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Chapter 7: Problem 66

The decomposition of ozone in the stratosphere can occur by the following two- step mechanism:Step I: \(\mathrm{Br}+\mathrm{O}_{3} \rightarrow \mathrm{BrO}+\mathrm{O}_{2}\) Step II: \(\mathrm{BrO}+\mathrm{O} \rightarrow \mathrm{Br}+\mathrm{O}_{2}\) Which species is an intermediate in this mechanism? a. \(\mathrm{BrO}\) b. \(\mathrm{Br}\) c. \(\mathrm{O}_{3}\) d. \(\mathrm{O}\)

Short Answer

Expert verified
The intermediate is \(\mathrm{BrO}\).

Step by step solution

01

Understanding the Concept of Reaction Intermediates

In a chemical reaction mechanism, an intermediate is a species that is produced in one step of the mechanism and consumed in a subsequent step. Intermediates do not appear in the overall balanced equation for the reaction because they are not stable products, just temporary forms.
02

Analyzing the Mechanism Steps

Let's look at each reaction in the given mechanism:- Step I: \(\mathrm{Br} + \mathrm{O}_{3} \rightarrow \mathrm{BrO} + \mathrm{O}_{2}\)- Step II: \(\mathrm{BrO} + \mathrm{O} \rightarrow \mathrm{Br} + \mathrm{O}_{2}\)In these steps, \(\mathrm{BrO}\) is formed in Step I and used up in Step II.
03

Identifying the Intermediate

The species \(\mathrm{BrO}\) is produced in the first step and consumed in the second step, making it an intermediate. It does not appear in the overall balanced equation, as it is not a final product or reactant.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Chemical Reaction Intermediates
In many chemical reactions, particularly complex ones, certain species appear temporarily during the course of the reaction. These are known as 'intermediates.' An intermediate is formed in one step and consumed in a subsequent step. It plays a critical role in the reaction mechanism but does not appear in the net overall reaction.
Here is how intermediates function in the context of reaction mechanisms:
  • They help bridge the gap between reactants and products.
  • Intermediates are not stable; they do not accumulate during the reaction.
  • Identifying intermediates is essential to understanding the full reaction pathway.
In the decomposition of ozone, which involves steps forming and consuming species, the intermediate "BrO" stands out as it is produced in one step and then used up in the next.
Ozone Decomposition
The decomposition of ozone ( "O_3" ) in the stratosphere is a significant process, impacting both atmospheric chemistry and the Earth's climate. Ozone acts as a protective layer by absorbing ultraviolet radiation. Understanding how it breaks down is crucial.
  • Ozone molecules can be broken down into oxygen molecules ( "O_2" ) and atomic oxygen ( "O" ).
  • This process can be initiated by various catalysts, including halogens like bromine ( "Br" ).
  • The decomposition of ozone through these reactions contributes to what is known as the 'ozone hole.'
The destruction of ozone is highly related to human activities that release ozone-depleting compounds into the atmosphere, prompting a need for careful study and regulation.
Stratospheric Chemistry
The stratosphere is the second major layer of Earth's atmosphere, situated above the troposphere. This region is vital for its role in hosting the ozone layer, which shields the Earth from harmful ultraviolet radiation. Stratospheric chemistry is largely centered around the reactions involving ozone and various atmospheric pollutants.
  • The sun鈥檚 ultraviolet light plays a crucial role in splitting oxygen molecules to form ozone.
  • Reactions in the stratosphere are driven by complex photochemical processes.
  • Pollutants such as chlorofluorocarbons (CFCs) and bromine compounds can lead to the depletion of ozone.
Understanding these chemical processes is essential to predict and mitigate the impacts of anthropogenic emissions on our atmosphere.
Stepwise Reaction Analysis
Reaction mechanisms can often be broken down into multiple steps, each involving distinct reactions and species. Stepwise reaction analysis is examining these individual stages to understand the overall mechanism.
  • Each step typically involves the transition of one or more intermediates.
  • Analyzing steps helps determine the rate-determining step, which controls the speed of the overall reaction.
  • By breaking down a complex mechanism into steps, chemists can predict the behavior and interaction of compounds under study.
In ozone decomposition, the two-step mechanism clarifies the role of intermediates and provides a clear picture of how catalysts like bromine function to facilitate the reaction.

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Most popular questions from this chapter

What happens when the temperature of a reaction system is increased by \(10^{\circ} \mathrm{C}\) ? a. The effective number of collisions between the molecules possessing certain threshold energy increases atleast by \(100 \%\). b. The total number of collisions between reacting molecules increases atleast by \(100 \%\) c. The activation energy of the reaction is increased d. The total number of collisions between reacting molecules increases merely by \(1-2 \%\).

Consider the chemical reaction, \(\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightarrow 2 \mathrm{NH}_{3}(\mathrm{~g})\) The rate of this reaction can be expressed in terms of time derivatives of concentration of \(\mathrm{N}_{2}(\mathrm{~g}), \mathrm{H}_{2}\) (g) or \(\mathrm{NH}_{3}(\mathrm{~g})\). Identify the correct relationship amongst the rate expressions. a. rate \(=-\mathrm{d}\left[\mathrm{N}_{2}\right] / \mathrm{dt}=-1 / 3 \mathrm{~d}\left[\mathrm{H}_{2}\right] / \mathrm{dt}=1 / 2 \mathrm{~d}\left[\mathrm{NH}_{3}\right] / \mathrm{dt}\) b. rate \(=-\mathrm{d}\left[\mathrm{N}_{2}\right] / \mathrm{dt}=-3 \mathrm{~d}\left[\mathrm{H}_{2}\right] / \mathrm{dt}=2 \mathrm{~d}\left[\mathrm{NH}_{3}\right] / \mathrm{dt}\) c. rate \(=-\mathrm{d}\left[\mathrm{N}_{2}\right] / \mathrm{dt}=-1 / 3 \mathrm{~d}\left[\mathrm{H}_{2}\right] / \mathrm{dt}=2 \mathrm{~d}\left[\mathrm{NH}_{3}\right] / \mathrm{dt}\) d. rate \(=-\mathrm{d}\left[\mathrm{N}_{2}\right] / \mathrm{dt}=\mathrm{d}\left[\mathrm{H}_{2}\right] / \mathrm{dt}=\mathrm{d}\left[\mathrm{NH}_{3}\right] / \mathrm{dt}\)

The first order isomerisation reaction: Cyclopropane \(\rightarrow\) propene, has a rate constant of \(1.10 \times 10^{-4} \mathrm{~s}^{-1}\) at \(470^{\circ} \mathrm{C}\) and an activation energy of \(264 \mathrm{~kJ} / \mathrm{mol}\). What is the temperature of the reaction when the rate constant is equal to \(4.36 \times 10^{-3} \mathrm{~s}^{-1}\) ? a. \(240^{\circ} \mathrm{C}\) b. \(150^{\circ} \mathrm{C}\) c. \(540^{\circ} \mathrm{C}\) d. \(450^{\circ} \mathrm{C}\)

The basic theory of Arrhenius equation is that (1) Activation energy and pre exponential fact are always temperature independent (2) The number of effective collisions is prop tional to the number of molecule above a cert threshold energy. (3) As the temperature increases, the number molecules with energies exceeding the thresh energy increases. (4) The rate constant in a function of tempe ture a. 2,3 and 4 b. 1,2 and 3 c. 2 and 3 d. 1 and 3

The rate constant of a reaction is given by In \(\mathrm{k}\left(\mathrm{sec}^{-1}\right)\) \(=14.34-\left(1.25 \times 10^{4}\right) / \mathrm{T}\) What will be the energy of activation? a. \(24.83 \mathrm{kcal} \mathrm{mol}^{-1}\) b. \(49.66 \mathrm{kcal} \mathrm{mol}^{-1}\) c. \(12.42 \mathrm{kcal} / \mathrm{mol}\) d. none
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