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JoVE Lab Manual
Lab: Chemistry

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Oxidation and Reduction

Some chemical reactions can be classified as reduction-oxidation reactions, or redox reactions. Oxidation is the process of an atom losing one or more electrons, and reduction is the process of an atom gaining one or more electrons.

Oxidation States

Each atom in a molecule has its own oxidation state or oxidation number. The oxidation state describes how oxidized a molecule is relative to its free elemental form. The oxidation state is expressed as the charge that an atom would have if each of its bonds to other elements were purely ionic. This means that the electrons in the bond are assigned to the more electronegative atom. The oxidation state of an atom in its free elemental form is defined as 0.

There are a few rules that are followed to determine oxidation state. Elements in Group I and Group II typically have oxidation states of +1 and +2, respectively. Hydrogen and oxygen typically have oxidation states of +1 and -2, respectively, and halogens usually have an oxidation state of -1. In addition, the oxidation states of the atoms in a molecule always add up to the charge on the molecule. Thus, the oxidation state of an atom not listed above can often be deduced. For example, consider carbon dioxide (CO2), which is a neutral molecule. If each of the two oxygen molecules contributes -2, carbon’s oxidation state must be +4 to cancel out the -4 from the oxygens.

For a more general approach, draw the Lewis structure of the molecule, identify the bonds between different atoms, and assign each bond to the more electronegative atom. Then, count the number of electrons on each atom, with each bond contributing two electrons. Subtract the number of electrons that are currently on the atom from the standard number of valence electrons for that atom to get the oxidation number.

Consider carbon dioxide again. Each oxygen has two lone pairs of electrons and is connected to the central carbon by a double bond. Oxygen is more electronegative than carbon, so each C=O bond, which accounts for four electrons, is assigned to its oxygen. Thus, each oxygen is assigned a total of eight electrons (four from the lone pairs and four from the double bond), and carbon is assigned none. The default number of valence electrons for oxygen is six, so the oxidation number for each oxygen is 6 – 8 = -2. The default number of valence electrons for carbon is four, so the oxidation number for carbon is 4 – 0 = +4.

Redox Reactions

Not all chemical reactions are classified as a redox reaction. A redox reaction is any reaction in which there is a change in an atom's oxidation state. Thus, to check whether a reaction is a redox reaction, determine the oxidation states of each atom in the reactants and products and look for any changes.

Many redox reactions involve a transfer of electrons directly from one molecule to another. In those reactions, if a molecule gains an electron, another molecule must lose an electron. One simple way to remember the definitions of oxidation and reduction is through the phrase OIL-RIG, which stands for: Oxidation Is Losing – Reduction Is Gaining.

The molecule gaining an electron is being reduced, but it is called an oxidant or oxidizing agent because it is oxidizing the other molecule. Similarly, the molecule that loses an electron is being oxidized, but it is called a reductant or reducing agent because it reduces the other molecule.

There are four major reaction types that typically involve redox processes.

  1. Single Displacement Reaction: An atom displaces another atom that is part of a compound, and replaces it.
  2. Combustion Reaction: A compound is reduced by a strong oxidant, typically oxygen gas. Combustion reactions that occur between hydrocarbons and organic compounds typically produce carbon dioxide and water.
  3. Synthesis Reaction: Two reactants form a single product.
  4. Decomposition Reaction: A single reactant breaks into two or more products.


1. Harris, D. C. (2015). Quantitative Chemical Analysis. New York, NY: W. H. Freeman and Company.

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