Homemade Battery Cells

Build a progression of galvanic cells — lemon, salt-water, and the classic Daniell¹ cell — and measure how the electrochemical series sets their voltage

Beginner🕐45 min●Low hazardelectrochemistry

Summary

Every battery is a redox² reaction with the electrons routed through a wire. In this experiment you build a series of galvanic cells, from a lemon to the 19th-century Daniell cell, and measure their voltages with a multimeter. You’ll be able to predict which metal pairing gives the biggest voltage from the electrochemical series, and explain the difference between what a cell produces on paper and what it delivers under load.

History

Alessandro Volta³ stacked zinc and copper discs separated by brine-soaked card in 1800 to make the first battery — the “voltaic pile.” The Daniell cell (1836) fixed the pile’s rapid voltage drop and became the first reliable battery, powering telegraph networks for decades. The cells here retrace that history with a multimeter standing in for the galvanometers of the period.

Hazards & preparation

Note

The core cells here are low-hazard. Read the note for any optional variation.

• Copper sulfate (Daniell cell) is toxic and must not go down the drain — collect and dispose per the Safety page.
• Using a salt (NaCl) electrolyte with a driven current can evolve small amounts of chlorine; prefer sodium sulfate to avoid it, and work with ventilation.
• The optional aluminium–air cell uses sodium hydroxide, which is corrosive — gloves and goggles.

Part A — The Lemon Cell

Materials: a lemon, a zinc strip or galvanised nail, a copper strip or pipe, a multimeter (and an LED).

Steps:

1. Push the zinc and copper electrodes ~2 cm apart into the lemon — close but not touching.
2. Measure open-circuit voltage (red lead to copper, black to zinc): typically 0.8–1.0 V.
3. Wire two or three lemons in series (copper of one to zinc of the next) to exceed ~1.8 V and light an LED.

The lemon’s citric acid is the electrolyte. Internal resistance is high, so current is only a few milliamps — enough to prove the principle, not to do work.

Part B — Salt-Water Cell and the Electrochemical Series

Materials: saturated NaCl or Na₂SO₄ solution, a zinc strip, a carbon rod (from a dead battery) plus other metal strips, a container, multimeter.

Steps:

1. Dissolve ~30 g salt in 100 mL water.
2. Immerse zinc and carbon a few centimetres apart; measure the voltage (~0.6–0.8 V).
3. Swap the carbon for copper, iron, and aluminium in turn and record each voltage.

The further apart two metals sit in the electrochemical series, the higher the voltage — you’re measuring that series directly.

Part C — The Daniell Cell

The Daniell cell separates the two half-reactions into their own compartments, so it holds its voltage far better than a single-pot cell.

Materials: two jars; ~1 mol/L copper sulfate (25 g CuSO₄·5H₂O per 100 mL) with a copper strip; ~1 mol/L zinc sulfate (or Na₂SO₄) with a zinc strip; a salt bridge (filter paper or a U-tube of saturated Na₂SO₄); multimeter.

Steps:

1. Copper strip in the copper-sulfate jar (cathode, +); zinc strip in the sulfate jar (anode, −).
2. Bridge the jars with the salt bridge, dipping into both.
3. Measure open-circuit voltage (~1.0–1.1 V). Add a 100–1000 Ω load and measure the current.

Over hours the blue copper solution fades as copper plates out and the zinc strip visibly dissolves — the redox reaction made concrete.

Recharging it: the Daniell reaction is reversible, so the cell can be recharged. Disconnect the load and connect a DC supply set a little above the cell’s own voltage (~1.3–1.5 V), with its negative terminal on the zinc and positive on the copper. This drives both reactions backward — zinc ions plate back onto the zinc electrode while copper redissolves into the blue solution. Charging is simply electrolysis⁴, the exact reverse of discharge and the same process as Copper Electroplating. In practice the recharge is only partial: the replated zinc tends to grow as loose, spongy dendrites rather than a smooth strip, and any mixing of the two solutions plates stray copper onto the zinc — which is why a home Daniell cell is a lovely demonstration of reversibility but a poor practical rechargeable battery.

What you should see

The lemon manages ~0.9 V but sags under any load; the salt-water cell ranks your metals by voltage; the Daniell cell holds a steady ~1.1 V and, given time, shows copper depositing while zinc erodes. Wiring cells in series adds their voltages — the route to lighting an LED or running a clock.

Symptom	Likely cause	Fix
Very low or zero voltage	Electrodes touching, or same metal both sides	Separate them; use two different metals
Voltage collapses under load	High internal resistance (normal for lemon)	Wire cells in series; use the Daniell cell
Daniell voltage drifts down fast	Solutions mixing / bad salt bridge	Refresh the salt bridge; keep compartments separate
LED won’t light	Below its forward voltage	Series-wire more cells to exceed ~1.8 V

The Science

A galvanic cell has two half-cells, each a metal in an electrolyte. The more reactive metal (anode⁵) is oxidised, pushing electrons into the wire; at the cathode a reduction consumes them. The voltage is set by the difference in standard electrode potential:

\[E^\circ_\text{cell} = E^\circ_\text{cathode} - E^\circ_\text{anode}\]

Electrode	Half-reaction	E° (V)
Zn²⁺/Zn	Zn → Zn²⁺ + 2e⁻	−0.76
Fe²⁺/Fe	Fe → Fe²⁺ + 2e⁻	−0.44
2H⁺/H₂	2H⁺ + 2e⁻ → H₂	0.00
Cu²⁺/Cu	Cu²⁺ + 2e⁻ → Cu	+0.34
Ag⁺/Ag	Ag⁺ + e⁻ → Ag	+0.80

Zinc (−0.76) paired with copper (+0.34) gives 0.34 − (−0.76) = 1.10 V — the Daniell cell’s value. For the circuit to work, ions must also move: a salt bridge lets them migrate between compartments without the solutions mixing.

Real voltage under load is always lower than the open-circuit value because of internal resistance. That’s the lemon’s weakness — its resistance is hundreds of ohms, so even a small current drags the voltage down.

Questions to Explore

1. Why does zinc–copper give ~1.1 V regardless of the electrolyte? What sets the number?

   Hint / answer

   The voltage is the difference in standard electrode potentials of the two metals (Cu +0.34, Zn −0.76 → 1.10 V). The electrolyte carries the ions but doesn’t change which two half-reactions are running.

2. Why does the lemon cell’s voltage collapse when you connect an LED, but the Daniell cell’s barely moves?

   Hint / answer

   The lemon has high internal resistance, so drawing current drops most of the voltage inside the cell. The Daniell cell has low internal resistance and generous electrolyte, so it sustains voltage under a modest load.

3. What is the salt bridge actually doing? Why does the cell stall without it?

   Hint / answer

   It completes the circuit for ions, letting charge balance as metal ions form and deposit. Without it, charge builds up in each compartment and the reaction quickly stops.

4. Can you recharge a Daniell cell — and why isn’t it a good rechargeable battery?

   Hint / answer

   Yes: the reaction is reversible, so driving current backward with a power supply replates zinc and redissolves copper. But the recharged zinc grows as spongy dendrites, and mixing of the two solutions deposits stray copper on the zinc, so it loses capacity fast — fine as a demonstration, poor as a real battery.

Going further

• Measure internal resistance. From open-circuit voltage \(V_\text{OC}\) and loaded voltage \(V_L\) across \(R_L\): \(r = R_L\,(V_\text{OC}/V_L - 1)\).
• Temperature and concentration. Chill vs warm a lemon cell, or run the Daniell cell at 0.1, 0.5, and 1.0 mol/L, and watch voltage and resistance shift.
• Recharge the Daniell cell. Drive it backward with a DC supply set just above its own voltage and watch zinc replate — charging is electrolysis, the exact reverse of discharge.
• In the Electrochemistry track: now drive the reverse of a cell with a power supply in Copper Electroplating, or store charge without chemistry in the Supercapacitor.

Footnotes

1. John Frederic Daniell — British chemist (1790–1845) who invented the Daniell cell, the first reliable battery, in 1836.

2. Redox — A reaction in which electrons transfer from one species to another, coupling oxidation with reduction.

3. Alessandro Volta — Italian physicist (1745–1827) who invented the first battery, the voltaic pile, in 1800.

4. Electrolysis — Driving a non-spontaneous reaction by passing an electric current through an electrolyte.

5. Anode / cathode — The two electrodes of an electrochemical cell — oxidation happens at the anode, reduction at the cathode.