Everything about Electrolysis Of Water totally explained
Electrolysis of water is the decomposition of
water (H
2O) into
oxygen (O
2) and
hydrogen gas (H
2) due to an
electric current being passed through the water. This
electrolytic process is used in some industrial applications when hydrogen is needed.
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.
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.
Electrolysis of
pure water is very slow, and can only occur due to the
self-ionization of water. Pure water has an
electrical conductivity about one millionth that of seawater. It is sped up dramatically by adding an
electrolyte (such as a
salt, an
acid or a
base).
Historically, the first known electrolysis of water was done by
William Nicholson and
Anthony Carlisle in about 1800.
Equations
In the water at the negatively charged cathode, a
reduction reaction takes place, with electrons (e
−) from the cathode being given to hydrogen cations to form hydrogen gas (the half reaction balanced with acid):
» Cathode (reduction): 2(
aq) + 2e
− → (
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[[water|]](
l) → (
g) + 4(
aq) + 4e
−;
The same half reactions can also be balanced with base as listed below. Not all 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[[water|]](
l) + 2e
− → (
g) + 2(
aq);
» Anode (oxidation): 4(
aq) → (
g) + 2[[water|]](
l) + 4e
−;
Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:
» Overall reaction: 2[[water|]](
l) → 2(
g) + (
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.
Thermodynamics of the process
Decomposition of pure
water into
hydrogen and
oxygen at
standard temperature and pressure isn't favorable in
thermodynamical terms. This is because, E(cell)=E(Oxidation) + E(Reduction).
If E(cell) < 0, reaction isn't favorable.
»
Thus, the standard potential of the water electrolysis cell is 1.23 V at 25 °C.
The positive voltage indicates the
Gibbs Free Energy for electrolysis of water is greater than zero for these reactions. This can be found using the
Nernst Equation at equilibrium. The reaction can't occur without adding necessary energy, usually supplied by an external electrical power source but also possible with thermal energy.
Electrolyte selection
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've
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's no competitor for the H
+ created by disassociating water. The most commonly used
anion is
sulfate (SO
42-), as it's very difficult to oxidize, with the standard potential for oxidation of this ion to the
peroxydisulfate ion being −0.22 volts.
Strong acids such as
sulfuric acid (H
2SO
4), and strong bases such as
potassium hydroxide (KOH), and
sodium hydroxide (NaOH) are frequently used as electrolytes.
Techniques
Proton Exhange Membrane Water Electrolysis
A
proton conducting membrane can also be used as to separate the
anode from the
cathode chamber. The typical membrane used is
NAFION and can be
catalyst coated to form a HYDRion® membrane.
Fundamental Demonstration
Two
leads, running from the terminals of a battery, are placed in a cup of water with a quantity of electrolyte added to establish conductivity. Hydrogen and Oxygen gases will stream from the oppositely charged
electrode. Oxygen will collect at the
anode and hydrogen will collect at the
cathode.
Hofmann voltameter
The Hofmann voltameter is often used as a small-scale electrolytic cell. It consists of three joined upright cylinders. The inner cylinder is open at the top to allow the addition of
water and the
electrolyte. A
platinum electrode is placed at the bottom of each of the two side cylinders, connected to the positive and negative terminals of a source of
electricity. When current is run through the hofmann voltameter, gaseous
oxygen forms at the
anode and gaseous
hydrogen at the
cathode. Each gas displaces water and collects at the top of the two outer tubes, where it can be drawn off with a stopcock.
Industrial electrolysis
Many industrial electrolysis cells are very similar to
Hofmann voltameters, with complex platinum plates or honeycombs as electrodes. Generally the only time hydrogen is intentionally produced from electrolysis is for specific point of use application such as is the case with
oxyhydrogen torches or when extremely high purity hydrogen or oxygen is desired. The vast majority of hydrogen is produced from hydrocarbons and as a result contains trace amounts of
carbon monoxide among other impurities. The carbon monoxide impurity can be detrimental to various systems including many
fuel cells.
High-temperature electrolysis
High-temperature electrolysis (also HTE or steam electrolysis) is a method currently being investigated for water electrolysis with a
heat engine. High temperature electrolysis is more efficient than traditional room-temperature electrolysis because some of the energy is supplied as heat, which is cheaper than electricity, and because the electrolysis reaction is more efficient at higher temperatures.
Applications
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.
2 NaCl + 2 H
2O → Cl
2 + H
2 + 2 NaOH
The electrolysis of
brine, 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 H
2 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 H
2 with CO
2 byproduct.
There is some speculation about future development of hydrogen as an energy carrier in a
hydrogen economy, although the rapid evolution of electric
battery technology makes overall efficiency a major consideration.
Efficiency
Water electrolysis doesn't 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 hasn't 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 advance.
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 50–70%, while the theoretical maximum efficiency of the electrolysis of water is between 80–94%.
These values refer only to the efficiency of converting electrical energy into hydrogen's chemical energy. The energy lost in generating the electricity isn't 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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