Electrochemistry Made Simple: Galvanic Cells, Nernst Equation and Faraday's Laws (NEET & JEE)
Two beakers, a wire, a salt bridge, and a voltmeter needle that moves on its own: electrochemistry looks like a lab demo, yet it hands out some of the most predictable marks in Class 12 Chemistry. Once you see that every question is just one of three formulas wearing a different costume, the chapter stops feeling heavy.
In every cell, oxidation happens at the anode and reduction at the cathode.
E°cell = E°cathode − E°anode, always using reduction potentials.
The Nernst equation tells you the voltage once concentrations are no longer 1 M.
ΔG° = −nFE° links voltage to spontaneity and equilibrium.
Electrolysis numericals reduce to charge = current × time and moles of electrons.
What Is a Galvanic Cell?
A galvanic cell converts the energy of a spontaneous redox reaction into electrical energy. The two half-reactions sit in separate beakers, connected by an external wire for electrons and a salt bridge for ions.
The anode is where oxidation occurs and electrons are released; the cathode is where reduction occurs and electrons are consumed. Electrons flow in the wire from anode to cathode, while current is conventionally drawn the opposite way. The salt bridge keeps each beaker electrically neutral so the reaction can continue.
Calculating Cell Potential Without Mistakes
The standard cell potential is E°cell = E°cathode − E°anode, and both values come from the table of standard reduction potentials. Never flip the sign of the anode value yourself; the subtraction already does it.
A positive E°cell means the reaction is spontaneous as written. If you get a negative number, you have either swapped the electrodes or the reaction runs the other way. Also remember that E° is an intensive property: multiplying a half-reaction by 2 does not double its potential, although it does change n.
Pro Tip: Before touching numbers, write the two half-reactions and ask which species is the stronger oxidising agent. The one with the higher reduction potential is the cathode. Students who drill this one-line decision until it is automatic rarely lose marks on sign errors, which is the commonest slip in this chapter.
The Nernst Equation: When Concentrations Are Not 1 M
Standard potentials assume 1 M solutions at 298 K. For anything else, use E = E° − (0.059/n) log Q at 298 K, where Q is the reaction quotient built from concentrations of ions (pure solids and metals are left out).
As the cell runs, Q rises and E falls. When E = 0 the cell is dead and has reached equilibrium, which gives the useful link E° = (0.059/n) log K. In a concentration cell, both electrodes are the same metal, so E° = 0 and the voltage comes entirely from the concentration difference.
A quick sense check helps: if the product-side concentration is higher than standard, Q is greater than 1 and the voltage should drop below E°. If your answer rises, recheck whether you inverted Q.
Linking Voltage, Gibbs Energy and Equilibrium
The master relation is ΔG° = −nFE°, with F ≈ 96,500 C per mole of electrons. It converts a voltage into the maximum useful work the cell can deliver, and it connects to K through ΔG° = −RT ln K.
So a positive E° means negative ΔG° and K greater than 1. Because ΔG depends on n, doubling the balanced equation doubles ΔG° even though E° stays put. That contrast, E° is intensive but ΔG° is extensive, is a favourite trap in assertion-reason questions.
Electrolysis and Faraday's Laws
In an electrolytic cell, an external source forces a non-spontaneous reaction, and the polarity flips in meaning: the anode is still where oxidation happens, but it is now the positive terminal. Faraday's first law says the mass deposited is proportional to the charge passed.
In practice: charge Q = I × t, moles of electrons = Q / 96,500, and moles of metal = moles of electrons / n. For a mixture of ions or water, the species with the higher reduction potential is discharged first at the cathode, and the one with the lower potential is oxidised first at the anode. Check for overpotential and electrode type only when a question hints at it.
A Study Routine for This Chapter
Memorise the five or six core formulas once, with units, on a single page.
Solve ten cell-potential problems where only the identity of the electrodes changes.
Solve ten Nernst problems, deliberately varying n and the concentration ratio.
Finish with Faraday numericals, then log every error with its cause in your error log.
Electrochemistry pairs well with conductance (Kohlrausch's law, molar conductivity), which sits in the same unit. Revise them together so the formulas do not blur.
Frequently Asked Questions
Is electrochemistry important for NEET and JEE Main?
Yes. It sits in the Class 12 physical chemistry portion and usually contributes a few questions across both exams, mixing conceptual and numerical types. Because its formulas are few and fixed, it is one of the safer chapters to score in.
What is the difference between a galvanic and an electrolytic cell?
A galvanic cell produces electricity from a spontaneous reaction, with the cathode as the positive terminal. An electrolytic cell uses external electricity to drive a non-spontaneous reaction, and the anode is positive. Oxidation is at the anode in both.
When should I use the Nernst equation?
Use it whenever concentrations or pressures differ from standard conditions, or when a question asks for the potential at a particular stage of a running cell. At 298 K the simplified form with 0.059/n is usually expected.
Why is E° of a half-reaction not multiplied when I double the equation?
Potential is a measure of energy per unit charge, so it does not scale with the amount of reaction. Only n and ΔG° change when you multiply the equation.
How do I know which ion is discharged first in electrolysis?
At the cathode, the species with the higher reduction potential is reduced first. At the anode, the species with the lower reduction potential, which is easier to oxidise, is oxidised first.
Have a doubt on Nernst or Faraday problems? Leave your prep questions in the comments, and share this post with a fellow aspirant who keeps losing marks on signs.
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