Chemistry Labs
Upper secondary · 15 min

A galvanic cell: from Daniell to Nernst

Build a zinc–copper cell, watch electrons flow, and see how the voltage depends on the metals and on concentration.

3D galvanic cell: two beakers with metal electrodes joined by a wire and a salt bridge, electrons moving along the wire, voltage shown in the corner.

Goal

Measure EE for four electrode pairs and test the Nernst equation by changing the ratio of ion concentrations.

Apparatus and reagents

Two half-cells connected by a wire and a salt bridge, for example Zn\ce{Zn} in ZnX2+\ce{Zn^2+} and Cu\ce{Cu} in CuX2+\ce{Cu^2+}; a voltmeter.

Procedure

  1. Keep Zn / Cu and read E∘E^\circ at lg⁡Q=0\lg Q = 0.
  2. Move lg⁡Q\lg Q to +2+2 and to −2-2 and record EE each time.
  3. Slide “extent of reaction” from 0 to 1 and watch the electrodes and the solutions.
  4. Change the electrode pair to Zn / Ag, Cu / Ag and Fe / Cu and compare the voltages.

What to observe

  • Zn / Cu gives 1.10 V, Zn / Ag 1.56 V, Fe / Cu 0.78 V and Cu / Ag 0.46 V at lg Q = 0.
  • E changes by about 0.030 V for every unit of lg Q (1.16 V at −2, 1.04 V at +2 for Zn / Cu).
  • The negative electrode gets thinner, the positive one thicker, and the blue color of Cu²⁺ fades.

Explanation

At the negative electrode (anode) the metal is oxidized, Zn→ZnX2++2 eX−\ce{Zn -> Zn^2+ + 2e-}; at the positive electrode (cathode) ions are reduced, CuX2++2 eX−→Cu\ce{Cu^2+ + 2e- -> Cu}. Electrons flow through the wire from − to +, and the salt bridge lets anions move toward the anode and cations toward the cathode so that charge stays balanced. At 25 °C, E=E∘−0.05916nlg⁡QE = E^\circ - \frac{0.05916}{n}\lg Q, where QQ is the ion-concentration ratio for the cell reaction; E∘E^\circ is the difference of the standard electrode potentials, Ecathode∘−Eanode∘E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}.

History of the experiment

In 1836 John Frederic Daniell built the first durable galvanic cell: a zinc electrode in dilute acid (later zinc sulfate) and a copper electrode in copper sulfate, separated by a porous partition. Unlike the erratic Volta pile, the Daniell cell delivered a steady ~1.1 V and powered the early telegraph network for decades. Faraday's electrolysis laws (1834) and Nernst's concentration-dependent potential (1889) then turned such cells into quantitative tools.

Chemists behind it

Michael FaradayWalther Nernst

Related topics

Galvanic cells, electrode potentials, Nernst equationElectrolysis, corrosion and metal protection

Virtual experiment: a simplified model to build intuition. It does not replace real lab work or safety training; never repeat chemistry at home without supervision.