Unit 5 Homework
Chapter 14 read pages 622-641
1) Consider the reaction 4PH3(g) → P4(g) + 6H2(g) If, in a certain experiment, over a specific
time period, 0.0048 mole of PH3 is consumed in a 2.0 L container each second of reaction, what
are the rates of production of P4 and H2 in this experiment?
2) At 40oC, hydrogen peroxide solution will decompose according to the following reaction:
2H2O2(aq) → 2H2O(l) + O2(g)
The following data were collected for the concentration of H2O2 at various times:
Time (s)
[H2O2] (mol/L)
0
1.000
2.16 x 104
0.500
4
4.32 x 10
0.250
a. Calculate the average rate of decomposition of H2O2 between 0 and 2.16 x 104 s. Use this rate
to calculate the average rate of production of O2(g) over the same period.
b. What are these rates for time period 2.16 x 104s and 4.32 x 104s.
3) What are the units for each of the following if the concentrations are expressed in moles per liter
and the time is in seconds?
a. Rate of chemical reaction.
b. Rate constant for a zero-order rate law.
c. Rate constant for a first-order rate law.
d. Rate constant for a second-order rate law.
e. Rate constant for a third-order rate law.
4) The reaction: 2NO(g) + Cl2(g) → 2NOCl(g) was studied at -10oC. The following results were
obtained:
[NO]0 (mol/L)
[Cl2]0 (mol/L)
Initial Rate (mol/L . min)
0.10
0.10
0.18
0.10
0.20
0.36
0.20
0.20
1.45
a. What is the rate law?
b. What is the value of the rate constant?
5) The decomposition of nitrosyl chloride was studied 2NOCl(g) → 2NO(g) + Cl2(g) The
following data were obtained
[NOCl]0 (molecules/cm3)
Initial Rate (molecules/cm3 . sec)
16
3.0 x 10
5.98 x 104
2.0 x 1016
2.66 x 104
16
1.0 x 10
6.64 x 103
4.0 x 1016
1.06 x 105
a. What is the rate law?
b. Calculate the value of the rate constant.
c. Calculate the value of the rate constant when concentration is given in moles per liter.
6) The reaction I-(aq) + OCl-(aq) → IO-(aq) + Cl-(aq) was studied and the following data
obtained.
[I-]0 (mol/L)
[OCl-] (mol/L)
Initial rate (mol/L . sec)
0.12
0.18
7.91 x 10-2
0.060
0.18
3.95 x 10-2
0.030
0.090
9.88 x 10-3
0.24
0.090
7.91 x 10-2
a. What is the rate law?
b. Calculate the value of the rate constant.
c. Calculate the initial rate for an experiment where both I- and OCl- are initially present at 0.15 mol/L.
7) The rate of reaction between hemoglobin (Hb) and carbon monoxide (CO) was studied at 20 oC.
The following data were collected with all concentration units in mol/L.
[Hb]0 (mol/L)
[CO]0 (mol/L)
Initial rate (mol/L. s)
2.21
1.00
0.619
4.42
1.00
1.24
4.42
3.00
3.71
a.
b.
c.
d.
Determine the orders of this reaction with respect to Hb and CO.
Determine the rate law.
Calculate the value of the rate constant.
What would be the initial rate for an experiment with [HB]0 = 3.36 mol/L and [CO]0 =
2.40 mol/L?
8) The decomposition of H2O2 was studied and the following data were obtained at a particular
temperature
Time (s)
[H2O2] (mol/L)
0
1.00
120
0.91
300
0.78
600
0.59
1200
0.37
1800
0.22
2400
0.13
3000
0.082
3600
0.050
a.
b.
c.
d.
9)
Determine the rate law
Determine the value of the rate constant.
Determine the integrated rate law.
Calculate the concentration of hydrogen peroxide at 4000. S after the start of the reaction.
The rate law for the decomposition of phosphine (PH3) is Rate = k[PH3].
It takes 120. S for 1.oo M PH3 to decrease to decrease to 0.250 M. How much time is
required for 2.00 M PH3 to decrease to a concentration of 0.350 M?
10) The rate of the reaction NO2(g) + CO(g) → NO(g) + CO2(g) depends only on the
concentration of nitrogen dioxide at temperatures below 225oC. At a temperature below 225oC
the following data were collected:
Time (s)
0
1.20 x 103
3.00 x 103
4.50 x 103
9.00 x 103
1.80 x 103
[NO2] (mol/L)
0.500
0.444
0.381
0.340
0.250
0.174
a. Determine the rate law.
b. Determine the integrated rate law and the value of the rate constant
c. Calculate [NO2] at 2.70 x 104 s after the start of the reaction.
11) The decomposition of ethanol (C2H5OH) on an alumina (Al2O3) surface
C2H5OH(g) → C2H4(g) + H2O(g)
was studied at 600 K. Concentration vs. time data were collected for this reaction, and a plot
of [A] vs. time resulted in a straight line with a slope of -4.00 x 10-5 mol/L . s.
a. Determine the rate law, the integrated rate law, and the value of the rate constant for this
reaction.
b. If the initial concentration of C2H5OH was 1.25 x 10-2 M, calculate the half-life for this
reaction.
c. How much time is required for all the 1.25 x 10-2 M C2H5OH to decompose?
12) The reaction A → B e + C is known to be zero order in A and to have a rate constant of
5.0 x 10-2 mol/L . s at 25oC. An experiment was run at 25oC where [A]0 is 1.0 x 10-3 M.
a. Write the integrated rate law for this reaction.
b. Calculate the half-life for this reaction.
c. Calculate the concentration of B after 5.0 x 10-3 s has elapsed assuming [B]0 = 0.
13) A certain first-order reaction is 45.0% complete in 65 s. What are the values of the rate
constant and the half-life for his process.
14) A certain substance, initially at 0.10 M in solution, decomposes by second-order kinetics. If
the rate constant for this process is 0.40 L/mol . min, how much time is required for the
concentration to reach 0.020 M?
15) For the reaction A → products, successive half-lives are observed to be 10.0, 20.0, and 40.0
min. for an experiment in which [A]0 = 0.10 M. Calculate the concentration at the following
times.
a. 80.0 min
b. 30.0 min
16) The experimentally determined rate law for the decomposition of hydrogen peroxide is rate =
k[H2O2] a possible mechanism is:
H2O2 → 2OH
H2O2 + OH → H2O + HO2
HO2 + OH → H2O + O2
What is the rate determining step?
Identify any intermediates.
What is the overall balanced equation for this reaction?
17) The mechanism for the gas-phase reaction of nitrogen dioxide with carbon monoxide to form
nitric oxide and carbon dioxide is thought to be:
NO2 + NO2 → NO3 + NO
Slow
NO3 + CO → NO2 + CO2
Fast
Write the rate law for this reaction.
What is the overall balanced equation for the reaction?
18) A proposed mechanism for a reaction is
C4H9Br → C4H9+ + BrSlow
C4H9+ + H2O → C4H9OH2+
Fast
+
+
C4H9OH2 + H2O → C4H9OH + H3O Fast
Write the rate law expected for this mechanism. What is the overall balanced equation for this
reaction? What are the intermediates in the proposed mechanism?
19) One mechanism for the destruction of ozone in the upper atmosphere is:
O3(g) + NO(g) → NO2(g) + O2(g)
Slow
NO2(g) + O(g) → NO(g) O2(g)
Fast
Overall reaction:
O3(g) + O(g) → 2O2(g)
a. Which species is a catalyst?
b. Which species is an intermediate?
20) The reaction
NO2(g) + CO(g) → NO(g) + CO2(g) has the energy profile shown on slide
35 of the PowerPoint. Which of the following mechanisms for this reaction is consistent with
the energy profile?
a. NO2(g) + NO2(g) → NO3(g) + NO(g)
Fast
NO3((g) + CO(g) → NO2(g) + CO2(g)
Slow
b. NO2(g) + NO2(g) → NO3(g) + NO(g)
NO3((g) + CO(g) → NO2(g) + CO2(g)
Draw the energy profile and indicate on it:
a. The position and identities of the reactant and products.
b. The activation energy for the overall reaction.
c. E for the reaction.
d. The position and identity of the intermediate.
e. Write the rate law for this reaction.
Slow
Fast
Answers
1) P4: 6.0 x 10-4 mol/L . s
H2: 3.6 x 10-3 mol/L . s
2) a. rate of decomposition of H2O2 = 2.31 x 10-5 mol/L.s
rate of product of O2 = 1.16 X 10-5 mol/L.s
b. rate of decomposition of H2O2 = 1.16 x 10-5 mol/L.s
rate of product of O2 = 5.80 X 10-6 mol/L.s
3) a. mol/L. s b. mol/L. s c. s-1 d. L/mol . s e. L2/mol2 . s
4) a. rate = k[NO]2[Cl2]
b.1.8 x 102 L2/mol2 . min
2
5) a. rate = k[NOCl]
b. 6.6 x 10-29 cm3/ molecule . s
c. 4.0 x 10-8 L/mol.s
6) a. rate = k[I-][OCl-]
b. 3.7 L/mol.s
c. 0.083 mol/L. s
7) a. first order in Hb and first order in CO
b. rate = k[Hb][CO]
.
c. 0.280 L/mol s
d. 2026 mol/L. s
8) a. Rate = k[H2O2]; b. k= 8.3 x 10-4 s-1 c. ln[H2O2] = -kt + ln[H2O2]0 d. 0.037M
9) a. Rate = k[NO2]2 b.
1 = kt + 1 ; k = 2.08 x 10-4 l/mol.s
c. 0.037 M
[NO2]
[NO2]0
10)
a. rate = k; [C2H5OH] = -kt + [C2H5OH]0; k = 4.00 x 10-5 mol/L.s b. 156 s c. 313 s
11)
a. [A] = -kt + [A]0
b. 1.0 x 10-2 s
c. 2.5 x 10-4 M
12)
− s-1 ; 75 s
13)
150. s
14)
1.0 x 102 min
15)
a. 1.1 x 10-2 M
b. 0.025 M
0
