Balancing of Redox Reactions

Oxidation-Reduction Reactions, often known as redox reactions, occur when one reactant is oxidised while another is reduced at the same time. This discussion will show you how to balance an equation of redox reaction.

Redox Reactions: How to Spot Them

Identifying whether or not a redox process is an oxidation-reduction reaction is the first step in balancing it. This necessitates the change of oxidation states of one or more species during the process. The redox reaction will include both a reduction and an oxidation component to maintain charge neutrality in the sample. To make the reaction easier to grasp, they are frequently split into two hypothetical half-reactions. This necessitates determining which elements are oxidised and which are reduced. Consider the following reaction:

Cu(s) + 2Ag+ (aq) → Cu2+ (aq)+ 2Ag(s)

Splitting the equation into two hypothetical half-reactions is the first step in determining whether the reaction is a redox reaction or not. Let’s start with the copper atoms and their half-reaction:

Cu(s) →  Cu2+(aq)

Because copper is a pure element, its oxidation state on the left side is 0 (zero). On the right hand side of the equation, copper’s oxidation state is +2. As the oxidation states advance from 0 in Cu to +2 in Cu2+, the copper in this half-reaction is oxidised. Consider the atoms of silver:

2Ag+(aq) →  2Ag(s)

Silver’s oxidation state on the left side is +1 in this half-reaction. Because silver is a pure element, its oxidation state on the right is 0 (zero). The reduction half-reaction occurs when the oxidation state of silver goes from +1 to 0.

As a result, this is a redox reaction since both reduction and oxidation half-reactions take place (via the transfer of electrons, that are not explicitly shown in equations 2). Once the reaction has been validated, it is typically required to balance it (the reaction in equation 1 is already balanced), which can be done in two ways because the reaction might occur in neutral, acidic, or basic conditions.

Balancing Redox Reaction

Redox reaction balancing is slightly more complicated than conventional reaction balancing, although it still follows a pretty straightforward set of rules. One significant distinction is the requirement to know the half-reactions of the reactants involved; a half-reaction table is quite helpful in this regard. Half-reactions are frequently useful since they can be combined to form a complete net equation. Although the half-reactions must be known in order to complete a redox reaction, they may often be calculated without the assistance of a half-reaction table. The acidic and basic solution examples explain this. For aqueous reactions under acidic or basic circumstances, extra rules must be followed in addition to the standard principles for neutral conditions.

We all clearly know about the method of balancing  a redox reaction by half reaction method. So let us discuss the oxidation number method.

The oxidation number method and the half-reaction method have no real differences. It’s basically a different technique of keeping track of the electrons that are moved during a reaction.

Recognizing the oxidation and reduction parts of a redox equation is the only surefire method to balance it. Then you strike a balance by ensuring that the electron loss and gain are equal.You determine the oxidation numbers of all atoms using the oxidation number approach. The atoms that have altered are then multiplied by small whole numbers. You’re making the entire electron loss equal to the total electron gain. The remaining atoms are then balanced.

For a really simple equation that you could probably balance in your head, here’s how the oxidation number approach works:

                                                          Zn + HCl →  ZnCl2 + H2

Step 1: Determine which atoms change their oxidation number.

LHS: Zn = 0 ; H = +1 ; Cl = -1

RHS: Zn = +2 ; Cl = -1 ; H = 0

So the following are the changes in oxidation number:

Zn : 0 → +2 ; change = +2

H : +1 → 0 ; change = -1

Step 2: Equalise the oxidation number changes.

Every Zn atom has lost two electrons, whereas every H atom has gained one. For every atom of Zn, you’ll need two atoms of H. We have total changes of +2 and -2 as a result of this.

Step 3. Insert coefficients to get these numbers

1Zn +2HCl  →  1ZnCl2 + 1H2

So the balanced equation is given as

Zn + 2HCl  →  ZnCl2 +H2

Conclusion

An oxidation-reduction process is a chemical reaction in which electrons are transferred between two substances. The loss of electrons in oxidation is called oxidation, while the gain of electrons is called reduction. The reducing agent loses electrons and is oxidised, whereas the oxidising agent receives electrons and is reduced. Because the reduction of one molecule causes the oxidation of another, oxidation and reduction will always occur simultaneously. The reducing agent will shift to a more positive or neutral state, whereas the oxidising agent will shift to a more negative state (less positive). To balance the redox reaction, the half-reaction approach is utilised, which requires both mass and charge to be balanced.