# Chapter 17 Thermochemical Equation Which Practice Exam Questions

Chapter 17 – Practice Exam QuestionsMultiple Choice
1. What effect does an increase in temperature have on the entropy of a system?
a. The is no effect on the entropy of the system by increasing its temperature and
the number of microstates remains the same.
b. By increasing the temperature of a system, the particles have more energy and
hence the number of microstates increases and the entropy increases.
c. By increasing the temperature of a system, the particles have more energy and
hence the number of microstates decreases and the entropy increases.
d. By increasing the temperature of a system, there are more available energy
levels and hence the number of microstates increases and the entropy increases.
e. By increasing the temperature of a system, there are less energy levels and
hence the number of microstates decreases and the entropy increases.
2. The thermochemical equation which is associated with G of , the standard enthalpy of
formation, for glucose, C6H12O6(s), is:
a. 6 C (s, graphite) + 6 H2O (l) → C6H12O6 (s)
b. 6 C (s, graphite) + 12 H (g) + 6 O (g) → C6H12O6 (s)
c. 6 C (s, graphite) + 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)
3. All of the following are correct except:
a. Liquids have more entropy than their solids.
b. Solutions have more entropy than the solids and solvents before dissolution.
c. Gases have more entropy than their liquids.
d. Noble gases have more entropy than gaseous molecules that have similar molar
masses.
e. Entropy of a substance increases as its temperature increases.
4. The free energy change for a reaction is 298 kJ. The reaction is therefore:
a. exothermic
b. irreversible
c. spontaneous
d. endothermic
e. non-spontaneous
5. Which property associated with a chemical reaction is dependent on how the reaction is
carried out, not on just its initial and final states?
a. S
b. H
c. E
d. w
e. G
6. Which reaction is accompanied by an increase in entropy?
a. C12H20(l) + 17 O2(g)  12 CO2(g) + 10 H2O(l)
b. NH4Cl(s)  NH3(g) + HCl(g)
c. 2 C2H2(g) + 5 O2(g)  4 CO2(g) + 2 H2O(s)
d. Ba(OH)2(s) + 2 HCl(g)  BaCl22H2O(s)
e. (CH3)2CO(l) + 4 O2(g)  3 CO2(g) + 3 H2O(l)
7. Which set has the species listed in order of increasing standard entropy, S°?
a. CaO(s) < H2O(l) < Mg(s) < CO(g) b. CO(g) < CaO(s) < Mg(s) < H2O(l) c. Mg(s) < CaO(s) < H2O(l) < CO(g) d. Mg(s) < CO(g) < CaO(s) < H2O(l) e. Mg(s) < H2O(l) < CaO(s) < CO(g) 8. In which one of the following cases is the sample with higher absolute entropy listed first? a. 1 mol Br2 (l) > 1 mol Br2 (g)
b. 1 mol Br2 (g) > 2 mol Br (g)
c. 1 mol Br2 (g, 100°C) > 1 mol Br2 (g, 200°C)
H
d. 1 mol
H
H
e. 1 mol
H
H
C
C
C
C
H
H
H
C
H
H3C
H
H
C
C
H
H
H
C
C
C
C
H
C
H
C
H > 1 mol
H
C
C
H3C
H
H
H
CH3
H H
H
H
> 1 mol
H
C
C
H
H
H
H
H
C
C
H H H
H
C
C
H
H
9. The rusting of iron is described by the following chemical reaction:
2 Fe(s) + 3/2 O2(g) → Fe2O3(s)
ΔH° = -824.2 kJ, ΔS° = -582.8 J/K
For what temperature range is the rusting of iron spontaneous?
a. At all temperatures
b. At temperatures above 1.4 K
c. At temperatures below 1.4 K
d. At temperatures above 1400 K
e. At temperatures below 1400 K
10. Ethene (ethylene, C2H4) can dimerize to produce cyclobutane (C4H8) in the presence of
ultraviolet light:
H
H
H
C
C
+
C
H
H
H
h
C
H
H
H
H
C
C
H
C
C
H
H
H
H
H
The average bond energies of C–H, C–C, and C=C are 414 kJ/mol, 347 kJ/mol and 611
kJ/mol, respectively. For this reaction,
a. ΔH° = ΔS° = 0
b. ΔH° < 0 and ΔS° < 0 c. ΔH° < 0 and ΔS° > 0
d. ΔH° > 0 and ΔS° < 0 e. ΔH° > 0 and ΔS° > 0
11. Which of the following is TRUE if ln K is zero?
a. ΔrG° is positive and the reaction is spontaneous in the forward direction.
b. ΔrG° is negative and the reaction is spontaneous in the forward direction.
c. ΔrG° is negative and the reaction is spontaneous in the reverse direction.
d. ΔrG° is positive and the reaction is spontaneous in the reverse direction.
e. ΔrG° is zero and the reaction is at equilibrium.
12. Consider the following reaction:
HNO2 (aq) + H2O (l)
NO 2 (aq) + H3O+ (aq)
ΔG° = +17.9 kJ
When initial concentrations are [HNO2] = 1.0 M and
[ NO 2 ] = [H3O+] = 1.0 x 10-5, calculate ∆G in kJ at
298 K.
13. An important reagent in organic chemistry is thionyl chloride,
SOCl2. One synthetic route for its production is represented by
the following reaction:
SO2(g) + 2 HCl(g) → SOCl2(g) + H2O(g)
What is the standard entropy change for this reaction given the
following gas-phase data:
S°(SO2) = 248.2 J K mol–1
S°(HCl) = 186.9 J K mol–1
–1
S°(SOCl2) = 309.7 J K mol
S°(H2O) = 188.8 J K mol–1
14. Calculate ΔG for the following reaction, at 298 K, under the
conditions shown below.
2Hg(g)  O2 (g)  2HgO(s)
 G° = -180.8 kJ
p(Hg) = 0.025 bar, p(O2) = 0.037 bar
15. Given the following equation,
N2O(g) + NO2(g) → 3 NO(g)
ΔrG° = -23.0 kJ mol-1
calculate ΔrG° for the following reaction:
9NO(g) → 3N2O(g) + 3NO2(g)
16. Estimate ΔrG° for the following reaction at 449.0 K.
CH2O(g) + 2H2(g) → CH4(g) + H2O(g)
ΔrH°= -94.9 kJ mol-1; ΔrS°= -224.2 J K-1 mol-1
17. Use Hess’s law to calculate ΔrG° using the following information:
CO(g) → C(s) + 1/2 O2(g)
ΔrG° = ?
CO2(g) → C(s) + O2(g)
ΔrG° = +394.4 kJ mol-1
CO(g) + 1/2 O2(g) → CO2(g) ΔrG° = -257.2 kJ mol-1
18. Calculate ΔrG° at 298 K using the following information:
2HNO3(aq) + NO(g) → 3NO2(g) + H2O(l)
ΔrG° = ?
ΔfH° (kJ mol-1) -207.0
91.3
33.2
S°(J K-1 mol-1) 146.0
210.8
240.1
-285.8
70.0
19. Use the free energies of formation given below to calculate the equilibrium constant (K)
for the following reaction at 298 K:
ΔfG° (kJ mol-1)
2HNO3(aq) + NO(g) → 3 NO2(g) + H2O(l)
K=?
-110.9
87.6
51.3
-237.1
20. One of the more important industrial chemicals is hydrogen. One process for hydrogen
production is called “steam reforming”, in which hydrocarbons react with water to give
hydrogen and CO. The equation of reaction for reforming methane is written below.
CH4 (g) + H2O (g) → CO (g) + 3 H2 (g)
CH4 (g)
H2O (g)
H2 (g)
Chemical
CO (g)
ΔG˚ (kJ/mol)
-50.5
-228.7
-137.2
0
ΔH˚ (kJ/mol)
-74.6
-241.8
-110.5
0
-1
-1
S˚ (J K mol )
186.3
188.8
197.7
130.7
a. Calculate the free energy change for this reaction under standard conditions.
b. Estimate at what temperature(s), if any, is the process spontaneous?
c. If the partial pressure of the CO and H2O products is 0.25 bar, what is the minimum partial
pressures of the reactants CH4 and H2O required for the reaction to be spontaneous in the
forward reaction (assume CH4 and H2O are added in a 1:1 mixture)?
Multiple Choice
1. B
2. C
3. D
4. E
5. D
6. B
7. C
8. D
9. E
10. B
11. E
12. –39.1 kJ
13. –123.5 J/K
14. -154 kJ
15. 69.0 kJ mol-1
16. 5.8 kJ mol-1
17. 137.2 kJ mol-1
18. 50.8 kJ mol-1
19. 1.15 × 10-9
20. A) 142.5 kJ/mol; B) The reaction is spontaneous at temperatures that are higher
than 968.5 K (695.3˚C); C) pCH 4 = pH 2O = 1.8 x 1011 bar
Chapter 13 – Exam-type Questions
Multiple Choice
1. Which one of the following lowers the activation energy (barrier) of a reaction?
c. removing products
d. raising the temperature
e. none of the above; the activation barrier is not affected by any of these changes
2. From the stoichiometry of the reaction
2 I– (aq) + H2O2 (aq) + 2 H3O+ (aq)  I2 (s) + 4 H2O (l)
What can be said about rate law for this reaction?
a. The rate law is predicted to have a molecularity of five.
b. The rate law is predicted to be second order in [I–].
c. The rate law is predicted to be first-order in [H2O2].
d. Statements a, b, and c are all true.
e. The rate law cannot be determined from the stoichiometry of the reaction.
3. A reaction has the rate law Rate = k[A][B]2. Which one of the following changes will cause
the rate to increase the most?
a. An increase in the size of the reaction container
c. The fast removal of the products from the reaction mixture.
d. The lowering of temperature
e. Tripling [B]
4. For a reaction that follows the general rate law, Rate = k[A][B]2, what will happen to the
rate of reaction if the concentration of B is increased by a factor of 3.00?
a. The rate will decrease by a factor of 1/9.00.
b. The rate will decrease by a factor of 1/3.00.
c. The rate will increase by a factor of 3.00.
d. The rate will increase by a factor of 9.00.
e. The rate will increase by a factor of 6.00.
5. The rate of a reaction usually increases as temperature increases. The most important
reason for this is that increasing the temperature:
a. Makes the reaction more exothermic.
b. Changes the mechanism of the reaction.
c. Increases the activation energy (barrier) for the reaction.
d. Decreases the activation energy (barrier) for the reaction.
e. Increases the fraction of reactants that have kinetic energies that are greater than
the activation energy (barrier).
6. Which of the following statements is FALSE?
a. The average rate of a reaction decreases during a reaction.
b. It is not possible to determine the rate of a reaction from its balanced equation.
c. The rate of zero-order reactions is not dependent on concentration.
d. The half-life of a first-order reaction is dependent on the initial concentration of
reactant.
e. None of the above statements is false.
7. Which of the following reactions would you predict to have the smallest orientation
factor?
a. X2 + Y2 → 2XY
b. NOCl2 + NO → 2 NOCl
c. N2 + O2 → 2NO
d. N + O2 → NO2
e. All of these reactions should have nearly identical orientation factors.
8. A reaction has the rate law Rate = k[A][B]2. Which one of the following changes will cause
the rate to increase the most?
a. An increase in the size of the reaction container
c. The fast removal of the products from the reaction mixture.
d. The lowering of temperature
e. Tripling [B]
9. For the hypothetical chemical reaction A + B → C a plot of [A] versus time was found
to give a straight line with a negative slope. What is the order of the reaction with
respect to A?
a. zero
b. first
c. second
d. third
e. not possible to determine the order from this information
10. Identify a homogeneous catalyst.
a. SO2 over vanadium (V) oxide
b. Pd in H2 gas
c. Pt with methane
d. H2SO4 with concentrated HCl
e. N2 and H2 catalyzed by Fe
11. Write a balanced reaction for which the following rate relationships are true.
1   N 2 O5  1   NO 2    O 2 
Rate = 
=
=
2
t
4 t
t
12. Determine the rate law for the following reaction using the following initial rate data:
2 NO + 2 H2  N2 + 2 H2O
Experimen Initial [NO] Initial [H2] Initial Rate
t
1
0.20
0.12
0.066
2
0.20
0.24
0.13
3
0.20
0.36
0.20
4
0.40
0.12
0.26
5
0.60
0.12
0.59
13. Iodine-131 is used in the form of sodium iodide to treat cancer of the thyroid. Iodine131 has a half-life of 8.05 days. If you begin with 25.0 mg of radioactive 131I, what mass
will remain after 32.2 days (about a month; nuclear decay is a first order process)?
14. The following is the rate law for the reaction: 2 N2O5(g) → 4 NO2 (g) + O2 (g)
O2 
 k N 2 O5 
t
At the time when N2O5 is being consumed at a rate of 3.0×10-4 M/s, what is the rate at
which NO2 is being formed?
15. Consider the aqueous phase reaction between dichromate anions and iron(II) cations:
14 H3O+(aq) + Cr2 O 72 (aq) + 6 Fe2+ (aq) → 2 Cr3+ (aq) + 21 H2O (l)
What is the rate of increase of Cr3+ concentration expressed in terms of changing
H3O+ concentration?
16. The second-order decomposition of HI has a rate constant of 1.80 × 10-3 L mol-1 s-1.
How much HI remains after 27.3 s if the initial concentration of HI is 4.78 mol L-1?
17. A reaction is found to have a rate constant of 3.36 × 104 L mol-1 s-1 at 344 K and a rate
constant of 7.69 L mol-1 s-1 at 219 K. Determine the activation energy for this reaction.
18. In the reaction profile diagram below, label A, B, C, D, and E.
A
B
C
D
E
19. A reaction’s rate is monitored at various temperatures. Plotting a graph of 1/T vs lnk
results in the graph below. Determine the activation energy of the reaction in units of
kJ/mol.
-1
ln(k)
-1.5
-2
-2.5
y = -5640x + 16.6
R² = 0.9999
-3
-3.5
0.0031
0.0032
0.0033
0.0034
0.0035
1/T (/K)
20. The following data were obtained at 25°C for the reaction:
OH   aq   CH 3Br  aq   CH 3OH  aq   Br   aq 
[CH3Br]
[OH–]
Production rate of CH3Br
(mol/L)
(mol/L)
(mol L–1 min–1)
1
0.200
0.200
0.015
2
0.400
0.200
0.030
3
0.400
0.400
0.030
a. Determine the rate law for this reaction, and calculate the value of the rate constant, k
(include the correct units).
b. The proposed reaction mechanism for this reaction is
Experiment
Step 1:
CH3Br  CH3  Br 

CH3  OH  CH3OH
Step 2:
From the answer you obtained in question (a) above, which step is the rate-determining
step? Explain your reasoning using the data above.
c. (1 point) List any catalysts and/or reaction intermediates that appear in the proposed
mechanism:
Multiple Choice
1. B
2. E
3. E
4. D
5. E
6. D
7. B
8. E
9. A
10. D
11. 2 N2O5  4 NO2  O2
12. Rate = k[NO]2[H2]
13. 1.56 mg
14. 6.0×10–4 M/s
 Cr 3
1  H 3O 

15.
t
7
t
16. 3.87 mol L-1
17. 42.0 kJ mol-1
18. A – reactants; B – activation energy; C – activation energy of reverse reaction; D –
enthalpy of reaction; D – products.
19. 46.9 kJ/mol
20. a. Rate = 0.075 min–1 [CH3Br]; b. Step 1 with a valid explanation; c. catalyst(s):
NONE intermediate(s) CH3+

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