# Why does chromium metal not react with a solution of Mg^2+ even though their standard cell potentials produce a net positive voltages?

Write net equations for the spontaneous redox reactions that occur during the following or NONE if there is no extensive reaction. Use the lowest possible coefficients. Include states-of-matter under the given conditions in your answer.

## My effort:

Chromium metal is added to a solution of $$\ce{Mg^2+}$$

\begin{align} \ce{2Cr^3+ (aq) + 6e- &-> 2Cr(s)}\\ \ce{3Mg(s) &-> 3Mg^2+ (aq) + 6e-} \end{align}

The net equation is $$\ce{2Cr^3+ (aq) + 3Mg(s) -> 2Cr(s) + 3Mg^2+ (aq)}$$

Standard reduction potential of $$\ce{Cr}$$ is $$\pu{-0.74 V}$$ and the standard reduction potential of $$\ce{Mg}$$ is $$\pu{-2.36 V}$$.

$$\ce{Cr}$$ reaction is cathode reaction and therefore it is $$\ce{Cr^3+}$$ getting reduced, $$\ce{Mg}$$ is anode reaction and therefore it is $$\ce{Mg(s) -> Mg^2+ (aq) + 2e^-}.$$

$$E^\circ_\mathrm{cell} = \pu{-0.74 V} - (\pu{-2.36 V}) = \pu{1.62 V}$$

Yet apparently my homework here says that that is wrong. Does anyone see anything wrong with how I am going about this problem?

• It's the magnesium ion that is being converted into magnesium metal. $\ce{Mg^{2+}}$ is being reduced at the cathode. Apr 10, 2016 at 15:17
• In other words, all your equations are written backwards. Mar 14, 2019 at 15:24
• Metals that could theoretically reduce magnesium ions would rather reduce water to hydrogen. Feb 16, 2022 at 17:07