Mr. Dehne
Name: _____________________________________
Date: _______________
Per#: ___
AP Chem
Kinetics
Topic #5
All answers must be written on another sheet of paper and for problems where calculations are necessary, you must show work.
Kinetics WS#1 (Reaction Rates/Rate Laws from Experimental Data: Initial Rates Method)
1. The reaction 2NO(g) + Cl2(g) → 2NOCl(g) was studied at -10oC. The following results were obtained where
Rate = -Δ[Cl2]
Δt
[NO]0 (mol/L)
[Cl2]0 (mol/L)
Initial Rate (mol/L●s)
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?
2. The decomposition of nitrosyl chloride was studied: 2NOCl(g) ↔ 2NO(g) + Cl 2(g). The following data were obtained where
Rate = -Δ[NOCl]
Δt
3
[NOCl]0 (molecules/cm ) Initial Rate (molecules/cm3●s)
3.0x1016
5.98x104
2.0x1016
2.66x104
16
1.0x10
6.64x103
16
4.0x10
1.06x105
a. What is the rate law?
c. Calculate the rate constant for the concentrations given in moles per liter.
b. Calculate the rate constant.
3. Consider the reaction 4PH3(g) → P4(g) + 6H2(g). If, in a certain experiment, over a specific time period, 0.0048 mol PH3 is
consumed in a 2.0L container each second of reaction, what are the rates of the production of P 4 and H2 in this experiment?
4. What are the units for each of the following if the concentrations are expressed in moles per liter and the time in seconds?
a. Rate of a chemical reaction.
d. Rate constant for a second-order rate law.
b. Rate constant for a zero-order rate law. e. Rate constant for a third-order rate law. (Look in notes)
c. Rate constant for a first-order rate law.
5. The rate of the reaction between hemoglobin (Hb) and carbon monoxide (CO) was studied at 20 oC. The following data were
collected with all concentration units in μmol/L. (A hemoglobin concentration of 2.21 μmol/L is equal to 2.21x10-6mol/L.)
[Hb]0(μmol/L)
[CO]0 (μmol/L)
Initial Rate (μmol/L●s)
2.21
1.00
0.619
4.42
1.00
1.24
4.42
3.00
3.71
a. Determine the orders of this reaction with respect to Hb and CO.
b. Determine the rate law.
c. Calculate the value of the rate constant.
d. What would be the initial rate for an experiment with [Hb]0 = 3.36 μmol/L and [CO]0 = 2.40 μmol/L?
6. The decomposition of hydrogen peroxide was studied, and the following data were obtained at a particular temperature:
Time (s)
[H2O2] (mol/L)
0
1.00
120 ± 1
0.91
300 ± 1
0.78
600 ± 1
0.59
1
1200 ± 1
0.37
1800 ± 1
0.22
2400 ± 1
0.13
3000 ± 1
0.082
3600 ± 1
0.050
Assuming that Rate = -Δ[H2O2] determine the rate law, the integrated rate law, and the value of the rate constant. Calculate
Δt
[H2O2] at 4000.s after the start of the reaction.
7. It took 143s for 50.0% of a particular substance to decompose. If the initial concentration was 0.060M and the decomposition
reaction follows second-order kinetics, what is the value of the rate constant?
8. For the reaction A→ products, successive half-loves are observed to be 10.0, 20.0, and 40min for an experiment in which [A] 0 =
0.10M. Calculate the concentrations of A at the following times.
a. 80.0min
b. 30.0min.
9. The rate of the reaction NO2(g) + CO(g) → NO(g) + CO2(g) depends only on the concentration of nitrogen dioxide below 225 oC.
At a temperature below 225oC, the following date were collected:
Time(s)
[NO2] (mol/L)
0
0.500
1.20x103
0.444
3.00x103
0.381
4.50x103
0.340
9.00x103
0.250
1.80x104
0.174
Determine the rate law, the integrated law, and the value of the rate constant. Calculate [NO 2] at 2.70x104s after the start of the
reaction.
10. The decomposition of ethanol (C2H5OH) on an alumina (Al2O3) surface C2H5OH(g) →C2H4(g) + H2O(g) was studied at 600K.
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 4.00x10-5mol/L●s.
a. Determine the rate law, the integrated rate law, and the value of the rare constant for this reaction.
b. If the initial concentration of C2H5OH was 1.25x10-2M, calculate the half-life for this reaction.
c. How much time is required for all the 1.25x10 -2M C2H5OH to decompose?
11. The dimerization of butadiene 2C4H6(g) → C8H12(g) was studied at 500.K, and the following data were obtained:
Time(s)
[C4H6] (mol/L)
195
1.6x10-2
604
1.5x10-2
1246
1.3x10-2
2180
1.1x10-2
6210
0.68x10-2
Assuming that Rate = - Δ[C4H6]/Δt determine the form of the rate law, the integrated rate law, and the rate constant for this
reaction. (These are actual experimental data, so they may not give a perfectly straight line.)
12. Experimental data for the reaction A → 2B + C have been plotted in the following three different ways (with concentration units
in mol/L):
What is the order of the reaction with respect to A and what is the initial concentration of A?
2
13. The reaction A → B + C is known to be zero order in A and to have a rate constant of 5.0x10 -2mol●L-1●s-1 at 25oC. An
experiment was run at 25oC where [A]o = 1.0x10-3M.
a. Write the integrated rate law for this reaction.
b. Calculate the half-life for the reaction.
c. Calculate the concentration of B after 5.0x10-3s has elapsed.
14. A first-order reaction is 38.5% complete in 480. S.
a. Calculate the rate constant.
b. What is the value of the half-life?
c. How long will it take for the reaction to go to 25%, 75%, and 95% completion?
15. It took 143s for 50% of a particular substance to decompose. If the initial concentration was 0.060M and the decomposition
reaction follows second-order kinetics, what is the value of the rate constant?
16. For the reaction A → products, successive half-lives are observed to be 10.0, 20.0, and 40.0min for an experiment in which [A] o =
0.10M. Calculate the concentration of A at the following times
a. 80.0min
b. 30.0min
Kinetics WS#2 (Reaction Mechanisms)
17. Write the rate laws for the following elementary reactions.
a. CH3NC(g) → CH3CN(g)
b. O3(g) + NO(g) → O2(g) + NO2(g)
c. O3(g) → O2(g) + O(g)
d. O3(g) + O(g) → 2O2(g)
18. A proposed mechanism for a reaction is
C4H9Br → C4H9+ + Br(slow)
C4H9+ + H2O → C4H9OH2+
(fast)
C4H9OH2+ + H2O → C4H9OH + H3O+
(fast)
Write the rate law expected for this mechanism. What is the overall balanced equation for the reaction? What are the
intermediates in the proposed mechanism?
Kinetics WS#3 (Temperature Dependence of Rate Constants and the Collision Model/Catalysts)
19. For the following reaction profile, indicate
a. The positions of reactants and products.
b. The activation energy.
c. ΔE for the reaction
E
reaction coordinate
20. The activation energy for the reaction NO2(g) + CO(g) →NO(g) + CO2(g) is 125kJ/mol, and ΔE for the reaction is -216kJ/mol.
What is the activation energy for the reverse reaction [NO(g) + CO2(g) → NO2(g) + CO(g)]?
21. The rate constant for the gas-phase decomposition of N2O5, N2O5 → 2NO2 + ½ O2 has the following temperature dependence:
T(K)
k(s-1)
338
4.9x10-3
318
5.0x10-4
298
3.5x10-5
Make the appropriate graph using these data, and determine the activation energy for this reaction.
22. At 25oC the first-order rate constant for a reaction is 2.0x103s-1. The activation energy is 15.0kJ/mol. What is the value of the rate
constant at 75oC?
23. A certain reaction has an activation energy of 54.0kJ/mol. As the temperature is increased from 22 oC to a higher temperature, the
rate constant increases by a factor of 7.00. Calculate the higher temperature.
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24. Which of the following reactions would you expect to proceed at a faster rate at room temperature? Why? (Hint: Think of which
reaction would have the lower activation energy.)
2Ce4+(aq) + Hg22+(aq) → 2Ce3+(aq) + 2Hg2+(aq)
H3O+(aq) + OH-(aq) → 2H2O(l)
25. 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?
c. Ea for the uncatalyzed reaction O3(g) + O(g) → 2O2(g) is 14.0kJ. Ea for the same reaction when catalyzed is 11.9kJ.
What is the ratio of the rate constant for the catalyzed reaction to that for the uncatalyzed reaction at 25oC? Assume that
the frequency factor A is the same for each reaction.
Extra Credit
1.
Answer Key.
1. (a) rate = k[NO]2[Cl2]
(b) 1.8x102 L2/mol2●s
2
2. (a) rate = k[NOCl]
(b) 6.6x10-29 cm3/molecules●s
(c) 4.0x10-8L/mol●s
3. P4 – 6.0x10-4mol/L
H2 – 3.6x10-3mol/L
4. (a) mol/L●s
(b) mol/L●s
(c) s-1
(d) L/mol●s
(e) L2/mol2●s
st
5. (a) 1 order for Hb and CO (b) rate = k[Hb][CO]
(c) 0.288L/μmol●s
(d) 2.26μmol/L●s
-4 -1
6. rate = k[H2O2], ln[H2O2] = -kt + ln[H2O2]0, k = 8.3x10 s , 0.037M
7. 0.12L/mol●s
8. (a) 1.1x10-3M
(b) 0.025M
9. rate = k[NO2]2, 1/[NO2] = kt + 1/[NO2]0, k = 2.08x10-4L/mol●s, 0.131M
10. (a) rate = k, [C2H5OH] = -kt + [C2H5OH]0, slope = -k, then k = 4.00x10-5mol/L●s (b) 156s
(c) 313s
11. ? 33
12. ? 35
13. ? 37
14. ? 39
15. ? 41
16. ? 43
17. (a) rate = k[CH3NC]
(b) rate = k[O3][NO]
(c) rate = k[O3]
(d) rate = k[O3][O]
18. Rate = k[C4H9Br], C4H9Br + 2H2O → C4H9OH + Br- + H3O+, intermediates - C4H9+ and C4H9OH2+
19.
E
reactants
20.
21.
22.
23.
24.
25.
Ea
products
ΔE
Reaction Coordinate
? 51
Graph ln(k) vs. 1/T is linear with slope = -Ea/R = -1.2x104K, Ea = 1.0x102kJ/mol
4.8x103s-1
? 57
? 59
? 61
4