Electrolysis
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:

electric current measured in amperes or amps

time measured in seconds

the number of electrons required to produce or consume 1 mole of the substance
Amps, Time, Coulombs,
Faradays, and Moles of Electrons
Three equations relate these quantities:

amperes x time = Coulombs

96,485 coulombs = 1 Faraday

1 Faraday = 1 mole of electrons
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.
Top
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::

Write the balanced halfreactions involved.

Calculate the number of moles of electrons that were transferred.

Calculate the number of moles of substance that was produced/consumed at
the electrode.

Convert the moles of substance to desired units of measure.
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 25^{o}C and 1 atm) that
is produced during this time.

Write the halfreactions that take place at the anode and at the cathode.
anode (oxidation): 2 Cl^{} Cl_{2}(g)
+ 2 e^{}
cathode (reduction) Fe^{3+} + 3 e^{} Fe(s)

Calculate the number of moles of electrons.

Calculate the moles of iron and of chlorine produced using the number of
moles of electrons calculated and the stoichiometries from the balanced
halfreactions. According to the equations, three moles of electrons
produce one mole of iron and 2 moles of electrons produce 1 mole of chlorine
gas.

Calculate the mass of iron using the molar mass and calculate the volume
of chlorine gas using the ideal gas law (PV = nRT).
Top
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:

Find the quantity of substance produced/consumed in moles.

Write the balanced halfreaction involved.

Calculate the number of moles of electrons required.

Convert the moles of electrons into coulombs.

Calculate the time required.
Example: How long must a 20.0 amp current flow through a solution
of ZnSO_{4} in order to produce 25.00 g of Zn metal.
Top
Calculating
the Current Required
To determine the amount of current necessary to produce a known quantity
of substance in a given amount of time:

Find the quantity of substance produced/or consumed in moles.

Write the equation for the halfreaction taking place.

Calculate the number of moles of electrons required.

Convert the moles of electrons into coulombs of charge.

Calculate the current required.
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)?

Calculate the number of moles of H_{2}. (Remember, at STP, 1 mole
of any gas occupies 22.4 L.)

Write the equation for the halfreaction that takes place.
Hydrogen is produced during the reduction of water at the cathode.
The equation for this halfreaction is:
4 e^{} + 4 H_{2}O(l) 2 H_{2}(g) + 4
OH^{}(aq)

Calculate the number of moles of electrons. According to the stoichiometry
of the equation, 4 mole of e^{} are required to produce 2 moles
of hydrogen gas, or 2 moles of e^{}'s for every one mole of hydrogen
gas.

Convert the moles of electrons into coulombs of charge.

Calculate the current required.
Top