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Chapter 7: Problem 65

A mechanism for a naturally occurring reaction that destroys ozone is: Step I: \(\mathrm{O}_{3}(\mathrm{~g})+\mathrm{HO}(\mathrm{g}) \rightarrow \mathrm{HO}_{2}(\mathrm{~g})+\mathrm{O}_{2}\) Step II: \(\mathrm{HO}_{2}(\mathrm{~g})+\mathrm{O}(\mathrm{g}) \rightarrow \mathrm{HO}(\mathrm{g})+\mathrm{O}_{2}(\mathrm{~g})\) Which species is an intermediate? a. \(\mathrm{O}\) b. \(\mathrm{O}_{3}\) c. HO d. \(\mathrm{HO}_{2}\)

Short Answer

Expert verified
The intermediate is \( \mathrm{HO}_{2} \) (option d).

Step by step solution

01

Definition Identification

In chemical reactions, intermediates are species that are formed in one step and consumed in another. They appear in the reaction mechanism but are neither the starting reactants nor the final products.
02

Analyze Step I

In the first step, \( \mathrm{O}_{3} \text{ and } \mathrm{HO} \) react to form \( \mathrm{HO}_{2} \text{ and } \mathrm{O}_{2} \).Here, \( \mathrm{HO}_{2} \) is produced.
03

Analyze Step II

In the second step, \( \mathrm{HO}_{2} \text{ and } \mathrm{O} \) react to produce \( \mathrm{HO} \text{ and } \mathrm{O}_{2} \).Here, \( \mathrm{HO}_{2} \) is consumed.
04

Identify the Intermediate

\( \mathrm{HO}_{2} \) is formed in Step I and consumed in Step II, making it an intermediate. Thus, option d, \( \mathrm{HO}_{2} \), is the intermediate species in the reaction mechanism.

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Key Concepts

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

Chemical Intermediates
In the world of chemistry, understanding chemical intermediates is crucial for grasping how reactions proceed from start to finish. Intermediates are special types of species that are produced and then later consumed as the reaction progresses. They sort of "hang out" briefly during the reaction process, making them neither reactants at the start nor products at the finish, but rather key players along the way.
In the example mechanism for ozone depletion, the species HO\(_2\) is identified as the intermediate. It is first formed in the initial step by the reaction between ozone (\(\mathrm{O}_3\)) and hydroxyl radical (\(\mathrm{HO}\)), and then consumed in the next step when reacting with an oxygen atom (\(\mathrm{O}\)). This makes it a textbook case of a chemical intermediate.
Understanding intermediates can help chemists to design reactions better and can also offer insights into how to control reaction rates.
Ozone Depletion
Ozone depletion is a significant environmental concern because the ozone layer protects Earth from harmful ultraviolet radiation. The natural mechanism involving ozone mentioned here is one way this depletion can occur.
In the process, ozone (\(\mathrm{O}_3\)) undergoes interactions that lead to its breakdown. In the first step, ozone reacts with \(\mathrm{HO}\) radicals, leading to the formation of \(\mathrm{HO}_2\) and \(\mathrm{O}_2\). Then, \(\mathrm{HO}_2\), the intermediary, reacts with an oxygen atom to regenerate \(\mathrm{HO}\) and produce more \(\mathrm{O}_2\).
This cycle illustrates how ozone can be systematically reduced in the atmosphere through such mechanisms, causing thinning of the protective layer. While this process occurs naturally, it is exacerbated by man-made chemicals, leading to increased environmental and health risks.
Step-by-step Reaction Analysis
Understanding a reaction step-by-step allows chemists to dissect complex processes into manageable parts. This way, each ingredient and its transformation become more transparent, making it easier to predict outcomes and troubleshoot problems.
In the presentation of the reaction mechanism for ozone destruction, each step needs careful breakdown: - **Step I:** The combination of \(\mathrm{O}_3\) and \(\mathrm{HO}\) leads to the formation of \(\mathrm{HO}_2\) and \(\mathrm{O}_2\). This is the formation phase of the intermediate.
- **Step II:** Here, \(\mathrm{HO}_2\), the intermediate, reacts with oxygen (\(\mathrm{O}\)) to produce \(\mathrm{HO}\) and \(\mathrm{O}_2\). In this step, the intermediate is consumed to form the final products.
This detailed step-by-step analysis enables a clear understanding of each species' role and fate, solidifying comprehension of reaction mechanisms.

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Most popular questions from this chapter

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

For a first order reaction, which of the following are not correct? a. \(t_{7 / 8}=2 t_{34}\) b. \(t_{3 / 4}=2 t_{1 / 2}\) c. \(t_{15 / 16}=4 t_{1 / 2}\) d. \(t_{15 / 6}=3 t_{3 / 4}\)

For a first order reaction \(\mathrm{A} \rightarrow \mathrm{P}\), the temperature (T) dependent rate constant (k) was found to follow the equation \(\log \mathrm{k}=-(2000) \mathrm{i} / \mathrm{T}+6.0 .\) The pre-exponential factor A and the activation energy Ea, respectively, are a. \(1.0 \times 10^{6} \mathrm{~s}^{-1}\) and \(9.2 \mathrm{~kJ} \mathrm{~mol}^{-1}\) b. \(6.0 \mathrm{~s}^{-1}\) and \(16.6 \mathrm{~kJ} \mathrm{~mol}^{-1}\) c. \(1.0 \times 10^{6} \mathrm{~s}^{-1}\) and \(16.6 \mathrm{~kJ} \mathrm{~mol}^{-1}\) d. \(1.0 \times 10^{6} \mathrm{~s}^{-1}\) and \(38.3 \mathrm{~kJ} \mathrm{~mol}^{-1}\)

At \(380^{\circ} \mathrm{C}\), half life period for the first order decomposition of \(\mathrm{H}_{2} \mathrm{O}_{2}\) is \(360 \mathrm{~min}\). The energy of activation of the reaction is \(200 \mathrm{~kJ} \mathrm{~mol}^{-1} .\) Calculate the time required for \(75 \%\) decomposition at \(450^{\circ} \mathrm{C}\) if half life for decomposition of \(\mathrm{H}_{2} \mathrm{O}_{2}\) is \(10.17 \mathrm{~min}\) at \(450^{\circ} \mathrm{C}\). a. \(20.4 \mathrm{~min}\) b. \(408 \mathrm{~min}\) c. \(10.2 \mathrm{~min}\) d. none

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 \(1-2 \%\).
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