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Oxidation-Reduction (Redox) reactions. Valence. Oxidation number. Oxidizing and reducing agents.

Def. Species. Atoms, molecules, molecular fragments, ions, etc. subjected to a chemical process or to a measurement.

Def. Radical. A group of elements which usually behaves as if it were a single element in the formation of a compound. It is a molecular fragment that carries a valence and combines with atoms to form compounds. Examples are found in Table 1: C2H3O2, CO3, OH, NO3, SO4, etc. Radicals are typically highly reactive.

Formation of compounds. In forming a compound, two or more atoms of different elements combine together to form a molecule. An atom of one element may combine with one or more atoms of certain other elements to form a compound, but only certain other elements, not any other elements. For example, atoms of iron will combine with atoms of oxygen to form an iron oxide. But atoms of iron will not combine with atoms of copper, or of lead, or of zinc, or atoms of many other elements. What is it that determines what elements will combine with each other to form compounds and what elements won’t combine with each other? What is it that holds the various atoms in a molecule together to create a distinct unit (the molecule) with its own distinct properties? What are the forces that hold the atoms of a molecule together and where do they come from?

Chemical bonding. The mechanism by which two elements (or an element and radical) unite to form a compound is generally considered to be the following: the electrons in the outer shell of one atom are either transferred to another atom or shared with another atom so as to make the outer shells of both atoms as complete as possible. Thus in the case of the compound H2O two hydrogen atoms each share their single electron with an oxygen atom (whose outer shell contains 6 electrons). By this sharing, the outer shell of the oxygen atom gains two electrons and thus becomes completed with 8 electrons, providing chemical stability.

Valence (or Oxidation number). When the atoms of two different elements combine, each atom either gains or loses electrons. The valence of an element is the net charge on an atom as given by the algebraic sum of the positive (proton) charges and the negative (electron) charges. When an atom is in its pure state its valence is 0 since it has the same number of electrons as protons. When an element is in a compound its valence will be either positive or negative depending on the number of electrons that it lost or gained in forming the chemical bond.

Syn. Valence number, Oxidation number, Oxidation state

Two categories of elements. The mechanism that determines how atoms combine to form molecules thus involves interactions involving the outer shell electrons of the atoms. We now consider two categories of elements:

1. Category A elements. Elements with one or two electrons in their outer orbit (or shell). This group includes the most of the metals. Copper, for example, has one electron in its outer orbit. Iron has two. Elements with one or two electrons in their outer orbits tend to have them stolen away from them by other elements when forming a molecule. If they have their outer orbit electrons stolen away from them, they are said to be oxidized and their valence increases. (the term “oxidize” deriving from the fact that the element oxygen often steals electrons in this way). If copper has its outer orbit electron stolen,

            Cu →Cu+ + e-

The valence of Cu is 0 and of Cu+ is +1. The valence corresponds to the net charge on the Cu+ ion (the Cu+ ion consists of the Cu atom minus an electron — it has n protons and n-1 electrons).


The valences of Category A elements are all positive when they are in compounds. Category A elements correspond to those elements shown in the upper half of Table 1.

2. Category B elements. Elements and radicals whose outer orbit lacks one or two electrons. This group includes oxygen and other nonmetals and many radicals. Elements whose outer orbit lacks one or two electrons tend to steal electrons from other elements in order to complete their outer orbit. Oxygen unites with many metals, and in doing so, steals their outer orbit electrons. If an element steals electrons from another element in the process of forming a molecule, it is said to oxidize the element and to be itself reduced. Its valence drops (becomes more negative due to increased electrons). For the case of oxygen,

            O + 2e- → O-2

The valence of O is 0 and of O-2 is -2.

The valences of Category B elements are all negative when they are in compounds. Category B elements correspond to those elements shown in the lower half of Table 1.

In general, Category A elements tend to combine with Category B elements with the Category B elements utilizing the outer orbit electrons of the Category A elements – one element donates electrons and the other receives them. A Category A element bonds together with a Category B element to the benefit of both (creating a more stable state for both).


Def. Oxidation-Reduction reaction. Any reaction in which oxidation and reduction take place.

An oxidation-reduction reaction distinguishes itself from other reactions by the fact that in an oxidation-reduction reaction a change of valence occurs in one or more elements. If an atom changes in valence, the reaction is an oxidation-reduction reaction.

Def. Redox reaction. An Oxidation-Reduction reaction.

Def. Oxidizing agent. An oxidizing agent is the element or compound in an oxidation-reduction (redox) reaction that accepts an electron from another species. Because the oxidizing agent is gaining electrons (and is thus often called an electron acceptor), it is said to have been reduced.

Syn. Oxidant, oxidizer

In most cases the oxidizing agent is oxygen or a compound containing oxygen. An oxidation/reduction process is then a process in which oxygen gains electrons and the substance that is oxidized loses electrons. Consider the following: Oxygen acts on metals such as tin, lead, copper, zinc, and iron to form oxides of the metals. Oxygen is then the oxidizing agent. In the process oxygen gains electrons and its valence is lowered; the metals lose electrons and their valence is raised.

Example of an oxidation-reduction reaction. Consider the following reaction in which sulfur combines with oxygen to form sulfur dioxide

            0__0___+4 _-2

1)        S + O2 → SO2

where we have written the valence above the symbols to show what is happening. Here the sulfur atom goes from a valence of 0 to a valence of +4. It changes from sulfur in the pure state with a valence of 0 to a sulfur ion with a valence of four in the molecule. Meanwhile, oxygen goes from a valence of 0 to a valence of -2. It changes from oxygen in the pure state with a valence of 0 to an oxygen ion with a valence of -2.

In this reaction oxygen is the oxidizing agent and oxidizes the sulfur. At the same time, sulfur can be viewed as a reducing agent which reduces the oxygen. An oxidizing agent steals electrons from the element it oxidizes and raises its valence. A reducing agent gives electrons to another element and reduces its valence.

The following are some examples of oxidation-reduction reactions where represents a gaseous state and represents a solid:

Oxidation-Reduction reactions

C + O2 → CO2

CO2 + C → 2 CO 

Fe2O3 + 3 CO → 2 Fe + 3 CO2

S + O2 → SO2

Mg + 2HNO3 → Mg(NO3)2 + H2

3 Cu + 8 HNO3 →3 Cu(NO3)2 + 2 NO + 4H2O

H2SO4 → H2O + SO2+ (O)

Ca + 2 H2O → Ca(OH)2 + H2

H2O2 + H2S → 2 H2O + S

2 H2S + 3 O2 → 2 SO2 + 2 H2O

2 H2S + 2 O2 → SO2 + 2 H2O + S

2 H2S + O2 → 2 H2O + 2 S

SnCl2 + 2 HgCl2 → SnCl4 + Hg2Cl2

SnCl2 + HgCl2 → SnCl4 + Hg

4 Cu + O2 → 2 Cu2O

Example of an oxidation-reduction reaction not involving oxygen. Stannous chloride is a good reducing agent. If we add stannous chloride to a solution of mercuric chloride, we have an interesting example of oxidation in which there is no transfer of oxygen at all. By varying the amount used, we can reduce the mercuric chloride to mercurous chloride, or even to metallic mercury. The equations follow:

            SnCl2 + 2 HgCl2 → SnCl4 + HgCl2

            SnCl2 + HgCl2 → SnCl4 + Hg

In each case the mercuric chloride acted as an oxidizing agent, and the stannous chloride as a reducing agent. The tin, which has a valence of two in stannous chloride, is oxidized to the stannic condition, in which the valence of tin is four. At the same time, the mercury, which has a valence of two in mercurous chloride, is reduced until its valence is one in mercurous chloride, or zero in metallic mercury. The ionic changes may be expressed as follows for the second equation:

            Sn++ + Mg++ → Sn++++ + Hg0

                        Source: Dull, Brooks, Metcalfe. Modern Chemistry. pp. 527,528


            Oxidation: valence of oxidized element increases

            Reduction: valence of reduced element decreases

In the equation

            SnCl2 + HgCl2 → SnCl4 + Hg

the valence of Sn goes from +2 to +4 and the valence of Mg drops from +2 to 0. Sn has been oxidized and Mg reduced. The mercuric chloride (HgCl2) acts as an oxidizing agent because it causes oxidation of Sn, and the stannous chloride (SnCl2) acts as a reducing agent because it causes reduction of Hg.

To find out what oxidizes what and what reduces what, write the valence of each element above it so you can see what is happening.

           +2 -1 __ +2 -1 __ +4 -1 __ 0

            SnCl2 + HgCl2 → SnCl4 + Hg

Changes in valence:

            Sn: +2 → +4

            Hg: +2 → 0

1. Consider the reaction

            C + O2 → CO2

            0 __ 0 __ +4 -2

            C + O2 → CO2

Changes in valence:

            C: 0 → +4

            O: 0 → -2


The carbon has been oxidized and the oxygen reduced. The oxygen has acted as an oxidizing agent on the carbon, causing an increase in its valence of 4. The carbon has acted as a reducing agent on the oxygen, causing a drop in its valence of 2.

2. Consider the reaction

            Fe2O3 + 3 CO → 2 Fe + 3 CO2

            +3 -2 __ +2 -2 ___ 0 ___ +4 -2

            Fe2O3 + 3 CO → 2 Fe + 3 CO2

Changes in valence: 

            Fe: +3 → 0

            C: +2 → +4

The carbon has been oxidized and the iron reduced. The iron oxide (Fe2O3) has acted as an oxidizing agent on the carbon, causing an increase in its valence of 2. The carbon monoxide (CO) has acted as a reducing agent on the iron, causing a drop in its valence of 3.

The above reaction is one in which carbon monoxide is employed as a reducing agent to remove the oxygen from hot iron ore (Fe2O3).

Note. This is presumably the source of the term “reducing agent”. Here the carbon monoxide preforms the function of reducing iron ore to iron. A substance that reduces a metallic ore to the basic metal is viewed as “reducing” it.

Oxidizing agents, reducing agents. In general, in an oxidation-reduction reaction, some element or elements are oxidized and others are reduced. Reactants containing elements that are oxidized are said to be the reducing agents and reactants containing elements that are reduced are said to be the oxidizing agents.

A substance is said to be a good oxidizing agent if it tends to oxidize other substances (i.e. if it tends to oxidize some element in the substance). A substance is said to be a good reducing agent if it tends to reduce other substances.

Good oxidizing agents:


            ■ hydrogen peroxide, H2O2

            ■ nitric acid, HNO3

            ■ potassium chlorate, KClO3

            ■ barium peroxide, Ba(NO3)3

            ■ chlorine, Cl2

            ■ cupric oxide, CuO

            ■ nitrogen dioxide, NO2

            ■ ozone, O3

            ■ sodium chlorate, NaClO3

            ■ sodium peroxide, Na2O2

            ■ sulfuric acid, H2SO4

Good reducing agents:

            ■ calcium, Ca

            ■ carbon monoxide, CO

            ■ charcoal, C

            ■ coke, C

            ■ hydrogen sulfide, H2S

            ■ oxalic acid, (COOH)2

            ■ stannous chloride, SnCl2

Ref. Dull, Brooks, Metcalfe. Modern Chemistry. pp. 579, 581, oxidizing agents, reducing agents

Oxides of oxygen: tin, lead, copper, zinc, iron, carbon, sulfur, phosphorus

stannous oxide, SnO, SnO2

lead oxide, PbO, Pb2O, PbO2, Pb3O4

copper oxide, Cu2O

zinc oxide, ZnO

iron oxide, FeO, Fe2O3, Fe3O4

carbon oxide, CO, CO2

sulfur oxide, SO2, SO3

phosphorus oxide, PO4, P2O3, P2O5

Rules for Assigning Oxidation Numbers

•The oxidation number of an atom is zero in a neutral substance that contains atoms of only one element. Thus, the atoms in O2, O3, P4, S8, and aluminum metal all have an oxidation number of 0.

•The oxidation number of monatomic ions is equal to the charge on the ion. The oxidation number of sodium in the Na+ ion is +1, for example, and the oxidation number of chlorine in the Cl- ion is -1.


•The oxidation number of hydrogen is +1 when it is combined with a nonmetal. Hydrogen is therefore in the +1 oxidation state in CH4, NH3, H2O, and HCl.

•The oxidation number of hydrogen is -1 when it is combined with a metal. Hydrogen is therefore in the -1 oxidation state in LiH, NaH, CaH2, and LiAlH4.

•The metals in Group IA form compounds (such as Li3N and Na2S) in which the metal atom is in the +1 oxidation state.

•The elements in Group IIA form compounds (such as Mg3N2 and CaCO3) in which the metal atom is in the +2 oxidation state.


•Oxygen usually has an oxidation number of -2. Exceptions include molecules and polyatomic ions that contain O-O bonds, such as O2, O3, H2O2, and the O22- ion.


•The nonmetals in Group VIIA often form compounds (such as AlF3, HCl, and ZnBr2) in which the nonmetal is in the -1 oxidation state.

•The sum of the oxidation numbers of the atoms in a molecule is equal to the charge on the molecule.

•The most electronegative element in a compound has a negative oxidation number.

                                    Source: The Bodner Group, Purdue University

Balancing oxidation-reduction (redox) equations

Conditions that must be met by a balanced redox equation:

Condition 1. As with all equations, redox or not, the number of atoms of each element appearing on the right side of the equation must be the same as the number of atoms of that element on left side of the equation.

Example. In the equation

            2 H2S + O2 → 2 H2O + 2 S

there are 4 H atoms, 2 S atoms and 2 O atoms on both sides of the equation.


Condition 2. In the equation, the following must hold: 

            ∑ valence increases of oxidized element atoms = ∑ valence decreases of reduced element atoms

Example. Consider the correctly balanced equation 

1)        Fe2O3 + 3 CO → 2 Fe + 3 CO2

            +3 -2 __ +2 -2 ___ 0 ___ +4 -2

            Fe2O3 + 3 CO → 2 Fe + 3 CO2

Changes in valence: 

            Fe: +3 → 0                    valence change = 3 - 0 =3

            C: +2 → +4                  valence change = 4 - 2 = 2

The oxidized element is C. There are three atoms of C, each of which are oxidized by an amount of 2. Thus

            ∑ valence increases of oxidized element atoms = 3×2 = 6

The reduced element is Fe. There are two atoms of Fe, each of which are reduced by an amount of 3. Thus

            ∑ valence decreases of reduced element atoms = 3×2 = 6

Steps in balancing an equation

1. Write the skeleton equation i.e.

            Fe2O3 + CO → Fe + CO2

2. Determine which elements are oxidized and which are reduced and rewrite the equation with the valences above the elements

            +3 -2 _ +2 -2 _ 0 __ +4 -2

            Fe2O3 + CO → Fe + CO2

3. Try to find coefficients that will satisfy conditions 1 and 2 above.

Oxidation-reduction equations of the combination (synthesis), decomposition, and single replacement types are rather easily balanced. Some more complex oxidation-reduction equations are more difficult to balance.


  Dull, Brooks, Metcalfe. Modern Chemistry.

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