Balancing Redox reaction

Oxidation-Reduction Reactions, or redox reactions, are reactions in which one reactant is oxidized and one reactant is reduced simultaneously. This module demonstrates how to balance various redox equations.

Redox Reactions identification

The first step in balancing any redox reaction is determining whether or not it is even an oxidation-reduction reaction, which requires that species exhibits changing oxidation states during the reaction. To maintain charge neutrality in the sample, the redox reaction will entail both a reduction component and an oxidation components and is often separated into independent two hypothetical half-reactions to aid in understanding the reaction. This requires identifying which element is oxidized and which element is reduced. For example, consider this reaction:

Cu(s)+2Ag+(aq)→Cu2+(aq)+2Ag(s)

The first step in determining whether the reaction is a redox reaction is to splitting the equation into two hypothetical half-reactions. Let's start with the half-reaction involving the copper atoms:

Cu(s)→Cu2+(aq)(2a)(2a)Cu(s)→Cu2+(aq)

The oxidation state of copper on the left side is 0 because it is an element on its own. The oxidation state of copper on the right hand side of the equation is +2. The copper in this half-reaction is oxidized as the oxidation states increases from 0 in Cu to +2 in Cu2+. Now consider the silver atoms.

2Ag+(aq)→2Ag(s)(2b)(2b)2Ag+(aq)→2Ag(s)

In this half-reaction, the oxidation state of silver on the left side is a +1. The oxidation state of silver on the right is 0 because it is an element on its own. Because the oxidation state of silver decreases from +1 to 0, this is the reduction half-reaction.

The examples of redox reactions in everyday life are as follows:

photosynthesis,

respiration,

biological processes,

decay,

corrosion/rusting,

combustion and batteries.

Types of Redox Reactions

1. Decomposition Reaction

2. Combination Reaction

3. Displacement Reaction

4. Disproportionation Reactions