Electrochemistry Complete Revision: Galvanic Cells, Nernst Equation, Electrolysis, and Conductance

Chemical reactions and electrical energy share a direct, measurable relationship governed by electron transfer and ion mobility. Spontaneous redox processes generate electric current through galvanic arrangements, whereas electrolytic setups consume external electrical energy to force non-spontaneous transformations.

Mastering standard electrode potentials, the Nernst equation, Faraday's laws, and electrolytic conductance reveals how chemical equilibria translate into measurable voltage and current. These foundational principles govern everything from industrial electrolysis and corrosion prevention to battery design and analytical solution chemistry.

What is Electrochemistry?

Chemical reactions and electrical energy share a direct, measurable relationship that depends on electron transfer and ion movement. Spontaneous redox processes generate electric current in galvanic cells, while electrolytic cells use external electrical energy to force non-spontaneous changes. Standard electrode potentials, the Nernst equation, Faraday's laws and electrolytic conductance explain how chemical equilibria turn into voltage and current. These ideas are used in industrial electrolysis, corrosion prevention, battery design and solution chemistry.

Types Of Electrochemical Cells

Electrochemical cells are divided into galvanic and electrolytic systems based on whether a spontaneous redox reaction produces electrical energy or external electric current drives a chemical change. 

In every redox cell:

• Oxidation means loss of electrons.

• Reduction means gain of electrons.

• Oxidation occurs at the anode.

• Reduction occurs at the cathode.

Galvanic Cell And Daniell Cell

A galvanic cell converts chemical energy into electrical energy. The Daniell cell is a zinc-copper galvanic cell containing:

• Zinc electrode: Zn(s) | Zn2+(aq)

• Copper electrode: Cu2+(aq) | Cu(s)

At the zinc electrode:

Zn(s) becomes Zn2+(aq) and releases two electrons.

This is oxidation. The zinc rod becomes thinner as zinc ions enter the solution.

At the copper electrode:

Cu2+(aq) accepts two electrons and forms Cu(s).

This is reduction. Copper deposits on the copper rod, making it thicker.

Electrons flow through the external wire from the zinc electrode to the copper electrode.

Net Cell Reaction

Write the anodic and cathodic reactions, balance the electrons, and add them:

Zn(s) + Cu2+(aq) gives Zn2+(aq) + Cu(s)

The electrons cancel because two electrons are lost and two are gained.

Electrodes And Their Types

An electrode is an arrangement at which oxidation or reduction occurs. A galvanic cell generally contains two electrodes.

Metal-Metal Ion Electrode

A solid metal rod is immersed in a solution containing ions of the same metal.

Examples:

• Zn(s) | Zn2+(aq)

• Cu(s) | Cu2+(aq)

(*Memory Tip: A metal rod is placed in a solution of its own ions.)*

Gas-Gas Ion Electrode

A gas is passed into a solution containing ions related to that gas. Platinum is used because it is chemically inert, conducts electrons, and provides a solid surface.

For a hydrogen electrode:

H2(g) | H+(aq) | Pt

Platinum does not participate in the reaction.

Metal-Metal Insoluble Salt Electrode

This electrode contains:

1. A metal

2. A layer of its insoluble salt

3. A solution containing the salt’s anion

Example:

Ag | AgCl | Cl-

(*Memory Tip: Metal + its insoluble salt + solution containing the salt’s anion.)*

Redox Electrode

A redox electrode contains ions of the same element in different oxidation states.

Example:

• Fe2+

• Fe3+

The conversion may occur as:

Fe2+ loses one electron to form Fe3+, or Fe3+ gains one electron to form Fe2+.

A platinum wire provides the surface and transfers electrons.

(*Memory Tip: The same element must be present in two or more oxidation states.)*

Anode, Cathode, And Salt Bridge

In an electrochemical cell, redox half-reactions are split between the anode and cathode, while a salt bridge preserves electrical neutrality by facilitating counter-ion migration. 

Functions Of A Salt Bridge

• Maintains electrical neutrality.

• Completes the internal circuit.

• Prevents direct mixing of the two solutions.

• Allows ionic movement between the half-cells.

Suitable electrolytes include KCl, KNO3, NH4Cl, and NH4NO3. Their cations and anions should have nearly similar ionic mobilities.

At the zinc half-cell, negative ions from the salt bridge neutralise the excess positive charge produced by Zn2+. At the copper half-cell, positive ions neutralise the excess negative charge formed when Cu2+ is consumed.

A separate salt bridge is not required when both half-cells contain the same electrolyte and common ions.

Electron And Current Flow

An electrochemical circuit completes its loop through complementary charge carriers, driving electrons through the external wire and migrating ions across the internal solution in reverse conventional directions. 

(*Memory Tip: In the external circuit, electrons move from anode to cathode, while current moves in the opposite direction.)*

Cell Representation And EMF

A Daniell cell is represented as:

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

The single line shows a phase boundary, while the double line represents the salt bridge.

The electromotive force, or EMF, is the potential difference between the electrodes under open-circuit conditions, when no current flows.

Ecell = Ecathode - Eanode

When reduction potentials are given:

E°cell = E°cathode, reduction - E°anode, reduction

Both values must be used as reduction potentials.

Electrode Potential And Standard Hydrogen Electrode

When a metal is placed in a solution containing its ions, a potential difference develops between the metal and the solution. This is called electrode potential.

Oxidation and reduction potentials have equal magnitudes but opposite signs:

Eoxidation = -Ereduction

A higher reduction potential indicates a greater tendency to undergo reduction. A higher oxidation potential indicates a greater tendency to undergo oxidation.

The standard hydrogen electrode, or SHE, is the reference electrode. Under standard conditions:

• Ion concentration is 1 M.

• Gas pressure is approximately 1 bar or 1 atm.

• Temperature is 298 K.

By convention:

E°SHE = 0 V

When two reduction potentials are compared:

• Higher reduction potential means cathode.

• Lower reduction potential means anode.

Nernst Equation

For non-standard conditions at 298 K:

E = E° - (0.059/n) log Q

Here:

• E is the non-standard potential.

• E° is the standard potential.

• n is the number of electrons transferred.

• Q is the reaction quotient.

For a complete cell, write the net reaction first and then construct Q.

Pure solids and pure liquids are omitted from Q because their activities are taken as 1. Aqueous ions are represented by their concentrations, and gases by their pressures.

Equilibrium Constant

At equilibrium:

• Ecell = 0

• Q = K

Therefore:

E°cell = (0.059/n) log K

This relation connects standard EMF with the equilibrium constant.

Electrolytic Cell And Electrolysis

An electrolytic cell uses an external battery to force a non-spontaneous redox reaction.

These reaction names remain unchanged in both cell types, but the electrode charges reverse.

For molten sodium chloride:

• At the cathode:

Na+ gains an electron and forms sodium.

• At the anode:

Chloride ions lose electrons and form chlorine gas.

In aqueous solutions, water ions also compete. For aqueous zinc sulfate with inert electrodes:

• H+ is reduced at the cathode, producing hydrogen.

• OH- is oxidised at the anode, producing oxygen.

• Zinc is not deposited under the stated conditions.

Faraday’s Laws Of Electrolysis

• First Law

The mass deposited is directly proportional to the charge passed:

m = ZQ

Since Q = It:

m = ZIt

Here, Z is the electrochemical equivalent.

Z = Equivalent weight / Faraday constant

Faraday constant:

F = 96500 C per mole of electrons

• Second Law

When the same charge passes through different electrolytes:

m1 / m2 = Equivalent weight of substance 1 / Equivalent weight of substance 2

Conductance And Conductivity

Electrolytic conductance occurs due to the flow of ions.

• Metallic conductance occurs through electrons.

• Electrolytic conductance occurs through ions.

Conductance and resistance are related as:

G = 1/R

Conductivity and resistivity are related as:

kappa = 1/rho

Resistance is:

R = rho × l/A

Conductivity is:

kappa = G × l/A

The quantity l/A is called the cell constant.

Effect Of Dilution

On dilution:

• Specific conductivity decreases because ions per unit volume decrease.

• Molar conductivity increases because ions become more widely separated and move more freely.

Molar conductivity is:

Lambda m = kappa × 1000 / M

For strong electrolytes, molar conductivity increases gradually on dilution. For weak electrolytes, it increases sharply because dilution also increases their degree of dissociation.

Kohlrausch’s Law

At infinite dilution, ions are far apart, and interionic attraction becomes negligible. Each ion contributes independently to the total molar conductivity.

This is Kohlrausch’s law.

• For weak electrolytes:

Degree of dissociation = Molar conductivity at given concentration / Molar conductivity at infinite dilution

• For a weak acid:

Ka = concentration × degree of dissociation squared

Electrochemistry Important Questions

Practising a few well-chosen questions shows how well you have understood each concept. The questions here cover the points that appear most often in tests, with short answers for quick revision.

Question 1

In a cell, the following reactions take place:

Fe²⁺ → Fe³⁺ + e⁻

E°(Fe³⁺/Fe²⁺) = 0.77 V

2I⁻ → I₂ + 2e⁻

E°(I₂/I⁻) = 0.54 V

The standard electrode potential for the spontaneous reaction in the cell is x × 10⁻² V at 298 K. The value of x is _____ (nearest integer).

Question 2

Given that:

• E°(O₂/H₂O) = +1.23 V

• E°(S₂O₈²⁻/SO₄²⁻) = +2.05 V

• E°(Br₂/Br⁻) = +1.09 V

• E°(Au³⁺/Au) = +1.40 V

The strongest oxidising agent is:

(1) Br₂

(2) Au³⁺

(3) S₂O₈²⁻

(4) O₂

Question 3

The standard reduction potentials at 298 K for the following half-cells are given:

• Zn²⁺(aq) + 2e⁻ ⇌ Zn(s); E° = −0.762 V

• Cr³⁺(aq) + 3e⁻ ⇌ Cr(s); E° = −0.740 V

• 2H⁺(aq) + 2e⁻ ⇌ H₂(g); E° = 0.000 V

• Fe³⁺(aq) + e⁻ ⇌ Fe²⁺(aq); E° = +0.770 V

Which is the strongest reducing agent?

(1) Zn(s)

(2) Cr(s)

(3) H₂(g)

(4) Fe²⁺(aq)

[IIT JEE 1981]

Question 4

Given that the standard electrode potentials (E°) of the Cu²⁺/Cu and Cu⁺/Cu couples are 0.34 V and 0.522 V, respectively, the standard electrode potential (E°) of the Cu²⁺/Cu⁺ couple is:

(1) −0.158 V

(2) 0.182 V

(3) +0.158 V

(4) −0.182 V

[7 Jan, 2020 (Shift-I)]

Question 5

The hydrogen electrode is dipped in a solution of pH = 3 at 25°C. The potential of the electrode will be _____ × 10⁻² V.

Given: 2.303 RT/F = 0.059 V

[27 Jan, 2024 (Shift-II)]

Question 6

At 298 K, the standard reduction potential for the Cu²⁺/Cu electrode is 0.34 V.

Given:

• Ksp [Cu(OH)₂] = 1 × 10⁻²⁰

• 2.303 RT/F = 0.059 V

The reduction potential at pH = 14 for the above couple is (−)x × 10⁻² V.

The value of x is _____.

[13 April, 2023 (Shift-II)]

Key Tips to Revise Electrochemistry

Short, regular revision works better than long study sessions for this chapter. These tips help you avoid common mistakes in Electrochemistry exams.

• Remember that oxidation always occurs at the anode and reduction at the cathode, in both types of cells.

• Learn the electrode charges: in a galvanic cell the anode is negative, and in an electrolytic cell it is positive.

• Use reduction potentials for both electrodes when writing E°cell = E°cathode - E°anode.

• Write the net reaction before building Q, and leave out pure solids and pure liquids.

• Keep the direction of electron flow and current flow separate in your notes.

• Write the formulas for Faraday's laws, conductance and Kohlrausch's law on one revision sheet.

• Solve numericals on the Nernst equation and Faraday's laws every few days, and keep units in each step. Quick formula sheets are available on PW's study resources.

Once cells, electrode potentials, electrolysis and conductance are linked, Electrochemistry becomes easier to revise. Regularly practise numerical problems, memorise key formulas and solve important questions to strengthen your concepts and improve exam readiness. PW's Class 12 online courses offer concept classes and practice questions that can support your chemistry revision.

Related Resources

• Class 12 Chemistry Notes

• Chemistry Formulas

• Science Formulas

• Class 12 Online Course

• Class 12 Maths Notes