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Chapter 10: Problem 16

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

Short Answer

Expert verified
In the given redox reaction \(4 \mathrm{Fe} + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3\), the oxidizing agent is O2, and the reducing agent is Fe.

Step by step solution

01

Write down the given reaction

The given redox reaction is: \(4 \mathrm{Fe} + 3 \mathrm{O}_2 \rightarrow 2 \mathrm{Fe}_2 \mathrm{O}_3\)
02

Determine the oxidation states of the elements before and after the reaction

Before the reaction, the oxidation state of each element is: - Fe: 0 (as it is in its pure elemental form) - O: 0 (as it is in its pure elemental form) After the reaction, the oxidation state of each element in the Fe2O3 compound is: - Fe: +3 (because it is bound to 3 oxygen atoms, each with an oxidation state of -2) - O: -2 (typical oxidation state of oxygen in compounds)
03

Identify the element being oxidized and the element being reduced

Comparing the oxidation states before and after the reaction, we can see that: - Fe has increased its oxidation state from 0 to +3, which means it has lost electrons and is being oxidized. - O has decreased its oxidation state from 0 to -2, which means it has gained electrons and is being reduced.
04

Identify the oxidizing agent and the reducing agent

The oxidizing agent is the substance that causes the oxidation of another substance (i.e., it gains electrons from the other substance). In this case, it is the O2 that causes the oxidation of Fe, so O2 is the oxidizing agent. The reducing agent is the substance that causes the reduction of another substance (i.e., it loses electrons to the other substance). In this case, it is the Fe that causes the reduction of O2, so Fe is the reducing agent.
05

Conclusion

In the given redox reaction, the oxidizing agent is O2, and the reducing agent is Fe.

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Most popular questions from this chapter

Which metal is most easily oxidized: \(\mathrm{Pt}, \mathrm{Hg}, \mathrm{Fe}\), \(\mathrm{Mg}, \mathrm{Zn} ?\)

You are trapped on a desert island with plenty of water (both fresh and salt), a drinking glass, some wire, a radio, and no batteries. You do have a tin cup, a tube of toothpaste containing stannous fluoride \(\left(\mathrm{Sn} \mathrm{F}_{2}\right.\), a source of \(\mathrm{Sn}^{2+}\) ions), a silver pendant, and undeveloped black-andwhite film (such film has silver bromide, \(\mathrm{AgBr}\), in it, a source of \(\mathrm{Ag}^{+}\) ions). (a) How would you use the above materials to construct a battery? Show how with a diagram, including an arrow over the wire to show which way the electrons flow. (You can make a salt bridge by soaking a sock in salt water and then dipping one end of the sock in one cell and the other end in the other cell.) (b) Which metal would be eaten away? Explain. (c) Which is the oxidizing agent? (d) Which is the reducing agent?

What happens when you place an active metal in a solution of ions of a less active metal?

The reaction \(\mathrm{Cu}^{2+}+\mathrm{I}_{2} \rightarrow \mathrm{Cu}^{+}+2 \mathrm{I}^{-}\) is not possible as written. Assign an oxidation number to each atom and explain what is wrong with this reaction.

Indicate whether each reaction is a redox reaction. If it is, which atom gets oxidized and which atom gets reduced? Consult the shortcut rules. \(\mathrm{P}_{4}+6 \mathrm{Br}_{2} \rightarrow 4 \mathrm{PBr}_{3}\) (Hint: \(\mathrm{Br}\) is more electronegative than P.)
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