Galvanic Cells & Cell Potential

How a spontaneous redox reaction, split into two beakers, becomes a working battery — and how to predict its voltage.

High schoolIntro Gen ChemUni Year 1
⏱️ About 20 min

Drop a strip of zinc into copper sulfate and the zinc slowly dissolves while copper plates out — a spontaneous reaction that just wastes its energy as heat. But separate the two halves into different beakers and force the electrons to travel through a wire to get from one to the other, and that same reaction lights a bulb. That's a battery.

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The big idea: A galvanic (voltaic) cell splits a spontaneous redox reaction into two half-cells. Oxidation happens at the ANODE, reduction at the CATHODE; electrons flow from anode to cathode through the external wire, while ions flow through the salt bridge to keep each side neutral. The push behind the electrons is the cell potential, E°cell.
🎯 By the end, you'll be able to
  • Identify the anode and cathode and the reaction happening at each
  • Trace where electrons flow (the wire) and where ions flow (the salt bridge)
  • Explain the job of the salt bridge and the sign of each electrode
  • Calculate E°cell from standard reduction potentials and judge spontaneity
📎 Helpful to know first

Splitting a reaction to harvest its electrons

In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, zinc is oxidized and copper(II) is reduced. If the zinc metal and the copper ions touch, electrons jump directly and the energy is lost as heat. A galvanic cell (also called a voltaic cell) keeps them apart in two half-cells and forces the electrons to travel through an external wire — where we can put them to work.

One beaker holds a zinc electrode in Zn²⁺ solution; the other holds a copper electrode in Cu²⁺ solution. A wire connects the electrodes, and a salt bridge connects the solutions.

🔑 An Ox, Red Cat
The two electrodes have fixed jobs: An Ox — Anode is Oxidation — and Red Cat — Reduction at the Cathode. This is true for every electrochemical cell. In our example, the zinc electrode is the anode (Zn → Zn²⁺ + 2e⁻) and the copper electrode is the cathode (Cu²⁺ + 2e⁻ → Cu).
⚠️ The #1 misconception: what moves where
Electrons travel through the external WIRE, from anode to cathode — never through the salt bridge. The salt bridge carries ions, not electrons: current in the solution and bridge is carried by moving ions, because there are no free electrons swimming in solution. If you remember one thing, remember this: electrons in the wire, ions in the bridge.

What the salt bridge is for

As zinc dissolves, the anode beaker fills with extra positive Zn²⁺ and would build up positive charge; the cathode beaker loses Cu²⁺ and would go negative. Either imbalance would instantly stop the electron flow. The salt bridge (a tube of inert ions like K⁺ and NO₃⁻) fixes this: negative ions drift toward the anode and positive ions toward the cathode, keeping both beakers neutral so current keeps flowing.

✨ Electrode signs (and why they flip later)
In a galvanic cell the anode is negative and the cathode is positive — the anode is where electrons are pumped out into the wire. Keep this straight now, because in an electrolytic cell the signs reverse (anode becomes positive). The jobs — oxidation at the anode, reduction at the cathode — never change; only the signs do.

Cell potential: the electrical push

Each half-reaction has a standard reduction potential, E°, measured in volts against a common reference (the standard hydrogen electrode). The bigger the E°, the more that species 'wants' to be reduced. The overall cell potential is the difference between the two:

\[ E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}} \]
Both values are standard REDUCTION potentials. A positive E°cell means the cell reaction is spontaneous.
✨ Positive voltage = spontaneous cell
A positive E°cell means the reaction runs on its own — that's what makes a working galvanic cell. It ties back to energy: E°cell relates to the Gibbs free energy by ΔG° = −nFE°cell, so E°cell > 0 corresponds to ΔG° < 0 (spontaneous). A negative E°cell would need to be forced — that's electrolysis, two lessons from now.
📝 Worked example: Build a cell from zinc and copper. Given E°(Zn²⁺/Zn) = −0.76 V and E°(Cu²⁺/Cu) = +0.34 V, find E°cell and identify each electrode.
  1. The more positive potential is reduced (it's the cathode): copper, at +0.34 V. The other is oxidized (anode): zinc.
  2. So cathode = Cu (Cu²⁺ + 2e⁻ → Cu); anode = Zn (Zn → Zn²⁺ + 2e⁻).
  3. E°cell = E°cathode − E°anode = (+0.34) − (−0.76).
  4. = 0.34 + 0.76 = +1.10 V.
✓ E°cell = +1.10 V. It's positive, so the cell is spontaneous — the copper electrode is the cathode (+), the zinc electrode the anode (−).
✏️ Practice: A cell uses silver and zinc. Standard reduction potentials: E°(Ag⁺/Ag) = +0.80 V, E°(Zn²⁺/Zn) = −0.76 V. What is the standard cell potential E°cell, in volts? (Silver is reduced; zinc is oxidized.)
V
Solution
  1. Cathode (reduction) is the more positive half: silver, +0.80 V. Anode is zinc, −0.76 V.
  2. E°cell = E°cathode − E°anode = (+0.80) − (−0.76).
  3. = 0.80 + 0.76 = +1.56 V. Positive, so the cell is spontaneous.

Check your understanding

1. In a galvanic cell, electrons travel from the anode to the cathode through the:
Electrons flow through the external wire. The salt bridge carries ions, not electrons — there are no free electrons moving through the solution.
2. Which statement about a galvanic cell is correct?
An Ox, Red Cat: oxidation at the anode, reduction at the cathode. Electrons move through the wire, ions through the salt bridge.
3. E°cell for a cell is calculated as −0.25 V. What does this tell you?
A negative E°cell means the reaction is not spontaneous in that direction — it would have to be driven by an external power source (electrolysis).
4. What carries the current inside the solution and salt bridge?
Inside solutions, current is carried by the movement of ions. Free electrons only travel through the metal wire and electrodes.
🔎 Related lessons
Previous: Balancing Redox Reactions Next: The Nernst Equation Electrolysis & Faraday's Laws
✅ Key takeaways
  • A galvanic (voltaic) cell turns a spontaneous redox reaction into electricity using two half-cells.
  • An Ox, Red Cat: oxidation at the anode, reduction at the cathode.
  • Electrons flow anode → cathode through the external WIRE; ions move through the salt bridge.
  • In a galvanic cell the anode is negative and the cathode is positive (this flips in electrolysis).
  • E°cell = E°cathode − E°anode; a positive E°cell means a spontaneous cell (ΔG° = −nFE°cell).
➡️ E°cell assumes standard conditions — every solution at 1 M. Real batteries run down as concentrations change. The Nernst equation, next, tells you exactly how the voltage shifts away from E° as a cell operates.
Want to test yourself on this? Try the Chemistry practice test →