One of the notable features of the transition element is the great variety of oxidation states it may show in its compounds. The elements which give the greatest number of oxidation states occur in or near the middle of the series. Manganese, for example exhibits all the oxidation states from +2 to +7. The lesser number of oxidation states at the extreme ends is either due to the loss or sharing of a few electrons (Sc, Ti) or too many d electrons of higher valence. Early in the first series scandium (II) is virtually unknown and titanium (IV) which is more stable than Ti (III) or Ti (II). On the other end, the only oxidation state of zinc is +2.

Fe shows extra stability as \(Fe^{3+}\) due to

\(d^5\)

The extra stability of \(Fe^{3+}\) is due to its \(d^{5}\) configuration. This is because \(d^{5}\) is a half-filled subshell, which is particularly stable due to Hund's rule. Hund's rule states that when orbitals of equal energy are available, electrons will fill them so that each orbital has the maximum number of unpaired electrons allowed by Pauli's exclusion principle. In the case of \(Fe^{3+}\), the \(d^{5}\) configuration has three unpaired electrons, which is the maximum number of unpaired electrons allowed for a (d) subshell. This makes the \(d^{5}\) configuration particularly stable.

The other options are not correct. \(d^{10}\) is a full subshell, which is also stable, but it is not as stable as a half-filled subshell. \(s^0\) and \(s^2\) are not stable configurations for iron, because they would require iron to lose or gain electrons, which would be a very difficult process.

So the answer is 2. \(d^{5}\).