Chapter 10: Thermochemistry

10.1 Energy Basics [Go to section 10.1]

  1. If 14.5 kJ of heat were added to 485 g of liquid water (specific heat of water is 4.184 J/g °C), how much would its temperature increase?
  2. Calculate the heat capacity, in joules and in calories per degree, of the following:
    1. 28.4 g of water
    2. 1.00 oz of lead
  3. How much heat, in kilojoules and in kilocalories, must be added to a 75.0g iron block with a specific heat of 0.449 J/g °C to increase its temperature from 25 °C to its melting temperature of 1535 °C?
  4. How much would the temperature of 275 g of water increase if 36.5 kJ of heat were added?
  5. A piece of unknown substance weighs 44.7 g and requires 2110 J to increase its temperature from 23.2 °C to 89.6 °C
    1. What is the specific heat of the substance?
    2. If it is one of the substances found in Table 10.1.1, what is its likely identity?
  6. An iron kettle weighs 0.96 kg.
    1. What is the heat capacity of the kettle?
    2. How much heat is required to increase the temperature of this kettle from 21.0 °C to 100.0 °C?
    3. How much heat is required to heat this kettle from 21.0 °C to 100.0 °C if it contains 2.05 L of water (density of 0.997 g/mL and a specific heat of 4.184 J/g °C)?

[reveal-answer q=”11″]Show Selected Solutions[/reveal-answer]
[hidden-answer a=”11″]

  1. 7.15 °C
  2. 50.8 kJ and 12.2 Calories
  3. The answers are as follows:
    1. 0.711 J/g °C
    2. silicon


[/hidden-answer]

10.2 Calorimetry [Go to section 10.2]

  1. An “instant ice pack” utilizes the dissolution of solid ammonium nitrate in water. When this dissolution occurs the solution grows colder. (When 4.00 g of solid [latex]\ce{NH4NO3}[/latex] dissolves in 50.0 g of water at 23.1 °C in a calorimeter, the temperature decreases to 19.4 °C.)How would the amount of heat absorbed by the dissolution change if the heat capacity of the calorimeter were taken into account?
  2. Would the amount of heat absorbed by the dissolution in problem 9 appear greater, lesser, or remain the same if the experimenter used a calorimeter that was a poorer insulator than a coffee cup calorimeter? Explain your answer.
  3. Water is used as a coolant for hot engines. Water coming from a hot engine has a temperature of 250 °F then drops to 170 °F after passing through the radiator. What amount of heat was transferred from the engine to the surroundings by 1.25 gallons of water with a specific heat of 4.184 J/g °C.
  4. How much cold water (temperature 15 °C), in mL, must I add to 180 mL of hot coffee (temperature 98 °C) to cool it to 59 °C? Assume that coffee and water have a density of 1.0 g/mL and the same specific heat.
  5. Adding [latex]\ce{Ba(OH)2}[/latex] to [latex]\ce{NH4SCN}[/latex] in a calorimeter is an endothermic reaction which can represented by the following equation [latex]\ce{Ba(OH)2} + \ce{2NH4SCN} \longrightarrow \ce{Ba(SCN)2} + \ce{2NH4OH}[/latex]. The temperature of 100. g of water in a calorimeter drops by 3.00 °C if we add 2.95 g of [latex]\ce{Ba(OH)2}[/latex] to 1.41 g of [latex]\ce{NH4SCN}[/latex]. Assuming the specific heat of the solution and products is 4.20 J/g °C, calculate the approximate amount of heat absorbed by the reaction.
  6. A 56-g silver spoon (specific heat 0.235J/g °C) at 24 °C is placed in 180 mL (180 g) of tea at 76 °C, heat transfers from one to the other until the temperature of the two become equal.
    1. What is the final temperature when the two become equal? Assume that tea has the same specific heat as water.
    2. Another student solved this problem and got an answer of 18 °C. Explain why this is clearly an incorrect answer.
  7. From question 9 what if the 4.00 g of [latex]\ce{NH4NO3}[/latex] were dissolved in 150.0 g of water under the same conditions, how much would the temperature change? Is this answer what you expected and if so, why?
  8. A 43.2g piece of metal at 90.0 °C is placed in 100 g of water at 22.0 °C contained in a calorimeter like that shown in Figure 5.12. At thermal equilibrium the temperature of the water is 24.6 °C. How much heat did the metal give up to the water? What is the identity of the metal?
  9. When a 2.450-g sample of toluene, [latex]\ce{C7H8}[/latex], is burned in a bomb calorimeter, the temperature increases from 23.4 °C to 25.8 °C. The heat capacity of the calorimeter is 435 J/°C, and it contains 570. mL of water. How much heat was produced by the combustion of the toluene sample?
  10. A 0.650 g sample of [latex]\ce{NaBr}[/latex] is added to 75.0 g of water in a calorimeter. The temperature of the solution in the calorimeter decreases by 1.00 °C, what is the amount of heat involved in dissolving [latex]\ce{NaBr}[/latex]? (assume the heat capacity of the solution is 4.18 J/g °C) Is this reaction endothermic or exothermic and how do you know?
  11. If the amount of protein recommended in a 2000 Calorie daily diet is 100 g, then what percent of calories in this diet would be supplied by protein if the average number of Calories for protein is 4.0 Calories/g?
  12. When 50.0 g of 0.200 M [latex]\ce{NaCl}[/latex](aq) at 24.1 °C is added to 100.0 g of 0.100 M [latex]\ce{AgNO3}[/latex](aq) at 24.1 °C in a calorimeter, the temperature increases to 25.2 °C as [latex]\ce{AgCl}[/latex](s) forms. Assuming the specific heat of the solution and products is 4.20 J/g °C, calculate the approximate amount of heat in joules produced.
  13. If an 8 ounce serving of diet cola has 2 Calories then what is the mass of carbohydrates per ounce in the diet cola if a carbohydrate contains 4.1 Calories/gram. Assume a density of 1.0 g/mL for the cola.
  14. The reaction of 50 mL of acid and 50 mL of base described in Example 10.2.3 increased the temperature of the solution by 6.9 degrees. How much would the temperature have increased if 100 mL of acid and 100 mL of base had been used in the same calorimeter starting at the same temperature of 22.0 °C? Explain your answer.
  15. Tator Tot Hotdish contains 20 g of carbohydrates, 4 grams of protein and 10 grams of fat per serving. What is the Calorie content of a serving of Tater Tot Hotdish if the Calories for fat is 9.1 Calories/g, for carbohydrates is 4.1 Calories/g, and for protein is 4.1 Calories/g?
  16. When 1.0 g of fructose, [latex]\ce{C6H_{12}O6}[/latex](s), a sugar commonly found in fruits, is burned in oxygen in a bomb calorimeter, the temperature of the calorimeter increases by 1.58 °C. If the heat capacity of the calorimeter and its contents is 9.90 kJ/°C, what is q for this combustion?
  17. Energy for your home is often produced via electric generator. The generator is driven by steam produced from various materials, one of which is coal. 2.00 grams of coal are burned in a bomb calorimeter with a heat capacity of 12.3 kJ/°C and the temperature increases by 1.65 °C. How many pounds of coal must be fed into the burner to produce 21,000 kJ?
  18. A teaspoon of the carbohydrate sucrose (common sugar) contains 16 Calories (16 kcal). What is the mass of one teaspoon of sucrose if the average number of Calories for carbohydrates is 4.1 Calories/g?
  19. What is the mass of fat, in both grams and pounds if a can of crisco containing 6300 Calories? There are 9.1 Calories/g in fat.
  20. Which is the least expensive source of energy in kilojoules per dollar: a box of breakfast cereal that weighs 32 ounces and costs $4.23, or a liter of isooctane (density, 0.6919 g/mL) that costs $0.45? Compare the nutritional value of the cereal with the heat produced by combustion of the isooctane under standard conditions. A 1.0-ounce serving of the cereal provides 130 Calories.

[reveal-answer q=”11″]Show Selected Solutions[/reveal-answer]
[hidden-answer a=”11″]

  1. The amount of heat absorbed by the dissolution would be larger if the heat capacity of the calorimeter were taken into account. Since the calorimeter would also contribute some amount of heat to the dissolution. We would see not only the water/solution grow colder but the calorimeter as well.
  2. 5.3 × 102 kJ
  3. 1.26 kJ
  4. 1.2 °C since the mass and the heat capacity of the solution is approximately equal to that of the water, the two-fold increase in the amount of water leads to a two-fold decrease of the temperature change.
  5. 6.77 kJ
  6. 20%
  7. 0.061 g/ounce
  8. 1.9 × 102 Calories
  9. 4.56 pounds
  10. 692 grams or 1.52 pounds.


[/hidden-answer]

10.3 Enthalpy [Go to section 10.3]

  1. If dissolving 100 grams of ammonium chloride produces 27.69 kJ of heat then calculate the enthalpy of solution (ΔH for the dissolution) per mole of [latex]\ce{NH4Cl}[/latex].
  2. Using the data in the check your learning section of Example 10.2.3, calculate ΔH in kJ/mol of [latex]\ce{AgNO3}[/latex](aq) for the reaction: [latex]\ce{NaCl}(aq)+\ce{AgNO3}(aq)\longrightarrow\ce{AgCl}(s)+\ce{NaNO3}(aq)[/latex].
  3. The combustion of methane reaction is as follows: [latex]\ce{CH4}(g) + \ce{2O2}(g) \longrightarrow \ce{CO2}(g) + \ce{2H2O}(g) \text{ } \Delta \text{H } = -74.48[/latex] kJ/mol. A 1.347 g sample of methane is burned a bomb calorimeter with an oxygen environment. The temperature of the bomb calorimeter went from 24.32 °C to 25.46 °C. What is the heat capacity of the calorimeter and its contents?
  4. Calculate ΔH for the reaction described by the equation. [latex]\ce{Ba(OH)2}\cdot \ce{8H2O}(s)+\ce{2NH4SCN}(aq)\longrightarrow\ce{Ba(SCN)2}(aq)+\ce{2NH3}(aq)+\ce{10H2O}(l)[/latex]
  5. Radiators use hot water flowing through them to heat a house. What mass of water would provide the same quantity of heat when cooled from 97.0 to 28.0 °C, as the heat provided when 80 g of steam is cooled from 120 °C to 100 °C.
  6. How much heat is produced by combustion of 125 g of methanol under standard state conditions?
  7. Using the first two reactions determine the ΔH for the third reaction.
    [latex]\begin{array}{lr} \ce{P4}(s) + 6\ce{Cl2}(g) \longrightarrow \ce{4PCl3}(g) & \Delta H = -2439 \text{ kJ} \\ \ce{4PCl5}(g) \longrightarrow \ce{P4}(s) + \ce{10Cl2}(g) & \Delta H = 3438 \text{ kJ} \\ \ce{PCl5}(g) \longrightarrow \ce{PCl3}(g) + \ce{Cl2}(g) & \Delta H = ?? \end{array}[/latex]
  8. What mass of carbon monoxide must be burned to produce 175 kJ of heat under standard state conditions?
  9. Nitroglycerine is dangerous because its decomposition reaction is extremely fast and it produces several moles of gas which expand quickly. The reaction is as follows:[latex]\ce{4C3H5N3O9}(l) \longrightarrow \ce{12CO2}(g) + \ce{10H2O}(g) + \ce{6N2}(g) + \ce{O2}(g)[/latex] ΔH = -5678 KJWhat is the energy released per pound of nitroglycerine? If we calculated the standard enthalpy of formation for nitroglycerine what would we except the ΔH formation for [latex]\ce{O2}[/latex] to be?
  10. How much heat is produced when 100 mL of 0.250 M [latex]\ce{HCl}[/latex] (density, 1.00 g/mL) and 200 mL of 0.150 M [latex]\ce{NaOH}[/latex] (density, 1.00 g/mL) are mixed?[latex]\ce{HCl}(aq)+\ce{NaOH}(aq)\longrightarrow\ce{NaCl}(aq)+\ce{H2O}(l)\Delta{H}_{298}^{\circ }=-58\text{ kJ}[/latex] If both solutions are at the same temperature and the heat capacity of the products is 4.19 J/g °C, how much will the temperature increase? What assumption did you make in your calculation?
  11. Johnny prepares terrain for his Mazes and Monsters game by pouring dental plaster into dungeon molds. The plaster is comprised of [latex]\ce{CaSO4}[/latex][latex]\cdot[/latex][latex]\ce{4H2O}[/latex] in a specific crystal form and as the plaster sets it forms [latex]\ce{CaSO4}[/latex][latex]\cdot[/latex][latex]\ce{2H2O}[/latex] in a different crystal structure. A mole of dental plaster produces 2875 kJ of heat. If Johnny mixes 25.0 grams of dental plaster with water for his next set of molds how much heat is evolved?
  12. Before the introduction of chlorofluorocarbons, sulfur dioxide (enthalpy of vaporization, 6.00 kcal/mol) was used in household refrigerators. What mass of [latex]\ce{SO2}[/latex] must be evaporated to remove as much heat as evaporation of 1.00 kg of [latex]\ce{CCl2F2}[/latex] (enthalpy of vaporization is 17.4 kJ/mol)? The vaporization reactions for [latex]\ce{SO2}[/latex] and [latex]\ce{CCl2F2}[/latex] are [latex]\ce{SO2}(l)\longrightarrow\ce{SO2}(g)[/latex] and [latex]\ce{CCl2F}(l)\longrightarrow\ce{CCl2F2}(g)[/latex], respectively.
  13. Acetylene ([latex]\ce{C2H2}[/latex]) is used as a fuel in torches for welding. Acetylene gas reacts (combusts) with oxygen gas to produce [latex]\ce{CO2}[/latex] gas and water vapor. Write out the balanced reaction equation showing the reaction of acetylene with oxygen and its products. If burning 10.0 grams of acetylene produces 499.5 joules of heat then what is the enthalpy of the balanced combustion of acetylene reaction?
  14. Joseph Priestly prepared oxygen in 1774 by heating red mercury(II) oxide with sunlight focused through a lens. How much heat is required to decompose exactly 1 mole of red [latex]\ce{HgO}[/latex](s) to [latex]\ce{Hg}[/latex](l) and [latex]\ce{O2}[/latex](g) under standard conditions?
  15. You are investigating an alternative fuel for combustion engines in cars. Octane ([latex]\ce{C8H18}[/latex], ΔH°formation = -250kJ/mol) is the current fuel, and it reacts with oxygen to produce [latex]\ce{CO2}[/latex] and [latex]\ce{H2O}[/latex]. One alternative fuel you are looking at is liquid hydrazine ([latex]\ce{N2H4}[/latex]) which when it reacts with oxygen gas produces [latex]\ce{N2}(g)[/latex] and [latex]\ce{H2O}(g)[/latex]. Using standard ΔH found in Standard Thermodynamic Properties for Selected Substances calculate the heat produced from 1.00 gram of octane and from 1.00 gram of hydrazine. Which would be a better fuel source for your car? Is the amount of heat produced the only thing you would need to consider using this as an alternative fuel source for automobiles?
  16. How many kilojoules of heat will be released when exactly 1 mole of iron, [latex]\ce{Fe}[/latex], is burned to form [latex]\ce{Fe2O3}[/latex](s) at standard state conditions?
  17. Ethene, [latex]\ce{C2H4}[/latex] is a byproduct from the fractional distillation of petroleum. Synthetic ethanol can be manufactured by the hydration (reaction with water) of ethene with steam. Balance the reaction of ethene and steam to make ethanol vapor. Using the data in the table in Standard Thermodynamic Properties for Selected Substances, calculate ΔH° for the reaction.
  18. Both graphite and diamond burn. [latex]\ce{C}(s,\text{diamond})+\ce{O2}(g)\longrightarrow\ce{CO2}(g)[/latex] For the conversion of graphite to diamond:[latex]\ce{C}(s,\text{graphite})\longrightarrow\ce{C}(s,\text{diamond})\Delta{H}_{298}^{\circ }=1.90\text{ kJ}[/latex] Which produces more heat, the combustion of graphite or the combustion of diamond?
  19. Butane is a hydrocarbon that is sometimes used as a fuel.
    1. What is the balanced equation for the reaction of butane with oxygen and what are the products?
    2. What volume of air would be required to combust 20.00 grams of butane. Assume standard temperature and pressure. Air is 21.0 percent oxygen by volume. (Hint: use the information that 1.00 L of air at 25 °C and 1.00 atm contains 0.275 g of [latex]\ce{O2}[/latex] per liter.)
    3. The heat of combustion of butane is -2,657 kJ/mol. Assuming that all of the heat released in burning 20.0 grams of butane is transferred to 3.00 kilograms of water, calculate the increase in temperature of the water.
  20. Which produces more heat? [latex]\ce{Os}(s)\longrightarrow \ce{2O2}(g)\longrightarrow\ce{OsO4}(s)[/latex] or [latex]\ce{Os}(s)\longrightarrow \ce{2O2}(g)\longrightarrow\ce{OsO4}(g)[/latex] for the phase change [latex]{\ce{OsO}}_{4}\left(s\right)\longrightarrow{\ce{OsO}}_{4}\left(g\right)\Delta\text{H}=56.4\text{ kJ}[/latex]
  21. Find the ΔH of the following reaction: [latex]\ce{C}[/latex](s, gr) + [latex]\ce{O2}(\text{g}) \longrightarrow \ce{CO2}(\text{g})[/latex]. Given the following data:
    [latex]\ce{SrO}(\text{s}) + \ce{CO2}(\text{g}) \longrightarrow \ce{SrCO3}(\text{s})[/latex] ΔH = −234 kJ
    [latex]2\ce{SrO}(\text{s}) \longrightarrow 2\ce{Sr}(\text{s}) + \ce{O2}(\text{g})[/latex] ΔH = +1184 kJ
    [latex]\ce{2SrCO3}(\text{s}) \longrightarrow 2\ce{Sr}(\text{s}) + 2\ce{C}(\text{s, gr}) + 3\ce{O2}(\text{g})[/latex] ΔH = +2440 kJ
  22. Using the data in Standard Thermodynamic Properties for Selected Substances, calculate the standard enthalpy change for each of the following reactions:
    1. [latex]\ce{N2}(g)+\ce{O2}(g)\longrightarrow \ce{2NO}(g)[/latex]
    2. [latex]\ce{Si}(s)+\ce{2Cl2}(g)\longrightarrow\ce{SiCl4}(g)[/latex]
    3. [latex]\ce{Fe2O}(s)+\ce{3H2}(g)\longrightarrow \ce{2Fe}(s)+\ce{3H2O}(l)[/latex]
    4. [latex]\ce{2LiOH}(s)+\ce{CO2}(g)\longrightarrow\ce{Li2CO3}(s)+\ce{H2O}(g)[/latex]
  23. Using the enthalpy changes under standard state conditions determine the enthalpy change for these metal preparation reactions.
    1. [latex]\ce{2Al2O3}(s) \Longrightarrow 4\ce{Al}(s) + 3\ce{O2}(g)[/latex]
    2. [latex]\ce{Cr2O3}(s) + 3\ce{H2}(g) \Longrightarrow \ce{2Cr}(s) + \ce{3H2O}(l)[/latex]
    3. [latex]\ce{6Fe3C}(s) + 3\ce{H2}(g) \Longrightarrow 18\ce{Fe}(s) + \ce{C6H6}(g)[/latex]
  24. Which of the enthalpies of combustion in Table 10.3.1 are also standard enthalpies of formation?
  25. Find ΔH for the following reaction [latex]\ce{Co3O4}(s) \Longrightarrow \ce{3Co}(s) + \ce{2O2}(g)[/latex] using the information below
    [latex]\ce{Co}(s) + \frac{1}{2} \ce{O2}(g) \Longrightarrow \ce{CoO}(s) \Delta H = -237.9 \text{ kJ}[/latex]
    [latex]6ce{CoO}(s) + \ce{O2}(g) \Longrightarrow \ce{2Co3O4}(s) \Delta H = -3247.44[/latex] kJ
  26. Calculate the enthalpy of combustion of propane, [latex]\ce{C3H8}[/latex](g), for the formation of [latex]\ce{H2O}[/latex](g) and [latex]\ce{CO2}[/latex](g). The enthalpy of formation of propane is −104 kJ/mol.
  27. Looking at the combustion of hexane and octane, which substance would produce the most heat per gram burned?
  28. Water gas, a mixture of [latex]\ce{H2}[/latex] and CO, is an important industrial fuel produced by the reaction of steam with red hot coke, essentially pure carbon: [latex]\ce{C}\left(s\right)+{\ce{H}}_{2}\ce{O}\left(g\right)\longrightarrow\ce{CO}\left(g\right)+{\ce{H}}_{2}\left(g\right).[/latex]
    1. Assuming that coke has the same enthalpy of formation as graphite, calculate [latex]\Delta{H}_{298}^{\circ }[/latex] for this reaction.
    2. Methanol, a liquid fuel that could possibly replace gasoline, can be prepared from water gas and additional hydrogen at high temperature and pressure in the presence of a suitable catalyst: [latex]2{\ce{H}}_{2}\left(g\right)+\ce{CO}\left(g\right)\longrightarrow{\ce{CH}}_{3}\ce{OH}\left(g\right)[/latex]. Under the conditions of the reaction, methanol forms as a gas. Calculate [latex]\Delta{H}_{298}^{\circ }[/latex] for this reaction and for the condensation of gaseous methanol to liquid methanol.
    3. Calculate the heat of combustion of 1 mole of liquid methanol to [latex]\ce{H2O}[/latex](g) and [latex]\ce{CO2}[/latex](g).
  29. Using the standard enthalpies of formation, which of the following fuels produces the greatest amount of heat per gram when combusted under standard conditions, [latex]\ce{CO, CH4}, \text{ or } \ce{C2H2}[/latex]?
  30. Ethanol, [latex]\ce{C2H5OH}[/latex], is used as a fuel for motor vehicles, particularly in Brazil.
    1. Write the balanced equation for the combustion of ethanol to [latex]\ce{CO2}[/latex](g) and [latex]\ce{H2O}[/latex](g), and, using the data in Standard Thermodynamic Properties for Selected Substances, calculate the enthalpy of combustion of 1 mole of ethanol.
    2. The density of ethanol is 0.7893 g/mL. Calculate the enthalpy of combustion of exactly 1 L of ethanol.
    3. Assuming that an automobile’s mileage is directly proportional to the heat of combustion of the fuel, calculate how much farther an automobile could be expected to travel on 1 L of gasoline than on 1 L of ethanol. Assume that gasoline has the heat of combustion and the density of n–octane, [latex]\ce{C8H_{18}}[/latex] [latex]\left(\Delta{H}_{\text{f}}^{\circ }=-208.4\text{ kJ/mol}\right)[/latex] density = 0.7025 g/mL).
  31. [latex]\ce{2Fe}(\text{s}) + \frac{3}{2} \ce{O2}(\text{g}) \text{ } \square \text{ } \ce{Fe2O3}(\text{s}) \Delta \text{Hf˚} = -824[/latex] kJ/mol. How much heat, in joules, is produced when 1.47 g of iron metal reacts with excess oxygen gas under standard conditions?
  32. The oxidation of the sugar glucose, [latex]\ce{C6H_{12}O6}[/latex], is described by the following equation: [latex]{\ce{C}}_{6}{\ce{H}}_{12}{\ce{O}}_{6}\left(s\right)+6{\ce{O}}_{2}\left(g\right)\longrightarrow 6{\ce{CO}}_{2}\left(g\right)+6{\ce{H}}_{2}\ce{O}\left(l\right)\Delta\text{H}=-2816\text{ kJ}[/latex] The metabolism of glucose gives the same products, although the glucose reacts with oxygen in a series of steps in the body.
    1. How much heat in kilojoules can be produced by the metabolism of 1.0 g of glucose?
    2. How many Calories can be produced by the metabolism of 1.0 g of glucose?

[reveal-answer q=”11″]Show Selected Solutions[/reveal-answer]
[hidden-answer a=”11″]

  1. 14.8 kJ mol-1
  2. 5.5 kJ/ºC
  3. 10.34 g
  4. ΔH = 285.8 kJ
  5. 2839 kJ, enthalpy for [latex]\ce{O2}[/latex] would be zero because it is an element in its naturally occurring state.
  6. 418 kJ of heat is evolved.
  7. [latex]\ce{2C2H2}(\text{g}) + \ce{5O2}(\text{g}) \longrightarrow 4\ce{CO2}(\text{g}) + 2\ce{H2O}(\text{g})[/latex]
    [latex]\Delta H_{\text{rxn}} = -2599[/latex] kJ
  8. 1.0 g of octane produes 44.46 kJ of heat, 1.0 gram of hydrazine produces 16.68 kJ. On the assumption that the best fuel is the one that gives off the most heat, octane is the prime candidate. Other things must be considered, however. For example, the moles of gaseous product formed are related to the specific impulse of the fuel, toxicity of products, cost, and ability to contain original fuel (stability and corrosiveness).
  9. -45.38 kJ
  10. The answers are as follows:
    1. [latex]2\ce{C4H_{10}} + 13\ce{O2} \longrightarrow 8\ce{O2} + 10\ce{H2O}[/latex]
    2. Mass of [latex]\ce{O2}[/latex] = 71.7 g, Volume of [latex]\ce{O2}[/latex] = 261 L, and Volume of Air = 1242 L
    3. ΔT = 73.0 °C
  11. -394 kJ
  12. The answers are as follows:
    1. 3352 kJ
    2. 282.21 kJ
    3. -67.673 kJ
  13. -910.02 kJ
  14. hexane
  15. [latex]\ce{CH4}[/latex]
  16. 10,800 J


[/hidden-answer]

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