Brown - Chemistry: The Central Science 15th Edition - Chapter 14

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Ch.14 - Chemical Kinetics

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Brown 15th Edition

Ch.14 - Chemical Kinetics

Problem 58a
Indicate whether each statement is true or false. (a) If you measure the rate constant for a reaction at different temperatures, you can calculate the overall enthalpy change for the reaction.
Problem 58b
Indicate whether each statement is true or false. (b) Exothermic reactions are faster than endothermic reactions.
Problem 58c
Indicate whether each statement is true or false. (c) If you double the temperature for a reaction, you cut the activation energy in half.
Problem 59
Based on their activation energies and energy changes and assuming that all collision factors are the same, rank the following reactions from slowest to fastest. (a) Ea = 45 kJ>mol; E = -25 kJ>mol (b) Ea = 35 kJ>mol; E = -10 kJ>mol (c) Ea = 55 kJ>mol; E = 10 kJ>mol
Problem 61
(a) A certain first-order reaction has a rate constant of 2.75 * 10^-2 s^-1 at 20 _x001E_C. What is the value of k at 60 _x001E_C if Ea = 75.5 kJ/mol? (b) Another first-order reaction also has a rate constant of 2.75 * 10^-2 s^-1 at 20 _x001E_C. What is the value of k at 60 _x001E_C if Ea = 125 kJ/mol?
Problem 62
Understanding the high-temperature behavior of nitrogen oxides is essential for controlling pollution generated in automobile engines. The decomposition of nitric oxide (NO) to N2 and O2 is second order with a rate constant of 0.0796 M-1s-1 at 737 _x001E_C and 0.0815 M-1s-1 at 947 _x001E_C. Calculate the activation energy for the reaction.
Problem 63
The rate of the reaction CH3COOC2H5(aq) + OH-(aq) → CH3COO-(aq) + C2H5OH(aq) was measured at several temperatures, and the following data were collected: Temperature (°C) k (M⁻¹ s⁻¹) 15 0.0521 25 0.101 35 0.184 45 0.332. Calculate the value of Ea by constructing an appropriate graph.
Problem 64
The temperature dependence of the rate constant for a reaction is tabulated as follows: Temperature (K) k 1M 1 s1 2 600 0.028 650 0.22 700 1.3 750 6.0 800 23 Calculate Ea and A.
Problem 65b
(b) What is the difference between a unimolecular and a bimolecular elementary reaction?
Problem 66
(a) Can an intermediate appear as a reactant in the first step of a reaction mechanism? (b) On a reaction energy profile diagram, is an intermediate represented as a peak or a valley?
Problem 67a
What is the molecularity of each of the following elementary reactions? Write the rate law for each. (a) Cl₂(g) → 2 Cl(g)
Problem 67b
What is the molecularity of each of the following elementary reactions? Write the rate law for each. (b) OCl^-(aq + H₂O(l) → HOCl(aq) + OH^-(aq)
Problem 67c
What is the molecularity of each of the following elementary reactions? Write the rate law for each. (c) NO(g) + Cl₂(g) → NOCl₂(g)
Problem 68a,b
What is the molecularity of each of the following elementary reactions? Write the rate law for each.
(a) 2 NO(g) → N₂O₂(g)
(c) SO₃(g) → SO₂(g) + O(g)
Problem 68b
What is the molecularity of each of the following elementary reactions? Write the rate law for each. (b)
Problem 69a
(a) Based on the following reaction profile, how many intermediates are formed in the reaction A→D?
Problem 69c
(c) Which step is the fastest?
Problem 70a
Consider the following energy profile.
(a) How many elementary reactions are in the reaction mechanism?
Problem 70b
Consider the following energy profile.
(b) How many intermediates are formed in the reaction?
Problem 70c
Consider the following energy profile.
(c) Which step is rate limiting?
Problem 72a
The decomposition of hydrogen peroxide is catalyzed by iodide ion. The catalyzed reaction is thought to proceed by a two-step mechanism:
H₂O₂(aq) + I^-(aq) → H₂O(l) + IO^-(aq) (slow)
IO^-(aq) + H₂O₂(aq) → H₂O(l) + O₂(g) + I^-(aq) (fast)
(a) Write the chemical equation for the overall process.
Problem 72b
The decomposition of hydrogen peroxide is catalyzed by iodide ion. The catalyzed reaction is thought to proceed by a two-step mechanism:
H₂O₂(aq) + I^-(aq) → H₂O(l) + IO^-(aq) (slow)
IO^-(aq) + H₂O₂(aq) → H₂O(l) + O₂(g) + I^-(aq) (fast)
(b) Identify the intermediate, if any, in the mechanism.
Problem 72c
The decomposition of hydrogen peroxide is catalyzed by iodide ion. The catalyzed reaction is thought to proceed by a two-step mechanism:
H₂O₂(aq) + I^-(aq) → H₂O(l) + IO^-(aq) (slow)
IO^-(aq) + H₂O₂(aq) → H₂O(l) + O₂(g) + I^-(aq) (fast)
(c) Assuming that the first step of the mechanism is rate determining, predict the rate law for the overall process.
Problem 73a
The reaction 2 NO1g2 + Cl21g2¡2 NOCl1g2 was performed and the following data were obtained under conditions of constant 3Cl24:
(a) Is the following mechanism consistent with the data? NO1g2 + Cl21g2ΔNOCl21g2 1fast2 NOCl21g2 + NO1g2¡2 NOCl1g2 1slow2
Problem 74a
You have studied the gas-phase oxidation of HBr by O₂: 4 HBr(g) + O₂(g) → 2 H₂O(g) + 2 Br₂(g)
You find the reaction to be first order with respect to HBr and first order with respect to O₂. You propose the following mechanism:
HBr(g) + O₂(g) → HOOBr(g)
HOOBr(g) + HBr(g) → 2 HOBr(g)
HOBr(g) + HBr(g) → H₂O(g) + Br₂(g)
(a) Confirm that the elementary reactions add to give the overall reaction.
Problem 74b
You have studied the gas-phase oxidation of HBr by O₂: 4 HBr(g) + O₂(g) → 2 H₂O(g) + 2 Br₂(g)
You find the reaction to be first order with respect to HBr and first order with respect to O₂. You propose the following mechanism:
HBr(g) + O₂(g) → HOOBr(g)
HOOBr(g) + HBr(g) → 2 HOBr(g)
HOBr(g) + HBr(g) → H₂O(g) + Br₂(g)
(b) Based on the experimentally determined rate law, which step is rate determining?
Problem 75a,b
(a) What is a catalyst? (b) What is the difference between a homogeneous and a heterogeneous catalyst?
Problem 75c
(c) Do catalysts affect the overall enthalpy change for a reaction, the activation energy, or both?
Problem 76a
(a) Most commercial heterogeneous catalysts are extremely finely divided solid materials. Why is particle size important?
Problem 78
In solution, chemical species as simple as H+ and OH- can serve as catalysts for reactions. Imagine you could measure the [H+] of a solution containing an acid-catalyzed reaction as it occurs. Assume the reactants and products themselves are neither acids nor bases. Sketch the [H+] concentration profile you would measure as a function of time for the reaction, assuming t = 0 is when you add a drop of acid to the reaction.