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Chapter 5: Principles of Chemical Reactivity: Energy and Chemical Reactions Study Questions: Practicing Skills
Book Title: Chemistry & Chemical Reactivity
Printed By: Cerrie Rogers (cerrie.rogers@concordia.ca)
© 2019 Cengage Learning, Cengage Learning
Chapter Review
Study Questions: Practicing Skills
denotes challenging questions. Odd-numbered questions have fully worked solutions in
the Student Solutions Manual.
Energy: Some Basic Principles
(See Section 5.1.)
1. Define the terms system and surroundings. What does it mean to say that a
system and its surroundings are in thermal equilibrium?
2. What determines the directionality of energy transfer as heat?
3. Identify whether the following processes are exothermic or endothermic. Is the
sign on
positive or negative?
a. combustion of methane
b. melting of ice
c. raising the temperature of water from
d. heating
to form
to
and
4. Identify whether the following processes are exothermic or endothermic. Is the
sign on
positive or negative?
a. the reaction of
and
b. cooling and condensing gaseous
c. cooling a soft drink from
d. heating
to form
to form liquid
to
and
Specific Heat Capacity
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(See Section 5.2 and Examples 5.1 and 5.2.)
5. The molar heat capacity of mercury is
heat capacity of this metal in
. What is the specific
?
6. The specific heat capacity of benzene
molar heat capacity (in
is
. What is its
)?
7. The specific heat capacity of copper metal is
is required to heat
. How much energy
g of copper from
to
?
8. How much energy as heat is required to raise the temperature of
water from
to
mL of
? (Density of water at this
.)
-g sample of iron is
9. The initial temperature of a
absorbs
kJ of energy as heat, what is its final temperature?
10. After absorbing
kJ of energy as heat, the temperature of a
block of copper is
11. A
. If the sample
-kg
. What was its initial temperature?
-g sample of copper at
g of water at
is dropped into a beaker containing
. What is the final temperature when thermal
equilibrium is reached?
12. One beaker contains
g of water at
g of water at
, and a second beaker contains
. The water in the two beakers is mixed. What is the
final water temperature?
13. A
-g sample of gold at some temperature was added to
The initial water temperature was
g of water.
, and the final temperature was
. If the specific heat capacity of gold is
, what was the
initial temperature of the gold sample?
14. When
g of water at a temperature of
is mixed with
g of water
at an unknown temperature, the final temperature of the resulting mixture is
. What was the initial temperature of the second sample of water?
15. A
-g piece of zinc is heated to
into a beaker containing
in boiling water and then dropped
g of water at
come to thermal equilibrium, the temperature is
. When the water and metal
. What is the specific
heat capacity of zinc?
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16. A
-g piece of molybdenum, initially at
water at
, is dropped into
g of
. When the system comes to thermal equilibrium, the
temperature is
. What is the specific heat capacity of molybdenum?
Changes of State
(See Section 5.3 and Examples 5.3 and 5.4.)
17. How much energy is evolved as heat when
ice? (The heat of fusion of water is
18. The energy required to melt
a mass of
L of water at
J/g.)
g of ice at
g and a tray contains
is
J. If one ice cube has
ice cubes, what quantity of energy is
required to melt a tray of ice cubes to form liquid water at
19. How much energy is required to vaporize
boiling point,
solidifies to
?
g of benzene,
, at its
? (The heat of vaporization of benzene is
20. Chloromethane,
kJ/mol.)
, arises from microbial fermentation and is found
throughout the environment. It is also produced industrially, is used in the
manufacture of various chemicals, and has been used as a topical anesthetic.
How much energy is required to convert
point,
g of liquid to a vapor at its boiling
? (The heat of vaporization of
21. The freezing point of mercury is
is released to the surroundings if
is
kJ/mol.)
. What quantity of energy, in joules,
mL of mercury is cooled from
to
and then frozen to a solid? (The density of liquid mercury is
. Its specific heat capacity is
and its heat of fusion is
J/g.)
22. What quantity of energy, in joules, is required to raise the temperature of
g of tin from room temperature,
, to its melting point,
, and
then melt the tin at that temperature? (The specific heat capacity of tin is
, and the heat of fusion of this metal is
23. Ethanol,
, boils at
to raise the temperature of
J/g.)
. How much energy, in joules, is required
kg of ethanol from
to the boiling point
and then to change the liquid to vapor at that temperature? (The specific heat
capacity of liquid ethanol is
, and its enthalpy of vaporization is
J/g.)
24. A
-mL sample of benzene at
was cooled to its melting point,
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, and then frozen. How much energy was given off as heat in this
process? (The density of benzene is
g/mL, its specific heat capacity is
, and its heat of fusion is
J/g.)
Heat, Work, and Internal Energy
(See Section 5.4 and Example 5.5.)
25. As a gas cools, it is compressed from
pressure of
L to
L under a constant
Pa. Calculate the work (in J) required to compress the
gas.
26. A balloon expands from
pressure of
L to
L as it is heated under a constant
Pa. Calculate the work (in J) done by the balloon on
the environment.
27. A balloon does
J of work on the surroundings as it expands under a
constant pressure of
Pa. What is the change in volume (in L) of the
balloon?
28. As the gas trapped in a cylinder with a movable piston cools,
kJ of work
is done on the gas by the surroundings. If the gas is at a constant pressure of
Pa, what is the change of volume (in L) of the gas?
29. When
J of energy in the form of heat is transferred from the environment
to a gas, the expansion of the gas does
J of work on the environment.
What is the change in internal energy of the gas?
30. The internal energy of a gas decreases by
kJ when it transfers
kJ of
energy in the form of heat to the surroundings.
a. Calculate the work done by the gas on the surroundings.
b. Does the volume of gas increase or decrease?
31. A volume of
L of argon gas is confined in a cylinder with a movable
piston under a constant pressure of
Pa. When
kJ of energy in
the form of heat is transferred from the surroundings to the gas, the internal
energy of the gas increases by
kJ. What is the final volume of argon gas
in the cylinder?
32. Nitrogen gas is confined in a cylinder with a movable piston under a constant
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Pa. When
pressure of
J of energy in the form of heat is
transferred from the gas to the surroundings, its volume decreases by
L.
What is the change in internal energy of the gas?
Enthalpy Changes
(See Section 5.5 and Example 5.6.)
33. Nitrogen monoxide, a gas recently found to be involved in a wide range of
biological processes, reacts with oxygen to give brown
gas.
Is this reaction endothermic or exothermic? What is the enthalpy change if
g of
is converted completely to
34. Calcium carbide,
?
, is manufactured by the reaction of
with carbon
at a high temperature. (Calcium carbide is then used to make acetylene.)
Is this reaction endothermic or exothermic? What is the enthalpy change if
g of
is allowed to react with an excess of carbon?
35. Isooctane (2,2,4-trimethylpentane), one of the many hydrocarbons that make
up gasoline, burns in air to give water and carbon dioxide.
What is the enthalpy change if you burn
36. Acetic acid,
L of isooctane
?
, is made industrially by the reaction of methanol and
carbon monoxide.
What is the enthalpy change for producing
L of acetic acid
by this reaction?
Calorimetry
(See Section 5.6 and Examples 5.7 and 5.8.)
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37. Assume you mix
mL of
M
with
mL of
M
in a
coffee-cup calorimeter. The following reaction occurs:
The temperature of both solutions before mixing was
, and it rises to
after the acid–base reaction. What is the enthalpy change for the
reaction per mole of
? Assume the densities of the solutions are all
g/mL and the specific heat capacities of the solutions are
38. You mix
mL of
M
with
mL of
M
.
in a coffee-cup
calorimeter, and the temperature of both solutions rises from
mixing to
after the reaction.
What is the enthalpy of reaction per mole of
the solutions are all
are
before
? Assume the densities of
g/mL, and the specific heat capacities of the solutions
.
39. A piece of titanium metal with a mass of
g is heated in boiling water to
and then dropped into a coffee-cup calorimeter containing
water at
g of
. When thermal equilibrium is reached, the final temperature is
. Calculate the specific heat capacity of titanium.
40. A piece of chromium metal with a mass of
to
g is heated in boiling water
and then dropped into a coffee-cup calorimeter containing
water at
g of
. When thermal equilibrium is reached, the final temperature is
. Calculate the specific heat capacity of chromium.
41. Adding
g of
to
g of water in a coffee-cup calorimeter
(with stirring to dissolve the salt) resulted in a decrease in temperature from
to
. Calculate the enthalpy change for dissolving
in water, in kJ/mol. Assume the solution (whose mass is
specific heat capacity of
g) has a
. (Cold packs take advantage of the fact
that dissolving ammonium nitrate in water is an endothermic process.)
A cold pack uses the endothermic enthalpy of solution of
ammonium nitrate.
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© Cengage Learning/Charles D. Winters
42. You should use care when dissolving
in water because the process is
highly exothermic. To measure the enthalpy change,
was added (with stirring) to
g of concentrated
g of water in a coffee-cup
calorimeter. This resulted in an increase in temperature from
to
. Calculate the enthalpy change for the process
, in kJ/mol.
43. Sulfur (
g) was burned in a constant-volume calorimeter with excess
. The temperature increased from
heat capacity of
Calculate
to
. The bomb has a
J/K, and the calorimeter contained
per mole of
g of water.
formed for the reaction
.
Sulfur burns in oxygen with a bright blue flame to give
© Cengage Learning/Charles D. Winters
44. Suppose you burned
g of
volume calorimeter to give
in an excess of
.
The temperature of the calorimeter, which contained
from
to
in a constant-
g of water, increased
. The heat capacity of the bomb is
J/K. Calculate
per mole of carbon.
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45. Suppose you burned
g of benzoic acid,
, in a constant-
volume calorimeter and found that the temperature increased from
. The calorimeter contained
capacity of
J/K. Calculate
to
g of water, and the bomb had a heat
per mole of benzoic acid.
, occurs naturally in many berries. Its
Benzoic acid,
heat of combustion is well known, so it is used as a standard to
calibrate calorimeters.
46. A
-g sample of glucose,
, was burned in a constant-volume
calorimeter. The temperature rose from
contained
is
to
. The calorimeter
g of water, and the bomb had a heat capacity of
J/K. What
per mole of glucose?
47. An “ice calorimeter” can be used to determine the specific heat capacity of a
metal. A piece of hot metal is dropped onto a weighed quantity of ice. The
energy transferred from the metal to the ice can be determined from the
amount of ice melted. Suppose you heated a
-g piece of silver to
and then dropped it onto ice. When the metal’s temperature had dropped to
, it is found that
g of ice had melted. What is the specific heat
capacity of silver?
48. A
-g piece of platinum was heated to
in a boiling water bath and
then dropped onto ice. (See Study Question 47.) When the metal’s
temperature had dropped to
, it was found that
g of ice had melted.
What is the specific heat capacity of platinum?
Hess’s Law
(See Section 5.7 and Example 5.9.)
49. The enthalpy changes for the following reactions can be measured:
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a. Use these values and Hess’s law to determine the enthalpy change for
the reaction
b. Draw an energy level diagram that shows the relationship between the
energy quantities involved in this problem.
50. The enthalpy changes of the following reactions can be measured:
a. Use these values and Hess’s law to determine the enthalpy change for
the reaction
b. Draw an energy level diagram that shows the relationship between the
energy quantities involved in this problem.
51. Enthalpy changes for the following reactions can be determined
experimentally:
Use these values to determine the enthalpy change for the formation of
from the elements (an enthalpy change that cannot be measured directly
because the reaction is reactant-favored).
52. You wish to know the enthalpy change for the formation of liquid
from
the elements.
The enthalpy change for the formation of
from the elements can be
determined experimentally, as can the enthalpy change for the reaction of
with more chlorine to give
:
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Use these data to calculate the enthalpy change for the formation of
of
mol
from phosphorus and chlorine.
Standard Enthalpies of Formation
(See Section 5.7 and Example 5.10.)
53. Write a balanced chemical equation for the formation of
elements in their standard states. Find the value for
from the
for
in
Appendix L.
54. Write a balanced chemical equation for the formation of
elements in their standard states. Find the value for
from the
for
in
Appendix L.
55.
a. Write a balanced chemical equation for the formation of
from
and
for
mol of
in their standard states. (Find the value for
in Appendix L.)
b. What is the standard enthalpy change if
g of chromium is oxidized to
?
56.
a. Write a balanced chemical equation for the formation of
mol of
from the elements in their standard states. (Find the value for
for
in Appendix L.)
b. What is the standard enthalpy change for the reaction of
mol of
with oxygen?
57. Use standard enthalpies of formation in Appendix L to calculate enthalpy
changes for the following:
a.
g of white phosphorus burns, forming
b.
mol of
c.
g of
decomposes to
is formed from
and
and excess
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g of iron is oxidized with oxygen to
d.
58. Use standard enthalpies of formation in Appendix L to calculate enthalpy
changes for the following:
a.
g of sulfur burns, forming
b.
mol of
c.
g of
decomposes to
is formed from
d.
and
and excess
mol of carbon is oxidized to
59. The first step in the production of nitric acid from ammonia involves the
oxidation of
.
a. Use standard enthalpies of formation to calculate the standard enthalpy
change for this reaction.
b. How much energy is evolved or absorbed as heat in the oxidation of
g of
?
60. The Romans used calcium oxide,
, to produce a strong mortar to build
stone structures. Calcium oxide was mixed with water to give
reacted slowly with
in the air to give
, which
.
a. Calculate the standard enthalpy change for this reaction.
b. How much energy is evolved or absorbed as heat if
reacts with a stoichiometric amount of
kg of
?
61. The standard enthalpy of formation of solid barium oxide,
, is
kJ/mol, and the standard enthalpy of formation of barium peroxide,
, is
kJ/mol.
a. Calculate the standard enthalpy change for the following reaction. Is the
reaction exothermic or endothermic?
b. Draw an energy level diagram that shows the relationship between the
enthalpy change of the decomposition of
to
and
and the
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and
enthalpies of formation of
.
62. An important step in the production of sulfuric acid is the oxidation of
to
.
Formation of
from the air pollutant
is also a key step in the formation
of acid rain.
a. Use standard enthalpies of formation to calculate the enthalpy change
for the reaction. Is the reaction exothermic or endothermic?
b. Draw an energy level diagram that shows the relationship between the
enthalpy change for the oxidation of
formation of
and
to
and the enthalpies of
.
63. The enthalpy change for the oxidation of naphthalene,
, is measured
by calorimetry.
Use this value, along with the standard enthalpies of formation of
and
, to calculate the enthalpy of formation of naphthalene, in kJ/mol.
64. The enthalpy change for the oxidation of styrene,
, is measured by
calorimetry.
Use this value, along with the standard enthalpies of formation of
and
, to calculate the enthalpy of formation of styrene, in kJ/mol.
Chapter 5: Principles of Chemical Reactivity: Energy and Chemical Reactions Study Questions: Practicing Skills
Book Title: Chemistry & Chemical Reactivity
Printed By: Cerrie Rogers (cerrie.rogers@concordia.ca)
© 2019 Cengage Learning, Cengage Learning
© 2020 Cengage Learning Inc. All rights reserved. No part of this work may by reproduced or used in any form or by any means graphic, electronic, or mechanical, or in any other manner - without the written permission of the copyright holder.
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