Schooling · Class 12 · Chemistry · Electrochemistry · Chapter on a page
CBSE Class 12 · Chemistry · NCERT Chemistry-I
Electrochemistry: chapter on a page
A cut of Shishya's notes for this chapter: the first paragraph of the introduction and the key ideas, formulas and facts, common mistakes and quick revision lists, word for word. Nothing is added or reworded; the worked examples are left out (they are on the chapter page).
About the notes: Written by Shishya's AI on 27 Sept 2026 from the chapter's title and class level, in Shishya's own words — not a copy or summary of the textbook. Read the official chapter for the book's own text, activities and exercises.
This note's current text has no applied fact-check on record.
Mind map
Drawn from the same notes by a fixed rule (no AI): each label is the item's own words — its bold term, its label before a colon, a short formula as written, or its first words.
The mind map as a list
- Key ideas
- A galvanic (voltaic) cell converts…
- The standard electrode potential…
- The standard emf of a cell is…
- The Nernst equation relates cell emf…
- Gibbs energy change and emf are…
- Faraday's first law states that the…
- Faraday's second law states that for…
- Formulas and facts
- Standard cell emf
- Nernst equation at 298 K
- ΔG° = −nFE°
- E = 0
- m = ZIt
- Λₘ = κ / c
- Kohlrausch's law
- Worked examples
- Calculating the emf of a Daniell…
- Mass deposited using Faraday's law
- Using Kohlrausch's law
- Watch out for
- Adding electrode potentials instead…
- Forgetting to convert time to…
- Using concentration of solid or pure…
- Confusing conductivity (κ) with…
- Ignoring the sign of E° when…
What this chapter is about
Electrochemistry studies the relationship between chemical energy and electrical energy. It explains how chemical reactions can produce electric current (as in a battery) and how electric current can drive non-spontaneous chemical reactions (as in electroplating or the extraction of metals). This chapter builds on your understanding of redox reactions from earlier classes and connects thermodynamic ideas like Gibbs energy to the measurable quantity called electromotive force (emf).
Key ideas
- A galvanic (voltaic) cell converts chemical energy into electrical energy through a spontaneous redox reaction; an electrolytic cell uses external electrical energy to drive a non-spontaneous reaction.
- The standard electrode potential (E°) of a half-cell is measured against the Standard Hydrogen Electrode (SHE), which is assigned a potential of exactly 0 V at 298 K, 1 bar H₂ and 1 M H⁺.
- The standard emf of a cell is E°(cell) = E°(cathode) − E°(anode); a positive value means the reaction is spontaneous under standard conditions.
- The Nernst equation relates cell emf to ion concentrations: E = E° − (RT / nF) ln Q, or at 298 K, E = E° − (0.0591 / n) log₁₀ Q, where Q is the reaction quotient and n is the number of electrons transferred.
- Gibbs energy change and emf are linked by ΔG = −nFE; at equilibrium, E = 0 and ΔG = 0.
- Faraday's first law states that the mass of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed.
- Faraday's second law states that for the same quantity of electricity, the masses of different substances deposited are proportional to their equivalent masses (molar mass / number of electrons involved).
- Conductivity (κ) of an electrolyte solution and molar conductivity (Λₘ = κ / c, where c is concentration in mol/m³) describe how well a solution conducts; Λₘ increases on dilution for both strong and weak electrolytes.
Formulas and facts to remember
- Standard cell emf: E°(cell) = E°(cathode) − E°(anode).
- Nernst equation at 298 K: E = E° − (0.0591 V / n) log₁₀ Q.
- Relation of Gibbs energy to emf: ΔG° = −nFE°, where F = 96485 C mol⁻¹ (Faraday constant).
- At equilibrium: E = 0, so E° = (0.0591 / n) log₁₀ K (K is equilibrium constant).
- Faraday's first law: m = ZIt, where Z = M / nF (electrochemical equivalent), I is current (A), t is time (s).
- Molar conductivity: Λₘ = κ / c; unit is S cm² mol⁻¹ when κ is in S cm⁻¹ and c in mol cm⁻³.
- Kohlrausch's law: The limiting molar conductivity of an electrolyte equals the sum of limiting molar conductivities of its cation and anion: Λ°ₘ = λ°₊ + λ°₋.
- One faraday of electricity (96485 C) deposits one mole of a monovalent ion or half a mole of a divalent ion.
Common mistakes
- Adding electrode potentials instead of subtracting → always use E°(cell) = E°(cathode) − E°(anode).
- Forgetting to convert time to seconds when using Faraday's law → convert hours or minutes to seconds before calculating charge.
- Using concentration of solid or pure liquid in the reaction quotient Q → only aqueous or gaseous species appear; activities of pure solids and liquids are taken as 1.
- Confusing conductivity (κ) with molar conductivity (Λₘ) → κ is for the solution, Λₘ is per mole of electrolyte.
- Ignoring the sign of E° when identifying anode and cathode → the electrode with lower (more negative) E° acts as the anode in a galvanic cell.
Quick revision
- Galvanic cell: spontaneous reaction produces current; electrolytic cell: current drives non-spontaneous reaction.
- E°(cell) = E°(cathode) − E°(anode); positive E° means spontaneous.
- Nernst equation at 298 K: E = E° − (0.0591/n) log Q.
- ΔG° = −nFE° links thermodynamics to electrochemistry.
- Mass deposited = (M × I × t) / (n × F) (Faraday's first law rearranged).
- Kohlrausch's law: Λ°ₘ = λ°₊ + λ°₋; useful for weak electrolytes.