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Balancing of Redox Reactions for JEE

Redox Reactions - A Brief Description

Last updated date: 21st Mar 2023
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Redox reactions are oxidation-reduction processes where both oxidation and reduction occur simultaneously in a single reaction, and the oxidation states of the reactants change when reactants are converted into products. Redox is a term that describes the reduction-oxidation process. Oxidation and reduction are the two subtypes of redox reactions.

Oxidation is defined as the loss of electrons and the resulting increase in the oxidation state of a given reactant. Reduction is the process of gaining electrons and decreasing the oxidation state of a reactant. In a chemical reaction, the material that is being oxidised  is the reducing agent, and the substance that is being reduced  is the oxidising agent. Some examples of oxidising agents are KMnO4,K2Cr2O7,HNO3,KClO3. Some examples of reducing agents are FeCl2, FeSO4, SnCl2, Hg2Cl2.

Balancing of Redox Reactions

A chemical equation that is balanced correctly represents the quantities of reactants and products in chemical processes. The Law of Conservation of Mass states that no mass is created or lost during a normal chemical reaction. This means that each element must have the same number of atoms on both sides of a chemical equation. Additionally, the charges on the reactant side should add up to be the same as that on the product side. When these two requirements are satisfied, the equation is regarded as being balanced. Redox reaction balance is therefore of extreme significance. The redox reaction can be balanced in two ways. The change in oxidation number of the oxidising and reducing agents is one technique, while the other is based on separating the redox reaction into two half processes, one of reduction and the other of oxidation. The method in which the redox reaction is separated into two half processes is called the ion electron method.

Balancing Redox Reactions by the Ion-Electron Method

To balance equation some basic steps should be followed:

• Write the basic equation containing the oxidising and reducing agents, as well as  the products in their ionic form.  Assign each atom its oxidation number in these ionic reactions.

• The equation must be separated into two half reactions. The oxidation half and the reduction half. The oxidation half has the part of the reaction in which reactants oxidise to form the product. The reduction half is the part of reaction in which reactants get reduced to form a product.

• Separately, balance all the atoms in each half reaction (apart from O and H). This is called mass balance.

• Add H2O to balance the oxygen atoms by ion electron method in acidic medium, and H+ to balance the hydrogen atoms.

• The atoms must first be balanced for reactions in the basic medium, just as they would be in an acidic solution. Add OH- ion to both sides of the half-reaction for each H+ ion after that. H+ and OH- should always be combined to form H2O whenever they are on the same side.

• To balance the charges in each half-reaction, add electrons to one side. In a reduction half-reaction, electrons are added to the reactant side, while in an oxidation half-reaction, they are added to the product side. If enough electrons are transferred to one side of the half-reaction, the overall charge on that side should equal the charge on the other side. This procedure is called a charge balance.

• Multiply the half-reactions through by the appropriate coefficients to balance the number of electrons as needed.

• Add the two half-reactions, and check the final equation by inspection. Reduce the balanced equation to lowest terms (coefficients).

Balancing by Ion Electron Method Example

• Balance the reaction in the basic medium.

$\mathrm{MnO}_{4}^{-}(a q)+I^{-}(a q) \rightarrow \mathrm{MnO}_{2}(s)+I_{2}(s)$

• Assign oxidation number to ions.

$\overset{+7}{MnO_4^-}~(aq)+\overset{-1}{I^-}~(aq)\rightarrow \overset{+4}{MnO_2}~(s)+\overset{0}{I_2}~(s)$

• Oxidation and reduction halves

\begin{align} &I^{-}(a q) \rightarrow I_{2}(s) \\ &\mathrm{MnO}_{4}^{-}(a q) \rightarrow \mathrm{MnO}_{2}(a q) \end{align}

• Start balancing atoms on each half

\begin{align} &2 \mathrm{I}^{-}(a q) \rightarrow I_{2}(s) \\ &\mathrm{MnO}_{4}^{-}(a q) \rightarrow \mathrm{MnO}_{2}(a q) \end{align}

• Add electrons to balance the charge

\begin{align} &2 I^{-}(a q) \rightarrow I_{2}(s)+2 e^{-} \\ &\mathrm{MnO}_{4}{ }^{-}(a q)+3 e^{-} \rightarrow \mathrm{MnO}_{2}(a q) \end{align}

• Balance charge in the reduction half reaction by adding 4 hydroxide ions to the product side.

\begin{align} &2 \mathrm{I}^{-}(a q) \rightarrow \mathrm{I}_{2}(s)+2 e^{-} \\ &\mathrm{MnO}_{4}^{-}(a q)+3 e^{-} \rightarrow \mathrm{MnO}_{2}(a q)+4 \mathrm{OH}^{-} \end{align}

• To balance O atoms, add 2 water molecules  to the reactant side.

\begin{align} &2 \mathrm{I}^{-}(a q) \rightarrow \mathrm{I}_{2}(s)+2 e^{-} \\&\mathrm{MnO}_{4}^{-}(a q)+3 e^{-}+2 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{MnO}_{2}(a q)+4 \mathrm{OH}^{-} \end{align}

• Equalise the number of electrons in both reactions.

\begin{align}&6 \mathrm{I}^{-}(a q) \rightarrow 3 \mathrm{I}_{2}(s)+6 e^{-} \\&2 \mathrm{MnO}_{4}^{-}(a q)+6 e^{-}+4 \mathrm{H}_{2} \mathrm{O} \rightarrow 2 \mathrm{MnO}_{2}(a q)+8 \mathrm{OH}^{-} \end{align}

• Add both half reactions to get the final reaction.

$2 \mathrm{MnO}_{4}{ }^{-}(a q)+6 \mathrm{I}^{-}(a q)+4 \mathrm{H}_{2} \mathrm{O}(l) \rightarrow 2 \mathrm{MnO}_{2}(s)+3 \mathrm{I}_{2}(s)+8 \mathrm{OH}^{-}$

Conclusion

In chemical reactions, a balanced chemical equation is important in describing the quantities of reactants and products accurately. According to the Law of Conservation of Mass, no mass can be generated or destroyed. This indicates that both sides of a chemical equation must have the same amount of atoms of each element. The redox reaction can be balanced in two ways. The change in oxidation number of the oxidising and reducing agents is one technique called oxidation number method, while the other is based on separating the redox reaction into two half processes, one of reduction and the other of oxidation. The method in which the redox reaction is separated into two half processes is called the ion electron method.

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FAQs on Balancing of Redox Reactions for JEE

1. Define briefly  the oxidation number method of balancing reactions.

The oxidation number approach is based on the difference between the oxidising and reducing agents oxidation numbers. It is important, as with any reaction, to write accurate compositions and formulas. For example: The number of electrons gained during the reduction reaction equals the number of electrons lost during the oxidation reaction, according to this technique. H+, H2O and OH- are used in balancing redox reactions similarly as half reaction methods.

2. What is redox titration?

A redox reaction between the titrant and the analyte is used to determine the concentration of a given analyte in a laboratory setting. A potentiometer or a redox indicator may be required for these types of titrations.

The titrant and the analyte undergo an oxidation-reduction reaction, which is used in redox titration. One of the most common laboratory methods for determining the quantity of unknown analytes is redox titration. The titration of potassium permanganate (KMnO4) against oxalic acid (C2H2O4) is an example of a redox titration.