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Chapter 10: Problem 62

In the upper atmosphere, sunlight can convert oxygen to ozone: \(2 \mathrm{O}_{2} \rightarrow \mathrm{O}_{3}+\mathrm{O}\) Ozone Is this a redox reaction? Completely justify your answer.

Short Answer

Expert verified
No, this is not a redox reaction. The oxidation states of oxygen in the reactants (O₂) and the products (O₃ and O) are all 0, and no changes in oxidation states occur during this reaction.

Step by step solution

01

Identify the initial and final oxidation states of the elements involved

Start by identifying the initial and final oxidation states for oxygen in the reactants (O₂) and in the products (O₃ and O). The oxygen molecule (O₂) has an oxidation state of 0, since it is in its elemental form. For ozone (O₃), we will need to determine the oxidation states of the individual oxygen atoms. For atomic oxygen (O), the oxidation state is also 0, as it is the elemental form of oxygen.
02

Determine the oxidation states of oxygen in ozone (O₃)

Let's assign oxidation state variable x to each oxygen atom in O₃, and set up the equation: 3x = 0 (since the overall charge of the O₃ molecule is 0) Solving this equation for x yields: x = 0 So, each oxygen atom in ozone (O₃) has an oxidation state of 0.
03

Compare the initial and final oxidation states of the elements

Now that we have found the oxidation states of oxygen in the reactants and the products, we can compare them: Oxygen gas (O₂): Oxidation state 0 Ozone (O₃): Oxidation state 0 Atomic oxygen (O): Oxidation state 0
04

Determine if the reaction is a redox reaction

A reaction is classified as a redox reaction if there is a change in the oxidation states of the elements involved. In this case, the oxidation states of oxygen in the reactants (O₂) and the products (O₃ and O) remain unchanged at 0. Therefore, this reaction is not a redox reaction.

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Most popular questions from this chapter

For each redox reaction, indicate which substance is the oxidizing agent and which is the reducing agent. \(4 \mathrm{Fe}+3 \mathrm{O}_{2} \rightarrow 2 \mathrm{Fe}_{2} \mathrm{O}_{3}\)

Consider the perchlorate \(\left(\mathrm{ClO}_{4}^{-}\right)\) ion. (a) Assign oxidation states to all of the atoms in perchlorate. (b) Explain why the chlorine atom in perchlorate does not follow the halide rule. (c) Do you think perchlorate would be a powerful oxidizing agent or a powerful reducing agent? Explain. (d) Would it be possible to have an even higher oxidation state for the \(C 1\) atom in some other compound? Explain. (Hint: Chlorine is in group VIIA).

Suppose an atom has an oxidation state of \(+3\). (a) Would more or fewer electrons be assigned to this atom by oxidation-state bookkeeping than are present on a free atom of that element? (b) How many more or fewer?

You are trying to determine the identity of an unknown metal. You place a strip of it in a solution of \(\mathrm{Mg}^{2+}\), and no reaction occurs. You then place a strip of it in a solution of \(\mathrm{Zn}^{2+}\), and zinc metal plates out on the strip. Name one possibility for the identity of the unknown metal.

Which ion is most difficult to reduce: \(\mathrm{Mn}^{2+}, \mathrm{Hg}^{2+}, \mathrm{Fe}^{3+}, \mathrm{Mg}^{2+}, \mathrm{Li}^{+} ?\)
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