Revision: Redox Reactions and Electrochemistry JEE Main 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 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.

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

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.

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

"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.

"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 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.

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.

The electrode at which the oxidation occur is called anode.

The electrode at which the reduction occur is called cathode.

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

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.

It is the inverse of resistance R and may be simply defined as the speed through which current flows in a conductor.

c = `1/R = A/(pl)`

k = `A/l`

Here k is the specific conductance. The SI unit of conductance is Siemens, which is denoted by the symbol ‘S’ and is equal to ohm⁻¹ or Ω⁻¹.

Molar conductivity is the conductance of a volume of solution containing 1 mole of dissolved electrolyte when placed between two parallel electrodes 1 cm apart and large enough to contain between them all the solution.

The conductivity, which is shown by all the ions when 1 mol of electrolyte is dissolved in the solution, is called molar conductivity; it is expressed by ∧_m (lambda). If 1 mol of electrolyte is present in V_m cm³ of electrolyte solution, then ∧_m = κ × V

= `(kappa xx 1000)/"Molarity" = (kappa xx 1000)/M`

Its unit is ohm⁻¹ cm² mol⁻¹ or S cm² mol⁻¹.

The limiting molar conductivity of an electrolyte is defined as its molar conductivity when the concentration of the electrolyte in the solution approaches zero.

When the concentration of an electrolytic solution placed between electrodes of a conductivity cell placed at a unit distance having an area of cross-section sufficient to accommodate enough volume of solution containing one mole of electrolyte approaches zero, then the conductance of the solution is known as limiting molar conductivity.

It is the process of decomposition of an electrolyte in the molten or aqueous state by discharge of ions at the electrodes on the passage of an electric current.

Electrolysis :
It is the process of decomposition of the electrolyte in the molten or aqueous state by discharge of ions at the elctrodes on passage of an electric current.

It is the process of decomposition of a chemical compound in aqueous solutions or in a molten state accompanied by a chemical change using direct electric current. 

Electroplating is a process in which a thin film of a metal like gold, silver, nickel, chromium, etc. gets deposited on another metallic article with the help of electricity.

A fuel cell is a galvanic cell in which the reactants are not placed within the cell, but are continuously supplied from outside, where one reactant acts as a fuel (such as hydrogen or methanol) and the other as an oxidant (such as oxygen).

Corrosion is the gradual damage of metals caused by their reaction with components of the atmosphere, such as oxygen and moisture.

Kohlrausch law states that at infinite dilution of the solution, each ion of electrolyte migrates independently of its co-ions and contribute independently to the total molar conductivity irrespective of the nature of other ion.

Kohlrausch’s law states that the molar conductivity of an electrolyte at infinite dilution is the same as the sum of the anions' and cations' limited molar conductivities.

`∧_m^° = v_+  λ_+^° + v_-  λ_-^°`

Here `λ_+^°` and `λ_-^°` are limiting molar conductivities of cations and anions.

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?

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:

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

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.

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.

Electrical conductance and resistance:

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

K = Conductivity

G = Conductance

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

R = Resistance

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

Reactions

• Anode:
  2H₂ + 4OH⁻ → 4H₂O + 4e⁻
• Cathode:
  O₂ + 4H₂O + 4e⁻ → 4OH⁻
• Overall reaction:
  2H₂ + O₂ → 2H₂O

Applications

• Spacecraft (electric power)
• Power generators (homes, hospitals)
• Automobiles (experimental)
• Clean energy for industries

Drawbacks

• Hydrogen gas is hazardous
• High cost of hydrogen preparation