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Chapter 11
1.
The reaction for the Haber process, the industrial production of ammonia, is
N2(g) + 3 H2(g) → 2 NH3(g)
Assume that under certain laboratory conditions ammonia is produced at the rate
of 6.29  10–5 mol L–1 s–1. At what rate is nitrogen consumed? At what rate is
hydrogen consumed?
2.
The following data were obtained in the decomposition of H2O2(aq) to O2(g) and
H2O(). The rate at which oxygen gas was produced was measured. (No oxygen
was present initially.)
(a) Calculate the average rate in mL/min for the first 3.3 minutes.
(b) Calculate the average rate in mL/min for the first 6.9 minutes.
3.
The reaction of CO(g) + NO2(g) is second-order in NO2 and zero-order in CO at
temperatures less than 500 K.
(a) Write the rate law for the reaction.
(b) How will the reaction rate change if the NO2 concentration is halved?
(c) How will the reaction rate change if the concentration of CO is doubled?
4.
One reaction that destroys O3 molecules in the stratosphere is
NO + O3 → NO2 + O2
When this reaction was studied in the laboratory, it was found to be first order
with respect to both NO and O3, with a rate constant of 1.9  104 L mol–1 s–1. If
[NO] = 1.2  10–5 mol L–1 and [O3] = 2.0  10–5 mol L–1, what is the rate of this
reaction?
5.
The rate of the decomposition of hydrogen peroxide, H2O2, depends on the
concentration of iodide ion present. The rate of decomposition was measured at
constant temperature and pressure for various concentrations of H2O2 and of KI.
The data appear below. Determine the order of reaction for each substance, write
the rate law, and evaluate the rate constant.
6.
The following experimental data were obtained for the reaction
2A+3B→C+2D
Determine the reaction order for each reactant and the value of the rate constant.
7.
Rate data were obtained at 25°C for the following reaction. What is the rate law
for this reaction?
A+2B→C+2D
Expt. Initial [A] Initial [B]
(mol L–1) (mol L–1)
Initial Rate of
Formation of C
(mol L–1 min–1)
1
0.10
0.10
3.0  10–4
2
0.30
0.30
9.0  10–4
3
0.10
0.30
3.0  10–4
4
0.20
0.40
6.0  10–4
8.
A possible reaction for the degradation of the pesticide DDT to a less harmful
compound was simulated in the laboratory. The reaction was found to be firstorder, with k = 4.0  10–8 s–1 at 25°C. What is the half-life for the degradation of
DDT in this experiment, in years?
9.
A substance undergoes first-order decomposition. After 40.0 min at 500°C, only
12.5% of the original sample remains. What is the half-life of the decomposition?
If the original sample weighed 243 g, how much would remain after 2.00 hr?
10.
The labels on most pharmaceuticals state that the medicine should be stored in a
cool, dark place. In the context of this chapter, explain why this is sound advice.
11.
The table below presents measured rate constants for the reaction of NO with
ozone at three temperatures. From these data, determine the activation energy of
the reaction in kJ/mol. (Assume the temperatures all have two significant
figures.)
12.
Many reactions occur in the formation of photochemical smog, including the
reaction of ozone with various volatile organic chemicals, or VOCs. The table
below shows the Arrhenius expression for four such reactions.
(a) Which reaction is likely to be the fastest at 310 K?
(b) Which reaction has the lowest activation energy?
(c) Which has the highest activation energy?
Show the details of your calculation.
13.
What is the rate law for each of the following elementary reactions?
(a) Cl(g) + ICl(g) → I(g) + Cl2(g)
(b) O(g) + O3(g) → 2 O2(g)
(c) 2 NO2(g) → N2O4(g)
14.
HBr is oxidized in the following reaction:
4 HBr(g) + O2(g) → 2 H2O(g) + 2 Br2(g)
A proposed mechanism is
HBr + O2 → HOOBr
(slow)
HOOBr + HBr → 2 HOBr
(fast)
HOBr + HBr → H2O + Br2
(fast)
(a) Show that this mechanism can account for the correct stoichiometry.
(b) Identify all intermediates in this mechanism.
(c) What is the molecularity of each elementary step?
(d) Write the rate expression for each elementary step.
(e) Identify the rate-determining step.
15.
The following statements relate to the reaction for the formation of HI:
H2(g) + I2(g) → 2 HI(g)
Rate = k[H2][I2]
Determine which of the following statements are true. If a statement is false,
indicate why it is incorrect.
(a) The reaction must occur in a single step.
(b) This is a second-order reaction overall.
(c) Raising the temperature will cause the value of k to decrease.
(d) Raising the temperature lowers the activation energy for this reaction.
(e) If the concentrations of both reactants are doubled, the rate will double.
(f) Adding a catalyst in the reaction will cause the initial rate to increase.
16.
Experiments show that the reaction of nitrogen dioxide with fluorine
2 NO2(g) + F2(g) → 2 FNO2(g)
has the rate law
Rate = k[NO2][F2]
The reaction is thought to occur in two steps.
Step 1: NO2(g) + F2(g) → FNO2(g) + F(g)
Step 2: NO2(g) + F(g) → FNO2(g)
(a) Show that the sum of this sequence of reactions gives the balanced equation
for the overall reaction.
(b) Which step is rate determining?

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