Water splitting

Diagram of the chemical equation of the electrolysis of water, a form of water splitting.

Water splitting is the general term for a chemical reaction in which water is separated into oxygen and hydrogen. Efficient and economical water splitting would be a key technological component of a hydrogen economy. Various techniques for water splitting have been issued in water splitting patents in the United States.[1] In photosynthesis, water splitting donates electrons to the electron transport chain in photosystem II.

Electrolysis

Main article: Electrolysis of water
Electrolyser front with electrical panel in foreground

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. In chemistry and manufacturing, electrolysis is a method of separating chemically bonded elements and compounds by passing an electric current through them. One use of electrolysis of water or artificial photosynthesis (photoelectrolysis in a photoelectrochemical cell) is to produce hydrogen. Recently, researchers have shown that water splitting can be broken into two discrete steps using polyoxometalate based redox mediators.[2]

Power to gas production schemes, the excess power or off peak power created by wind generators or solar arrays is used for load balancing of the energy grid by storing and later injecting the hydrogen into the natural gas grid.

Electrolysis of water ship Hydrogen Challenger

Production of hydrogen from water requires large amounts of energy and is uncompetitive with its production from coal or natural gas. Potential electrical energy supplies include hydropower, wind turbines, or photovoltaic cells. Usually, the electricity consumed is more valuable than the hydrogen produced so this method has not been widely used. Other potential energy supplies include heat from nuclear reactors and light from the sun. Hydrogen can also be used to store renewable generated electricity when it is not needed (like the wind blowing at night) and then used to meet power needs during the day or to fuel vehicles. The storable quality of hydrogen helps make hydrogen an enabler of the wider use of renewables,[3] and internal combustion engines. (See hydrogen economy.)

High pressure electrolysis

When water is pressurized and then electrolysis is conducted at those high pressures, the produced hydrogen gas is pre-compressed at around 120–200 bar (1740–2900 psi).[4] By pre-pressurising the hydrogen in the electrolyser energy is saved as the need for an external hydrogen compressor is eliminated, the average energy consumption for internal compression is around 3%.[5] The energy required to compress water is very much less than that required to compress hydrogen gas.

High-temperature electrolysis

When the energy supplied is in the form of heat (originating from solar thermal, or nuclear), the best path to the production of hydrogen is through high-temperature electrolysis (HTE). In contrast with low-temperature electrolysis, HTE of water converts more of the initial heat energy into chemical energy (hydrogen), potentially doubling efficiency to about 50%. Because some of the energy in HTE is supplied in the form of heat, less of the energy must be converted twice (from heat to electricity, and then to chemical form), and so the process is more efficient.

HTE processes are generally only considered in combination with a nuclear heat source, because the other non-chemical form of high-temperature heat (concentrating solar thermal) is not consistent enough to bring down the capital costs of the HTE equipment. Research into HTE and high-temperature nuclear reactors may eventually lead to a hydrogen supply that is cost-competitive with natural gas steam reforming. HTE has been demonstrated in a laboratory, but not at a commercial scale.

Photoelectrochemical water splitting

Using electricity produced by photovoltaic systems potentially offers the cleanest way to produce hydrogen. Again, water is broken down into hydrogen and oxygen by electrolysis, but the electrical energy is obtained by a photoelectrochemical cell (PEC) process. The system is also named artificial photosynthesis.[6][7][8][9]

Photocatalytic water splitting

The conversion of solar energy to hydrogen by means of water splitting process is one of the most interesting ways to achieve clean and renewable energy. However, if this process is assisted by photocatalysts suspended directly in water rather than a photovoltaic or an electrolytic system, the reaction takes place in one step, it therefore can be more efficient.[10][11]

Radiolysis

Nuclear radiation routinely breaks water bonds, in the Mponeng gold mine, South Africa, researchers found in a naturally high radiation zone, a community dominated by a new phylotype of Desulfotomaculum, feeding on primarily radiolytically produced H2.[12] Spent nuclear fuel/"nuclear waste" is also being looked at as a potential source of hydrogen.

Photobiological water splitting

Biological hydrogen can be produced in an algae bioreactor. In the late 1990s it was discovered that if the algae are deprived of sulfur it will switch from the production of oxygen, i.e. normal photosynthesis, to the production of hydrogen. It seems that the production is now economically feasible by surpassing the 7–10 percent energy efficiency (the conversion of sunlight into hydrogen) barrier.[13] with a hydrogen production rate of 10-12 ml per liter culture per hour.[14]

Thermal decomposition of water

Main article: Thermochemical cycle

Thermal decomposition, also called thermolysis, is defined as a chemical reaction whereby a chemical substance breaks up into at least two chemical substances when heated. At elevated temperatures water molecules split into their atomic components hydrogen and oxygen. For example, at 2200 °C about three percent of all H2O molecules are dissociated into various combinations of hydrogen and oxygen atoms, mostly H, H2, O, O2, and OH. Other reaction products like H2O2 or HO2 remain minor. At the very high temperature of 3000 °C more than half of the water molecules are decomposed, but at ambient temperatures only one molecule in 100 trillion dissociates by the effect of heat.

Thermal water splitting has been investigated for hydrogen production since the 1960s.[15] The high temperatures needed to obtain substantial amounts of hydrogen impose severe requirements on the materials used in any thermal water splitting device. For industrial or commercial application, the material constraints have limited the success of applications for hydrogen production from direct thermal water splitting and with few exceptions most recent developments are in the area of the catalysis and thermochemical cycles.

Nuclear-thermal

See also: NGNP

Some prototype Generation IV reactors, such as the HTTR, operate at 850 to 1000 degrees Celsius, considerably hotter than existing commercial nuclear power plants. General Atomics predicts that hydrogen produced in a High Temperature Gas Cooled Reactor (HTGR) would cost $1.53/kg. In 2003, steam reforming of natural gas yielded hydrogen at $1.40/kg. At 2005 gas prices, hydrogen cost $2.70/kg. Hence, just within the United States, a savings of tens of billions of dollars per year is possible with a nuclear-powered supply. Much of this savings would translate into reduced oil and natural gas imports.

One side benefit of a nuclear reactor that produces both electricity and hydrogen is that it can shift production between the two. For instance, the plant might produce electricity during the day and hydrogen at night, matching its electrical generation profile to the daily variation in demand. If the hydrogen can be produced economically, this scheme would compete favorably with existing grid energy storage schemes. What is more, there is sufficient hydrogen demand in the United States that all daily peak generation could be handled by such plants.[16]

Recent research on the hybrid thermoelectric Copper-chlorine cycle has focused on a cogeneration system using the waste heat from nuclear reactors, specifically the CANDU supercritical water reactor.[17]

Solar-thermal

The high temperatures necessary to split water can be achieved through the use of concentrating solar power. Hydrosol-2 is a 100-kilowatt pilot plant at the Plataforma Solar de Almería in Spain which uses sunlight to obtain the required 800 to 1,200 °C to split water. Hydrosol II has been in operation since 2008. The design of this 100-kilowatt pilot plant is based on a modular concept. As a result, it may be possible that this technology could be readily scaled up to megawatt range by multiplying the available reactor units and by connecting the plant to heliostat fields (fields of sun-tracking mirrors) of a suitable size.[18]

An interesting approach to solar thermal hydrogen production is proposed by H2 Power Systems.[19] Material constraints due to the required high temperatures above 2200 °C are reduced by the design of a membrane reactor with simultaneous extraction of hydrogen and oxygen that exploits a defined thermal gradient and the fast diffusion of hydrogen. With concentrated sunlight as heat source and only water in the reaction chamber, the produced gases are very clean with the only possible contaminant being water. A "Solar Water Cracker" with a concentrator of about 100 m² can produce almost one kilogram of hydrogen per sunshine hour.

Water Splitting by Coordination Compounds

Huge energy requirements, the expected shortage of petroleum in the future and quick rise in pollution are the problems that need to be addressed by putting more efforts into investigating clean and sustainable energy resources. In pursuit of such energy sources, efforts are being put in light driven splitting of water into O2 and H2 in an attempt to convert solar energy into fuel.[20][21][22] Water oxidation (2H2O → 4H+ +4e− +O2) is the first important step providing the necessary electrons and protons for the next step (proton reduction) in which hydrogen production takes place in a catalysis reaction by a proton reduction catalyst. 

                                              2H2O + hʋ → 2H2 + O2                   ∆G°=4.92eV (113 kcal mol−1)
                                              4H+ + 4e− → 2H2
                                              2H2O → 4H+ +4e− +O2

The water oxidation step has been considered the bottleneck of this process, so the designing of highly active and robust water oxidation catalysts (WOCs) is an important step in the development of light-driven water splitting.[23][24][25] Water oxidation catalysts minimize the overpotential and increase the reaction rate. An ideal WOC is required to possess low overpotential, high stability, high activity/efficiency, low toxicity and low cost.

General steps in water oxidation catalysis

First step in the water oxidation catalysis by WOCs is the transfer of electron from the catalyst to the oxidizing agent, which can be enhanced by introducing strong electron donating ligands in the catalyst. Second step involves the water oxidation for which the redox potential of the catalyst should have a high oxidizing power at the highest oxidation state. The third step which is the O-O bond formation has been investigated to proceed by two mechanisms: water nucleophilic attack (WNA), in which water attacks the oxo group of the metal complex, resulting in two electron reduction of the metal and O2 formation; and interaction of two M-O species (12M) which leads to the formation of peroxo-intermediate. The last step is the release of the molecular oxygen which is a slow step.

Water splitting in photosynthesis

In nature, the solar energy is converted to chemical energy by the process of photosynthesis producing energy essential for the survival of life on earth. Photosynthesis works with two photosystems: Photosystems I (P700) and Photosystem II (P680). When photosystem I gets photo-excited, electron transfer reactions gets initiated, which results in reduction of a series of electron acceptors, eventually reducing NADP+ to NADPH and PS I is oxidized. The oxidized photosystem I captures electrons from photosystem II through a series of steps involving agents like plastoquinone, cytochromes and plastocyanine. The photosystem II then brings about water oxidation resulting in evolution of oxygen, the reaction being catalyzed by CaMn4O5 clusters embedded in complex protein environment; the complex is known as oxygen evolving complex (OEC).[26][27][28]

Water splitting by some metal complexes

Water splitting by coordination complexes got its inspiration from the natural photosynthesis. These coordination complexes are used as catalysts in water splitting or water oxidation and hence being called water oxidation catalysts.[29] There are a number of synthetic metal complexes used for this purpose; a brief review is given below.

Water splitting by Ruthenium complexes

Among a number of WOCs known to us, ruthenium based complexes are known to exhibit high turnover numbers in catalytic activity. A number of mononuclear Ru catalysts have been reported to be efficient catalysts, which have caused us to believe that multicentered catalysts are not necessary for robust catalytic activity. Tanaka, Meyer, Llobet, and their co-workers have reported the synthesis of some ruthenium-aqua complexes which are active in water oxidation, and the catalytic mechanisms have been studied extensively mainly based on the six-coordinate ruthenium models.[30][31][32] Llobet and coworkers investigated [(Ru(terpy)(H2O))2(µ-bpp))]3+ (where terpy=2,2':6',2-terpyridine; bpp=2,6-bis(pyridyl)pyrazolate)[31] A number of non-aqua complexes (trans-[Ru(pbn)(4-R-py)2(OH2)]2+ where (pbn = 2,2'-(4-(tert-butyl)pyridine-2,6-diyl)bis(1,8-naphthyridine); py = pyridine; R = CH3, CF3 and N(CH3)2) have been reported by Thummel et al.[33] and shown to be highly active in the process. Duan et al.[34] synthesized a mononuclear [RuIII(bda)(pic)2] complex 1 (pic= 4-picoline) and isolated a rare 7-coordinated RuIV dimeric intermediate 2 (µ-(HOHOH) [RuIV(bda)(pic)2]2(PF6)3•2H2O) during its water oxidation catalytic reaction. The intermediate was found to have a bridging ligand [HOHOH]- in addition to two water molecules, which appears in the form of H-bonding with the intermediate. Strong electron donating ability of carboxylate groups in the complex causes more lowering of the oxidation potential of the complex 1 as compared with the complex [RuII(L′)(pic)2]2+ (L′= 2,9-di(pyrid-2′-yl)-1,10-phenanthroline) reported by Thummel et al.[35] Chemical investigation of the catalytic process involved a huge liberation of oxygen when the complex was added to an aqueous triflic acid solution with a CeIV oxidant. Duan et al.[36] also reported the synthesis of another mononuclear ruthenium complex [Ru(bda)(isoq)2] (3; H2bda= 2,2′-bipyridine-6,6′-dicarboxylic acid; isoq= isoquinoline) which has improved catalytic activity (TOF= 303 s−1) towards water oxidation as compared with another structurally similar mononuclear [Ru(bda)(pic)2] complex 1. Highly polarizable π-system of isoquinolines on the axial positions in 3 reduces the barrier for radical dimerization of Ru—O because of non-covalent attractive interactions between them as compared with the 4-picolines in complex 2 which make it a better catalyst. The catalytic activity is comparable to that of the photosynthetic photosystem II, the driving force being CeIV instead of visible light for the complex.

Water splitting by Nickel complexes

Nickel based homogenous catalysts are hard to find in the literature but they have been used as alkaline electrolyzers on the commercial scale.[37] A Ni-based oxide film has been synthesized which liberates oxygen in quasi-neutral conditions at an overpotential of ~425mV and shows long lasting stability.[38] The X-ray spectroscopy revealed the presence of di-µ-oxide bridging between NiIII/NiIV ions but no evidence of mono-µ-oxide bridging was found between the ions.[39] Similar structures can be found in Co-WOC films and Mn-WOC catalysts.[40][41]

Water splitting by Cobalt complexes

Cobalt salts can be used as WOCs,[42] but their water oxidation capability decreases with the passage of time because of the precipitation of the salt from the homogenous solutions, resulting in turbidity. Kanan and Nocera[43] have demonstrated the formation of a dark thin film over the surface of indium tin oxide which was used an electrode, along with bubbling in the solution. These bubbles were confirmed to be of oxygen. Co-WOCs have been shown to work very efficiently at a pH 7.0. Heterogenous cobalt oxide(Co3O4) have been investigated to work on the same pattern as other cobalt salts.[44] No mechanism has been reported so far for these WOCs. A homogenous cobalt polyoxametalate complex [Co4(H2O)2(α-PW9O34)2]10- have been described to be highly efficient as compared to the reported heterogenous cobalt complexes reported by Yin et al.[45] Zidki et al.[46] prepared stable WOCs by adsorbing CoII as Co(OH)2 clusters on silica nanoparticles which exhibited a high water oxidation activity. A homogeneous WOC [Co(Py5)(H2O)](ClO4)2 was reported by Waylenko and co-workers.[47] The complex undergoes a proton-coupled electron transfer to form a stable [CoIII--OH]2+ specie which on further oxidation forms a CoIV intermediate. The intermediate formed reacts with water to liberate O2.

Water splitting by Iron complexes

Julio and co-workers[48] demonstrated that common iron complexes catalyze water oxidation reactions to produce O2. A water soluble complex [Fe(OTf)2(Me2Pytacn)] (Pytacn=pyridine triazacyclononane; OTf= CF3SO−
3) was reported to be a very efficient catalyst for water oxidation when an oxidant was added and hence liberated oxygen bubbles. The concentration of the catalyst and the oxidant were found to be strongly affecting the oxidation process. Catalytic activity of the complexes [FeCl2(mcp)], [Fe(OTf)2(mcp)], [Fe(OTf)2(bpbp)], [Fe(OTf)2(mep)], and [Fe(OTf)2(tpa)] (mcp=N,N′-dimethyl-N,N′-bis(2-pyridylmethyl)-cyclohexane-1,2-diamine; bpbp= N,N′-bis(2-pyridylmethyl)-2,2′-bipyrrolidine; mep= N,N′-dimethyl-N,N′-bis-(2-pyridylmethyl)-ethane-1,2-diamine; tpa= tris-(2-pyridylmethyl)amine) were studied which showed them to be active catalysts. All these complexes have two accessible sites for coordination, cis to each other. The complexes were found to undergo degradation in a few hours of activity because of strong acidic and oxidizing environment. Another complex [(Fe(mcp))2(µ-O)(µ-OH)](OTf)2 was found to show low catalytic activity. [Fe(OTf)2(tmc)] and [Fe(NCCH3)(MePy2CH-tacn)](OTf)2 (tmc= 1,4,8,11-tetramethyl-1,4,8,11-tetraazacyclotetradecane; tacn= 1,4,7-triazacyclononane) were found to be inactive, both possess two trans and only one coordination site. So it could be referred that the presence of cis reactive sites is a key factor for water oxidation catalysis in the aforementioned complexes.

Water splitting by Iridium complexes

Iridium oxide is a stable bulk WOC catalyst with low overpotential.[49] McDaniel et al.synthesized analogues of [Ir(ppy)2(OH2)2]+ (ppy = 2-phenylpyridine) with high turnover numbers, but low catalytic rates. Replacing ppy of the complex with Cp* (C5Me5) resulted in an increase in the catalytic activity but decreased the turnover number.[50] Water nucleophilic attack on Ir=O species was found to be responsible for the O2 formation.[51]

So it could be concluded that a number of complexes of Iron, ruthenium, cobalt, iridium and many others can be used for water oxidation in artificial photosynthesis. The use of these complexes is reliable in terms of high energy production and is more environment-friendly than the commonly used fuel. Several of these WOCs have good TONs and TOFs but unfortunately none have reached to the point of using on large scale as they lose their activity after some time due to multiple reasons or because they are slow to react. Intense research is required in the field to design the ideal WOCs that can be used on the commercial scale.

Chemical production

A variety of materials react with water or acids to release hydrogen. Such methods are non-sustainable. In terms of stoichiometry, these methods resemble the steam reforming process. The great difference between such chemical methods and steam reforming (which is also a "chemical method"), is that the necessary reduced metals do not exist naturally and require considerable energy for their production. For example, in the laboratory strong acids react with zinc metal in Kipp's apparatus.

In the presence of sodium hydroxide, aluminium and its alloys react with water to generate hydrogen gas.[52][53] Unfortunately, due to its energetic inefficiency, aluminium is expensive and usable only for low volume hydrogen generation. Also high amounts of waste heats must be disposed.

Although other metals can perform the same reaction, aluminium is among the most promising materials for future development[54] because it is safer, cheaper and easier to transport than some other hydrogen storage materials like sodium borohydride.

The initial reaction (1) consumes sodium hydroxide and produces both hydrogen gas and an aluminate byproduct. Upon reaching its saturation limit, the aluminate compound decomposes (2) into sodium hydroxide and a crystalline precipitate of aluminum hydroxide. This process is similar to the reactions inside an aluminium battery.

(1) Al + 3 H2O + NaOH → NaAl(OH)4 + 1.5 H2
(2) NaAl(OH)4 → NaOH + Al(OH)3

Overall:

Al + 3 H2O → Al(OH)3 + 1.5 H2

In this process, aluminium functions as a compact hydrogen storage material because 1 kg of aluminum can produce up to 0.111 kg of hydrogen (or 11.1%) from water. When employed in a fuel cell, that hydrogen can also produce electricity, recovering half of the water previously consumed.[55] The U.S. Department of Energy has outlined its goals for a compact hydrogen storage device[56] and researchers are trying many approaches, such as by using a combination of aluminum and NaBH4, to achieve these goals.[57]

Since the oxidation of aluminum is exothermic, these reactions can operate under mild temperatures and pressures, providing a stable and compact source of hydrogen. This chemical reduction process is specially suitable for back-up, remote or marine applications. While the passivation of aluminum would normally slow this reaction considerably,[58] its negative effects can be minimized by changing several experimental parameters such as temperature, alkali concentration, physical form of the aluminum, and solution composition.

Research

Research is being conducted over photocatalysis,[59] the acceleration of a photoreaction in the presence of a catalyst. Its comprehension has been made possible ever since the discovery of water electrolysis by means of the titanium dioxide. Artificial photosynthesis is a research field that attempts to replicate the natural process of photosynthesis, converting sunlight, water and carbon dioxide into carbohydrates and oxygen. Recently, this has been successful in splitting water into hydrogen and oxygen using an artificial compound called Nafion.[60]

High-temperature electrolysis (also HTE or steam electrolysis) is a method currently being investigated for the production of hydrogen from water with oxygen as a by-product. Other research includes thermolysis on defective carbon substrates, thus making hydrogen production possible at temperatures just under 1000 °C.[61]

The iron oxide cycle is a series of thermochemical processes used to produce hydrogen. The iron oxide cycle consists of two chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. All other chemicals are recycled. The iron oxide process requires an efficient source of heat.

The sulfur-iodine cycle (S-I cycle) is a series of thermochemical processes used to produce hydrogen. The S-I cycle consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. All other chemicals are recycled. The S-I process requires an efficient source of heat.

More than 352 thermochemical cycles have been described for water splitting or thermolysis.,[62][63] These cycles promise to produce hydrogen oxygen from water and heat without using electricity.[64] Since all the input energy for such processes is heat, they can be more efficient than high-temperature electrolysis. This is because the efficiency of electricity production is inherently limited. Thermochemical production of hydrogen using chemical energy from coal or natural gas is generally not considered, because the direct chemical path is more efficient.

For all the thermochemical processes, the summary reaction is that of the decomposition of water:

All other reagents are recycled. None of the thermochemical hydrogen production processes have been demonstrated at production levels, although several have been demonstrated in laboratories.

There is also research into the viability of nanoparticles and catalysts to lower the temperature at which water splits.[65][66]

Recently Metal-Organic Framework (MOF)-based materials have been shown to be a highly promising candidate for water splitting with cheap, first row transition metals.;[67][68]

Research is concentrated on the following cycles:[64]

Thermochemical cycle LHV Efficiency Temperature (°C/F)
Cerium(IV) oxide-cerium(III) oxide cycle (CeO2/Ce2O3) ? %2,000 °C (3,630 °F)
Hybrid sulfur cycle (HyS) 43%900 °C (1,650 °F)
Sulfur iodine cycle (S-I cycle) 38%900 °C (1,650 °F)
Cadmium sulfate cycle 46%1,000 °C (1,830 °F)
Barium sulfate cycle 39%1,000 °C (1,830 °F)
Manganese sulfate cycle 35%1,100 °C (2,010 °F)
Zinc zinc-oxide cycle (Zn/ZnO) 44%1,900 °C (3,450 °F)
Hybrid cadmium cycle 42%1,600 °C (2,910 °F)
Cadmium carbonate cycle 43%1,600 °C (2,910 °F)
Iron oxide cycle (Fe3O4/FeO)42%2,200 °C (3,990 °F)
Sodium manganese cycle 49%1,560 °C (2,840 °F)
Nickel manganese ferrite cycle 43%1,800 °C (3,270 °F)
Zinc manganese ferrite cycle 43%1,800 °C (3,270 °F)
Copper-chlorine cycle (Cu-Cl) 41%550 °C (1,022 °F)

Patents

Atmospheric electricity utilization for the chemical reaction in which water is separated into oxygen and hydrogen. (Image via: Vion, US patent 28793. June 1860.)

See also

References

  1. ↑ Patent Database Search Results: ttl/"water splitting" in US Patent Collection
  2. ↑ Decoupled catalytic hydrogen evolution from a molecular metal oxide redox mediator in water splitting B. Rausch, M. D. Symes, G. Chisholm, L. Cronin, Science , 2014 , 345 , 1326-1330.
  3. ↑ Electrolysis of Water
  4. ↑ 2001-High pressure electrolysis - The key technology for efficient H.2
  5. ↑ 2003-PHOEBUS-Pag.9
  6. ↑ Electrode lights the way to artificial photosynthesis
  7. ↑ Solar-Power Breakthrough: Researchers have found a cheap and easy way to store the energy made by solar power
  8. ↑ http://swegene.com/pechouse-a-proposed-cell-solar-hydrogen.html
  9. ↑ del Valle, F.; Ishikawa, A.; Domen, K.; Villoria De La Mano, J.A.; Sánchez-Sánchez, M.C.; González, I.D.; Herreras, S.; Mota, N.; Rivas, M.E. (May 2009). "Influence of Zn concentration in the activity of Cd1-xZnxS solid solutions for water splitting under visible light". Catalysis Today. 143 (1–2): 51–59. doi:10.1016/j.cattod.2008.09.024.
  10. ↑ del Valle, F.; et al. (Jun 2009). "Water Splitting on Semiconductor Catalysts under Visible-Light Irradiation". Chemsuschem. CHEMSUSCHEM. 2 (6): 471–485. doi:10.1002/cssc.200900018. PMID 19536754.
  11. ↑ del Valle, F.; et al. (2009). "Photocatalytic water splitting under visible Light: concept and materials requirements". Advances in Chemical Engineering. Advances in Chemical Engineering. 36: 111–143. doi:10.1016/S0065-2377(09)00404-9. ISBN 9780123747631.
  12. ↑ Li-Hung Lin; Pei-Ling Wang; Douglas Rumble; Johanna Lippmann-Pipke; Erik Boice; Lisa M. Pratt; Barbara Sherwood Lollar; Eoin L. Brodie; Terry C. Hazen; Gary L. Andersen; Todd Z. DeSantis; Duane P. Moser; Dave Kershaw & T. C. Onstott (2006). "Long-Term Sustainability of a High-Energy, Low-Diversity Crustal Biome". Science. 314 (5798): 479–82. Bibcode:2006Sci...314..479L. doi:10.1126/science.1127376. PMID 17053150.
  13. ↑ DOE 2008 Report 25 %
  14. ↑ Biohydrogen study in Thailand
  15. ↑ "Thermochemical hydrogen production: past and present". Int J Hydrogen Energy. 26: 185pp. 2001. doi:10.1016/S0360-3199(00)00062-8.
  16. ↑ http://www.dis.anl.gov/ceeesa/documents/NuclearHydrogen_ANL0530Final.pdf
  17. ↑ Naterer, G. F.; et al. (2009). "Recent Canadian Advances in Nuclear-Based Hydrogen Production and the Thermochemical Cu-Cl Cycle". International Journal of Hydrogen Energy. 34 (7): 2901–2917. doi:10.1016/j.ijhydene.2009.01.090.
  18. ↑ http://www.dlr.de/en/desktopdefault.aspx/tabid-1/86_read-14380/
  19. ↑ http://www.h2powersystems.com
  20. ↑ Eisenberg, R.; Gray, H. B., Preface on Making Oxygen. Inorganic Chemistry 2008, 47, 1697-1699.
  21. ↑ Sun, L.; Hammarstrom, L.; Akermark, B.; Styring, S., Towards artificial photosynthesis: ruthenium-manganese chemistry for energy production. Chemical Society Reviews 2001, 30, 36-49.
  22. ↑ Gust, D.; Moore, T. A.; Moore, A. L., Solar Fuels via Artificial Photosynthesis. Accounts of Chemical Research 2009, 42, 1890-1898.
  23. ↑ Meyer, T. J., Chemical approaches to artificial photosynthesis. Accounts of Chemical Research 1989, 22, 163-170.
  24. ↑ Balzani, V.; Credi, A.; Venturi, M., Photochemical Conversion of Solar Energy. ChemSusChem 2008, 1, 26-58.
  25. ↑ Sala, X.; Romero, I.; Rodríguez, M.; Escriche, L.; Llobet, A., Molecular Catalysts that Oxidize Water to Dioxygen. Angewandte Chemie International Edition 2009, 48, 2842-2852.
  26. ↑ Yano, J.; Kern, J.; Sauer, K.; Latimer, M. J.; Pushkar, Y.; Biesiadka, J.; Loll, B.; Saenger, W.; Messinger, J.; Zouni, A.; Yachandra, V. K., Where Water Is Oxidized to Dioxygen: Structure of the Photosynthetic Mn(4)Ca Cluster. Science (New York, N.Y.) 2006, 314, 821-825.
  27. ↑ Barber, J., Crystal Structure of the Oxygen-Evolving Complex of Photosystem II. Inorganic Chemistry 2008, 47, 1700-1710.
  28. ↑ Haumann, M.; Liebisch, P.; Müller, C.; Barra, M.; Grabolle, M.; Dau, H., Biochemistry: Photosynthetic O2 formation tracked by time-resolved x-ray experiments. Science 2005, 310, 1019-1021.
  29. ↑ Gratzel, M., Photoelectrochemical cells. Nature 2001, 414, 338-344.
  30. ↑ Wada, T.; Tsuge, K.; Tanaka, K., Electrochemical Oxidation of Water to Dioxygen Catalyzed by the Oxidized Form of the Bis(ruthenium – hydroxo) Complex in H2O. Angewandte Chemie International Edition 2000, 39, 1479-1482.
  31. 1 2 Sens, C.; Romero, I.; Rodríguez, M.; Llobet, A.; Parella, T.; Benet-Buchholz, J., A New Ru Complex Capable of Catalytically Oxidizing Water to Molecular Dioxygen. Journal of the American Chemical Society 2004, 126, 7798-7799.
  32. ↑ Liu, F.; Concepcion, J. J.; Jurss, J. W.; Cardolaccia, T.; Templeton, J. L.; Meyer, T. J., Mechanisms of Water Oxidation from the Blue Dimer to Photosystem II. Inorganic Chemistry 2008, 47, 1727-1752.
  33. ↑ Zong, R.; Thummel, R. P., A New Family of Ru Complexes for Water Oxidation. Journal of the American Chemical Society 2005, 127, 12802-12803.
  34. ↑ Duan, L.; Fischer, A.; Xu, Y.; Sun, L., Isolated Seven-Coordinate Ru(IV) Dimer Complex with [HOHOH]− Bridging Ligand as an Intermediate for Catalytic Water Oxidation. Journal of the American Chemical Society 2009, 131, 10397-10399.
  35. ↑ Zhang, G.; Zong, R.; Tseng, H.-W.; Thummel, R. P., Ru(II) Complexes of Tetradentate Ligands Related to 2,9-Di(pyrid-2‘-yl)-1,10-phenanthroline. Inorganic Chemistry 2008, 47, 990-998.
  36. ↑ Duan, L.; Bozoglian, F.; Mandal, S.; Stewart, B.; Privalov, T.; Llobet, A.; Sun, L., A molecular ruthenium catalyst with water-oxidation activity comparable to that of photosystem II. Nat Chem 2012, 4, 418-423.
  37. ↑ Singh, R.; Pandey, J.; Singh, N.; Lal, B.; Chartier, P.; Koenig, J.-F., Sol-gel derived spinel M x Co 3− x O 4 (M= Ni, Cu; 0≤ x≤ 1) films and oxygen evolution. Electrochimica Acta 2000, 45, 1911-1919.
  38. ↑ Dincă, M.; Surendranath, Y.; Nocera, D. G., Nickel-borate oxygen-evolving catalyst that functions under benign conditions. Proceedings of the National Academy of Sciences 2010, 107, 10337-10341.
  39. ↑ Risch, M.; Klingan, K.; Heidkamp, J.; Ehrenberg, D.; Chernev, P.; Zaharieva, I.; Dau, H., Nickel-oxido structure of a water-oxidizing catalyst film. Chemical communications 2011, 47, 11912-11914.
  40. ↑ Zaharieva, I.; Najafpour, M. M.; Wiechen, M.; Haumann, M.; Kurz, P.; Dau, H., Synthetic manganese–calcium oxides mimic the water-oxidizing complex of photosynthesis functionally and structurally. Energy & Environmental Science 2011, 4, 2400-2408.
  41. ↑ Kanan, M. W.; Yano, J.; Surendranath, Y.; Dincă, M.; Yachandra, V. K.; Nocera, D. G., Structure and Valency of a Cobalt−Phosphate Water Oxidation Catalyst Determined by in Situ X-ray Spectroscopy. Journal of the American Chemical Society 2010, 132, 13692-13701.
  42. ↑ Brunschwig, B. S.; Chou, M. H.; Creutz, C.; Ghosh, P.; Sutin, N., Mechanisms of water oxidation to oxygen: cobalt(IV) as an intermediate in the aquocobalt(II)-catalyzed reaction. Journal of the American Chemical Society 1983, 105, 4832-4833.
  43. ↑ Kanan, M. W.; Nocera, D. G., In Situ Formation of an Oxygen-Evolving Catalyst in Neutral Water Containing Phosphate and Co2+. Science 2008, 321, 1072-1075.
  44. ↑ Harriman, A.; Pickering, I. J.; Thomas, J. M.; Christensen, P. A., Metal oxides as heterogeneous catalysts for oxygen evolution under photochemical conditions. Journal of the Chemical Society, Faraday Transactions 1: Physical Chemistry in Condensed Phases 1988, 84, 2795-2806.
  45. ↑ Yin, Q.; Tan, J. M.; Besson, C.; Geletii, Y. V.; Musaev, D. G.; Kuznetsov, A. E.; Luo, Z.; Hardcastle, K. I.; Hill, C. L., A fast soluble carbon-free molecular water oxidation catalyst based on abundant metals. Science 2010, 328, 342-345.
  46. ↑ Zidki, T.; Zhang, L.; Shafirovich, V.; Lymar, S. V., Water Oxidation Catalyzed by Cobalt(II) Adsorbed on Silica Nanoparticles. Journal of the American Chemical Society 2012, 134, 14275-14278.
  47. ↑ Wasylenko, D. J.; Ganesamoorthy, C.; Borau-Garcia, J.; Berlinguette, C. P., Electrochemical evidence for catalyticwater oxidation mediated by a high-valent cobalt complex. Chemical communications 2011, 47, 4249-4251.
  48. ↑ Fillol, J. L.; Codolà, Z.; Garcia-Bosch, I.; Gómez, L.; Pla, J. J.; Costas, M., Efficient water oxidation catalysts based on readily available iron coordination complexes. Nat Chem 2011, 3, 807-813.
  49. ↑ Youngblood, W. J.; Lee, S.-H. A.; Maeda, K.; Mallouk, T. E., Visible Light Water Splitting Using Dye-Sensitized Oxide Semiconductors. Accounts of Chemical Research 2009, 42, 1966-1973.
  50. ↑ Hull, J. F.; Balcells, D.; Blakemore, J. D.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H., Highly Active and Robust Cp* Iridium Complexes for Catalytic Water Oxidation. Journal of the American Chemical Society 2009, 131, 8730-8731.
  51. ↑ Blakemore, J. D.; Schley, N. D.; Balcells, D.; Hull, J. F.; Olack, G. W.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H., Half-Sandwich Iridium Complexes for Homogeneous Water-Oxidation Catalysis. Journal of the American Chemical Society 2010, 132, 16017-16029.
  52. ↑ Belitskus, David (August 1970). "Reaction of Aluminum with Sodium Hydroxide Solution as a Source of Hydrogen". Journal of the Electrochemical Society. Pennington, New Jersey: ECS. 117 (8): 1097–1099. doi:10.1149/1.2407730. ISSN 0013-4651.
  53. ↑ Soler, Lluís; Macanás, Jorge; Muñoz, Maria; Casado, Juan (2007). "Aluminum and aluminum alloys as sources of hydrogen for fuel cell applications". Journal of Power Sources. Elsevier. 169 (1): 144–149. doi:10.1016/j.jpowsour.2007.01.080.
  54. ↑ Wang, H.Z.; Leung, D.Y.C.; Leung, M.K.H.; Ni, M. (2008). "A review on hydrogen production using aluminum and aluminum alloys". Renewable and Sustainable Energy Reviews. Elsevier. 13 (4): 845–853. doi:10.1016/j.rser.2008.02.009.
  55. ↑ Amendola, Steven C.; Binder, Michael; Kelly, Michael T.; Petillo, Phillip J.; Sharp-Goldman, Stefanie L. (2000). "A Novel Catalytic Process for Generating Hydrogen Gas from Aqueous Borohydride Solutions". In Grégoire Padró, Catherine E.; Lau, Francis. Advances in Hydrogen Energy. New York: Kluwer Academic Publishers. pp. 69–86. doi:10.1007/0-306-46922-7_6. ISBN 978-0-306-46922-0.
  56. ↑ http://www.sc.doe.gov/bes/hydrogen.pdf
  57. ↑ Soler, Lluís; Macanás, Jorge; Muñoz, Maria; Casado, Juan (2007). "Synergistic hydrogen generation from aluminum, aluminum alloys and sodium borohydride in aqueous solutions". International Journal of Hydrogen Energy. Elsevier. 32 (18): 4702–4710. doi:10.1016/j.ijhydene.2007.06.019. ISSN 0360-3199.
  58. ↑ Stockburger, D.; Stannard, J.H.; Rao, B.M.L.; Kobasz, W.; Tuck, C.D. (1992). Corrigan, Dennis A.; Srinivasan, Supramaniam, eds. Hydrogen storage materials, batteries, and electrochemistry. Pennington, New Jersey: ECS. pp. 431–444. ISBN 978-1-56677-006-4. OCLC 25662899.
  59. ↑ Strategies for the Development of Visible-light-driven Photocatalysts for Water Splitting Akihiko Kudo, Hideki Kato1 and Issei Tsuji Chemistry Letters Vol. 33 (2004), No. 12 p.1534
  60. ↑
  61. ↑ Kostov, M. K.; Santiso, E. E.; George, A. M.; Gubbins, K. E. & Nardelli, M. Buongiorno (2005). "Dissociation of Water on Defective Carbon Substrates" (PDF). Physical Review Letters. Retrieved 2007-11-05.
  62. ↑ 353 Thermochemical cycles
  63. ↑ UNLV Thermochemical cycle automated scoring database (public)
  64. 1 2 Development of solar-powered thermochemical production of hydrogen from water
  65. ↑ Naoptek
  66. ↑ http://www.treehugger.com/files/2008/07/hydrogen-production-breakthrough-from-mit-a-giant-leap.php
  67. ↑ Das; et al. (2013). Angew. Chem. Int. Ed. 52: 7224–7227. doi:10.1002/anie.201301327 http://onlinelibrary.wiley.com/doi/10.1002/anie.201301327/abstract. Missing or empty |title= (help)
  68. ↑ Hansen; Das. "Text". Energy & Environ Sci.

External links

This article is issued from Wikipedia - version of the 10/9/2016. The text is available under the Creative Commons Attribution/Share Alike but additional terms may apply for the media files.