Why do Cu⁺ ions spontaneously form copper metal and Cu²⁺ ions in solution?

Question

The standard electrode potentials for three reactions involving copper and copper ions are:

$$ \begin{align} \ce{Cu^2+(aq) + e- &-> Cu+(aq)} &\quad E^\circ &= \pu{+0.15 V} \\ \ce{Cu^2+(aq) + 2e- &-> Cu(s)} &\quad E^\circ &= \pu{+0.34 V} \\ \ce{Cu+(aq) + e- &-> Cu(s)} &\quad E^\circ &= \pu{+0.52 V} \\ \end{align} $$

Which statement is correct?

A. $\ce{Cu^2+}$ ions are a better oxidizing agent than $\ce{Cu+}$ ions.
B. Copper metal is a better reducing agent than $\ce{Cu+}$ ions.
C. $\ce{Cu+}$ ions will spontaneously form copper metal and $\ce{Cu^2+}$ ions in solution.
D. Copper metal can be spontaneously oxidized by $\ce{Cu^2+}$ ions to form $\ce{Cu+}$ ions.

I thought that because copper is more likely to get oxidised than $\ce{Cu+},$ it is a better reducing agent and thus B is the answer. But the given answer is C. Why?

See related https://chemistry.stackexchange.com/questions/115660/disproportionation-of-2-ko%e2%82%82-2-h%e2%82%82o-%e2%86%92-2-koh-h%e2%82%82o%e2%82%82-o%e2%82%82/115662#comment218026_115662 — Poutnik, May 21 '19 at 10:56
@Poutnik - thanks, yes I did. But that does not explain why in terms of the standard electrode potentials involved. I'm having a hard time understanding how you could infer that from the SEPs given. Do you have any idea? Thanks in advance.! — Mayuri Vaish, May 21 '19 at 11:02
Did you compare standard potentials of all 3 half reactions ? — Poutnik, May 21 '19 at 11:04
Strongly related; https://chemistry.stackexchange.com/questions/42855/why-cu-is-unstable-in-aqueous-medium — Nilay Ghosh, May 21 '19 at 11:23
Take it as if Cu(I) had been 2 different compounds in one. One as oxidizing agent, the other as reduction agent. Compare potentials for their half reaction. — Poutnik, May 21 '19 at 11:38
@Poutnik I got it!!! Thank you!!! Essentially, both reactions involving Cu+ have the lowest and highest SEPs respectively - so they have to be the strongest oxidizing/reducing agents - thus undergoing disproportionation. :))) — Mayuri Vaish, May 21 '19 at 12:09

score 3 · Answer 1 · edited Jun 02 '22 at 03:25

Both reactions involving $\ce{Cu+}$ have the lowest and highest standard electrode potentials, respectively. Thus, $\ce{Cu+}$ has to be the strongest oxidizing/reducing agents, and therefore it undergoes disproportionation.

Because the formation of $\ce{Cu}$ has a higher SEP than $\ce{Cu+},$ it is more likely to get reduced and therefore is a better oxidizing agent — not a better reducing agent — than $\ce{Cu+}$ (which is more likely to get oxidized as it has a higher SEP).

Why do Cu⁺ ions spontaneously form copper metal and Cu²⁺ ions in solution?

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