AP Chemistry
Chapter 14. Chemical Kinetics
Chapter 14. Chemical Kinetics
Sample Exercise 14.1 (p. 576)
For the reaction pictured at the bottom p. 575 in textbook, calculate the average rate at which A disappears over
the time interval from 20 s to 40 s.
(1.2 x 10-2 M/s)
Practice Exercise 14.1
From the reaction pictured, calculate the average rate of appearance of B over the time interval from 0 to 40 s.
(1.8 x 10-2 M/s)
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AP Chemistry
Chapter 14. Chemical Kinetics
Average Rate vs. Instantaneous Rate
Sample Exercise 14.2 (p. 578)
Using the reaction graphed above, calculate the instantaneous rate of disappearance of C4H9Cl at t = 0 (the initial
rate).
(2.0 x 10-4 M/s)
Practice Exercise 14.2
For the reaction graphed above, calculate the instantaneous rate of disappearance of C4H9Cl at t = 300 s.
(1.1 x 10-4 M/s)
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AP Chemistry
Chapter 14. Chemical Kinetics
Reaction Rates and Stoichiometry
Sample Exercise 14.3 (p. 579)
a) How is the rate at which ozone disappears related to the rate at which oxygen appears in in the
following equation? 2 O3(g) Æ 3 O2(g)
(- 1 Δ[O3] = 1 Δ[O2])
2 Δt
3 Δt
b)
If the rate at which O2 appears, Δ[O2]/Δt, is 6.0 x 10-5 M/s at a particular instant, at what rate is O3
disappearing at this same time, - Δ[O3]/Δt?
(4.0 x 10-5 M/s)
Practice Exercise 14.3
The decomposition of N2O5 proceeds according to the following equation:
2 N2O5(g) Æ 4 NO2(g) + O2(g)
If the rate of decomposition of N2O5 at a particular instant in a reaction vessel is 4.2 x 10-7 M/s, what is the rate
of appearance of
a) NO2 (8.4 x 10-7 M/s)
b) O2 (2.1 x 10-7 M/s)
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AP Chemistry
Chapter 14. Chemical Kinetics
14.3 Concentration and Rate
Sample Exercise 14.4 (p. 582)
Consider a reaction A + B Æ C for which rate = k[A][B]2. Each of the following boxes represents a reaction
mixture in which A is shown as red spheres and B as blue ones. Rank these mixtures in order of increasing rate
of reaction.
(2 < 1 < 3)
Practice Exercise 14.4
Assuming that the rate = k[A][B], rank the mixtures represented above in order of increasing rate.
(2 = 3 < 1)
Sample Exercise 14.5 (p. 583)
a) What are the overall reaction orders for the reactions described in the following equations:
i)
2 N2O5(g) Æ 4 NO2(g) + O2(g)
Rate = k[N2O5]
ii)
CHCl3(g) + Cl2(g) Æ CCl4(g) + HCl(g)
Rate = k[CHCl3][Cl2]1/2
b) What are the units of the rate constant for the rate law for Equation (i)?
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AP Chemistry
Chapter 14. Chemical Kinetics
Practice Exercise 14.5
a) What is the reaction order of the reactant H2 in Equation (iii)?
iii)
H2(g) + I2(g) Æ 2 HI
Rate = k[H2][I2]
(1)
b) What are the units of the rate constant for Equation (iii)?
(M-1s-1)
Sample Exercise 14.6 (p. 584)
The initial rate of a reaction A + B Æ C was measured for several different starting concentrations of
A and B, and the results are as follows:
Using these data, determine
a) the rate law for the reaction
(k[A]2)
(4.0 x 10-3 M-1s-1)
b) the magnitude of the rate constant
c) the rate of the reaction when [A] = 0.050 M and [B] = 0.100 M.
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(1.0 x 10-5 M/s)
AP Chemistry
Chapter 14. Chemical Kinetics
Practice Exercise 14.6
The following data were measured for the reaction of nitric oxide with hydrogen:
2 NO(g) + 2 H2(g) Æ N2(g) + 2 H2O(g)
a) Determine the rate law for this reaction.
(k[NO]2[H2])
b) Calculate the rate constant.
(1.2 M-2s-1)
c) Calculate the rate when [NO] = 0.050 M and [H2] = 0.150 M.
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(4.5 x 10-4 M/s)
AP Chemistry
Chapter 14. Chemical Kinetics
14.4 The Change of Concentration with Time
Sample Exercise 14.7 (p. 586)
The first-order rate constant for the decomposition of a certain insecticide in water at 12oC is 1.45 yr-1. A
quantity of this insecticide is washed into a lake on June 1, leading to a concentration of 5.0 x 10-7 g/cm3 of
water. Assume that the average temperature of the lake is 12oC.
a) What is the concentration of the insecticide on June 1 of the following year?
(1.2 x 10-7 g/cm3)
b) How long will it take for the concentration of the insecticide to drop to 3.0 x 10-7 g/cm3?
(0.35 yr)
Practice Exercise 14.7
The decomposition of dimethyl ether, CH3OCH3, at 510oC is a first-order process with a rate constant of 6.8 x
10-4 s-1:
CH3OCH3 (g) Æ CH4(g) + H2(g) + CO(g)
If the initial pressure of CH3OCH3 is 135 torr, what is its partial pressure after 1420 s?
(51 torr)
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AP Chemistry
Chapter 14. Chemical Kinetics
Second-Order Reactions
Sample Exercise 14.8 (p. 588)
The following data were obtained for the gas phase decomposition of nitrogen dioxide at 300°C:
NO2(g) Æ NO(g) + 1/2 O2(g)
f
Is the reaction first or second order in NO2? (Hint: use data and plots below.)
(2nd order)
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AP Chemistry
Chapter 14. Chemical Kinetics
Practice Exercise 14.8
Consider again the decomposition of NO2 discussed above. The reaction is second order in NO2 with k = 0.543
M-1s-l. If the initial concentration of NO2 in a closed vessel is 0.0500 M, what is the remaining concentration
after 0.500 hr?
(1.0 x 10-3 M)
Sample Exercise 14.9 (p. 590)
From the figure below, estimate the half-life of C4H9Cl with water.
(~340 s)
Practice Exercise 14.9
a) Calculate t1/2 for the decomposition of the insecticide described in Sample Exercise 14.7.
(0.478 yr = 1.51 x 107 s)
b) How long does it take for the concentration of the insecticide to reach one-quarter of the initial value?
(two half-lives = 2(0.478 yr) = 0.956 yr)
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AP Chemistry
Chapter 14. Chemical Kinetics
14.5 Temperature and Rate
Sample Exercise 14.10 (p. 595)
Consider a series of reactions having the following energy profiles:
Assuming that all three reactions have nearly the same frequency factors, rank the reactions from slowest to
fastest.
( (2)< (3) < (1) )
Practice Exercise 14.10
Imagine that these reactions are reversed. Rank these reverse reactions from slowest to fastest.
( (2) < (1) < (3) )
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AP Chemistry
Chapter 14. Chemical Kinetics
Sample Exercise 14.11 (p. 596)
The following table shows the rate constants for the rearrangement of methyl isonitrile at various temperatures:
a) From these data, calculate Ea for the reaction. (160 kJ/mol)
(Hint: make a table of T, 1/T and ln k)
b) What is the value of the rate constant at 430.0 K? (1.0 x 10-6 s-1)
Practice Exercise 14.11
Using the data in Sample Exercise 14.11, above, calculate the rate constant for the rearrangement of methyl
isonitrile at 280oC.
(2.2 x 10-2 s-1)
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AP Chemistry
Chapter 14. Chemical Kinetics
14.6 Reaction Mechanisms
Sample Exercise 14.12 (p. 598)
It has been proposed that the conversion of ozone into O2 proceeds via two elementary steps:
O3(g) Æ O2(g) + O(g)
O3(g) + O(g) Æ 2 O2(g)
a) Describe the molecularity of each step in this mechanism.
b) Write the equation for the overall reaction.
c) Identify the intermediate(s).
Practice Exercise 14.12
For the reaction
Mo(CO)6 + P(CH3)3 Æ Mo(CO)5P(CH3)3 + CO
the proposed mechanism is
Mo(CO)6 Æ Mo(CO)5 + CO
Mo(CO)5 + P(CH3)3 Æ Mo(CO)5P(CH3)3
a) Is the proposed mechanism consistent with the equation for the overall reaction?
b)
Identify the intermediates.
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AP Chemistry
Chapter 14. Chemical Kinetics
Sample Exercise 14.13 (p. 600)
If the following reaction occurs in a single elementary step, predict the rate law:
H2(g) + Br2(g) Æ 2 HBr(g)
Practice Exercise 14.13
Consider the following reaction: 2 NO(g) + Br2(g) Æ 2 NOBr(g).
a) Write the rate law for the reaction, assuming it involves a single elementary step.
b) Is a single-step mechanism likely for this reaction? Why or why not?
Rate Laws for Multistep Mechanisms
Sample Exercise 14.14 (p. 602)
The decomposition of nitrous oxide, N2O, is believed to occur by a two-step mechanism:
Step 1:
N2O(g) Æ N2(g) + O(g) (slow)
Step 2: N2O(g) + O(g) Æ N2(g) + O2(g) (fast)
a)
Write the equation for the overall reaction.
b) Write the rate law for the overall reaction.
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AP Chemistry
Chapter 14. Chemical Kinetics
Practice Exercise 14.14
Ozone reacts with nitrogen dioxide to produce dinitrogen pentoxide and oxygen:
O3(g) + 2 NO2(g) Æ N2O5(g) + O2(g)
The reaction is believed to occur in two steps:
O3(g) + NO2(g) Æ NO3(g) + O2(g)
NO3(g) + NO2(g) Æ N2O5(g)
The experimental rate law is rate = k[O3][NO2]. What can you say about the relative rates of the two steps of the
mechanism?
Sample Exercise 14.15 (p. 604)
Show that the following mechanism for the equation 2 NO(g) + Br2(g) Æ 2 NOBr(g) also produces a rate law
consistent with the experimentally observed one:
k1
Step 1: NO(g) + NO(g) ⎯⎯→ N2O2(g)
←⎯
⎯⎯
k −1
Step 2: N2O2(g) + Br2(g) Æ 2 NOBr(g)
(fast equilibrium)
(slow)
Practice Exercise 14.15
The first step of a mechanism involving the reaction of bromine is
Br2(g)
k1
⎯⎯→
←⎯
⎯⎯
k −1
2 Br(g) (fast equilibrium)
What is the expression relating the concentration of Br(g) to that of Br2(g)?
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AP Chemistry
Chapter 14. Chemical Kinetics
Sample Integrative Exercise 14: Putting Concepts Together (p. 611)
Formic acid (HCOOH) decomposes in the gas phase at elevated temperatures as follows:
HCOOH(g) Æ CO2(g) + H2(g)
The decomposition reaction is determined to be first order. A graph of the partial pressure of HCOOH versus
time for decomposition at 838 K is shown as the red curve in the figure below. When a small amount of solid
ZnO is added to the reaction chamber, the partial pressure of acid versus time varies as shown by the blue curve
in the figure below.
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AP Chemistry
Chapter 14. Chemical Kinetics
a) Estimate the half-life and first-order rate constant for formic acid decomposition.
b) What can you conclude from the effect of added ZnO on the decomposition of formic acid?
c) The progress of the reaction was followed by measuring the partial pressure of formic acid vapor at
selected times. Suppose that, instead, we had plotted the concentration of formic acid in units of mol/L.
What effect would this have had on the calculated value of k?
d) The pressure of formic acid vapor at the start of the reaction is 3.00 x 102 torr. Assuming constant
temperature and ideal-gas behavior, what is the pressure in the system at the end of the reaction? If the
volume of the reaction chamber is 436 cm3, how many moles of gas occupy the reaction chamber at the
end of the reaction?
e) The standard heat of formation of formic acid vapor is ΔHof = -378.6 kJ/mol. Calculate ΔHo for the
overall reaction. Assuming that the activation energy (Ea) for the reaction is 184 kJ/mol, sketch an
approximate energy profile for the reaction, and label Ea, ΔHo, and the transition state.
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