Consider the following mechanism: (1) \(\mathrm{ClO}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{HClO}(a q)+\mathrm{OH}^{-}(a q) \quad\) [fast (2) \(\mathrm{I}^{-}(a q)+\mathrm{HClO}(a q) \longrightarrow \mathrm{HIO}(a q)+\mathrm{Cl}^{-}(a q)\) [slow] (3) \(\mathrm{OH}^{-}(a q)+\mathrm{HIO}(a q) \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{IO}^{-}(a q)\) [fast (a) What is the overall equation? (b) Identify the intermediate(s), if any. (c) What are the molecularity and the rate law for each step? (d) Is the mechanism consistent with the actual rate law: Rate = \(k\left[\mathrm{ClO}^{-}\right]\left[\mathrm{I}^{-}\right] ?\)

Chemical kinetics explores the rates at which chemical reactions occur and the factors that influence these rates. It plays a crucial role in understanding how fast a reaction takes place and what conditions can alter its speed. In the given mechanism, we have three steps happening at different rates. To determine the overall reaction rate, we need to consider the slowest step since it limits the rate of the entire process.

• **Reaction rate**: The speed at which reactants turn into products.
• **Rate-determining step**: The slowest step in a reaction mechanism, which dictates the overall rate.
• **Factors affecting reaction rate**: Concentration of reactants, temperature, surface area, and presence of a catalyst.

By concentrating on the reaction kinetics, we can understand how the individual steps contribute to the overall reaction mechanism.

Rate laws are mathematical expressions that describe the relationship between the concentration of reactants and the rate of a chemical reaction. They are determined through experimental observation and can be inferred from the rate-determining step in a reaction mechanism. The general form of a rate law is:

• **General rate law**: \text{Rate} = k[A]^m[B]^n
• **k**: The rate constant, which is specific to a reaction at a given temperature.
• **m and n**: The reaction orders, indicating how the rate is affected by the concentration of reactants.

For the given mechanism, the experimental rate law is Rate = k[\(\text{ClO}^-\)][\(\text{I}^-\)]. This rate law is consistent with the slow step of the mechanism (Step 2), where the rate is primarily dependent on the concentrations of \(\text{I}^-\) and \(\text{HClO}\). By applying the steady-state approximation, we can relate the concentration of intermediates to the rate law and verify its consistency.

Molecularity refers to the number of reactant molecules involved in an elementary step of a reaction mechanism. Unlike reaction orders, molecularity is a theoretical concept used to describe individual steps rather than the overall reaction. The three types of molecularity are:

• **Unimolecular**: Involves a single reactant molecule decomposing or rearranging (e.g., \(\text{A} \rightarrow \text{products}\)).
• **Bimolecular**: Involves two reactant molecules colliding to form products (e.g., \(\text{A} + \text{B} \rightarrow \text{products}\)).
• **Termolecular**: Involves three reactant molecules colliding simultaneously (very rare due to low probability).

In the given mechanism:

• **Step 1**: Bimolecular, involving \(\text{ClO}^-\) and \(\text{H}_2\)O.
• **Step 2**: Bimolecular, involving \(\text{I}^-\) and \(\text{HClO}\).
• **Step 3**: Bimolecular, involving \(\text{OH}^-\) and \(\text{HIO}\).

Understanding molecularity helps us discern why certain steps proceed at different rates and how collisions between molecules lead to product formation.

Reaction intermediates are species that are formed during the course of a reaction mechanism but do not appear in the overall balanced equation. They are produced in one step and consumed in subsequent steps. Identifying intermediates gives insight into the step-by-step progress of a reaction. In the given mechanism, we identify the following intermediates:

• **\(\text{HClO}\)**: Formed in Step 1 and consumed in Step 2.
• **\(\text{OH}^-\)**: Formed in Step 1 and consumed in Step 3.
• **\(\text{HIO}\)**: Formed in Step 2 and consumed in Step 3.

Intermediates are crucial for applying the steady-state approximation, where we assume the concentration of intermediates remains relatively constant over the course of the reaction. This simplification allows us to solve for the overall rate law more easily. By understanding the role of intermediates, we gain a clearer picture of how reactants are converted to products through a series of elementary steps.