Electrons down a wire: how the Daniell cell becomes a battery

How to use this simulation

1. Read the equation and the live EMF on the voltmeter at the top. With the circuit closed, yellow electrons run down the wire from Zn to Cu.
2. In the lower plot watch the rod masses: zinc goes down, copper goes up. Open the circuit and everything stops, but the voltage stays.
3. Now move the Zn²⁺ and Cu²⁺ sliders. Lowering Cu²⁺ lowers the EMF, lowering Zn²⁺ raises it. That is the Nernst equation at work.
4. Run the cell for a while and the reactants dwindle, Q rises and the EMF slowly sinks. Raise the simulation speed to see it sooner.

What happens inside a torch

Open a torch battery and there is no machine inside, only chemicals. Yet press the switch and the bulb lights. At one end a substance wants to give up electrons, at the other a substance wants to take them. Give the electrons a wire and they will run through it, and a bulb on the way lights up.

A galvanic cell is exactly that arrangement. A reaction wants to happen anyway, but we do not let it happen in one beaker. Oxidation takes place in one, reduction in the other, and electrons are forced to travel through the wire. That journey is the electricity.

In the simulation the zinc rod gives up electrons and the Cu²⁺ in the copper solution takes them. The difference in pull between the two ends is what the voltmeter shows as 1.10 V.

Parts of the Daniell cell

John Daniell built this cell in 1836. Every part has a specific job.

Two half-reactions, one overall reaction

At the zinc rod, Zn atoms give up electrons and leave as Zn²⁺ ions, so the rod thins and loses mass. The electrons travel down the wire to the copper rod. There, Cu²⁺ ions from the solution take them and become Cu atoms that stick to the rod, so it thickens and gains mass.

Add the half-reactions and you get the overall reaction, Zn + Cu²⁺ → Zn²⁺ + Cu. Both halves move the same number of electrons, so n = 2.

What if there were no salt bridge

As Zn²⁺ builds up, the left beaker turns positive. As Cu²⁺ drains away, the right beaker is left with surplus negative sulfate ions. This charge imbalance would stop the flow, because each new electron would have more to push against.

The salt bridge fixes this. K⁺ cations drift toward the cathode and NO₃⁻ anions toward the anode. Both beakers stay neutral, so the circuit is complete and the electrons keep flowing.

Close the circuit in the simulation and you will see blue particles move right and red particles move left in the bridge. Open the circuit and they stop.

Standard potentials and cell EMF

Every metal and ion pair has a standard reduction potential E°, measured against the hydrogen electrode. Zn²⁺/Zn has E° = -0.76 V and Cu²⁺/Cu has E° = +0.34 V. The one with the higher E° is keener to take electrons, so it is the cathode.

Cell EMF = E°(cathode) − E°(anode) = 0.34 − (-0.76) = 1.10 V. It is positive, so the reaction is spontaneous. A positive EMF also means a negative ΔG°: ΔG° = −nFE° = −-212.3 kJ.

Pair any metal with another and E° tells you which is the anode. Here each is paired with copper:

The Nernst equation: concentration moves the EMF

The standard value applies when both ion concentrations are 1 M. Real cells differ, and then the Nernst equation gives the EMF. Q is the reaction quotient, products over reactants, here [Zn²⁺] ÷ [Cu²⁺].

At 25 °C, 2.303RT ÷ F is 0.0592 V. For n = 2 the coefficient becomes 0.0592 ÷ 2 = 0.0296.

A larger Q (more Zn²⁺ or less Cu²⁺) lowers the EMF, a smaller Q raises it, and Q = 1 gives E°. Compare the table below with the sliders in the simulation.

How much the rods change: Faraday’s arithmetic

The charge that flows through the wire equals the moles of electrons moved. One mole of electrons carries 96,485 C (the Faraday constant, F). By Zn → Zn²⁺ + 2e⁻, two moles of electrons dissolve one mole of Zn and deposit one mole of Cu.

In the simulation each beaker holds 100 mL. After 10 mmol of reaction, 0.654 g of Zn has dissolved, 0.636 g of Cu has plated, and 1,930 C has passed. Then [Zn²⁺] is 1.10 M, [Cu²⁺] is 0.90 M, and the EMF has fallen to 1.097 V.

Try it yourself in the simulation

Run these one at a time, and write down a prediction before each one.

• Experiment 1: leave both sliders at 10⁰ M (1 M) and close the circuit. The EMF is about 1.10 V, and it sinks slightly as the cell runs.
• Experiment 2: open the circuit. Electrons and ions both stop, the masses stay put, yet the voltmeter still shows a voltage.
• Experiment 3: drag [Cu²⁺] down to −2 (0.01 M). The EMF drops to about 1.041 V because Q = 100.
• Experiment 4: drop [Zn²⁺] to 0.1 M and keep [Cu²⁺] at 1 M. The EMF rises to 1.130 V because Q is small.
• Experiment 5: keep [Cu²⁺] very low and speed up the clock. The Cu²⁺ runs out quickly, the cell reads “spent”, and the masses stop changing.

Solved problems

Every number here is computed from the simulation’s own model, so you can check it against the screen.

Mistakes everyone makes

These are the places that cost the most marks.

• Mixing up anode and cathode: in a galvanic cell the anode is negative and the cathode positive. In electrolysis it is the other way round.
• Subtracting in the wrong order: cathode minus anode, with both as reduction potentials.
• Multiplying E° by the coefficients: potential is intensive, so doubling the equation does not change it.
• Thinking electrons cross the salt bridge: electrons use the wire, ions use the bridge.
• Putting solid metals in Q: Zn and Cu are solids, so Q holds only the ion concentrations.
• Getting the sign of log Q wrong: Q = [products] ÷ [reactants], and E = E° − …

Galvanic cells in real life

The Daniell cell is a museum piece, but its principle runs inside every battery today.

• Torch and clock batteries: zinc-carbon and alkaline cells work on exactly this principle.
• Phone batteries: a lithium-ion cell is a galvanic cell you can recharge.
• Car batteries: a lead-acid battery is several cells joined in series.
• Rust protection: galvanised iron is coated with zinc, which becomes the anode and corrodes first, sparing the iron.
• Fuel cells: electricity made straight from hydrogen and oxygen.

Exam corner (SSC, HSC)

Try these yourself first, then check the answers.

• Short questions: What is a galvanic cell? Give two jobs of the salt bridge. What happens at the anode and the cathode?
• Numerical: standard EMF of the Zn-Cu cell? Answer: 1.10 V.
• Numerical: [Zn²⁺] = 0.10 M, [Cu²⁺] = 1.00 M. EMF? Answer: 1.130 V.
• Numerical: 0.50 A for 30 minutes. How much Zn dissolves? Answer: 0.305 g.
• Reasoning: why does the voltage of a Zn-Cu cell fall with time? Zn²⁺ rises, Cu²⁺ falls, Q rises.
• Notation: write the Zn-Cu cell. Answer: Zn | Zn²⁺ ‖ Cu²⁺ | Cu.

Revision: at a glance

Read this list the night before the exam.

• A galvanic cell makes electricity from a spontaneous redox reaction.
• Oxidation at the anode (negative), reduction at the cathode (positive); electrons run anode to cathode through the wire.
• The salt bridge moves ions to balance charge and complete the circuit.
• E°cell = E°cathode − E°anode; 1.10 V for Zn-Cu.
• Nernst: E = E° − (0.0592 ÷ n) log Q; a larger Q means a lower EMF.
• Zn loses mass and Cu gains it; the amounts follow m = I t M ÷ (n F).

Frequently asked questions