Energy And Chemical Change Exam Prep Chapter.6

Chemistry: Molecular Nature of Matter, 8e (Jespersen)

Chapter 6 Energy and Chemical Change

1) Which is a unit of energy, but is not the SI unit of energy?

A) joule

B) newton

C) pascal

D) watt

E) calorie

Diff: 1

Section: 6.1

2) Which is a unit of energy?

A) pascal

B) newton

C) joule

D) watt

E) ampere

Diff: 1

Section: 6.1

3) Chemical energy is

A) the kinetic energy resulting from violent decomposition of energetic chemicals.

B) the heat energy associated with combustion reactions.

C) the electrical energy produced by fuel cells.

D) the potential energy that resides in chemical bonds.

E) the energy living plants receive from solar radiation.

Diff: 1

Section: 6.1

4) Calculate the kinetic energy (KE) of an object that has a mass of 5.00 × 102 g and is traveling in a straight line with a speed of 50.0 m s−1. (1 J = 1 kg m2s−2.)

A) 0.625 kJ

B) 1.25 kJ

C) 2.5 kJ

D) 6.25 kJ

E) 25 kJ

Diff: 2

Section: 6.1

5) Calculate the kinetic energy (KE) of an object that has a mass of 9.00 × 102 g and is traveling in a straight line with a speed of 4.0 × 101 m s−1. (1 J = 1 kg m2s−2)

A) 0.72 kJ

B) 1.44 kJ

C) 2.88 kJ

D) 16.2 kJ

E) 18 kJ

Diff: 2

Section: 6.1

6) Calculate the kinetic energy (KE) of an object that has a mass of 1.200 × 103 g and is traveling in a straight line with a speed of 5.0 × 101 m s−1. (1 J = 1 kg m2s−2)

A) 1.5 kJ

B) 3.0 kJ

C) 6.0 kJ

D) 36 kJ

E) 300 kJ

Diff: 2

Section: 6.1

7) Calculate the kinetic energy (KE) of an object that has a mass of 2.45 kg and is traveling in a straight line with a speed of 12.0 m s−1. (1 J = 1 kg m2s−2)

A) 414 J

B) 353 J

C) 36.0 J

D) 176 J

E) 465 J

Diff: 2

Section: 6.1

8) How many kilojoules are equivalent to 8.18 kilocalories?

A) 1.96 kJ

B) 1,955 kJ

C) 8,180 kJ

D) 34,200 kJ

E) 34.2 kJ

Diff: 1

Section: 6.1

9) How many kilocalories are equivalent to 18.9 kilojoules?

A) 79.1 kcal

B) 4.52 kcal

C) 9.03 kcal

D) 7.91 kcal

E) 34.2 kcal

Diff: 1

Section: 6.1

10) Which statement is true?

A) Molecules in gases possess kinetic energy because they are in constant motion, while molecules in liquids and solids are not in constant motion and hence possess no kinetic energy.

B) Molecules in gases and liquids possess kinetic energy because they are in constant motion, while molecules in solids are not in constant motion and hence possess no kinetic energy.

C) Molecules in gases, liquids and solids possess kinetic energy because they are in constant motion.

D) Polyatomic molecules possess kinetic energy in the liquid and gaseous states because the atoms can move about in the molecule even if the molecule cannot move.

E) Because solids are rigid, their molecules do not possess kinetic energy unless the solid is melted.

Diff: 2

Section: 6.2

11) For a chemical reaction, where the internal energy is given the symbol E,

A) Efinal signifies the internal energy of the reactants.

B) Einitial signifies the internal energy of the products.

C) ΔE = Eproducts − Ereactants

D) ΔE is positive if energy is released to the surroundings.

E) ΔE is positive if energy is released by the chemical reaction.

Diff: 2

Section: 6.2

12) Which statement is incorrect?

A) Heat can be considered the energy transferred between objects with different temperatures.

B) Internal energy is the sum of the energies of all the individual particles in a particular sample of matter.

C) If a system absorbs energy, its internal energy increases.

D) Kinetic molecular theory is related to the total molecular kinetic energy.

E) If the Kelvin temperature is doubled, the average kinetic energy is also doubled.

Diff: 2

Section: 6.2

13) Which statement is true?

A) A state function is one whose value for a system depends on the method of preparation of the reactants and products.

B) A state function is one whose value for a system is determined by the difference in temperature of the system, and not on the pressure of the system.

C) A state function is one whose value for the system is determined by only the pressure of the system, and not on the temperature of the system.

D) A state function is one whose value for a system is determined by the temperature of the system, and not on the composition of the system.

E) A state function is one whose value for a system is determined by the composition of the system, the volume, the temperature, and the pressure.

Diff: 2

Section: 6.2

14) A freshly baked pie is placed near an open window to cool. Which of the following statements best describes this situation?

A) The pie is the system and loses heat to the surroundings.

B) The pie is the system and gains heat from the surroundings.

C) The pie is the surroundings and gains heat from the system.

D) The pie is the surroundings and loses heat to the system.

E) The pie is the surroundings and neither gains nor loses heat.

Diff: 1

Section: 6.3

15) A system that does not allow the transfer of mass but does allow the transfer of thermal energy would best be classified as

Hint: Consider this from the thermodynamic perspective.

A) an open system.

B) a closed system.

C) an isolated system.

D) an adiabatic system.

E) an isobaric system.

Diff: 3

Section: 6.3

16) A system that allows the transfer of mass and allows the transfer of thermal energy would best be classified as

A) an open system.

B) a closed system.

C) an isolated system.

D) an adiabatic system.

E) an isobaric system.

Diff: 2

Section: 6.3

17) A certain oil used in industrial transformers has a density of 1.068 g mL−1 and a specific heat of 1.628 J g−1 °C−1. Calculate the heat capacity of one gallon of this oil. (1 gallon = 3.785 liters)

A) 0.3747 kJ °C−1

B) 0.4027 kJ °C−1

C) 2.483 kJ °C−1

D) 5.770 kJ °C−1

E) 6.581 kJ °C−1

Diff: 2

Section: 6.3

18) A certain oil used in industrial transformers has a density of 1.086 g mL−1 and a specific heat of 1.826 J g−1 °C−1. Calculate the heat capacity of one gallon of this oil. (1 gallon = 3.785 liters)

A) 0.4442 kJ °C−1

B) 0.5239 kJ °C−1

C) 2.251 kJ °C−1

D) 6.364 kJ °C−1

E) 7.506 kJ °C−1

Diff: 2

Section: 6.3

19) A 500.0 gram sample of aluminum is initially at 25.0 °C. It absorbs 32.60 kJ of heat from its surroundings. What is its final temperature, in °C? (specific heat = 0.9930 J g−1 °C−1 for aluminum)

A) 40.4 °C

B) 64.7 °C

C) 65.7 °C

D) 89.7 °C

E) 90.7 °C

Diff: 2

Section: 6.3

20) A 113.25 gram sample of gold is initially at 100.0 °C. It gains 20.00 J of heat from its surroundings. What is its final temperature? (specific heat of gold = 0.129 J g−1 °C−1)

A) 98.6 °C

B) −98.6 °C

C) 101.4 °C

D) −101.4 °C

E) 96.6 °C

Diff: 2

Section: 6.3

21) A 225.0 gram sample of copper absorbs 735 J of heat from its surroundings. What is the temperature change for copper sample? (specific heat = 0.387 J g−1 °C−1 for copper)

A) 64.0 °C

B) 8.44 °C

C) 92.2 °C

D) 117.3 °C

E) 156.7 °C

Diff: 2

Section: 6.3

22) A 350.0 gram sample of copper initially at 25.0 °C absorbs 12.50 kJ of heat from its surroundings. What is its final temperature? (specific heat = 0.387 J g−1 °C−1 for copper)

A) 38.8 °C

B) 67.2 °C

C) 92.2 °C

D) 117.3 °C

E) 156.7 °C

Diff: 2

Section: 6.3

23) A bomb calorimeter consists of metal parts with a heat capacity of 850.0 J °C−1 and 1.100 × 103 grams of oil with a specific heat of 2.184 J g−-1 °C−1. What is the heat capacity, in joules per degree, of the entire calorimeter?

A) 1354 J °C−1

B) 1952 J °C−1

C) 2956 J °C−1

D) 3252 J °C−1

E) 4259 J °C−1

Diff: 2

Section: 6.3

24) A bomb calorimeter consists of metal parts with a heat capacity of 925.0 J °C−1 and 1.100 × 103 grams of oil with a specific heat of 2.814 J g−1 °C−1. What is the heat capacity, in joules per degree, of the entire calorimeter?

Hint: Determine the heat capacity of the oil and the metal separately first.

A) 1321 J °C−1

B) 2028 J °C−1

C) 3703 J °C−1

D) 4020 J °C−1

E) 5698 J °C−1

Diff: 3

Section: 6.3

25) A bomb calorimeter consists of metal parts with a heat capacity of 950.0 J °C−1 and 8.50 × 102 grams of oil with a specific heat of 2.418 J g−1 °C−1. Calculate the amount of heat energy required, in kJ, to raise the temperature of the calorimeter from 25.00 °C to 31.60 °C.

Hint: Determine the amount of energy required to raise the temperature of oil and metal separately first.

A) 4.91 kJ

B) 11.9 kJ

C) 19.8 kJ

D) 20.8 kJ

E) 28.7 kJ

Diff: 3

Section: 6.3

26) A bomb calorimeter consists of metal parts with a heat capacity of 925.0 J °C−1 and 1.100 × 103 grams of oil with a specific heat of 2.184 J g−1 °C−1. Calculate the heat required, in kJ, to raise the temperature of the calorimeter from 24.40 °C to 29.75 °C.

Hint: Determine the amount of energy required to raise the temperature of oil and metal separately first.

A) 0.827 kJ

B) 7.64 kJ

C) 17.8 kJ

D) 23.7 kJ

E) 99.0 kJ

Diff: 3

Section: 6.3

27) A 113.25 gram sample of gold is initially at 100.0 °C. It loses 20.00 J of heat to its surroundings. What is its final temperature? (specific heat of gold = 0.129 J g−1 °C−1)

A) 98.6 °C

B) −98.6 °C

C) 94.6 °C

D) −94.6 °C

E) 96.6 °C

Diff: 2

Section: 6.3

28) A 25.00 gram gold ingot and a 30.00 gram block of copper are placed in 100.00 grams of water. If the initial temperatures of the gold, copper, and water were 95.0 °C, 85.0 °C, and

25.0 °C, respectively, what would the final temperature of the entire system be? The specific heats of gold, copper, and liquid water are 0.129, 0.387, and 4.18 J g−1 °C−1, respectively.

Hint: Set up your equation for each metal and water individually before combining to determine the final temperature of the entire system.

A) 26.0 °C

B) 28.2 °C

C) 23.1 °C

D) −27.1 °C

E) 27.1 °C

Diff: 3

Section: 6.3

29) A 25.00 gram pellet of lead (specific heat = 0.128 J g−1 °C−1) at 25 °C is added to 95.3 g of boiling water (specific heat of 4.18 J g−1 °C−1) at 100 °C in an insulated cup. What is the expected final temperature of the water?

A) 26.6 °C

B) 62.5 °C

C) 84.4 °C

D) 99.4 °C

E) 100.6 °C

Diff: 2

Section: 6.3

30) A 55.00 gram pellet of lead at 25 °C is added to 58.5 g of boiling water (specific heat of 4.18 J g−1 °C−1) at 100 °C in an insulated cup. If the final temperature of the water in the cup is

97.9 °C, what is the specific heat of lead?

A) 17.8 J g−1 °C−1

B) 0.128 J g−1 °C−1

C) 4.17 J g−1 °C−1

D) 22.2 J g−1 °C−1

E) 0.372 J g−1 °C−1

Diff: 2

Section: 6.3

31) A sample of chromium weighing 254 g was initially at a temperature of 25.88 °C. It required 843 joules of heat energy to increase the temperature to 32.75 °C. What is the molar heat capacity of the chromium?

Hint: It might help to calculate the specific heat capacity first.

A) 21.6 J mol−1 °C−1

B) 25.1 J mol−1 °C−1

C) 33.2 J mol−1 °C−1

D) 37.3 J mol−1 °C−1

E) 17.4 J mol−1 °C−1

Diff: 3

Section: 6.3

32) A coffee cup calorimeter contains 480.0 grams of water at 25.00 °C. To it are added:

380.0 grams of water at 53.5 °C

525.0 grams of water at 65.5 °C

Assuming the heat absorbed by the coffee cup is negligible, calculate the expected final temperature of the water. The specific heat of water is 4.184 J g−1 °C−1.

Hint: Take this problem one step (one water addition) at a time.

A) 38.2 °C

B) 48.2 °C

C) 67.6 °C

D) 88.7 °C

E) 94.4 °C

Diff: 3

Section: 6.3

33) A coffee cup calorimeter contains 525.0 grams of water at 25.0 °C. To it are added:

350.0 grams of water at 48.3 °C

480.0 grams of water at 63.8 °C

Neglecting the heat absorbed by the coffee cup, calculate the final temperature of the water. The specific heat of water is 4.184 J g−1 °C−1.

Hint: Take this problem one step (one water addition) at a time.

A) 39.6 °C

B) 45.7 °C

C) 44.8 °C

D) 66.7 °C

E) 92.4 °C

Diff: 3

Section: 6.3

34) A constant pressure calorimeter consists of metal parts with a heat capacity of 850.0 J °C−1 and 1.050 × 103 grams of oil with a specific heat of 2.148 J g−1 °C−1. Both are at 24.50 °C. A 5.00 × 102 g copper slug, at 220.0 °C is added. What is the final temperature? Specific heat of Cu = 0.3874 J g−1 °C−1.

Hint: Take this problem one step at a time. Look at the energy required to raise the temperature of oil and metal separately first before considering the entire system.

A) 33.4 °C

B) 36.0 °C

C) 36.8 °C

D) 89.7 °C

E) 120.5 °C

Diff: 3

Section: 6.3

35) A constant pressure calorimeter has metal parts (heat capacity of 850.0 J °C−1) and 1.100 × 103 grams of oil (specific heat = 2.184 J g−1 °C−1), both at 24.50°C. Adding a 4.60 × 102 g slug, at 240.0°C, caused the temperature to rise to 32.5 °C. Find the specific heat of the metal.

Hint: Consider the amount of energy that was required to raise the temperature of the entire system.

A) 0.236 J g−1 °C−1

B) 0.273 J g−1 °C−1

C) 0.309 J g−1 °C−1

D) 0.357 J g−1 °C−1

E) 2.28 J g−1 °C−1

Diff: 3

Section: 6.3

36) A constant pressure calorimeter has metal parts (heat capacity of 925.0 J °C−1) and 1.100 × 103 grams of oil (specific heat = 2.824 J g−1 °C−1), both at 25.40°C. Adding a 5.50 × 102 g slug at 220.0°C, caused the temperature to rise to 35.2 °C. Find the specific heat of the metal.

Hint: Consider the amount of energy that was required to raise the temperature of the entire system.

A) 0.365 J g−1 °C−1

B) 0.389 J g−1 °C−1

C) 0.395 J g−1 °C−1

D) 0.551 J g−1 °C−1

E) 1.20 J g−1 °C−1

Diff: 3

Section: 6.3

37) During an exothermic chemical reaction,

A) a system becomes warmer and the chemical substances undergo an increase in potential energy.

B) a system becomes warmer and the chemical substances undergo a decrease in potential energy.

C) a system becomes cooler and the chemical substances undergo an increase in potential energy.

D) a system becomes cooler and the chemical substances undergo a decrease in potential energy.

E) a system becomes warmer and additional heat is gained from the surroundings.

Diff: 2

Section: 6.4

38) During an endothermic chemical reaction,

A) a system becomes warmer and the chemical substances undergo an increase in potential energy.

B) a system becomes warmer and the chemical substances undergo a decrease in potential energy.

C) a system becomes cooler and the chemical substances undergo an increase in potential energy.

D) a system becomes cooler and the chemical substances undergo a decrease in potential energy.

E) a system becomes warmer and additional heat is gained from the surroundings.

Diff: 2

Section: 6.4

39) Which statement is generally true?

A) A chemical reaction involves only the making of chemical bonds.

B) A chemical reaction involves only the breaking of chemical bonds.

C) Breaking weak chemical bonds require a relatively large amount of energy.

D) When bonds break in chemical reactions, the potential energy of the system tends to increase.

E) When bonds break in chemical reactions, the potential energy of the system tends to decrease.

Diff: 2

Section: 6.4

40) For a change in a system that takes place at constant pressure, which statement below is true?

A) ΔH = ΔE

B) ΔH = qp − P ΔV

C) ΔH = ΔE − qp

D) ΔH = qp

E) ΔE = qp

Diff: 1

Section: 6.5

41) For a chemical reaction taking place at constant pressure, which one of the following is true?

A) ΔHsystem = (Kinetic Energy)system + (Potential Energy)system

B) ΔHsystem = (Kinetic Energy)system − (Potential Energy)system

C) ΔHsystem = ΔEsystem − qp

D) ΔHsystem = ΔEsystem + PΔVsystem

E) ΔHsystem = ΔEsystem + qp

Diff: 2

Section: 6.5

42) An endothermic reaction is one in that there is

A) a positive value for the work done by the system (w > 0 joules).

B) a negative value for the work done by the system (w < 0 joules).

C) a negative value for ΔH (ΔH < 0 joules).

D) a positive value for ΔH (ΔH > 0 joules).

E) a negative value for ΔE (ΔE > 0 joules).

Diff: 2

Section: 6.5

43) In the course of measuring fuel content values, a reaction for the conversion of crude oil fuel into water and carbon dioxide is carried out in two steps

Crude fuel oil + oxygen → CO(g) + H2O

CO(g) + oxygen → CO2(g)

The net reaction taking place is: crude fuel oil + oxygen → CO2(g) + H2O. A large fraction of the raw material is converted in one step, while the second step is to collect the fraction that was just partially burned the first time. For the overall or net process, which statement below is always true?

A) ΔH is independent of the time interval between the two steps, but dependent on the fraction that had to be converted in two steps.

B) ΔH is dependent on the time interval between the two steps, but dependent on the fraction that had to be converted in two steps.

C) ΔH is independent of the time interval between the two steps and also independent of the fraction that had to be converted in two steps.

D) ΔH is dependent on the time interval between the two steps, but independent of the fraction that had to be converted in two steps.

E) ΔH is independent of the time interval between the two steps, but dependent on the time required for completion of the entire process.

Diff: 2

Section: 6.5

44) A chemical reaction took place in a 6 liter cylindrical enclosure fitted with a piston (like the cylinder in an internal combustion engine). Over the time required for the reaction to be completed, the volume of the system changed from 0.400 liters to 3.20 liters. Which of the following statements below is true?

A) Work was performed on the system.

B) Work was performed by the system.

C) The internal energy of the system increased.

D) The internal energy of the system decreased.

E) The internal energy of the system remained unchanged.

Diff: 2

Section: 6.5

45) A chemical reaction took place in a 5 liter cylindrical enclosure fitted with a piston (like the cylinder in an internal combustion engine). Over the time required for the reaction to be completed, the volume of the system changed from 1.40 liters to 3.70 liters. Which of the following statements below is true?

A) The enthalpy of the system remained unchanged.

B) The enthalpy of the system decreased.

C) The enthalpy of the system increased.

D) Work was performed by the system.

E) Work was performed on the system.

Diff: 2

Section: 6.5

46) A closed, uninsulated system was fitted with a movable piston. The introduction of 430 J of heat caused the system to expand, doing 238 J of work in the process against a constant pressure of 101 kPa (kilopascals). What is the value of ΔE for this process?

A) (430 + 238) joules

B) (430 − 238) joules

C) (238 − 430) joules

D) 430 joules

E) (−238 − 430) joules

Diff: 2

Section: 6.5

47) A closed, uninsulated system was fitted with a movable piston. Introduction of 430 J of heat caused the system to expand, doing 238 J of work in the process against a constant pressure of 101 kPa (kilopascals). What is the value of ΔH for this process?

A) (430 + 238) joules

B) (430 − 238) joules

C) (238 − 430) joules

D) 430 joules

E) (−238 − 430) joules

Diff: 2

Section: 6.5

48) A closed, uninsulated system was fitted with a movable piston. Introduction of 483 J of heat caused the system to expand, doing 320 J of work in the process against a constant pressure of 101 kPa (kilopascals). What is the value of ΔE for this process?

A) (483 + 320) joules

B) (483 − 320) joules

C) (320 − 483) joules

D) 483 joules

E) (−320 − 483) joules

Diff: 2

Section: 6.5

49) A closed, uninsulated system was fitted with a movable piston. Introduction of 483 J of heat caused the system to expand, doing 320 J of work in the process against a constant pressure of 101 kPa (kilopascals). What is the value of ΔH for this process?

A) (483 + 320) joules

B) (483 − 320) joules

C) (320 − 483) joules

D) 483 joules

E) (−320 − 483) joules

Diff: 2

Section: 6.5

50) For the reaction, D2(s) + 2 AX(g) → A2(g) + 2 DX(g) taking place in an insulated system, the enthalpy of the reactants is lower than that of the products. Which one of the following is true for the system?

A) The energy of the system decreases as the reactants are converted to products.

B) The energy of the system increases as the reactants are converted to products.

C) The total energy of the system decreases as the reactants are converted to products.

D) The total mass of the system decreases as the reactants are converted to products.

E) The total mass of the system increases as the reactants are converted to products.

Diff: 2

Section: 6.6

51) When pure sodium hydroxide is dissolved in water, heat is evolved. In a laboratory experiment to measure the molar heat of solution of sodium hydroxide, the following procedure was followed. To a calorimeter containing 3.00 × 102 g of water at 20.00 °C, 10.65 g of NaOH, also at 20.00 °C was added. The temperature of the solution, which was monitored by a digital thermometer with negligible heat capacity, increased to 28.50 °C. If the specific heat of the mixture is 4.184 J g−1 °C−1, and the small heat capacity of the calorimeter is ignored, what is the heat evolved, per mole of sodium hydroxide?

Hint: Do not ignore the specific heat of water given in this problem and be sure to convert NaOH to moles before proceeding.

A) −37.4 kJ

B) −41.5 kJ

C) −45.5 kJ

D) −90.5 kJ

E) −153 kJ

Diff: 3

Section: 6.6

52) When pure sulfuric acid is dissolved in water, heat is evolved. In a laboratory experiment to measure the molar heat of solution of sulfuric acid, the following procedure was followed. To a calorimeter containing 3.00 × 102 g of water at 20.00 °C, 10.65 g of H2SO4, also at 20.00 °C was added. The temperature of the solution, which was monitored by a digital thermometer with negligible heat capacity, increased to 26.55 °C. If the specific heat of the mixture is

4.184 J g−1 °C−1, and the small heat capacity of the calorimeter is ignored, what is the heat evolved, per mole of sulfuric acid?

Hint: Do not ignore the specific heat of water given in this problem and be sure to convert H2SO4 to moles before proceeding.

A) −27.4 kJ

B) −72.8 kJ

C) −78.4 kJ

D) −84.6 kJ

E) −292 kJ

Diff: 3

Section: 6.6

53) When 0.250 moles of LiCl are added to 200.0 g of water in a constant pressure calorimeter a temperature change of +11.08 °C is observed. Given that the specific heat of the resulting solution is 4.184 J g−1 °C and we can ignore the small amount of energy absorbed by the calorimeter, what is the molar enthalpy of solution (ΔHsol) for LiCl?

Hint: Do not ignore the specific heat of water in this problem.

A) 37.1 kJ/mol

B) −185.4 kJ/mol

C) −37.1 kJ/mol

D) 18.5 kJ/mol

E) −18.5 kJ/mol

Diff: 3

Section: 6.6

54) What would be the "standard state" for acetic acid in solution?

A) A solution with a concentration of 1.000 M.

B) A solution at 1.000 bar of pressure.

C) A solution at 1.000 Pascal of pressure.

D) A solution at 298 K.

E) A solution that is in the solid state.

Diff: 1

Section: 6.7

55) What would be the "standard state" for hydrogen gas at room temperature?

A) A gas sample with a concentration of 1.000 M.

B) A gas sample at 1.000 atm of pressure.

C) A gas sample at 1.000 Pascal of pressure.

D) A liquid solution at 298 K.

E) A liquid solution at 1.000 atm.

Diff: 1

Section: 6.7

56) When nitrogen gas reacts with hydrogen gas to form ammonia, 92.38 kJ of heat are given off for each mole of nitrogen gas consumed, under constant pressure and standard conditions. What is the correct value for the standard enthalpy of reaction in the thermochemical equation below when 0.750 mol of hydrogen reacts?

N2(g) + 3H2(s) → 2 NH3(g)

A) +34.5 kJ

B) −98.3 kJ

C) +59.2 kJ

D) −59.2 kJ

E) −23.1 kJ

Diff: 2

Section: 6.7

57) When aluminum metal reacts with iron(III) oxide to form aluminum oxide and iron metal, 429.6 kJ of heat are given off for each mole of aluminum metal consumed, under constant pressure and standard conditions. What is the correct value for the standard enthalpy of reaction in the thermochemical equation below?

2 Al(s) + Fe2O3(s) → 2 Fe(s) + Al2O3(s)

A) +429.6 kJ

B) −429.6 kJ

C) +859.2 kJ

D) −859.2 kJ

E) −1289 kJ

Diff: 2

Section: 6.7

58) The thermochemical equation for the reaction between dinitrogen monoxide and oxygen to produce nitrogen dioxide is shown below. Write the thermochemical equation for the reaction when 1.00 mole of nitrogen dioxide is formed.

2N2O(g) + 3O2(g) → 4NO2(g) ΔH° = −28.0 kJ

A) N2O(g) + 3O2(g) → NO2(g) ΔH° = −28.0 kJ

B) N2O(g) + O2(g) → NO2(g) ΔH° = −28.0 kJ

C) 2N2O(g) + 3O2(g) → 4NO2(g) ΔH° = −56.0 kJ

D) ½ N2O(g) + ¾ O2(g) → NO2(g) ΔH° = −7.00 kJ

E) ½ N2O(g) + O2(g) → NO2(g) ΔH° = −14.0 kJ

Diff: 2

Section: 6.7

59) The combustion of butane, C4H10, is given as:

2 C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(l), ΔH° = −5,314 kJ.

How many grams of butane must be reacted by this reaction to release 15,285 kJ of heat?

Hint: Work backwards with moles potentially if you are unsure where to start.

A) 167.2 g

B) 83.62 g

C) 668.8 g

D) 333.7g

E) 33.09 g

Diff: 3

Section: 6.7

60) Propane is often used to heat homes. The combustion of propane follows the following reaction:

C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g), ΔH° = −2044 kJ.

How many grams of propane must be reacted by this reaction to release 7563 kJ of heat?

Hint: Work backwards with moles potentially if you are unsure where to start.

A) 3.70 g

B) 44.1 g

C) 81.6 g

D) 243.4 g

E) 162.8 g

Diff: 3

Section: 6.7

61) For the reaction below:

CaO(s) + H2O(l) → Ca(OH)2(s) ΔH° = −64.8 kJ.

How many grams of CaO must be reacted by this reaction to release 1050 kJ of heat?

Hint: Work backwards with moles potentially if you are unsure where to start

A) 16.2 g

B) 907 g

C) 1817 g

D) 454 g

E) 56.1 g

Diff: 3

Section: 6.7

62) When nitrogen gas reacts with hydrogen gas to form ammonia, 92.38 kJ of heat is given off for each mole of nitrogen gas consumed, under constant pressure and standard conditions. What is the value of ΔH° for the reverse of the reaction shown?

N2(g) + 3H2(s) → 2 NH3(g) ΔH° = −92.38 kJ

A) +34.5 kJ

B) −46.19 kJ

C) +59.2 kJ

D) −59.2 kJ

E) +92.38 kJ

Diff: 1

Section: 6.8

63) Consider the following thermochemical equation:

2NO(g) + O2(g) → 2 NO2(g) ΔH° = −113.2 kJ

Calculate H° for the reaction below:

4 NO2(g) → 4NO(g) + 2O2(g) ΔH° = ??

A) +334.5 kJ

B) −146.19 kJ

C) +226.4 kJ

D) −509.2 kJ

E) +192.38 kJ

Diff: 1

Section: 6.8

64) Determine the enthalpy change, ΔH, for the reaction, N2(g) + 2H2(g) → N2H4(l), given the following thermochemical equations:

N2H4(l) + O2(g) → 2H2O(l) + N2(g) ΔH = −622.0 kJ

H2(g) + ½ O2(g) → H2O(l) ΔH = −285.9 kJ

A) −151.7 kJ

B) −236.2 kJ

C) +106.1 kJ

D) +50.2 kJ

E) +567.4 kJ

Diff: 2

Section: 6.8

65) Determine the enthalpy change, ΔH, for the reaction, W(s) + C(s) → WC(s), given the following thermochemical equations:

2W(s) + 3O2(g) → 2WO3(s) ΔH = −1680.8 kJ

C(s) + O2(g) → CO2(g) ΔH = −393.5 kJ

2WC(s) + 5O2(g) → 2WO3(s) + 2CO2(g) ΔH = −2391.4 kJ

Hint: Pay careful attention to your signs. If you reverse an equation remember to change the sign appropriately.

A) +33.3 kJ

B) −38.2 kJ

C) +106.1 kJ

D) −52.9 kJ

E) +177.4 kJ

Diff: 3

Section: 6.8

66) Determine the enthalpy change, ΔH, for the reaction, CS2(l) + 3O2(g) → CO2(g) + 2SO2(g), given the following thermochemical equations:

C(s) + O2(g) → CO2(g) ΔH = −393.5 kJ

S(s) + O2(g) → SO2(g) ΔH = −296.8 kJ

C(s) + 2S(s) → CS2(l) ΔH = 87.9 kJ

Hint: Pay careful attention to your signs. If you reverse an equation remember to change the sign appropriately.

A) +778.2 kJ

B) −602.4 kJ

C) −1075 kJ

D) −778.2 kJ

E) +602.4 kJ

Diff: 3

Section: 6.8

67) Determine the standard enthalpy change, ΔH°, for the reaction,

CCl4(l) + 4HCl(g) → CH4(g) + 4Cl2(g), given the following thermochemical equations:

2HCl(g) → H2(g) + Cl2(g) ΔH° = 184.6 kJ

C(s) + 2Cl2(g) → CCl4(l) ΔH° = −139 kJ

CH4(g) → C(s) + 2H2(g) ΔH° = 74.8 kJ

Hint: Pay careful attention to your signs. If you reverse an equation remember to change the sign appropriately.

A) +55.3 kJ

B) −187 kJ

C) +101 kJ

D) −179 kJ

E) +433 kJ

Diff: 3

Section: 6.8

68) The thermochemical equation that is associated with ΔH°f, the standard enthalpy of formation for HCl(g), is

A) H(g) + Cl(g) → HCl(g).

B) H2(g) + Cl2(g) → 2 HCl(g).

C) ½ H2(g) + ½ Cl2(g) → HCl(g).

D) H2(g) + Cl2(l) → 2 HCl(g).

E) ½ H2(g) + ½ Cl2(l) → HCl(g).

Diff: 1

Section: 6.9

69) The thermochemical equation that is associated with ΔH°f, the standard enthalpy of formation for H2O(g), is

A) 2 H(g) + O(g) → H2O(g).

B) H2O2(l) + ½O(g) → H2O(g).

C) 2 H2(g) + O2(g) → 2 H2O(g).

D) H2(g) + ½ O2(g) → H2O(g).

E) 2 H(g) + ½ O2(g) → H2O(g).

Diff: 1

Section: 6.9

70) The thermochemical equation that is associated with ΔH°f, the standard enthalpy of formation for glucose, C6H12O6(s), is

A) 6 C(s) + 6 H2O(l) → C6H12O6(s).

B) 6 C(s) + 12 H(g) + 6 O(g) → C6H12O6(s).

C) 6 C(s) + 6 H2(g) + 3 O2(g) → C6H12O6(s).

D) 2 C2H5OH(l) + 2 CO2(g) → C6H12O6(s).

E) 6 C(g) + 6 H2(g) + 3 O2(g) → C6H12O6(s).

Diff: 2

Section: 6.9

71) The thermochemical equation that is associated with ΔH°f, the standard enthalpy of formation for acetic acid, C2H4O2(l), is

A) C2(s) + 4 H(g) + O2(g) → C2H4O2(l).

B) 2 C(g) + 4 H(g) + 2 O(g) → C2H4O2(l).

C) C2(s) + 2 H2(g) + O2(g) → C2H4O2(l).

D) 2 C(s) + 2 H2(g) + O2(g) → C2H4O2(l).

E) C(s) + H2(g) + O2(g) → C2H4O2(l).

Diff: 2

Section: 6.9

72) The thermochemical equation that is associated with ΔH°f, the standard enthalpy of formation for urea, CO(NH2)2(s), is

A) CO(g) + 2 NH3(g) → CO(NH2)2(s) + H2(g).

B) CO(g) + 2 H2(g) + N2(g) → CO(NH2)2(s).

C) C(s) + O(g) + N2(g) + 2 H2(g) → CO(NH2)2(s).

D) C(s) + ½ O2(g) + N2(g) + 2 H2(g) → CO(NH2) 2(s).

E) C(s) + ½ O2(g) + 2 NH2(g) → CO(NH2) 2(s).

Diff: 2

Section: 6.9

73) Given the equation for a hypothetical reaction, 3A + 4B → 4C + 7D, and the following standard enthalpies of formation,

ΔH°f: A: +15.7 kJ mol−1 B: 86.4 kJ mol−1 C: −52.7 kJ mol−1 D: −71.6 kJ mol−1

calculate the standard enthalpy of reaction, in kJ, for the reaction shown.

A) −53.6 kJ

B) −413.5 kJ

C) −515.6 kJ

D) −853.6 kJ

E) −908.4 kJ

Diff: 2

Section: 6.9

74) Given the equation for a hypothetical reaction, 5A + 3B → 7C + 3D, and the following standard enthalpies of formation,

ΔH°f : A: −15.7 kJ mol−1 B: −86.4 kJ mol−-1 C: −52.7 kJ mol−1 D: −71.6 kJ mol−-1

what is the standard enthalpy of reaction, in kJ for the reaction shown?

A) +26.6 kJ

B) −53.6 kJ

C) −198.8 kJ

D) −246.0 kJ

E) −413.5 kJ

Diff: 2

Section: 6.9

75) Given the equation for the reaction, CO2(g) + 4H2(g) → CH4(g) + 2H2O(g), and the following standard enthalpies of formation,

ΔH°f : CO2(g): −393.5 kJ mol−1

CH4(g): −74.8 kJ mol−1

H2O(g): −241.8 kJ mol−1

H2O(l): −285.8 kJ mol−1

what is the standard enthalpy of reaction, in kJ for the reaction shown?

A) −164.9 kJ

B) +76.9 kJ

C) −164.5 kJ

D) +978.3 kJ

E) +995.9 kJ

Diff: 2

Section: 6.9

76) Ammonia gas reacts with molecular oxygen to produce nitrogen dioxide gas and water vapor. Given the following standard enthalpies of formation, ΔH°f:

NH3 (g): −80.3 kJ mol−1

NO2 (g): +33.2 kJ mol−1

H2O(g): −241.8 kJ mol−1

H2O(l): −285.8 kJ mol−1

What is the standard enthalpy of reaction, in kJ for the reaction shown?

Hint: Write the balanced reaction and be sure to pay attention to your coefficients when calculating the standard enthalpy of the overall reaction.

A) −172.3 kJ

B) −128.3 kJ

C) +157.5 kJ

D) −996.8 kJ

E) +1003.8 kJ

Diff: 3

Section: 6.9

77) The standard enthalpy of combustion for xylene, C8H10(l), is −3908 kJ mol−1. Using this information and the standard enthalpies of formation of the following, ΔH°f:

H2O(l) = −285.9 kJ mol−1; CO2(g) = −393.5 kJ mol−1, calculate the standard enthalpy of formation of C8H10(l), in kJ mol−1.

Hint: To begin, determine the thermochemical equation associated with the standard enthalpy for the formation of xylene.

A) -669.5 kJ

B) +3228.6 kJ

C) -3228.6 kJ

D) +4587.4 kJ

E) +8485.5 kJ

Diff: 3

Section: 6.9

78) A chemical compound has a molecular weight of 89.05 g/mole. 1.400 grams of this compound underwent complete combustion under constant pressure conditions in a calorimeter with a heat capacity of 2.980 × 103 J °C−1. The temperature went up by 11.95 degrees. Calculate the standard heat of combustion of the compound.

Hint: Remember to use moles as an intermediary when determining the standard heat of combustion.

A) 35.6 kJ mol−1

B) 686.2 kJ mol−1

C) 1681 kJ mol−1

D) 1886 kJ mol−1

E) 2265 kJ mol−1

Diff: 3

Section: 6.9

79) The standard heat of combustion for naphthalene, C10H8(s), is −5156.8 kJ mol−1. Using this information and the standard enthalpies of formation, ΔH°f: H2O(l) = −285.9 kJ mol−1; CO2(g) = −393.5 kJ mol−1, calculate the standard enthalpy of formation of C10H8(s), in kJ mol−1.

Hint: Write out an equation and remember products minus reactants for your calculations.

A) +78.2 kJ

B) +935.9 kJ

C) −1065.4 kJ

D) +3619.7 kJ

E) −10235.4 kJ

Diff: 3

Section: 6.9

80) Complete combustion of hydrocarbons or compounds with only C, H, and O gives CO2 and H2O as the only products. If carried out under standard conditions, the CO2 is a gas and the H2O is a liquid.

Given these standard enthalpies of combustion:

C2H4(g) = −1411.08 kJ mol−1 C2H2(g) = −1299.65 kJ mol−1

H2(g) = −285.90 kJ mol−1 C(s) = −393.50 kJ mol−1

calculate ΔH°reaction for the reaction, C2H2(g) + H2(g) → C2H4(g)

Hint: Remember it is always products minus reactants when performing enthalpy calculations.

A) −174.47 kJ

B) +397.33 kJ

C) −961.47 kJ

D) −2424.83 kJ

E) −2996.63 kJ

Diff: 3

Section: 6.9

81) Complete combustion of hydrocarbons or compounds with only C, H, and O gives CO2 and H2O as the only products. If carried out under standard conditions, the CO2 is a gas and the H2O is a liquid. Given these standard enthalpies of combustion: C6H12(l) = −3919.86 kJ mol−1, C6H6(l) = −3267.80 kJ mol−1, H2(g) = −285.90 kJ mol−1, C(s) = −393.50 kJ mol−1, calculate ΔH°reaction for the reaction,

C6H6(l) + 3 H2(g) → C6H12(l)

Hint: Remember it is always products minus reactants when performing enthalpy calculations.

A) −205.64 kJ

B) +366.16 kJ

C) +759.66 kJ

D) +2155.36 kJ

E) +5684.36 kJ

Diff: 3

Section: 6.9

82) The energy transferred between objects caused by differences in their temperatures is called ________.

Diff: 1

Section: 6.2

83) In the equation for the determination of the kinetic energy of an object, KE = mv2, the 'm' in the equation represents its ________.

Diff: 1

Section: 6.1

84) In order for a smaller motorcycle to have the same kinetic energy as a large SUV, the SUV would have to be traveling ________ than the motorcycle.

Diff: 2

Section: 6.1

85) What is the mass of an object that has a kinetic energy of 1 J travelling at a speed of 1 m s−1? ________.

Diff: 2

Section: 6.1

86) The sum of the energies of all the individual particles in a sample of a substance is its ________.

Diff: 1

Section: 6.2

87) When the Kelvin temperature of a sample of molecules increases, its average kinetic energy ________.

Diff: 1

Section: 6.2

88) A 500.0 gram sample of water has an initial temperature of 25.0 °C. It absorbs 50.0 kJ of heat from its surroundings. What is its final temperature, in °C?

Specific heat of water = 4.184 J g−1 °C−−1.

Diff: 2

Section: 6.3

89) A sample of nickel weighing 425 grams was initially at a temperature of 26.20 °C. It required 975 joules of heat energy to increase the temperature to 31.55 °C. What is the specific heat of the nickel?

Diff: 2

Section: 6.3

90) A sample of zinc weighing 425 grams was initially at a temperature of 25.40 °C. It required 1360 joules of heat energy to increase the temperature to 33.70 °C. What is the molar heat capacity of the zinc?

Hint: Determine how many moles of zinc you have when doing this problem.

Diff: 3

Section: 6.3

91) Reactions that consume energy are said to be ________.

Diff: 1

Section: 6.4

92) In a chemical equation, heat can be written as a product. A chemical reaction in which heat is a product is described as ________.

Diff: 1

Section: 6.4

93) In an exothermic reaction, energy is ________ the system.

Diff: 1

Section: 6.4

94) In an endothermic reaction, energy is ________ the system.

Diff: 1

Section: 6.4

95) The heat released or absorbed in a chemical reaction is called the ________.

Diff: 1

Section: 6.5

96) In a gas-phase chemical reaction performed at constant volume, the heat absorbed by the insulated calorimeter was calculated to be 29.3 kJ. What is qv for the reaction?

Diff: 2

Section: 6.5

97) In a gas-phase chemical reaction performed at constant atmospheric pressure, the heat absorbed by the insulated calorimeter was calculated to be 19.8 kJ. What is qp for the reaction?

Diff: 2

Section: 6.5

98) When a chemical reaction occurs in a bomb calorimeter, the change in volume for the system, V, is equal to ________.

Diff: 2

Section: 6.6

99) For a chemical reaction in which the change is endothermic, the sign of the enthalpy change, H, is ________.

Diff: 2

Section: 6.6

100) At constant pressure, the difference between the enthalpy change, H, and the internal energy change, ΔE, is equal to ________.

Diff: 2

Section: 6.6

101) When 0.250 moles of KCl are added to 200.0 g of water in a constant pressure calorimeter a temperature change is observed. Given that the specific heat of the resulting solution is 4.184 J g−1 °C, the molar heat of solution of KCl is +17.24 kJ/mol, and that we can ignore the small amount of energy absorbed by the calorimeter, the observed temperature change should be ________.

Hint: Make sure you convert to moles within the problem.

Diff: 3

Section: 6.6

102) The thermochemical equation for the reaction between methane and oxygen to produce carbon dioxide and water is:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(l), ΔH° = −890.5 kJ.

What is the corresponding thermochemical equation for this reaction when 1 mol of oxygen reacts?

Diff: 2

Section: 6.7

103) The thermochemical equation for the reaction between sulfur and oxygen to produce sulfur dioxide is:

S(s) + O2(g) → SO2(g), ΔH° = −298 kJ.

What is the corresponding thermochemical equation for this reaction when 2 mol of sulfur react?

Diff: 2

Section: 6.7

104) The thermochemical equation for the reaction between hydrazine, N2H4, and dinitrogen tetroxide is given as:

2N2H4(l) + N2O4(l) → 3N2(g) + 4H2O(g), ΔH° = −1049 kJ.

What is the corresponding thermochemical equation for this reaction when 6 mol of nitrogen are formed?

Diff: 2

Section: 6.7

105) The thermochemical equation for the reaction of sulfur dioxide, SO2, with oxygen is given as:

2SO2(g) + O2(g) → 2SO3(g), ΔH° = −198 kJ.

What is the change in enthalpy when 3 moles of SO2 react by this reaction?

Diff: 2

Section: 6.7

106) The thermochemical equation for the reaction of sulfur dioxide, SO2, with oxygen is given as:

2SO2(g) + O2(g) → 2SO3(g), ΔH° = -198 kJ.

What is the change in enthalpy when 5 moles of SO2 react with 2 moles of oxygen by this reaction?

Hint: Use the coefficients in the equation and the mole ratio shown during your calculation.

Diff: 3

Section: 6.7

107) For the reaction of graphite with oxygen the reaction is given as:

C(graphite) + O2(g) → 2CO2(g), ΔH° = −393 kJ.

How many grams of graphite must be reacted by this reaction to release 225 kJ of heat?

Hint: Use your coefficients from the equation and convert to moles as an intermediary when performing your calculations.

Diff: 3

Section: 6.7

108) The combustion of butane, C4H10, is given as:

2 C4H10(g) + 13O2(g) → 8CO2(g) + 8H2O(l), ΔH° = −5,314 kJ.

How many grams of butane must be reacted by this reaction to release 10,525 kJ of heat?

Hint: Use your coefficients from the equation and convert to moles as an intermediary when performing your calculations.

Diff: 3

Section: 6.7

109) The thermochemical equation for the reaction of sulfur dioxide, SO2, with oxygen is given as:

2SO2(g) + O2(g) → 2SO3(g), ΔH° = −198 kJ.

How much energy is given off when 300 g of SO2 is burned?

Hint: Use your coefficients from the equation and convert to moles as an intermediary when performing your calculations.

Diff: 3

Section: 6.7

110) For the reaction, N2(g) + 3 H2(g) → 2 NH3(g), ΔH° = −92.38 kJ per mole of nitrogen gas.

What is the value of ΔH° for the reaction, NH3(g) → 1/2 N2(g) + 3/2 H2(g)?

Diff: 2

Section: 6.8

111) For the reaction, 3 N2(g) + H2(g) → 2 HN3(g), ΔH° = +264 per mole of hydrogen gas.

What is the value of ΔH° for the reaction, HN3(g) → 1/2 H2(g) + 3/2 N2(g)?

Diff: 2

Section: 6.8

112) Given the thermochemical equation, 2 M2O5(s) → 4 MO2(s) + O2(g) with a standard enthalpy of reaction = +74.2 kJ, calculate the value for the standard enthalpy of reaction in the thermochemical equation, 2 MO2(s) + 1/2 O2(g) → M2O5(s)

Diff: 2

Section: 6.8

113) Given the thermochemical equation, 3 M(s) + 3 O2(g) → 3 MO2(s) with a standard enthalpy of reaction = −1443 kJ, calculate the value for ΔH°reaction for the reaction:

MO2(s) → M(s) + O2(g).

Diff: 2

Section: 6.8

114) Use these reactions and standard enthalpies, ΔH°

2 ZbO(s) + 1/2 O2(g) → Zb2O3(s) −128.0 kJ

2 ZbO(s) + 1 1/2 O2(g) → Zb2O5(s) −344.5 kJ

to find the value for

Zb2O3(s) + O2(g) → Zb2O5(s)

Diff: 2

Section: 6.8

115) Use these reactions and standard enthalpies, ΔH°

2 ZbO(s) + 1/2 O2(g) → Zb2O3(s) −128.0 kJ

2 ZbO(s) + 1 1/2 O2(g) → Zb2O5(s) −344.5 kJ

to find the value for

Zb2O3(s) + Zb2O5(s) → 4 ZbO(s) + 2 O2(g)

Diff: 2

Section: 6.8

116) Using the equation shown, 7A + 5B → 3C + 4D, and the standard enthalpies of formation,

ΔHf °:

A: 15.7 kJ mol−1 B: 86.4 kJ mol−1 C: −52.7 kJ mol−1 D: −71.6 kJ mol−1

calculate ΔH°reaction in kJ for the hypothetical reaction above.

Diff: 2

Section: 6.9

117) Using the standard enthalpies of formation, ΔH°f:

CO(g) = -110.5 kJ mol−1 CO(NH2)2(s) = −333.19 kJ mol−1 NH3(g) = −46.19 kJ mol−1

calculate ΔH°reaction for

CO(NH2)2(s) + H2(g) → 2NH3(g) + CO(g)

Diff: 2

Section: 6.9

118) Complete combustion of hydrocarbons or compounds with C, H, and O as the only elements give CO2 and H2O as the only products. If carried out under standard conditions, the CO2 is a gas and the H2O is a liquid. Given these standard enthalpies of combustion:

C2H4(g) = −1411.08 kJ mol−1, C6H12(l) = −3919.86 kJ mol−1, H2(g) = −285.90 kJ mol−1,

C(s) = −393.50 kJ mol−1, calculate ΔH°reaction for the process: 3 C2H4(g) → C6H12(l).

Hint: Remember it is always products minus reactants when performing enthalpy calculations.

Diff: 3

Section: 6.9

119) Complete combustion of hydrocarbons or compounds with C,H, and O as the only elements give CO2 and H2O as the only products. If carried out under standard conditions, the CO2 is a gas and the H2O is a liquid. Given these standard enthalpies of combustion:

CH3CHO(l) = −1166.37 kJ mol−1, C6H12O3(l) = −3340.34 kJ mol−1, H2(g) = −285.90 kJ mol−1,

C(s) = −393.50 kJ mol−1, calculate ΔH°reaction for the process, 3 CH3CHO(l) → C6H12O3(l).

Remember it is always products minus reactants when performing enthalpy calculations.

Diff: 3

Section: 6.9

120) Complete combustion of hydrocarbons or compounds with C, H, and O as the only elements give CO2 and H2O as the only products. If carried out under standard conditions, the CO2 is a gas and the H2O is a liquid. Given these standard enthalpies of combustion:

C2H2(g) = −1299.65 kJ mol−1, C6H6(l) = −3267.80 kJ mol−1, H2(g) = −285.90 kJ mol−1,

C(s) = −393.50 kJ mol−1, calculate ΔH°reaction for the process, 3 C2H2(g) → C6H6(l).

Remember it is always products minus reactants when performing enthalpy calculations.

Diff: 3

Section: 6.9

121) Using the standard enthalpies of formation, ΔH°f:

H2O(l) = −285.9 kJ mol−1; C2H4(g) = 52.284 kJ mol−1; C2H5OH(l) = −277.63 kJ mol−1

calculate ΔH°reaction for

C2H4(g) + H2O(l) → C2H5OH(l)

Diff: 2

Section: 6.9

122) Using the standard enthalpies of formation, ΔH°f:

NO(g) = 90.37 kJ mol−1 NO2(g) = 33.8 kJ mol−1

SO2(g) = −296.9 kJ mol−1 SO3(g) = −395.2 kJ mol−1

calculate ΔH°reaction for

SO2(g) + NO2(g) → SO3(g) + NO(g)

Diff: 2

Section: 6.9

123) Using the standard enthalpies of formation, ΔH°f:

CO(g) = −110.5 kJ mol−1 CO2(g) = −393.5 kJ mol−1

SO2(g) = −296.9 kJ mol−1 SO3(g) = −395.2 kJ mol−1

calculate ΔH°reaction for

CO2(g) + SO2(g) → SO3(g) + CO(g)

Diff: 2

Section: 6.9

124) Using the standard enthalpies of formation, ΔH°f:

B2O3(s) = −1,273.5 kJ mol−-1 B5H9(s) = 73.2 kJ mol−-1 H2O(g) = −241.8 kJ mol−1

calculate how much energy would be given off when 104.4 g of B5H9(s) burns in an oxygen environment to produce B2O3(s) and H2O(g).

Hint: Write out the combustion equation and balance it.

Diff: 3

Section: 6.9

125) The statement, "the total energy of the universe is constant", is a logical extension of the law of conservation of energy.

Diff: 1

Section: 6.1

126) The potential energy of an object can be changed to kinetic energy.

Diff: 1

Section: 6.1

127) Heat is energy that is transferred between particles that have different kinetic energies.

Diff: 1

Section: 6.2

128) Temperature is a measure of the average potential energy of the particles in a substance.

Diff: 2

Section: 6.2

129) A property, like energy, that depends only on an object's current state is called a state function.

Diff: 1

Section: 6.2

130) The concept of an average kinetic energy suggests that there is a range of different kinetic energies among the particles of a substance.

Diff: 1

Section: 6.2

131) A system cannot allow the transfer of thermal energy but can allow the transfer of mass.

Diff: 2

Section: 6.3

132) Isolated systems do not exchange mass or energy with the surroundings.

Diff: 1

Section: 6.3

133) As heat is transferred to ice cubes, its water molecules gain kinetic energy, causing the ice to melt.

Diff: 1

Section: 6.3

134) Processes that occur within closed systems are called adiabatic.

Diff: 2

Section: 6.3

135) The temperature change, measured in Kelvins, experienced by an object, is directly proportional to the heat it absorbs.

Diff: 2

Section: 6.3

136) If a 10 g block of metal is placed in a certain volume of water, and the metal loses 25 J of heat, then the water can only gain 22 J of heat.

Diff: 2

Section: 6.3

137) Any chemical reaction in which heat is a product is described as exothermic.

Diff: 1

Section: 6.4

138) Heat energy can always be quantitatively converted into various other forms of energy.

Diff: 2

Section: 6.4

139) In an endothermic reaction, the temperature of the system generally tends to decrease as the reaction proceeds.

Diff: 2

Section: 6.4

140) Acid-base neutralization reactions involving strong acids like HCl, and strong bases like NaOH, are generally endothermic reactions.

Diff: 1

Section: 6.4

141) The first law of thermodynamics is expressed in the equation, ΔH° = q + w.

Diff: 2

Section: 6.5

142) The first law of thermodynamics is expressed in the equation, ΔE = Kinetic Energy - Potential Energy.

Diff: 2

Section: 6.5

143) Based on the terms and symbols used in the discussions on thermochemistry, the expression, is correct.

Diff: 2

Section: 6.5

144) For a chemical reaction taking place at constant pressure in which all reactants and products are solids or liquids, ΔE ≈ qp.

Diff: 2

Section: 6.5

145) In a bomb calorimeter, no expansion work is done, so PΔV must be equivalent to, or, greater than a value of 1 at all times.

Diff: 2

Section: 6.6

146) The heat of reaction, measured in a bomb calorimeter, is the heat of reaction at constant volume and is equivalent to ΔE = qv.

Diff: 2

Section: 6.6

147) A coffee-cup calorimeter can be used to measure the heat of reaction at a constant volume.

Diff: 2

Section: 6.6

148) All substances in their "standard state" are at 1.000 atm of pressure and 298 K.

Diff: 2

Section: 6.7

149) The following thermochemical equation indicates that 184.8 kJ are released, if two moles of hydrogen react with nitrogen to form ammonia, N2(g) + 3 H2(g) → 2 NH3(g), ΔH° = -92.38 kJ.

Diff: 2

Section: 6.7

150) The following thermochemical equation indicates that 258.9 kJ are absorbed in the reaction, if one mole of hydrogen reacts with oxygen to form water, H2(g) + 1/2 O2(g) → H2O(l), ΔH° = -258.9 kJ

Diff: 2

Section: 6.7

151) It is possible for the combustion of 1 mol of methane to have different values of ΔH°, if the water produced in the combustion reaction is in the liquid or gaseous state at 25 °C.

Diff: 2

Section: 6.7

152) The combustion of butane, C4H10, is given as: 2 C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(l).

This reaction has the same heat of reaction as: 2 C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(g).

Hint: Look at the phases for each species.

Diff: 3

Section: 6.7

153) The value of ΔH° for any reaction that can be written in three steps, is equal to the sum of the values of ΔH° of each of the individual steps.

Diff: 2

Section: 6.8

154) When a thermochemical equation is reversed, the sign of ΔH° for the same reaction is also reversed.

Diff: 1

Section: 6.8

155) The formulas of substances cancelled from both sides of a thermochemical equation must be for the substance even if it is in different physical states.

Diff: 2

Section: 6.8

156) Combustion reactions generally have positive values for ΔH°.

Diff: 2

Section: 6.9

157) In combustion reactions, the carbon in the substance that is the fuel, is converted to carbon dioxide gas, while the hydrogen is converted to water.

Diff: 2

Section: 6.9

158) The standard heat of combustion of a substance is the amount of heat released when 1 mole of the substance is burned in pure oxygen gas, with all substances at 25°C at a pressure of 1 bar.

Diff: 2

Section: 6.9

159) Water has a high specific heat capacity that allows it to retain its temperature for long periods of time and helps it to resist sudden temperature changes.

Diff: 1

Section: Chemistry Outside the Classroom 6.1

160) It is known that water has a high specific heat capacity. Due to its high water content, the human body is therefore able to adjust to large and sudden changes in the outside temperature.

Diff: 2

Section: Chemistry Outside the Classroom 6.1

161) The high heat capacity of water ensures that cities and towns near large bodies of water will have cooler summers and milder winters compared to other places that are located inland.

Diff: 2

Section: Chemistry Outside the Classroom 6.1

162) The heat generated in a typical chemical reaction can be removed by pumping cold water around the outside of the vessel in which the reaction occurs.

Diff: 2

Section: Chemistry and Current Affairs 6.2

163) The hazards associated with highly exothermic chemical reactions can be alleviated by calculating the heat of the reaction from standard heats of formation, knowing how fast the heat is released, and the rate at which it can be removed from the reaction vessel.

Diff: 2

Section: Chemistry and Current Affairs 6.2

164) At elevated temperatures, any volatile substances in a chemical reaction can be quickly converted to gases, which tend to decrease the pressure within a reaction vessel.

Diff: 2

Section: Chemistry and Current Affairs 6.2

165) When the rate at which a chemical reaction generates heat is faster than the rate at which the equipment can remove the heat, the reaction is considered endothermic in nature, and the heat is used to generate more of the products.

Diff: 2

Section: Chemistry and Current Affairs 6.2

166) A large oceangoing container vessel (200,000 tons) was drifting at just 1.5 miles per hour on the tide when it struck the bridge pylon. Last year a 4500 pound speedboat struck the same pylon while going 48.5 miles per hour. Which possessed more kinetic energy? Use 1 pound = 0.4536 kg = 5.000 × 10−4 ton, 1 mile = 1.609 km.

Hint: The formula for kinetic energy is KE = mv2.

Diff: 3

Section: 6.1

167) Explain how it is possible for a person to freeze to death when the temperature of their surroundings is above their body temperature.

Hint: Focus on the kinetic energy of individual particles.

Diff: 3

Section: 6.2

168) A constant pressure calorimeter has metal parts (heat capacity of 850.0 J °C−1) and 1020 grams of oil (specific heat = 2.248 J g−1 °C−1), both at 24.50 °C. Two metal slugs, one a 460.0 g piece of cobalt and the other a 360.0 g piece of cadmium were removed from an oven that was maintained at 240.0 °C and added to the calorimeter. If no heat was lost to the surroundings, what would be the final temperature in the calorimeter? (cadmium, molar heat capacity = 25.34 J mol−1 °C−1; cobalt, molar heat capacity = 25.12 J mol−1 °C−1)

Hint: Remember to use moles as an intermediary and pay careful attention to units throughout the problem.

Diff: 3

Section: 6.3

169) A volume of 500.0 mL of 0.220 M HCl(aq) was added to a high quality constant-pressure calorimeter containing 500.0 mL of 0.200 M NaOH(aq). Both solutions have a density of and a specific heat of 4.184 J g−1 °C−1. The calorimeter had a heat capacity of 850.0 J °C−1. The temperature of the entire system rose from 25.60 °C to 26.70 °C. Calculate the heat of reaction, in kJ, per mole of NaOH(aq).

Hint: Consider your reaction taking place and pay careful attention to your mole values.

Diff: 3

Section: 6.6

170) Propane is often used to heat homes. The combustion of propane follows the following reaction:

C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g), ΔH° = −2044 kJ.

If during a winter month a houses furnace runs for one hour a day, and the furnace puts out 1,667 BTU's per minute, how many pounds of propane are needed for that month, assuming a 30 day month? Hint: 1 BTU = 1055 Joules and 1 kg = 2.2 pounds.

Hint: Use the coefficients in the equation and be sure to convert your answer to the correct units.

Diff: 3

Section: 6.7

171) A volume of 600.0 mL of 0.240 M perchloric acid, HClO4(aq) was added to a high quality constant-pressure calorimeter containing 400.0 mL of 0.300 M KOH(aq). Both solutions have a density of 1.000 g mL−1 and a specific heat of 4.184 J g−1 °C−1. The calorimeter had a heat capacity of 950.0 J °C−1. The temperature of the entire system rose from 25.30 °C to 26.59 °C. Calculate the heat of reaction, in kJ, per mole of KOH(aq).

Hint: Consider your reaction taking place and pay careful attention to your mole values.

Diff: 3

Section: 6.6

172) The formulas of ethylene, water, and ethanol suggests that the reaction,

C2H4(g) + H2O(g) → C2H5OH(l)

could be made to occur under the correct conditions. This is a wild idea, but Mike says it's just a matter of the right catalyst combination, reactor temperature, and pressure. Using the values in the table below, what is the calculated value of ΔH° for Mike's proposed reaction?

C(s) + 2 H2(g) → CH4(g) ΔH° = −74.848 kJ

H2(g) + ½ O2(g) → H2O(g) ΔH° = −241.8 kJ

H2(g) + ½ O2(g) → H2O(l) ΔH° = −285.9 kJ

C(s) + 2 H2(g) + ½ O2(g) → CH3OH(l) ΔH° = −238.6 kJ

2 C(s) + 3 H2(g) + ½ O2(g) → C2H5OH(l) ΔH° = −277.63 kJ

2 C(s) + 2 H2(g) → C2H4(g) ΔH° = +52.284 kJ

Hint: Be sure to adjust your signs whenever you reverse a reaction.

A) −39.0 kJ

B) −44.0 kJ

C) +44.0 kJ

D) −88.1 kJ

E) +88.1 kJ

Diff: 3

Section: 6.8

173) Some workers in central research were talking around the lunch table. They claim to have an idea for a process that might possibly convert C2H4(g) to C2H5OH(l), an alternative fuel. It would require two steps:

NaOH(s) + C2H4(g) NaC2H5O(s)

NaC2H5O(s) + H2O(g) → C2H5OH(l) + NaOH(aq)

It is known that for the reaction, NaOH(s) → NaOH(aq), ΔH° = −44.505 kJ. Using the values below, what is the calculated value of ΔH° for the proposed reaction?

Hint: Be sure to adjust your signs whenever you reverse a reaction.

C(s) + 2 H2(g) → CH4(g) ΔH° = −74.848 kJ

H2(g) + ½ O2(g) → H2O(g) ΔH° = −241.8 kJ

H2(g) + ½ O2(g) → H2O(l) ΔH° = −285.9 kJ

C(s) + 2 H2(g) + ½ O2(g) → CH3OH(l) ΔH° = −238.6 kJ

2 C(s) + 3 H2(g) + ½ O2(g) → C2H5OH(l) ΔH° = −277.63 kJ

2 C(s) + 2 H2(g) → C2H4(g) ΔH° = +52.284 kJ

Diff: 3

Section: 6.8

174) A check of the formulas of methane, water, and methanol suggests that the reaction,

CH4(g) + H2O(g) → CH3OH(l) + H2(g)

could be made to occur under the correct conditions. This is a wild idea, but the chemist thinks it is just a matter of the right catalyst combination, reactor temperature, and pressure. Using the values in the table below, what is the calculated value of ΔH° for Mike's proposed reaction?

C(s) + 2 H2(g) → CH4(g) ΔH° = −74.848 kJ

H2(g) + ½ O2(g) → H2O(g) ΔH° = −241.8 kJ

H2(g) + ½ O2(g) → H2O(l) ΔH° = −285.9 kJ

C(s) + 2 H2(g) + ½ O2(g) → CH3OH(l) ΔH° = −238.6 kJ

2 C(s) + 3 H2(g) + ½ O2(g) → C2H5OH(l) ΔH° = −277.63 kJ

2 C(s) + 2 H2(g) → C2H4(g) ΔH° = +52.284 kJ

Hint: Be sure to adjust your signs whenever you reverse a reaction.

A) -78.0 kJ

B) +78.0 kJ

C) -122.1 kJ

D) +122.1 kJ

E) +368.9 kJ

Diff: 3

Section: 6.8

175) The standard enthalpy of combustion for oxalic acid, H2C2O4(s), is −251.9 kJ mol−1. Using this data and the standard enthalpies of formation,

ΔH°f H2O(l) = −285.9 kJ mol−1 CO2(g) = −393.5 kJ mol−1

calculate the standard enthalpy of formation of H2C2O4(s), in kJ mol−1.

Hint: Pay careful attention to your signs when calculating the enthalpy.

Diff: 3

Section: 6.9

176) The standard enthalpy of combustion for ethylene glycol, C2H6O2(l), is −1179.5 kJ mol−1. Using this information and the standard enthalpies of formation, ΔH°f:

H2O(l) = −285.9 kJ mol−1; CO2(g) = −393.5 kJ mol−1

calculate the standard enthalpy of formation of C2H6O2(l), in kJ mol−1.

Hint: Pay careful attention to your signs when calculating the enthalpy.

Diff: 3

Section: 6.9

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