How do you account for the strong reducing power of lithium in aqueous solution?

Strong reducing power of lithium in aqueous solution can be understood in terms of electrode potential. Electrode potential is a measure of the tendency of an element to lose electrons in the aqueous solution. It mainly depends upon the following three factors i.e.,
(i) $\mathrm{Li}\left(\mathrm{s}\right)\underset{\mathrm{Enthalpy}}{\overset{\mathrm{Sublimation}}{\to }}\mathrm{Li}\left(\mathrm{g}\right)$
(ii) $\mathrm{Li}\left(\mathrm{g}\right)\underset{\mathrm{Enthalpi}}{\overset{\mathrm{Ionisation}}{\to }}{\mathrm{Li}}^{+}\left(\mathrm{g}\right)+{\mathrm{e}}^{-}$
(iii)
With the small size of its ion, lithium has the highest hydration enthalpy. However, ionisation enthalpy of Li is highest among alkali metals but hydration enthalpy predominates over ionisation enthalpy.
Therefore, lithium is the strongest reducing agent in aqueous solution mainly because of its high enthalpy of hydration.