An oxidation state is a concept used in chemistry to track the transfer or sharing of electrons during chemical bond formation. It represents the hypothetical charge an atom would possess if the electrons in every bond were assigned completely to the more electronegative atom. Although not a true physical charge, the oxidation state is a convenient tool for understanding chemical reactions, particularly those involving electron transfer. While many elements, such as the alkali metals in Group 1, have a fixed oxidation state (nearly always +1), a significant number of elements exhibit variable oxidation states, meaning they can adopt two or more different charges when forming compounds.
What It Means to Have Variable Oxidation States
An element with a variable oxidation state can participate in different chemical environments by losing or sharing a varying number of electrons. This flexibility allows the element to form a variety of compounds with distinct chemical and physical properties.
Iron is a common example, frequently found in two states: Iron(II) and Iron(III). In Iron(II) oxide (\(\text{FeO}\)), the iron atom has an oxidation state of +2, having formally lost two electrons. Conversely, in Iron(III) oxide (\(\text{Fe}_2\text{O}_3\)), the iron atom exhibits a +3 state, having lost three electrons.
This difference in electron loss dictates the molecular formula and structure of the resulting compound. Copper also demonstrates this variability, forming compounds with an oxidation state of +1 (like \(\text{Cu}_2\text{O}\)) and others where it is +2 (like \(\text{CuO}\)). This versatility enables these elements to act as catalysts or participate in complex biological systems, often requiring the ability to switch between states.
The Underlying Reason: Electron Configuration
The ability to exhibit multiple oxidation states results directly from an element’s electron configuration, specifically the energetic proximity of different electron sublevels. Elements with variable states possess valence electrons in more than one orbital type, such as \(s\) and \(d\) orbitals, that are energetically similar. When an atom forms a bond, the electrons lost first are usually the outermost \(s\)-electrons, as they are the easiest to remove.
In transition metals, the \(4s\) orbital is filled before the \(3d\) orbital, but the \(4s\) electrons are lost first during ionization, often resulting in the common +2 oxidation state. Since the inner \(3d\) electrons are close in energy to the \(4s\) electrons, a small additional energy input can cause one or more \(d\)-electrons to be lost as well. The sequential removal of these \(d\)-electrons leads to a series of progressively higher oxidation states.
Manganese, for example, can exhibit states from +2 up to +7, depending on how many \(4s\) and \(3d\) electrons are involved in bonding. The presence of partially filled \(d\)-orbitals provides numerous possibilities for electron loss, generating the wide range of states observed. The relative stability of these different states is influenced by configurations that result in half-filled or fully-filled \(d\)-subshells.
A different mechanism governs variable oxidation states in the heavier elements of the \(p\)-block. In elements like Tin and Lead, the outermost \(s\)-electrons are held more tightly due to the inert pair effect. This effect makes it energetically favorable for these elements to lose only their \(p\)-electrons, resulting in a lower oxidation state (e.g., +2 for Lead). A higher energy input is then required to involve the \(s\)-electrons, leading to a higher state (e.g., +4 for Lead).
Identifying the Groups on the Periodic Table
The elements that most prominently display multiple oxidation states are the Transition Metals, which occupy the \(d\)-block in the center of the periodic table. Nearly every element from Group 3 through Group 11 possesses this characteristic; Manganese exhibits the largest number of known oxidation states. This consistent variability is a hallmark of their chemistry, enabling them to form compounds with vibrant colors and function effectively as catalysts.
The Inner Transition Metals (Lanthanides and Actinides, or \(f\)-block elements) also show multiple oxidation states, though their variability is less pronounced than the transition metals. Many Lanthanides prefer the +3 state but can exhibit +2 or +4 states due to the participation of \(f\)-electrons. The Actinides show greater variability than the Lanthanides, particularly lighter elements like Uranium, due to the close energy levels of their \(5f\), \(6d\), and \(7s\) orbitals.
Variable oxidation states are also common among the heavier elements of the \(p\)-Block, specifically in Groups 13 through 16. Elements like Thallium, Tin, and Lead frequently show two distinct stable states that differ by two (e.g., +2 and +4 for Tin). This pattern contrasts sharply with the \(s\)-block metals of Groups 1 and 2, which are characterized by their singular, fixed oxidation states of +1 and +2, respectively.