The oxidation number, often called the oxidation state, is a number assigned to an element in a compound representing the degree of oxidation or electron loss. This concept is a formalism: a hypothetical charge an atom would possess if all its bonds were completely ionic, meaning electrons were fully transferred. It is a tool used in chemistry to track the movement of electrons during oxidation-reduction (redox) processes. Unlike fixed properties like atomic mass or atomic number, the oxidation state is not listed as a single, fixed value on a standard periodic table. It is a calculated value that depends entirely on the specific compound and its bonding partner.
The Difference Between Oxidation States and Fixed Properties
The periodic table is a repository for fundamental, fixed properties of elements, such as the atomic number, which defines the element, and atomic mass. Elements also have a consistent electron configuration in their ground state. Oxidation states, however, are variable and conditional, changing based on the element’s chemical environment.
An element’s oxidation state is a measure of its electron gain or loss when it forms a chemical bond. Because the same element can bond with many different partners, it can exhibit multiple oxidation states, which is why a single number cannot be printed on the table. Some periodic tables may list a set of common oxidation states, but these are possibilities, not inherent values. The periodic table serves as a guide for predicting this variable number.
Predicting Oxidation States Using Group Location
The periodic table’s layout is structured to help predict an element’s most likely oxidation states for the main group elements. Elements in Group 1, the Alkali Metals, always have an oxidation state of +1 in their compounds. Similarly, elements in Group 2, the Alkaline Earth Metals, are consistently found with an oxidation state of +2.
For Group 17, the Halogens, the most common oxidation state is -1 when bonded to a less electronegative element. Fluorine, the most electronegative element, is nearly always found with an oxidation state of -1. Elements in Group 18, the Noble Gases, have a full outer shell and are generally unreactive, resulting in an oxidation state of 0.
Nonmetals in Groups 14 through 16 show more variability but follow a predictable pattern. When bonding with metals, they tend to achieve negative states, such as -3 for Group 15 and -2 for Group 16, by gaining electrons. However, when these nonmetals bond with more electronegative elements like oxygen or fluorine, they take on positive oxidation states.
The Variability of Transition Metals
The elements in the D-block, known as the transition metals, present an exception to the predictable rules of the main groups. These metals can exhibit multiple, variable oxidation states, often differing by only one electron. This variability prevents their oxidation numbers from being easily determined solely by their group number.
For example, Iron (Fe) commonly forms compounds in both the +2 and +3 oxidation states, while Copper (Cu) often appears as +1 or +2. This behavior is due to the small energy difference between the electrons in the outermost s-orbital and the inner d-orbital. Both sets of electrons are available to be lost during chemical reactions, allowing the atom to achieve various degrees of electron loss.
Transition metals generally lose the two electrons from their outermost s-orbital first, making +2 a common oxidation state for many of them. Subsequent higher oxidation states, such as the +7 seen in Manganese (Mn), result from the involvement of the inner d-orbital electrons. This complexity is also seen in the F-block elements (Lanthanides and Actinides), whose common oxidation state is +3, but which also exhibit other states.
Oxidation State Versus Valence Electrons
A common point of confusion is the distinction between an element’s oxidation state and its valence electrons, as both relate to bonding capacity. Valence electrons are the actual electrons present in the outermost electron shell of a neutral atom. For main group elements, the number of valence electrons is a fixed property determined by the element’s group number on the periodic table.
The oxidation state, by contrast, is the hypothetical charge assigned to the atom after it has formed a compound, assuming a complete transfer of electrons has occurred. While the number of valence electrons determines the potential for an atom to form bonds, the oxidation state is the result of how those bonds are formed. The oxidation state is a calculated formalism for electron bookkeeping, whereas the number of valence electrons is a physical characteristic of the atom itself.
Even though an element like carbon has a fixed four valence electrons, its oxidation state can range from -4 in methane (CH4) to +4 in carbon dioxide (CO2), demonstrating the variable, calculated nature of the oxidation number. The periodic table directly lists or implies the fixed number of valence electrons, but it only provides the framework used to calculate the resulting oxidation state.