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ChemWiki: The Dynamic Chemistry E-textbook > Analytical Chemistry > Electrochemistry > Redox Chemistry > Oxidation-Reduction Reactions

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Oxidation-Reduction Reactions

Oxidation-reduction reactions (or redox) reactions, are a type of chemical reaction that involves a transfer of electrons between two species. An oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion changes by gaining or losing an e-. They occur every day and are vital to some of the basic functions of life. Some examples include photosynthesis, respiration, combustion, and corrosion or rusting. To understand these types of reactions, you must first understand oxidation numbers or states.

Rules for Assigning Oxidation States

The Oxidation State (OS) corresponds to the number of electrons, e-, that an atom loses, gains, or appears to use when joining with other atoms in compounds. When determining the OS of an atom there are seven guidelines to follow:

  1. The OS of an individual atom is 0.
  2. The total OS of all atoms in: a neutral species is 0 and in an ion is equal to the ion charge.
  3. Group 1 metals have an OS of +1 and group 2 an OS of +2
  4. The OS of fluorine is -1, when in compounds
  5. Hydrogen generally has an OS of +1 in compounds
  6. Oxygen generally has an OS of -2 in compounds
  7. In binary metal compounds, group 17 elements have an OS of -1, group 16 of -2, and group 15 of -3.

(Note: The sum of the OSs is equal to zero for neutral compounds and equal to the charge for polyatomic ion species.)


Example 1: Assigning OSs

Determine the OSs of the elements in the following reactions:

  1. Fe(s) + O2(g) → Fe2O3(g)
  2. Fe2+
  3. Ag(s) + H2S → Ag2S(g) + H2(g)



  1. Fe and O2 are free elements, therefore they have an OS of "0" according to Rule #1. The product has a total OS equal to "0" and following Rule #6, O has an OS of -2, which means Fe has an OS of +3.
  2. The OS of Fe corresponds to its charge, therefore the OS is +2.
  3. Ag has an OS of 0, H has an OS of +1 according to Rule #5, S has an OS of -2 according to Rule #7 and hence Ag in Ag2S has an OS of +1.
Example 2: Assigning OSs

Determine the OS of the bold element in each of the following:

  1. Na3PO3
  2. H2PO4-



  1. The oxidation numbers of Na and O are +1 and -2. Since sodium phosphite is neutral, the sum of the oxidation numbers must be zero. Letting x be the oxidation number of phosphorus then, 0= 3(+1) + x + 3(-2). x=oxidation number of P= +3. 
  2. Hydrogen and oxygen have oxidation numbers of +1 and -2. The ion has a charge of -1, so the sum of the oxidation numbers must be -1. Letting y be the oxidation number of phosphorus, -1= y + 2(+1) +4(-2), y= oxidation number of P= +5.
Example 3: Identifying Reduced and Oxidized Elements

Determine which element is oxidized and which element is reduced in the following reactions (be sure to include the OS of each): 

  1. Zn + 2H+ → Zn2+ + H2
  2. 2Al + 3Cu2+→2Al3+ +3Cu
  3. CO32+ 2H+→ CO2 + H2O



  1. Zn is oxidized (Oxidation number: 0 → +2); H+ is reduced (Oxidation number: +1 → 0)
  2. Al is oxidized (Oxidation number: 0 → +3); Cu2+ is reduced (+2 → 0)
  3. This is not a redox type because each element has the same oxidation number in both reactants and products: O= -2, H= +1, C= +4.

(For a more in depth look see oxidation numbers).

An atom is oxidized when it oxidation number increases, the reducing agent, and an atom is reduced when its oxidation number decreases, the oxidizing agent. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound).

Oxidation-Reduction Reactions

Redox reactions are comprised of two parts, a reduced half and an oxidized half, that always occur together. The reduced half gains electrons and the oxidation number decreases, while the oxidized half losses electrons and the oxidation number increases. Simple ways to remember this are the mnemonic devices OIL RIG meaning "oxidation is loss" and "reduction is gain" or LEO says GER meaning "loss of e- = oxidation" and "gain of e- = reduced." There is no net change in the number of electrons in a redox reaction. Those given off in the oxidation half reaction are taken on by another species in the reduction half reaction.

The two species that exchange electrons in a redox reaction are given special names. The ion or molecule that accepts electrons is called the oxidizing agent; by accepting electrons it brings about the oxidation of another species. Conversely, the species that donates electrons is called the reducing agent; when reaction occurs it reduces the other species. In other words, what is oxidized is the reducing agent and what is reduced is the oxidizing agent. (Note: the oxidizing and reducing agents can be the same element or compound This will be further discussed under Types of Redox Reactions: Disproportionation).

A good example of a redox reaction is the thermite reaction in which iron atoms of ferric oxide lose (or give up) O atoms to Al atoms, producing Al2O3.

Fe2O3(s) + 2Al(s) → Al2O3(s) + 2Fe(l)

Another example of the redox reaction is the reaction between Zinc and Copper sulfate.


Example 4: Identifying Oxidized Elements

Using the equations from the previous examples determine what is oxidized?

Zn + 2H+ → Zn2+ + H2


The OS of H goes from +1 to 0 and the OS of Zn goes from 0 to 2+. Hence, Zn is oxidized and acts as the reducing agent.

Example 5: Identifying Reduced Elements

What is reduced species in this reaction?

Zn +​ 2H+ → Zn2+ + H2


The OS of H goes from +1 to 0 and the OS of Zn goes from 0 to 2+. Hence, H+ ion is reduced and acts as the oxidizing agent.

Types of Redox Reactions

Combination reactions are some of the simplest redox reactions and as the name suggests involves the "combining" of elements to form a chemical compound. As usual, oxidation and reduction occur together. General Equation:

A + B → AB

Example 6: Combination Reaction

Equation: H2 + O2 → H2O
Calculation: 0 + 0 → (2)(+1) + (-2) = 0
Explanation: In this equation both H2 and O2 are free elements and following Rule #1, their OS is "0." The product is H2O, which has a total OS of "0." According to Rule #6, the OS of oxygen is usually -2. So, the OS of H2 must be +1.

Decomposition reactions are the reverse of combination reactions, meaning they are the breakdown of a chemical compound into the individual elements. General Equation:

AB → A + B

Example 7: Decomposition Reaction

Equation: H2O → H2 + O2
Calculation: (2)(+1) + (-2) = 0 → 0 + 0
Explanation: In this equation the water is "decomposed" into a Hydrogen and Oxygen. Similar to the previous sample the H2O has a total OS of "0," thus according to Rule #6 the OS of oxygen is usually -2 so the OS of H2 must be +1.

A single replacement reaction involves the "replacing" of an element in the reactants with another element in the products.General Equation:

A + BC → AB + C

Example 8: Single Replacement Reaction


Cl2 + NaBr → NaCl + Br2

Calculation: (0) + ((+1) + (-1) = 0) -> ((+1) + (-1) = 0) + 0
Explanation: In this equation Br is replaced with Cl and Cl is reduced, while Br is oxidized.

A double replacement reaction is similar to a double replacement reaction, but involves "replacing" two elements in the reactants, with two in the products. General Equation:

AB + CD → AD + CB

Example 9: Double Replacement Reaction

Equation: Fe2O3 + HCl → FeCl3 + H2O
Explanation: In this equation Fe and H trade places and oxygen and chlorine trade places.

Combustion reactions always involve oxygen, in the form of O2 and are almost always exothermic, meaning they produce heat. General Equation:

CxHy + O2 → CO2 + H2O


Disproportionation Reactions: In some redox reactions substances can be both oxidized and reduced. These are known as disproportionation reactions with a general equation:

2A → A' + A"


Example 10: Disproportionation Reaction

Disproportionation reactions have some practical significance in everyday life including the reaction of hydrogen peroxide, H2O2 poured over a cut. This a decomposition reaction of hydrogen peroxide, which produces oxygen and water. Oxygen is present in all parts of the chemical equation and as a result it is both oxidized and reduced. Reaction:

2H2O2(aq) → 2H2O(l) + O2(g)

Explanation: In the reactants H has an O has an OS of -1, which changes to -2 for the product, H2O (reduced) and 0 for the product, O2 (oxidized).


  • http://www.youtube.com/watch?v=yp60-oVxrT4
  • Remember the 7 Rules of OSs (these are vital to understanding redox reactions)
  • Oxidation signifies a loss of electron and reduction signifies a gain of electrons.
  • Balancing redox reactions is an important step that changes in neutral, basic, and acidic solutions.
  • Remember the various types of redox reactions
    • Combination and Decomposition
    • Displacement Reactions (Single and Double)
    • Combustion
    • Disproportionation
  • The oxidizing agent undergoes reduction and the reducing agent undergoes oxidation.


  1. Petrucci, et al. General Chemistry: Principles & Modern Applications. 9th ed. Upper Saddle River, New Jersey: Pearson/Prentice Hall, 2007.
  2. Sadava, et al. Life: The Science of Biology. 8th ed. New York, NY. W.H. Freeman and Company, 2007


  • Christopher Spohrer (UCD), Christina Breitenbuecher (UCD), Luvleen Brar (UCD)
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