# Electrolysis of Water

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Electrolysis of Water
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. An electrical power source is
connected to two electrodes, or two plates (typical y made from some inert metal such as
platinum or stainless steel) which are placed in the water.
Hydrogen wil appear at the cathode (the negatively charged electrode, where electrons enter
the water), and oxygen wil appear at the anode (the positively charged electrode). Assuming
ideal faradaic efficiency, the amount of hydrogen generated is twice the number of moles of
oxygen, and both are proportional to the total electrical charge conducted by the solution.
However, in many cel s competing side reactions dominate, resulting in different products and
less than ideal faradaic efficiency. Electrolysis of pure water requires excess energy in the
form of overpotential to overcome various activation barriers.
Without the excess energy the electrolysis of pure water occurs very slowly or not at al . This
is in part due to the limited self-ionization of water. Pure water has an electrical conductivity
about one mil ionth that of seawater. Many electrolytic cells may also lack the requisite
electrocatalysts.
Know More About :- Net Ionic Equations

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Equations Electrolysis of water :- In pure water at the negatively charged cathode, a
reduction reactioOverall reaction: n takes place, with electrons (e-) from the cathode being
given to hydrogen cations to form hydrogen gas (the half reaction balanced with acid):
Reduction at cathode: 2 H+(aq) + 2e- H2(g)
At the positively charged anode, an oxidation reaction occurs, generating oxygen gas and
giving electrons to the anode to complete the circuit:
Anode (oxidation): 2 H2O(l) O2(g) + 4 H+(aq) + 4e-
The same half reactions can also be balanced with base as listed below. Not al half reactions
must be balanced with acid or base. Many do, like the oxidation or reduction of water listed
here. To add half reactions they must both be balanced with either acid or base.
Cathode (reduction): 2 H2O(l) + 2e- H2(g) + 2 OH-(aq)
Anode (oxidation): 4 OH- (aq) O2(g) + 2 H2O(l) + 4 e-
Combining either half reaction pair yields the same overall decomposition of water into oxygen
and hydrogen:
Overall reaction: 2 H2O(l) 2 H2(g) + O2(g)
The number of hydrogen molecules produced is thus twice the number of oxygen molecules.
Assuming equal temperature and pressure for both gases, the produced hydrogen gas has
therefore twice the volume of the produced oxygen gas.
The number of electrons pushed through the water is twice the number of generated hydrogen
molecules and four times the number of generated oxygen molecules.

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Thermodynamics of the process :- Decomposition of pure water into hydrogen and oxygen
at standard temperature and pressure is not favorable in thermodynamic terms.
Anode (oxidation): 2 H2O(l) O2(g) + 4 H+(aq) + 4e- Eo
ox = -1.23 V (Eo
red = 1.23 )[4])
Cathode (reduction): 2 H+(aq) + 2e- H2(g) Eo
red = 0.00 V
Thus, the standard potential of the water electrolysis cel is -1.23 V at 25 C at pH 0 (H+ = 1.0
M). The potential is changed to -0.82 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.

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