Galvanic Cell Virtual Laboratory

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Galvanic Cell Virtual Lab EMF & Thermodynamics

Electrode Potentials & Nernst Equation for JEE & NEET

Electrode Interface Active www.chemca.in

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1. Half-Cell Setup

Compartment A (Left Beaker) Zinc (Zn²⁺/Zn) [E° = -0.76 V] Iron (Fe²⁺/Fe) [E° = -0.44 V] Nickel (Ni²⁺/Ni) [E° = -0.25 V] Copper (Cu²⁺/Cu) [E° = +0.34 V] Silver (Ag⁺/Ag) [E° = +0.80 V]

Compartment B (Right Beaker) Zinc (Zn²⁺/Zn) [E° = -0.76 V] Iron (Fe²⁺/Fe) [E° = -0.44 V] Nickel (Ni²⁺/Ni) [E° = -0.25 V] Copper (Cu²⁺/Cu) [E° = +0.34 V] Silver (Ag⁺/Ag) [E° = +0.80 V]

Salt Bridge Electrolyte KCl in Agar-Agar KNO₃ in Agar-Agar NH₄NO₃ in Agar-Agar

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2. Ion Concentrations

Anode Ion [M²⁺] 1.000 M

Cathode Ion [M²⁺] 1.000 M

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3. Thermodynamic Live Log

>> Galvanic circuit initialized. Voltmeter calibrated.

Standard reduction potentials read. Computing spontaneous configuration.

Digital Voltmeter 1.100 V

Gibbs Free Energy -212.3 kJ

Zn / Zn²⁺

Cu / Cu²⁺

Left Compartment (Anode) Zn (s) → Zn²⁺ (aq) + 2e⁻

Right Compartment (Cathode) Cu²⁺ (aq) + 2e⁻ → Cu (s)

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4. Nernst Plotter (E vs. log Q)

E°(cell) Intercept

Cell Potential $E_{\text{cell}}$ Standard State intercept

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Balanced Cell Reactions

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JEE/NEET Electro Quiz

Score: 0/5

📖 Theoretical Foundations: Electrochemistry & Nernst Equation (NCERT Sync)

Spontaneity & Cell Potential

A cell is spontaneous if its Gibbs Free Energy change ($\Delta G$) is negative. This requires a positive cell electromotive force ($E_{\text{cell}}$) in accordance with:

$$\Delta G = -nFE_{\text{cell}}$$

Where $n$ represents the moles of electrons exchanged per unit reaction, and $F$ is the Faraday constant ($96485\text{ C mol}^{-1}$).

Mathematical Nernst Formulation

Standard potentials are defined under standard conditions ($1\text{ M}$, $298\text{ K}$, $1\text{ bar}$). To calculate potentials at non-standard concentrations, the Nernst Equation is used:

$$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{2.303 RT}{nF} \log Q$$

At $298\text{ K}$, the constant prefix simplifies to $0.0591\text{ V}$, making: $$E_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0591}{n} \log \frac{[\text{Anode Ion}]^x}{[\text{Cathode Ion}]^y}$$

Salt Bridge & Liquid Junction

The salt bridge prevents direct mixing of compartments while closing the circuit. It contains an inert electrolyte gel (e.g., $\text{KCl}$ or $\text{KNO}_3$).

To prevent the build-up of a liquid junction potential, the mobilities of the cation and anion must be nearly identical. Thus, ions like $\text{K}^+$ and $\text{NO}_3^-$ are ideally suited.

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