Revision: Class 11 >> Redox Reactions and Electrochemistry NEET (UG) Redox Reactions and Electrochemistry

Any reaction that involves both oxidation and reduction occurring simultaneously is called an oxidation-reduction reaction or simply a redox reaction.

or

The chemical reaction in which both oxidation and reduction occur simultaneously is called a redox reaction.

The species which gets itself oxidised and reduce another species is called reducing agent.

or

A substance which involves an increase in the oxidation number of one or more of its elements. A reducing agent helps reduce the other substance by being oxidised.

\[\begin{aligned} & \mathrm{ZnO}+\mathrm{C}\longrightarrow\mathrm{Zn}+\mathrm{CO} \\ & \mathrm{Reducing} \\ & \mathrm{agent} \end{aligned}\]

The species which gets itself reduced and oxidise another species is called oxidising agent.

or

A substance which involves a decrease in the oxidation number of one or more of its elements. An oxidising agent helps oxidise the other substance by being reduced itself.

\[\begin{aligned} & \mathrm{S}+6\mathrm{HNO}_{3}\longrightarrow\mathrm{H}_{2}\mathrm{SO}_{4}+2\mathrm{H}_{2}\mathrm{O}+6\mathrm{NO}_{2} \\ & \mathrm{Oxidising} \\ & \mathrm{agent} \end{aligned}\]

"Reduction" is defined as the removal of oxygen/ electronegative element from a substance or the addition of hydrogen/Electron/ electropositive element to a substance.

or

A process involving decrease in oxidation number by gain of electrons.

"Oxidation" is defined as the addition of oxygen/electronegative element to a substance or the removal of hydrogen/ Electron/ electropositive element from a substance.

or

A process involving an increase in oxidation number by the loss of electrons.

Electrochemistry is the branch of chemistry that deals with the production of electricity from the energy released during spontaneous reactions and the use of electrical energy to drive non-spontaneous reactions.

Oxidation number (also called oxidation state) is the charge that an atom of an element appears to have when present in a combined state with other atoms. It is a hypothetical charge assigned by assuming all bonds are ionic — atoms in real molecules like H₂O do not actually carry these charges.

It states that at infinite dilution, the molar conductance of an electrolyte is the sum of molar conductances of its ions with molar conductance of each ion multiplied with the number of ions present in the formula of the electrolyte.

Fuel cells are the galvanic cells in which the energy of combustion of fuels like hydrogen, methanol, etc., is directly converted into electrical energy.

The electrode at which the oxidation occur is called anode.

The electrode at which the reduction occur is called cathode.

The limiting molar conductivity of an electrolyte (i.e., molar conductivity at infinite dilution) is the sum of the limiting ionic conductivities of the cation and the anion, each multiplied by the number of ions present in one formula unit of electrolyte.

\[\Lambda_{m}\operatorname{for}A_{x}B_{y}=x\mathrm{l}_{(A)}{}^{0y+}+yl_{(B)}{}^{0x-}\]

Where l^°_A  and l^°_B are the limiting ionic conductivities of the cation and anion.

Faraday's laws of electrolysis are quantitative laws that describe the effects of electrolysis. They are:

First Law: The mass of any substance deposited or liberated at any electrode is directly proportional to the quantity of electricity passed through the electrolyte.

W ∝ Q 

W = ZQ (Z = electrochemical equivalent)

W = ZIt

\[(Z=\frac{\text{Equivalent Weight}}{96500})\]

Second Law: When the same quantity of electricity is passed through solutions of different electrolytes connected in series, the weight of substances produced at the electrodes is directly proportional to their equivalent weights.

\[\frac{\text{Weight of }M_1\mathrm{~Deposited}}{\text{Weight of }M_2\mathrm{~Deposited}}=\frac{\text{Eq. wt. of }M_1}{\text{Eq. wt. of }M_2}\cdots\cdots\]

Redox Reactions:

• A substance that oxidises another substance (and is itself reduced) is called an oxidising agent.
• A substance that reduces another substance (and is itself oxidised) is called a reducing agent.

What is Oxidation and Reduction?

Electrical conductance and resistance:

\[\mathrm{K}=\mathrm{G}\frac{l}{A}\]

K = Conductivity

G = Conductance

\[\mathrm{G}=\mathrm{}\frac{1}{R}\]

R = Resistance

\[\mathrm{K}=\mathrm{}\frac{l}{RA}\]

Rules for Assigning Oxidation Numbers:

• Oxidation number of N can be −3 (bonded to less electronegative atoms) or +3 (bonded to more electronegative atoms)
• Oxidation number of halogens is always −1 in metal halides
• In interhalogen compounds, the more electronegative halogen gets the oxidation number of −1
• Oxidation number of metals in amalgams and carbonyls is zero (e.g., Fe in [Fe(CO)₅] = 0)
• In complex ions, the algebraic sum of oxidation numbers of all atoms = net charge on the ion
• Oxidation number can be positive, negative, zero, a whole number, or a fraction
• Oxidation number greater than +6 or less than −4 is unusual — double-check for errors

Stock Notation

Variable oxidation states are indicated using Roman numerals in parentheses after the element symbol:

Two methods are used to balance redox reactions:

Method 1: Oxidation Number Method

The change in oxidation number is used to balance electron gain and loss.

Steps (Acidic Medium):

1. Write the skeleton equation; balance all atoms except O and H first
2. Identify which atoms change oxidation number; calculate the net increase and decrease
3. Multiply coefficients to make total increase in oxidation number = total decrease
4. Balance O atoms by adding H₂O to the side with fewer O atoms
5. Balance H atoms by adding H⁺ ions
6. Check that all atoms and charges are balanced

Method 2: Ion Electron Method (Half-Reaction Method)

The reaction is split into two half-reactions (oxidation and reduction) which are balanced separately and then combined.

Steps:

1. Write the redox reaction in ionic form
2. Split into oxidation half-reaction and reduction half-reaction
3. Balance atoms in each half-reaction (except O and H first)
4. Balance O by adding H₂O; balance H by adding H⁺ (acidic) or OH⁻ (basic)
5. Balance charge by adding electrons to the appropriate side
6. Equalise electrons transferred — multiply one or both half-reactions by suitable factors so electrons cancel
7. Add both half-reactions; cancel identical species on both sides
8. Check that all atoms and charges are balanced

For Basic Medium (Ion Electron Method):
After balancing in acidic conditions:

• Add OH⁻ ions to both sides equal to the number of H⁺ ions
• H⁺ + OH⁻ → H₂O (combine)
• Eliminate H₂O molecules appearing on both sides
• Final check: all elements and charges must balance

Two methods are used to balance redox reactions:

Method 1: Oxidation Number Method

The change in oxidation number is used to balance electron gain and loss.

Steps (Acidic Medium):

1. Write the skeleton equation; balance all atoms except O and H first
2. Identify which atoms change oxidation number; calculate the net increase and decrease
3. Multiply coefficients to make total increase in oxidation number = total decrease
4. Balance O atoms by adding H₂O to the side with fewer O atoms
5. Balance H atoms by adding H⁺ ions
6. Check that all atoms and charges are balanced

Method 2: Ion Electron Method (Half-Reaction Method)

The reaction is split into two half-reactions (oxidation and reduction) which are balanced separately and then combined.

Steps:

1. Write the redox reaction in ionic form
2. Split into oxidation half-reaction and reduction half-reaction
3. Balance atoms in each half-reaction (except O and H first)
4. Balance O by adding H₂O; balance H by adding H⁺ (acidic) or OH⁻ (basic)
5. Balance charge by adding electrons to the appropriate side
6. Equalise electrons transferred — multiply one or both half-reactions by suitable factors so electrons cancel
7. Add both half-reactions; cancel identical species on both sides
8. Check that all atoms and charges are balanced

For Basic Medium (Ion Electron Method):
After balancing in acidic conditions:

• Add OH⁻ ions to both sides equal to the number of H⁺ ions
• H⁺ + OH⁻ → H₂O (combine)
• Eliminate H₂O molecules appearing on both sides
• Final check: all elements and charges must balance

Electrolysis of NaCl

1. Molten NaCl:

• Oxidation: Cl⁻ → Cl₂ (gas)

• Reduction: Na⁺ → Na (metal)

• Products: Na (cathode), Cl₂ (anode)

2. Aqueous NaCl:

• Oxidation: Cl⁻ → Cl₂

• Reduction: H₂O → H₂ + OH⁻

• Products: H₂ (cathode), Cl₂ (anode), NaOH formed

The Nernst equation is used to calculate the electrode or cell potential under non-standard conditions.

\[E_{cell}=E_{cell}^\circ-\frac{RT}{nF}\ln Q\]

At 25°C, it becomes:

\[E_{cell}=E_{cell}^\circ-\frac{0.0591}{n}\log Q\]

Where E°cell is the standard cell potential, n is the number of electrons transferred, and Q is the reaction quotient.

The equation helps determine the direction and spontaneity of a reaction:

• Ecell > 0 → spontaneous
• Ecell = 0 → equilibrium (Q = K)

It also relates to Gibbs energy:

ΔG = −nFE_cell

Thus, the Nernst equation is important for electrochemical calculations and equilibrium analysis.