John Frederic Daniell and the Daniell Cell: A Battery That Finally Lasted

The chemistry student's first galvanic cell is almost always the zinc-copper cell, drawn with two beakers, a salt bridge and a voltmeter showing about 1.1 volts. This textbook cell is named after John Frederic Daniell, an English chemist and meteorologist who invented it in 1836. Volta's pile of 1800 had given the first current, but it was unreliable: its voltage faded quickly because bubbles of hydrogen coated the copper. Daniell designed a cell in which no gas was produced, and which could supply a steady current for hours. It became the standard battery for the early telegraph and the standard laboratory source of EMF for decades. Understanding the Daniell cell is necessary for electrochemistry, a high-weight topic in JAM, NET and GATE.

The Scientific Problem Before the Discovery

The simple voltaic cell, with zinc and copper in dilute sulphuric acid, suffers from two main faults:

• Polarisation: hydrogen formed at the copper electrode (2H⁺ + 2e⁻ → H₂) collects as a layer of gas bubbles. Hydrogen is a poor conductor, it increases internal resistance and it sets up a back EMF, so the current drops quickly.
• Local action: impurities in commercial zinc form tiny local cells, so zinc dissolves even when no current is drawn, wasting the metal.

Scientists such as William Sturgeon had attempted to improve matters by amalgamating zinc with mercury to reduce local action. But the problem of polarisation remained. Telegraphy and electrochemical research both needed a battery that gave a constant voltage.

The Discovery: Step by Step

1. Two solutions. Daniell used two different electrolytes, kept apart: zinc in dilute sulphuric acid (later zinc sulphate), and copper in copper(II) sulphate solution.
2. Separation by a porous barrier. The two solutions were separated by a porous partition, such as unglazed earthenware or an animal membrane, which allows ions to pass but prevents quick mixing. In his first design, the zinc rod stood in a porous pot placed inside a copper vessel containing the copper sulphate.
3. Removing the hydrogen problem. At the copper electrode, copper ions are reduced instead of hydrogen ions. Thus, a solid copper deposit forms in place of a gas, and polarisation is avoided.
4. Constant voltage. The cell gave a relatively steady EMF of about 1.1 V while in use.
5. Publication. He described the cell in 1836 in the Philosophical Transactions of the Royal Society.

The Chemistry Behind It

Half-reactions

Anode (oxidation, negative):    Zn(s)  →  Zn²⁺(aq) + 2e⁻
Cathode (reduction, positive):  Cu²⁺(aq) + 2e⁻  →  Cu(s)
Overall:                        Zn(s) + Cu²⁺(aq)  →  Zn²⁺(aq) + Cu(s)

The overall reaction is a displacement reaction: zinc, being more reactive, displaces copper from its salt. In the cell, the two half-reactions are separated so that electrons must travel through the external wire, from zinc to copper, doing electrical work.

Cell notation and EMF

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
E°cell = E°(Cu²⁺/Cu) − E°(Zn²⁺/Zn) = 0.34 − (−0.76) = +1.10 V

The single vertical lines mean a phase boundary, and the double line means a salt bridge or porous partition. The anode is written on the left.

Role of the salt bridge or porous pot

As the cell runs, Zn²⁺ ions build up in the anode compartment and Cu²⁺ ions are removed in the cathode compartment. Without any ion movement, charge would build up and the cell would stop. The salt bridge (a U-tube with KCl or KNO₃ in agar) or the porous pot lets ions migrate to maintain electrical neutrality: anions move toward the anode and cations toward the cathode. The salt in a bridge must have ions with nearly equal mobilities (K⁺ and Cl⁻), to reduce the liquid-junction potential.

Nernst equation for the Daniell cell

E = E° − (0.0592/2) log ( [Zn²⁺] / [Cu²⁺] )     at 298 K

If [Zn²⁺] = [Cu²⁺], E = E° = 1.10 V. If the cell is run until equilibrium, E = 0. The equilibrium constant is found from log K = nE°/0.0592 = 2 × 1.10 / 0.0592 ≈ 37, so K ≈ 10³⁷, showing that the reaction goes almost to completion. Further, ΔG° = −nFE° = −2 × 96500 × 1.10 ≈ −212 kJ per mole of reaction.

Applications Today

• Teaching and research: the Daniell cell is the model for studying electrode potentials, the Nernst equation and thermodynamics of cells.
• Telegraph and early electrical industry: Daniell cells provided power for the early telegraph and for electroplating in the 1800s.
• Concept of separated half-cells: all modern batteries, from lead-acid to lithium-ion, have separate electrode reactions joined by an ion-conducting electrolyte or separator.
• Electroplating and refining: reverse of the cell; copper is refined by electrolysis in copper sulphate solution.
• Corrosion protection: galvanising and sacrificial anodes use the same relative reactivity of zinc and iron.

Exam Corner

• Daniell cell: Zn | ZnSO₄ || CuSO₄ | Cu; E°cell = 1.10 V; zinc is anode (oxidation), copper is cathode (reduction).
• Electrons flow from Zn to Cu in the external circuit; conventional current from Cu to Zn.
• If external opposing potential > 1.10 V, the cell reaction is reversed (electrolytic cell).
• Salt bridge: completes the circuit, maintains neutrality and minimises junction potential.
• ΔG° = −nFE° and log K = nE°/0.0592 at 298 K.

Common Misconceptions

• "The cathode is always positive." In a galvanic cell the cathode is positive; in an electrolytic cell it is negative.
• "Electrons flow through the salt bridge." Only ions move through the bridge; electrons move through the wire.
• "The Daniell cell is rechargeable." It is not designed to be recharged; it is a primary cell, though passing opposite current can reverse its reaction.

FAQ

Why does the Daniell cell not polarise?

Because the cathode reaction forms solid copper and not hydrogen gas, so there are no insulating bubbles.

What happens to the voltage as the cell runs?

[Zn²⁺] rises and [Cu²⁺] falls, so by the Nernst equation E decreases slowly, reaching zero at equilibrium.

Can a salt bridge be made with any salt?

No. The salt must not react with the electrolytes, and its ions should have similar mobilities, so KCl, KNO₃ or NH₄NO₃ are used.

Why is zinc the anode?

Zinc has a more negative reduction potential than copper, so it is more easily oxidised.

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