Electrochemistry
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Direct answer
Electrochemistry converts chemical energy to electrical energy in a galvanic cell, whose driving force is the EMF: E°cell = E°cathode - E°anode using standard reduction potentials, and a positive E°cell marks a spontaneous cell because delta G° = -n F E°cell. The Nernst equation, E_cell = E°cell - (0.059/n) log Q at 25 °C, tracks EMF as concentrations change, while conductance and Faraday's laws handle electrolysis, where an external current drives a non-spontaneous change.
What you must remember
- E°cell = E°cathode - E°anode; for the zinc-copper Daniell cell, E°cell = 0.34 - (-0.76) = 1.10 V; galvanic cells require a positive E°cell.
- delta G° = -n F E°cell with F = 96500 C per mole, and delta G° = -2.303 R T log K, linking EMF to the equilibrium constant.
- Nernst at 25 °C: E_cell = E°cell - (0.059/n) log Q; for a metal electrode M^n+ + ne → M, E = E° + (0.059/n) log[M^n+]; concentration cells give E = (0.059/n) log(c2/c1).
- Molar conductivity lambda_m = 1000 × kappa / c; on dilution kappa falls but lambda_m rises; for a weak electrolyte alpha = lambda_m / lambda°_m.
- Kohlrausch's law: lambda°_m = nu+ lambda°+ + nu- lambda°-, assembling limiting molar conductivities (as for acetic acid) from strong-electrolyte data.
- Faraday's first law: mass deposited = (equivalent mass × I × t)/96500; one faraday deposits one mole of a monovalent metal.
- Lead-acid battery: Pb and PbO2 in H2SO4; discharge forms PbSO4 on both plates and consumes acid, so density falls; about 2 V per cell. Dry cell is primary; Ni-Cd and lead-acid secondary; the H2-O2 fuel cell gives electricity directly with water as product.
- Corrosion is electrochemical: iron dissolves at anodic spots while oxygen is reduced nearby; galvanising or a sacrificial magnesium anode protects it.
Common confusion
The commonest error is electrode bookkeeping: a galvanic cell's anode is negative and cathode positive, an electrolytic cell's anode is positive — oxidation always sits at the anode either way. Students invert the Nernst log or drop its minus sign, and mix the dilution trends: conductivity falls (fewer ions per volume) while molar conductivity rises. In Faraday problems, substituting molar for equivalent mass is the standard slip.
Exam-focused takeaway
JEE Main tests the toolkit numerically: EMF from standard potentials, Nernst calculations, mass deposited by a current, conductivity conversions. JEE Advanced prefers reasoning on the same laws — concentration cells, Kohlrausch applications to weak electrolytes, electrolysis of brine or copper sulphate with products at both electrodes, and corrosion written as paired half-reactions. Write both half-reactions and the electron direction before computing.
Frequently asked questions
What does the Nernst equation calculate?
The cell EMF under non-standard conditions: E_cell = E°cell - (0.059/n) log Q at 25 °C.
How are E°cell and K related?
Through delta G° = -nFE°cell = -2.303 RT log K, so a large positive EMF means a huge equilibrium constant.
Why does molar conductivity increase on dilution?
Weak electrolytes ionise more completely and interionic attraction weakens, so conducting power per mole rises even as solution conductivity falls.
What is Kohlrausch's law used for?
Adding limiting ionic molar conductivities of the constituent ions gives lambda°_m of any electrolyte — the indirect route for weak electrolytes like acetic acid.
What are the products of electrolysis of aqueous NaCl?
Hydrogen at the cathode, chlorine at the anode, sodium hydroxide left behind — the chlor-alkali outcome.
Why is the anode negative in a galvanic cell but positive in electrolysis?
Both anodes host oxidation; a galvanic anode feeds electrons to the circuit, while in electrolysis the supply pulls electrons out of it.