Electrochemistry Class 12 Notes | CBSE Chemistry Revision

Key Definitions in Electrochemistry

Electrochemistry

The branch of chemistry that deals with the study of production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.

Electrochemical Cell (Galvanic/Voltaic Cell)

A device that converts chemical energy released during a spontaneous redox reaction into electrical energy.

Electrolytic Cell

A device that uses electrical energy to drive a non-spontaneous redox reaction.

Electrode Potential

The potential difference developed between an electrode and the electrolyte solution when it is in contact with its own ions.

Standard Electrode Potential (E°)

The electrode potential measured at standard conditions (1 M concentration for ions, 1 atm pressure for gases, 298 K temperature).

Standard Hydrogen Electrode (SHE)

A reference electrode with a standard electrode potential arbitrarily assigned as zero volts, used to measure the standard electrode potentials of other electrodes.

Nernst Equation

An equation that relates the electrode potential or cell potential to the concentrations of the species involved in the electrode reaction at any given temperature. E = E° - (RT/nF)lnQ or E = E° - (0.0592/n)logQ at 298 K.

Molar Conductivity (Λm)

The conductivity of an electrolyte solution containing one mole of the electrolyte placed between two electrodes separated by unit distance and having sufficient area to contain all the electrolyte. Λm = κ / C, where κ is conductivity and C is molar concentration.

Limiting Molar Conductivity (Λ°m)

The molar conductivity of an electrolyte when the concentration approaches zero (infinite dilution).

Faraday's Laws of Electrolysis

Laws stating the quantitative relationships between the amount of substance produced/consumed at electrodes and the quantity of electricity passed during electrolysis.

Understanding Electrochemical Cells and Redox Reactions

At the heart of electrochemistry lies the redox reaction, a chemical process involving simultaneous oxidation and reduction. Oxidation is the loss of electrons, while reduction is the gain of electrons. In an electrochemical cell, these half-reactions occur at separate electrodes. The anode is the electrode where oxidation takes place, and it is negatively charged in a galvanic cell. The cathode is where reduction occurs, and it is positively charged in a galvanic cell. Electrons flow from the anode to the cathode through an external circuit, generating electrical current.

A crucial component of a galvanic cell is the salt bridge. This U-shaped tube containing an inert electrolyte (like KCl or KNO₃) allows the migration of ions between the two half-cells, thereby completing the internal circuit and maintaining electrical neutrality. Without a salt bridge, charge would build up in the half-cells, quickly stopping the flow of electrons. The overall cell potential (E_cell) is the difference between the reduction potentials of the cathode and the anode (E_cell = E_cathode - E_anode). A positive E_cell indicates a spontaneous reaction, characteristic of galvanic cells. Conversely, in electrolytic cells, an external power source drives a non-spontaneous reaction, making the anode positive and the cathode negative.

Must Remember: Key Principles for Electrochemistry

• Anode Oxidation, Cathode Reduction: Remember 'AN OX' (Anode Oxidation) and 'RED CAT' (Reduction Cathode).
• Electron Flow: Electrons always flow from Anode to Cathode in the external circuit.
• Salt Bridge Function: Completes the circuit by allowing ion migration and maintains electrical neutrality in half-cells.
• Spontaneity and Gibbs Energy: A spontaneous cell reaction has ΔG° < 0 and E°_cell > 0. The relationship is ΔG° = -nFE°_cell.
• Nernst Equation Significance: Used to calculate cell potential under non-standard conditions (i.e., when concentrations are not 1 M).
• Kohlrausch's Law: For weak electrolytes, limiting molar conductivity can be calculated from the limiting ionic conductivities of its constituent ions (Λ°m = v₊λ°₊ + v₋λ°₋). This is crucial for determining dissociation constants.
• Conductivity vs. Molar Conductivity: Conductivity (κ) is for a specific volume, while molar conductivity (Λm) is for one mole of electrolyte.
• Effect of Dilution: For strong electrolytes, Λm increases slightly with dilution; for weak electrolytes, it increases sharply due to increased dissociation.
• Faraday's Laws: 1st Law: m ∝ Q; 2nd Law: m₁/m₂ = E₁/E₂ (where E is equivalent mass). Q = It (Charge = Current × Time).
• Corrosion: An electrochemical process where metals are oxidized in the presence of moisture and oxygen (e.g., rusting of iron).

Galvanic vs. Electrolytic Cells

Working of a Daniel Cell (A Type of Galvanic Cell)

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2. —
3. —
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Worked Examples & Calculations

• 1. Calculating Standard Cell Potential (E°_cell) Given: E°(Zn²⁺/Zn) = -0.76 V, E°(Cu²⁺/Cu) = +0.34 V. Question: Calculate the standard cell potential for a Daniel cell. Solution: In a Daniel cell, Zn is oxidized (anode) and Cu²⁺ is reduced (cathode). E°_cell = E°_cathode - E°_anode = E°(Cu²⁺/Cu) - E°(Zn²⁺/Zn) E°_cell = (+0.34 V) - (-0.76 V) = 0.34 V + 0.76 V = 1.10 V.
• 2. Applying Nernst Equation Consider the cell: Zn(s) | Zn²⁺(0.1 M) || Cu²⁺(0.01 M) | Cu(s). E°_cell = 1.10 V. Question: Calculate the cell potential (E_cell) at 298 K. Solution: The overall reaction is Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). n = 2. Nernst Equation: E_cell = E°_cell - (0.0592/n)log([Zn²⁺]/[Cu²⁺]) E_cell = 1.10 - (0.0592/2)log(0.1/0.01) E_cell = 1.10 - (0.0296)log(10) E_cell = 1.10 - 0.0296 = 1.0704 V.

Exam Tip: Avoiding Common Traps!

Always pay close attention to the sign conventions for electrode potentials and the stoichiometry (n-value) in the Nernst equation and Faraday's laws. A common mistake is to confuse anode/cathode polarity between galvanic and electrolytic cells. Remember: 'AN OX, RED CAT' holds for both in terms of process, but the sign of the electrode is reversed. In galvanic cells, the anode is negative; in electrolytic cells, the anode is positive (connected to the positive terminal of the battery). Also, ensure you use consistent units for conductivity (S cm⁻¹ or S m⁻¹) and concentration (mol L⁻¹ or mol m⁻³) when calculating molar conductivity.

Practice Questions with Solutions

• Q: What is the main purpose of a salt bridge in an electrochemical cell? A: To complete the electrical circuit and maintain electrical neutrality in the half-cells by allowing ion migration.
• Q: How does molar conductivity change for a strong electrolyte with increasing dilution? A: It increases slightly due to reduced interionic attraction and increased mobility of ions.
• Q: For a spontaneous reaction in an electrochemical cell, what is the sign of ΔG° and E°_cell? A: ΔG° is negative, and E°_cell is positive.
• Q: State Faraday's first law of electrolysis in terms of mass and charge. A: The mass of a substance deposited or liberated at any electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

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