The $Mn^{3+}$ ion is unstable in solution and undergoes disproportionation to give $Mn^{2+}$,$MnO_2$,and $H^{+}$ ion. Write a balanced ionic equation for the reaction.

(N/A) The disproportionation reaction of $Mn^{3+}$ involves both oxidation and reduction of the same species.
$1. \text{ Oxidation Half-Reaction (O.H.R.): } Mn^{3+}_{(aq)} \rightarrow MnO_{2(s)}$
$2. \text{ Reduction Half-Reaction (R.H.R.): } Mn^{3+}_{(aq)} \rightarrow Mn^{2+}_{(aq)}$
Balancing the $O$.$H$.$R$.:
$Mn^{3+}_{(aq)} + 2H_2O_{(l)} \rightarrow MnO_{2(s)} + 4H^{+}_{(aq)} + e^-$
Balancing the $R$.$H$.$R$.:
$Mn^{3+}_{(aq)} + e^- \rightarrow Mn^{2+}_{(aq)}$
Adding both half-reactions to cancel the electrons:
$2Mn^{3+}_{(aq)} + 2H_2O_{(l)} \rightarrow MnO_{2(s)} + Mn^{2+}_{(aq)} + 4H^{+}_{(aq)}$