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Electrolysis - Electrolytic Cells

Key Concepts

Electrolysis: the process in which an electric current is used to bring about a chemical reaction which does not occur spontaneously.
Electrolytic Cell: converts electrical energy into chemical energy.
Electrodes : conductors used to permit the flow of electrons in an electrochemical cell.
One electrode is the anode, the other is the cathode.
Anode: Oxidation occurs at the anode.
Anions (negatively charged ions) migrate to the anode.
Anode is positive.
Anode disintegrates.
Cathode: Reduction occurs at the cathode.
Cations (positively charged ions) migrate to the cathode.
Cathode is negative.
Solid deposits at the cathode.
Electron flow: from anode to cathode.
Electrons flow from positive to negative.
If more than one reaction is possible, the reaction with the lowest E^o will occur.
Non-spontaneous reaction: E^o for the electrolytic cell is negative.
Applied emf must be greater than the emf for the cell, ie greater than E^o.
Mass of substance produced electrolytically is proportional to the quantity of electricity flowing.

Electrolytic Cell

A supply of electricity is required to drive the electrolytic cell reactions.

Electrolytic Cell Reactions

anode: X^- ---> X + e E^o₁

cathode: M⁺ + e ---> M E^o₂

cell: X^- + M⁺ ---> X + M E^o_cell = -ve

E^o_cell = E^o₁ + E^o₁ = negative
EMF (volts) required to drive this non-spontaneous reaction is greater than E^o_cell

Example

Consider the following electrochemical cell:

The half-cell equations are:
Cu²⁺ + 2e ---> Cu_(s) E^o = +0.34V
and
Ag⁺ + e ---> Ag_(s) E^o = +0.80V

Whether copper is oxidised or reduced depends on what the ? in the diagram is.

? = voltmeter, galvanometer, or metal wire

If we replace the ? in the diagram with a voltmeter or galvanometer to complete the circuit the reaction would proceed spontaneously:

anode (oxidation): Cu_(s) ---> Cu²⁺ + 2e E^o = -0.34V

cathode (reduction): 2Ag⁺ + 2e ---> 2Ag_(s) E^o = +0.80V

cell: Cu_(s) + 2Ag⁺ ---> Cu²⁺ + 2Ag_(s) E^o = -0.34 + 0.80 = +0.46V

Electrons would spontaneously flow from the negative copper anode to the positive silver cathode generating electricity which would be measured by the voltmeter or galvanometer.
The copper anode would disintegrate and silver would be deposited on the silver cathode.
This type of electrochemical cell is known as a galvanic or voltaic cell.

? = battery or other electricity supply with emf > 0.46V

If we replace the ? in the diagram with a battery supplying more than 0.46V then the battery has greater electron pushing power than the spontaneous reaction in the electrochemical cell, forcing the reactions to go in the opposite direction to those above.

cathode (reduction): Cu²⁺ + 2e ---> Cu_(s) E^o = +0.34V

anode (oxidation): 2Ag_(s) ---> 2Ag⁺ + 2e E^o = -0.80V

cell: Cu²⁺ + 2Ag_(s) ---> Cu_(s) + 2Ag⁺ E^o = +0.34 + -0.80 = -0.46V

Electrons are pulled out of the positive silver anode and flow to the negative copper cathode.
This reaction is not generating electricity it is using electricity!
The silver anode disintengrates while copper deposits on the copper cathode.
This type of electrochemical cell is known as an electrolytic cell.

Uses of Electrolysis

Recharging rechargable batteries
Plating one metal, eg, silver, gold or chromium, onto another metal (electroplating)
Production of sodium and chlorine from molten (fused) sodium chloride (Downs Cell)
Production of chlorine gas and sodium hydroxide from concentrated aqueous sodium chloride (Nelson or Diaphragm Cell)
Extraction of aluminium (Hall-Heroult Cells) and copper from their ores (electrowinning)
Refining of copper (electrorefining)

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