Question

Why is it easier to oxidise `Fe^(2+)` to `Fe^(3+)` than it is to oxidise `Mn^(2+)` to `Mn^(3+)`?

Answer 1

Well, consider the NEUTRAL electron configurations:

`"Fe": [Ar] 3d^6 4s^2`
`"Mn": [Ar] 3d^5 4s^2`

The `4s` orbital is higher in energy in these atoms, so it is ionized first:

`"Fe"^(2+): [Ar] 3d^6`
`"Mn"^(2+): [Ar] 3d^5`

Drawn out:

`"Fe"^(2+): ul(uarr darr)" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))`

`"Mn"^(2+): ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))" "ul(uarr color(white)(darr))`

A single oxidation is the act of singly ionizing:

`"M"^(2+) -> "M"^(3+) + e^(-)`

The electron to be removed from `Fe^(2+)` is paired, having charge repulsions (making it easier to remove, i.e. the ionization energy is smaller).

Therefore, it is easier to ionize `"Fe"^(2+)` than `"Mn"^(2+)`.