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. $2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)$

For each equation, assign oxidation numbers to the elements. If the oxidation numbers change from reactants to products, the equation is a redox reaction. a. \(CH_4(g)+H_2O(g) \rightarrow CO(g)+3H_2(g)\) - Assign oxidation numbers: C: -4 to +2, H: +1 to 0, O: -2 - Oxidation numbers change, so this is a redox reaction. b. \(2AgNO_3(aq)+Cu(s) \rightarrow Cu(NO_3)_2(aq)+2Ag(s)\) - Assign oxidation numbers: Ag: +1 to 0, N: +5, O: -2, Cu: 0 to +2 - Oxidation numbers change, so this is a redox reaction. c. \(Zn(s)+2HCl(aq) \rightarrow ZnCl_2(aq)+H_2(g)\) - Assign oxidation numbers: Zn: 0 to +2, H: +1 to 0, Cl: -1 - Oxidation numbers change, so this is a redox reaction. d. \(2H^+(aq)+2CrO_4^{2-}(aq) \rightarrow Cr_2O_7^{2-}(aq)+H_2O(l)\) - Assign oxidation numbers: H: +1, Cr: +6 to +6, O: -2 - Oxidation numbers do not change, so this is not a redox reaction.

For each redox reaction, identify the species that cause the oxidation and reduction by determining which species gained or lost electrons. a. \(CH_4(g)+H_2O(g) \rightarrow CO(g)+3H_2(g)\) - Oxidation: C is oxidized from -4 to +2 (loses electrons) - Reduction: H is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Water - Reducing agent: Methane b. \(2AgNO_3(aq)+Cu(s) \rightarrow Cu(NO_3)_2(aq)+2Ag(s)\) - Oxidation: Cu is oxidized from 0 to +2 (loses electrons) - Reduction: Ag is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Silver nitrate - Reducing agent: Copper c. \(Zn(s)+2HCl(aq) \rightarrow ZnCl_2(aq)+H_2(g)\) - Oxidation: Zn is oxidized from 0 to +2 (loses electrons) - Reduction: H is reduced from +1 to 0 (gains electrons) - Oxidizing agent: Hydrochloric acid - Reducing agent: Zinc

The redox reactions are a, b, and c. The oxidation and reduction species, along with their corresponding agents, are as follows: a. Oxidizing agent: Water; Reducing agent: Methane b. Oxidizing agent: Silver nitrate; Reducing agent: Copper c. Oxidizing agent: Hydrochloric acid; Reducing agent: Zinc