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Chapter 4: Problem 55

Classify the following redox reactions. (a) \(2 \mathrm{H}_{2} \mathrm{O}_{2} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\) (b) \(\mathrm{Mg}+2 \mathrm{AgNO}_{3} \longrightarrow \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}\) (c) \(\mathrm{NH}_{4} \mathrm{NO}_{2} \longrightarrow \mathrm{N}_{2}+2 \mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{H}_{2}+\mathrm{Br}_{2} \longrightarrow 2 \mathrm{HBr}\)

Short Answer

Expert verified
Reaction (a) is a disproportionation redox reaction, reactions (b) and (d) are standard redox reactions, and reaction (c) is also a disproportionation redox reaction.

Step by step solution

01

Identify and classify the reaction (a)

For the reaction \(2 \mathrm{H}_{2} \mathrm{O}_{2} \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}+\mathrm{O}_{2}\), each oxygen in the \(\mathrm{H}_{2}\mathrm{O}_{2}\) changes from an oxidation state of -1 (in \(\mathrm{H}_{2} \mathrm{O}_{2}\)) to -2 (in \(\mathrm{H}_{2}\mathrm{O}\)) or 0 (in \(\mathrm{O}_{2}\)), showing that oxygen is both oxidized and reduced. Thus it is a disproportionation redox reaction.
02

Identify and classify the reaction (b)

For the reaction \(\mathrm{Mg}+2 \mathrm{AgNO}_{3} \longrightarrow \mathrm{Mg} \left(\mathrm{NO}_{3}\right)_{2}+2 \mathrm{Ag}\), Mg changes from an oxidation state of 0 to +2, indicating that it is oxidized, and Ag changes from +1 to 0, showing that it is reduced. Thus it is a redox reaction.
03

Identify and classify the reaction (c)

For the reaction \(\mathrm{NH}_{4} \mathrm{NO}_{2} \longrightarrow \mathrm{N}_{2}+2 \mathrm{H}_{2} \mathrm{O}\), the nitrogen in \(\mathrm{NH}_{4} \mathrm{NO}_{2}\) goes from -3 to 0 (in \(\mathrm{N}_{2}\)) indicating that it is oxidized and nitrogen in \(\mathrm{NO}_{2}\) goes from +3 to 0 (in \(\mathrm{N}_{2}\)) indicating that it is reduced. This is also a disproportionation redox reaction.
04

Identify and classify the reaction (d)

For the reaction \(\mathrm{H}_{2}+\mathrm{Br}_{2} \longrightarrow 2 \mathrm{HBr}\), the hydrogen is oxidized (goes from 0 to +1) and the bromine is reduced (goes from 0 to -1), this also is a redox reaction.

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Most popular questions from this chapter

Barium hydroxide, often used to titrate weak organic acids, is obtained as the octahydrate, \(\mathrm{Ba}(\mathrm{OH})_{2} \cdot 8 \mathrm{H}_{2} \mathrm{O}\). What mass of \(\mathrm{Ba}(\mathrm{OH})_{2} \cdot 8 \mathrm{H}_{2} \mathrm{O}\) would be required to make \(500.0 \mathrm{~mL}\) of a solution that is \(0.1500 \mathrm{M}\) in hydroxide ions?

A \(46.2-\mathrm{mL}, 0.568 M\) calcium nitrate \(\left[\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}\right]\) solution is mixed with \(80.5 \mathrm{~mL}\) of \(1.396 \mathrm{M}\) calcium nitrate solution. Calculate the concentration of the final solution.

Give a chemical explanation for each of the following: (a) When calcium metal is added to a sulfuric acid solution, hydrogen gas is generated. After a few minutes, the reaction slows down and eventually stops even though none of the reactants is used up. (b) In the activity series, aluminum is above hydrogen, yet the metal appears to be unreactive toward steam and hydrochloric acid. (c) Sodium and potassium lie above copper in the activity series. In your explanation, discuss why \(\mathrm{Cu}^{2+}\) ions in a \(\mathrm{CuSO}_{4}\) solution are not converted to metallic copper upon the addition of these metals. (d) A metal M reacts slowly with steam. There is no visible change when it is placed in a pale green iron(II) sulfate solution. Where should we place \(\mathrm{M}\) in the activity series? (e) Before aluminum metal was obtained by electrolysis, it was produced by reducing its chloride \(\left(\mathrm{AlCl}_{3}\right)\) with an active metal. What metals would you use to produce aluminum in that way?

Describe how you would prepare the following compounds: (a) \(\mathrm{Mg}(\mathrm{OH})_{2},\) (b) \(\mathrm{AgI}\), (c) \(\mathrm{Ba}_{3}\left(\mathrm{PO}_{4}\right)_{2}\).

A 22.02-mL solution containing \(1.615 \mathrm{~g} \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) is mixed with a 28.64-mL solution containing \(1.073 \mathrm{~g}\) \(\mathrm{NaOH}\). Calculate the concentrations of the ions remaining in solution after the reaction is complete. Assume volumes are additive.
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