Galvanic and electrolytic cells are two major types of electrochemical cells that connect chemical reactions with electrical energy. A galvanic cell uses a spontaneous redox reaction to produce an electric current, like a battery. An electrolytic cell uses an external power source to force a nonspontaneous redox reaction, such as plating a metal or splitting molten salt.
Comparing them helps students keep track of electron flow, electrode signs, and energy changes.
Understanding Chemistry: Galvanic vs Electrolytic Cells
A cell works because oxidation and reduction are separated into different places. At one electrode, atoms or ions lose electrons. Those electrons enter the solid electrode and move through a wire.
At the other electrode, particles in the solution gain the arriving electrons. This separation is important because it gives electrons a path through the outside circuit. In a galvanic setup, the two solutions are often linked by a salt bridge or a porous barrier.
It does not carry electrons through the liquid. It carries ions.
Negative ions move toward the side where positive ions are building up, while positive ions move toward the other side. Without this ion movement, charge would build up quickly and the current would stop.
The names anode and cathode describe the reaction taking place, not the electrical charge of an electrode. This is the source of many mistakes. In a galvanic cell, electrons leave the anode, so that electrode is negative.
They arrive at the cathode, so that electrode is positive. In an electrolytic cell, a power supply pulls electrons away from the anode and pushes electrons toward the cathode. The anode is therefore positive and the cathode is negative.
Electron flow in the wire always goes from the oxidation site toward the reduction site. Conventional current is defined in the opposite direction, which can add another layer of confusion in circuit diagrams.
Cell voltage tells how strongly a reaction can push electrons through a circuit under particular conditions. Tables usually list reduction potentials, so students must be careful about which half reaction is being reversed. A metal with a stronger tendency to lose electrons can act as the oxidation electrode when paired with a suitable ion.
Standard values assume fixed concentrations, pressure, and temperature. Real cells rarely stay at those conditions. As reactants are used up and products collect, the voltage changes.
Internal resistance, poor electrical contacts, and slow reactions at electrode surfaces can lower the measured voltage. A battery may still contain chemicals that could react, yet fail to deliver useful current because these practical limits become too large.
Electrolytic cells show why the surrounding material matters. During copper plating, copper ions gain electrons at the object being coated, forming a metal layer. The layer can become rough or uneven if the current is too high, the surface is dirty, or ion supply near the object is too slow.
In water based solutions, water may react instead of the ion students expect. For example, attempting to collect a very reactive metal from an aqueous solution may produce hydrogen gas instead. Molten salts avoid water but require high temperatures.
Rechargeable batteries use related ideas when a charger forces chemical changes in reverse. When studying any cell, first mark where oxidation happens, then track electrons, then track ions, and finally determine how the power source or chemical reaction provides the energy.
Key Facts
- Oxidation occurs at the anode in both galvanic and electrolytic cells.
- Reduction occurs at the cathode in both galvanic and electrolytic cells.
- Galvanic cell: spontaneous reaction, chemical energy becomes electrical energy, Ecell > 0.
- Electrolytic cell: nonspontaneous reaction, electrical energy becomes chemical energy, Ecell < 0 without the power source.
- Standard cell potential: E°cell = E°cathode - E°anode.
- Gibbs free energy and cell voltage are related by ΔG = -nFEcell.
Vocabulary
- Galvanic cell
- A cell that converts energy from a spontaneous redox reaction into electrical energy.
- Electrolytic cell
- A cell that uses electrical energy from an external source to drive a nonspontaneous redox reaction.
- Anode
- The electrode where oxidation occurs and electrons are produced or removed from a substance.
- Cathode
- The electrode where reduction occurs and electrons are gained by a substance.
- Salt bridge
- A pathway containing ions that completes the circuit and keeps charge balanced between half-cells.
Common Mistakes to Avoid
- Calling the anode always negative is wrong because the anode is negative in a galvanic cell but positive in an electrolytic cell.
- Calling the cathode always positive is wrong because the cathode is positive in a galvanic cell but negative in an electrolytic cell.
- Reversing oxidation and reduction is wrong because oxidation always occurs at the anode and reduction always occurs at the cathode, no matter which type of cell is used.
- Forgetting charge balance in the solutions is wrong because ion movement through a salt bridge or electrolyte is needed to prevent charge buildup and keep current flowing.
Practice Questions
- 1 A galvanic cell uses Zn(s) | Zn2+(aq) and Cu2+(aq) | Cu(s). If E°red for Cu2+/Cu is +0.34 V and E°red for Zn2+/Zn is -0.76 V, calculate E°cell.
- 2 For the cell reaction involving 2 moles of electrons and Ecell = 1.10 V, calculate ΔG using ΔG = -nFEcell and F = 96485 C/mol e-.
- 3 A student sees bubbles forming at one electrode in a cell connected to a battery. Explain how to decide whether the cell is galvanic or electrolytic and identify which electrode is the anode.