Chemistry I Accelerated
Study Guideline - Chapter 17
Thermochemistry
____________________________________________________
By the end of this chapter the skills you should be able to demonstrate
are:
1. Determine the heat of a chemical reaction in which a specified
amount of substance is involved.
2. State the reason enthalpy changes in a chemical reaction.
3. Calculate enthalpies of formation.
4. Calculate changes in heat of reaction (enthalpy).
5. Describe and give examples of changes in entropy.
6. Calculate changes in entropy.
7. State two reasons reactions occur.
8. Relate Gibbs free energy to the spontaneity of reactions and to
equilibrium.
9. Calculate Gibbs free energy changes involving enthalpy and
entropy.
Suggested Problems: p.535-537: #56, 59, 61, 66, 68, 73, 74, 82
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Calorimetry of Reactions
An insulated container
that is used to make
heat measurements is
called a calorimeter.
Answer the questions
below based on the
figure shown to the
right.
1. Was the reaction that occurred in the reaction chamber, endothermic or exothermic? How can you tell?
2. How much heat was involved in the reaction within the reaction chamber in joules?
3. If the reaction involved 10.4 moles of reactant, how many kilojoules of heat were involved per mole of
reactant?
page 2
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Heat of Reaction - Enthalpy (ΔH)
1. Define Enthalpy Change
2. Energy changes are associated with every chemical reaction. When chemical bonds are _____________
(formed/broken), energy is consumed. When chemical bonds are _______________ (formed/broken),
energy is given off. The net sum of the energy changes related to these bond rearrangements is called the
enthalpy change and has the symbol ________________.
3. When energy is produced in a chemical reaction, the reaction is ____________ (endothermic/exothermic).
The change in enthalpy, __________, may be written on the right side of the equation. __________ is
usually reported in kilocalories or kilojoules per mole of one of the reactants.. For instance in the
equation:
N2 (g) + 3H2 (g) - - > 2NH3 (g) + 92 kJ
when 1 mole of __________ reacts with 3 moles of ________ to form _______ moles of ammonia,
_______ kJ of energy is released to the surrounding
4. When energy is absorbed in a reaction, the reaction is called ______________ (endothermic/exothermic).
The change in enthalpy, ΔH, is written on the reactants side of the equation. The ΔH again may be
reported in kilojoules per mole of one of the reactants or products. For example, in the equation:
67.8 kJ + N2 (g) + 2O2 (g) - - > 2NO2 (g)
when one mole of _______ gas reacts with _______ of oxygen gas, ____________ kJ of energy is
absorbed from the surroundings and stored in ________ moles of the product.
5. For exothermic reactions, ΔH, values are ___________ (positive/negative). For endothermic reactions,
ΔH, values are ___________ (positive/negative). For the reaction in question 3, ΔH = ____________ kJ
per mole of ammonia gas. For the reaction in question 4, ΔH = __________ kJ per mole of nitrogen
dioxide gas.
page 3
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
6. The graph below is a potential energy diagram for the hypothetical reaction A+B -> C+D
a. Is the forward reaction endothermic or exothermic? Calculate the value of ΔH for this reaction
b. Is the reverse reaction endothermic or exothermic? What is the value of ΔH for this reaction
c. What is the value of potential energy of the activated complex?
d. Calculate the activation energy for the forward reaction?
400
7. On the graph to the right, draw a potential
energy diagram for the following reaction:
300
Q + R --> S + T
given the following: potential energy of Q + R
PE
is 150 kJ; potential energy of S + T is 250 kJ;
potential energy of the activated complex is
375 kJ
200
100
Reaction Coordinate
b. Is the forward reaction endothermic or exothermic? Calculate the value of ΔH for this reaction
c. Calculate the activation energy for the forward reaction?
page 4
ChemIAcc-22Thermochemistry WS
140
8. On the graph to the right, draw a potential
energy diagram for the following reaction:
100
E + F --> G + H
given the following: potential energy of E + F
Dr. Corell - Chemistry I Accelerated
PE
is 72 kcal; potential energy of G + H is 112
kcal; activation energy of the forward reaction
is 58 kcal
60
20
Reaction Progress
b. What is the total potential energy of the activated complex?
c. Is the forward reaction endothermic or exothermic? Calculate the value of ΔH for this reaction
d. What is the activation energy of the reverse reaction?
Standard Heat of Formation
When one mole of a compound is formed from its elements, the heat of reaction is called the
standard heat of formation. For example, when one mole of carbon dioxide is formed from its elements,
carbon and oxygen, the heat of the reaction, ΔH, is –393.7 kJ. Thus the value of the standard heat of
0
formation , with the symbol ΔHf , for CO2 = –393.7 kJ. You will need a reference table of standard heats of
formation found starting on the next page to answer the following questions.
0
1. The value of ΔHf for magnesium oxide is _____________ kJ /mol. This means for the reaction:
____________(s) + _____________(g) - - - > ____ MgO(s)
the enthalpy of 1 mole of the product, _______, is 601.6 kJ_________(greater/less) than the sum of the
enthalpies of the reactants, _______ mole of magnesium metal and _______ mole of oxygen gas.
0
2. The value of ΔHf for ethene is __________ kcal/mole. This means that for the reaction
____________(s) + _____________(g) - - - > C2 H4 (g)
the enthalpy of 1 mole of the product, _______, is 52.3 kJ _________(greater/less) than the sum of the
enthalpies of the reactants, _______ mole(s) of solid carbon and _______ mole(s) of hydrogen gas.
page 5
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Molar Thermodynamic Properties of Pure Substances
0
Formula
Ag (c)
Ag2 CO3 (c)
Ag2 O (c)
0
ΔH f
kJ/mol
0.0
-505.8
0
0
ΔG f
kJ/mol
0.0
S
J/K mol
42.55
Cp
J/K mol
25.351
-436.8
167.4
112.26
-31.05
-11.20
121.3
65.86
Ag2 S (c, argentite)
AgCN (c)
AgCNS (c)
AgCl (c, cerargyrite)
AgBr (c)
AgI (c)
-32.59
146.0
87.9
-127.068
-100.37
-61.83
-40.67
156.9
101.39
-109.789
-96.90
-66.19
144.01
107.19
131.0
96.2
107.1
115.5
76.53
66.73
63.0
50.79
52.38
56.82
AgNO3 (c)
-124.39
-33.47
140.92
93.05
-879.0
---
---
Ag3 PO4 (c)
---
Ag2 CrO4 (c)
-731.74
-641.76
217.6
142.26
Ag2 SO4 (c)
Al (c)
-715.88
0.0
-618.41
0.0
200.4
28.33
131.38
24.35
Al(OH) 3
-1276.0
---
---
---
AlCl3 (c)
-704.2
-628.8
110.67
91.84
AlCl3 (g)
-583.2
---
---
---
Al2 O3 (c, alumina, alpha)
B (c)
-1675.7
0.0
-1582.3
0.0
50.92
5.86
79.04
11.09
BF3 (g)
-1137.0
-1120.35
254.01
50.46
BaCO3 (c, witherite)
-1216.3
-1137.6
112.1
85.35
BaC2 O4 (c)
-1368.6
---
---
---
BaCrO4 (c)
-1446.0
-1345.22
158.6
---
BaF2 (c)
-1207.1
-1156.8
96.36
71.21
BaSO4 (c)
Bi (c)
-1473.2
0.0
-1362.2
0.0
132.2
56.74
101.75
25.52
Bi2 S3 (c)
-143.1
-140.6
200.4
122.2
0.0
0.0
152.231
75.689
Br2 (g)
C (c, graphite)
C (c, diamond)
C (g)
CO (g)
30.907
0.0
1.895
716.682
-110.525
3.110
0.0
2.900
671.257
-137.168
245.463
5.740
2.377
158.096
197.674
36.02
8.527
6.113
20.838
29.42
CO2 (g)
-393.509
-394.359
213.74
37.11
COCl2 (g, phosgene)
-218.8
-204.6
283.53
57.66
Br2 (l)
CH4 (g, methane)
-74.81
-50.72
186.264
35.309
C2 H2 (g, ethyne)
226.73
209.20
200.94
43.93
page 6
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
0
0
0
0
ΔH f
kJ/mol
ΔG f
kJ/mol
S
J/K mol
C2 H4 (g, ethene)
52.25
68.12
219.45
43.56
C2 H6 (g, ethane)
-84.68
-32.82
229.60
52.63
C3 H6 (g, propene)
20.2
62.72
266.9
64.0
C3 H8 (g, propane)
-104.5
-23.4
269.9
7.0
C4 H1 0 (g, n-butane)
-126.5
-17.15
310.1
97.4
C5 H1 2 (g, n-pentane)
-146.5
-8.37
348.9
120.2
C8 H1 8 (g, octane)
-208.5
16.40
466.7
189.0
CH3 OCH3 (g, dimethyl ether)
-184.05
-112.59
266.38
64.39
CH3 OH (g, methanol)
-200.66
-162.00
239.70
43.89
CH3 OH (l, methanol)
-238.66
-166.36
126.8
81.6
C2 H5 OH (g, ethanol)
-235.10
-168.49
282.70
65.44
C2 H5 OH (l, ethanol)
-277.69
-174.78
160.7
111.46
CH3 COOH (l, acetic acid)
-484.51
-389.9
159.8
124.3
CH3 CHO (l, acetaldehyde)
-192.30
-128.20
160.2
---
-80.83
-57.37
234.58
40.75
CHCl3 (g, chloroform)
-103.14
-70.34
295.71
65.69
CCl4 (l, carbon tetrachloride)
-135.44
-65.27
216.40
131.75
Formula
CH3 Cl (g)
Cp
J/K mol
C6 H6 (g, benzene)
82.9
129.7
269.2
81.6
C6 H6 (l, benzene)
49.0
124.7
172.0
132.0
C6 H1 2 (l, cyclohexane)
CaO (c)
-156.3
-635.09
26.7
-604.03
204.4
39.75
157.7
42.80
Ca(OH)2 (c)
-986.09
-898.49
83.39
87.49
CaCO3 (c)
-1206.92
-1128.79
92.9
81.88
CaC2 O4 (c)
CaF2 (c)
-1360.6
-1219.6
---1167.3
--68.87
--67.03
Ca3 (PO4 ) 2 (c)
-4109.9
-3884.7
240.91
231.58
CaSO4 (c, anhydrite)
Cd (c)
Cd (g)
-1434.11
0.0
2623.64
-1321.79
0.0
---
106.7
51.76
---
99.66
25.98
---
-560.7
-161.9
-473.6
-156.5
96.0
64.9
0.0
0.0
223.066
33.907
ClO2 (g)
Cu (c)
102.5
0.0
120.5
0.0
256.84
33.150
41.97
24.35
CuC2 O4 (c)
---
-661.8
---
---
-1051.4
-893.6
186.2
---
Cd(OH)2 (c)
Cds (c)
Cl2 (g)
CuCO3 .Cu(OH)2 (c, malachite)
page 7
-----
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
0
Formula
0
ΔH f
kJ/mol
ΔG f
kJ/mol
Cu2 O (c, cuprite)
CuO (c, tenorite)
-168.6
-157.3
-146.0
-129.7
Cu(OH)2 (c)
-449.8
0
S
J/K mol
0
Cp
J/K mol
93.14
42.63
63.64
42.30
---
---
---
-79.5
-53.1
-86.2
-53.6
120.9
66.5
76.32
47.82
F2 (g)
Fe (c)
FeO (c, wuestite)
0.0
0.0
-266.27
0.0
0.0
-245.12
202.78
27.28
57.49
31.30
25.10
48.12
Fe2 O3 (c, hematite)
-824.2
-742.2
87.40
103.85
Fe3 O4 (c, magnetite)
-1118.4
-1015.4
146.4
-823.0
-696.5
106.7
---
Fe3 C (c, cementite)
25.1
20.1
104.6
105.9
FeCO3 (c, siderite)
FeS (c, pyrrhotite)
-740.57
-100.0
-666.67
-100.0
92.9
60.29
82.13
50.54
FeS2 (c, pyrite)
-178.2
-166.9
52.93
62.17
FeSO4 (c)
-928.4
-820.8
-2581.5
---
Cu2 S (c, chalcocite)
CuS (c, covellite)
Fe(OH)3 (c)
Fe2 (SO4 ) 3 (c)
H2 (g)
107.5
143.43
100.58
---
---
0.0
0.0
130.684
28.824
H2 O (g)
-241.818
-228.572
188.825
33.577
H2 O (l)
-285.830
-237.129
69.91
75.291
H2 O2 (g)
-136.31
-105.57
232.7
43.1
H2 O2 (l)
-187.78
-120.35
109.6
89.1
H2 S (g)
-20.63
-33.56
105.79
34.23
H2 SO4 (l)
HF (g)
HCl (g)
HBr (g)
HI (g)
HCN (g)
Hg (l)
-813.989
-271.1
-92.307
-36.40
26.48
135.1
0.0
-690.003
-273.2
-95.299
-53.45
1.70
124.7
0.0
156.904
173.779
186.908
198.695
206.594
201.78
76.02
138.91
29.133
29.12
29.142
29.158
35.86
27.983
HgCl2 (c)
-224.3
-178.6
146.0
---
Hg2 Br2 (c)
-206.90
-181.075
218.0
---
Hg2 Cl2 (c)
HgS (c, red)
HgS (c, black)
-265.22
-58.2
-53.6
-210.745
-50.6
-47.7
192.5
82.4
88.3
--48.41
---
Hg2 SO4 (c)
-743.12
-625.815
200.66
131.96
I 2 (c)
0.0
0.0
116.135
54.438
I 2 (g)
62.438
19.327
260.69
36.90
page 8
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
0
0
0
0
Formula
ICl (g)
K (c)
KF (c)
KCl (c)
KBr (c)
KI (c)
ΔH f
kJ/mol
17.78
0.0
-567.27
-436.747
-393.798
-327.900
ΔG f
kJ/mol
-5.46
0.0
-537.75
-409.14
-380.66
-324.892
S
J/K mol
247.551
64.18
66.57
82.59
95.90
106.32
Cp
J/K mol
35.56
29.58
49.04
51.30
52.30
52.93
KClO4 (c)
-432.75
303.09
151.0
112.38
KNO3 (c)
Mg (c)
-494.63
0.0
-394.86
0.0
133.05
32.68
96.40
24.89
MgF2 (c)
-1123.4
-1070.2
57.24
61.59
MgCO3 (c, magnesite)
-1095.8
-1012.1
65.7
75.52
Mg(OH)2 (c)
Mn (c)
-924.54
0.0
-833.51
0.0
63.18
32.01
77.03
26.32
MnO2 (c)
MnS (c, green)
-520.03
-214.2
-465.14
-218.4
53.05
78.2
54.14
49.96
N2 (g)
0.0
0.0
191.61
29.125
-46.11
-16.45
192.45
35.06
-314.43
90.25
-202.87
86.55
94.6
210.761
84.1
29.844
NO2 (g)
33.18
51.31
240.06
37.20
N2 O (g)
82.05
104.20
219.85
38.45
N2 O4 (g)
9.16
97.89
304.29
77.28
N2 O4 (l)
-19.50
97.54
209.2
142.7
N2 O5 (g)
11.3
355.7
84.5
N2 O5 (c)
NOCl (g)
NOBr (g)
Na (c)
NaF (c)
NaCl (c)
NaBr (c)
NaI (c)
-43.1
51.71
82.17
0.0
-573.647
-411.153
-361.062
-287.78
178.2
261.69
273.66
51.21
51.46
72.13
86.82
98.53
143.1
44.69
45.48
28.24
46.86
50.50
51.38
52.09
112.30
NH3 (g)
NH4 Cl (c)
NO (g)
115.1
113.9
66.08
82.42
0.0
-543.494
-384.138
-348.983
-286.06
Na2 CO3 (c)
-1130.68
-1044.44
134.98
NaNO2 (c)
-358.65
-284.55
103.8
---
NaNO3 (c)
-467.85
-367.00
116.52
92.88
Na2 O (c)
NiS (c)
-414.22
-82.0
-375.46
-79.5
75.06
52.97
69.12
47.11
205.138
29.355
238.93
41.09
39.20
23.840
O2 (g)
O3 (g, ozone)
P (c)
0.0
0.0
142.7
0.0
163.2
0.0
page 9
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
0
Formula
0
ΔH f
kJ/mol
0
ΔG f
kJ/mol
S
J/K mol
0
Cp
J/K mol
PH3 (g)
5.4
13.4
210.23
37.11
PCl3 (g)
-287.0
-267.8
311.78
71.84
PCl5 (g)
Pb (c)
-374.9
0.0
-305.0
0.0
364.58
64.81
112.80
26.44
PbBr2 (c)
-278.9
-261.92
161.5
80.12
PbCl2 (c)
PbO (c, red)
PbO (c, yellow, litharge)
-359.41
-218.99
-217.32
-314.10
-189.93
-187.89
-136.0
66.5
68.70
--45.81
45.77
PbO2 (c)
-277.4
-217.33
68.6
64.64
Pb3 O4 (c)
-718.4
-601.2
211.3
Pb(OH)2 (c)
PbS (c, galena)
---100.4
-452.2
-98.7
--91.2
--49.50
PbSO4 (c)
S (c)
-919.94
0.0
-813.14
0.0
148.57
31.80
103.207
22.64
291.82
97.28
SF6 (g)
-1209.0
-1105.3
146.9
SO2 (g)
-296.830
-300.194
248.22
39.87
SO3 (g)
-395.72
-371.06
256.76
50.67
SO3 (l)
-441.04
-373.75
113.8
---
SO2 Cl2 (g)
Sn (c)
SnO (c)
-364.0
0.0
-285.8
-320.0
0.0
-256.9
311.94
51.55
56.5
77.0
26.99
44.31
SnO2 (c, cassiterite)
SnS (c)
Tl (c)
W (c)
-580.7
-100.0
0.0
0.0
-519.6
-98.3
0.0
0.0
52.3
77.0
64.18
32.64
52.59
49.25
26.32
24.27
WO2 (c)
-589.69
-533.89
50.54
56.11
WO3 (c)
-842.87
-764.03
75.90
73.76
0.0
-348.28
-205.98
0.0
-318.30
-201.29
41.63
43.64
57.7
25.40
40.25
46.03.
Zn (c)
ZnO (c)
ZnS (c)
3. Using the standard heats of formation, write the equation for the formation of ethyne from its elements
including the heat as one of the components of the reaction. Is the reaction endothermic or exothermic?
page 10
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
4. Using the standard heats of formation, write the equation for the formation of sulfur dioxide from its
elements including the heat as one of the components of the reaction. Is the reaction endothermic or
exothermic?
5. Using the standard heats of formation, write the equation for the formation of silver chloride from its
elements including the heat as one of the components of the reaction. Is the reaction endothermic or
exothermic?
6a. Using the standard heats of formation, write the equation for the formation of methane from its elements
including the heat as one of the components of the reaction. Is the reaction endothermic or exothermic?
b. Calculate the heat of reaction (ΔH) when 14.0 g of methane is produced from its elements?
7. Using the standard heat of formation, calculate the heat of reaction (ΔH) when 112.0 grams of sodium
chloride is made from its elements. Is energy absorbed or released? How can you tell?
page 11
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
8. Using the standard heats of formation, calculate the heat of reaction (ΔH) in calories when 110.0 grams of
potassium bromide is made from its elements.
9. Calculate the heat of reaction (ΔH), in kJ for the decomposition of Carbon Dioxide gas into Carbon
Monoxide gas and Oxygen gas. Is the reaction endothermic or exothermic? How can you tell?
10. Calculate the heat of reaction (ΔH), in kJ for the reaction where Nitrogen Dioxide gas decomposes to
Nitrogen Monoxide gas and Oxygen gas. Is the reaction endothermic or exothermic? How can you tell?
11. Calculate the heat of reaction (ΔH), in kJ for the reaction where Nitrogen Monoxide gas combines with
Hydrogen gas to form ammonia gas and Oxygen gas. Is the reaction endothermic or exothermic? How
can you tell?
3
12. The combustion of 17.8 L of ethene at STP liberates 1.104 x 10 kJ. Calculate the heat of reaction
(ΔH) (kJ/mol).
page 12
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Hess’s Law
The heat change associated with a reaction can be calculated if information on other
reactions containing the same reactants or products is available.
1.
Calculate the heat of reaction (ΔH) for the reaction: A2 + B - - - > A2B given the following
information:
AB + A - - - > A2 + B
ΔH = +27.1 kJ/mol
A2B - - - > AB + A
ΔH = –30.4 kJ/mol
2.
Is the reaction endothermic or exothermic? How can you tell?
3.
What would be the value for ΔH for the reverse reaction: A2B - - - > A2 + B?
4.
What would be the value of ΔH for the forward reaction if 3 moles of each reactant and product were
involved?
5.
Calculate the ΔH for the reaction sulfur dioxide gas plus hydrogen gas plus oxygen gas yields liquid
sulfuric acid given the following information:
Solid sulfur plus oxygen gas yields sulfur dioxide gas
ΔH = – 297.1 kJ
Hydrogen gas plus solid sulfur plus oxygen gas yields liquid sulfuric acid
ΔH = – 811.7 kJ
6.
Calculate the ΔH for the reaction in terms of moles of propane for the combustion of propane given the
following information:
Solid carbon plus hydrogen gas yields propane gas
ΔH = – 103.8 kJ
Solid carbon plus oxygen gas yields carbon dioxide gas
ΔH = – 393.5 kJ
Hydrogen gas plus 1/2 O2 (g) yields water vapor
page 13
ΔH = – 241.8 kJ
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
7. Calculate the ΔH for the reaction in terms of moles of ammonia for the reaction of ammonia gas plus
oxygen gas yields nitrogen monoxide gas water vapor given the following information:
1/2 N2 (g) + 1/2 O2 (g) - - > NO(g)
ΔH = + 90.4 kJ/mol
1/2 N2 (g) + 3/2 H2 (g) - - > NH3 (g)
H2 (g) + 1/2 O2 (g) - - > H2 O(g)
ΔH = – 46.0 kJ/mol
ΔH = – 393.3 kJ/mol
Entropy
On the line at the left, rank each item in this list from 1 to 3, with 1 being the item with the
greatest entropy and 3 being the item with the least entropy.
____ 1. an ice cube.
____ 2. a solution of NaOH and water
____ 3. a flask of air
On the line at the left, write the sign (pos. or neg.) of the entropy change (ΔS) for the listed
reaction or event. If there is not enough information to determine the sign of ΔS, write the
letter N on the line.
____ 1. AB + CD - - - - > AC + BD where AB and CD are both solids and AC and BD are both
liquids.
____ 2. AB + CD - - - - > AC + BD where AB and CD are both gases and AC and BD are both
liquids.
____ 3. Dry ice (solid CO2) is exposed to room temperature air.
____ 4. Steam condenses on a lid covering a pot of boiling water.
____ 5. Solid Ammonium Chloride decomposes to gaseous Ammonia and gaseous Hydrochloric Acid.
____ 6. Table salt dissolves in water to form an aqueous solution.
____ 7. Water becomes ice cubes in the freezer
____ 8. Water is heated to a boil and becomes steam.
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Enthalpy and Entropy
Complete each of the following sentences in the space provided by filling in the appropriate
word or phrase from the list below:
Endothermic reaction
Free Energy of Formation
Heat of reaction
enthalpy
Free Energy
Hess’s Law
entropy
Gibbs Equation
Kelvin Temperature
exothermic reaction
Heat of Formation
standard heat of formation
The heat content of a substance is called its _________________ . The change in this quantity that
occurs during a chemical reaction is called the ______ ___ _____________, ΔH. If the chemical change is
the production of a compound from its elements, this quantity is called the ______ ____ ______________.
At 298K and 101.3 kPa , this quantity is called the _____________ ______ ____ _______________ . The
sign of the quantity ΔH is positive in the case of an __________________ _______________. It is
negative in the case of an __________________ _______________.
When a reaction can be expressed as the algebraic sum of two or more other reactions, its heat of
reaction is equal to the algebraic sum of the heats of reaction of these other reactions. This relationship is
called ____________ __________.
A measure of the randomness of a system is its _______________. If, in any reaction, the change in this
quantity is multiplied by the ____________ ___________________ and subtracted from ΔH for the
reaction, the result is called the __________ _____________ of the reaction. If the reaction is the
production of a compound from its elements, the above result is called the _________ ___________ ______
________________ of the compound. This relationship is expressed in the __________ _____________.
Spontaneous Reactions
1. Entropy changes can be represented by the equation: ΔS = Sproducts – Sreactants
What is the meaning of this equation?
2. Describe how the two factors, entropy change and enthalpy change, determine whether a reaction will
proceed spontaneously.
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
3. The algebraic combination of these two factors is given as the equation for the free energy change, ΔG.
Complete this equation:
ΔG = _______ – _________
4. Reactions are spontaneous when the value for ΔG is _______________ (positive/negative).
5. When ΔH is negative and ΔS is positive, the sign of ΔG must be ______________. A negative value for
ΔH and positive value for ΔS apply to a reaction that _________ (is/is not) spontaneous at any
temperature. Therefore, both factors that determine spontaneity are __________________
(favorable/unfavorable). When ΔH < 0, the chemical change moves the system to a ___________
(lower/higher) enthalpy; energy is ___________ (released/absorbed) . When ΔS > 0, the chemical
change moves the system to a __________ (greater/lesser) entropy; the system becomes __________
(more/less) disordered.
6. When ΔH is positive and ΔS is negative, the sign of ΔG must be ______________. A positive value for
ΔH and negative value for ΔS apply to a reaction that _________ (is/is not) spontaneous at any
temperature. Therefore, both factors that determine spontaneity are __________________
(favorable/unfavorable). When ΔH > 0, the chemical change moves the system to a ___________
(lower/higher) enthalpy; energy is ___________ (released/absorbed) . When ΔS < 0, the chemical
change moves the system to a __________ (greater/lesser) entropy; the system becomes __________
(more/less) disordered.
7. When ΔH is positive and ΔS is positive, or when ΔH is negative and ΔS is negative, the sign of ΔG
depends on the magnitude of the contribution to ΔG by the ______________ (entropy/enthalpy) change
which, in turn is determined by the __________________ (pressure/temperature) of the system.
Entropy Change (ΔS) and Gibbs Free Energy Calculations
0
0
Refer to reference tables for the values of S and ΔHf
1. Given the formation of gaseous hydrosulfuric acid from its elements, calculate the ΔS for this reaction.
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
2. Calculate the heat of reaction (ΔH), in kJ for the reaction where Carbon Monoxide gas combines with
solid Iron(III) oxide to form solid Iron and Carbon Dioxide gas. Is the reaction endothermic or
exothermic? How can you tell?
3. Calculate the entropy change (ΔS) for the previous reaction at 298K. Does the entropy change by itself
favor the forward reaction? Why or why not?
4. Is the previous reaction spontaneous at 298K? How can you tell?
5. Is the formation of solid Calcium Oxide from its elements at 25°C spontaneous? ΔS is 0.0259 kcal/K,
and ΔH is +151.8 kcal. How can you tell?
6.
For the reaction 1/2 I2 (g) + 1/2 Cl2 (g) - - - > ICl(g)
reaction spontaneous at 25°C?
7.
For the previous reaction, calculate in degrees Celsius, the temperature at which ΔG = 0.
page 17
ΔH = 17.6 kJ/mole; ΔS = 77.4 J/K; Is this
ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Challenge Problems – Thermodynamics
1. Solar power is proposed as a solution to our energy problems. One place that solar power is used
extensively is Arizona because there the energy of solar radiation is equal to 8.40 x 106 kJ/m2. Much of
our electricity in this country is generated in coal fire power plants. How many tons of coal (carbon)
would need to be combusted to generate this same amount of energy? Note: 454 grams = One pound.
2. At what temperature does ΔG0 become zero for the reaction where solid calcium carbonate decomposes
to solid calcium oxide and gaseous carbon dioxide. ΔH0f CaCO3 = –1207.1 kJ/mol; ΔH0f CaO =
–635.0 kJ/mol; S0 CaCO3 = 93 J/(mol x K); S0 CaO = 38.1 J/(mol x K).
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
3. The combustion of sucrose (C12H22O11) produces 5790 kJ/mole of heat energy at 25°C. 25% of the heat
energy liberated is available to do work (the rest is converted to heat – sweat). If you use 500 Calories
every hour of vigorous exercise, how many hours do you have to exercise to burn off 1.00 pound of
sucrose. Also remember that a large C Calorie equals 1000 small c calories.
4. Is the following reaction spontaneous at 25°C? Oxygen can be made in the lab by reacting solid sodium
peroxide with water vapor to produce the oxygen gas and solid sodium hydroxide. Standard Enthalpy of
formation for sodium peroxide is –510.90 kJ/mol; Standard Enthalpy of formation for sodium
hydroxide is –425.609 kJ/mol; Standard Entropy for sodium peroxide is 95.00 J/(mol x K); Standard
Entropy for Sodium hydroxide is 64.46 J/(mol x K). Why or why not? Is there any temperature at
which this reaction is spontaneous?
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
5. The swimming pool at HS North is 20.0 yds. wide by 50.0 yds. long. At the shallow end it is 1.00 meter
deep. The pool then has a steady and continuous decline until it reaches the far end where it is 3.50 m
deep. The HS North pool is heated by heat transfer released by burning propane which is stored in a
tank outside the school at a temperature of 5.50°C and a pressure of 8935.0 kPa. On the first day after
winter break the temperature of the pool has dropped to 11.0°C. Before gym classes can be held the
pool must be heated to 23.5°C. If propane sells for $4.75/gallon (Summer 2008 prices) using the
additional data below, calculate how much it will cost the WWP district to heat the pool to the desired
temperature.
ΔH0f propane = –74.86 kJ/mol
ΔH0f water = –241.8 kJ/mol
ΔH0f carbon dioxide = –393.5 kJ/mol
1.06 Liters = 1 US Quart
1.00 inch = 2.54 cm
4 quarts = 1 gallon
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
6. Calculate the Enthalpy change associated with the next two reactions (a + b) given the following data:
2 O3 (g) - - - - - - - - > 3 O2 (g)
ΔH = –427 kJ
O2 (g) - - - - - - - - > 2 O (g)
ΔH = +495 kJ
NO (g) + O3 (g) - - - - - - - - > NO2 (g) + O2 (g)
ΔH = –199 kJ
Fe2O3 (g) + 3 CO (g) - - - - - - - - > 2 Fe (s) + 3 CO2 (g)
ΔH =
–23 kJ
3 Fe2O3 (g) + CO (g) - - - - - - - - > 2 Fe3O4 (s) + CO2 (g)
ΔH =
–39 kJ
Fe3O4 (s) + CO (g) - - - - - - - - > 3 FeO (g) + CO2 (g)
ΔH = +18 kJ
a. NO (g) + O (g) - - - - - - - - > NO2 (g)
b. FeO (s) + CO (g) - - - - - - - - > Fe (s) + CO2 (g)
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ChemIAcc-22Thermochemistry WS
Dr. Corell - Chemistry I Accelerated
Thermodynamics Crossword Puzzle
ACROSS
1.
4.
6.
7.
8.
9.
12.
14.
15.
18.
20.
22.
23.
DOWN
th
1/1000 of a gram (abbrev.)
A sudden decrease of temperature
4.184 of 3 DOWN
Symbol for element #2
Smallest particle of an element
Conditions of electrons in metals
Study of energy transformations
Prefix meaning 1000
Gibbs quantity of maximum possible work
(2 words)
What a substance does when it releases heat
of fusion.
Homogeneous mixture
Element with symbol Sn
Measure of Randomness
6.02257 x 1023
Device used to measure heat changes
Unit of energy
Symbol for Change
Word for + sign
Energy associated with melting (3 words)
Symbol for element 9
Measure of average kinetic energy
Ability to do work
Same as 6 ACROSS
Thaw
Eating plan carried out to reduce the number of
6 ACROSS
19. A metal that will cause copper metal to be
precipitated from a solution of copper (II) ions
1.
2.
3.
4.
5.
7.
9.
10.
11.
13.
16.
17.
21. Noble Gas
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
20
18
19
21
22
23
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