Concept Core
From galvanic cells to electrolysis — the electrochemistry framework.
Galvanic Cell — The Electrochemistry Setup
A galvanic cell converts chemical energy to electrical energy. Key components:
Anode (−): Oxidation occurs. Loses electrons: Zn → Zn²⁺ + 2e⁻
Cathode (+): Reduction occurs. Gains electrons: Cu²⁺ + 2e⁻ → Cu
Salt bridge: maintains electrical neutrality between the two half-cells.
E°cell = E°cathode − E°anode = E°reduction(cathode) − E°reduction(anode)
Standard Electrode Potentials
All E° values are standard reduction potentials (vs SHE at 25°C, 1 M, 1 atm).
Higher E°: better oxidising agent (cathode). Lower E°: better reducing agent (anode).
SHE: E°(H⁺/H₂) = 0 V (by definition)
E°cell = E°(reduction at cathode) − E°(reduction at anode)
Spontaneous cell: E°cell > 0 → ΔG° < 0
Nernst Equation
Cell potential under non-standard conditions:
E_cell = E°cell − (RT/nF) ln Q
E_cell = E°cell − (0.0592/n) log Q (at 25°C)
n = number of electrons transferred. Q = reaction quotient. At equilibrium: E = 0 → E°cell = (0.0592/n) log K
Relationship: ΔG°, E°, and K
ΔG° = −nFE°cell
ΔG° = −RT ln K
nFE°cell = RT ln K
F = Faraday constant = 96485 C/mol ≈ 96500 C/mol
Spontaneous: E°cell > 0 ↔ ΔG° < 0 ↔ K > 1
Electrolysis — Faraday's Laws
First Law: Mass deposited ∝ charge passed: m = ZQ = ZIt
Second Law: Same charge deposits masses in proportion to their equivalent weights.
m = (M × I × t) / (n × F)
where M = molar mass, n = electrons per ion
Conductance & Kohlrausch's Law
Conductance: G = 1/R. Molar conductivity Λm = κ/C (κ = specific conductance).
Kohlrausch's Law: At infinite dilution, Λ°m = Σ λ°(ions). Used to find Λ°m of weak electrolytes from strong ones.
Molar conductivity increases with dilution (more ions dissociate for weak electrolytes).
Formula Vault
All electrochemistry formulas for EAPCET.
Cell EMF
E°cell = E°cat − E°an
Both in standard reduction potential
ΔG° and E°cell
ΔG° = −nFE°cell
F = 96500 C/mol
Nernst Equation
E = E° − (0.0592/n) log Q
At 25°C; Q = reaction quotient
E° and K
log K = nE°/0.0592
At equilibrium E = 0
Faraday's Law
m = MIt/(nF)
M = molar mass; n = electrons per ion
Charge Passed
Q = I × t
Q in coulombs; I in amperes; t in seconds
Specific Conductance
κ = G × (l/A)
l/A = cell constant
Molar Conductivity
Λm = κ × 1000/C
C in mol/L; κ in S/cm
Worked Examples
5 problems — E°cell, Nernst, Faraday, ΔG, and a Daniel cell trap.
EasyCalculate E°cell for Zn-Cu Daniel cell▾
Calculate E°cell for the cell Zn|Zn²⁺||Cu²⁺|Cu. Given E°(Zn²⁺/Zn) = −0.76 V, E°(Cu²⁺/Cu) = +0.34 V.
1
Zn is oxidised (anode), Cu²⁺ is reduced (cathode).
2
E°cell = E°cathode − E°anode = 0.34 − (−0.76) = 0.34 + 0.76 = 1.10 V
✓ E°cell = 1.10 V
EasyFind ΔG° for a cell with E°cell = 1.10 V, n = 2▾
Calculate ΔG° for a galvanic cell with E°cell = 1.10 V and n = 2.
1
ΔG° = −nFE°cell = −2 × 96500 × 1.10 = −2 × 96500 × 1.10
2
= −212300 J/mol = −212.3 kJ/mol
✓ ΔG° = −212.3 kJ/mol (spontaneous)
MediumFind mass of Cu deposited by 2A current for 1 hour▾
Calculate the mass of copper deposited when a current of 2 A is passed through CuSO₄ solution for 1 hour. (M = 64, n = 2)
1
Charge Q = I × t = 2 × 3600 = 7200 C
2
m = MQ/(nF) = 64 × 7200/(2 × 96500) = 460800/193000 = 2.39 g
✓ Mass of Cu deposited = 2.39 g
EAPCET LevelUse Nernst equation: E°cell = 1.10 V, [Zn²⁺] = 0.1 M, [Cu²⁺] = 0.01 M▾
Calculate E_cell for the Daniel cell when [Zn²⁺] = 0.1 M, [Cu²⁺] = 0.01 M. E°cell = 1.10 V, n = 2.
1
Nernst equation: E = E° − (0.0592/n) log Q
2
Q = [Zn²⁺]/[Cu²⁺] = 0.1/0.01 = 10 (products/reactants for cell reaction)
3
E = 1.10 − (0.0592/2) log 10 = 1.10 − 0.0296 × 1 = 1.07 V
✓ E_cell = 1.07 V
Trap QuestionThe species with higher reduction potential is always the cathode — always?▾
In a cell with E°(Ag⁺/Ag) = +0.80 V and E°(Zn²⁺/Zn) = −0.76 V, a student says 'Ag is always the cathode.' Is this correct under all conditions?
1
Under standard conditions: E°cell = 0.80 − (−0.76) = 1.56 V > 0 → Ag is cathode. Correct here.
2
But under non-standard conditions: Nernst equation can change actual E values.
3
If [Ag⁺] is extremely low, E_Ag can drop below E_Zn, reversing the cell direction.
4
Conclusion: At standard conditions, higher E° = cathode. At non-standard conditions, use Nernst equation to find actual E before deciding.
✓ Under standard conditions yes; under non-standard conditions, use Nernst equation to verify
Mistake DNA
4 electrochemistry errors from EAPCET distractor analysis.
🔌
Anode-Cathode Sign Confusion
In galvanic cells, anode is negative (−) and cathode is positive (+). In electrolytic cells, it's reversed.
❌ Wrong
Galvanic cell:
Anode = positive ✗
Cathode = negative ✗
✓ Correct
Galvanic: Anode(−), Cathode(+) ✓
Electrolytic: Anode(+), Cathode(−) ✓
Oxidation always at anode
Memory: 'An Ox, Red Cat' — Oxidation at Anode, Reduction at Cathode. This holds for BOTH cell types.
📐
E°cell = E°anode − E°cathode (Signs Swapped)
The formula is E°cell = E°cathode − E°anode, both as reduction potentials.
❌ Wrong
E°cell = E°anode − E°cathode
= (−0.76) − 0.34 = −1.10 V ✗
(sign error)
✓ Correct
E°cell = E°cathode − E°anode
= 0.34 − (−0.76) = +1.10 V ✓
Spontaneous when positive
Both electrode potentials are quoted as STANDARD REDUCTION POTENTIALS. Subtract the anode from the cathode. Positive result = spontaneous cell.
⏱️
Forgetting to Convert Hours to Seconds in Faraday's Law
Q = It requires I in amperes and t in seconds. Students forget to convert hours/minutes to seconds.
❌ Wrong
2A for 1 hour:
m = M×2×1/(nF) ✗
(t=1, should be 3600)
✓ Correct
t = 1h = 3600 s
Q = 2 × 3600 = 7200 C ✓
m = M×7200/(n×96500) ✓
Faraday's Law: Q = It where t is in SECONDS. Always convert: 1 min = 60 s; 1 hour = 3600 s.
🔢
Nernst Q: Writing Concentration of Solids/Pure Liquids
Pure solids and liquids have activity = 1 and are NOT included in Q for Nernst equation.
❌ Wrong
Zn + Cu²⁺ → Zn²⁺ + Cu:
Q = [Zn²⁺][Cu]/([Zn][Cu²⁺]) ✗
(Zn and Cu are solids)
✓ Correct
Q = [Zn²⁺]/[Cu²⁺] ✓
Solids and pure liquids
are excluded from Q
In Q (and K) expressions, only aqueous ions and gases appear. Solid Zn and Cu metal have activity = 1 — they don't contribute to Q.
Chapter Intelligence
Electrochemistry is heavily numerical — practise Faraday calculations and Nernst equation.
EAPCET Weightage (2019–2024)
Faraday's law (electrolysis)~7 ΔG° and E°cell relation~5 Conductance & Kohlrausch~3 High-Yield PYQ Patterns
E°cell from standard potentialsMass deposited by Faraday's lawNernst equation E calculationΔG° = −nFE° numericallog K from E°cellIdentify anode and cathode
Exam Strategy
- For E°cell: always identify which electrode is anode (lower E°/more negative) and which is cathode (higher E°/more positive). Then E°cell = E°cat − E°an.
- For Faraday's Law: Q = It (seconds!), then m = MQ/(nF). The two most common errors: wrong time units, wrong n value.
- Nernst equation: at 25°C, the coefficient is 0.0592/n. Calculate Q from the cell reaction (products/reactants of the cell reaction, not just the solution concentrations).
- ΔG° = −nFE°cell. If E°cell > 0, ΔG° < 0 → spontaneous. This is a very fast MCQ connection question.
- Links to Equilibrium: nFE° = RT ln K → log K = nE°/0.0592. This bridges electrochemistry and equilibrium.