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Chapter 14: Problem 18

(a) What are the units usually used to express the rates of reactions occurring in solution? (b) As the temperature increases, does the reaction rate increase or decrease? (c) As a reaction proceeds, does the instantaneous reaction rate increase or decrease?

Short Answer

Expert verified
(a) The units usually used to express rates of reactions in solutions are molarity per unit time (M/s or M/min). (b) As the temperature increases, the reaction rate generally increases due to higher kinetic energy and more frequent and effective collisions between reacting molecules. (c) As a reaction proceeds, the instantaneous reaction rate typically decreases because the concentration of reactants decreases, leading to less frequent collisions between reacting molecules or particles.

Step by step solution

01

Question (a): Units for Reaction Rates

The units most commonly used to express the rates of reactions occurring in solutions are molarity per unit time (M/s or M/min).
02

Question (b): Effect of Temperature on Reaction Rate

As the temperature of the reaction system increases, the reaction rate usually increases. This is because as the temperature increases, the (average) kinetic energy of molecules or particles in the system also increases. This leads to more frequent collisions between them, and it also results in a higher probability of collisions having enough energy to overcome the activation energy barrier, leading to a faster reaction rate.
03

Question (c): Instantaneous Reaction Rate as Reaction Proceeds

As a reaction proceeds, the instantaneous reaction rate typically decreases. This is primarily because the concentration of reactants decreases as the reaction progresses since the reactants get used up to form products. Lower reactant concentrations lead to less frequent collisions between reacting molecules or particles, lowering the overall rate of reaction over time.

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Most popular questions from this chapter

You perform a series of experiments for the reaction $\mathrm{A} \rightarrow 2 \mathrm{~B}$ and find that the rate law has the form, rate \(=k[\mathrm{~A}]^{x} .\) Determine the value of \(x\) in each of the following cases: (a) The rate increases by a factor of \(6.25,\) when \([\mathrm{A}]_{0}\) is increased by a factor of \(2.5 .(\mathbf{b})\) There is no rate change when \([\mathrm{A}]_{0}\) is increased by a factor of \(4 .(\mathbf{c})\) The rate decreases by a factor of \(1 / 2,\) when \([\mathrm{A}]_{0}\) is cut in half.

The decomposition of hydrogen peroxide is catalyzed by iodide ion. The catalyzed reaction is thought to proceed by a two-step mechanism: $$ \begin{aligned} \mathrm{H}_{2} \mathrm{O}_{2}(a q)+\mathrm{I}^{-}(a q) & \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{IO}^{-}(a q)(\text { slow }) \\ \mathrm{IO}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}_{2}(a q) & \longrightarrow \mathrm{H}_{2} \mathrm{O}(l)+\mathrm{O}_{2}(g)+\mathrm{I}^{-}(a q) \quad(\text { fast }) \end{aligned} $$ (a) Write the chemical equation for the overall process. (b) Identify the intermediate, if any, in the mechanism. (c) Assuming that the first step of the mechanism is rate determining, predict the rate law for the overall process.

(a) In which of the following reactions would you expect the orientation factor to be more important in leading to reaction: $\mathrm{O}_{3}+\mathrm{O} \longrightarrow 2 \mathrm{O}_{2}\( or \)\mathrm{NO}+\mathrm{NO}_{3} \longrightarrow 2 \mathrm{NO}_{2} ?$ (b) What is related to the orientation factor? Which, smaller or larger ratio of effectively oriented collisions to all possible collisions, would lead to a smaller orientation factor?

(a) Develop an equation for the half-life of a zero-order reaction. (b) Does the half-life of a zero-order reaction increase, decrease, or remain the same as the reaction proceeds?

Many metallic catalysts, particularly the precious-metal ones, are often deposited as very thin films on a substance of high surface area per unit mass, such as alumina \(\left(\mathrm{Al}_{2} \mathrm{O}_{3}\right)\) or silica \(\left(\mathrm{SiO}_{2}\right) .(\mathbf{a})\) Why is this an effective way of utilizing the catalyst material compared to having powdered metals? (b) How does the surface area affect the rate of reaction?
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