Electrochemistry Made Simple: Cells, EMF, Nernst and Electrolysis
A friendly, NCERT-aligned walkthrough of Class 12 Electrochemistry: how galvanic cells turn reactions into electricity, standard electrode potentials and the SHE, calculating cell EMF, the Nernst equation with a worked numerical, conductance of solutions, and electrolysis with Faraday's laws.
By the PadhoDost Team · 📖 8 min read · Updated 4 August 2026
Part of Class 12 (CBSE) prep🧠 A dam that lights up your village
Imagine two ponds at different heights connected by a channel. Water rushes from the higher pond to the lower one, and on its way it spins a turbine that lights up your village. In electrochemistry, electrons are the water and the 'height difference' is the electrical push we call potential. A galvanic cell is simply a clever setup that lets electrons flow downhill through a wire, doing useful work like running your calculator or torch, instead of just fizzing away uselessly.
Padho, dost! Electrochemistry is the study of the give-and-take between chemical reactions and electricity. Some reactions release energy that we capture as electric current (that is a galvanic or voltaic cell, like a battery). Others need electricity pushed into them to happen (that is electrolysis). The whole chapter rests on one simple idea from redox: oxidation is loss of electrons, reduction is gain of electrons. Every reaction here is just electrons changing owners.
Galvanic cells: making electricity from a reaction
Take the classic Daniell cell. A zinc rod sits in ZnSO4 solution, a copper rod sits in CuSO4 solution, and a salt bridge connects the two beakers. Zinc is more reactive, so it gives up electrons: Zn goes to Zn2+ + 2e-. This is oxidation, and it happens at the ANODE (the negative terminal in a galvanic cell). Those electrons travel through the external wire to the copper, where Cu2+ + 2e- becomes Cu. This is reduction, at the CATHODE (the positive terminal). The salt bridge (often KCl or KNO3 set in agar jelly) keeps both solutions electrically neutral by letting ions flow, completing the circuit.
Standard electrode potential and cell EMF
Every electrode has a tendency to get reduced, measured as its standard reduction potential (E-degree) in volts, under standard conditions: 298 K, 1 bar pressure for any gas, and 1 M concentration for ions in solution. We cannot measure a single electrode alone, so we compare everything to the Standard Hydrogen Electrode (SHE), whose potential is defined as exactly 0.00 V. A more positive E-degree means the species is more easily reduced. The standard cell potential is then found from the two electrode potentials.
| Electrode (reduction half-reaction) | E-degree / V |
|---|---|
| Zn2+ + 2e- to Zn | -0.76 |
| 2H+ + 2e- to H2 (SHE) | 0.00 |
| Cu2+ + 2e- to Cu | +0.34 |
| Ag+ + e- to Ag | +0.80 |
| F2 + 2e- to 2F- | +2.87 |
For the Daniell cell: E-degree(cell) = E-degree(Cu) - E-degree(Zn) = (+0.34) - (-0.76) = +1.10 V. A positive cell potential means the reaction is spontaneous (feasible), which links to Gibbs energy through delta-G-degree = -nF E-degree(cell), where n is the moles of electrons transferred and F is Faraday's constant, 96500 C/mol. A positive E-degree therefore gives a negative delta-G-degree, confirming the cell works on its own.
The Nernst equation: when conditions are not standard
Real cells rarely have exactly 1 M solutions, and the voltage changes as the cell discharges and concentrations shift. The Nernst equation corrects the standard potential for the actual concentrations. At 298 K, plugging in R, T and F and converting to log base 10 gives the handy exam-ready form.
📝 Worked example: Nernst equation for a Daniell cell
Cell: Zn(s) | Zn2+ (0.1 M) || Cu2+ (0.01 M) | Cu(s). Find E(cell) at 298 K.
Overall reaction: Zn + Cu2+ to Zn2+ + Cu, so n = 2 electrons transferred.
Standard EMF: E-degree(cell) = 0.34 - (-0.76) = +1.10 V.
Reaction quotient Q = [Zn2+]/[Cu2+] = 0.1 / 0.01 = 10.
Nernst: E(cell) = 1.10 - (0.0591/2) log(10).
log(10) = 1, so the correction term = (0.0591/2)(1) = 0.02955 V.
E(cell) = 1.10 - 0.0296 = 1.0704 V, approximately 1.07 V.
Note: the solids Zn and Cu do not appear in Q; only ion concentrations do.
Conductance of electrolytic solutions
Electrolytes conduct electricity through moving ions. Conductivity (kappa) is the conductance of a cube of solution 1 cm on each side (that is, unit length and unit cross-sectional area). But to compare different electrolytes fairly, we use molar conductivity, which accounts for concentration. On dilution, molar conductivity INCREASES because the same amount of electrolyte spreads into more solution and its ions move more freely. For strong electrolytes (like KCl) it rises only slightly and follows Debye-Huckel-Onsager behaviour; for weak electrolytes (like acetic acid) it shoots up steeply near infinite dilution as more molecules ionise.
Electrolysis and Faraday's laws
In electrolysis we FORCE a non-spontaneous reaction by pushing current from an external source. Here the anode is positive and the cathode is negative (the signs flip versus a galvanic cell), but oxidation still occurs at the anode and reduction at the cathode. Faraday's first law says the amount of substance deposited or liberated is directly proportional to the charge passed. Charge Q = current (I) x time (t), and 1 mole of electrons carries 96500 C.
📝 Worked example: Faraday's first law
Question: A current of 5 A flows for 30 minutes through CuSO4 solution. How much copper is deposited at the cathode? (Molar mass of Cu = 63.5 g/mol)
Step 1 - charge: Q = I x t = 5 A x (30 x 60 s) = 5 x 1800 = 9000 C.
Step 2 - electrons: Cu2+ + 2e- to Cu, so 2 mol electrons (2 x 96500 = 193000 C) deposit 1 mol (63.5 g) of Cu.
Step 3 - proportion: mass = (63.5 / 193000) x 9000.
Step 4: mass = 63.5 x 9000 / 193000 = 571500 / 193000 = 2.96 g of copper.
Answer: about 2.96 g of copper is deposited at the cathode.
Quick revision checklist
- ✓Galvanic cell = spontaneous reaction makes electricity; anode is negative, cathode is positive.
- ✓Electrolytic cell = electricity forces a reaction; anode is positive, cathode is negative. In both, oxidation is at the anode.
- ✓E-degree(cell) = E-degree(cathode) - E-degree(anode); a positive value means spontaneous, and delta-G-degree = -nF E-degree(cell).
- ✓SHE is the reference electrode with E-degree = 0.00 V.
- ✓Nernst at 298 K: E = E-degree - (0.0591/n) log Q; voltage falls as products build up.
- ✓Molar conductivity rises on dilution; use Kohlrausch's law to get a weak electrolyte's limiting value.
- ✓Faraday: Q = I x t, and 96500 C = 1 mole of electrons.
⚡ Quick check
In the Daniell cell Zn | Zn2+ || Cu2+ | Cu, if the concentration of Cu2+ ions is DECREASED (keeping Zn2+ the same), what happens to the cell EMF?
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