What happens to the electrons in an oxidation reaction?

In this tutorial, you will learn what a redox reaction is, the different parts of such a reaction, as well as how to recognize and write redox reactions. You will also learn the difference between oxidation and reduction, and the definition of oxidation.

Topics Covered in Other Articles

Vocabulary

• Oxidation: a type of chemical reaction where one or more electrons are lost.
• Oxidation State / Number: a number assigned to an atom describing its degree of oxidation, meaning how many electrons it has gained or lost.
• Reduction: a type of chemical reaction where one or more electrons are gained.
• Oxidation Reduction Reaction: a chemical reaction where oxidation and reduction occurs simultaneously

What are redox reactions?

Redox is a shorthand for reduction-oxidation, meaning that a redox reaction is one in which both a reduction reaction and an oxidation reaction takes place at once. It is also shorthand for oxidation reduction reaction. Let’s discuss these two components separately, then circle back to how they combine in a full redox reaction.

Reduction:

Reduction happens when an atom gains one or more electrons during a chemical reaction. That means that its oxidation number decreases. This is because an electron has a negative charge, thus when an atom gains an electron, it gains a negative charge, which decreases the oxidation number. This could look like an atom going from X2+ to X1+, or X0 to X1-, for example. This can help us remember what reduction is, as the oxidation number gets reduced.

Note: the species getting reduced is called the oxidizing agent. This may be a little confusing at first, but remember the oxidizing agent is the species that causes another to become oxidized—since redox reactions happen in pairs of reduction and oxidation, the species causing oxidation is the one getting reduced.

Let’s look at some examples of reduction half-reactions. These are called half-reactions because they make up half of a full redox reaction:

Cu2+ (aq) → Cu (s)

F2 (g) → 2F– (g)

Oxidation Definition:

There are three different oxidation definitions. “The process or result of oxidizing or being oxidized”, “the gain of oxygen, such as when an element combines with oxygen to form its oxide, like rusting”, and lastly “the loss of electrons during a reaction, which results an atom increasing its oxidation state”. The last definition is the one you should remember for chemistry.

Oxidation happens when an atom loses one or more electrons during a chemical reaction, meaning that its oxidation number increases. This is because the atom loses the negative charge of the electron, which is similar to gaining a positive charge, increasing the oxidation number. This could look like an atom going from X1- to X0, or X0 to X1+, for example. The phrase “oxidation” is used because historically, the first redox reactions observed were ones involving oxygen. If it helps, however, we can think of it as oxidation because the oxidation number increases/becomes more positive.

When bleach or hydrogen peroxide turns something white, we say that the substance is oxidized. The hypochlorite ion in the bleach, or the peroxide ion, gains electrons, and the substance that turns white (by default the reducing agent) loses electrons.

Note: The reducing agent is the species getting oxidized. This is because the reducing agent is the species causing another to become reduced. In a redox reaction, the species getting oxidized causes the other species to get reduced, through the exchange of electrons.

Let’s look at some examples of oxidation half-reactions:

2I– (aq) → I2 (s)

Zn (s) → Zn2+ (aq)

Oxidation vs Reduction

Remember OIL RIG: Oxidation Is Loss of electrons – Reduction Is Gain of electrons. Alternatively, you can also remember “LEO GER” (imagine a lion saying “Ger!!!!). LEO means “Loose electrons, oxidation”. GER means “Gain electrons, reduction”.

Redox reactions: involve both

Redox reactions put the two half-reactions (one reduction, one oxidation) together into a complete equation. The electrons lost in the oxidation half-reaction are then gained in the reduction half-reaction. In this manner, a redox reaction is a chemical reaction wherein electrons are transferred between two species.

Oxidation Examples

Rust is the classic example of oxidation. Iron metal is oxidized to brown iron (III) oxide. Other oxidation examples include zinc metal displacing silver or copper in solution, the zinc is oxidized to the zinc (II) ion. And burning anything in oxygen, for example sugar or magnesium, is an oxidation reaction. Bleaching something, like hair, blue jeans, or glowing tonic water, is also an oxidation-reduction reaction – the sodium hypochlorite in bleach is the oxidizer.

Examples of Redox Reactions

Let’s take a look at some examples of full redox reactions

CO2 (g) + H2 (g) → CO (g) + H2O (g)

CH4 (g) + O2 (g) → CO2 (g) + H2O (g)

Zn + CuCl2 → ZnCl2 + Zn

Na + Cl → NaCl

How do you recognize redox reactions?

The easiest and primary way to recognize a redox reaction is by look for changes in the oxidation states of atoms from the reactants to the products. One species should have its oxidation number decrease from reactants to products (this, again, is reduction). Another species in the same reaction should have its oxidation number increased from reactants to products (this, again, is oxidation). Just remember – the number of electrons lost should be equal to the number of electrons gained.

Reminder: an atom’s oxidation number (also sometimes called the oxidation state) is a measure of how many electrons it has gained or lost. That is, an oxidation number of zero indicates a neutral atom. Similarly, gaining electrons decreases the oxidation number, since electrons are negative and thus add a negative charge. Losing electrons increases the oxidation number, since the negative charge is being lost. Read more about oxidation states here!

Common types of redox reactions

Aside from using oxidation states to recognize redox reactions, there are other clues to when a redox reaction is happening; namely, there are a few different types of reactions where reduction and oxidation often both take place.

• Combination reactions: these reactions combine elements to create a compound, generally taking the form of A + B → AB
• Decomposition reactions: these reactions are the reverse of combination reactions; a compound decomposes into its elemental parts. These generally take the form of AB → A + B
• Combustion reactions: these reactions involve a fuel (usually organic) and oxygen as reactants, and result in water and carbon dioxide, and occasionally another organic product such as nitrogen. The general form is X + O2 → CO2 + H2O. This equation, of course, needs to balanced according to what X, the organic fuel, is.
• Displacement reactions with an oxidation state change. These reactions can either come in the form of single replacement, where one element in a compound is replaced by another one, or in the form of a double replacement, where an element in each of the two different reactants gets replaced. Effectively, swapped. The general form of a single replacement reaction is A + BC → B + AC. The general form of a double replacement reaction is AB + CD → AD + CB.

See our tutorial on balancing redox reactions for more information and further examples!

These are usually NOT redox reactions:

• Acid-base reactions
• Simple precipitation reactions
• Double displacement reactions with no oxidation state change

Video of a simple redox reaction

Let’s look at an oxidation reduction reaction. In this video of the elephant toothpaste reaction, manganese in potassium permanganate is reduced from the +7 to the +4 and +2 oxidation states, while the oxygen in hydrogen peroxide is oxidized from the -1 to the 0 oxidation state (to elemental oxygen).

Oxidation & Reduction – Further Reading

Two key types of chemical reactions are oxidation and reduction. Oxidation doesn't necessarily have anything to do with oxygen. Here's what it means and how it relates to reduction.

• Oxidation occurs when an atom, molecule, or ion loses one or more electrons in a chemical reaction.
• When oxidation occurs, the oxidation state of the chemical species increases.
• Oxidation doesn't necessarily involve oxygen! Originally, the term was used when oxygen caused electron loss in a reaction. The modern definition is more general.

Oxidation is the loss of electrons during a reaction by a molecule, atom or ion.
Oxidation occurs when the oxidation state of a molecule, atom or ion is increased. The opposite process is called reduction, which occurs when there is a gain of electrons or the oxidation state of an atom, molecule, or ion decreases.

An example of a reaction is that between hydrogen and fluorine gas to form hydrofluoric acid:

H2 + F2 → 2 HF

In this reaction, hydrogen is being oxidized and fluorine is being reduced. The reaction may be better understood if it is written in terms of two half-reactions.

H2 → 2 H+ + 2 e-

F2 + 2 e- → 2 F-

Note there is no oxygen anywhere in this reaction!

An older meaning of oxidation was when oxygen was added to a compound. This was because oxygen gas (O2) was the first known oxidizing agent. While the addition of oxygen to a compound typically meets the criteria of electron loss and an increase in the oxidation state, the definition of oxidation was expanded to include other types of chemical reactions.

A classic example of the old definition of oxidation is when iron combines with oxygen to form iron oxide or rust. The iron is said to have oxidized into rust. The chemical reaction is:

2 Fe + O2 → Fe2O3

The iron metal is oxidized to form the iron oxide known as rust.

Electrochemical reactions are great examples of oxidation reactions. When a copper wire is placed into a solution that contains silver ions, electrons are transferred from the copper metal to the silver ions. The copper metal is oxidized. Silver metal whiskers grow onto the copper wire, while copper ions are released into the solution.

Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)

Another example of oxidation where an element combines with oxygen is the reaction between magnesium metal and oxygen to form magnesium oxide. Many metals oxidize, so it's useful to recognize the form of the equation:

2 Mg (s) + O2 (g) → 2 MgO (s)

Once the electron was discovered and chemical reactions could be explained, scientists realized oxidation and reduction occur together, with one species losing electrons (oxidized) and another gaining electrons (reduced). A type of chemical reaction in which oxidation and reduction occurs is called a redox reaction, which stands for reduction-oxidation.

The oxidation of a metal by oxygen gas could then be explained as the metal atom losing electrons to form the cation (being oxidized) with the oxygen molecule gaining electrons to form oxygen anions. In the case of magnesium, for example, the reaction could be rewritten as:

2 Mg + O2 → 2 [Mg2+][O2-]

comprised of the following half-reactions:

Mg → Mg2+ + 2 e-

O2 + 4 e- → 2 O2-

Oxidation in which oxygen is involved is still oxidation according to the modern definition of the term. However, there is another old definition involving hydrogen which may be encountered in organic chemistry texts. This definition is the opposite of the oxygen definition, so it may cause confusion. Still, it's good to be aware. According to this definition, oxidation is the loss of hydrogen, while reduction is the gain of hydrogen.

For example, according to this definition, when ethanol is oxidized into ethanal:

CH3CH2OH → CH3CHO

Ethanol is considered oxidized because it loses hydrogen. Reversing the equation, ethanal can be reduced by adding hydrogen to it to form ethanol.

So, remember the modern definition of oxidation and reduction concern electrons (not oxygen or hydrogen). One way to remember which species is oxidized and which is reduced is to use OIL RIG. OIL RIG stands for Oxidation Is Loss, Reduction Is Gain.

• Haustein, Catherine Hinga (2014). K. Lee Lerner and Brenda Wilmoth Lerner (eds.). Oxidation–Reduction Reaction. The Gale Encyclopedia of Science (5th ed.). Farmington Hills, MI: Gale Group.
• Hudlický, Miloš (1990). Oxidations in Organic Chemistry. Washington, D.C.: American Chemical Society. p. 456. ISBN 978-0-8412-1780-5.