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

First, we should assign oxidation states to each element in the reactants and products of the given reactions: a. \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)\) C: -4 to +2, H: +1 to 0, O: -2 to -2 b. \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)\) Ag: +1 to 0, N: +5 to +5, O: -2 to -2, Cu: 0 to +2 c. \(\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)\) Zn: 0 to +2, Cl: -1 to -1, H: +1 to 0 d. \(2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) H: +1 to +1, Cr: +6 to +6, O: -2 to -2

Now, let's identify which of the given reactions are redox reactions and determine their oxidizing agent, reducing agent, species being oxidized, and species being reduced: a. \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)\) Redox reaction (yes) Oxidizing agent: \(\mathrm{H}_{2}\mathrm{O}(g)\) Reducing agent: \(\mathrm{CH}_{4}(g)\) Species being oxidized: C in \(\mathrm{CH}_{4}(g)\) Species being reduced: O in \(\mathrm{H}_{2}\mathrm{O}(g)\) b. \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)\) Redox reaction (yes) Oxidizing agent: \(\mathrm{AgNO}_{3}(a q)\) Reducing agent: \(\mathrm{Cu}(s)\) Species being oxidized: \(\mathrm{Cu}(s)\) Species being reduced: \(\mathrm{Ag}^{+}(a q)\) c. \(\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)\) Redox reaction (yes) Oxidizing agent: \(\mathrm{HCl}(a q)\) Reducing agent: \(\mathrm{Zn}(s)\) Species being oxidized: \(\mathrm{Zn}(s)\) Species being reduced: \(\mathrm{H}^{+}(a q)\) d. \(2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) Redox reaction (no) There is no change in the oxidation state of any elements involved. In summary, reactions a, b, and c are redox reactions, and for each reaction, we have identified the oxidizing agent, reducing agent, species being oxidized, and species being reduced. Reaction d is not a redox reaction.