Electrolysis involves passing an electric current through either a molten salt or an ionic solution.  The ions are "forced" to undergo either oxidation (at the anode) or reduction (at the cathode).  Most electrolysis problems are really stoichiometry problems with the addition of an amount of electric current.  The quantities of substances produced or consumed by the electrolysis process is dependent upon the following:

Amps, Time, Coulombs, Faradays, and Moles of Electrons

Three equations relate these quantities:

The thought process for interconverting between amperes and moles of electrons is:

amps & time <----> Coulombs <----> Faradays <----> moles of electrons

Use of these equations are illustrated in the following sections.


Calculating the Quantity of Substance Produced or Consumed

To determine the quantity of substance either produced or consumed during electrolysis given the time a known current flowed::

Example:   A 40.0 amp current flowed through molten iron(III) chloride for 10.0 hours (36,000 s).  Determine the mass of iron and the volume of chlorine gas (measured at 25oC and 1 atm) that is produced during this time.
anode (oxidation):  2 Cl- --> Cl2(g) + 2 e-

cathode (reduction)  Fe3+ + 3 e- -->  Fe(s)

Finding mass of Fe and volume of Cl<sub>2</sub >


Calculating the Time Required

To determine the quantity of time required to produce a known quantity of a substance given the amount of current that flowed:

Example:  How long must a 20.0 amp current flow through a solution of ZnSO4 in order to produce 25.00 g of Zn metal.
Finding the time


Calculating the Current Required

To determine the amount of current necessary to produce a known quantity of substance in a given amount of time:

Example:  What current is required to produce 400.0 L of hydrogen gas, measured at STP, from the electrolysis of water in 1 hour (3600 s)?

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