Honors Chemistry
Ch 13 Quiz
Name ________________________
1. Given the following balanced equation, determine the rate of reaction with respect to [SO 2].
2 SO2 (g) + O2(g) → 2 SO3(g)
1 Δ[SO 2 ]
2 Dt
1 Δ[SO2 ]
b. Rate = 
2 t
Δ[SO 2 ]
c. Rate = 
Dt
a. Rate = 
d. Rate = 
2 Δ[SO 2 ]
t
e. It is not possible to determine
without more information.
2. Identify the methods used to monitor a reaction as it occurs in the reaction flask.
a. polarimeter
d. none of the above
b. spectrometer
e. all of the above
c. pressure measurement
3. Write a balanced reaction for which the following rate relationships are true.
Rate = 
1 Δ[N 2O5 ] 1 Δ[NO 2 ] Δ[O 2 ]
=
=
2
t
4 t
t
a. 2 N2O5 → 4 NO2 + O2
b. 4 NO2 + O2 → 2 N2O5
c. 2 N2O5 → NO2 + 4 O2
d.
e.
1
1
NO2 + O2 →
N2O5
4
2
1
1
N2O5 →
NO2 + O2
2
4
4. Given the following balanced equation, determine the rate of reaction with respect to [NOCl]. If the rate
of Cl2 loss is 4.84 × 10-2 M/s, what is the rate of formation of NOCl?
2 NO(g) + Cl2(g) → 2 NOCl(g)
a. 4.84 × 10-2 M/s
d. 9.68 × 10-2 M/s
-2
b. 2.42 × 10 M/s
e. 1.61 × 10-2 M/s
c. 1.45 × 10-1 M/s
5. What is the overall order of the following reaction, given the rate law?
NO(g) + O3(g) → NO2(g) + O2(g)
Rate = k[NO][O3]
a. 1st order
1
d. 1 order
b. 2nd order
2
c. 3rd order
e. 0th order
6. What are the units of k in a zero order reaction?
a.
M
s
d.
b. Ms
c. M-1s-1
e.
7. What are the units of k in the following rate law?
a.
1
Ms2
b.
c.
M2 s
M2
Rate = k[X][Y]2
1
M 2s
M2
s
s
d.
M2
s
e.
1
M 3s
8. Given the following rate law, how does the rate of reaction change if the concentration of Y is doubled?
……Rate = k [X][Y]2
a.
b.
c.
d.
e.
9.
The rate of reaction will increase by a factor of 2.
The rate of reaction will increase by a factor of 4.
The rate of reaction will increase by a factor of 5.
The rate of reaction will decrease by a factor of 2.
The rate of reaction will remain unchanged.
Given the following rate law, how does the rate of reaction change if the concentration of Y is doubled?
Rate = k [X]2[Y]3
a.
b.
c.
d.
e.
The rate of reaction will increase by a factor of 9.
The rate of reaction will increase by a factor of 2.
the rate of reaction will increase by a factor of 8.
The rate of reaction will increase by a factor of 4.
The rate of reaction will remain unchanged.
10. Determine the rate law and the value of k for the following reaction using the data provided.
CO(g) + Cl2(g) → COCl2(g)
[CO]i (M)
[Cl2]i (M) Initial Rate (M-1s-1)
0.25
0.40
0.696
0.25
0.80
1.97
0.50
0.80
3.94
a. Rate = 11 M-3/2s-1 [CO][Cl2]3/2
d. Rate = 4.4 M-1/2s-1 [CO][Cl2]1/2
b. Rate = 36 M-1.8s-1 [CO][Cl2]2.8
e. Rate = 18 M-3/2s-1 [CO]2[Cl2]1/2
c.
Rate = 17 M-2s-1 [CO][Cl2]2
11. Which of the following represents the integrated rate law for a second-order reaction?
d. k = Ae(-Ea/RT)
[A]t
a. ln
= - kt
E 1
k
[A]o
e. ln 2 = a   + ln A
k1
1
1
b.
= kt
[A]t [A]o
c.
R T
[A]t - [A]o = - kt
12. Which of the following represents the integrated rate law for a zeroth-order reaction?
d. k = Ae(-Ea/RT)
[A]t
a. ln
= - kt
E 1
k
[A]o
e. ln 2 = a   + ln A
b.
1
1
= kt
[A]t [A]o
c.
[A]t - [A]o = - kt
k1
R T
13. What data should be plotted to show that experimental concentration data fits a first-order reaction?
a. 1/[reactant] vs. time
d. ln(k) vs. 1/T
b. [reactant] vs. time
e. ln(k) vs. Ea
c. ln[reactant] vs. time
2
14. Which of the following represents the equation for a first-order half-life?
1
k[A]o
a. t 1/2 =
b. t 1/2 =
c.
t 1/2 =
0.693
k
[A]o
d. t 1/2 =
2k
[A]o
e. t 1/2 =
k
[A]o
2k
15. 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 are 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 statements are FALSE.
16. How many half-lives are required for the concentration of reactant to decrease to 25% of its original
value?
a. 1
c. 1.5
e. 2
b. 3
d. 2.5
17. The rate constant for a second-order reaction is 0.54 M-1s-1. What is the half-life of this reaction if the
initial concentration is 0.27 M?
a. 2.0 s
c. 0.25 s
e. 1.3 s
b. 5.0 s
d. 6.9 s
18. The half life for the decay of radium is 1620 years. What is the rate constant for this first-order process?
a. 4.28 × 10-4 yr-1
d. 8.91 × 10-4 yr-1
b. 1.12 × 10-4 yr-1
e. 6.17 × 10-4 yr-1
c. 2.33 × 10-4 yr-1
f.
19. The first-order decomposition of N2O at 1000 K has a rate constant of 0.76 s-1. If the initial
concentration of N2O is 10.9 M, what is the concentration of N2O after 9.6 s?
a. 7.4 × 10-3 M
d. 3.6 × 10-3 M
b. × 10-3 M
e. 8.7 × 10-3 M
-3
c. 1.4 × 10 M
20. Derive an expression for a "1/3-life" for a first-order reaction.
ln
a.
b.
c.
[A]t
[A]o
1.099
k
3
e.
k
d.
-k
0.462
k
[A]t
ln
[A]o
-3k
3
21. The second-order decomposition of NO2 has a rate constant of 0.255 M-1s-1. How much NO2
decomposes in 4.00 s if the initial concentration of NO2 (1.00 L volume) is 1.33 M?
a. 1.8 mol
d. 0.77 mol
b. 0.85 mol
e. 0.56 mol
c. 0.48 mol
22. For a reaction, what generally happens if the temperature is increased?
a. a decrease in k occurs, which results in a faster rate
b. a decrease in k occurs, which results in a slower rate
c. an increase in k occurs, which results in a faster rate
d. an increase in k occurs, which results in a slower rate
e. there is no change with k or the rate
23. Which of the following reactions would you predict to have the smallest orientation factor?
a. X2 + Y2 → 2 XY
d. N + O2 → NO2
b. NOCl2 + NO → 2 NOCl
e. All of these reactions should have
nearly identical orientation factors.
c. N2 + O2 → 2 NO
24. Which rate law is termolecular?
a. rate = k[A][B]2
b. rate = k [A][B]
c. rate = k [A]
d. rate = k [A][B][C][D]
e. rate = k [A]2
25. Identify the rate-determining step.
a. the slowest step
b. the faster step
c. the fast step
d. always the last step
e. always the second step
26. A reaction is followed and found to have a rate constant of 3.36 × 104 M-1s-1 at 344 K and a rate
constant of 7.69 M-1s-1 at 219 K. Determine the activation energy for this reaction.
a. 23.8 kJ/mol
d. 12.5 kJ/mol
b. 42.0 kJ/mol
e. 58.2 kJ/mol
c. 11.5 kJ/mol
27. A reaction is found to have an activation energy of 108 kJ/mol. If the rate constant for this reaction is
4.60 × 10-6 s-1 at 275 K, what is the rate constant at 366 K?
a. 12 s-1
d. 5.4 × 10-5 s-1
b. 1.7 s-1
e. 1.9 × 10-4 s-1
c. 0.58 s-1
28. Given the following proposed mechanism, predict the rate law for the overall reaction.
2NO2 + Cl2 → 2NO2Cl (overall reaction)
Mechanism
NO2 + Cl2 → NO2Cl + Cl
slow
NO2 + Cl → NO2Cl
fast
a. Rate = k[NO2][Cl2]
d. Rate = k[NO2Cl][Cl]
e. Rate = k[NO2Cl]2
b. Rate = k[NO2]2[Cl2]
c.
Rate = k[NO2][Cl]
29. Identify an 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
4
30. Is the activation energy for a forward reaction the same as the activation energy for the reverse of the
same reaction? Use a sketch of a reaction energy diagram.
Answer: The relationship between the activation energy of the forward and reverse reactions is related
by the energy difference between the reactants and products. If the reactants are higher in energy than
the products, the activation energy for the forward reaction will be smaller than the activation energy for
the reverse reaction and vice versa. (A sketch showing this relationship should be included).
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