Assuming that the combustion of hydrogen gas provides three times as much energy per

Calculate DE for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system.

You are calculating DE in a chemistry problem. What if you confuse the system and the surroundings? How would this affect the magnitude of the answer you calculate? The sign?

Calculate the work associated with the expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm.

A balloon is being inflated to its full extent by heating the air inside it. In the final stages of this process, the volume of the balloon changes from 4.00 3 106 L to 4.50 3 106 L by the addition of 1.3 3 108 J of energy as heat. Assuming that the balloon expands against a constant pressure of 1.0 atm, calculate DE for the process. (To convert between L ? atm and J, use 1 L ? atm 5 101.3 J.)

When 1 mole of methane (CH4) is burned at constant pressure, 890 kJ of energy is released as heat. Calculate DH for a process in which a 5.8-g sample of methane is burned at constant pressure.

When 1.00 L of 1.00 M Ba(NO3)2 solution at 25.08C is mixed with 1.00 L of 1.00 M Na2SO4 solution at 25.08C in a calorimeter, the white solid BaSO4 forms, and the temperature of the mixture increases to 28.18C. Assuming that the calorimeter absorbs only a negligible quantity of heat, the specific heat capacity of the solution is 4.18 J/8C ? g, and the density of the final solution is 1.0 g/mL, calculate the enthalpy change per mole of BaSO4 formed.

It has been suggested that hydrogen gas obtained by the decomposition of water might be a substitute for natural gas (principally methane). To compare the energies of combustion of these fuels, the following experiment was carried out using a bomb calorimeter with a heat capacity of 11.3 kJ/8C. When a 1.50-g sample of methane gas was burned with excess oxygen in the calorimeter, the temperature increased by 7.38C. When a 1.15-g sample of hydrogen gas was burned with excess oxygen, the temperature increase was 14.38C. Compare the energies of combustion (per gram) for hydrogen and methane.

What if Hesss law were not true? What are some possible repercussions this would have?

Diborane (B2H6) is a highly reactive boron hydride that was once considered as a possible rocket fuel for the U.S. space program. Calculate DH for the synthesis of diborane from its elements, according to the equation 2B1s2 1 3H2 1g2 h B2H6 1g2 using the following data: Reaction DH (a) 2B1s2 1 1 3 2O2 1g2 h B2O3 1s2 21273 kJ (b) B2H6 1g2 1 3O2 1g2 h B2O3 1s2 1 3H2O1g2 22035 kJ (c) H2 1g2 1 1 1 2 O2 1g2 h H2O1l2 2286 kJ (d) H2O1l2 h H2O1g2 44 kJ

Using the standard enthalpies of formation listed in Table 6.2, calculate the standard enthalpy change for the overall reaction that occurs when ammonia is burned in air to form nitrogen dioxide and water. This is the first step in the manufacture of nitric acid. 4NH3 1g2 1 7O2 1g2 h 4NO2 1g2 1 6H2O1l2

Using enthalpies of formation, calculate the standard change in enthalpy for the thermite reaction: 2Al1s2 1 Fe2O3 1s2 h Al2O3 1s2 1 2Fe1s2 This reaction occurs when a mixture of powdered aluminum and iron(III) oxide is ignited with a magnesium fuse.

For DHreaction calculations, we define DH8f for an element in its standard state as zero. What if we define DH8f for an element in its standard state as 10 kJ/mol? How would this affect your determination of DHreaction? Provide support for your answer with a sample calculation.

Until recently, methanol (CH3OH) was used as a fuel in high-performance engines in race cars. Using the data in Table 6.2, compare the standard enthalpy of combustion per gram of methanol with that per gram of gasoline. Gasoline is actually a mixture of compounds, but assume for this problem that gasoline is pure liquid octane (C8H18)

Compare the energy available from the combustion of a given volume of methane and the same volume of hydrogen at the same temperature and pressure.

Assuming that the combustion of hydrogen gas provides three times as much energy per gram as gasoline, calculate the volume of liquid H2 (density 5 0.0710 g/mL) required to furnish the energy contained in 80.0 L (about 20 gal) of gasoline (density 5 0.740 g/mL). Calculate also the volume that this hydrogen would occupy as a gas at 1.00 atm and 258C.

Define the following terms: potential energy, kinetic energy, path-dependent function, state function, system, surroundings.

Consider the following potential energy diagrams for two different reactions. Reactants Products Potential energy Products Reactants Potential energy a b Which plot represents an exothermic reaction? In plot a, do the reactants on average have stronger or weaker bonds than the products? In plot b, reactants must gain potential energy to convert to products. How does this occur?

What is the first law of thermodynamics? How can a system change its internal energy, E? What are the sign conventions for thermodynamic quantities used in this text?

When a gas expands, what is the sign of w? Why? When a gas contracts, what is the sign of w? Why? What are the signs of q and w for the process of boiling water?

What is the heat gained/released at constant pressure equal to (qP 5 ?)? What is the heat gained/released at constant volume equal to (qV 5 ?)? Explain why DH is obtained directly from a coffee-cup calorimeter, whereas DE is obtained directly from a bomb calorimeter

High-quality audio amplifiers generate large amounts of heat. To dissipate the heat and prevent damage to the electronic components, heat-radiating metal fins are used. Would it be better to make these fins out of iron or aluminum? Why? (See Table 6.1 for specific heat capacities.)

Explain how calorimetry works to calculate DH or DE for a reaction. Does the temperature of the calorimeter increase or decrease for an endothermic reaction? For an exothermic reaction? Explain

What is Hesss law? When a reaction is reversed, what happens to the sign and magnitude of DH for that reversed reaction? When the coefficients in a balanced reaction are multiplied by a factor n, what happens to the sign and magnitude of DH for that multiplied reaction?

Define the standard enthalpy of formation. What are standard states for elements and for compounds? Using Hesss law, illustrate why the formula DHreaction 5 SnpDHf 1products2 2SnrDHf 1reactants2 works to calculate DH8 for a reaction.

What are some of the problems associated with the worlds dependence on fossil fuels? What are some alternative fuels for petroleum products?

Objects placed together eventually reach the same temperature. When you go into a room and touch a piece of metal in that room, it feels colder than a piece of plastic. Explain.

What is meant by the term lower in energy? Which is lower in energy, a mixture of hydrogen and oxygen gases or liquid water? How do you know? Which of the two is more stable? How do you know?

A fire is started in a fireplace by striking a match and lighting crumpled paper under some logs. Explain all the energy transfers in this scenario using the terms exothermic, endothermic, system, surroundings, potential energy, and kinetic energy in the discussion.

Liquid water turns to ice. Is this process endothermic or exothermic? Explain what is occurring using the terms system, surroundings, heat, potential energy, and kinetic energy in the discussion.

Consider the following statements: Heat is a form of energy, and energy is conserved. The heat lost by a system must be equal to the amount of heat gained by the surroundings. Therefore, heat is conserved. Indicate everything you think is correct in these statements. Indicate everything you think is incorrect. Correct the incorrect statements and explain.

Consider 5.5 L of a gas at a pressure of 3.0 atm in a cylinder with a movable piston. The external pressure is changed so that the volume changes to 10.5 L. a. Calculate the work done, and indicate the correct sign. b. Use the preceding data but consider the process to occur in two steps. At the end of the first step, the volume is 7.0 L. The second step results in a final volume of 10.5 L. Calculate the work done, and indicate the correct sign. c. Calculate the work done if after the first step the volume is 8.0 L and the second step leads to a volume of 10.5 L. Does the work differ from that in part b? Explain

In Question 6 the work calculated for the different conditions in the various parts of the question was different even though the system had the same initial and final conditions. Based on this information, is work a state function? a. Explain how you know that work is not a state function. b. Why does the work increase with an increase in the number of steps? c. Which two-step process resulted in more work, when the first step had the bigger change in volume or when the second step had the bigger change in volume? Explain.

Explain why oceanfront areas generally have smaller temperature fluctuations than inland areas.

Hesss law is really just another statement of the first law of thermodynamics. Explain

In the equation w 5 2PDV, why is there a negative sign?

Consider an airplane trip from Chicago, Illinois, to Denver, Colorado. List some path-dependent functions and some state functions for the plane trip.

How is average bond strength related to relative potential energies of the reactants and the products?

Assuming gasoline is pure C8H18(l), predict the signs of q and w for the process of combusting gasoline into CO2(g) and H2O(g).

What is the difference between DH and DE?

The enthalpy change for the reaction CH4 1g2 1 2O2 1g2 h CO2 1g2 1 2H2O1l2 is 2891 kJ for the reaction as written. a. What quantity of heat is released for each mole of water formed? b. What quantity of heat is released for each mole of oxygen reacted?

For the reaction HgO1s2 S Hg1l2 1 1 2O2 1g2, DH 5 190.7 kJ: a. What quantity of heat is required to produce 1 mole of mercury by this reaction? b. What quantity of heat is required to produce 1 mole of oxygen gas by this reaction? c. What quantity of heat would be released in the following reaction as written? 2Hg1l2 1 O2 1g2 h 2HgO1s

The enthalpy of combustion of CH4(g) when H2O(l) is formed is 2891 kJ/mol and the enthalpy of combustion of CH4(g) when H2O(g) is formed is 2803 kJ/mol. Use these data and Hesss law to determine the enthalpy of vaporization for water.

The enthalpy change for a reaction is a state function and it is an extensive property. Explain

Standard enthalpies of formation are relative values. What are DHf 8 values relative to?

The combustion of methane can be represented as follows: Reactants Step 1 (a) Step 2 (c) Elements Products Ha = 75 kJ (b) Hb = 0 kJ (d) Hd = 572 kJ Hc = 394 kJ a. Use the information given above to determine the value of DH for the combustion of methane to form CO2(g) and 2H2O(l). b. What is DHf for an element in its standard state? Why is this? Use the figure above to support your answer.c. How does DH for the reaction CO2(g) 1 2H2O(l) S CH4(g) 1 O2(g) compare to that of the combustion of methane? Why is this?

Why is it a good idea to rinse your thermos bottle with hot water before filling it with hot coffee?

Photosynthetic plants use the following reaction to produce glucose, cellulose, and so forth: 6CO2 1g2 1 6H2O1l2 88888n C6H12O6 1s2 1 6O2 1g2 How might extensive destruction of forests exacerbate the greenhouse effect?

What is incomplete combustion of fossil fuels? Why can this be a problem?

Explain the advantages and disadvantages of hydrogen as an alternative fuel.

Calculate the kinetic energy of a baseball (mass 5 5.25 oz) with a velocity of 1.0 3 102 mi/h.

Which has the greater kinetic energy, an object with a mass of 2.0 kg and a velocity of 1.0 m/s or an object with a mass of 1.0 kg and a velocity of 2.0 m/s?

Consider the following diagram when answering the questions below. A B A B Initial Final Held in place a. Compare balls A and B in terms of potential energy in both the initial and final setups. b. Ball A has stopped moving in the figure on the right above, but energy must be conserved. What happened to the potential energy of ball A?

Consider the accompanying diagram. Ball A is allowed to fall and strike ball B. Assume that all of ball As energy is transferred to ball B at point I, and that there is no loss of energy to other sources. What is the kinetic energy and the potential energy of ball B at point II? The potential energy is given by PE 5 mgz, where m is the mass in kilograms, g is the gravitational constant (9.81 m/s2 ), and z is the distance in meters.

A gas absorbs 45 kJ of heat and does 29 kJ of work. Calculate DE

A system releases 125 kJ of heat while 104 kJ of work is done on it. Calculate DE.

Calculate DE for each of the following. a. q 5 247 kJ, w 5 188 kJ b. q 5 182 kJ, w 5 247 kJ c. q 5 147 kJ, w 5 0 d. In which of these cases do the surroundings do work on the system?

A system undergoes a process consisting of the following two steps: Step 1: The system absorbs 72 J of heat while 35 J of work is done on it. Step 2: The system absorbs 35 J of heat while performing 72 J of work. Calculate DE for the overall process.

If the internal energy of a thermodynamic system is increased by 300. J while 75 J of expansion work is done, how much heat was transferred and in which direction, to or from the system?

Calculate the internal energy change for each of the following. a. One hundred (100.) joules of work is required to compress a gas. At the same time, the gas releases 23 J of heat. b. A piston is compressed from a volume of 8.30 L to 2.80 L against a constant pressure of 1.90 atm. In the process, there is a heat gain by the system of 350. J. c. A piston expands against 1.00 atm of pressure from 11.2 L to 29.1 L. In the process, 1037 J of heat is absorbed.

A sample of an ideal gas at 15.0 atm and 10.0 L is allowed to expand against a constant external pressure of 2.00 atm at a constant temperature. Calculate the work in units of kJ for the gas expansion. (Hint: Boyles law applies.)

A piston performs work of 210. L ? atm on the surroundings, while the cylinder in which it is placed expands from 10. L to 25 L. At the same time, 45 J of heat is transferred from the surroundings to the system. Against what pressure was the piston working?

Consider a mixture of air and gasoline vapor in a cylinder with a piston. The original volume is 40. cm3 . If the combustion of this mixture releases 950. J of energy, to what volume will the gases expand against a constant pressure of 650. torr if all the energy of combustion is converted into work to push back the piston?

As a system increases in volume, it absorbs 52.5 J of energy in the form of heat from the surroundings. The piston is working against a pressure of 0.500 atm. The final volume of the system is 58.0 L. What was the initial volume of the system if the internal energy of the system decreased by 102.5 J?

A balloon filled with 39.1 moles of helium has a volume of 876 L at 0.08C and 1.00 atm pressure. The temperature of the balloon is increased to 38.08C as it expands to a volume of 998 L, the pressure remaining constant. Calculate q, w, and DE for the helium in the balloon. (The molar heat capacity for helium gas is 20.8 J/C # mol.)

One mole of H2O(g) at 1.00 atm and 100.8C occupies a volume of 30.6 L. When 1 mole of H2O(g) is condensed to 1 mole of H2O(l) at 1.00 atm and 100.8C, 40.66 kJ of heat is released. If the density of H2O(l) at this temperature and pressure is 0.996 g/cm3 , calculate DE for the condensation of 1 mole of water at 1.00 atm and 100.8C.

One of the components of polluted air is NO. It is formed in the high-temperature environment of internal combustion engines by the following reaction: N2 1g2 1 O2 1g2 h 2NO1g2 DH 5 180 kJ Why are high temperatures needed to convert N2 and O2 to NO?

The reaction SO3 1g2 1 H2O1l2 h H2SO4 1aq2 is the last step in the commercial production of sulfuric acid. The enthalpy change for this reaction is 2227 kJ. In designing a sulfuric acid plant, is it necessary to provide for heating or cooling of the reaction mixture? Explain.

Are the following processes exothermic or endothermic? a. When solid KBr is dissolved in water, the solution gets colder. b. Natural gas (CH4) is burned in a furnace. c. When concentrated H2SO4 is added to water, the solution gets very hot. d. Water is boiled in a teakettle.

Are the following processes exothermic or endothermic? a. the combustion of gasoline in a car engine b. water condensing on a cold pipe c. CO2 1s2 h CO2 1g2 d. F2 1g2 h 2F1g2

The overall reaction in a commercial heat pack can be represented as 4Fe1s2 1 3O2 1g2 h 2Fe2O3 1s2 DH 5 21652 kJ a. How much heat is released when 4.00 moles of iron are reacted with excess O2? b. How much heat is released when 1.00 mole of Fe2O3 is produced? c. How much heat is released when 1.00 g iron is reacted with excess O2? d. How much heat is released when 10.0 g Fe and 2.00 g O2 are reacted?

Consider the following reaction: 2H2 1g2 1 O2 1g2 h 2H2O1l2 DH 5 2572 kJ a. How much heat is evolved for the production of 1.00 mole of H2O(l)? b. How much heat is evolved when 4.03 g hydrogen are reacted with excess oxygen? c. How much heat is evolved when 186 g oxygen are reacted with excess hydrogen?d. The total volume of hydrogen gas needed to fill the Hindenburg was 2.0 3 108 L at 1.0 atm and 258C. How much heat was evolved when the Hindenburg exploded, assuming all of the hydrogen reacted? 47. Consider the combustion of propane: C3H8 1g2 1 5O2 1g2 h 3CO2 1g2 1 4H2O1l2 DH 5 22221 kJ Assume that all the heat in Example 6.3 comes from the combustion of propane. What mass of propane must be burned to furnish this amount of energy assuming the heat transfer process is 60.% efficient?

Consider the following reaction: CH4 1g2 1 2O2 1g2 h CO2 1g2 1 2H2O1l2 DH 5 2891 kJ Calculate the enthalpy change for each of the following cases: a. 1.00 g methane is burned in excess oxygen. b. 1.00 3 103 L methane gas at 740. torr and 258C are burned in excess oxygen.

For the process H2O(l) h H2O(g) at 298 K and 1.0 atm, DH is more positive than DE by 2.5 kJ/mol. What does the 2.5 kJ/mol quantity represent?

For the following reactions at constant pressure, predict if DH . DE, DH , DE, or DH 5 DE. a. 2HF1g2 h H2 1g2 1 F2 1g2 b. N2 1g2 1 3H2 1g2 h 2NH3 1g2 c. 4NH3 1g2 1 5O2 1g2 h 4NO1g2 1 6H2O1g2

Consider the substances in Table 6.1. Which substance requires the largest amount of energy to raise the temperature of 25.0 g of the substance from 15.08C to 37.08C? Calculate the energy. Which substance in Table 6.1 has the largest temperature change when 550. g of the substance absorbs 10.7 kJ of energy? Calculate the temperature change

The specific heat capacity of silver is 0.24 J/8C ? g. a. Calculate the energy required to raise the temperature of 150.0 g Ag from 273 K to 298 K. b. Calculate the energy required to raise the temperature of 1.0 mole of Ag by 1.08C (called the molar heat capacity of silver). c. It takes 1.25 kJ of energy to heat a sample of pure silver from 12.08C to 15.28C. Calculate the mass of the sample of silver

A 5.00-g sample of one of the substances listed in Table 6.1 was heated from 25.28C to 55.18C, requiring 133 J to do so. Which substance was it?

It takes 585 J of energy to raise the temperature of 125.6 g mercury from 20.08C to 53.58C. Calculate the specific heat capacity and the molar heat capacity of mercury.

A 30.0-g sample of water at 280. K is mixed with 50.0 g water at 330. K. Calculate the final temperature of the mixture assuming no heat loss to the surroundings

A biology experiment requires the preparation of a water bath at 37.08C (body temperature). The temperature of the cold tap water is 22.08C, and the temperature of the hot tap water is 55.08C. If a student starts with 90.0 g cold water, what mass of hot water must be added to reach 37.08C?

A 5.00-g sample of aluminum pellets (specific heat capacity 5 0.89 J/8C ? g) and a 10.00-g sample of iron pellets (specific heat capacity 5 0.45 J/8C ? g) are heated to 100.08C. The mixture of hot iron and aluminum is then dropped into 97.3 g water at 22.08C. Calculate the final temperature of the metal and water mixture, assuming no heat loss to the surroundings.

Hydrogen gives off 120. J/g of energy when burned in oxygen, and methane gives off 50. J/g under the same circumstances. If a mixture of 5.0 g hydrogen and 10. g methane is burned, and the heat released is transferred to 50.0 g water at 25.08C, what final temperature will be reached by the water?

A 150.0-g sample of a metal at 75.08C is added to 150.0 g H2O at 15.08C. The temperature of the water rises to 18.38C. Calculate the specific heat capacity of the metal, assuming that all the heat lost by the metal is gained by the water.

A 110.-g sample of copper (specific heat capacity 5 0.20 J/8C ? g) is heated to 82.48C and then placed in a container of water at 22.38C. The final temperature of the water and copper is 24.98C. What is the mass of the water in the container, assuming that all the heat lost by the copper is gained by the water?

In a coffee-cup calorimeter, 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M HCl are mixed to yield the following reaction: Ag1 1aq2 1 Cl2 1aq2 h AgCl1s2 The two solutions were initially at 22.608C, and the final temperature is 23.408C. Calculate the heat that accompanies this reaction in kJ/mol of AgCl formed. Assume that the combined solution has a mass of 100.0 g and a specific heat capacity of 4.18 J/8C ? g.

In a coffee-cup calorimeter, 100.0 mL of 1.0 M NaOH and 100.0 mL of 1.0 M HCl are mixed. Both solutions were originally at 24.68C. After the reaction, the final temperature is 31.38C. Assuming that all the solutions have a density of 1.0 g/cm3 and a specific heat capacity of 4.18 J/8C ? g, calculate the enthalpy change for the neutralization of HCl by NaOH. Assume that no heat is lost to the surroundings or to the calorimeter.

A coffee-cup calorimeter initially contains 125 g water at 24.28C. Potassium bromide (10.5 g), also at 24.28C, is added to the water, and after the KBr dissolves, the final temperature is 21.18C. Calculate the enthalpy change for dissolving the salt in J/g and kJ/mol. Assume that the specific heat capacity of the solution is 4.18 J/8C ? g and that no heat is transferred to the surroundings or to the calorimeter.

In a coffee-cup calorimeter, 1.60 g NH4NO3 is mixed with 75.0 g water at an initial temperature of 25.008C. After dissolution of the salt, the final temperature of the calorimeter contents is 23.348C. Assuming the solution has a heat capacity of 4.18 J/8C ? g and assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of NH4NO3 in units of kJ/mol.

Consider the dissolution of CaCl2: CaCl2 1s2 h Ca21 1aq2 1 2Cl2 1aq2 DH 5 281.5 kJ An 11.0-g sample of CaCl2 is dissolved in 125 g water, with both substances at 25.08C. Calculate the final temperature of the solution assuming no heat loss to the surroundings and assuming the solution has a specific heat capacity of 4.18 J/8C ? g.

Consider the reaction 2HCl1aq2 1 Ba1OH2 2 1aq2 h BaCl2 1aq2 1 2H2O1l2 DH 5 2118 kJ Calculate the heat when 100.0 mL of 0.500 M HCl is mixed with 300.0 mL of 0.100 M Ba(OH)2. Assuming that the temperature of both solutions was initially 25.08C and that the final mixture has a mass of 400.0 g and a specific heat capacity of 4.18 J/8C ? g, calculate the final temperature of the mixture.

The heat capacity of a bomb calorimeter was determined by burning 6.79 g methane (energy of combustion 5 2802 kJ/ mol CH4) in the bomb. The temperature changed by 10.88C. a. What is the heat capacity of the bomb? b. A 12.6-g sample of acetylene, C2H2, produced a temperature increase of 16.98C in the same calorimeter. What is the energy of combustion of acetylene (in kJ/mol)?

The combustion of 0.1584 g benzoic acid increases the temperature of a bomb calorimeter by 2.548C. Calculate the heat capacity of this calorimeter. (The energy released by combustion of benzoic acid is 26.42 kJ/g.) A 0.2130-g sample of vanillin (C8H8O3) is then burned in the same calorimeter, and the temperature increases by 3.258C. What is the energy of combustion per gram of vanillin? Per mole of vanillin?

The enthalpy of combustion of solid carbon to form carbon dioxide is 2393.7 kJ/mol carbon, and the enthalpy of combustion of carbon monoxide to form carbon dioxide is 2283.3 kJ/ mol CO. Use these data to calculate DH for the reaction 2C1s2 1 O2 1g2 h 2CO1g2

Combustion reactions involve reacting a substance with oxygen. When compounds containing carbon and hydrogen are combusted, carbon dioxide and water are the products. Using the enthalpies of combustion for C4H4 (22341 kJ/mol), C4H8 (22755 kJ/mol), and H2 (2286 kJ/mol), calculate DH for the reaction C4H4 1g2 1 2H2 1g2 h C4H8 1g2

Given the following data (g) (g) H = 92 kJ (g) (g) H = 484 kJ N H O (g) (g) + + calculate DH for the reaction (g) + (g) (g) + (g)

On the basis of the enthalpy change, is this a useful reaction for the synthesis of ammonia?

Given the following data 2ClF1g2 1 O2 1g2 h Cl2O1g2 1 F2O1g2 DH 5 167.4 kJ 2ClF3 1g2 1 2O2 1g2 h Cl2O1g2 1 3F2O1g2 DH 5 341.4 kJ 2F2 1g2 1 O2 1g2 h 2F2O1g2 DH 5 243.4 kJ calculate DH for the reaction ClF1g2 1 F2 1g2 h ClF3 1g2

Given the following data 2O3 1g2 h 3O2 1g2 DH 5 2427 kJ O2 1g2 h 2O1g2 DH 5 1495 kJ NO1g2 1 O3 1g2 h NO2 1g2 1 O2 1g2 DH 5 2199 kJ calculate DH for the reaction NO1g2 1 O1g2 h NO2 1g

Calculate DH for the reaction N2H4 1l2 1 O2 1g2 h N2 1g2 1 2H2O1l2 given the following data: 2NH3 1g2 1 3N2O1g2 h 4N2 1g2 1 3H2O1l2 DH 5 21010. kJ N2O1g2 1 3H2 1g2 h N2H4 1l2 1 H2O1l2 DH 5 2317 kJ 2NH3 1g2 1 1 2O2 1g2 h N2H4 1l2 1 H2O1l2 DH 5 2143 kJ H2 1g2 1 1 2O2 1g2 h H2O1l2 DH 5 2286 kJ

Given the following data Ca1s2 1 2C1graphite2 h CaC2 1s2 DH 5 262.8 kJ Ca1s2 1 1 2O2 1g2 h CaO1s2 DH 5 2635.5 kJ CaO1s2 1 H2O1l2 h Ca1OH2 2 1aq2 DH 5 2653.1 kJ C2H2 1g2 1 5 2O2 1g2 h 2CO2 1g2 1 H2O1l2 DH 5 21300. kJ C1graphite2 1 O2 1g2 h CO2 1g2 DH 5 2393.5 kJ calculate DH for the reaction CaC2 1s2 1 2H2O1l2 h Ca1OH2 2 1aq2 1 C2H2 1g2

Given the following data P4 1s2 1 6Cl2 1g2 h 4PCl3 1g2 DH 5 21225.6 kJ P4 1s2 1 5O2 1g2 h P4O10 1s2 DH 5 22967.3 kJ PCl3 1g2 1 Cl2 1g2 h PCl5 1g2 DH 5 284.2 kJ PCl3 1g2 1 1 2 O2 1g2 h Cl3PO1g2 DH 5 2285.7 kJ calculate DH for the reaction P4O10 1s2 1 6PCl5 1g2 h 10Cl3PO1g2

Give the definition of the standard enthalpy of formation for a substance. Write separate reactions for the formation of NaCl, H2O, C6H12O6, and PbSO4 that have DH8 values equal to DHf 8 for each compound

Write reactions for which the enthalpy change will be a. DHf 8 for solid aluminum oxide. b. the standard enthalpy of combustion of liquid ethanol, C2H5OH(l). c. the standard enthalpy of neutralization of sodium hydroxide solution by hydrochloric acid.d. DHf 8 for gaseous vinyl chloride, C2H3Cl(g). e. the enthalpy of combustion of liquid benzene, C6H6(l). f. the enthalpy of solution of solid ammonium bromide.

Use the values of DHf 8 in Appendix 4 to calculate DH8 for the following reactions. a. (g) (g) (g) (g) (g) N H O C ++ + b. Ca3 1PO42 2 1s2 1 3H2SO4 1l2 h 3CaSO4 1s2 1 2H3PO4 1l2 c. NH3 1g2 1 HCl1g2 h NH4Cl1s

Use the values of DHf 8 in Appendix 4 to calculate DH8 for the following reactions. (See Exercise 79.) a. (l) ( + g) (g) + (g) b. SiCl4 1l2 1 2H2O1l2 h SiO2 1s2 1 4HCl1aq2 c. MgO1s2 1 H2O1l2 h Mg1OH2 2 1s2

The Ostwald process for the commercial production of nitric acid from ammonia and oxygen involves the following steps: 4NH3 1g2 1 5O2 1g2 h 4NO1g2 1 6H2O1g2 2NO1g2 1 O2 1g2 h 2NO2 1g2 3NO2 1g2 1 H2O1l2 h 2HNO3 1aq2 1 NO1g2 a. Use the values of DHf 8 in Appendix 4 to calculate the value of DH8 for each of the preceding reactions. b. Write the overall equation for the production of nitric acid by the Ostwald process by combining the preceding equations. (Water is also a product.) Is the overall reaction exothermic or endothermic?

Calculate DH8 for each of the following reactions using the data in Appendix 4: 4Na1s2 1 O2 1g2 h 2Na2O1s2 2Na1s2 1 2H2O1l2 h 2NaOH1aq2 1 H2 1g2 2Na1s2 1 CO2 1g2 h Na2O1s2 1 CO1g2 Explain why a water or carbon dioxide fire extinguisher might not be effective in putting out a sodium fire.

The reusable booster rockets of the space shuttle use a mixture of aluminum and ammonium perchlorate as fuel. A possible reaction is 3Al1s2 1 3NH4ClO4 1s2 h Al2O3 1s2 1 AlCl3 1s2 1 3NO1g2 1 6H2O1g2 Calculate DH8 for this reaction.

The space shuttle Orbiter utilizes the oxidation of methylhydrazine by dinitrogen tetroxide for propulsion: 4N2H3CH3 1l2 1 5N2O4 1l2 h 12H2O1g2 1 9N2 1g2 1 4CO2 1g2 Calculate DH8 for this reaction.

Consider the reaction 2ClF3 1g2 1 2NH3 1g2 h N2 1g2 1 6HF1g2 1 Cl2 1g2 DH 5 21196 kJ Calculate DH8f for ClF3(g).

The standard enthalpy of combustion of ethene gas, C2H4(g), is 21411.1 kJ/mol at 298 K. Given the following enthalpies of formation, calculate DHf 8 for C2H4(g). CO2 1g2 2393.5 kJ/mol H2O1l2 2285.8 kJ/mol

Water gas is produced from the reaction of steam with coal: C1s2 1 H2O1g2 h H2 1g2 1 CO1g2 Assuming that coal is pure graphite, calculate DH8 for this reaction

Syngas can be burned directly or converted to methanol. Calculate DH8 for the reaction CO1g2 1 2H2 1g2 h CH3OH1l2

Ethanol (C2H5OH) has been proposed as an alternative fuel. Calculate the standard enthalpy of combustion per gram of liquid ethanol

Methanol (CH3OH) has also been proposed as an alternative fuel. Calculate the standard enthalpy of combustion per gram of liquid methanol, and compare this answer to that for ethanol in Exercise 89.

Some automobiles and buses have been equipped to burn propane (C3H8). Compare the amounts of energy that can be obtained per gram of C3H8(g) and per gram of gasoline, assuming that gasoline is pure octane, C8H18(l). (See Example 6.11.) Look up the boiling point of propane. What disadvantages are there to using propane instead of gasoline as a fuel?

Acetylene (C2H2) and butane (C4H10) are gaseous fuels with enthalpies of combustion of 249.9 kJ/g and 249.5 kJ/g, respectively. Compare the energy available from the combustion of a given volume of acetylene to the combustion energy from the same volume of butane at the same temperature and pressure.

Assume that 4.19 3 106 kJ of energy is needed to heat a home. If this energy is derived from the combustion of methane (CH4), what volume of methane, measured at STP, must be burned? (DH8combustion for CH4 5 2891 kJ/mol)

The complete combustion of acetylene, C2H2(g), produces 1300. kJ of energy per mole of acetylene consumed. How many grams of acetylene must be burned to produce enough heat to raise the temperature of 1.00 gal water by 10.08C if the process is 80.0% efficient? Assume the density of water is 1.00 g/cm3.

It has been determined that the body can generate 5500 kJ of energy during one hour of strenuous exercise. Perspiration is the bodys mechanism for eliminating this heat. What mass of water would have to be evaporated through perspiration to rid the body of the heat generated during 2 hours of exercise? (The heat of vaporization of water is 40.6 kJ/mol.)

One way to lose weight is to exercise! Walking briskly at 4.0 miles per hour for an hour consumes about 400 kcal of energy. How many hours would you have to walk at 4.0 miles per hour to lose one pound of body fat? One gram of body fat is equivalent to 7.7 kcal of energy. There are 454 g in 1 lb.

Three gas-phase reactions were run in a constant-pressure piston apparatus as shown in the following illustration. For each reaction, give the balanced reaction and predict the sign of w (the work done) for the reaction. 1 atm S O a. 1 atm Cl C O b. 1 atm 1 atm O N c. 1 atm 1 atm If just the balanced reactions were given, how could you predict the sign of w for a reaction?

Nitrogen gas reacts with hydrogen gas to form ammonia gas . Consider the reaction between nitrogen and hydrogen as depicted below: 1 atm a. Draw what the container will look like after the reaction has gone to completion. Assume a constant pressure of 1 atm. b. Is the sign of work positive or negative, or is the value of work equal to zero for the reaction? Explain your answer

Combustion of table sugar produces CO2(g) and H2O(l). When 1.46 g table sugar is combusted in a constant-volume (bomb) calorimeter, 24.00 kJ of heat is liberated. a. Assuming that table sugar is pure sucrose, C12H22O11(s), write the balanced equation for the combustion reaction. b. Calculate DE in kJ/mol C12H22O11 for the combustion reaction of sucrose. c. Calculate DH in kJ/mol C12H22O11 for the combustion reaction of sucrose at 258C.

Consider the following changes: a. N2 1g2 h N2 1l2 b. CO1g2 1 H2O1g2 h H2 1g2 1 CO2 1g2 c. Ca3P2 1s2 1 6H2O1l2 h 3Ca1OH2 2 1s2 1 2PH3 1g2 d. 2CH3OH1l2 1 3O2 1g2 h 2CO2 1g2 1 4H2O1l2 e. I2 1s2 h I2 1g2 At constant temperature and pressure, in which of these changes is work done by the system on the surroundings? By the surroundings on the system? In which of them is no work done?

Consider the following cyclic process carried out in two steps on a gas: Step 1: 45 J of heat is added to the gas, and 10. J of expansion work is performed. Step 2: 60. J of heat is removed from the gas as the gas is compressed back to the initial state. Calculate the work for the gas compression in Step 2

Calculate DH8 for the reaction 2K1s2 1 2H2O1l2 h 2KOH1aq2 1 H2 1g2 A 5.00-g chunk of potassium is dropped into 1.00 kg water at 24.08C. What is the final temperature of the water after the preceding reaction occurs? Assume that all the heat is used to raise the temperature of the water. (Never run this reaction. It is very dangerous; it bursts into flame!)

The enthalpy of neutralization for the reaction of a strong acid with a strong base is 256 kJ/mol water produced. How much energy will be released when 200.0 mL of 0.400 M HNO3 is mixed with 150.0 mL of 0.500 M KOH?

When 1.00 L of 2.00 M Na2SO4 solution at 30.08C is added to 2.00 L of 0.750 M Ba(NO3)2 solution at 30.08C in a calorimeter, a white solid (BaSO4) forms. The temperature of the mixture increases to 42.08C. Assuming that the specific heat capacity of the solution is 6.37 J/8C ? g and that the density of the final solution is 2.00 g/mL, calculate the enthalpy change per mole of BaSO4 formed

If a student performs an endothermic reaction in a calorimeter, how does the calculated value of DH differ from the actual value if the heat exchanged with the calorimeter is not taken into account?

In a bomb calorimeter, the reaction vessel is surrounded by water that must be added for each experiment. Since the amount of water is not constant from experiment to experiment, the mass of water must be measured in each case. The heat capacity of the calorimeter is broken down into two parts: the water and the calorimeter components. If a calorimeter contains 1.00 kg water and has a total heat capacity of 10.84 kJ/8C, what is the heat capacity of the calorimeter components?

The bomb calorimeter in Exercise 106 is filled with 987 g water. The initial temperature of the calorimeter contents is 23.328C. A 1.056-g sample of benzoic acid (DEcomb 5 226.42 kJ/g) is combusted in the calorimeter. What is the final temperature of the calorimeter contents?

Consider the two space shuttle fuel reactions in Exercises 83 and 84. Which reaction produces more energy per kilogram of reactant mixture (stoichiometric amounts)?

Consider the following equations: 3A 1 6B h 3D DH 5 2403 kJ/mol E 1 2F h A DH 5 2105.2 kJ/mol C h E 1 3D DH 5 164.8 kJ/mol Suppose the first equation is reversed and multiplied by 1 6, the second and third equations are divided by 2, and the three adjusted equations are added. What is the net reaction and what is the overall heat of this reaction?

Given the following data Fe2O3 1s2 1 3CO1g2 h 2Fe1s2 1 3CO2 1g2 DH 5 223 kJ 3Fe2O3 1s2 1 CO1g2 h 2Fe3O4 1s2 1 CO2 1g2 DH 5 239 kJ Fe3O4 1s2 1 CO1g2 h 3FeO1s2 1 CO2 1g2 DH 5 118 kJ calculate DH8 for the reaction FeO1s2 1 CO1g2 h Fe1s2 1 CO2 1g2

At 298 K, the standard enthalpies of formation for C2H2(g) and C6H6(l) are 227 kJ/mol and 49 kJ/mol, respectively. a. Calculate DH8 for C6H6 1l2 h 3C2H2 1g2 b. Both acetylene (C2H2) and benzene (C6H6) can be used as fuels. Which compound would liberate more energy per gram when combusted in air?

Using the following data, calculate the standard heat of formation of ICl(g) in kJ/mol: Cl2 1g2 h 2Cl1g2 DH 5 242.3 kJ I2 1g2 h 2I1g2 DH 5 151.0 kJ ICl1g2 h I1g2 1 Cl1g2 DH 5 211.3 kJ I2 1s2 h I2 1g2 DH 5 62.8 kJ

A sample of nickel is heated to 99.88C and placed in a coffeecup calorimeter containing 150.0 g water at 23.58C. After the metal cools, the final temperature of metal and water mixture is 25.08C. If the specific heat capacity of nickel is 0.444 J/8C ? g, what mass of nickel was originally heated? Assume no heat loss to the surroundings.

Quinone is an important type of molecule that is involved in photosynthesis. The transport of electrons mediated by quinone in certain enzymes allows plants to take water, carbon dioxide, and the energy of sunlight to create glucose. A 0.1964-g sample of quinone (C6H4O2) is burned in a bomb calorimeter with a heat capacity of 1.56 kJ/8C. The temperature of the calorimeter increases by 3.28C. Calculate the energy of combustion of quinone per gram and per mole

Calculate DH8 for each of the following reactions, which occur in the atmosphere. a. C2H4 1g2 1 O3 1g2 h CH3CHO1g2 1 O2 1g2 b. O3 1g2 1 NO1g2 h NO2 1g2 1 O2 1g2 c. SO3 1g2 1 H2O1l2 h H2SO4 1aq2 d. 2NO1g2 1 O2 1g2 h 2NO2 1g2

Consider a balloon filled with helium at the following conditions. 313 g He 1.00 atm 1910. L Molar Heat Capacity 5 20.8 J/8C ? mol The temperature of this balloon is decreased by 41.68C as the volume decreases to 1643 L, with the pressure remaining constant. Determine q, w, and DE (in kJ) for the compression of the balloon.

In which of the following systems is(are) work done by the surroundings on the system? Assume pressure and temperature are constant. a. 2SO2 1g2 1 O2 1g2 h 2SO3 1g2 b. CO2 1s2 h CO2 1g2 c. 4NH3 1g2 1 7O2 1g2 h 4NO2 1g2 1 6H2O1g2 d. N2O4 1g2 h 2NO2 1g2 e. CaCO3 1s2 h CaCO1s2 1 CO2 1g

Which of the following processes are exothermic? a. N2 1g2 h 2N1g2 b. H2O1l2 h H2O1s2 c. Cl2 1g2 h 2Cl1g2 d. 2H2 1g2 1 O2 1g2 h 2H2O1g2 e. O2 1g2 h 2O1g2

Consider the reaction B2H6 1g2 1 3O2 1g2 h B2O3 1s2 1 3H2O1g2 DH 5 22035 kJ Calculate the amount of heat released when 54.0 g of diborane is combusted.

A swimming pool, 10.0 m by 4.0 m, is filled with water to a depth of 3.0 m at a temperature of 20.28C. How much energy is required to raise the temperature of the water to 24.68C?

In a coffee-cup calorimeter, 150.0 mL of 0.50 M HCl is added to 50.0 mL of 1.00 M NaOH to make 200.0 g solution at an initial temperature of 48.28C. If the enthalpy of neutralization for the reaction between a strong acid and a strong base is 256 kJ/mol, calculate the final temperature of the calorimeter contents. Assume the specific heat capacity of the solution is 4.184 J/g ? 8C and assume no heat loss to the surroundings.

Calculate DH for the reaction N2H4 1l2 1 O2 1g2 h N2 1g2 1 2H2O1l2 given the following data: Equation DH (kJ) 2NH3 1g2 1 3N2O1g2 h 4N2 1g2 1 3H2O1l2 21010 N2O1g2 1 3H2 1g2 h N2H4 1l2 1 H2O1l2 2317 2NH3 1g2 1 1 2 O2 1g2 h N2H4 1l2 1 H2O1l2 2143 H2 1g2 1 1 2 O2 1g2 h H2O1l2 2286

Which of the following substances have an enthalpy of formation equal to zero? a. Cl2(g) b. H2(g) c. N2(l) d. Cl(g)

Consider 2.00 moles of an ideal gas that are taken from state A (PA 5 2.00 atm, VA 5 10.0 L) to state B (PB 5 1.00 atm, VB 5 30.0 L) by two different pathways: a VC 5 30.0 L PC 5 2.00 atmb a VA 5 10.0 L PA 5 2.00 atmb a VB 5 30.0 L PB 5 1.00 atmb a VD 5 10.0 L PD 5 1.00 atmb These pathways are summarized on the following graph of P versus V: P (atm) 0 V (L) 1 2 10 20 30 A B 1 C D 4 3 2 Calculate the work (in units of J) associated with the two pathways. Is work a state function? Explain.

Calculate w and DE when 1 mole of a liquid is vaporized at its boiling point (80.8C) and 1.00 atm pressure. DHvap for the liquid is 30.7 kJ/mol at 80.8C.

The sun supplies energy at a rate of about 1.0 kilowatt per square meter of surface area (1 watt 5 1 J/s). The plants in an agricultural field produce the equivalent of 20. kg sucrose (C12H22O11) per hour per hectare (1 ha 5 10,000 m2 ). Assuming that sucrose is produced by the reaction 12CO2 1g2 1 11H2O1l2 h C12H22O11 1s2 1 12O2 1g2 DH 5 5640 kJ calculate the percentage of sunlight used to produce the sucrosethat is, determine the efficiency of photosynthesis

The best solar panels currently available are about 19% efficient in converting sunlight to electricity. A typical home will use about 40. kWh of electricity per day (1 kWh 5 1 kilowatt hour; 1 kW 5 1000 J/s). Assuming 8.0 hours of useful sunlight per day, calculate the minimum solar panel surface area necessary to provide all of a typical homes electricity. (See Exercise 126 for the energy rate supplied by the sun.)

On Easter Sunday, April 3, 1983, nitric acid spilled from a tank car near downtown Denver, Colorado. The spill was neutralized with sodium carbonate: 2HNO3 1aq2 1 Na2CO3 1s2 h 2NaNO3 1aq2 1 H2O1l2 1 CO2 1g2 a. Calculate DH8 for this reaction. Approximately 2.0 3 104 gal nitric acid was spilled. Assume that the acid was an aqueous solution containing 70.0% HNO3 by mass with a density of 1.42 g/cm3 . What mass of sodium carbonate was required for complete neutralization of the spill, and what quantity of heat was evolved? (DHf 8 for NaNO3(aq) 5 2467 kJ/mol) b. According to The Denver Post for April 4, 1983, authorities feared that dangerous air pollution might occur during the neutralization. Considering the magnitude of DH8, what was their major concern?

A piece of chocolate cake contains about 400 Calories. A nutritional Calorie is equal to 1000 calories (thermochemical calories), which is equal to 4.184 kJ. How many 8-in-high steps must a 180-lb man climb to expend the 400 Cal from the piece of cake? See Exercise 28 for the formula for potential energy.

The standard enthalpy of formation of H2O(l) at 298 K is 2285.8 kJ/mol. Calculate the change in internal energy for the following process at 298 K and 1 atm: H2O1l2 h H2 1g2 1 1 2O2 1g2 DE 5 ? (Hint: Using the ideal gas equation, derive an expression for work in terms of n, R, and T.)

You have a 1.00-mole sample of water at 230.8C and you heat it until you have gaseous water at 140.8C. Calculate q for the entire process. Use the following data. Specific heat capacity of ice 5 2.03 J/C # g Specific heat capacity of water 5 4.18 J/C # g Specific heat capacity of steam 5 2.02 J/C # g H2O1s2 h H2O1l2 DHfusion 5 6.02 kJ/mol 1at 0C2 H2O1l2 h H2O1g2 DHvaporization 5 40.7 kJ/mol 1at 100.C2

A 500.0-g sample of an element at 1958C is dropped into an icewater mixture; 109.5 g ice melts and an icewater mixture remains. Calculate the specific heat of the element. See Exercise 131 for pertinent information.

The preparation of NO2(g) from N2(g) and O2(g) is an endothermic reaction: N2 1g2 1 O2 1g2 h NO2 1g2 1unbalanced2 The enthalpy change of reaction for the balanced equation (with lowest whole-number coefficients) is DH 5 67.7 kJ. If 2.50 3 102 mL N2(g) at 100.8C and 3.50 atm and 4.50 3 102 mL O2(g) at 100.8C and 3.50 atm are mixed, what amount of heat is necessary to synthesize the maximum yield of NO2(g)?

Nitromethane, CH3NO2, can be used as a fuel. When the liquid is burned, the (unbalanced) reaction is mainly CH3NO2 1l2 1 O2 1g2 h CO2 1g2 1 N2 1g2 1 H2O1g2 a. The standard enthalpy change of reaction (DH8rxn) for the balanced reaction (with lowest whole-number coefficients) is 21288.5 kJ. Calculate DHf 8 for nitromethane. b. A 15.0-L flask containing a sample of nitromethane is filled with O2 and the flask is heated to 100.C. At this temperature, and after the reaction is complete, the total pressure of all the gases inside the flask is 950. torr. If the mole fraction of nitrogen (xnitrogen) is 0.134 after the reaction is complete, what mass of nitrogen was produced?

A cubic piece of uranium metal (specific heat capacity 5 0.117 J/8C ? g) at 200.08C is dropped into 1.00 L deuterium oxide (heavy water, specific heat capacity 5 4.211 J/8C ? g) at 25.58C. The final temperature of the uranium and deuterium oxide mixture is 28.58C. Given the densities of uranium (19.05 g/cm3 ) and deuterium oxide (1.11 g/mL), what is the edge length of the cube of uranium?

Consider a sample containing 5.00 moles of a monatomic ideal gas that is taken from state A to state B by the following two pathways: Pathway one: PA 5 3.00 atm 1 PC 5 3.00 atm VA 5 15.0 L 88n VC 5 55.0 L 2 PB 5 6.00 atm 88n VB 5 20.0 L Pathway two: PA 5 3.00 atm 3 PD 5 6.00 atm VA 5 15.0 L 88n VD 5 15.0 L 4 PB 5 6.00 atm 88n VB 5 20.0 L For each step, assume that the external pressure is constant and equals the final pressure of the gas for that step. Calculate q, w, DE, and DH for each step in kJ, and calculate overall values for each pathway. Explain how the overall values for the two pathways illustrate that DE and DH are state functions, whereas q and w are path functions. Hint: In a more rigorous study of thermochemistry, it can be shown that for an ideal gas: DE 5 nCvDT and DH 5 nCpDT where Cv is the molar heat capacity at constant volume and Cp is the molar heat capacity at constant pressure. In addition, for a monotomic ideal gas, Cv 5 3 2R and Cp 5 5 2R

A gaseous hydrocarbon reacts completely with oxygen gas to form carbon dioxide and water vapor. Given the following data, determine DHf 8 for the hydrocarbon: DHreaction 5 22044.5 kJ/mol hydrocarbon DHf 1CO22 5 2393.5 kJ/mol DHf 1H2O2 5 2242 kJ/mol Density of CO2 and H2O product mixture at 1.00 atm, 200.8C 5 0.751g/L. The density of the hydrocarbon is less than the density of Kr at the same conditions.