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Chapter 18: Problem 19

Specify which of the following equations represent oxidation鈥搑eduction 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) 鈫� 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) 鈫� 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) 鈫� 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) 鈫� 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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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Oxidizing Agent
In chemistry, an oxidizing agent, or oxidant, is a substance that causes another substance to lose electrons. When a substance loses electrons, it is said to be oxidized. The oxidizing agent itself gains electrons in this process. This means that in a reaction, the oxidizing agent is reduced.
  • It helps another substance to oxidize by accepting electrons from it.
  • The oxidizing agent achieves this typically by adding oxygen to the other substance or removing hydrogen.
For example, in the reaction \[ \text{Cu}(s) + \text{AgNO}_3(aq) \rightarrow \text{Cu(NO}_3\text{)}_2(aq) + \text{Ag}(s) \] Silver nitrate (\( \text{AgNO}_3 \)) serves as an oxidizing agent. Here, silver ions (\( \text{Ag}^+ \)) accept electrons from copper, thereby causing copper to be oxidized to copper(II). Silver nitrate is reduced to metallic silver in the process.
Reducing Agent
The reducing agent, sometimes called a reductant, is a compound that donates electrons to another chemical species in a redox reaction. When it donates electrons, the reducing agent itself is oxidized. It helps another substance to be reduced by losing electrons.
  • Reducing agents facilitate the gaining of electrons by the oxidized reactant.
  • Typical reducing agents include metals and certain anions like \( \text{Fe}^{2+} \)and \( \text{C}\text{l}^- \).
In the reaction \[ \text{Zn}(s) + 2 \text{HCl}(aq) \rightarrow \text{ZnCl}_2(aq) + \text{H}_2(g) \] zinc (\( \text{Zn} \)) acts as the reducing agent. It loses electrons and thus is oxidized to \( \text{Zn}^{2+} \). Meanwhile, the hydrogen ions in \( \text{HCl} \) gain electrons and are reduced to form hydrogen gas.
Oxidation States
Oxidation states (also known as oxidation numbers) help determine how many electrons are lost or gained by an atom in a chemical reaction. They show the hypothetical charge that an atom would have if all bonds were 100% ionic.
  • Oxidation states allow us to identify which atoms are oxidized or reduced in reactions.
  • A change in the oxidation state of an element usually indicates a redox reaction.
For instance, in the reaction\[ \text{CH}_4(g) + \text{H}_2\text{O}(g) \rightarrow \text{CO}(g) + 3 \text{H}_2(g) \]carbon changes its oxidation state from \( -4 \) in methane to \( +2 \) in carbon monoxide. Hence, carbon is oxidized, indicating a redox process.
Redox Reaction Identification
Redox, or oxidation-reduction reactions, involve the transfer of electrons between two species. Identifying these reactions is key in understanding many chemical processes.
  • A redox reaction involves a change in oxidation state between the reactants and products.
  • The species whose oxidation state increases is oxidized, while the one whose oxidation state decreases is reduced.
For example, in\[ \text{2AgNO}_3 (aq) + \text{Cu} (s) \rightarrow \text{Cu(NO}_3\text{)}_2 (aq) + \text{2Ag} (s) \]copper is oxidized from \( 0 \) to \( +2 \), while silver is reduced from \( +1 \) to \( 0 \). To determine if a reaction is redox, look for changes in the oxidation states of atoms within the reactants and products.

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Most popular questions from this chapter

Estimate \(\mathscr{E}^{\circ}\) for the half-reaction $$2 \mathrm{H}_{2} \mathrm{O}+2 \mathrm{e}^{-} \longrightarrow \mathrm{H}_{2}+2 \mathrm{OH}^{-}$$ given the following values of \(\Delta G_{\mathrm{f}}^{\circ} :\) $$\quad\quad\quad \mathrm{H}_{2} \mathrm{O}(l)=-237 \mathrm{kJ} / \mathrm{mol}$$ $$\mathrm{H}_{2}(g)=0.0$$ $$\quad\quad\quad \mathrm{OH}^{-}(a q)=-157 \mathrm{kJ} / \mathrm{mol}$$ $$\quad \mathrm{e}^{-}=0.0$$ Compare this value of \(\mathscr{E}^{\circ}\) with the value of \(\mathscr{E}^{\circ}\) given in Table 18.1

Sketch the galvanic cells based on the following half- reactions. Show the direction of electron flow, show the direction of ion migration through the salt bridge, and identify the cathode and anode. Give the overall balanced equation, and determine \(\mathscr{E}^{\circ}\) for the galvanic cells. Assume that all concentrations are 1.0 \(M\) and that all partial pressures are 1.0 atm. a. \(\mathrm{Cl}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Cl}^{-} \quad \mathscr{E}^{\circ}=1.36 \mathrm{V}\) \(\mathrm{Br}_{2}+2 \mathrm{e}^{-} \rightarrow 2 \mathrm{Br}^{-} \quad \mathscr{E}^{\circ}=1.09 \mathrm{V}\) b. \(\mathrm{MnO}_{4}^{-}+8 \mathrm{H}^{+}+5 \mathrm{e}^{-} \rightarrow \mathrm{Mn}^{2+}+4 \mathrm{H}_{2} \mathrm{O} \quad \mathscr{E}^{\circ}=1.51 \mathrm{V}\) \(\mathrm{IO}_{4}^{-}+2 \mathrm{H}^{+}+2 \mathrm{e}^{-} \rightarrow \mathrm{IO}_{3}^{-}+\mathrm{H}_{2} \mathrm{O} \quad \quad \mathscr{E}^{\circ}=1.60 \mathrm{V}\)

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 \(\operatorname{In}^{3+}(a q) ?\)

A chemist wishes to determine the concentration of \(\mathrm{CrO}_{4}^{2-}\) electrochemically. A cell is constructed consisting of a saturated calomel electrode (SCE; see Exercise 111\()\) and a silver wire coated with \(\mathrm{Ag}_{2} \mathrm{CrO}_{4}\) . The \(\mathscr{E}^{\circ}\) value for the following half-reaction is 0.446 \(\mathrm{V}\) relative to the standard hydrogen electrode: $$\mathrm{Ag}_{2} \mathrm{CrO}_{4}+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Ag}+\mathrm{CrO}_{4}^{2-}$$ a. Calculate \(\mathscr{E}_{\text { cell } \text { and }} \Delta G\) at \(25^{\circ} \mathrm{C}\) for the cell reaction when \(\left[\mathrm{CrO}_{4}^{2-}\right]=1.00 \mathrm{mol} / \mathrm{L}\) . b. Write the Nernst equation for the cell. Assume that the SCE concentrations are constant. c. If the coated silver wire is placed in a solution (at \(25^{\circ} \mathrm{C} )\) in which \(\left[\mathrm{CrO}_{4}^{2-}\right]=1.00 \times 10^{-5} M,\) what is the expected cell potential? d. The measured cell potential at \(25^{\circ} \mathrm{C}\) is 0.504 \(\mathrm{V}\) when the coated wire is dipped into a solution of unknown \(\left[\mathrm{CrO}_{4}^{2-}\right] .\) What is \(\left[\mathrm{CrO}_{4}^{2-}\right]\) for this solution? e. Using data from this problem and from Table \(18.1,\) calculate the solubility product \(\left(K_{\mathrm{sp}}\right)\) for \(\mathrm{Ag}_{2} \mathrm{CrO}_{4}\).

The overall reaction and standard cell potential at \(25^{\circ} \mathrm{C}\) for the rechargeable nickel-cadmium alkaline battery is \(\mathrm{Cd}(s)+\mathrm{NiO}_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) \(\mathrm{Ni}(\mathrm{OH})_{2}(s)+\mathrm{Cd}(\mathrm{OH})_{2}(s) \qquad \mathscr{E}^{\circ}=1.10 \mathrm{V}\) For every mole of Cd consumed in the cell, what is the maximum useful work that can be obtained at standard conditions?
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