Discover

'Electrolysis of water' is an electrolytic process which decomposes water into oxygen and hydrogen gas due to the flow of electric current. A DC voltage source, such as a battery, is commonly used to induce the flow of electrical current. The voltage of the battery creates a current in the water that is equal to the voltage of the battery divided by the resistance of the water, as per Ohm's law. For water to conduct a substantial electric current, an electrolyte is required to reduce resistance. An electrolysis cell can consist of an electrode or parallel plate design. The former utilizes two or more electrodes, (usually an inert metal such as platinum), submerged in water with electrolyte. The latter utilizes two or more plates, also usually an inert metal, with water situated between them, also with an electrolyte added.
The electric current disassociates water molecules into hydroxide (OH⁻) and hydrogen (H⁺) ions.
In the electrolytic cell, at the cathode (negatively charged electrode), hydrogen ions accept electrons in a reduction reaction that forms hydrogen gas:
: Cathode (reduction): 2(l) + 2e⁻ → (g) + 2(aq)
At the anode (positively charged electrode), hydroxide ions undergo an oxidation reaction and give up electrons to the anode to complete the circuit and form oxygen gas:
: Anode (oxidation): 2(l) → (g) + 4(aq) + 4e⁻
hence decomposing water into oxygen and hydrogen;
: Overall reaction: 2(l) → 2(g) + (g)
The number of hydrogen molecules produced is therefore twice the amount of oxygen molecules. Assuming equal temperature and pressure for both gases, the hydrogen gas has twice the quantity of moles as oxygen.

Contents

Spontaneity of the process

Electrolyte selection

Techniques

Fundamental Application

Hofmann voltameter

Industrial electrolysis

Electrolysis in nature

High-temperature electrolysis

Applications

Efficiency

See also

References

Spontaneity of the process

Decomposition of water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamical terms, as half of the reaction's standard potential are negative values...
:$mbox\{Anode\; (oxidation):\; \}2H\_\{2\}O(l)$
ightarrow O_{2}(g) + 4H^{+}(aq) + 4e^{-}qquad E^{o}_{ox}=-1.23 V,
:$mbox\{Cathode\; (reduction):\; \}2H\_\{2\}O(l)\; +\; 2e^\{-\}$
ightarrow H_{2}(g) + 2OH^{-}(aq)qquad E^{o}_{red}=-0.83 V,
... On the other hand, Gibbs free energy for the process at standard conditions is a higher positive value, about $474.4\; kJ,$. Those considerations makes the process ''"impossible"'' to occur without adding electrolytes in the solution.

Electrolyte selection

As pure water conducts electricity very poorly, a water-soluble electrolyte must be added to establish substantial conductivity. The electrolyte dissolves and disassociates into cations and anions (positive and negative ions) that carry the current. Electrolytes are normally acids, bases, or salts.
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 SO₄^2-, as it is very difficult to oxidize.
Standard potential for oxidation of this ion to the peroxydisulfate ion is −0.22 volts.
:$mbox\{Anode\; (oxidation):\; \}\; 2\; SO^\{2-\}\_\{4\}$
ightarrow S_{2}O^{2-}_{6} + 2e^{-} qquad E^{o}_{ox}=-0.22 V ,
Frequently used electrolytes:
Strong acids such as Sulphuric acid (H₂SO₄), and strong bases such as Potassium Hydroxide (KOH), and Sodium Hydroxide (NaOH) are frequently used as electrolytes.

Techniques

Fundamental Application

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

Main articles: 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. Hydrogen gas is usually created and collected on site for use in other chemical processes, although in case of refineries it then makes more sense to produce it from natural gas. It can also be produced as a by-product, for example in brine electrolysis. Electrolysis could be used in a hydrogen economy to produce hydrogen from e.g. solar power.

Electrolysis in nature

Plants electrolyze water in the process of photosynthesis utilizing a naturally occurring catalyst.
2 H₂O + 2 NADP+ + 2 ADP + 2 Pi + light → 2 NADPH + 2 H+ + 2 ATP + O₂

High-temperature electrolysis

Main articles: 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, 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. There is some speculation about future development of hydrogen as an energy carrier, although the rapid evolution of electric battery technology makes overall efficiency a major consideration. Hydrogen fuel injection is also a potentially viable application.

Efficiency

The energy efficiency of water electrolysis varies widely. Some report 50–70%[1], while others report 80–94%.[2] 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 25–45%.[3]

See also

★ Electrochemistry

★ Electrolysis

★ Hydrogen production

★ Gas cracker

References

★ Electrolysis of Water

★ Electrolysis of Water