Question

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.

Short answer

Fe(III) is more stable than Fe(II) due to a half-filled stable 3d5 subshell. In contrast, oxidizing Ni(II) to Ni(III) or Co(II) to Co(III) does not afford a half-filled or fully filled d shell, making Ni(III) and Co(III) less stable and harder to prepare.

Step-by-step solution

Understand electron configurations

The first step is to consider the electron configurations for Fe (iron), Ni (nickel), and Co (cobalt). Iron's electron configuration ends with 3d6 4s2, nickel's with 3d8 4s2, and cobalt's with 3d7 4s2. When they lose two electrons to become Fe(II), Ni(II), and Co(II), their configurations become 3d6, 3d8, and 3d7, respectively.

Recognize the stability of half-filled d subshells

We must understand that half-filled d subshells (having 5 electrons) and completely filled d subshells (having 10 electrons) are particularly stable. This is due to the symmetry and exchange energy benefits in these configurations. Iron, after losing the third electron to become Fe(III), achieves a half-filled stable 3d5 configuration.

Compare the ease of further oxidation

Fe(III) is more stable than Fe(II) because it attains the half-filled 3d5 configuration. However, oxidizing Ni(II) (3d8) to Ni(III) or Co(II) (3d7) to Co(III) disturbs their stable configurations without achieving a half-filled or fully filled d shell. This renders nickel(III) and cobalt(III) less stable and more difficult to prepare.

Conclude with the inherent stability

Because iron(III) has a half-filled 3d subshell, it has the inherent stability that the nickel(III) and cobalt(III) ions lack. As a result, nickel(II) and cobalt(II) are more resistant to oxidation than iron(II), making the creation of nickel(III) and cobalt(III) more challenging.

Key concepts

Oxidation States

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

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.

Subshell Stability

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).