UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
1.
The average rate of disappearance of ozone in the reaction
2O3(g)  3O2(g) is found to be 9.0  10–3 atm over a
The following data were obtained at 55°C.
[(CH3)3CBr]0 [OH–]0
Initial Rate
certain interval of time. What is the rate of appearance of
O2 during this interval?
A)
B)
C)
D)
E)
3.
Exp. (mol/L)
(mol/L)
(mol/L s)
1
0.10
0.10
1.0  10–3
2
0.20
0.10
2.0  10–3
3
0.10
0.20
1.0  10–3
4
0.30
0.20
?
What will the initial rate (in mol/L  s) be in Experiment 4?
A) 3.0  10–3
B) 6.0  10–3
C) 9.0  10–3
D) 18  10–3
E) none of these
1.3  10–2 atm/s
9.0  10–3 atm/s
6.0  10–3 atm/s
3.0  10–5 atm/s
2.7  10–5 atm/s
Consider the reaction 2H2 + O2  2H2O
What is the ratio of the initial rate of the appearance of
water to the initial rate of disappearance of oxygen?
A) 1 : 1 B) 2 : 1 C) 1 : 2 D) 2 : 2 E) 3 : 2
4.
9.
Consider the reaction: 4NH3 + 7O2  4NO2 + 6H2O
At a certain instant the initial rate of disappearance of the
oxygen gas is X. What is the value of the appearance of
water at the same instant?
A) 1.2 X
B) 1.1 X
C) 0.86 X
D) 0.58 X
E) cannot be determined from the data
5.
Consider the reaction X  Y + Z
Which of the following is a possible rate law?
A) Rate = k[X]
D) Rate = k[X][Y]
B) Rate = k[Y]
E) Rate = k[Z]
C) Rate = k[Y][Z]
6.
Consider the following rate law: Rate = k[A]n[B]m
How are the exponents n and m determined?
A) By using the balanced chemical equation
B) By using the subscripts for the chemical formulas
C) By using the coefficients of the chemical formulas
D) By educated guess
E) By experiment
7.
10.
Initial Rate
1  1018 1  1018
2.0  1016
2  1018 1  1018
8.0  1016
3  1018 1  1018
18.0  1016
1  1018 2  1018
4.0  1016
1  1018 3  1018
6.0  1016
Which of the following is the correct rate law?
A) Rate = k[NO][O2]
D) Rate = k[NO]2
B) Rate = k[NO][O2]2
C) Rate = k[NO]2[O2]
8.
Tabulated below are initial rate data for the reaction
2Fe(CN)63– + 2I–  2Fe(CN)64– + I2
Run [Fe(CN)63–]0 [I–]0 [Fe(CN)64-]0 [I2]0Rate
(M/s)
1
0.01 0.01
0.01 0.01
1  10–5
2
0.01 0.02
0.01 0.01
2  10–5
3
0.02 0.02 0.01 0.01
8  10–5
4
0.02 0.02 0.02 0.01
8  10–5
5
0.02 0.02 0.02 0.02 8  10–5
The experimental rate law is:
A) rate = k[Fe(CN)63–]2[I–]2[Fe(CN)64–]2[I2]
B) rate = k[Fe(CN)63–]2[I–][Fe(CN)64–][I2]
The following data were obtained for the reaction of NO
with O2. Concentrations are in molecules/cm3 and rates are
in molecules/cm3  s.
[NO]0
[O2]0
For a reaction in which A and B react to form C, the
following initial rate data were obtained:
[A]
[B]
Initial Rate [C]
(mol/L)
(mol/L)
(mol/L  s)
0.2
0.2
0.50
0.4
0.2
2.00
0.8
0.2
8.00
0.2
0.4
1.00
0.2
0.8
2.00
What is the rate law for the reaction?
A) Rate = k[A][B]
B) Rate = k[A]2[B]
C) Rate = k[A][B]2
D) Rate = k[A]2[B]2
E) Rate = k[A]3
C)
rate = k[Fe(CN)63–)]2[I–]
rate = k[Fe(CN)63–][I–]2
D)
E)
11.
E) Rate = k[NO]2[O2]2
rate = k[Fe(CN)63–][I–] [Fe(CN)64–]
Tabulated below are initial rate data for the reaction
2Fe(CN)63– + 2I–  2Fe(CN)64– + I2
Run
1
2
3
4
5
The reaction of (CH3)3CBr with hydroxide ion proceeds
with the formation of (CH3)3COH.
(CH3)3CBr(aq) + OH–(aq)  (CH3)3COH(aq) + Br–(aq)
Page 1
[Fe(CN)63–]0 [I–]0 [Fe(CN)64–]0
0.01
0.01
0.02
0.02
0.02
0.01
0.02
0.02
0.02
0.02
0.01
0.01
0.01
0.02
0.02
0.01
0.01
0.01
0.01
0.02
[I2]0 Rate
(M/s)
1  10–5
2  10–5
8  10–5
8  10–5
8  10–5
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
The value of k is:
A) 107 M–5 s–1
B) 103 M–3 s–1
C) 10 M–2 s–1
D) 50 M–2 s–1
E) none of these
III.
IV.
18.
C) rate = k[I–][H+]
D) rate = k[H2O2][H+]
E) rate = k[H2O2][I–]
A general reaction written as 2A + 2B  C + 2D is studied and yields
the following data:
[A]0
[B]0
Initial [C]/t
12.
What is the order of the reaction with respect to B?
A) 1 B) 4 C) 3 D) 2 E) 0
What is the order of the reaction with respect to A?
A) 1 B) 4 C) 3 D) 2 E) 0
14.
What is the overall order of the reaction?
A) 1 B) 4 C) 3 D) 2 E) 0
15.
What are the proper units for the rate constant for the
reaction?
A) s–1
B) mol L–1 s–1
C) L mol–1 s–1
D) L3 mol–3 s–1
E) L2 mol–2 s–1
17.
19.
The average value for the rate constant k (without units) is
A) 2710
B) 2.74 × 104
C) 137
D) 108
E) none of these
20.
Two mechanisms are proposed:
I. H2O2 + I–  H2O + OI–
0.000040 mol/L  s
0.000160 mol/L  s
0.000040 mol/L  s
13.
16.
OI– + H+  HOI
HOI + I– + H+  I2 + H2O
II.
I2 + I–  I3–
H2O2 + I– + H+  H2O + HOI
HOI + I– + H+  I2 + H2O
I2 + I–  I3–
Which of the following describes a potentially correct
mechanism?
A) Mechanism I with the first step the rate determining
step.
B) Mechanism I with the second step the rate determining
step.
C) Mechanism II with the first step rate determining.
D) Mechanism II with the second step rate determining.
E) None of these could be correct.
What is the numerical value of the rate constant?
A) 0.000040
B) 0.000160
C) 0.0040
D) 0.0160
E) 4.0  10–7
21.
A first-order reaction is 35% complete at the end of 55
minutes. What is the value of the rate constant?
A) 1.9 × 10–3 min-1
D) 7.8 × 10–3 min–1
B) 36 min–1
E) none of these
C) 89 min–1
Use the following to answer questions 22-23:
For the first of the reactions in the table of data, how
many seconds would it take for [A] to decrease to 0.050 M?
A) 1200
B) 1700
C) 170
D) 2500
E) 250
The following initial rate data were found for the reaction
2MnO4– + 5H2C2O4 + 6H+  2Mn2+ + 10CO2 + 8H2O
[MnO4–]0
1  10–3
2  10–3
2  10–3
2  10–3
Use the following to answer questions 18-20:
22.
Consider the following data concerning the equation:
H2O2 + 3I– + 2H+  I3– + 2H2O
I.
II.
The rate law for this reaction is
A) rate = k[H2O2][I–][H+]
B) rate = k[H2O2]2[I–]2[H+]2
Use the following to answer questions 12-17:
0.100 M 0.100 M
0.200 M 0.100 M
0.100 M 0.200 M
0.200 M 1.00 × 10–3 M 1.00 × 10–2 M 0.542 M/sec
0.400 M 1.00 × 10–3 M 2.00 × 10–2 M 1.084 M/sec
[H2O2]
[I– ]
[H+]
rate
–4
0.100 M 5.00 × 10 M 1.00 × 10–2 M 0.137 M/sec
0.100 M 1.00 × 10–3 M 1.00 × 10–2 M 0.268 M/sec
[H2C2O4]0
1  10–3
1  10–3
2  10–3
2  10–3
[H+]0
1.0
1.0
1.0
2.0
2  10–4
8  10–4
1.6  10–3
1.6  10–3
Which of the following is the correct rate law?
A) Rate = k[MnO4–]2[H2C2O4]5[H+]6
B) Rate = k[MnO4–]2[H2C2O4][H+]
C) Rate = k[MnO4–][H2C2O4][H+]
D) Rate = k[MnO4–]2[H2C2O4]
E) Rate = k[MnO4–]2[H2C2O4]2
Page 2
Initial Rate (M/s)
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
23.
1.0  10–4
What is the value of the rate constant?
A) 2  105 M  s–1
B) 2  105 M-2  s–1
C) 200 M–1  s–1
D) 200 M–2  s–1
E) 2  10–4 M  s–1
28.
24.
(mol/L)
6.4  10–3
12.8  10–3
6.4  10–3
The rate law is
A) Rate = k[H2SeO3][H+][I–]
D) Rate = k[H2SeO3]2[H+][I–]
E) Rate = k[H2SeO3][H+]2[I–]3
Initial [H2] Disappearance of NO
(mol/L)
2.2  10–3
2.2  10–3
4.5  10–3
3.36  10–7
C) Rate = k[H2SeO3][H+][I–]2
The following questions refer to the reaction between nitric oxide
and hydrogen
2NO + H2  N2O + H2O
Initial [NO]
4.0  10–2
B) Rate = k[H2SeO3][H+]2[I–]
Use the following to answer questions 24-27:
Exp
1
2
3
1.0  10–2
(mol/L  s)
2.6  10–5
1.0  10–4
5.1  10–5
29.
The numerical value of the rate constant is
A) 5.2  105
B) 2.1  102
C) 4.2
Initial Rate of
D) 1.9  10–6
E) none of these
Use the following to answer questions 30-33:
The following questions refer to the reaction shown below:
A + 2B  2AB
Initial [A] Initial [B] Disappearance of A
Exp
(mol/L)
(mol/L)
(mol/L  s)
1
0.16
0.15
0.08
2
0.16
0.30
0.30
3
0.08
0.30
0.08
What is the rate law for this reaction?
A) Rate = k[NO]
B) Rate = k[NO]2
C) Rate = k[NO]2[H2]
D) Rate = k[NO][H2]
E) Rate = k[N2O][H2O]
30.
What is the rate law for this reaction?
A) Rate = k[A][B]
B) Rate = k[A]2[B]
C) Rate = k[A][B]2
D) Rate = k[A]2[B]2
E) Rate = k[B]
31.
What is the magnitude of the rate constant for the
reaction?
A) 140 B) 79
C) 119 D) 164
E) 21
What are the units for the rate constant for this reaction?
A) L/mol  s
B) L2/mol2  s
C) mol/L  s
D) L3/mol3  s
E) mol3/L
25.
What is the magnitude of the rate constant for this
reaction?
A) 1150 B) 98
C) 542 D) 112 E) 289
26.
What are the units for the rate constant for this reaction?
A) L/mol  s
B) L2/mol2  s
C) mol/L  s
D) s–2
E) L–2
32.
What is the order of this reaction?
A) 3
B) 2
C) 1
D) 0
E) cannot be determined from the data
33.
What is the order of this reaction?
A) 4
B) 3
C) 2
D) 1
E) 0
34.
Initial rate data have been determined at a certain
temperature for the gaseous reaction
2NO + 2H2  N2 + 2H2O.
27.
Use the following to answer questions 28-29:
The reaction
H2SeO3(aq) 6I–(aq) + 4H+(aq)  2I3–(aq) + 3H2O(l) + Se(s)
was studied at 0°C by the method of initial rates:
[H2SeO3]0
[H+]0
[I–]0
1.0  10–4
2.0  10–4
3.0  10–4
1.0  10–4
1.0  10–4
1.0  10–4
2.0  10–2
2.0  10–2
2.0  10–2
4.0  10–2
1.0  10–2
2.0  10–2
2.0
2.0
2.0
2.0
2.0
4.0
 10–2
 10–2
 10–2
 10–2
 10–2
 10–2
Rate (mol/L s)
1.66  10–7
3.33  10–7
4.99  10–7
6.66  10–7
0.42  10–7
13.4  10–7
[NO]0
0.10
0.10
0.20
Page 3
[H2]0
0.20
0.30
0.20
Initial Rate (M/s)
0.0150
0.0225
0.0600
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
The numerical value of the rate constant is:
A) 7.5
B) 3.0  10–3
C) 380
D) 0.75
E) 3.0  10–4
35.
The following data were obtained at 25°C:
[A]0
[B]0
[C]0
0.1
0.2
0.3
0.3
0.4
0.2
0.6
0.4
0.2
0.3
0.4
0.1
0.6
0.2
0.2
What
A)
B)
C)
D)
E)
k.
D) The rate of the reaction increases with time.
E) A plot of 1/[HO2] versus time gives a straight line.
Use the following to answer questions 42-45:
The following questions refer to the gas-phase decomposition of
ethylene chloride.
C2H5Cl  products
Rate
0.063
0.084
0.168
0.021
0.168
Experiment shows that the decomposition is first order.
The following data show kinetics information for this reaction:
Time (s)
ln [C2H5Cl] (M)
1.0
2.0
is the correct rate law?
Rate = k[A][B][C]
Rate = k[A][B][C]2
Rate = k[A][C]
Rate = k[A]3[B]2[C]
Rate = k[A][C]2
Use the following to answer questions 36-38:
buffered solution. The following data were obtained:
Relative
Initial Rate
[Ce4+]0
[Ce3+]0
[Cr3+]0
1
2.0  10–3 1.0  10–2 3.0  10–2
2
4.0  10–3 2.0  10–2 3.0  10–2
4
4.0  10–3 1.0  10–2 3.0  10–2
16
8.0  10–3 2.0  10–2 6.0  10–2
37.
Determine the order in the rate law of the species Cr3+.
A) 1
B) 2
C) 3
D) -1
E) -2
39.
The rate expression for a particular reaction is rate =
k[A][B]2. If the initial concentration of B is increased from
0.1 M to 0.3 M, the initial rate will increase by which of the
following factors?
A) 2
B) 6
C) 12
D) 3
E) 9
41.
What is the rate constant for this decomposition?
A) 0.29/s B) 0.35/s C) 0.11/s D) 0.02/s E) 0.22/s
43.
What was the initial concentration of the ethylene
chloride?
A) 0.29/s B) 0.35/s C) 0.11/s D) 0.02/s E) 0.22/s
45. What is the time to half-life?
A) 0.7 s
B) 1.3 s
C) 8.9 s
D) 6.3 s
E) 2.2 s
Use the following to answer questions 46-47:
For a reaction: aA  Products, [A]o = 4.0 M, and the first two halflives are 34 and 68 minutes, respectively.
Determine the order in the rate law of the species Ce4+.
A) 1
B) 2
C) 3
D) -1
E) -2
Determine the order in the rate law of the species Ce3+.
A) 1
B) 2
C) 3
D) -1
E) -2
38.
42.
44. What would the concentration be after 5.0 seconds?
A) 0.13 M B) 0.08 M C) 0.02 M
D) 0.19 M E) 0.12 M
The oxidation of Cr3+ to CrO42– can be accomplished using Ce4+ in a
36.
–1.625
–1.735
46.
Calculate k (without units)
A) 2.0 × 10–2
B) 7.4 × 10–3
C) 5.9 × 10–2
D) 1.0 × 10–2
E) none of these
47.
Calculate [A] at t= 192 minutes.
A) 0.086 M
B) 0.00 M
C) 0.60 M
D) 1.4 M
E) none of these
48.
For which order reaction is the half life of the reaction
proportional to 1/k (k is the rate constant)?
A) zero order
B) first order
C) second order
D) all of the above
E) none of the above
The following data were collected for the decay of HO2
radicals:
Time [HO2]
Time
[HO2]
0s
1.0  1011 molec/cm3 14 s
1.25  1010 molec/cm3
2s
5.0  1010 molec/cm3 30 s
6.225  109 molec/cm3
6s
2.5  1010 molec/cm3
Which of the following statements is true?
A) The decay of HO2 occurs by a first-order process.
B) The half-life of the reaction is 2 ms.
C) A plot of ln [HO2] versus time is linear with a slope of –
Page 4
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
52.
The order of this reaction in N2O5 is
A)
B)
C)
D)
E)
53.
54.
The rate law for the reaction is Rate = k[A]x[B]y. What are
the values of x and y?
A) x = 0 y = 1
B) x = 1 y = 0
C) x = 1 y = 1
D) x = 2 y = 1
E) x = 1 y = 2
58.
What form will the pseudo-rate law have?
A) Rate = k´[A]x
B) Rate = k´[B]y
C) Rate = k´[A]x[B]y
D) Rate = kk´[A]x
E) Rate = kk´[B]y
59.
Determine the magnitude of the pseudo-rate constant (k´)
if the magnitude of X in the rate data is 0.00905.
A) 4.3  10–3
B) 1.2  10–2
C) 0.86
D) .31
E) 1.81  10–3
0
1
2
3
none of these
The concentration of O2 at t = 10. minutes is
A)
B)
C)
D)
E)
57.
2.0  10–4 mol/L
0.32  10–2 mol/L
0.16  10–2 mol/L
0.64  10–2 mol/L
none of these
The initial rate of production of NO2 for this reaction is
approximately
A) 6.4  10–4 mol/L  min
B) 3.2  10–4 mol/L  min
C) 1.24  10–2 mol/L  min
D) 1.6  10–4 mol/L  min
E) none of these
55.
56.
Use the following to answer questions 60-63:
The half-life of this reaction is approximately
A) 15 minutes
B) 18 minutes
C) 23 minutes
D) 36 minutes
E) 45 minutes
The reaction A  B + C is known to be zero order in A with a rate
constant of
5.0  10–2 mol/L  s at 25°C. An experiment was run at 25°C where
[A]0 = 1.0  10–3 M.
60.
The concentration N2O5 at 100 minutes will be
approximately
A) 0.03  10–2 mol/L
B) 0.06  10–2 mol/L
C) 0.10  10–2 mol/L
D) 0.01  10–2 mol/L
E) none of these
C)
Use the following to answer questions 57-59:
D)
E) [A]0 – [A] = kt
The following questions refer to the hypothetical reaction A + B 
products. The kinetics data given can be analyzed to answer the
questions.
[A]0
[B]0
Rate of decrease
(mol/L)
5.0
10.0
5.0
Time (s)
10.0
20.0
30.0
The integrated rate law is
A) [A] = kt
B) [A] – [A]0 = kt
(mol/L)
5.0
5.0
10.0
61.
After 5.0 minutes, the rate is
A) 5.0  10–2 mol/L  s
B) 2.5  10–2 mol/L  s
C) 1.2  10–2 mol/L  s
D) 1.0  10–3 mol/L  s
E) none of these
62.
The half-life for the reaction is
A) 1.0  10–2 s
B) 1.0  102 s
C) 5.0  10–2 s
D) 5.0  10–4 s
E) none of these
of [A] (M/s)
X
2X
2X
[B] (mol/L)
100
100
100
Page 5
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
63.
What is the concentration of B after 5  10–3 sec?
A) 5.0  10–5 M
B) 5.0  10–4 M
C) 7.5  10–4 M
D) 2.5  10–4 M
E) none of these
Use the following to answer questions 69-72:
Use the following to answer questions 64-65:
Consider the reaction
3A + B + C  D + E
where the rate law is defined as
The reaction
2NOBr  2NO + Br2
exhibits the rate law
An experiment is carried out where [B]0 = [C]0 = 1.00 M and [A]0 =
1.00  10–4 M.
where k = 1.0  10–5 M–1  s–1 at 25°C. This reaction is run where the
initial concentration of NOBr ([NOBr]0) is 1.00  10–1 M.
64.
65.
What is one half-life for this experiment?
A) 5.0  10–1 s
B) 6.9  104 s
C) 1.0  10–5 s
D) 1.0  106 s
E) none of these
The [NO] after 1.00 hour has passed is
A) 3.5  10–4 M
B) 9.9  10–3 M
C) 9.7  10–3 M
D) 1.0  10–3 M
E) none of these
69.
After 3.00 minutes, [A] = 3.26  10–5 M. The value of k is
A) 6.23  10–3 L3/mol3  s
B) 3.26  10–5 L3/mol3  s
C) 1.15  102 L3/mol3  s
D) 1.00  108 L3/mol3  s
E) none of these
70.
The half-life for this experiment is
A) 1.11  102 s
B) 87.0 s
C) 6.03  10–3 s
D) 117 s
E) none of these
71.
The concentration of C after 10.0 minutes is
A) 1.00 M
B) 1.10  10–5 M
C) 0.330 M
D) 0.100 M
E) none of these
72.
The concentration of A after 10.0 minutes is
A) 1.06  10–9 M
B) 2.38  10–6 M
C) 9.80  10–6 M
D) 1.27  10–5 M
E) none of these
Use the following to answer questions 66-67:
For the reaction A  Products, successive half-lives are observed to
be 10.0 min and 40.0 min. At the beginning of the reaction, [A] was
0.10 M.
66.
The reaction follows the integrated rate law
A) [A] = -kt + [A]0
B) ln [A] = –kt + ln [A]0
C)
Use the following to answer questions 73-75:
The reaction
AB+C
is second order in A. When [A]0 = 0.100 M, the reaction is 20.0%
D)
E) none of these
67.
68.
The numerical value of the rate constant is
A) 0.069
B) 1.0
C) 10.0
D) 5.0  10–3
E) none of these
complete in 40.0 minutes.
73.
The reaction
Page 6
Calculate the value of the rate constant (in L/min  mol).
A) 6.25  10–2
B) 5.58  10–3
C) 1.60  101
D) 1.00
E) none of these
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
74.
Calculate the half-life for the reaction.
A) 1.60  102 min
B) 1.11  101 min
C) 1.00  101 min
D) 1.00  102 min
E) none of these
75.
A first-order reaction is 40.% complete at the end of 50.
minutes. What is the value of the rate constant (in min–1)?
A) 1.8  10–2
B) 1.0  10–2
C) 1.2  10–2
D) 8.0  10–3
E) none of these
76.
variety of products. The reaction is first order in
Ru(NH3)63+ and has a half-life of 14 hours at 25°C. Under
these conditions, how long will it take for the [Ru(NH3)63+]
to decrease to 12.5% of its initial value?
A) 28 hours
B) 35 hours
C) 2.7 hours
D) 14 hours
E) 42 hours
80.
is made pseudo-first order in oxygen atoms by using a large
excess of ClO radicals. The rate constant for the reaction
is 3.5  10–11 cm3/molecule  s. If the initial concentration
of ClO is 1.0  1011 molecules/cm3, how long will it take for
the oxygen atoms to decrease to 10.% of their initial
concentration?
A) 2.4 s
B) 0.017 s
C) 3.2  10–3 s
D) 0.66 s
E) 23 s
The OH radical disproportionates according to the
elementary chemical reaction
OH + OH  H2O + O. This reaction is second order in OH.
The rate constant for the reaction is 2.0  10–12
cm3/molecule  s at room temperature. If the initial OH
concentration is 1.0  1013 molecules/cm3, what is the first
half-life for the reaction?
A) 20. s
B) 2.0  10–3 s
C) 0.050 s
D) 0.035 s
E) 12 s
77.
81.
The following data were obtained for the reaction 2A + B 
C where rate = d{A]/dt
[A](M)
[B](M)
Initial Rate(M/s)
0.100
0.0500
2.13  10–4
0.200
0.0500
4.26  10–4
0.300
0.100
2.56  10–3
Determine the value of the rate constant.
A) 0.426
B) 0.852
C) 0.0426
D) 0.284
E) none of these
82.
Determine the molecularity of the following elementary
reaction: O3  O2 + O.
At a particular temperature, N2O5 decomposes according
to a first-order rate law with a half-life of 3.0 s. If the
initial concentration of N2O5 is 1.0  1016 molecules/cm3,
what will be the concentration in molecules/cm3 after 10.0
s?
A) 9.9  1014
B) 1.8  1012
C) 7.3  109
D) 6.3  103
E) 9.4  102
78.
A)
B)
C)
D)
E)
The reaction
3NO  N2O + NO2
is found to obey the rate law Rate = k[NO]2. If the first
half-life of the reaction is found to be 2.0 s, what is the
length of the fourth half-life?
A) 2.0 s
B) 4.0 s
C) 8.0 s
D) 12.0 s
E) 16.0 s
79.
The elementary chemical reaction
O + ClO  Cl + O2
83.
unimolecular
bimolecular
termolecular
quadmolecular
the molecularity cannot be determined
The decomposition of ozone may occur through the twostep mechanism shown:
step 1
O3  O2 + O
step 2
O3 + O  2O2
The oxygen atom is considered to be a(n)
A) reactant
B) product
C) catalyst
D) reaction intermediate
E) activated complex
Use the following to answer questions 84-87:
In 6 M HCl, the complex ion Ru(NH3)63+ decomposes to a
Page 7
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
The following questions refer to the reaction 2A2 + B2  2C. The
2H2O occurs by the following series of steps:
following mechanism has been proposed:
step 1 (very slow)
A2 + B2  R + C
step 2 (slow)
84.
85.
86.
A2 + R  C
What is the molecularity of step 2?
A) unimolecular
B) bimolecular
C) termolecular
D) quadmolecular
E) the molecularity cannot be determined
Step 2.
constant k2)
Step 3.
Step 4.
Step 5.
Which step is "rate determining"?
A) both steps
B) step 1
C) step 2
D) a step that is intermediate to step 1 and step 2
E) none of these
According to collision theory, the activated complex that
forms in step 1 should have the following structure. (The
dotted lines represent partial bonds)
H3O2+ + I–  H2O + HOI (slow, rate
HOI + I–  OH– + I2 (fast, rate constant k3)
OH– + H+  H2O
(fast, rate constant k4)
I2 + I–  I3–
(fast, rate constant k5)
95.
Which of the steps would be called the rate-determining
step?
A) 1
B) 2
C) 3
D) 4
E) 5
96.
The rate constant k for the reaction would be given by
A) k = k2
B) k = k2k3
C) k = k2K
A)
D) k = k5
D)
E) k = Kk2k3k4k5
97.
B)
The rate law for the reaction would be:
A) [I3]/t = k[H2O2]
B) [I3]/t = k[H2O2][H+][I–]
E)
C) [I3]/t = k[H2O2][H+]
D) [I3]/t = k[H2O2][I–]
E) [I3]/t = k[H2O2][H+]2[I–]–3
98.
C)
87.
According to the proposed mechanism, what should the
overall rate law be?
A) rate = k[A2]2
The reaction
2A + B  C
has the following proposed mechanism:
B) rate = k[A2]
C) rate = k[A2][B2]
D) rate = k[A2][R]
E) rate = k[R]2
94.
If step 2 is the rate-determining step, then the rate of
formation of C should equal:
A) k[A]
B) k[A]2[B]
C) k[A]2[B]2
D) k[A][B]
E) k[A][B]2
If the reaction 2HI  H2 + I2 is second order, which of the
following will yield a linear plot?
A) log [HI] vs time
B) 1/[HI] vs time
C) [HI] vs time
D) ln [HI] vs time
Use the following to answer questions 95-97:
99.
Under certain conditions the reaction H2O2 + 3I– + 2H+  I3– +
Page 8
The reaction 2NO + O2  2NO2 obeys the rate law
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
B)
C)
D)
E)
Which of the following mechanisms is consistent with the
experimental rate law?
A) NO + NO  N2O2
(slow)
N2O2 + O2  2NO2
(fast)
102.
If the reaction were reversible, would the forward or the
reverse reaction have a higher activation energy?
A) The diagram shows no indication of any activation
energy.
B) The forward and reverse activation energies are equal.
C) The forward activation energy
D) The reverse activation energy
E) none of these
103.
What would happen if the kinetic energy of the reactants
was not enough to provide the needed activation energy?
A) The products would be produced at a lower energy
state.
B) The rate of the reaction would tend to increase.
C) The activated complex would convert into products.
D) The reactants would re-form.
E) The products would form at an unstable energy state.
B)
C)
D) O2 + O2  O2 + O2*
O2 + NO  NO2 + O
O + NO  NO2
point X
point Y
point Z
none of these
(slow)
(fast)
(fast)
E) none of these
Use the following to answer questions 100-102:
Use the following to answer questions 104-106:
The questions below refer to the following diagram:
The questions below refer to the following information:
The rate constant k for the reaction shown below is 2.6  10–8 L/mol
 s when the reaction proceeds at 300.0 K. The activation energy is
98000 J/mol. (The universal gas constant (R) is 8.314 J/mol  K)
2NOCl  2NO + Cl2
100.
101.
Why is this reaction considered to be exothermic?
A) Because energy difference B is greater than energy
difference C
B) Because energy difference B is greater than energy
difference A
C) Because energy difference A is greater than energy
difference C
D) Because energy difference B is greater than energy
difference C plus energy difference A
E) Because energy difference A and energy difference C
are about equal
104.
Determine the magnitude of the frequency factor for the
reaction.
A) 1.2  10–8
B) 4.6  10–9
C) 3.2  10–9
D) 2.7  10–8
E) 9.1  10–9
105.
If the temperature changed to 310 K the rate constant k
would change. The ratio of k at 310 K to k at 300.0 K is
closest to what whole number?
A) 1
B) 2
C) 3
D) 4
E) 5
106.
Using the following information determine the activation
energy for the reaction shown here:
2NO  N2 + O2
Temperature (K)
1400
1500
A) 3.2  104 J/mol
B) 9.5  106 J/mol
C) 2.8  104 J/mol
At what point on the graph is the activated complex
present?
A) point W
Page 9
Rate Constant
(L/mol  s)
0.143
0.659
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
D) 6.8  105 J/mol
E) 2.7  105 J/mol
107.
A)
B)
C)
D)
E)
The reaction 2H2O2  2H2O + O2 has the following
mechanism?
H2O2 + I–  H2O + IO–
H2O + IO–  H2O + O2 + I–
The catalyst in the reaction is:
A) H2O
112.
The rate constant k is dependent on
I. the concentration of the reactant.
II. the nature of the reactants.
III. the temperature.
IV. the order of the reaction.
A) none of these
B) one of these
C) two of these
D) three of these
E) all of these
113.
The rate law for a reaction is found to be Rate = k[A]2[B].
Which of the following mechanisms gives this rate law?
B) I–
C) H2O2
D) IO–
108.
When ethyl chloride, CH3CH2Cl, is dissolved in 1.0 M NaOH,
it is converted into ethanol, CH3CH2OH, by the reaction
CH3CH2Cl + OH–  CH3CH2OH + Cl–
At 25°C the reaction is first order in CH3CH2Cl, and the
higher, lower
higher, higher
lower, higher
lower, steady
higher, steady
rate constant is 1.0  10–3 s–1. If the activation
parameters are A = 3.4  1014 s–1 and Ea = 100.0 kJ/mol,
what will the rate constant be at 40°C?
A) 6.9  10–3 s–1
B) 1.7  102 s–1
C) 5.0  10–3 s–1
D) 2.0  10–3 s–1
E) 5.0  1014 s–1
109.
110.
111.
Which of the following statements best describes the
condition(s) needed for a successful formation for a
product according to the collision model?
A) The collision must involve a sufficient amount of energy,
provided from the motion of the particles, to overcome
the activation energy.
B) The relative orientation of the particles has little or no
effect on the formation of the product.
C) The relative orientation of the particles has an effect
only if the kinetic energy of the particles is below some
minimum value.
D) The relative orientation of the particles must allow for
formation of the new bonds in the product.
E) The energy of the incoming particles must be above a
certain minimum value and the relative orientation of
the particles must allow for formation of new bonds in
the product.
A)
B)
C)
D)
E)
I
II
III
two of these
none of these
Use the following to answer questions 114-116:
A reaction represented by the equation
3O2 (g)  2O3 (g)
was studied at a specific temperature and the following data were
collected:
time (seconds)
total pressure (atm)
0
1.000
46.89
0.9500
98.82
0.9033
137.9
0.8733
200.0
0.8333
286.9
0.7900
337.9
0.7700
511.3
0.7233
Which of the following statements is typically true for a
catalyst?
A) The concentration of the catalyst will go down as a
reaction proceeds.
B) The catalyst provides a new pathway in the reaction
mechanism.
C) The catalyst speeds up the reaction.
D) Two of these.
E) None of these.
The catalyzed pathway in a reaction mechanism has a
__________ activation energy and thus causes a
__________ reaction rate.
Page 10
114.
Which is the rate law for this reaction?:
115.
Which is the value of the rate constant?:
116.
How many seconds would it take for the total pressure to
be 0.7133 atm?
117.
The rate constant for a reaction at 40.0°C is exactly three
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
times that at 20.0°C. Calculate the Arrhenius energy of
activation for the reaction.
A) 3.00 kJ/mol
B) 366 kJ/mol
C) 41.9 kJ/mol
D) 3.20 kJ/mol
E) none of these
118.
heat is added to the reaction.
E) None of these statements is true.
126.
The value of the equilibrium constant, K, is dependent on
I. The temperature of the system.
II. The nature of the reactants and products.
III. The concentration of the reactants.
IV. The concentration of the products.
A) I, II
B) II, III
C) III, IV
D) It is dependent on three of the above choices.
E) It is not dependent on any of the above choices.
127.
Apply the law of mass action to determine the equilibrium
expression for
Determine (a) the rate equation and (b) the rate constant
for the hypothetical reaction
A + B  C given the following initial concentrations and
initial rate data.
[A]0
[B]0
Initial Rate
Run #
(1)
(2)
(3)
(mol/L)
0.100
0.100
0.200
(mol/L)
0.100
0.200
0.200
(mol/L  s)
0.18
0.36
1.44
A) 2[NO2][Cl2]/2[NO2Cl]
B) 2[NO2Cl]/2[NO2][Cl2]
Use the following to answer questions 119-121:
C) [NO2Cl]2/[NO2]2[Cl2]
D) [NO2]2[Cl2]/[NO2Cl]2
Use the potential energy diagram shown to answer the following:
E) [NO2Cl]2[NO2]2[Cl2]
Use the following to answer questions 128-130:
119.
Which letter shows the activation energy?
120.
Which letter shows the change in energy for the overall
reaction?
121.
Which letter shows the activation energy using a catalyst?
122.
Which of the following statements concerning equilibrium is
not true?
A) A system that is disturbed from an equilibrium
condition responds in a manner to restore equilibrium.
B) Equilibrium in molecular systems is dynamic, with two
opposing processes balancing one another.
C) The value of the equilibrium constant for a given
reaction mixture is the same regardless of the
direction from which equilibrium is attained.
D) A system moves spontaneously toward a state of
equilibrium.
E) The equilibrium constant is independent of
temperature.
123.
Consider the chemical system
K = 4.6  109 L/mol.
128.
How do the equilibrium concentrations of the reactants
compare to the equilibrium concentration of the product?
A) They are much smaller.
B) They are much bigger.
C) They are about the same.
D) They have to be exactly equal.
E) You can't tell from the information given.
129.
If the concentration of the product were to double, what
would happen to the equilibrium constant?
A) It would double its value.
B) It would become half its current value.
C) It would quadruple its value.
D) It would not change its value.
E) It would depend on the initial conditions of the product.
130.
Determine the equilibrium constant for the system
at 25°C. The concentrations are
shown here: [N2O4] = 4.27  10–2 M, [NO2] = 1.41  10–2 M
Which of the following statements is true?
A) When two opposing processes are proceeding at
identical rates, the system is at equilibrium.
B) Catalysts are an effective means of changing the
position of an equilibrium.
C) The concentration of the products equals that of
reactants and is constant at equilibrium.
D) An endothermic reaction shifts toward reactants when
A)
B)
C)
D)
E)
Page 11
0.33
3.0
0.66
0.05
0.0047
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
131.
At 500.0 K, one mole of gaseous ONCl is placed in a oneliter container. At equilibrium it is 9.0% dissociated
according to the equation shown here:
the right?
I. increasing the temperature
II. decreasing the temperature
III. increasing the volume
IV. decreasing the volume
V. removing some NH3
Determine the equilibrium constant.
A) 4.4  10–4
B) 2.2  102
C) 1.1  102
D) 2.2  10–4
E) 9.1  10–1
132.
133.
VIII. adding some N2
A) I, IV, VI, VII
B) II, III, V, VIII
C) I, VI, VIII
D) I, III, V, VII
E) II, IV, V, VIII
Consider the reaction
whose K =
54.8 at 425°C. If an equimolar mixture of reactants gives
the concentration of the product to be 0.50 M at
equilibrium, determine the concentration of the hydrogen.
A) 4.6  10–3 M
B) 6.8  10–2 M
C) 1.2  10–3 M
D) 9.6  10–2 M
E) 1.6  10–4 M
137.
Consider the gaseous reaction
expression for Kp in terms of K?
A) K(RT)
B) K/(RT)
C) K(RT)2
D) K/(RT)2
E) 1/K(RT)
134.
VI. adding some NH3
VII. removing some N2
A)
B)
C)
D)
E)
What is the
138.
Find the value of the equilibrium constant (K) (at 500 K) for
139.
at 600 K,
4.4  1043
9.8  1024
1.2  10–4
5.4  10–13
2.6  10–31
Given the equation:
The
equilibrium constant is 0.0150 at 115°C. Calculate Kp.
A) 0.0150
B) 0.478
C) 0.142
D) 1.41 × 10–4
E) none of these
The equilibrium constant K is 0.28 at 900°C. What is Kp at
this temperature?
A) 5.0  10–5
B) 4.0  10–5
C) 3.0  10–5
D) 2.0  10–5
E) 1.0  10–5
136.
Calculate Kp for
A)
B)
C)
D)
E)
Consider the following reaction:
is
1.00 – 2(0.123)
8.13
0.123
66.1
16.3
using the following data:
The value for Kp at 500 K is 1.5  10–5/atm2.
A) 7.5  10–2
B) 1.3  10–2
C) 9.6  10–2
D) 2.5  10–2
E) 6.0  10–2
135.
If the equilibrium constant for
0.123, then the equilibrium constant for
140.
Consider the following system at equilibrium:
For the reaction below, Kp = 1.16 at 800°C.
If a 20.0-gram sample of CaCO3 is put into a 10.0-liter
container and heated to 800°C, what percent of the CaCO3
Which of the following changes will shift the equilibrium to
will react to reach equilibrium?
Page 12
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
A)
B)
C)
D)
E)
141.
14.6%
65.9%
34.1%
100.0%
none of these
placed in a 5.00-L container.
147.
At equilibrium, the concentration of A is 0.40 mol/L. What
is the value of K?
A) 0.89
B) 1.80
C) 2.00
D) 3.00
E) none of these
148.
The value of K is 0.90. If 3.0 moles of A and 4.0 moles of B
had been placed in a 2.5-L container at the same
temperature, the equilibrium constant would be
A) 1.8
B) 0.45
C) 3.6
D) 0.22
E) 0.90
149.
A 10.0-g sample of solid NH4Cl is heated in a 5.00-L
At –80°C, K for the reaction
is 4.66  10–8. We introduce 0.050 mole of N2O4 into a 1.0L vessel at –80°C and let equilibrium be established. The
total pressure in the system at equilibrium will be:
A) 0.23 atm
B) 0.79 atm
C) 1.3 atm
D) 2.3 atm
E) none of these
142.
The reaction
has Kp = 45.9 at 763 K. A particular equilibrium mixture at
that temperature contains gaseous HI at a partial pressure
of 4.00 atm and hydrogen gas at a partial pressure of 0.200
atm. What is the partial pressure of I2?
A)
B)
C)
D)
E)
146.
container to 900°C. At equilibrium the pressure of NH3(g)
is 1.20 atm.
0.200 atm
0.436 atm
1.74 atm
0.574 atm
14.3 atm
The equilibrium constant, Kp, for the reaction is:
A) 1.20
B) 1.44
C) 2.40
D) 31.0
E) none of these
Consider the reaction:
150.
The following reaction is investigated (assume an ideal gas
mixture):
at constant temperature. Initially a container is filled with
pure SO3(g) at a pressure of 2 atm, after which equilibrium
is allowed to be reached. If y is the partial pressure of O2
at equilibrium, the value of Kp is:
Initially there are 0.10 moles of N2O and 0.25 moles of
A)
equilibrium?
A) 0.9
B) 0.04
C) 0.06
D) 0.02
E) none of these
N2H4, in a 10.0-L container. If there are 0.06 moles of
N2O at equilibrium, how many moles of N2 are present at
B)
C)
151.
At a certain temperature K for the reaction
is 7.5 liters/mole. If 2.0 moles of NO2 are placed in a 2.0-
D)
E) none of these
liter container and permitted to react at this temperature,
calculate the concentration of N2O4 at equilibrium.
A) 0.39 moles/liter
B) 0.65 moles/liter
C) 0.82 moles/liter
Use the following to answer questions 147-148:
For the reaction given below, 3.00 moles of A and 4.00 moles of B are
Page 13
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
D) 7.5 moles/liter
E) none of these
152.
Use the following to answer questions 157-158:
Consider the following reaction (assume an ideal gas mixture):
Initially 2.0 moles of N2(g) and 4.0 moles of H2(g) were
added to a 1.0-liter container and the following reaction
then occurred:
A 1.0-liter vessel was initially filled with pure NOBr, at a pressure of
4.0 atm, at 300 K.
157.
The equilibrium concentration of NH3(g) = 0.68 moles/liter
at 700°C. K at 700°C for the formation of ammonia is:
A) 3.6  10–3
B) 1.4  10–1
C) 1.1  10–2
D) 5.0  10–2
E) none of these
153.
A)
B)
C)
D)
E)
158.
Consider the reaction
At 1273 K the Kp
value is 167.5. What is the PCO at equilibrium if the PCO2
is 0.10 atm at this temperature?
A) 16.7 atm
B) 2.0 atm
C) 1.4 atm
D) 4.1 atm
E) 250 atm
154.
155.
156.
After equilibrium was established, the partial pressure of
NOBr was 2.5 atm. What is Kp for the reaction?
0.45
0.27
0.18
0.75
none of these
After equilibrium was reached, the volume was increased to
2.0 liters, while the temperature was kept at 300 K. This
will result in:
A) an increase in Kp.
B) a decrease in Kp.
C) a shift in the equilibrium position to the right.
D) a shift in the equilibrium position to the left.
E) none of these
Use the following to answer questions 159-160:
Nitric oxide, an important pollutant in air, is formed from the
elements nitrogen and oxygen at high temperatures, such as those
obtained when gasoline burns in an automobile engine. At 2000°C, K
for the reaction
Which of the following is true for a system whose
equilibrium constant is relatively small?
A) It will take a short time to reach equilibrium.
B) It will take a long time to reach equilibrium.
C) The equilibrium lies to the left.
D) The equilibrium lies to the right.
E) Two of these.
is 0.01.
159.
The reaction quotient for a system is 7.2  102. If the
equilibrium constant for the system is 36, what will happen
as equilibrium is approached?
A) There will be a net gain in product.
B) There will be a net gain in reactant.
C) There will be a net gain in both product and reactant.
D) There will be no net gain in either product or reactant.
E) The equilibrium constant will decrease until it equals
the reaction quotient.
Predict the direction in which the system will move to reach
equilibrium at 2000°C if 0.4 moles of N2, 0.1 moles of O2,
and 0.08 moles of NO are placed in a 1.0-liter container.
A) The system remains unchanged.
B) The concentration of NO will decrease; the
concentrations of N2 and O2 will increase.
C) The concentration of NO will increase; the
concentrations of N2 and O2 will decrease.
D) The concentration of NO will decrease; the
concentrations of N2 and O2 will remain unchanged.
E) More information is necessary.
Consider the following equilibrated system:
160.
If the Kp value
A 1-L container originally holds 0.4 mol of N2, 0.1 mol of O2,
and 0.08 mole of NO. If the volume of the container
holding the equilibrium mixture of N2, O2, and NO is
is 0.860, find the equilibrium pressure of the O2 gas if the
NO2 gas pressure is 0.520 atm and the PNO is 0.300 atm at
decreased to 0.5 L without changing the quantities of the
gases present, how will their concentrations change?
A) The concentration of NO will increase; the
concentrations of N2 and O2 will decrease.
equilibrium.
A) 1.49 atm
B) 0.78 atm
C) 0.40 atm
D) 0.99 atm
E) 2.58 atm
B) The concentrations of N2 and O2 will increase; and the
concentration of NO will decrease.
C) The concentrations of N2, O2, and NO will increase.
Page 14
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
D) The concentrations of N2, O2, and NO will decrease.
E) There will be no change in the concentrations of N2, O2,
166.Calculate the equilibrium concentration of NO(g).
A) 1.0 M
B) 1.6  10–5 M
C) 0.50 M
D) 6.2  10–4 M
E) 4.0  10–3 M
and NO.
161.
A sample of solid NH4NO3 was placed in an evacuated
container and then heated so that it decomposed
explosively according to the following equation:
167.
A)
B)
C)
D)
E)
At equilibrium the total pressure in the container was found
to be 3.20 atm at a temperature of 500°C. Calculate Kp.
A)
B)
C)
D)
E)
162.
Calculate the equilibrium concentration of Cl2(g).
4.10
1.23
2.56
4.85
1.14
1.6  10–5 M
1.0 M
0.50 M
6.2  10–4 M
4.0  10–3 M
Use the following to answer questions 168-170:
The questions below refer to the following system:
A 2-liter flask initially contains 1.2 mol of gas A and 0.60 mol of gas
C. Gas A decomposes according to the following reaction:
Consider the following reaction:
The equilibrium concentration of gas B is 0.20 mol/L.
Given 1.00 mole of HF(g), 0.500 mole of H2(g), and 0.750
168.
Determine the equilibrium concentration of gas A.
A) 0.8 M
B) 0.2 M
C) 0.6 M
D) 0.4 M
E) 0.3 M
169.
Determine the equilibrium concentration of gas C.
A) 0.3 M
B) 0.5 M
C) 0.9 M
D) 1.2 M
E) 1.5 M
170.
What is the equilibrium constant?
A) 3.6
B) 10.8
C) 1.2
D) 0.4
E) 6.3
171.
Nitrogen gas (N2) reacts with hydrogen gas (H2) to form
mole of F2(g) are mixed in a 5.00-L flask, determine the
reaction quotient, Q, and the net direction to achieve
equilibrium.
A) Q = 0.150; the equilibrium shifts to the right.
B) Q = 0.375; the equilibrium shifts to the left.
C) Q = 0.150; the equilibrium shifts to the left.
D) Q = 0.375; the equilibrium shifts to the right.
E) Q = 0.150; the system is at equilibrium.
163.
Equilibrium is reached in chemical reactions when:
A) the rates of the forward and reverse reactions become
equal.
B) the concentrations of reactants and products become
equal.
C) the temperature shows a sharp rise.
D) all chemical reactions stop.
E) the forward reaction stops.
Use the following to answer questions 164-167:
Consider the following equilibrium:
ammonia (NH3). At 200°C in a closed container, 1.0 atm of
nitrogen gas is mixed with 2.0 atm of hydrogen gas. At
equilibrium, the total pressure is 2.0 atm. Calculate the
partial pressure of hydrogen gas at equilbrium.
A) 2.0 atm
B) 0.50 atm
C) 1.5 atm
D) 0.0 atm
E) none of these
with K = 1.6  10–5. 1.00 mole of pure NOCl and 1.00 mole of pure Cl2
are replaced in a 1.00-L container.
164.
If x moles of NOCl react, what is the equilibrium
concentration of NO?
A) +x
B) +2x
C) –x
D) –2x
E) x2
172.
Given the equation
45.0
Page 15
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
mL of 0.050 M A is mixed with 25.0 mL 0.100 M B. At
equilibrium the concentration of C is 0.0410 M. Calculate K.
A) 7.3
B) 0.34
C) 0.040
D) 0.14
E) none of these
173.
E) There will be less of the hydrated cobalt ion at the new
equilibrium position.
177.
Given the reaction
E) Nothing will change.
You have the
gases A, B, C, and D at equilibrium. Upon adding gas A, the
value of K:
A) increases because by adding A, more products are
made, increasing the product to reactant ratio.
B) decreases because A is a reactant o the product to
reactant ratio decreases.
C) does not change because A does not figure into the
product to reactant ratio.
D) does not change as long as the temperature is constant.
E) depends on whether the reaction is endothermic or
exothermic.
174.
178.
The equilibrium system
has a very
small equilibrium constant:
K = 2.6  10–6. Initially 3 moles of A are placed in a 1.5-L
flask. Determine the concentration of C at equilibrium.
A) 0.011 M
B) 0.022 M
C) 0.033 M
D) 0.044 M
E) 2.0 M
The following questions refer to the equilibrium shown here:
The questions below refer to the following system:
Cobalt chloride is added to pure water. The Co2+ ions hydrate. The
hydrated form then reacts with the Cl– ions to set up the equilibrium
shown here:
Which statement below describes the change that the
system will undergo if hydrochloric acid is added?
A) It should become more blue.
B) It should become more pink.
C) The equilibrium will shift to the right.
D) The equilibrium will shift to the left.
E) Two of these.
176.
Which statement below describes the change that the
system will undergo if water is added?
A) More chlorine ions will be produced.
B) More water will be produced.
C) The equilibrium will shift to the right.
D) The color will become more blue.
Which statement below describes the change that the
system will undergo if acetone (whose density is lower than
water and is insoluble in water) is added?
A) The system will become pink on the top and blue on the
bottom.
B) The system will become blue on the top and pink on the
bottom.
C) The system will become intensely pink in the middle.
D) The system will become intensely blue on the top and
clear on the bottom.
E) The system will become intensely pink on the top and
clear on the bottom.
Use the following to answer questions 179-181:
179.
What would happen to the system if oxygen were added?
A) More ammonia would be produced.
B) More oxygen would be produced.
C) The equilibrium would shift to the right.
D) The equilibrium would shift to the left.
E) Nothing would happen.
180.
What would happen to the system if the pressure were
decreased?
A) Nothing would happen.
B) More oxygen would be produced.
C) The water vapor would become liquid water.
D) The ammonia concentration would increase.
E) The NO concentration would increase.
181.
For a certain reaction at 25.0°C, the value of K is 1.2 × 10–3.
At 50.0°C the value of K is 3.4 × 10–1. This means that the
reaction is
A) Exothermic
B) Endothermic
C) More information is needed
182.
Ammonia is prepared industrially by the reaction:
Use the following to answer questions 175-178:
175.
Which statement below describes the change that the
system will undergo if silver nitrate is added?
A) It should become more blue.
B) It should become more pink.
C) Water will be produced.
D) The silver ion will react with the CoCl42–.
For the reaction, H° = –92.2 kJ and K (at 25°C) = 4.0  108.
When the temperature of the reaction is increased to
500°C, which of the following is true?
A) K for the reaction will be larger at 500°C than at 25°C.
Page 16
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
B) At equilibrium, more NH3 is present at 500°C than at
25°C.
C) Product formation (at equilibrium) is not favored as the
temperature is raised.
D) The reaction of N2 with H2 to form ammonia is
Given the equation
particular temperature,
K = 1.6  104.
endothermic.
E) None of these is true.
188.
If you start with 2.0 M of chemical A, calculate the
equilibrium concentration of chemical C.
A) 8.3  10–3 M
B) 6.25  10–5 M
C) 2.0 M
D) 0.99 M
E) none of these
189.
If you mixed 5.0 mol B, 0.10 mol C, and 0.0010 mol A in a
one-liter container, which direction would the reaction
initially proceed?
A) To the left.
B) To the right.
C) The above mixture is the equilibrium mixture.
D) Cannot tell from the information given.
190.
At a higher temperature, K = 1.8  10–5. If you start with
2.0 M of chemical A, calculate the equilibrium concentration
of chemical C.
A) 6.0  10–3 M
B) 2.6  10–2 M
C) 1.0 M
D) 2.1  10–2 M
E) none of these
191.
Addition of chemical B to an equilibrium mixture of the
above will
A) cause [A] to increase.
B) cause [C] to increase.
C) have no effect.
D) cannot be determined.
E) none of these.
192.
Placing the equilibrium mixture in an ice bath (thus lowering
the temperature) will
A) cause [A] to increase.
B) cause [B] to increase.
C) have no effect.
D) cannot be determined.
E) none of these.
193.
Raising the pressure by lowering the volume of the
container will
A) cause [A] to increase.
B) cause [B] to increase.
C) have no effect.
D) cannot be determined.
E) none of these.
Use the following to answer questions 183-186:
Consider the following equilibrium:
183.
Addition of X2 to a system described by the above
equilibrium
A) will cause [H2] to decrease.
B) will cause [X2] to decrease.
C) will cause [H2X] to decrease.
D) will have no effect.
E) cannot possibly be carried out.
184.
Addition of argon to the above equilibrium
A) will cause [H2] to decrease.
B) will cause [X2] to increase.
C) will cause [H2X] to increase.
D) will have no effect.
E) cannot possibly be carried out.
185.
Increasing the pressure by decreasing the volume will cause
A) the reaction to occur to produce H2X.
B) the reaction to occur to produce H2 and X2.
C) the reaction to occur to produce H2 but no more X2.
D) no reaction to occur.
E) X2 to dissociate.
186.
Increasing the temperature will cause
A) the reaction to occur to produce H2X.
B) the reaction to occur to produce H2 and X2.
C) the reaction to occur to produce H2 but no more X2.
D) no reaction to occur.
E) an explosion.
187.
The value of equilibrium constant K is dependent on
I. the initial concentrations of the reactants.
II. the initial concentrations of the products.
III. the temperature of the system.
IV. the nature of the reactants and products.
A) I, II
B) II, III
C) III, IV
D) It is dependent on three of the above choices.
E) It is not dependent on any of the above choices.
Use the following to answer questions 188-193:
Page 17
At a
UNIT 6 Worksheet: Kinetics and Gas Phase Equilibria
Answer Key -- Unit 6 Worksheet
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
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55.
56.
A
A
B
C
A
E
C
A
B
C
C
E
D
D
C
C
D
E
A
A
D
D
B
C
E
B
A
E
A
D
A
D
A
A
E
B
D
A
E
E
E
C
E
A
D
B
C
D
C
B
A
B
C
A
C
B
57.
58.
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108.
109.
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111.
112.
113.
114.
C
A
E
E
A
A
D
D
A
C
B
C
C
B
A
D
A
A
B
C
A
E
E
D
B
A
D
B
B
B
C
D
A
A
B
B
C
B
B
C
B
E
B
B
C
D
D
C
D
E
B
A
E
D
C
C
B
115.
116.
117.
118.
119.
120.
121.
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Page 18
rate =k[O2]
k = 3.47 × 10–3 sec–1
567 sec
C
a) rate = k[A]2[B] b)
1.8  102 L2/mol2s
a
d
e
E
A
B
C
A
D
A
D
E
A
B
B
E
C
E
D
D
B
B
B
C
B
E
D
D
A
E
B
C
A
C
C
C
B
E
B
C
B
C
D
B
A
A
D
E
B
B
C
170.
171.
172.
173.
174.
175.
176.
177.
178.
179.
180.
181.
182.
183.
184.
185.
186.
187.
188.
189.
190.
191.
192.
193.
A
B
C
D
A
E
A
B
B
C
E
A
C
A
D
A
B
C
D
A
B
A
B
A