Oxidation Number

The Concept of Oxidation Number

The idea of oxidation number was developed as a practical bookkeeping tool to track electron shifts in reactions, especially those involving covalent compounds. Consider the formation of water:

In this reaction, hydrogen goes from a neutral state in $\text{H}_2$ to a positive state in $\text{H}_2\text{O}$, while oxygen goes from neutral to a dinegative state. We can think of this as an electron transfer from hydrogen to oxygen — but the transfer is only partial. It is better described as an electron shift rather than a complete loss and gain. The same idea applies to other covalent reactions:

To keep track of these shifts systematically, we use oxidation numbers. The method assumes, for bookkeeping purposes only, that there is a complete transfer of electrons from the less electronegative atom to the more electronegative atom. This allows us to rewrite the above reactions with charges shown:

Oxidation number denotes the oxidation state of an element in a compound. It is ascertained according to a set of rules formulated on the basis that the electron pair in a covalent bond belongs entirely to the more electronegative element. When a molecule or ion contains two or more atoms of the same element (as in $\text{Na}_2\text{S}_2\text{O}_3$ or $\text{Cr}_2\text{O}_7^{2-}$), the oxidation number of that element is taken as the average of the oxidation numbers of all its atoms.

Rules for Determining Oxidation Number

Rule 1: Free Elements

In elements in the free or uncombined state, each atom bears an oxidation number of zero.

Thus, each atom in $\text{H}_2$, $\text{O}_2$, $\text{Cl}_2$, $\text{O}_3$, $\text{P}_4$, $\text{S}_8$, $\text{Na}$, $\text{Mg}$, and $\text{Al}$ has an oxidation number of zero.

Rule 2: Monatomic Ions

For ions composed of only one atom, the oxidation number is equal to the charge on the ion.

In their compounds, all alkali metals have oxidation number $+1$, all alkaline earth metals have $+2$, and aluminium has $+3$ in all its compounds.

Rule 3: Oxygen

The oxidation number of oxygen in most compounds is $-2$. There are two kinds of exceptions:

First exception — Peroxides and superoxides: When oxygen atoms are directly linked to each other:

• In peroxides (e.g., $\text{H}_2\text{O}_2$, $\text{Na}_2\text{O}_2$), each oxygen atom is assigned an oxidation number of $-1$.
• In superoxides (e.g., $\text{KO}_2$, $\text{RbO}_2$), each oxygen atom is assigned an oxidation number of $-\frac{1}{2}$.

Second exception — Oxygen bonded to fluorine: In compounds like oxygen difluoride ($\text{OF}_2$) and dioxygen difluoride ($\text{O}_2\text{F}_2$), oxygen is assigned oxidation numbers of $+2$ and $+1$, respectively. The number assigned depends on the bonding state, but it is a positive figure.

Rule 4: Hydrogen

The oxidation number of hydrogen is $+1$, except when it is bonded to metals in binary compounds (compounds containing two elements). For example, in $\text{LiH}$, $\text{NaH}$, and $\text{CaH}_2$, its oxidation number is $-1$.

Rule 5: Halogens

Fluorine has an oxidation number of $-1$ in all its compounds. The other halogens (Cl, Br, I) also have an oxidation number of $-1$ when they occur as halide ions in their compounds. However, chlorine, bromine, and iodine have positive oxidation numbers when combined with oxygen — for example, in oxoacids and oxoanions.

Rule 6: Sum of Oxidation Numbers

The algebraic sum of the oxidation numbers of all the atoms in a compound must be zero. In a polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion.

For example, in the carbonate ion $(\text{CO}_3)^{2-}$, the sum of the oxidation numbers of three oxygen atoms and one carbon atom must equal $-2$.

Applying the Rules

By applying the above rules, we can find the oxidation number of any desired element in a molecule or ion. Metallic elements have positive oxidation numbers, while nonmetallic elements can have positive or negative oxidation numbers. Transition elements usually display several positive oxidation states.

The highest oxidation number of a representative element follows a pattern based on the periodic table:

• For groups 1 and 2: the highest oxidation number equals the group number.
• For groups 13–17: the highest oxidation number equals the group number minus 10 (following the long form of the periodic table).

This means the highest oxidation number generally increases across a period. For the third period:

Stock Notation

The oxidation number or state of a metal in a compound is sometimes presented using Stock notation, developed by the German chemist Alfred Stock. In this notation, the oxidation number is expressed by putting a Roman numeral in parenthesis after the symbol of the metal in the molecular formula.

This change in oxidation number implies a change in oxidation state, which helps identify whether a species is in its oxidised or reduced form. For example, Hg$_2$(I)Cl$_2$ is the reduced form of Hg(II)Cl$_2$. …