Galvanic and electrolytic cells both use redox reactions to transfer electrons, but they do opposite energy conversions. A galvanic cell produces electrical energy from a spontaneous chemical reaction. An electrolytic cell uses electrical energy to drive a nonspontaneous chemical reaction.
This cheat sheet helps students compare the two cell types quickly and avoid sign, electrode, and electron-flow errors.
The most important ideas are oxidation, reduction, electrode identity, electron flow, and cell potential. Oxidation always occurs at the anode, and reduction always occurs at the cathode. For any electrochemical cell, the standard cell potential is found using .
Spontaneity is connected to free energy by .
Key Facts
- Oxidation is loss of electrons, reduction is gain of electrons, and this can be remembered as OIL RIG.
- Oxidation always occurs at the anode, and reduction always occurs at the cathode in both galvanic and electrolytic cells.
- In a galvanic cell, the reaction is spontaneous, so and .
- In an electrolytic cell, the reaction is nonspontaneous, so an external power source is required and for the reaction as written.
- The standard cell potential is calculated with using reduction potentials.
- The relationship between cell potential and free energy is , where is moles of electrons and .
- The Nernst equation at is .
- In line notation, the anode is written on the left, the cathode is written on the right, and the salt bridge is shown with .
Vocabulary
- Galvanic cell
- A cell that converts chemical energy into electrical energy using a spontaneous redox reaction.
- Electrolytic cell
- A cell that uses electrical energy from an outside source to force a nonspontaneous redox reaction.
- Anode
- The electrode where oxidation occurs in any electrochemical cell.
- Cathode
- The electrode where reduction occurs in any electrochemical cell.
- Cell potential
- The voltage of an electrochemical cell, calculated under standard conditions by .
- Salt bridge
- A connection that allows ions to move between half-cells so charge does not build up.
Common Mistakes to Avoid
- Calling the anode positive in every cell is wrong because the anode is negative in a galvanic cell but positive in an electrolytic cell.
- Reversing oxidation and reduction is wrong because oxidation always happens at the anode and reduction always happens at the cathode.
- Multiplying reduction potentials when balancing electrons is wrong because values are intensive and do not change when half-reactions are multiplied.
- Using is wrong because standard cell potential must be calculated as .
- Forgetting the salt bridge is wrong because ion flow is needed to maintain electrical neutrality and keep current flowing.
Practice Questions
- 1 A cell has and . Calculate .
- 2 For a reaction with and , calculate using and .
- 3 At , calculate if , , and . Use .
- 4 Explain why a galvanic cell can power a device without an external battery, while an electrolytic cell cannot.
Understanding Galvanic vs Electrolytic Cells
A working cell must keep charge balanced inside the liquids while electrons travel through the wire. This is the job of a salt bridge in many galvanic setups. Ions in the bridge move into each half cell as the reaction proceeds.
If metal atoms at one electrode enter solution as positive ions, that solution becomes too positive unless negative ions move in. At the other electrode, positive ions are removed from solution as they become solid metal, so positive ions from the bridge may move toward that side.
Without this ion movement, charge buildup quickly stops the electron current. A porous barrier can do the same job when it permits ions to pass but limits direct mixing of the solutions.
Electrode labels and electrode signs are different ideas. The labels anode and cathode describe the chemical process occurring there. The signs describe the electrical conditions of the particular device.
In a battery delivering current, electrons leave one electrode and enter the other through the external circuit. The electron source has a negative sign. In an electrolysis apparatus, the power supply pulls electrons away from one electrode and pushes them toward the other.
This reverses the signs compared with a battery. Students often lose marks by memorizing that anode means negative.
That statement only fits a galvanic cell. Link the sign to the direction electrons are forced or allowed to move, not to the electrode name alone.
Electrolysis has many practical uses. Electroplating coats an object with a thin metal layer for appearance, corrosion resistance, or improved electrical contact. The object being coated is placed where metal ions gain electrons and become solid metal.
Producing aluminum from its ore requires huge electrolytic cells because aluminum ions are difficult to reduce. Charging a rechargeable battery uses an external source to restore chemical substances that were changed during discharge. These examples show why the required voltage is often greater than a simple calculation suggests.
Real cells have resistance in wires and solutions. Some reactions need an extra push at an electrode surface before they proceed at a useful rate. This extra required voltage is called overpotential.
Standard potentials describe carefully defined conditions, not every real situation. Concentrations change as a cell runs, gases may build up, and temperature may differ from the standard value. The reaction quotient summarizes the current amounts of reactants and products.
When products become more abundant, the cell voltage commonly falls for the forward reaction. When reactants are more concentrated, the forward voltage can rise. The Nernst equation puts this effect into numbers at a stated temperature.
First balance the overall redox equation. Then determine how many electrons are transferred in that balanced equation. Do not multiply a tabulated electrode potential when multiplying a half reaction.
Potential is an intensive quantity, unlike free energy. Finally, check whether your calculated sign agrees with the reaction direction and the physical setup.