# Balancing equations

## About Balancing Equations

The term oxidation was first used to mean the addition of oxygen to an element or compound, or the removal of hydrogen from a compound. Reduction meant the addition of hydrogen to an element or compound, or the removal of oxygen from a compound. Such definitions have been extended and now-a-days many oxidation-reduction, or redox, reactions are best interpreted in terms of transfer of electrons.Oxidation is the loss of electrons by an atom, ion or molecule.

Reduction is the gain of electrons by an atom, ion or molecule.

An oxidising agent takes electrons; it is an electron acceptor.

A reducing agent gives electrons; it is an electron donor.

When Fe2+(aq) ions are being oxidised they are acting as reducing agents, and when Fe3+(aq) ions are being reduced they are acting as oxidising agents.

### Oxidation Number or Oxidation State use in Balancing Equations

When an element is oxidised it must be acting as a reducing agent and it, therefore, loses electrons; when reduced, it gains electrons. The oxidation state or oxidation number of an element is the number of electrons it might be considered to have lost or gained.All elements in the elementary, uncombined state are given oxidation numbers of zero. When sodium, for example, is oxidised it loses one electron, and the Na+ ion is said to have an oxidation number of +1. Similarly, the Cu2+ and Al3+ ions have oxidation numbers of +2 and +3, whilst F- and O2- have oxidation numbers of - 1 and - 2. For simple ions, the oxidation number is equal to the ionic charge,

## Balancing of Redox Equations

### Oxidation Number / State Method for Balancing Equations

This method is based on the principle that the number of electrons lost in oxidation must be equal to the number of electrons gained in reduction. The steps to be followed are.

i) Write the equation (if it is not complete, then complete it) representing the chemical changes.

ii) By knowing oxidation numbers of elements, identify which atom(s) is(are) undergoing oxidation and reduction. Write down separate equations for oxidation and reduction.

iii) Add respective electrons on the right of oxidation reaction and on the left of reduction reaction. Care must be taken to ensure that the net charge on both the sides of the equation is same.

iv) Multiply the oxidation and reduction reactions by suitable numbers to make the number of electrons lost in oxidation reactions equal to the number of electrons gained in reduction reactions.

v) Transfer the coefficient of the oxidizing and reducing agents and their products to the main equation.

By inspection, arrive at the co-efficients of the species not undergoing oxidation or reduction.

### Half-Reaction or Ion-Electron Method Balancing Equations

This method involves the following steps :

i) Divide the complete equation into two half reactions, one representing oxidation and the other reduction.

ii) Balance the atoms in each half reaction separately according to the following steps:

a) First of all balance the atoms other than H and O.

b) In a reaction taking place in acidic or neutral medium, oxygen atoms are balanced by adding molecules of water to the side deficient in oxygen atoms while hydrogen atoms are balanced by adding H+ ions to the other side deficient in hydrogen atoms. On the other hand, in alkaline medium (OH-), every excess of oxygen atom on one side is balanced by adding one H2O to the same side and 2OH- to the other side. In case hydrogen is still unbalanced, then balance by adding one OH-, for every excess of H atom on the same side as the excess and one H2O on the other side.

c) Equalize the charge on both sides by adding a suitable number of electrons to the side deficient in negative charge.

iii) Multiply the two half reactions by suitable integers so that the total number of electrons gained in one half reaction is equal to the number of electrons lost in the other half reaction.

iv) Add the two balanced half equations and cancel any term common to both sides.

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