Electrolysis of water 3206764 224919193 2008-07-11T00:21:27Z Lightbot 7178666 Units/dates/other <!-- Deleted image removed: [[Image:Hoffman voltameter.jpg|210px|thumb|right|[[Hoffman voltameter]] used to electrolyze water.{{puic|Image:Hoffman voltameter.jpg|log=2008 April 26}}]] --> '''Electrolysis of water''' is the decomposition of [[water]] (H<sub>2</sub>O) into [[oxygen]] (O<sub>2</sub>) and [[hydrogen]] gas (H<sub>2</sub>) due to an [[electricity|electric current]] being passed through the water. This [[electrolysis|electrolytic process]] is used in some industrial applications when hydrogen is needed. 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. Hydrogen will appear at the [[cathode]] (the negatively charged electrode, where [[electron]]s 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 (chemistry)|base]]). Historically, the first known electrolysis of water was done by [[William Nicholson (chemist)|William Nicholson]] and [[Anthony Carlisle]] in about 1800. == Equations == In the water at the negatively charged cathode, a [[redox|reduction]] reaction takes place, with electrons (e<sup>−</sup>) from the cathode being given to hydrogen cations to form hydrogen gas (the half reaction balanced with acid): : Cathode (reduction): 2{{hydrogen||+}}(''aq'') + 2e<sup>−</sup> &rarr; {{hydrogen|2}}(''g'') ; &nbsp;&nbsp; At the positively charged anode, an [[redox|oxidation]] reaction occurs, generating oxygen gas and giving electrons to the cathode to complete the circuit: : Anode (oxidation): 2[[water|{{ce|H|2}}{{ce|O}}]](''l'') &rarr; {{oxygen|2}}(''g'') + 4{{hydrogen||+}}(''aq'') + 4e<sup>−</sup>; &nbsp;&nbsp; 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|{{ce|H|2}}{{ce|O}}]](''l'') + 2e<sup>−</sup> &rarr; {{hydrogen|2}}(''g'') + 2{{hydroxide||−}}(''aq''); &nbsp;&nbsp; : Anode (oxidation): 4{{hydroxide||−}}(''aq'') &rarr; {{oxygen|2}}(''g'') + 2[[water|{{ce|H|2}}{{ce|O}}]](''l'') + 4e<sup>−</sup>; &nbsp;&nbsp; Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen: : Overall reaction: 2[[water|{{ce|H|2}}{{ce|O}}]](''l'') &rarr; 2{{hydrogen|2}}(''g'') + {{oxygen|2}}(''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 conditions for temperature and pressure|standard temperature and pressure]] is not favorable in [[thermochemistry|thermodynamical]] terms. This is because, E(cell)=E(Oxidation) + E(Reduction). If E(cell) < 0, reaction is not favorable. :<math> \begin{align} \mbox{Anode (oxidation): } &&\mbox{2H}_2\mbox{O (l)} &\rightarrow \mbox{O}_2\mbox{ (g)} + 4\mbox{H}^+ \mbox{ (aq)} + 4e^{-} &&E^o_\text{ox}=-1.23 \mbox{ V}\\ \mbox{Cathode (reduction): }&&\mbox{2H}^+ \mbox{ (aq)} + 2e^{-} &\rightarrow \mbox{H}_2\mbox{ (g)} &&E^o_\text{red}=0.00 \mbox{ V}\\ \end{align} </math> Thus, the standard potential of the water electrolysis cell is 1.23&nbsp;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 cannot occur without adding necessary energy, usually supplied by an external electrical power source but also possible with thermal energy. [[Image:Elektrolyse1.jpg|thumb|200px|right|Hoffman voltameter connected to a [[direct current]] [[Switched-mode power supply|power source converter]].]] == Electrolyte selection == If the above described processes occur in pure water, H<sup>+</sup> [[cation]]s will accumulate at the anode and OH<sup>−</sup> [[anion]]s 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 [[diffusion|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 [[cation]]s and [[anion]]s; the anions rush towards the anode and neutralize the buildup of positively charged H<sup>+</sup> there; similarly, the cations rush towards the cathode and neutralize the buildup of negatively charged OH<sup>−</sup> there. This allows the continued flow of electricity.<ref>[[Linus Pauling]], ''General Chemistry'', Section 15-2. San Francisco, 1970</ref> 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<sup>+</sup> and are therefore suitable for use as electrolyte cations: [[Lithium|Li]]<sup>+</sup>, [[Rubidium|Rb]]<sup>+</sup>, [[Potassium|K]]<sup>+</sup>, [[Caesium|Cs]]<sup>+</sup>, [[Barium|Ba]]<sup>2+</sup>, [[Strontium|Sr]]<sup>2+</sup>, [[Calcium|Ca]]<sup>2+</sup>, [[Sodium|Na]]<sup>+</sup>, and [[Magnesium|Mg]]<sup>2+</sup>. [[Sodium]] and [[lithium]] are frequently used, as they form inexpensive, soluble salts. If an [[acid]] is used as the [[electrolyte]], the cation is H<sup>+</sup>, and there is no competitor for the H<sup>+</sup> created by disassociating water. The most commonly used [[anion]] is [[sulfate]] (SO<sub>4</sub><sup>2-</sup>), as it is 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<sub>2</sub>SO<sub>4</sub>), and strong bases such as [[potassium hydroxide]] (KOH), and [[sodium hydroxide]] (NaOH) are frequently used as electrolytes. == Techniques == === Fundamental Demonstration === Two [[wire|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]]. [[Image:Elektrolyse2.jpg|thumb|150px|Match test used to detect the presence of hydrogen gas.]] === Hofmann voltameter === {{main|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. <!-- Unsourced image removed: [[Image:Electrolosis-water.jpg|thumb|140px|right|Hofmann voltameter as is used in laboratories.]] --> === Industrial electrolysis === Many industrial electrolysis cells are very similar to [[Hofmann voltameter]]s, 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 cell]]s. === High-temperature electrolysis === {{main|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#Production|chlorine]]. 2 NaCl + 2 H<sub>2</sub>O → Cl<sub>2</sub> + H<sub>2</sub> + 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<sub>2</sub> 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<sub>2</sub> with CO<sub>2</sub> 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 (electricity)|battery]] technology makes overall efficiency a major consideration. == 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 potential]]s. 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 [[overpotential#Reaction Overpotential|reaction overpotential]] for the four electron oxidation of water to oxygen at the anode. An effective [[catalyst#Electrocatalysts|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 cell]]s. 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 [[overpotential#Reaction Overpotential|reaction overpotential]] by platinum or in theory a [[Hydrogenase|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%<ref>{{cite web |title=Chapter 3: Production of Hydrogen. Part 4: Production from electricity by means of electrolysis |work=HyWeb: Knowledge - Hydrogen in the Energy Sector |url=http://www.hyweb.de/Knowledge/w-i-energiew-eng3.html#3.4 |author=Werner Zittel |coauthors=Reinhold Wurster |publisher=Ludwig-Bölkow-Systemtechnik GmbH |date=1996-07-08 |}}</ref>, while the theoretical maximum efficiency of the electrolysis of water is between 80–94%.<ref>{{cite web |url= http://www.bellona.org/reports/hydrogen |title= Hydrogen--Status and Possibilities |author= Bjørnar Kruse |coauthors= Sondre Grinna, Cato Buch |date= 2002-02-13 |format= pdf |publisher= The Bellona Foundation |archiveurl= http://www.bellona.org/filearchive/fil_Hydrogen_6-2002.pdf |archivedate= 2002-02-13 |quote=Efficiency factors for PEM electrolysers up to 94% are predicted, but this is only theoretical at this time. }}</ref> 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%.<ref>{{cite web |url= http://www.uic.com.au/nip73.htm |title= Briefing Paper #73: Transport and the Hydrogen Economy |date= January 2008 |publisher= Australian Uranium Association |work= Uranium Information Centre's Briefing Papers |quote= [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%. }}</ref><!-- See Talk:Electrolysis for widely varying values. Probably needs clarification of what exactly is being measured. --> == See also == *[[Electrochemistry]] *[[Electrolysis]] *[[Hydrogen production]] *[[Noryl]] *[[Gas cracker]] == References == <references/> *{{cite web|title=Electrolysis of Water|work=Experiments on Electrochemistry|url=http://www.pc.chemie.uni-siegen.de/pci/versuche/english/v21-2.html|accessdate=November 20|accessyear=2005}} *{{cite web|title=Electrolysis of Water|work=Do Chem 044|url=http://chemmovies.unl.edu/Chemistry/DoChem/DoChem044.html|accessdate=November 20|accessyear=2005}} == External links == *[http://ec.europa.eu/research/energy/pdf/efchp_hydrogen1.pdf High pressure electrolyzer] {{electrolysis}} {{Spoken Wikipedia|Electrolysis of water.ogg|2006-06-21}} *[http://www.HGenerators.com electrolysis hydrogen gas-from-water systems and processes] *[http://www.HGasStations.com onsite remote electrolysis hydrogen gas-from-water stations for deployment anywhere] [[Category:Electrolysis]] [[Category:Hydrogen production]] [[ar:كهرلة]] [[fr:Électrolyse de l'eau]] [[it:Elettrolisi dell'acqua]]