Oxidation-Reduction Reaction

Oxidation numbers, sometimes called oxidation states, help chemists keep track of the numbers of electrons that surround each atom in a chemical reaction, and how they change in oxidation-reduction reactions. When an atom gains an electron (is reduced), its oxidation number is increased by one. There are some simple rules for assigning oxidation numbers to elements in chemical compounds. These rules are:

1. The oxidation number of an element, having neither gained nor lost any of its electrons, is zero. For example, the oxidation number of pure copper, Cu, is zero, as is the oxidation number of each oxygen atom in a molecule of oxygen, O₂.
2. The oxidation number of an elemental ion is the same as its charge. An ion of copper with a plus two charge, Cu²⁺, has an oxidation number of plus two. A fluoride ion, F^–, has an oxidation number of minus one.
3. Some elements almost always form compounds in which they have a particular oxidation number. Aluminum always forms a plus three ion and therefore exists in the plus three oxidation state in compounds. Sodium and other alkali metals almost always form a plus one ion; its oxidation state is plus one. Hydrogen can form compounds in which the hydrogen atom has an oxidation number of either plus one or minus one. When the hydrogen has an oxidation number of plus one, it is written on the left hand side of the chemical formula. If its oxidation number is minus one, it is written on the right hand side. Oxygen usually has a minus two oxidation number. Chlorine and other halogens usually take on a minus one charge. Other elements are not so predictable. Nitrogen can have oxidation numbers of +5,+4,+3, +2,+1, and -3.
4. The sum of the oxidation numbers in a neutral molecule or compound is zero. Table salt, with the chemical formula of sodium chloride, is made up of two ions, a positively charged sodium ion and a negatively charged chloride ion. A water molecule consists of two hydrogen atoms, each having an oxidation number of plus one, and an oxygen atom with an oxidation number of minus two.

It is often easier to follow oxidation-reduction reactions if they are split into two half reactions. One-half reaction indicates what is happening to the chemical substances and electrons in the oxidation portion of the reaction. The other half-reaction does the same for the reduction portion. The complete reaction is the sum of the two half reactions.

A useful tool for chemists is a table of standard reduction potentials. This table lists common half reactions, and assigns each a numerical value that indicates how easily the reduction reaction proceeds-that is, how eagerly electrons are accepted. A high standard reduction potential value indicates that the substance is easily reduced. A low standard reduction potential indicates that the substance is easily oxidized-it prefers to lose electrons. In general, a substance will oxidize something that has a lower reduction potential than it has. The halogens, chemical elements found in group 17 of the periodic table, are strong oxidizing agents because their atoms readily accept negative ions. The alkali metals such as sodium, found on the left side of the periodic table in group 1, are strong reducing agents because their atoms readily give up an electron, becoming positive ions. The arbitrary zero point for standard reduction potentials has been designated as this reaction:

This reaction has been assigned a potential of 0.000 volts under standard conditions. The standard reduction potential for fluorine gas is 2.890 volts while that for sodium metal is -2.714 volts.