ChemistryElectrochemistryClass 12

Galvanic cell, Nernst & electrolysis: interactive simulation

A galvanic cell converts chemical energy from a spontaneous redox reaction into electrical energy. In the classic Daniell cell, zinc is oxidized at the anode (Zn → Zn²⁺ + 2e⁻) and copper is reduced at the cathode (Cu²⁺ + 2e⁻ → Cu), yielding a standard EMF of 1.10 V.

Useful for: JEE Main, JEE Advanced, NEET, CBSE Boards · 4 modes in one lab

What you can explore

  1. 1

    Daniell cell (Zn-Cu)

    Zinc and copper half-cells with salt bridge ion flow and live voltmeter reading.
  2. 2

    Nernst equation & EMF

    Live plot of cell potential vs log Q with temperature and electron transfer controls.
  3. 3

    Faraday's electrolysis

    Electrolytic deposition of Cu, Ag or Al with current, time and deposited mass readouts.
  4. 4

    Conductance curves

    Titration conductance graphs for strong/weak acid-base combinations.

Try this

Slide the concentrations of Zn²⁺ and Cu²⁺ to see how the Nernst equation updates the cell potential in real time. Switch to Faraday’s mode to calculate mass deposited at the cathode given current and time.

  1. 1

    Keep [Cu²⁺] = 1 M and decrease [Zn²⁺] to 0.01 M.

    What you should see

    E_cell increases from 1.10 V to 1.16 V because Q falls, pushing the forward reaction harder.

  2. 2

    In Faraday electrolysis, select Silver (Ag⁺, n=1) vs Copper (Cu²⁺, n=2) at same current and time.

    What you should see

    Ag deposits about 3.4× more mass than Cu due to its high equivalent weight (108 g vs 31.77 g).

  3. 3

    In Conductance curves, choose HCl + NaOH.

    What you should see

    Conductance falls sharply before equivalence as fast H⁺ is replaced by slower Na⁺, then climbs sharply after equivalence due to excess fast OH⁻.

Formulas behind it

Nernst equation at 298 K

Ecell=Ecell∘−0.0591nlog⁡10QE_{\text{cell}} = E^\circ_{\text{cell}} - \frac{0.0591}{n} \log_{10} Q

Standard cell potential

Ecell∘=Ecathode∘−Eanode∘E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}

Faraday’s first law of electrolysis

w=Z⋅I⋅t=M⋅I⋅tn⋅96500w = Z \cdot I \cdot t = \frac{M \cdot I \cdot t}{n \cdot 96500}

Gibbs free energy and cell EMF

ΔG∘=−nFEcell∘=−2.303RTlog⁡10Keq\Delta G^\circ = -n F E^\circ_{\text{cell}} = -2.303 RT \log_{10} K_{\text{eq}}

Common mistakes

  • Writing [Reactants]/[Products] instead of [Products]/[Reactants] in reaction quotient Q.
  • Putting E°_cell = 0 at chemical equilibrium. At equilibrium, E_cell = 0, but E°_cell is a fixed thermodynamic constant.
  • Confusing the signs of electrodes: in a Galvanic cell the anode is negative, whereas in an Electrolytic cell the anode is positive.
Stuck on a galvanic cell, nernst & electrolysis question?Snap it and get a whiteboard video solution.

FAQ

Why is KCl used in the salt bridge?

Because the ionic mobilities of K⁺ and Cl⁻ are almost identical, preventing liquid junction potential from developing.

What happens when an external opposing voltage > 1.1 V is applied to a Daniell cell?

The cell stops behaving as a galvanic cell and reverses into an electrolytic cell: electrons flow from Cu to Zn.

What is 1 Faraday of charge?

The charge of 1 mole of electrons: 1 F ≈ 96,485 Coulombs, which deposits 1 gram-equivalent of any element.

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