NEET Crash Course Module - 61

chemca.in

Home Class XII Grand Test

Home › Class XII › NEET Rapid Revision › Galvanic Cell & Nernst Equation

NEET Crash Course • Module 61

Galvanic Cell & Nernst Equation

Convert chemical energy into electrical power. Master the Daniel cell, salt bridge traps, Gibbs free energy calculations, and the all-important Nernst equation mathematics.

By chemca Academic Team • Updated for NEET 2027

Module Focus: Spontaneous Power

A Galvanic Cell (or Voltaic Cell) is an electrochemical device that converts the chemical energy of a spontaneous redox reaction ($\Delta G < 0$) into electrical energy. By separating the oxidation half-reaction from the reduction half-reaction, the electrons are forced to travel through an external wire, creating a usable electric current. The Nernst Equation allows us to calculate exactly how much voltage this cell will produce under non-standard conditions.

1. Architecture of a Galvanic Cell

The Daniel Cell ($Zn / Cu$)

Notice the flow of electrons in the wire versus the flow of ions in the salt bridge.

The Ultimate Mnemonic: L-O-A-N

To perfectly recall the setup of a Galvanic cell, remember the left side:

• L = Left
• O = Oxidation
• A = Anode
• N = Negative

Consequently, the Right side is Reduction, Cathode, and Positive.

The Salt Bridge

A U-tube filled with an inert electrolyte (like $KCl, KNO_3, NH_4NO_3$) in an agar-agar gel.

Functions:

1. Completes the electrical circuit (inner circuit).
2. Maintains electrical neutrality in both half-cells (prevents accumulation of $Zn^{2+}$ at anode and $SO_4^{2-}$ at cathode, which would instantly stop the cell).

NEET Mega Trap: When does $KCl$ fail?

The electrolyte in the salt bridge MUST be highly inert. It cannot react with the ions in either half-cell.

If an electrode involves Silver ($Ag^+$), Lead ($Pb^{2+}$), or Thallium ($Tl^+$) ions, you CANNOT use $KCl$ in the salt bridge!
Reason: The $Cl^-$ from the salt bridge will immediately react to form an insoluble precipitate ($AgCl \downarrow, PbCl_2 \downarrow$), destroying the cell's function. We use $KNO_3$ or $NH_4NO_3$ instead.

IUPAC Cell Representation

Instead of drawing a beaker, we represent the cell in a single line using the $Anode || Cathode$ format (ABC: Anode-Bridge-Cathode).

Zn(s) | Zn²⁺(aq, 1M) || Cu²⁺(aq, 1M) | Cu(s)

2. Cell Potential ($E^\circ_{\text{cell}}$) & Gibbs Energy

The Standard Cell Potential ($E^\circ_{\text{cell}}$) is calculated entirely using Standard Reduction Potentials (SRP).

$E^\circ_{\text{cell}} = E^\circ_{\text{Cathode}} - E^\circ_{\text{Anode}}$

*Both must be Reduction Potentials

Spontaneity & Maximum Work

The electrical work done by a galvanic cell equals the decrease in Gibbs Free Energy.

$\Delta G^\circ = -n \cdot F \cdot E^\circ_{\text{cell}}$

Where $n$ = moles of electrons transferred, $F$ = Faraday's constant ($96487 \approx 96500 \text{ C/mol}$).

For a cell to be spontaneous (act as a galvanic cell):
$E^\circ_{\text{cell}}$ MUST be Positive (+) | $\Delta G^\circ$ MUST be Negative (-)

3. The Nernst Equation

Standard potentials ($E^\circ$) are measured at $1 \text{ M}$ concentration and $298 \text{ K}$. Walther Nernst provided an equation to calculate the cell potential ($E_{\text{cell}}$) at any concentration and temperature.

The Master Equation at 298 K ($25^\circ\text{C}$)

$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0591}{n} \log_{10} Q$

Where $Q$ is the Reaction Quotient. For the general reaction $aA + bB \rightarrow cC + dD$:
$Q = \frac{[C]^c [D]^d}{[A]^a [B]^b} = \frac{[\text{Products}]}{[\text{Reactants}]} = \frac{[\text{Anode Ions}]}{[\text{Cathode Ions}]}$

*Pure solids and liquids are assigned a concentration of 1.

NEET Hack: The Le Chatelier Approach

You don't always need to calculate the math. Think of $E_{\text{cell}}$ as the "forward driving force" of the reaction. If you do something that pushes the reaction FORWARD, $E_{\text{cell}}$ increases.

• Increase Reactant (Cathode ion) concentration: Reaction shifts forward $\rightarrow$ $E_{\text{cell}}$ INCREASES.
• Increase Product (Anode ion) concentration: Reaction shifts backward $\rightarrow$ $E_{\text{cell}}$ DECREASES.

4. Equilibrium & Concentration Cells

A. At Equilibrium

As the cell operates, reactant concentration drops and product concentration rises. Eventually, the forward and reverse rates become equal. The battery is "dead."

$E_{\text{cell}} = 0$

$\Delta G = 0$

$Q = K_c$ (Equilibrium Const.)

Substituting these into the Nernst equation gives the highly-tested equilibrium formula:

$E^\circ_{\text{cell}} = \frac{0.0591}{n} \log_{10} K_c$

B. Concentration Cells

A cell where both the anode and cathode are made of the exact same material, but the electrolyte concentrations are different.

The Silver Concentration Cell

Key Properties

• Because both electrodes are identical, $E^\circ_{\text{cell}} = 0$.
• The driving force is solely the tendency to equalize concentrations (entropy driven).

Spontaneity Condition

$E_{\text{cell}} = 0 - \frac{0.0591}{n} \log \left(\frac{C_1}{C_2}\right)$

For $E_{\text{cell}}$ to be positive, $\log(C_1/C_2)$ must be negative. Therefore, $C_1 < C_2$. The cell works only if Cathode concentration is higher than Anode concentration.

Target 180/180

NEET Grand Test: Electrochemistry

15 High-Yield Questions testing Nernst applications, salt bridge traps, and Gibbs Free Energy calculations.

All Quizzes NEET Rapid Revision Series Video Lectures

🎯 NEET 2027 Target 180

Join the Ultimate Chemistry Crash Course

Master Physical Chemistry, Kinetics, and Thermodynamics. Get access to our full suite of Rapid Revision modules, formula sheets, and mock tests specifically designed for the NTA NEET pattern.

Explore All NEET Modules →

© 2026 chemca.in. Empowering NEET Aspirants.

🏠 Home 📝 Quiz 📘 Revision Notes ⬇ Downloads 📚 Class Notes ▶ Video Lectures

Get new posts by email

Powered by

📚 Also Read

Lecture Notes