In chemistry and manufacturing, electrolysis is a method of using a direct electric current
to drive an otherwise non-spontaneous chemical reaction.
Electrolysis is commercially highly important as a stage in the separation of elements from
naturally occurring sources such as ores using an electrolytic cell.
The voltage that is needed for electrolysis to occur is called the decomposition potential.
History The word electrolysis comes from the Greek
ἤλεκτρον [ɛ̌ːlektron] "amber" and λύσις [lýsis] "dissolution".
1785 – Martinus van Marum's electrostatic generator was used to reduce tin, zinc, and
antimony from their salts using electrolysis.
1800 – William Nicholson and Anthony Carlisle, decomposed water into hydrogen and oxygen.
1807 – Potassium, sodium, barium, calcium and magnesium were discovered by Sir Humphry
Davy using electrolysis.
1875 – Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis.
1886 – Fluorine was discovered by Henri Moissan using electrolysis.
1886 – Hall-Héroult process developed for making aluminium
1890 – Castner-Kellner process developed for making sodium hydroxide
Overview Electrolysis is the passage of a direct electric
current through an ionic substance that is either molten or dissolved in a suitable solvent,
resulting in chemical reactions at the electrodes and separation of materials.
The main components required to achieve electrolysis are :
An electrolyte : a substance containing free ions which are the carriers of electric current
in the electrolyte.
If the ions are not mobile, as in a solid salt then electrolysis cannot occur.
A direct current supply : provides the energy necessary to create or discharge the ions
in the electrolyte.
Electric current is carried by electrons in the external circuit.
Two electrodes : an electrical conductor which provides the physical interface between
the electrical circuit providing the energy and the electrolyte
Electrodes of metal, graphite and semiconductor material are widely used.
Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte
and the cost of manufacture.
Process of electrolysis The key process of electrolysis is the interchange
of atoms and ions by the removal or addition of electrons from the external circuit.
The desired products of electrolysis are often in a different physical state from the electrolyte
and can be removed by some physical processes.
For example, in the electrolysis of brine to produce hydrogen and chlorine, the products
are gaseous.
These gaseous products bubble from the electrolyte and are collected.
2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2 A liquid containing mobile ions is produced
by: Solvation or reaction of an ionic compound
with a solvent to produce mobile ions An ionic compound is melted by heating
An electrical potential is applied across a pair of electrodes immersed in the electrolyte.
Each electrode attracts ions that are of the opposite charge.
Positively charged ions move towards the electron-providing cathode, whereas negatively charged ions move
towards the positive anode.
At the electrodes, electrons are absorbed or released by the atoms and ions.
Those atoms that gain or lose electrons to become charged ions pass into the electrolyte.
Those ions that gain or lose electrons to become uncharged atoms separate from the electrolyte.
The formation of uncharged atoms from ions is called discharging.
The energy required to cause the ions to migrate to the electrodes, and the energy to cause
the change in ionic state, is provided by the external source of electrical potential.
Oxidation and reduction at the electrodes Oxidation of ions or neutral molecules occurs
at the anode, and the reduction of ions or neutral molecules occurs at the cathode.
For example, it is possible to oxidize ferrous ions to ferric ions at the anode:
Fe2+ aq → Fe3+
aq + e– It is also possible to reduce ferricyanide
ions to ferrocyanide ions at the cathode: Fe(CN)3-
6 + e– → Fe(CN)4- 6
Neutral molecules can also react at either of the electrodes.
For example: p-Benzoquinone can be reduced to hydroquinone at the cathode:
+ 2 e– + 2 H+ → In the last example, H+ ions also take part
in the reaction, and are provided by an acid in the solution, or the solvent itself.
Electrolysis reactions involving H+ ions are fairly common in acidic solutions.
In alkaline water solutions, reactions involving OH- are common.
The substances oxidised or reduced can also be the solvent or the electrodes.
It is possible to have electrolysis involving gases.
Energy changes during electrolysis The amount of electrical energy that must
be added equals the change in Gibbs free energy of the reaction plus the losses in the system.
The losses can be arbitrarily close to zero, so the maximum thermodynamic efficiency equals
the enthalpy change divided by the free energy change of the reaction.
In most cases, the electric input is larger than the enthalpy change of the reaction,
so some energy is released in the form of heat.
In some cases, for instance, in the electrolysis of steam into hydrogen and oxygen at high
temperature, the opposite is true.
Heat is absorbed from the surroundings, and the heating value of the produced hydrogen
is higher than the electric input.
Related techniques The following techniques are related to electrolysis:
Electrochemical cells, including the hydrogen fuel cell, utilise differences in Standard
electrode potential in order to generate an electrical potential from which useful power
can be extracted.
Although related via the interaction of ions and electrodes, electrolysis and the operation
of electrochemical cells are quite distinct.
A chemical cell should not be thought of as performing "electrolysis in reverse".
Faraday's laws of electrolysis
First law of electrolysis In 1832, Michael Faraday reported that the
quantity of elements separated by passing an electric current through a molten or dissolved
salt is proportional to the quantity of electric charge passed through the circuit.
This became the basis of the first law of electrolysis:
Second law of electrolysis Faraday discovered that when the same amount
of electricity is passed through different electrolytes/elements connected in series,
the mass of substance liberated/deposited at the electrodes is directly proportional
to their equivalent weights.
Industrial uses
Production of aluminium, lithium, sodium, potassium, magnesium, calcium
Coulometric techniques can be used to determine the amount of matter transformed during electrolysis
by measuring the amount of electricity required to perform the electrolysis
Production of chlorine and sodium hydroxide Production of sodium chlorate and potassium
chlorate Production of perfluorinated organic compounds
such as trifluoroacetic acid Production of electrolytic copper as a cathode,
from refined copper of lower purity as an anode.
Electrolysis has many other uses: Electrometallurgy is the process of reduction
of metals from metallic compounds to obtain the pure form of metal using electrolysis.
For example, sodium hydroxide in its molten form is separated by electrolysis into sodium
and oxygen, both of which have important chemical uses.
Anodization is an electrolytic process that makes the surface of metals resistant to corrosion.
For example, ships are saved from being corroded by oxygen in the water by this process.
The process is also used to decorate surfaces.
Production of oxygen for spacecraft and nuclear submarines.
Electroplating is used in layering metals to fortify them.
Electroplating is used in many industries for functional or decorative purposes, as
in vehicle bodies and nickel coins.
Production of hydrogen for fuel, using a cheap source of electrical energy.
Electrolytic Etching of metal surfaces like tools or knives with a permanent mark or logo.
Electrolysis is also used in the cleaning and preservation of old artifacts.
Because the process separates the non-metallic particles from the metallic ones, it is very
useful for cleaning a wide variety of metallic objects, from old coins to even larger objects
including rusted cast iron cylinder blocks and heads when rebuilding automobile engines.
Competing half-reactions in solution electrolysis Using a cell containing inert platinum electrodes,
electrolysis of aqueous solutions of some salts leads to reduction of the cations and
oxidation of the anions.
However, with salts of some metals hydrogen is evolved at the cathode, and for salts containing
some anions oxygen is evolved at the anode.
In both cases this is due to water being reduced to form hydrogen or oxidised to form oxygen.
In principle the voltage required to electrolyze a salt solution can be derived from the standard
electrode potential for the reactions at the anode and cathode.
The standard electrode potential is directly related to the Gibbs free energy, ΔG, for
the reactions at each electrode and refers to an electrode with no current flowing.
An extract from the table of standard electrode potentials is shown below.
In terms of electrolysis, this table should be interpreted as follows:
Oxidized species with a more negative cell potential are more difficult to reduce than
oxidized species with a more positive cell potential.
For example it is more difficult to reduce a sodium ion to a sodium metal than it is
to reduce a zinc ion to a zinc metal.
Reduced species with a more positive cell potential are more difficult to oxidize than
reduced species with a more negative cell potential.
For example it is more difficult to oxidize sulfate anions than it is to oxidize bromide
anions.
Using the Nernst equation the electrode potential can be calculated for a specific concentration
of ions, temperature and the number of electrons involved.
For pure water: the electrode potential for the reduction
producing hydrogen is −0.41 V the electrode potential for the oxidation
producing oxygen is +0.82 V. Comparable figures calculated in a similar
way, for 1M zinc bromide, ZnBr2, are −0.76 V for the reduction to Zn metal and +1.10
V for the oxidation producing bromine.
The conclusion from these figures is that hydrogen should be produced at the cathode
and oxygen at the anode from the electrolysis of water which is at variance with the experimental
observation that zinc metal is deposited and bromine is produced.
The explanation is that these calculated potentials only indicate the thermodynamically preferred
reaction.
In practice many other factors have to be taken into account such as the kinetics of
some of the reaction steps involved.
These factors together mean that a higher potential is required for the reduction and
oxidation of water than predicted, and these are termed overpotentials.
Experimentally it is known that overpotentials depend on the design of the cell and the nature
of the electrodes.
For the electrolysis of a neutral sodium chloride solution, the reduction of sodium ion is thermodynamically
very difficult and water is reduced evolving hydrogen leaving hydroxide ions in solution.
At the anode the oxidation of chlorine is observed rather than the oxidation of water
since the overpotential for the oxidation of chloride to chlorine is lower than the
overpotential for the oxidation of water to oxygen.
The hydroxide ions and dissolved chlorine gas react further to form hypochlorous acid.
The aqueous solutions resulting from this process is called electrolyzed water and is
used as a disinfectant and cleaning agent.
Electrolysis of water
One important use of electrolysis of water is to produce hydrogen.
2 H2O(l) → 2 H2(g) + O2(g); E0 = -1.229 V
Hydrogen can be used as a fuel for powering internal combustion engines by combustion
or electric motors via hydrogen fuel cells.
This has been suggested as one approach to shift economies of the world from the current
state of almost complete dependence upon hydrocarbons for energy
The energy efficiency of water electrolysis varies widely.
The efficiency of an electrolyser is a measure of the enthalpy contained in the hydrogen,
compared with the input electrical energy.
Heat/enthalpy values for hydrogen are well published in science and engineering texts,
as 144 MJ/kg.
Note that fuel cells cannot utilise this full amount of heat/enthalpy, which has led to
some confusion when calculating efficiency values for both types of technology.
In the reaction, some energy is lost as heat.
Some reports quote efficiencies between 50% and 70% for alkaline electrolysers; however,
much higher practical efficiencies are available with the use of PEM and catalytic technology,
such as 95% efficiency.
In the US there is still an occasional erroneous tendency to use the 'Lower Heating Value'
for efficiencies.
This value does not represent the total amount of energy within the hydrogen, hence the efficiency
appears lower than when using the more accurately defined values.
The theoretical maximum considers the total amount of energy required for the formation
of the hydrogen and oxygen from water.
Note that, these values refer only to the efficiency of converting electrical energy
into hydrogen's chemical energy; the energy lost in generating the electricity is not
included.
NREL estimated that 1 kg of hydrogen could be produced by wind powered electrolysis for
between $5.55 in the near term and $2.27 in the long term.
About 4% of hydrogen gas produced worldwide is created by electrolysis, and normally used
onsite.
Hydrogen is used for the creation of ammonia for fertilizer via the Haber process, and
converting heavy petroleum sources to lighter fractions via hydrocracking.
Electrocrystallization A specialized application of electrolysis
involves the growth of conductive crystals on one of the electrodes from oxidized or
reduced species that are generated in situ.
The technique has been used to obtain single crystals of low-dimensional electrical conductors,
such as charge-transfer salts.
Experimenters Scientific pioneers of electrolysis include:
Pioneers of batteries: Alessandro Volta
Gaston Planté More recently, electrolysis of heavy water
was performed by Fleischmann and Pons in their famous experiment, resulting in anomalous
heat generation and the discredited claim of cold fusion.
See also
References