By considering electron configurations, suggest a reason why iron(III) is readily prepared from iron(II) but the conversion of nickel(II) and cobalt(II) into nickel(III) and cobalt(III) is much more difficult.

The concept of oxidation states is fundamental to the study of chemistry, particularly when examining reactions that involve a transfer of electrons. An oxidation state, often known as an oxidation number, is the charge that an atom would have if all bonds to atoms of different elements were completely ionic.

For transition metals like iron, nickel, and cobalt, multiple oxidation states are possible due to the variable manner in which electrons can be removed from the d and s orbitals. In the given exercise, iron tends to be more stable in the +3 oxidation state compared to its +2 state, while nickel and cobalt are more stable at +2.

This occurs because when transitioning to the higher oxidation states, iron(III) achieves a half-filled subshell, which is particularly stable. Nickel(III) and cobalt(III), on the other hand, do not reach this stability upon further oxidation, making their higher oxidation states more rare and difficult to achieve.

Transition metals are elements that have partially filled d or f subshells in any common oxidation state. They are known for displaying a wide range of oxidation states and unique chemistry.

The variability arises from the fact that the energy difference between their s and d orbitals is quite small, which allows electrons to be added to or removed from these orbitals quite readily. An interesting point to consider is the stability that certain electron configurations offer.

Stability Factors

Factors such as the attainment of a half-filled d subshell or completely filled d subshell confer extra stability to the transition metals in those oxidation states. For instance, iron in the +3 oxidation state (Fe(III)) has an electron configuration that gives it a half-filled d subshell (3d⁵), thereby making it particularly stable.

Electron configurations can suggest a lot about the stability of an element or ion. Within transition metals, the stability of certain subshells, specifically the d subshells, plays a significant role in the chemistry of these elements.

Subshells are considered stable when they are either half-filled or fully filled. A half-filled subshell has symmetry and minimizes repulsion between electrons, while a fully filled subshell maximizes symmetry and electron pairing. This stability is due to Hund's rule, which states that electrons will fill degenerate orbitals singly before pairing up, and maximizes the total spin. This leads to what is termed exchange energy, which contributes to the overall stability of the electron configuration.

As in the solution example, when iron changes from the +2 to the +3 oxidation state, it achieves a highly stable, half-filled d subshell, resulting in a lower energy state that nature prefers. In contrast, nickel(II) and cobalt(II) do not achieve such stable configurations upon further oxidation, which explains the difficulty in preparing nickel(III) and cobalt(III).