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Chapter 18: Problem 17

Specify which of the following equations represent oxidationreduction 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
In summary: a. Redox reaction. Oxidizing agent: H2O, Reducing agent: CH4, Species oxidized: C in CH4, Species reduced: O in H2O. b. Redox reaction. Oxidizing agent: AgNO3, Reducing agent: Cu, Species oxidized: Cu in Cu(s), Species reduced: Ag in AgNO3. c. Redox reaction. Oxidizing agent: HCl, Reducing agent: Zn, Species oxidized: Zn in Zn(s), Species reduced: H in HCl. d. Not a redox reaction.

Step by step solution

01

Calculate Oxidation States

Compare the oxidation states of each element before and after the reaction: - C in CH4: -4 - C in CO: +2 - H in H2O: +1 - H in H2: +1 Since there is an increase in the oxidation state of Carbon (from -4 to +2), this reaction is a redox process.
02

Identify the Oxidizing and Reducing Agents

- In CH4, Carbon is oxidized from -4 to +2. - In H2O, Oxygen is reduced from -2 to 0 in the product H2. So the oxidizing agent is H2O, the reducing agent is CH4, species being oxidized is C in CH4, and the species being reduced is O in H2O. #b. Identifying Reaction Type: 2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s)#
03

Calculate Oxidation States

Compare the oxidation states of each element before and after the reaction: - Ag in AgNO3: +1 - Ag in Ag(s): 0 - Cu in Cu(s): 0 - Cu in Cu(NO3)2: +2 Since there are changes in the oxidation states, this reaction is also a redox process.
04

Identify the Oxidizing and Reducing Agents

- In AgNO3, Silver (Ag) is reduced from +1 to 0. - In Cu(s), Copper is oxidized from 0 to +2. So the oxidizing agent is AgNO3, the reducing agent is Cu, the species being oxidized is Cu in Cu(s), and the species being reduced is Ag in AgNO3. #c. Identifying Reaction Type: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)#
05

Calculate Oxidation States

Compare the oxidation states of each element before and after the reaction: - Zn in Zn(s): 0 - Zn in ZnCl2: +2 - H in HCl: +1 - H in H2: 0 As there is a change in the oxidation states, this reaction is also a redox process.
06

Identify the Oxidizing and Reducing Agents

- In Zn(s), Zinc is oxidized from 0 to +2. - In HCl, Hydrogen is reduced from +1 to 0. So the oxidizing agent is HCl, the reducing agent is Zn, the species being oxidized is Zn in Zn(s), and the species being reduced is H in HCl. #d. Identifying Reaction Type: 2H+(aq) + 2CrO4^(2-)(aq) → Cr2O7^(2-)(aq) + H2O(l)#
07

Calculate Oxidation States

Compare the oxidation states of each element before and after the reaction: - H in H+: +1 - H in H2O: +1 - Cr in CrO4^(2-): +6 - Cr in Cr2O7^(2-): +6 As there is no change in the oxidation states of any element, this reaction is not a redox process.

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Most popular questions from this chapter

Balance the following oxidation-reduction reactions that occur in acidic solution using the half-reaction method. a. \(\mathrm{I}^{-}(a q)+\mathrm{ClO}^{-}(a q) \rightarrow \mathrm{I}_{3}^{-}(a q)+\mathrm{Cl}^{-}(a q)\) b. \(\mathrm{As}_{2} \mathrm{O}_{3}(s)+\mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{H}_{3} \mathrm{AsO}_{4}(a q)+\mathrm{NO}(g)\) c. \(\mathrm{Br}^{-}(a q)+\mathrm{MnO}_{4}^{-}(a q) \rightarrow \mathrm{Br}_{2}(l)+\mathrm{Mn}^{2+}(a q)\) d. \(\mathrm{CH}_{3} \mathrm{OH}(a q)+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q) \rightarrow \mathrm{CH}_{2} \mathrm{O}(a q)+\mathrm{Cr}^{3+}(a q)\)

Three electrochemical cells were connected in series so that the same quantity of electrical current passes through all three cells. In the first cell, \(1.15 \mathrm{~g}\) chromium metal was deposited from achromium(III) nitrate solution. In the second cell, \(3.15 \mathrm{~g}\) osmium was deposited from a solution made of \(\mathrm{Os}^{n+}\) and nitrate ions. What is the name of the salt? In the third cell, the electrical charge passed through a solution containing \(\mathrm{X}^{2+}\) ions caused deposition of \(2.11 \mathrm{~g}\) metallic \(\mathrm{X}\). What is the electron configuration of \(\mathrm{X}\) ?

An unknown metal \(\mathrm{M}\) is electrolyzed. It took \(74.1 \mathrm{~s}\) for a current of \(2.00 \mathrm{~A}\) to plate out \(0.107 \mathrm{~g}\) of the metal from a solution containing \(\mathrm{M}\left(\mathrm{NO}_{3}\right)_{3} .\) Identify the metal.

Balance the following oxidation-reduction reactions that occur in basic solution. a. \(\mathrm{Al}(s)+\mathrm{MnO}_{4}^{-}(a q) \rightarrow \mathrm{MnO}_{2}(s)+\mathrm{Al}(\mathrm{OH})_{4}^{-}(a q)\) b. \(\mathrm{Cl}_{2}(g) \rightarrow \mathrm{Cl}^{-}(a q)+\mathrm{OCl}^{-}(a q)\) c. \(\mathrm{NO}_{2}^{-}(a q)+\mathrm{Al}(s) \rightarrow \mathrm{NH}_{3}(g)+\mathrm{AlO}_{2}^{-}(a q)\)

The overall reaction and equilibrium constant value for a hydrogen-oxygen fuel cell at \(298 \mathrm{~K}\) is $$ 2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l) \quad K=1.28 \times 10^{83} $$ a. Calculate \(\mathscr{E}^{\circ}\) and \(\Delta G^{\circ}\) at \(298 \mathrm{~K}\) for the fuel cell reaction. b. Predict the signs of \(\Delta H^{\circ}\) and \(\Delta S^{\circ}\) for the fuel cell reaction. c. As temperature increases, does the maximum amount of work obtained from the fuel cell reaction increase, decrease, or remain the same? Explain.
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