Electrolysis of water: Difference between revisions
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==Principle== |
==Principle== |
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An electrical power source is connected to two [[electrode]]s, 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 [[ |
An electrical power source is connected to two [[electrode]]s, 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 [[anode]] (the negatively charged electrode, where [[electron]]s enter the water), and oxygen will appear at the [[cathode]] (the positively charged electrode). Assuming ideal [[faradaic efficiency]], the amount of hydrogen generated is twice the number of [[Mole (unit)|moles]] of oxygen, and both are [[Proportionality (mathematics)|proportional]] to the total [[electrical charge]] conducted by the solution. However, in many cells [[electrolysis#Competing half-reactions in solution electrolysis|competing side reactions]] dominate, resulting in different products and less than ideal faradaic efficiency. |
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[[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 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 cell]]s may also lack the requisite [[electrocatalyst]]s. The efficacy of electrolysis is increased through the addition of an [[electrolyte]] (such as a [[Salt (chemistry)|salt]], an [[acid]] or a [[base (chemistry)|base]]) and the use of [[electrocatalyst]]s. |
[[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 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 cell]]s may also lack the requisite [[electrocatalyst]]s. The efficacy of electrolysis is increased through the addition of an [[electrolyte]] (such as a [[Salt (chemistry)|salt]], an [[acid]] or a [[base (chemistry)|base]]) and the use of [[electrocatalyst]]s. |
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Revision as of 10:05, 10 September 2010
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.
Principle
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 anode (the negatively charged electrode, where electrons enter the water), and oxygen will appear at the cathode (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 cells 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 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.
Currently the electrolytic process is rarely used in industrial applications since hydrogen can be produced more affordably from fossil fuels.
History
Jan Rudolph Deiman and Adriaan Paets van Troostwijk used in 1789 an electrostatic machine to produce electricity which was discharged on gold electrodes in a Leyden jar with water.[1] In 1800 Alessandro Volta invented the voltaic pile, a few weeks later William Nicholson and Anthony Carlisle used it for the electrolysis of water. When Zénobe Gramme invented the Gramme machine in 1869 electrolysis of water became a cheap method for the production of hydrogen. A method of industrial synthesis of hydrogen and oxygen through electrolysis was developed by Dmitry Lachinov in 1888.[2]
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):
- 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 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 H2O(l) + 2e− → H2(g) + 2 OH-(aq)
- Anode (oxidation): 2 H2O (l) → O2(g) + 4 H++ 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.
Thermodynamics of the process
Decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamical terms.
- Anode (oxidation): 2 H2O(l) → O2(g) + 4 H+(aq) + 4e− Eo
ox = -1.23 V - Cathode (reduction): 2 H+(aq) + 2e− → H2(g) Eo
red = 0.00 V
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+ = 1.0×10−7 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.
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. The negative hydroxyl ions that approach the anode mostly combine with the positive hydronium ions (H3O+) to form water. The positive hydronium ions that approach the negative cathode mostly combine with negative hydroxyl ions to form water. Relatively few hydronium (hydroxyl) ions reach the cathode (anode). This can cause a concentration overpotential at both electrodes.
Pure water is a fairly good insulator since it has a low autoionization, Kw = 10 x 10-14 at room temperature and thus pure water conducts current poorly, 0.055 µS·cm−1. Unless a very large potential is applied to cause an increase in the autoionization of water the electrolysis of pure water proceeds very slowly limited by the overall conductivity.
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.[3]
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 (SO2−
4), as it is very difficult to oxidize, with the standard potential for oxidation of this ion to the peroxodisulfate ion being −2.05 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.
Techniques
Fundamental demonstration
Two leads, running from the terminals of a battery, are placed in a cup of water with a quantity of electrolyte to establish conductivity in the solution. Using NaCl as an electrolyte results in chlorine gas rather than oxygen due to a competing half-reaction. With the correct electrodes and correct electrolyte hydrogen and oxygen gases will stream from the oppositely charged electrodes. 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 hydrogen purity 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 pressure electrolysis
High pressure electrolysis is the electrolysis of water with a compressed hydrogen output around 120-200 Bar (1740-2900 psi).[4] By pressurising the hydrogen in the electrolyser the need for an external hydrogen compressor is eliminated, the average energy consumption for internal compression is around 3%.[5]
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.[6][7]
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. This is a prime example of a competing side reaction.
- 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.
Efficiency
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. Developing a cheap effective electrocatalyst for this reaction would be a great advance (see also[8]). In 2008, a group led by Daniel Nocera announced a potentially more efficient catalyst for artificial photosynthesis composed of a cobalt metal, phosphate, and an electrode.[9] Other researchers are pursuing carbon-based catalysts.[10] The reverse reaction, the reduction of oxygen to water, is responsible for the greatest loss of efficiency in fuel cells.
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%.[11][12][13] 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%,[14] although the inefficiencies of powerplants in turning heat into electrical energy is not usually included in efficiency, so the former measure of 50-80% efficient is probably a more realistic efficiency.
See also
References
- ^ Levie, R. de (1999-10). "The electrolysis of water". Journal of Electroanalytical Chemistry. 476 (1): 92–93. doi:10.1016/S0022-0728(99)00365-4. Retrieved 2009-12-30.
{{cite journal}}: Check date values in:|date=(help) - ^ Lachinov Dmitry Aleksandrovich at Great Cyrill and Methodius Encyclopedia Template:Ru icon
- ^ Linus Pauling, General Chemistry, Section 15-2. San Francisco, 1970
- ^ 2001-High pressure electrolysis - The key technology for efficient H.2
- ^ 2003-PHOEBUS-Pag.9
- ^ Hi2h2 - High temperature electrolysis using SOEC
- ^ WELTEMP-Water electrolysis at elevated temperatures
- ^ Kanan, M. W. (2008). "In Situ Formation of an Oxygen-Evolving Catalyst in Neutral Water Containing Phosphate and Co2+". Science. 321: 1072. doi:10.1126/science.1162018.
{{cite journal}}:|first2=missing|last2=(help) - ^ 'Major discovery' from MIT primed to unleash solar revolution July 31, 2008
- ^ Cheap Catalyst Could Turn Sunlight, Water Into Fuel July 31, 2008
- ^ Werner Zittel (1996-07-08). "Chapter 3: Production of Hydrogen. Part 4: Production from electricity by means of electrolysis". HyWeb: Knowledge - Hydrogen in the Energy Sector. Ludwig-Bölkow-Systemtechnik GmbH.
{{cite web}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - ^ Bjørnar Kruse (2002-02-13). "Hydrogen—Status and Possibilities" (PDF). The Bellona Foundation. Archived from the original (pdf) on 2002-02-13.
Efficiency factors for PEM electrolysers up to 94% are predicted, but this is only theoretical at this time.
{{cite web}}: Unknown parameter|coauthors=ignored (|author=suggested) (help) - ^ high-rate and high efficiency 3D water electrolysis
- ^ "Briefing Paper #73: Transport and the Hydrogen Economy". Uranium Information Centre's Briefing Papers. Australian Uranium Association. January 2008.
[Fuel cells] catalyse the oxidation of hydrogen directly to electricity at relatively low temperatures and the claimed theoretical efficiency of converting chemical to electrical energy to drive the wheels is about 60% (or more). However, in practice about half that has been achieved, except for the higher-temperature solid oxide fuel cells - 46%.
External links
- "Electrolysis of Water". Experiments on Electrochemistry. Retrieved November 20, 2005.
- "Electrolysis of Water". Do Chem 044. Retrieved November 20, 2005.
- EERE 2008 - 100 kgH2/day Trade Study
- NREL 2006 - Electrolysis technical report