: t.H = +90.7 kJ (a) Is heat absorbed or released in the course of this reaction? (b) Calculate the amount of heat transferred when 45.0 g of CH30H{g) is decomposed by this reaction at constant pressure. (c) For a given sample of CH30H, the enthalpy change on reaction is 25.8 kJ. How many grams of hydrogen gas are produced? What is the value of t.H for the reverse of the previous reaction? (d) How many kilojoules of heat are released when 50.9 g of CO(g) reacts completely with H2{g) to form CHPH{g) at constant pressure?

When solutions containing silver ions and chloride ions are mixed, silver chloride precipitates: Ag+(aq) + cqaq) -> AgCl(s) t.H = -65.5 kJ (a) Calculate t.H for production of 0.200 mol of AgCl by this reaction. (b) Calculate t.H for the production of 2.50 g of AgCI. (c) Calculate t.H when 0.150 mmol of AgCI dissolves in water.

At one time, a common means of forming small quantities of oxygen gas in the laboratory was to heat KC103: 2 KCI03(s) -> 2 KCl(s) + 3 02{g) t.H = -89.4 kJ For this reaction, calculate t.H for the formation of (a) 0.632 mol of 02 and (b) 8.57 g of KCI. (c) The decomposition of KC103 proceeds spontaneously when it is heated. Do you think that the reverse reaction, the formation of KC103 from KCl and 02, is likely to be feasible under ordinary conditions? Explain your answer.

Consider the combustion of liquid methanol, CH30H(/): CHPH(I) + 02{g) -> C02{g) + 2 H20( I) t.H = -726.5 kJ (a) What is the enthalpy change for the reverse reaction? (b) Balance the forward reaction with whole-number coefficients. What is D.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 H20{g) instead of H20(/), would you expect the magnitude of t.H to increase, decrease, or stay the same? Explain.

Consider the decomposition of liquid benzene, C6H6(/), to gaseous acetylene, C2H2{g): t.H = +630 kJ (a) What is the enthalpy change for the reverse reaction? (b) What is t.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 C6{g) were consumed instead of C6H6(1), would you expect the magnitude of t.H to increase, decrease, or stay the same? Explain.

(a) What are the units of molar heat capacity? (b) What are the units of specific heat? (c) If you know the specific heat of copper, what additional information do you need to calculate the heat capacity of a particular piece of copper pipe?

Two solid objects, A and 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 g water at 10.0 oc. Object A increases the water temperature by 3.50 oc; B increases the water temperature by 2.60 oc. (a) Which object has the larger heat capacity? (b) What can you say about the specific heats of A and 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 g of liquid water? (d) How many kJ of heat are needed to raise the temperature of 10.00 kg of liquid water from 24.6 oc to 46.2 C?

The specific heat of iron metal is 0.450 J/g-K. How many J of heat are necessary to raise the temperature of a 1.05-kg block of iron from 25.0 oc to 88.5 C?

The specific heat of ethylene glycol is 2.42 J/g-K. How many j of heat are needed to raise the temperature of 62.0 g of ethylene glycol from 13.1 oc to 40.5 C?

When a 9.55-g sample of solid sodium hydroxide dissolves in 100.0 g of water in a coffee-cup calorimeter (Figure 5.17), the temperature rises from 23.6 oc to 47.4 oc. Calculate !lH (in kjjmol NaOH) for the solution process NaOH(s) ----+ Na+(aq) + OH-(aq) 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 g of water in a coffee-cup calorimeter (Figure 5.17), the temperature drops from 23.0 oc to 18.4 C. Calculate !lH (in k)/mol NH4N03) for the solution process NH03(s) ----+ NH4 +(aq) + N03 -(aq) 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 (C6H402) is burned in a bomb calorimeter whose total heat capacity is 7.854 kjj"C. The temperature of the calorimeter increases from 23.44 oc to 30.57 oc. What is the heat of combustion per gram of quinone? Per mole of quinone?

A 1.800-g sample of phenol (C50H) was burned in a bomb calorimeter whose total heat capacity is 1 1.66 kjj"C. The temperature of the calorimeter plus contents increased from 21.36 oc to 26.37 oc. (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 (C6H1P6) is 15.57 k)/g. A 2.500-g sample of glucose is burned in a bomb calorimeter. The temperature of the calorimeter increased from 20.55 oc to 23.25 oc. (a) What is the total heat capacity of the calorimeter? (b) lf 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 (C6H5COOH) is 26.38 k)/g. A 1.640- g sample of benzoic acid is burned in a bomb calorimeter. The temperature of the calorimeter increases from 22.25 oc to 27.20 C. (a) What is the total heat capacity of the calorimeter? (b) A 1 .320-g sample of a new organic substance is com busted in the same calorimeter. The temperature of the calorimeter increases from 22.14 oc to 26.82 oc. 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?

Calculate the enthalpy change for the reaction P406(s) + 2 02(gl ----+ P4010(s) given the following enthalpies of reaction: P4(s) + 3 02(g) ----+ P406(s) P4(s) + 5 02(g) ----+ P4010(s) !lH = -1640.1 kj !lH = -2940.1 kj

From the enthalpies of reaction 2 H2(g) + 02(g) ----+ 2 HP(g) 3 02(g) ----+ 2 03(g) !lH = -483.6 kj !lH = +284.6 kj calculate the heat of the reaction

From the enthalpies of reaction Hz(g) + F2(g) ----+ 2 HF(g) C(s) + 2 F2(g) ----+ CF4(g) 2 C(s) + 2 H2(g) ----+ C2H4(g) !lH = -537 kj !lH = -680 kj !lH = +52.3 kj calculate flH for the reaction of ethylene with F2: C2H4(g) + 6 F2(g) ----+ 2 CF4(g) + 4 HF(g)

Given the data N2(g) + 02(g) ----+ 2 NO(g) 2 NO(g) + 02(g) ----+ 2 N02(g) 2 N20(g) ----+ 2 N2(g) + 02(g) !lH = + 180.7 kj !lH = -113.1 kj !lH = -163.2 kj use Hess's law to calculate !lH for the reaction N20(g) + NOz(g) ----+ 3 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, C6Hn06(s), llH'([C6Hn061

For each of the following compounds, write a balanced thermochemical equation depicting the formation of one mole of the compound from its elements in their standard states and use Appendix C to obtain the value of llHJ: (a) NH3(g), (b) S02(g), (c) RbCI03(s), (d) NH03(s).

Write balanced equations that describe the formation of the following compounds from elements in their stan dard states, and use Appendix C to obtain the values of their standard enthalpies of formation: (a) HBr(g), (b) AgN03(s), (c) Fep3(s), (d) CH3COOH(/).

The following is known as the thermite reaction [Figure 5.7(b)]: 2 Al(s) + Fep3(s) -----> Alp3(s) + 2 Fe(s) This highly exothermic reaction is used for welding massive units, such as propellers for large ships. Using standard enthalpies of formation in Appendix C, calculate !lH0 for this reaction.

Using values from Appendix C, calculate the standard enthalpy change for each of the following reactions: (a) 2 S02(g) + 02(g) -----> 2 S03(g) (b) Mg(OHh(s) -----> MgO(s) + H20(/) (c) N204(g) + 4 H2(g) -----> N2(g) + 4 H20(g) (d) SiC14(1) + 2 H20(/) -----> Si02(s) + 4 HCI(g)

Using values from Appendix C. calculate the value of llH0 for each of the following reactions: (a) 4 HBr(g) + 02(g) -----> 2 H20(/) + 2 Br2(/) (b) 2 Na(OH)(s) + S03(g) -----> Na2S04(s) + Hp(g) (c) CH4(g) + 4 Cl2(g) -----> CC14(1) + 4 HCI(g) (d) Fe203(s) + 6 HCI(g) -----> 2 FeCI3(s) + 3 Hp(g)

Complete combustion of 1 mol of acetone (C3H60) liberates 1790 kJ: C3H60(I) + 4 02(g) -----> 3 C02(g) + 3 H20(/) flW = -1790 kJ Using this information together with data from Appendix C, calculate the enthalpy of formation of acetone.

Calcium carbide (CaC2) reacts with water to form acetylene (C2H2) and Ca(OHh- From the following enthalpy of reaction data and data in Appendix C, calculate llHJ for CaC2(s): CaC2(s) + 2 H20(/) -----> Ca(OHh(s) + C2H2(g) flW = -127.2 kJ

Gasoline is composed primarily of hydrocarbons, including many with eight carbon atoms, called octanes. One of the cleanest-burning octanes is a compound called 2,3,4-trirnethylpentane, which has the following structural formula: CH3 CH3 CH3 I I I H3C-CH-CH-CH -CH3 The complete combustion of one mole of this compound to C02(g) and H20(g) leads to !lH0 = -5064.9 kJ/mol. (a) Write a balanced equation for the combustion of 1 mol of CsH!s(l). (b) Write a balanced equation for the formation of C8H18(1) from its elements. (c) By using the information in this problem and data in Table 5.3, calculate llHJ for 2,3,4-trimethylpentane.

Naphthalene (C10H8) is a solid aromatic compound often sold as mothballs. The complete combustion of this substance to yield C02(g) and H20(/) at 25 oc yields 5154 k)/mol. (a) Write balanced equations for the formation of naphthalene from the elements and for its combustion. (b) Calculate the standard enthalpy of formation of naphthalene.

Ethanol (C2H50H) is currently blended with gasoline as an automobile fuel. (a) Write a balanced equation for the combustion of liquid ethanol in air. (b) Calculate the standard enthalpy change for the reaction, assuming H20(g) as a product. (c) Calculate the heat produced per liter of ethanol by combustion of ethanol under constant pressure. Ethanol has a density of 0.789 g/mL. (d) Calculate the mass of C02 produced per kJ of heat emitted.

Methanol (CH30H) is used as a fuel in race cars. (a) Write a balanced equation for the combustion of liquid methanol in air. (b) Calculate the standard enthalpy change for the reaction, assuming H20(g) as a product. (c) Calculate the heat produced by combustion per liter of methanol. Methanol has a density of 0.791 g/mL. (d) Calculate the mass of C02 produced per kJ of heat emitted.

(a) What is meant by the term fuel value? (b) Which is a greater source of energy as food, 5 g of fat or 9 g of carbohydrate?

A serving of condensed cream of mushroom soup contains 7 g fat, 9 g carbohydrate, and 1 g protein. Estimate the number of Calories in a serving.

A pound of plain M&M candies contains 96 g fat, 320 g carbohydrate, and 21 g protein. What is the fuel value in kJ in a 42-g (about 1.5 oz) serving? How many Calories does it provide?

The heat of combustion of fructose, C6H1206, is -2812 kJ/mol. If a fresh golden delicious apple weighing 4.23 oz (120 g) contains 16.0 g of fructose, what caloric content does the fructose contribute to the apple?

The heat of combustion of ethanol, C2H50H(I), is -1367 kJ/mol. A batch of Sauvignon Blanc wine contains 10.6% ethanol by mass. Assuming the density of the wine to be 1.0 g/mL, what caloric content does the alcohol (ethanol) in a 6-oz glass of wine (177 mL) have?

The standard enthalpies of formation of gaseous propyne (C3H4), propylene (C3H6), and propane (C3Hs) are + 185.4, +20.4, and -103.8 k)/mol, respectively. (a) Calculate the heat evolved per mole on combustion of each substance to yield C02(g) and H20(g). (b) Calculate the heat evolved on combustion of 1 kg of each substance. (c) Which is the most efficient fuel in terms of heat evolved per unit mass?

It is interesting to compare the "fuel value" of a hydrocarbon in a world where fluorine rather than oxygen is the combustion agent. The enthalpy of formation of CF4(g) is -679.9 k)/mol. Which of the following two reactions is the more exothermic? CH4(g) + 2 02(g) -> C02(g) + 2 H20(g) CH4(g) + 4 F2(g) -> CF4(g) + 4 HF(g)

At 20 C (approximately room temperature) the average velocity of N2 molecules in air is 1050 mph. (a) What is the average speed in m/s? (b) What is the kinetic energy (in J) of an N2 molecule moving at this speed? (c) What is the total kinetic energy of 1 mol of N2 molecules moving at this speed?

Suppose an Olympic diver who weighs 52.0 kg executes a straight dive from a 10-m platform. At the apex of the dive, the diver is 10.8 m above the surface of the water. (a) What is the potential energy of the diver at the apex of the dive, relative to the surface of the water? (b) Assuming that all the potential energy of the diver is converted into kinetic energy at the surface of the water, at what speed in m/s will the diver enter the water? (c) Does the diver do work on entering the water? Explain.

When a mole of dry ice, C02(s), is converted to C02(g) at atmospheric pressure and -78 C, the heat absorbed by the system exceeds the increase in internal energy of the C02. Why is this so? What happens to the remaining energy?

An aluminum can of a soft drink is placed in a freezer. Later, you find that the can is split open and its contents frozen. Work was done on the can in splitting it open. Where did the energy for this work come from?

A sample of gas is contained in a cylinder-and-piston arrangement. It undergoes the change in state shown in the drawing. (a) Assume first that the cylinder and piston are perfect thermal insulators that do not allow heat to be transferred. What is the value of q for the state change? What is the sign of w for the state change? What can be said about fl for the state change? (b) Now assume that the cylinder and piston are made up of a thermal conductor such as a metal. During the state change, the cylinder gets warmer to the touch. What is the sign of q for the state change in this case? Describe the difference in the state of the system at the end of the process in the two cases. What can you say about the relative values of fl?

Limestone stalactites and stalagmites are formed in caves by the following reaction: Ca2+ (aq) + 2 HC03 -(aq) -> CaC03(s) + COz(g) + HzO(/) If 1 mol of CaC03 forms at 298 K under 1 atm pressure, the reaction performs 2.47 kJ of P-V work, pushing back the atmosphere as the gaseous C02 forms. At the same time, 38.95 kj of heat is absorbed from the environment. What are the values of flH and of fl for this reaction?

Consider the systems shown in Figure 5.9. In one case the battery becomes completely discharged by running the current through a heater, and in the other by running a fan. Both processes occur at constant pressure. In both cases the change in state of the system is the same: The battery goes from being fully charged to being fully discharged. Yet in one case the heat evolved is large, and in the other it is small. Is the enthalpy change the same in the two cases? If not, how can enthalpy be considered a state function? If it is, what can you say about the relationship between enthalpy change and q in this case, as compared with others that we have considered?

The enthalpy change for melting ice at 0 oc and constant atmospheric pressure is 6.01 kJ/mol. Calculate the quantity of energy required to melt a moderately large iceberg with a mass of 1.25 million metric tons. (A metric ton is 1000 kg.)

Comparing the energy associated with the rainstorm and that of a conventional explosive gives some idea of the immense amount of energy associated with a storm. (a) The heat of vaporization of water is 44.0 k)/mol. Calculate the quantity of energy released when enough water vapor condenses to form 0.50 inches of rain over an area of one square mile. (b) The energy released when one ton of dynamite explodes is 4.2 x 106 kj. Calculate the number of tons of dynamite needed to provide the energy of the storm in part (a).

A house is designed to have passive solar energy features. Brickwork incorporated into the interior of the house acts as a heat absorber. Each brick weighs approximately 1.8 kg. The specific heat of the brick is 0.85 J/g-K. How many bricks must be incorporated into the interior of the house to provide the same total heat capacity as 1.7 x 103 gal of water?

A coffee-cup calorimeter of the type shown in Figure 5.17 contains 150.0 g of water at 25.1 oc. A 121.0-g block of copper metal is heated to 100.4 oc by putting it in a beaker of boiling water. The specific heat of Cu(s) is 0.385 J/g-K. The Cu is added to the calorimeter, and after a time the contents of the cup reach a constant temperature of 30.1C. (a) Determine the amount of heat, in J, lost by the copper block. (b) Determine the amount of heat gained by the water. The specific heat of water is 4.18 J/g-K. (c) The difference between your answers for (a) and (b) is due to heat loss through the Styrofoam cups and the heat necessary to raise the temperature of the inner wall of the apparatus. The heat capacity of the calorimeter is the amount of heat necessary to raise the temperature of the apparatus (the cups and the stopper) by 1 K. Calculate the heat capacity of the calorimeter in J/K. (d) What would be the final temperature of the system if all the heat lost by the copper block were absorbed by the water in the calorimeter?

(a) When a 0.235-g sample of benzoic acid is combusted in a bomb calorimeter, the temperature rises 1.642 oc. When a 0.265-g sample of caffeine, CsH wOzN4, is burned, the temperature rises 1.525 C. Using the value 26.38 kJ/g for the heat of combustion of benzoic acid, calculate the heat of combustion per mole of caffeine at constant volume. (b) Assuming that there is an uncertainty of 0.002 oc in each temperature reading and that the masses of samples are measured to 0.001 g, what is the estimated uncertainty in the value calculated for the heat of combustion per mole of caffeine?

Meals-ready-to-eat (MREs) are military meals that can be heated on a flameless heater. The heat is produced by the following reaction: Mg(s) + 2 H20(l) -> Mg(OH)z(s) + H2(g). (a) Calculate the standard enthalpy change for this reaction. (b) Calculate the number of grams of Mg needed for this reaction to release enough energy to increase the temperature of 25 mL of water from 15 oc to 85 oc.

Burning methane in oxygen can produce three different carbon-containing products: soot (very fine particles of graphite), CO(g), and C02(g). (a) Write three balanced equations for the reaction of methane gas with oxygen to produce these three products. In each case assume that HzO(l) is the only other product. (b) Determine the standard enthalpies for the reactions in part (a). (c) Why, when the oxygen supply is adequate, is C02(g) the predominant carbon-containing product of the combustion of methane?

(a) Calculate the standard enthalpy of formation of gaseous diborane (BzH6) using the following thermochemical information: 4 B(s) + 3 02(g) -> 2 B203(s) !lW = -2509.1 kJ 2 H2(g) + 02(g) -> 2 H20(/) !lW = -571.7 kJ B2H6(g) + 3 02(g) -> B203(s) + 3 H20(/) !lW = -2147.5 kJ (b) Pentaborane (B5H9) is another boron hydride. What experiment or experiments would you need to perform to yield the data necessary to calculate the heat of formation of B5H9(/)? Explain by writing out and summing any applicable chemical reactions.

From the following data for three prospective fuels, calculate which could provide the most energy per unit volume: Density Molar Enthalpy at 20 oc of Combustion Fuel (gjcm3) kJ/mol Nitroethane, C2H5N02(1) 1.052 -1368 Ethanol, CzHsOH(I) 0.789 -1367 Methylhydrazine, CH6N2(1) 0.874 -1305

The hydrocarbons acetylene (C2H2) and benzene (C6H6) have the same empirical formula. Benzene is an "aromatic" hydrocarbon, one that is unusually stable because of its structure. (a) By using the data in Appendix C, determine the standard enthalpy change for the reaction 3 C2H2(g) -> C6H6(l). (b) Which has greater enthalpy, 3 mol of acetylene gas or 1 mol of liquid benzene? (c) Determine the fuel value in kJ/g for acetylene and benzene.

Ammonia (NH3) boils at -33 oc; at this temperature it has a density of 0.81 g/cm3 The enthalpy of formation of NH3(g) is -46.2 kJ/mol, and the enthalpy of vaporization of NH3(/) is 23.2 kJ/mol. Calculate the enthalpy change when 1 L of liquid NH3 is burned in air to give N2(g) and HzQ(g). How does this compare with !lH for the complete combustion of 1 L of liquid methanol, CHpH(I)? For CH30H(/), the density at 25 oc is 0.792 gjcm3 , and t.Hj equals -239 kjjmol.

Three common hydrocarbons that contain four carbons are listed here, along with their standard enthalpies of formation: Hydrocarbon 1,3-Butadiene 1-Butene n-Butane Formula C4H6(g) C4Hs(gl C4Hw(gl !lHJ (kJ/mol) 111.9 1.2 -124.7 (a) For each of these substances, calculate the molar enthalpy of combustion to C02(g) and H20(l). (b) Calculate the fuel value in kJ/g for each of these compounds. (c) For each hydrocarbon, determine the percentage of hydrogen by mass. (d) By comparing your answers for parts (b) and (c), propose a relationship between hydrogen content and fuel value in hydrocarbons.

The two common sugars, glucose (C1206) and sucrose (C12H22011), are both carbohydrates. Their standard enthalpies of formation are given in Table 5.3. Using these data, (a) calculate the molar enthalpy of combustion to C02(g) and H20(/) for the two sugars; (b) calculate the enthalpy of combustion per gram of each sugar; (c) determine how your answers to part (b) compare to the average fuel value of carbohydrates discussed in Section 5.8.

A 200-lb man decides to add to his exercise routine by walking up three flights of stairs (45 ft) 20 times per day. He figures that the work required to increase his potential energy in this way will permit him to eat an extra order of French fries, at 245 Cal, without adding to his weight. Is he correct in this assumption?

It is estimated that the net amount of carbon dimdde fixed by photosynthesis on the landmass of Earth is 5.5 X 1016 g/yr of C02. Assume that all this carbon is converted into glucose. (a) Calculate the energy stored by photosynthesis on land per year in kJ. (b) Calculate the average rate of conversion of solar energy into plant energy in MW (1 W = 1 J/s). A large nuclear power plant produces about 103 MW. The energy of how many such nuclear power plants is equivalent to the solar energy conversion?

Consider the combustion of a single molecule of CH4(g) forming H20( I) as a product. (a) How much energy, in J, is produced during this reaction? (b) A typical X-ray photon has an energy of 8 keV. How does the energy of combustion compare to the energy of the X-ray photon?

Consider the dissolving of NaCI in water, illustrated in Figure 4.3. Assume the system consists of 0.1 mol NaCI and 1 L of water. Considering that the NaCI readily dissolves in the water and that the ions are strongly stabilized by the water molecules, as shown in the figure, is it safe to conclude that the dissolution of NaCI in water results in a lower enthalpy for the system? Explain your response. What experimental evidence would you examine to test this question?

Consider the following unbalanced oxidation-reduction reactions in aqueous solution: Ag+ (aq) + Li(s) ---+ Ag(s) + u + (aq) Fe(s) + Na +(aq) ---+ Fe2 +(aq) + Na(s) K(s) + HzO(I) ---+ KOH(aq) + H2(g) (a) Balance each of the reactions. (b) By using data in Appendix C, calculate /lH0 for each of the reactions. (c) Based on the values you obtain for /lH0, which of the reactions would you expect to be thermodynamically favored? (That is, which would you expect to be spontaneous?) (d) Use the activity series to predict which of these reactions should occur. ax> (Section 4.4) Are these results in accord with your conclusion in part (c) of this problem?

Consider the following acid-neutralization reactions involving the strong base NaOH(aq): HN03(aq) + NaOH(aq) ---+ NaN03(aq) + HzO(I) HCl(aq) + NaOH(aq) ---+ NaCl(aq) + H20(/) NH4+(aq) + NaOH(aq) ----> NH3(aq) + Na+(aq) + H20(/) (a) By using data in Appendix C, calculate !lH0 for each of the reactions. (b) As we saw in Section 4.3, nitric acid and hydrochloric acid are strong acids. Write net ionic equations for the neutralization of these acids. (c) Compare the values of t.H0 for the first two reactions. What can you conclude? (d) In the third equation NH4 +(aq) is acting as an acid. Based on the value of /lH0 for this reaction, do you think it is a strong or a weak acid? Explain.

Consider two solutions, the first being 50.0 mL of 1.00 M CuS04 and the second 50.0 mL of 2.00 M KOH. When the two solutions are mixed in a constant-pressure calorimeter, a precipitate forms and the temperature of the mixture rises from 21.5 oc to 27.7 oc. (a) Before mixing, how many grams of Cu are present in the solution of CuS04? (b) Predict the identity of the precipitate in the reaction. (c) Write complete and net ionic equations for the reaction that occurs when the two solutions are mixed. (d) From the calorimetric data, calculate t.H for the reaction that occurs on mixing. Assume that the calorimeter absorbs only a negligible quantity of heat, that the total volume of the solution is 100.0 mL, and that the specific heat and density of the solution after mixing are the same as that of pure water.

The precipitation reaction between AgN03(aq) and NaCl(aq) proceeds as follows: AgN03(aq) + NaCl(aq) ----> NaN03(aq) + AgCl(s) (a) By using Appendix C, calculate !lH0 for the net ionic equation of this reaction. (b) What would you expect for the value of t.Ho of the overall molecular equation compared to that for the net ionic equation? Explain. (c) Use the results from (a) and (b) along with data in Appendix C to determine the value of t.HJ for AgN03(aq).

A sample of a hydrocarbon is combusted completely in 02(g) to produce 21.83 g C02(g), 4.47 g H20(g), and 311 kJ of heat. (a) What is the mass of the hydrocarbon sample that was combusted? (b) What is the empirical formula of the hydrocarbon? (c) Calculate the value of t.HJ per empirical-formula unit of the hydrocarbon. (d) Do you think that the hydrocarbon is one of those listed in Appendix C? Explain your answer.

The methane molecule, CH4, has the geometry shown in Figure 2.21. Imagine a hypothetical process in which the methane molecule is "expanded," by simultaneously extending all four C-H bonds to inlinity. We then have the process CH4(g) ----> C(g) + 4 H(g) (a) Compare this process with the reverse of the reaction that represents the standard enthalpy of formation. (b) Calculate the enthalpy change in each case. Which is the more endothermic process? What accounts for the difference in !lH0 values? (c) Suppose that 3.45 g CH4(g) is reacted with 1.22 g F2(g), forming CF4(g) and HF(g) as sole products. What is the limiting reagent in this reaction? If the reaction occurs at constant pressure, what amount of heat is evolved?