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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 Eo will occur.

  • Non-spontaneous reaction: Eo for the electrolytic cell is negative.

  • Applied emf must be greater than the emf for the cell, ie greater than Eo.

  • 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     Eo1
cathode: M+ + e ---> M     Eo2
 
cell: X- + M+ ---> X + M Eocell = -ve

Eocell = Eo1 + Eo1 = negative

EMF (volts) required to drive this non-spontaneous reaction is greater than Eocell

Example

Consider the following electrochemical cell:

The half-cell equations are:

Cu2+ + 2e ---> Cu(s)     Eo = +0.34V

and

Ag+ + e ---> Ag(s)       Eo = +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 Eo = -0.34V
cathode (reduction): 2Ag+ + 2e ---> 2Ag(s) Eo = +0.80V
 
cell: Cu(s) + 2Ag+ ---> Cu2+ + 2Ag(s) Eo = -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) Eo = +0.34V
anode (oxidation): 2Ag(s) ---> 2Ag+ + 2e Eo = -0.80V
 
cell: Cu2+ + 2Ag(s) ---> Cu(s) + 2Ag+ Eo = +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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Related AUS-e-TUTE Topics

Oxidation and Reduction

Calculating Electrochemical Cell EMF (voltage)

Faraday's Laws of Electroylsis

Aluminium Production: Hall-Heroult Cells

Sodium Production: Down's Cells

Production and Uses of Sodium Hydroxide (Electrolytic Processes)

Galvanic (Voltaic) Electrochemical Cells

Corrosion

Batteries

Writing cell half-equations

Oxidation Numbers (States)

Displacement Reactions (Activity Series)

Nernst Equation

 

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