Oxidation-Reduction Reactions — Definition & Examples

(Video) Redox Reactions

by Bozeman Science

This video provides an introduction to oxidation-reduction reactions. It begins with an overview of the problem solving strategy. The video goes over the rules for assigning oxidation numbers. Next, the video goes over several example reactions. The video ends with a discussion of real world applications of oxidation-reduction reactions.

What Are Redox Reactions?

An oxidation-reduction reaction, shortened as a redox reaction, is any chemical reaction in which electrons are transferred from one participant to another. In order for this to happen, one compound must lose electrons, which the other then gains. Therefore, we can think of oxidation-reduction reactions as the sum of two half reactions:

An oxidation reaction is any reaction where the reactant loses an electron.

For example, when Na becomes Na+, sodium loses an electron and is oxidized.

An reduction reaction is any reaction where the reactant gains an electron. As electrons have a negative charge, you can remember the fact that reduction is a gain because the charge on the species will be lowered or reduced.

For example, when Cl becomes Cl-, chlorine gains an electron, resulting in a lower charge, or reduction.

In order for chlorine to gain that electron, it needs to be taken from something else. Therefore, an oxidation half reaction and a reduction half reaction must always be combined to create the chemical equation for an oxidation-reduction reaction. In this case, Na + Cl → Na+ + Cl- becomes our overall redox reaction. As NaCl is an ionic compound, NaCl means the same thing as Na+ and Cl- individually, so this reaction can also be written as Na + Cl → NaCl

What are oxidation numbers?

In the above example, when Na and Cl react to form NaCl, Na is oxidized to Na+ and Cl is reduced to Cl-. As Na and Cl are both in their original forms as elements, they both begin with a charge of zero. The amount of electrons in Na relative to the normal charge of sodium then changes from 0 to +1, while the amount of electrons in Cl relative to the normal charge of chlorine changes from 0 to -1.

This amount of electrons relative to an atom's ordinary state is called an oxidation number. All elemental forms, or individual atoms and naturally occurring diatomic forms such as O2, H2, and N2, have an oxidation number of 0.

Since a change in oxidation number means a change in the amount of electrons, redox reactions can also be defined as any reaction in which at least one participating atom's oxidation number increases, and at least one's decreases.

Atoms do not need to become ions for their oxidation number to change. When a covalent bond is formed, the more electronegative atom's oxidation number is decreased by 1, while the other atom's oxidation number is increased by 1. Oxygen is the exception to this rule, and has an oxidation number of 2 in any covalent bond. Any neutrally charged molecule's overall oxidation charges will add to 0.

For example, the oxidation number on C in pure C is 0, as it is in its elemental form, but the oxidation number on C in CH4, or carbon with four hydrogens bonded around it, is +4, while the oxidation number of each H is -1. In CO2, the oxidation number on each O is -2, meaning the oxidation number on C is again +4.

Therefore, the reaction 2H2 + C → CH4 is also a redox reaction, where C is oxidized (as its oxidation number changes from 0 to +4) and H is reduced (as its oxidation number changes from 0 to -1.)