2.1.5 - Redox

Oxidative State

Oxidative state, or an elements oxidation number, is an artificial value applied to elements in a compound that aids with many calculations and overall understanding. Often, it is interchangeable with charge, but for polyatomic ions such as the manganate(VII) ion (MnO$\small_4^-$), oxidation number differs from charge, since it is assigned to manganese and oxygen separately.

Strictly, oxidation number can be thought of as the total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. If an atom ‘gains’ electrons, its number is negative, and if it ‘loses’ electrons then its number is positive.

Calculating oxidation numbers can be done by using the following rules all together.


  1. The oxidation number of an uncombined element is 0.
  2. Certain elements have fixed oxidation numbers. All group 1 elements are +1. All group 2 elements are +2. Hydrogen is always +1 except in hydrides (NaH, LiH, etc.), where it is -1. Fluorine is always –1. Oxygen is always –2 except in peroxides, where it is -1; superoxides, where it is -$\frac{1}{2}$; and when combined with fluorine, where it is -1. Chlorine is always –1 except when combined with fluorine and/or oxygen.
  3. The sum of oxidation numbers in a compound is always 0.
  4. The sum of oxidation numbers in an ion always adds up to the charge on the ion.

Often, the name of a compound will contain the oxidative state of one element. This is because its state is not fixed, and may vary depending on the compound it is in. For example, there are two chlorate ions, ClO$\small^-$ and ClO$\small_3^-$. In the former, chlorine has oxidation number +1, and in the latter, chlorine has oxidation number +5. Since otherwise the two ions would have the same name, they are distinguished by adding the oxidation number of the variable atom’s (chlorine) state in roman numerals in brackets:

  • chlorate(I) ion, ClO$\small^-$
  • chlorate(VI) ion, ClO$\small_3^-$

Redox Reactions

In most reactions, the oxidative state of atoms can change due to the exchange of electrons between them. These reactions are called redox reactions.

The word ‘redox’ is an amalgamation of the words reduction and oxidation, where reduction is the gain of electrons and decrease in the oxidative state; and oxidation is the loss of electrons and increase in the oxidative state. A redox reaction is one in which both reduction and oxidation occur.

An example of a redox reaction is the one between iron and hydrochloric acid, which is also an example of a metal and acid reacting to form a salt and hydrogen gas:

$$ \text{Fe} _\text{(s)} + \text{2HCl} _\text{(aq)} \rightarrow \text{FeCl$_2$} _\text{(aq)} + \text{H$_2$} _\text{(g)} $$

This table shows the oxidative states of each element before and after the reaction.

Element Oxidative state before Oxidative state after Change in Oxidative state Reduction or Oxidation?
Fe 0 +2 +2 Oxidation
H +1 0 -1 Reduction
Cl -1 -1 0 Neither

This shows that the reaction is indeed a redox reaction.


Another, special, example of a redox reaction is the formation of bleach, a reaction between chlorine gas and sodium hydroxide:

$$ \text{Cl$_2$} _\text{(g)} + \text{2NaOH} _\text{(aq)} \rightarrow \text{NaClO} _\text{(aq)} \text{NaCl} _\text{(aq)} + \text{H$_2$O} _\text{(l)} $$

where the only species (another term referring to either atom, group or compound) that changes is chlorine. It starts off with oxidation number 0, but ends up with +1 in sodium chlorate(I) and -1 in sodium chloride.

This type of redox, in which the same species is both oxidised and reduced is known as a disproportionation reaction.