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Chapter 14: Problem 66

According to collision theory, chemical reactions occur through molecular collisions. A unimolecular elementary process in a reaction mechanism involves dissociation of a single molecule. How can these two ideas be compatible? Explain.

Short Answer

Expert verified
Chemical reactions occur through molecular collisions according to the collision theory. Even a unimolecular elementary process, which involves the dissociation of a single molecule, can still be explained by this principle. The single molecule collides with other molecules or walls of a container and, provided it receives enough energy (activation energy) from the collision, it can dissociate. Thus, the two theories are compatible.

Step by step solution

01

Understanding Collision Theory

Collision theory postulates that for a reaction to occur, it is necessary for the reactive molecules to collide with one another. Not every collision causes a reaction; the molecules must possess a certain minimum amount of energy called the activation energy.
02

Understanding Unimolecular Elementary Process

Unimolecular elementary processes involve a single molecule. It could be that this molecule dissociates into two or more parts, changes structure, or reacts with light. It seems contrary to the collision theory as no collision appears to be involved in this process.
03

Relating Collision Theory and Unimolecular Processes

Considering a gas molecule in a container, it collides with other gas molecules as well as the walls of the container. While many of these collisions may not lead to a reaction, some can provide the molecule with sufficient energy (activation energy) to dissociate. This occurrence aligns with both the collision theory and the characteristic of a unimolecular process, making them compatible.

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Most popular questions from this chapter

One of the following statements is true and the other is false regarding the first-order reaction 2A \(\longrightarrow \mathrm{B}+\mathrm{C}\). Identify the true statement and the false one, and explain your reasoning. (a) The rate of the reaction decreases as more and more of \(\mathrm{B}\) and \(\mathrm{C}\) form. (b) The time required for one-half of substance \(A\) to react is directly proportional to the quantity of A present initially.

You want to test the following proposed mechanism for the oxidation of HBr. $$\begin{array}{c} \mathrm{HBr}+\mathrm{O}_{2} \stackrel{k_{1}}{\longrightarrow} \mathrm{HOOBr} \\\ \mathrm{HOOBr}+\mathrm{HBr} \stackrel{k_{2}}{\longrightarrow} 2 \mathrm{HOBr} \\\ \mathrm{HOBr}+\mathrm{HBr} \stackrel{k_{3}}{\longrightarrow} \mathrm{H}_{2} \mathrm{O}+\mathrm{Br}_{2} \end{array}$$ You find that the rate is first order with respect to HBr and to \(\mathrm{O}_{2}\). You cannot detect HOBr among the products. (a) If the proposed mechanism is correct, which must be the rate-determining step? (b) Can you prove the mechanism from these observations? (c) Can you disprove the mechanism from these observations?

The initial rate of the reaction \(A+B \longrightarrow C+D\) is determined for different initial conditions, with the results listed in the table. (a) What is the order of reaction with respect to A and to B? (b) What is the overall reaction order? (c) What is the value of the rate constant, \(k ?\) $$\begin{array}{llll} \hline \text { Expt } & \text { [A], M } & \text { [B], M } & \text { Initial Rate, M s }^{-1} \\ \hline 1 & 0.185 & 0.133 & 3.35 \times 10^{-4} \\ 2 & 0.185 & 0.266 & 1.35 \times 10^{-3} \\ 3 & 0.370 & 0.133 & 6.75 \times 10^{-4} \\ 4 & 0.370 & 0.266 & 2.70 \times 10^{-3} \\ \hline \end{array}$$

The mechanism proposed for the reaction of \(\mathrm{H}_{2}(\mathrm{g})\) and \(\mathrm{I}_{2}(\mathrm{g})\) to form \(\mathrm{HI}(\mathrm{g})\) consists of a fast reversible first step involving \(\mathrm{I}_{2}(\mathrm{g})\) and \(\mathrm{I}(\mathrm{g}),\) followed by a slow step. Propose a two-step mechanism for the reaction \(\mathrm{H}_{2}(\mathrm{g})+\mathrm{I}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{HI}(\mathrm{g}),\) which is known to be first order in \(\mathrm{H}_{2}\) and first order in \(\mathrm{I}_{2}.\)

In the reaction \(2 \mathrm{A}+\mathrm{B} \longrightarrow \mathrm{C}+3 \mathrm{D},\) reactant \(\mathrm{A}\) is found to disappear at the rate of \(6.2 \times 10^{-4} \mathrm{M} \mathrm{s}^{-1}.\) (a) What is the rate of reaction at this point? (b) What is the rate of disappearance of \(\mathrm{B}\) ? (c) What is the rate of formation of D?
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