Electrolysis of water
Electrolysis of water is the decomposition of water (H2O) into oxygen (O2) and hydrogen gas (H2) due to an
electric current being passed through the water. This electrolytic process is rarely used in industrial
applications since hydrogen can be produced more affordably from fossil fuels.
An electrical power source is connected to two electrodes, or two plates, (typically made from some inert
metal such as platinum or stainless steel) which are placed in the water. In a properly designed cell
Hydrogen will appear at the cathode (the negatively charged electrode, where electrons are pumped into the
water), and oxygen will appear at the anode (the positively charged electrode). The generated amount of
hydrogen is twice the amount of oxygen, and both are proportional to the total electrical charge that was sent
through the water. However, in many cells competing side reactions dominate resulting in different products.
Electrolysis of pure water requires a great deal of excess energy in the form of over potential to overcome
various activation barriers. Without the excess energy the electrolysis of pure water occurs very slowly if at
all. This is in part due to the limited self-ionization of water. Pure water has an electrical conductivity about
one millionth that of seawater. Many electrolytic cells may also lack the requisite electrocatalysts. The
efficacy of electrolysis is increased through the addition of an electrolyte (such as a salt, an acid or a base)
and the use of electrocatalysts.
Thermodynamics of the process
Decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not
favorable in thermodynamical terms.
Thus, the standard potential of the water electrolysis cell is -1.23 V at 25 °C at pH 0 (H+ = 1.0 M). It is also
-1.23 V at 25 °C at pH 7 (H+ = 1.0x10-7 M) based on the Nernst Equation.
The negative voltage indicates the Gibbs free energy for electrolysis of water is greater than zero for these
reactions. This can be found using the G=-nFE equation from chemical kinetics, where n is the moles of
electrons and F is the Faraday constant. The reaction cannot occur without adding necessary energy,
usually supplied by an external electrical power source.
If the above described processes occur in pure water, H+ cations will accumulate at the anode and OH−
anions will accumulate at the cathode. This can be verified by adding a pH indicator to the water: the water
near the anode is acidic while the water near the cathode is basic. These charged ions will repel the further
flow of electricity until they have diffused away, a slow process. This is why pure water conducts electricity
poorly and why electrolysis of pure water proceeds slowly.
If a water-soluble electrolyte is added, the conductivity of the water rises considerably. The electrolyte
disassociates into cations and anions; the anions rush towards the anode and neutralize the buildup of
positively charged H+ there; similarly, the cations rush towards the cathode and neutralize the buildup of
negatively charged OH− there. This allows the continued flow of electricity.
Care must be taken in choosing an electrolyte, since an anion from the electrolyte is in competition with the
hydroxide ions to give up an electron. An electrolyte anion with less standard electrode potential than
hydroxide will be oxidized instead of the hydroxide, and no oxygen gas will be produced. A cation with a
greater standard electrode potential than a hydrogen ion will be reduced in its stead, and no hydrogen gas
will be produced.
The following cations have lower electrode potential than H+ and are therefore suitable for use as
electrolyte cations: Li+, Rb+, K+, Cs+, Ba2+, Sr2+, Ca2+, Na+, and Mg2+. Sodium and lithium are frequently
used, as they form inexpensive, soluble salts.
If an acid is used as the electrolyte, the cation is H+, and there is no competitor for the H+ created by
disassociating water. The most commonly used anion is sulfate (SO42-), as it is very difficult to oxidize, with
the standard potential for oxidation of this ion to the peroxodisulfate ion being −0.22 volts.
Strong acids such as sulfuric acid (H2SO4), and strong bases such as potassium hydroxide (KOH), and
sodium hydroxide (NaOH) are frequently used as electrolytes.
A solid polymer electrolyte can also be used such as NAFION and when applied with a special catalyst on
each side of the membrane can efficiently split the water molecule with as little as 1.8 Volts.
About four percent of hydrogen gas produced worldwide is created by electrolysis. The majority of this
hydrogen produced through electrolysis is a side product in the production of chlorine. This is a prime
example of a competing side reactions.
2 NaCl + 2 H2O → Cl2 + H2 + 2 NaOH
The electrolysis of brine (saltwater), a water sodium chloride mixture, is only half the electrolysis of water
since the chloride ions are oxidized to chlorine rather than water being oxidized to oxygen. The hydrogen
produced from this process is either burned, used for the production of specialty chemicals, or various other
small scale applications.
The majority of hydrogen used industrially is derived from fossil fuels. One example is fossil fuel derived
hydrogen used for the creation of ammonia for fertilizer via the Haber process and for converting heavy
petroleum sources to lighter fractions via hydrocracking. The production of this hydrogen usually involves
the formation of synthesis gas a mixture of H2 and CO. Synthesis gas can be hydrogen enriched through
the water gas shift reaction. In this reaction the carbon monoxide is reacted with water to produce more H2
with CO2 byproduct.
Water electrolysis does not convert 100% of the electrical energy into the chemical energy of hydrogen. The
process requires more extreme potentials than what would be expected based on the cell's total reversible
reduction potentials. This excess potential accounts for various forms of overpotential by which the extra
energy is eventually lost as heat. For a well designed cell the largest overpotential is the reaction
overpotential for the four electron oxidation of water to oxygen at the anode. An effective electrocatalyst to
facilitate this reaction has not been developed. Platinum alloys are the default state of the art for this
oxidation. The reverse reaction, the reduction of oxygen to water, is responsible for the greatest loss of
efficiency in fuel cells. Developing a cheap effective electrocatalyst for this reaction would be a great
The simpler two-electron reaction to produce hydrogen at the cathode can be electrocatalyzed with almost
no reaction overpotential by platinum or in theory a hydrogenase enzyme. If other, less effective, materials
are used for the cathode then another large overpotential must be paid.
The energy efficiency of water electrolysis varies widely with the numbers cited below on the optimistic side.
Some report 50–80% 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. For instance, when
considering a power plant that converts the heat of nuclear reactions into hydrogen via electrolysis, the total
efficiency may be closer to 30–45%.
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