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Chapter 18: Problem 19

Specify which of the following equations represent oxidation–reduction reactions, and indicate the oxidizing agent, the reducing agent, the species being oxidized, and the species being reduced. a. $\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)$ b. $2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)$ c. $\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)$ d. $2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

Short Answer

Expert verified
Equation a: Redox, Reducing agent: CH4, Oxidizing agent: H2O, Species oxidized: CH4, Species reduced: H2O Equation b: Redox, Reducing agent: Cu, Oxidizing agent: AgNO3, Species oxidized: Cu, Species reduced: AgNO3 Equation c: Redox, Reducing agent: Zn, Oxidizing agent: HCl, Species oxidized: Zn, Species reduced: HCl Equation d: Not a redox reaction.

Step by step solution

01

Equation a: CH4 (g) + H2O (g) → CO(g) + 3H2 (g)

Step 1: Determine the oxidation state of each atom: - In CH4, C has an oxidation state of -4 and H has an oxidation state of +1. - In H2O, O has an oxidation state of -2 and H has an oxidation state of +1. - In CO, C has an oxidation state of +2 and O has an oxidation state of -2. - In H2, H has an oxidation state of +1. Step 2: Identify if any changes in those oxidation states occur: - For carbon: from -4 in CH4 to +2 in CO, there's a change in the oxidation state. - For hydrogen: the oxidation state remains at +1. - For oxygen: the oxidation state remains at -2. Since there's a change in the oxidation state of carbon, this reaction is a redox reaction. Step 3: Identify the species being oxidized and reduced, and the oxidizing and reducing agents: - Oxidation: C from CH4 (from -4 to +2), so species being oxidized is CH4. - Reduction: O from H2O (from -2 to -2), no change in oxidation state but required for the oxidation of C; species being reduced is H2O. - Reducing agent: CH4. - Oxidizing agent: H2O.
02

Equation b: 2AgNO3 (aq) + Cu (s) → Cu(NO3)2 (aq) + 2Ag (s)

Step 1: Determine the oxidation state of each atom: - In AgNO3, Ag has an oxidation state of +1, N has an oxidation state of +5, and O has an oxidation state of -2. - In Cu, the oxidation state is 0. - In Cu(NO3)2, Cu has an oxidation state of +2, N has an oxidation state of +5, and O has an oxidation state of -2. - In Ag, the oxidation state is 0. Step 2: Identify if any changes in those oxidation states occur: - For silver: from +1 in AgNO3 to 0 in Ag, there's a change in the oxidation state. - For copper: from 0 in Cu to +2 in Cu(NO3)2, there's a change in the oxidation state. - For nitrogen and oxygen: the oxidation states remain unchanged. Since there's a change in the oxidation state of silver and copper, this reaction is a redox reaction. Step 3: Identify the species being oxidized and reduced, and the oxidizing and reducing agents: - Oxidation: Cu (from 0 to +2), so species being oxidized is Cu. - Reduction: Ag from AgNO3 (from +1 to 0), species being reduced is AgNO3. - Reducing agent: Cu. - Oxidizing agent: AgNO3.
03

Equation c: Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g)

Step 1: Determine the oxidation state of each atom: - In Zn, the oxidation state is 0. - In HCl, H has an oxidation state of +1 and Cl has an oxidation state of -1. - In ZnCl2, Zn has an oxidation state of +2 and Cl has an oxidation state of -1. - In H2, H has an oxidation state of +1. Step 2: Identify if any changes in those oxidation states occur: - For zinc: from 0 in Zn to +2 in ZnCl2, there's a change in the oxidation state. - For hydrogen: from +1 in HCl to +1 in H2, no change in the oxidation state. - For chlorine: the oxidation state remains at -1. Since there's a change in the oxidation state of zinc, this reaction is a redox reaction. Step 3: Identify the species being oxidized and reduced, and the oxidizing and reducing agents: - Oxidation: Zn (from 0 to +2), so species being oxidized is Zn. - Reduction: H from HCl (from +1 to +1), no change in oxidation state but required for the oxidation of Zn; species being reduced is HCl. - Reducing agent: Zn. - Oxidizing agent: HCl.
04

Equation d: 2H+(aq) + 2CrO4(2-)(aq) → Cr2O7(2-)(aq) + H2O (l)

Step 1: Determine the oxidation state of each atom: - In H+, H has an oxidation state of +1. - In CrO4(2-), Cr has an oxidation state of +6, and O has an oxidation state of -2 - In Cr2O7(2-), Cr has an oxidation state of +6, and O has an oxidation state of -2. - In H2O, H has an oxidation state of +1 and O has an oxidation state of -2. Step 2: Identify if any changes in those oxidation states occur: - For all atoms involved (H, Cr, and O) their oxidation states remain unchanged. Since no changes in the oxidation states occur in this reaction, it is not a redox reaction. Therefore, no oxidizing or reducing agents can be identified.

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Most popular questions from this chapter

A disproportionation reaction involves a substance that acts as both an oxidizing and a reducing agent, producing higher and lower oxidation states of the same element in the products. Which of the following disproportionation reactions are spontaneous under standard conditions? Calculate $\Delta G^{\circ}\( and \)K\( at \)25^{\circ} \mathrm{C}$ for those reactions that are spontaneous under standard conditions. a. $2 \mathrm{Cu}^{+}(a q) \longrightarrow \mathrm{Cu}^{2+}(a q)+\mathrm{Cu}(s)$ b. $3 \mathrm{Fe}^{2+}(a q) \longrightarrow 2 \mathrm{Fe}^{3+}(a q)+\mathrm{Fe}(s)$ c. $\mathrm{HClO}_{2}(a q) \longrightarrow \mathrm{ClO}_{3}^{-}(a q)+\mathrm{HClO}(a q) \quad$ (unbalanced) Use the half-reactions: $\mathrm{ClO}_{3}^{-}+3 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}_{2}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.21 \mathrm{V}$ $\mathrm{HClO}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \longrightarrow \mathrm{HClO}+\mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.65 \mathrm{V}$

The following standard reduction potentials have been determined for the aqueous chemistry of indium: $$\operatorname{In}^{3+}(a q)+2 \mathrm{e}^{-} \longrightarrow \operatorname{In}^{+}(a q) \quad \mathscr{E}^{\circ}=-0.444 \mathrm{V}$$ $$\operatorname{In}^{+}(a q)+\mathrm{e}^{-} \longrightarrow \operatorname{In}(s) \qquad \quad \mathscr{E}^{\circ}=-0.126 \mathrm{V}$$ a. What is the equilibrium constant for the disproportionation reaction, where a species is both oxidized and reduced, shown below? $$3 \ln ^{+}(a q) \longrightarrow 2 \operatorname{In}(s)+\operatorname{In}^{3+}(a q)$$ b. What is \(\Delta G_{i}^{\circ}\) for \(\operatorname{In}^{+}(a q)\) if $\Delta G_{f}^{\circ}=-97.9 \mathrm{kJ} / \mathrm{mol}\( for \)\operatorname{In}^{3+}(a q) ?$

Assign oxidation numbers to all the atoms in each of the following: a. \(\mathrm{HNO}_{3} \quad\) g. \(\mathrm{PbSO}_{4}\) b. \(\mathrm{CuCl}_{2} \quad\) h. \(\mathrm{PbO}_{2}\) c. \(\mathrm{O}_{2} \quad\) i. \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) d. \(\mathrm{H}_{2} \mathrm{O}_{2} \quad\) j. \(\mathrm{CO}_{2}\) e. $\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6} \quad \mathrm{k} .\left(\mathrm{NH}_{0}\right)_{2} \mathrm{Ce}\left(\mathrm{SO}_{4}\right)_{3}$ f. \(\mathrm{Ag} \quad\) l. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)

Give the balanced cell equation and determine \(\mathscr{E}^{\circ}\) for the galvanic cells based on the following half-reactions. Standard reduction potentials are found in Table 18.1. a. $\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+14 \mathrm{H}^{+}+6 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cr}^{3+}+7 \mathrm{H}_{2} \mathrm{O}$ $\mathrm{H}_{2} \mathrm{O}_{2}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}$ b. \(2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{H}_{2}\) \(\mathrm{Al}^{3+}+3 \mathrm{e}^{-} \rightarrow \mathrm{Al}\)

When jump-starting a car with a dead battery, the ground jumper should be attached to a remote part of the engine block. Why?
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