Oxidation and reduction are two types of chemical reactions that go hand-in-hand to form redox reactions. Historically, oxidation was defined according to the behavior of oxygen in a reaction, but the modern definition is somewhat different. Here is the definition of oxidation, along with examples of oxidation reactions.
Key Takeaways: Oxidation
- Oxidation is the loss of electrons in a chemical reaction.
- Oxidation results in an increase in oxidation state.
- Oxidation and reduction always occur together in redox reactions.
- Oxidation does not require oxygen, although many oxidation reactions involve oxygen.
- Common examples of oxidation include rusting, combustion, and electrochemical reactions.
Oxidation Definition
Oxidation is the loss of electrons or increase in oxidation state of a molecule, atom, or ion in a chemical reaction. The opposite process is called reduction, which is a gain of electrons or the decrease in the oxidation state of a molecule, atom, or ion. In a redox reaction, one chemical species is reduced, while another is oxidized. Redox stands for reduction-oxidation.
Originally oxidation meant adding oxygen to a compound. This is because oxygen gas (O2) was the first known oxidizing agent. Adding oxygen to a compound typically meets the criteria of electron loss and an increase in the oxidation state, but the definition of oxidation was expanded to include other types of chemical reactions.
In organic chemistry, you might encounter a definition for oxidation involving hydrogen. This definition deals with protons (hydrogen ions) rather than electrons, 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.
An example is oxidation of ethanol 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.
Oxidation Examples
A classic example of oxidation occurs between iron and oxygen in moist air, forming iron oxide or rust. The iron is said to have oxidized into rust. The chemical reaction is:
4 Fe + 3 O2 + 6 H2O → 4 Fe(OH)3 or 2Fe2O3·6H2O
The iron metal is oxidized to form the iron oxide known as rust.
Fe → Fe2+ + 2 e−
Meanwhile, oxygen is reduced.
O2 + 4 e− + 2 H2O → 4 OH−
Another example of oxidation where an element combines with oxygen is the reaction between magnesium metal and oxygen to form magnesium oxide.
2 Mg (s) + O2 (g) → 2 MgO (s)
Many metals oxidize. Tarnish is the name given to the oxidation of silver. The green or blue patina that forms on copper in damp air is another example of oxidation.
But, oxidation does not require oxgen! 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 oxidized and fluorine is 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–
Electrochemical reactions are great examples of oxidation reactions. When a copper wire is placed into a solution that contains silver ions, electrons transfer 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)
Examples in Everyday Life
Oxidation reactions occur constantly in daily life. Examples include:
- Rusting: Iron reacts with oxygen and water to form hydrated iron(III) oxide.
- Combustion: Oxidizing fuels releases energy.
- Browning of fruit: Enzymatic oxidation causes cut apples and bananas to turn brown.
- Cellular respiration: Cells oxidize glucose to produce energy.
- Batteries: Controlled oxidation reactions generate chemical energy.
Oxidation is fundamental to energy production, corrosion, and many industrial processes.
How to Identify Oxidation in a Chemical Equation
Identify oxidation in a reaction by following these steps:
- Assign oxidation numbers to each element in the reactants and products.
- Compare the oxidation numbers before and after the reaction.
- Identify any element whose oxidation number increases.
An increase in oxidation number indicates oxidation.
Example
In the reaction:
2 Mg (s) + O₂ (g) → 2 MgO (s)
Magnesium has an oxidation number of 0 in Mg(s).
In MgO, magnesium has an oxidation number of +2.
Because magnesium’s oxidation number increases from 0 to +2, magnesium oxidizes.
Using OIL RIG to Remember Oxidation and Reduction
Remember, the modern definition of oxidation and reduction involves electrons (not oxygen or hydrogen). One way of remembering which species is oxidized and which is reduced is to use OIL RIG. OIL RIG stands for Oxidation Is Loss, Reduction Is Gain (of electrons).
| Feature | Oxidation | Reduction |
|---|---|---|
| Electrons | Loss of electrons | Gain of electrons |
| Oxidation State | Increases | Decreases |
| Oxygen | Often gained | Often lost |
| Hydrogen | Often lost | Often gained |
| Agent Role | Reducing agent | Oxidizing agent |
Oxidation State and Why It Matters
The oxidation state (oxidation number) of an atom represents its hypothetical charge if all bonds were completely ionic.
During oxidation:
- The oxidation state increases.
- The atom loses electron density.
Common oxidation number rules include:
- Elements in their pure form have an oxidation state of 0.
- Group 1 metals are +1 in compounds.
- Oxygen is usually −2 in compounds.
- Hydrogen is usually +1 when bonded to nonmetals.
Tracking oxidation numbers makes identifying which species is oxidized and which is reduced in a redox reaction easier.
Oxidizing and Reducing Agents
An oxidizing agent is a substance that causes another substance to be oxidized. In modern chemical terms, an oxidizing agent accepts electrons and is therefore reduced during the reaction.
A reducing agent is a substance that causes another substance to be reduced. A reducing agent donates electrons and is therefore oxidized during the reaction.
In other words:
- The oxidizing agent gains electrons.
- The reducing agent loses electrons.
For example, in the reaction:
Cu(s) + 2 Ag⁺(aq) → Cu²⁺(aq) + 2 Ag(s)
Copper loses electrons and oxidizes, so copper is the reducing agent.
Silver ions gain electrons and reduces, so silver ions are the oxidizing agent.
Historically, oxidizing agents were described as substances that added oxygen or removed hydrogen. While this is often true, the modern definition is based strictly on electron transfer.
Frequently Asked Questions
Is oxidation always reaction with oxygen?
No. Oxidation is loss of electrons, and it sometimes occurs without oxygen.
Can oxidation occur without reduction?
No. Oxidation and reduction always occur together in a redox reaction.
Is rusting the same as oxidation?
Yes. Rusting is a specific example of oxidation involving iron.
Does oxidation only involve metals?
No. Nonmetals and organic molecules also undergo oxidation.
Why does oxidation involve electron loss?
Oxidation increases oxidation state, which corresponds to a loss of electron density.
References
- 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. ISBN 978-0-8412-1780-5.
- IUPAC (1997). “Oxidation”. Compendium of Chemical Terminology (2nd ed.) (the “Gold Book”). ISBN 0-9678550-9-8. doi:10.1351/goldbook