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Chapter 4: Problem 51

Which element is oxidized, and which is reduced in the following reactions? (a) $\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)$ (b) $3 \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Al}(s) \longrightarrow$ $$ 3 \mathrm{Fe}(s)+2 \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q) $$ (c) $\mathrm{Cl}_{2}(a q)+2 \mathrm{NaI}(a q) \longrightarrow \mathrm{I}_{2}(a q)+2 \mathrm{NaCl}(a q)$ (d) $\mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(I)$

Short Answer

Expert verified
In the given reactions: (a) N is reduced and H is oxidized. (b) Fe is reduced and Al is oxidized. (c) Cl is reduced and I is oxidized. (d) S is oxidized and O is reduced.

Step by step solution

01

Assign Oxidation Numbers

First, assign oxidation numbers to each element in the reactants and products: In \(\mathrm{N}_{2}\), the oxidation number of N is 0. In \(\mathrm{H}_{2}\), the oxidation number of H is 0. In \(\mathrm{NH}_{3}\), the oxidation number of N is -3 and that of H is +1.
02

Identify Oxidation and Reduction

Compare the oxidation numbers of the elements in the reactants and products: N (oxidation number) changes from 0 to -3. H (oxidation number) changes from 0 to +1. In this reaction, N is reduced (its oxidation number decreases), and H is oxidized (its oxidation number increases). (b) \(3 \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Al}(s) \longrightarrow 3 \mathrm{Fe}(s)+2 \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}(a q)\)
03

Assign Oxidation Numbers

First, assign oxidation numbers to each element in the reactants and products: In \(\mathrm{Fe(NO}_3)_2\), the oxidation number of Fe is +2, N is +5, and O is -2. In \(\mathrm{Al}\), the oxidation number of Al is 0. In \(\mathrm{Fe}\), the oxidation number of Fe is 0. In \(\mathrm{Al(NO}_3)_3\), the oxidation number of Al is +3, N is +5, and O is -2.
04

Identify Oxidation and Reduction

Compare the oxidation numbers of the elements in the reactants and products: Fe (oxidation number) changes from +2 to 0. Al (oxidation number) changes from 0 to +3. In this reaction, Fe is reduced (its oxidation number decreases), and Al is oxidized (its oxidation number increases). (c) \(\mathrm{Cl}_{2}(a q)+2 \mathrm{NaI}(a q) \longrightarrow \mathrm{I}_{2}(a q)+2 \mathrm{NaCl}(a q)\)
05

Assign Oxidation Numbers

First, assign oxidation numbers to each element in the reactants and products: In \(\mathrm{Cl}_2\), the oxidation number of Cl is 0. In \(\mathrm{NaI}\), the oxidation number of Na is +1 and I is -1. In \(\mathrm{I}_2\), the oxidation number of I is 0. In \(\mathrm{NaCl}\), the oxidation number of Na is +1 and Cl is -1.
06

Identify Oxidation and Reduction

Compare the oxidation numbers of the elements in the reactants and products: Cl (oxidation number) changes from 0 to -1. I (oxidation number) changes from -1 to 0. In this reaction, Cl is reduced (its oxidation number decreases), and I is oxidized (its oxidation number increases). (d) \(\mathrm{PbS}(s)+4 \mathrm{H}_{2} \mathrm{O}_{2}(a q) \longrightarrow \mathrm{PbSO}_{4}(s)+4 \mathrm{H}_{2} \mathrm{O}(l)\)
07

Assign Oxidation Numbers

First, assign oxidation numbers to each element in the reactants and products: In \(\mathrm{PbS}\), the oxidation number of Pb is +2 and S is -2. In \(\mathrm{H}_2\mathrm{O}_2\), the oxidation number of H is +1 and O is -1. In \(\mathrm{PbSO}_4\), the oxidation number of Pb is +2, S is +6, and O is -2. In \(\mathrm{H}_2\mathrm{O}\), the oxidation number of H is +1 and O is -2.
08

Identify Oxidation and Reduction

Compare the oxidation numbers of the elements in the reactants and products: S (oxidation number) changes from -2 to +6. O (oxidation number) changes from -1 to -2. In this reaction, S is oxidized (its oxidation number increases), and O is reduced (its oxidation number decreases).

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Most popular questions from this chapter

(a) What volume of \(0.115 \mathrm{MHClO}_{4}\) solution is needed to neutralize \(50.00 \mathrm{~mL}\) of \(0.0875 \mathrm{MNaOH}\) ? (b) What volume of $0.128 \mathrm{MHCl}\( is needed to neutralize \)2.87 \mathrm{~g}$ of \(\mathrm{Mg}(\mathrm{OH})_{2} ?\) (c) If \(25.8 \mathrm{~mL}\) of an \(\mathrm{AgNO}_{3}\) solution is needed to precipitate all the \(\mathrm{Cl}^{-}\) ions in a \(785-\mathrm{mg}\) sample of \(\mathrm{KCl}\) (forming \(\mathrm{AgCl}\) ), what is the molarity of the \(\mathrm{AgNO}_{3}\) solution? (d) If \(45.3 \mathrm{~mL}\) of a 0.108 \(M\) HCl solution is needed to neutralize a solution of \(\mathrm{KOH}\), how many grams of KOH must be present in the solution?

True or false: (a) If a substance is oxidized, there must be more oxygen in the substance. (b) If a substance is oxidized, it must lose at least one electron and form an anion.

(a) A caesium hydroxide solution is prepared by dissolving \(3.20 \mathrm{~g}\) of \(\mathrm{CsOH}\) in water to make \(25.00 \mathrm{~mL}\) of solution. What is the molarity of this solution? (b) Then, the caesium hydroxide solution prepared in part (a) is used to titrate a hydroiodic acid solution of unknown concentration. Write a balanced chemical equation to represent the reaction between the caesium hydroxide and hydroiodic acid solutions. (c) If \(18.65 \mathrm{~mL}\) of the caesium hydroxide solution was needed to neutralize a $42.3 \mathrm{~mL}$ aliquot of the hydroiodic acid solution, what is the concentration (molarity) of the acid?

You know that an unlabeled bottle contains an aqueous solution of one of the following: \(\mathrm{AgNO}_{3}, \mathrm{CaCl}_{2},\) or \(\mathrm{Al}_{2}\left(\mathrm{SO}_{4}\right)_{3} . \mathrm{A}\) friend suggests that you test a portion of the solution with \(\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}\) and then with NaCl solutions. According to your friend's logic, which of these chemical reactions could occur, thus helping you identify the solution in the bottle? (a) Barium sulfate could precipitate. (b) Silver chloride could precipitate. (c) Silver sulfate could precipitate. (d) More than one, but not all, of the reactions described in answers a-c could occur. (e) All three reactions described in answers a-c could occur.

Suppose you have \(3.00 \mathrm{~g}\) of powdered zinc metal, \(3.00 \mathrm{~g}\) of powdered silver metal and \(500.0 \mathrm{~mL}\) of a \(0.2 \mathrm{M}\) copper(II) nitrate solution. (a) Which metal will react with the copper(II) nitrate solution? (b) What is the net ionic equation that describes this reaction? (c) Which is the limiting reagent in the reaction? (d) What is the molarity of \(\mathrm{Cu}^{2+}\) ions in the resulting solution?
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