Electrolysis
Definition and meaning of Electrolysis in chemistry.
Electrolysis is the process of using a direct electric current to drive a nonspontaneous chemical reaction, most commonly decomposing a compound into its elements or simpler substances.
In more detail
An external power source pushes electrons through the circuit, forcing reduction at the cathode and oxidation at the anode within an electrolyte (a molten salt or aqueous solution containing ions). Because the overall reaction would not occur on its own, the applied voltage must exceed the reaction's reverse cell potential to overcome the energy barrier. This makes electrolysis the reverse-in-spirit counterpart to a galvanic cell, which generates electricity from a spontaneous redox reaction. Industrially, electrolysis produces reactive elements such as hydrogen, chlorine, sodium, and aluminum, and underlies electroplating and electrorefining of metals.
Key facts
| Field | Physical Chemistry |
|---|---|
| Example reaction | 2H2O(l) → 2H2(g) + O2(g) |
| Cathode process | Reduction (gains electrons) |
| Anode process | Oxidation (loses electrons) |
Passing a direct current through water containing a small amount of electrolyte (to carry charge) decomposes it: 2H2O(l) → 2H2(g) + O2(g), producing hydrogen gas at the cathode and oxygen gas at the anode in a 2:1 volume ratio.
Frequently asked questions
How does electrolysis differ from a galvanic cell?
Electrolysis consumes electrical energy from an external source to force a nonspontaneous redox reaction, whereas a galvanic (voltaic) cell spontaneously generates electrical energy from a favorable redox reaction.
Why is electrolysis needed to produce metals like aluminum and sodium?
These metals are too reactive to be reduced from their ores by ordinary chemical reducing agents, so molten-salt electrolysis (e.g., the Hall-Héroult process for aluminum) is used to force reduction of the metal ion to the free metal.