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Chapter 14: Problem 70

(a) What is meant by the term molecularity? (b) Why are termolecular elementary reactions so rare? (c) What is an intermediate in a mechanism?

Short Answer

Expert verified
(a) Molecularity refers to the number of molecules involved in an elementary reaction step or single-event collisions between participating molecules. (b) Termolecular elementary reactions are rare due to the improbability of simultaneous collision between three molecules with the correct alignment, orientation, and sufficient energy to overcome the activation energy barrier. (c) An intermediate in a reaction mechanism is a transient species formed during one of the elementary reaction steps and consumed in a subsequent step, usually unstable and highly reactive, not observed directly in the overall balanced equation.

Step by step solution

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a) Molecularity Definition

The term "molecularity" refers to the number of molecules (atoms, ions, or radicals) involved in an elementary reaction step or single-event collisions between participating molecules. In other words, molecularity is the count of reactant particles in a given elementary reaction, which determines the reaction rate. Molecularity can be unimolecular (involving one molecule), bimolecular (involving two molecules), or termolecular (involving three molecules).
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b) Termolecular Reactions Rarity

The rarity of termolecular elementary reactions can be explained by the improbability of simultaneous collision between three molecules. For a termolecular reaction to take place, such collision should result in the correct alignment of the molecules, and also have the accurate orientation and sufficient energy to overcome the activation energy barrier. The chances of these factors aligning adequately in a simultaneous three-molecule collision are minimal compared to unimolecular or bimolecular reactions. Consequently, termolecular elementary reactions are considered rare.
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c) Intermediate in a Mechanism Definition

An intermediate in a reaction mechanism is a transient species formed during one of the elementary reaction steps and then consumed in a subsequent step. Intermediates are usually unstable and highly reactive, which leads to their fast decay and short lifetimes before reacting further. Unlike reactants or products, intermediates are not observed directly in the overall balanced equation of the reaction. Identifying and understanding intermediates involved in reaction mechanisms are crucial for gaining insight into the reaction's true pathway, rate laws, and kinetics.

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Most popular questions from this chapter

(a) What factors determine whether a collision between two molecules will lead to a chemical reaction? (b) According to the collision model, why does temperature affect the value of the rate constant? (c) Does the rate constant for a reaction generally increase or decrease with an increase in reaction temperature?

The rate of disappearance of HCl was measured for the following reaction: $$ \mathrm{CH}_{3} \mathrm{OH}(a q)+\mathrm{HCl}(a q) \longrightarrow \mathrm{CH}_{3} \mathrm{Cl}(a q)+\mathrm{H}_{2} \mathrm{O}(l) $$ The following data were collected: $$ \begin{array}{rl} \hline \text { Time (min) } & \text { [HCI] (M) } \\ \hline 0.0 & 1.85 \\ 54.0 & 1.58 \\ 107.0 & 1.36 \\ 215.0 & 1.02 \\ 430.0 & 0.580 \\ \hline \end{array} $$ (a) Calculate the average rate of reaction, in \(M / \mathrm{s}\), for the time interval between each measurement. (b) Calculate the average rate of reaction for the entire time for the data from \(t=0.0 \mathrm{~min}\) to \(t=430.0 \mathrm{~min} .\) (c) Graph [HCl] versus time and determine the instantaneous rates in \(M / \min\) and \(M / s\) at \(t=75.0 \mathrm{~min}\) and \(t=250\) min.

Indicate whether each statement is true or false. If it is false, rewrite it so that it is true. (a) If you measure the rate constant for a reaction at different temperatures, you can calculate the overall enthalpy change for the reaction. (b) Exothermic reactions are faster than endothermic reactions. (c) If you double the temperature for a reaction, you cut the activation energy in half.

The reaction \(2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{NO}_{2}(g)\) is second order in NO and first order in \(\mathrm{O}_{2}\). When [NO] \(=0.040 \mathrm{M}\) and \(\left[\mathrm{O}_{2}\right]=0.035 \mathrm{M},\) the observed rate of disappearance of \(\mathrm{NO}\) is \(9.3 \times 10^{-5} \mathrm{M} / \mathrm{s}\). (a) What is the rate of disappearance of \(\mathrm{O}_{2}\) at this moment? (b) What is the value of the rate constant? (c) What are the units of the rate constant? (d) What would happen to the rate if the concentration of NO were increased by a factor of \(1.8 ?\)

What is the molecularity of each of the following elementary reactions? Write the rate law for each. (a) \(2 \mathrm{NO}(g) \longrightarrow \mathrm{N}_{2} \mathrm{O}_{2}(g)\)
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