Hydrogen production from renewable sources: biomass and photocatalytic opportunities

Abstract

The demand for hydrogen over the coming decade is expected to grow for both traditional uses (ammonia, methanol, refinery) and running fuel cells. At least in the near future, this thirst for hydrogen will be quenched primarily through the reforming of fossil fuels. However, reforming fossil fuels emits huge amounts of carbon dioxide. One approach to reduce carbon dioxide emissions, which is considered first in this review, is to apply reforming methods to alternative renewable materials. Such materials might be derived from plant crops, agricultural residues, woody biomass, etc. Clean biomass is a proven source of renewable energy that is already used for generating heat, electricity, and liquid transportation fuels. Clean biomass and biomass-derived precursors such as ethanol and sugars are appropriate precursors for producing hydrogen through different conversion strategies. Virtually no net greenhouse gas emissions result because a natural cycle is maintained, in which carbon is extracted from the atmosphere during plant growth and released during hydrogen production. The second option explored here is hydrogen production from water splitting by means of the photons in the visible spectrum. The sun provides silent and precious energy that is distributed fairly evenly all over the earth. However, its tremendous potential as a clean, safe and economical energy source cannot be exploited unless it is accumulated or converted into more useful forms of energy. Finally, this review discusses the use of semiconductors, more specifically CdS and CdS-based semiconductors, which are able to absorb photons in the visible region of the spectrum. The energy stored within a semiconductor as electronic energy (electrons and holes) can be used to split water molecules by simultaneous reactions into H₂ and O₂. This conversion of solar energy into a clean fuel (H₂) is perhaps the greatest challenge for scientists in the 21st century.

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Rufino M. Navarro is a Tenured Scientist at the Institute of Catalysis and Petrochemistry of the National Council of Scientific Research (CSIC). His research focuses on heterogeneous catalysis applied to clean energy production: reforming of hydrocarbon, hydrogen technologies, renewable energies, thermochemical cycles and water splitting technologies.

Maria Cruz Sánchez is a PhD candidate in the Institute of Catalysis and Petrochemistry of the National Council for Scientific Research (CSIC). Her research activities are concentrated in the field of hydrogen production from bioalcohols.

Consuelo Alvarez-Galvan received her doctoral degree from the University of La Laguna and then moved to the Institute of Catalysis and Petrochemistry of the National Council for Scientific Research (CSIC) where she is as a Ramon y Cajal postdoctoral scientist. Her research interests include hydrocarbon reforming, thermochemical cycles, and catalytic combustion.

Fabian del Valle is a PhD candidate in the Institute of Catalysis and Petrochemistry of the National Council for Scientific Research (CSIC). His research activities are in the field of water splitting on semiconductor surfaces.

Jose L.G. Fierro is Research Professor at Institute of Catalysis and Petrochemistry of the National Council of Scientific Research (CSIC). His research interests include hydrogen technologies, renewable energy, natural gas conversion, chemical technology, solid state chemistry, environmental chemistry, and commodity and bulk chemicals production.

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Broader context

Hydrogen is a versatile energy carrier that is usually produced from steam reforming of carbon-containing sources with the coproduction of huge amounts of CO₂. To take full advantage of the environmental benefits of hydrogen, low carbon emitting, low polluting, low cost hydrogen production systems are needed. This review details the fundamentals of and describes some major achievements in the production of hydrogen from two renewable resources: biomass and water splitting with visible light. Particular emphasis is placed on biomass gasification under near- and supercritical water conditions, and also in semiconductors which split water under ambient temperature and pressure upon irradiation with light of wavelength in the visible spectrum.

1. Introduction

Hydrogen (H₂) is the most common element on earth, but it does not occur to a significant extent in elemental form. It is mostly present in water, biomass and hydrocarbons. Hydrogen gas is currently produced from a variety of primary sources, such as natural gas, naphtha, heavy oil, methanol, biomass, wastes, coal, solar, wind and nuclear.^1–4 It is a clean energy fuel because the chemical energy stored in the H–H bond is released when it combines with oxygen, yielding only water as a reaction product. Accordingly, a future energy infrastructure based on hydrogen has been perceived as an ideal long-term solution to energy-related environmental problems.^5–7 There is no doubt that hydrogen has the potential to provide a clean and affordable energy supply that can minimize our dependence on oil and therefore enhance the global economy and reduce environmental pollution.

It is generally understood that the renewable energy-based processes of hydrogen production (solar photochemical and photobiological water decomposition, electrolysis of water coupled with photovoltaic cells or wind turbines, etc.) are unlikely to involve significant reductions in hydrogen costs over the next ten-to-twenty years. Industry generates some 48 million metric tons of hydrogen globally each year from fossil fuel. Almost half of this hydrogen goes into making ammonia, while refineries use the second largest volume of hydrogen for chemical processes such as removing sulfur from gasoline and converting heavy hydrocarbons into gasoline and diesel fuel. Food producers add a small percentage of hydrogen to some edible oils through a catalytic hydrogenation process.

The demand for hydrogen is expected to grow over the next ten years, for both traditional uses, such as making ammonia, and running fuel cells. At least in the near future, this thirst for hydrogen will be quenched primarily through the use of fossil fuels. To make hydrogen, industry uses Steam Methane Reforming (SMR), which is the most widely used and most economical process for producing hydrogen.^1–4 Although SMR is a complex process involving many different catalytic steps, as long as natural gas (or CH₄) and hydrocarbon fuels remain at a low or even moderate price, SMR will continue to be the choice technology for the mass production of H₂. Following several decades of improvements in catalyst technology, substantial improvements have been introduced over the years. The SMR process also produces carbon monoxide and carbon dioxide, the primary greenhouse gas. Although this approach still generates pollution, these gases are released in a potentially more manageable way, as opposed to the case of millions of automobile engines.

Nonetheless, shedding the habit of fossil fuel entirely is the only way a wholesale shift to hydrogen will work in the long-term. One approach to this goal is to apply steam-reforming methods to alternative renewable materials. Such materials might be derived from crops. Not only do these biomass-conversion schemes turn waste into a valuable product, but, researchers say, there is another benefit: any carbon dioxide released in the processes could be soaked up by planting new crops to provide the required biomass. A biomass strategy of hydrogen generation could be a useful intermediate step between the current fossil fuel method and the dream of efficient water splitting. Still, any realistic contender for hydrogen generation must first replace the reforming of fossil fuel as the cheapest and most efficient process.

Although hydrogen production and storage/distribution infrastructures are commercially available in chemical and refining industries around the world, existing conversion and storage technologies are too expensive for widespread use in energy schemes. The development of hydrogen production as a realistic, viable energy option will require an unprecedented level of sustained and coordinated activities and, most importantly, the development of new concepts for the distributed production of hydrogen. Not only fossil fuel burning, which contributes to the greenhouse gas pool, but also the eventual depletion of the world's fossil-fuel reserves are threatening sustainable development.^8,9 Abundant, clean, and carbon-neutral hydrogen is widely believed to be the ultimate mobile energy carrier replacing gasoline, kerosene and diesel in internal combustion engines. The present status of hydrogen production from less costly and abundant biomass for producing low-cost hydrogen without net carbon emissions is reviewed.^10–14 In addition, this review includes advances in the fully renewable conversion of solar energy into hydrogen via the water splitting process assisted by photo-semiconductor catalysts.

2. Hydrogen production from carbon-neutral precursors

Hydrogen is currently produced almost entirely from natural gas, liquid hydrocarbons and coal. All these C-containing sources release massive amounts of CO₂ into the atmosphere during the production of hydrogen. Thus, renewable biomass, a product of photosynthesis, is an attractive alternative to fossil feedstocks, as it can be considered a CO₂ neutral precursor. Photosynthesis is the biological process that converts light energy into chemical energy and stores it in carbohydrates as “nCO₂ + nH₂O → [CH₂O]_n + nO₂”, and fixes atmospheric carbon into biomass. At the dawn of the 21st century, a combination of economic, technological, resource and political developments is driving the emergence of a new carbohydrate economy.^15,16

Biomass can be converted into hydrogen and other useful products through several thermochemical processes, such as liquefaction, pyrolysis and gasification. Among these, pyrolysis and gasification have received considerable attention in recent years. In addition, there are other biochemical processes for the conversion of biomass into hydrogen. A brief account of biological H₂ production and the future challenges in this field are also featured in this review.

2.1 Thermal pyrolysis

Pyrolysis is the thermal decomposition of solid biomass resources at a temperature of 650–800 K at 1–5 bar in the absence of air to yield liquid oils, solid charcoal and gaseous compounds. Pyrolysis can be classified into slow pyrolysis and fast pyrolysis (Fig. 1). Although most pyrolysis processes are designed for biofuel production, hydrogen can be produced directly through fast or flash pyrolysis if both high temperature and sufficient volatile phase residence time are provided for as follows (eqn 1):

C_xH_yO_z + heat → H₂ + CO + CH₄ + HCs (1)

where HCs stands for other hydrocarbons. The methane and hydrocarbons produced are reformed by steam to yield additional hydrogen (eqn 2). A water–gas shift reaction is used to increase hydrogen production (eqn 3):

CH₄ + HCs + H₂O → CO + H₂ (2)

CO + H₂O → CO₂ + H₂ (3)

Fig. 1 Sketch diagram of a fast pyrolysis plant. (R), reactor; (C), cyclone; (B), bio-oil. Adapted from ref. 21.

Lower process temperatures and longer vapour residence times favour the production of charcoal. High temperatures and longer residence times increase biomass conversion to gas, and moderate temperatures and short vapour residence times are optimum for producing liquid. Thus, fast pyrolysis for liquid production is currently of particular interest because liquids can be stored and transported more easily and at a lower cost than solid biomass.¹⁷ Besides the gaseous products, the bio-oil products can also be processed for hydrogen production.¹⁸ In general maximum hydrogen yield at about 90% is obtained with steam reforming over Ni-based catalysts and additional water–gas shift reaction. Pyrolysis reaction rates are accelerated by some chlorides and carbonates.¹⁹ Since it is difficult to gasify tar, extensive studies on the catalytic effect of inexpensive dolomite and CaO on the decomposition of HCs in tar have been carried out.²⁰ The beneficial effects of other catalysts, such as Ni,²¹ Y-type zeolite,²² carbonates,²³ and metal oxides²⁴ have also been investigated. Virtually any form of biomass can be considered for fast pyrolysis. While most work has been carried out on wood because of its consistency between tests, nearly one hundred different biomass types have been employed, ranging from agricultural wastes such as straw, olive pits and nut shells to energy crops, forestry wastes and solid wastes.

2.2 Steam/oxygen gasification

Biomass is also gasified at temperatures above 1000 K in the presence of oxygen and/or water. Under these conditions, biomass undergoes partial oxidation and/or steam reforming reactions yielding gas and char product. The char is subsequently reduced to form H₂, CO, CO₂ and CH₄. This conversion process can be expressed as:

C_xH_yO_z + H₂O + O₂ → H₂ + CO_x + CH₄ + HCs + char (4)

As the reaction products of biomass gasification are mainly gases, this process is more favourable for hydrogen production than pyrolysis. In order to optimize the process for hydrogen production, a number of efforts have been made to test hydrogen production from biomass gasification with various biomass types and at various operating conditions.

Most biomass gasification processes employ air as gasifying agent, which results in a low calorific value gas stream (3–5 MJ m⁻³). This gas can be used after cleaning in gas-fired engines or gas turbines. For gas turbines connected to a steam turbine, medium calorific value gas (10–15 MJ m⁻³) is more favourable than low calorific gas. Steam injection into the gas turbine combustion chamber requires at least medium calorific value gas. The production of methanol or hydrogen via biomass gasification or the use of producer gas in low-temperature fuel cells also require either gasifiers operating with highly-enriched oxygen or gasifiers using steam as a gasification medium to generate the necessary medium calorific value raw gas with high hydrogen content.

The gasification of wood and wood-type residues and waste in fixed bed or fluidised bed gasifiers with subsequent burning of the gas for heat production is the state-of-the-art. Significantly greater technical problems are posed by the gasification of straw and other solid agricultural feedstocks, which mostly have higher concentrations of chlorine, alkali, nitrogen and sulfur. The gasification of herbaceous biomass is still at an early stage of research and development. Intensified development efforts involving gasification technologies for herbaceous biomass feedstocks are desirable, as the potential supply of this group of fuels is comparatively large. Thorough gas cleaning and the adaptation of the gas from biomass gasification to the specific requirements of the gas utilisation systems are the prerequisites for gas use in gas-fired engines, gas turbines and fuel cells.

One of the major issues in biomass gasification is the tar formation that occurs during the process. The unwanted tar polymerizes to a more complex structure, which is not favourable for hydrogen production through steam reforming. Currently, three methods are available for minimising tar formation: (i) proper gasifier design; (ii) incorporation of catalysts; and (iii) control of operating variables. Regarding method (iii), the operating parameters, such as gasifying agent, temperature and residence time, are key factors in the formation and decomposition of tar. Tar can be thermally cracked at temperatures above 1273 K.²⁵ In the type (ii) method, the use of additives (such as dolomite, olivine and even char), also facilitates tar reduction.²⁶ Dolomite is particularly suited because 100% tar elimination can be achieved with this additive.²⁷ Catalysts also reduce tar content, but are particularly effective for improving gas product quality and conversion. Dolomite-loaded nickel catalysts and alkaline metal oxides are widely used as gasification catalysts.

2.2.1 Gasifiers. A number of different biomass gasifiers can be found in patent bibliographies. They can be grouped into three main types, as sketched in Fig. 2: (a) fluidized-bed gasifier; (b) downdraft gasifier; and (c) updraft gasifier.²⁸

Fig. 2 Sketch pictures of biomass reactors: (a), fluidized-bed gasifier; (b), downdraft gasifier; and (c), updraft gasifier (ref. 28).

In the fluidized-bed reactor (Fig. 2(a)) the biomass, which is previously reduced to a fine powder, air, steam, or oxygen enters through the bottom of the gasifier. A high linear velocity of the gas stream forces the fine particles of biomass upward through a bed of silica beads. Pyrolysis and char gasification take place in this process. This type of gasifier is suitable for large-scale applications and has a medium tar yield of around 10 g Nm⁻³. In the downdraft gasifier (Fig. 2(b)), the air or oxygen and biomass particles enter through the top of the reactor as a fine powder and flow downward, and the gas exits through the bottom of the reactor. The product gas contains the lowest concentration of particulates and tars of nearly 1 g Nm⁻³, which is a much lower level than in fluidized-bed reactor, because most of the tars are combusted. The flame temperature in this reactor is 1200–1600 K. This reactor configuration is ideal when clean gas is needed. The main disadvantage of this gasifier reactor is its low overall thermal efficiency, as well as the difficulty in handling ash content. In the updraft gasifier (Fig. 2(c)), biomass enters through the top and air/oxygen/steam flow upward from the bottom and the gas exits through the top. This reactor primarily forms tars at a very high level (in the order of 100 g Nm⁻³). The principal advantages of an updraft gasifier include: it is a mature technology for heat production, it can be used for small-scale applications, and it can handle feeds with high moisture content. On the other hand, the tar yield of this gasifier is very high and it has slagging potential.

2.2.2 Absorption enhanced reforming. The in situ capture of CO₂ during gasification is an especially attractive process since it allows for very high H₂-content with very low (near zero) CO₂ and tar contents in the gasification gas.^29,30 Attempts to use CaO in a CO₂ acceptor process were first conducted by Curran et al.³¹ and McCoy et al.³² In these studies, only one half of CO and CO₂ was immobilized in CaO. A new method that combines the gas production and separation reactions (Hydrogen Production by Reactions Integrated gasification, HyPr-Ring) in a single reactor was proposed by Lin et al.³³ In this process, the energy required for the endothermic reforming reactions is supplied by the heat of CO₂ absorption. Wang and Takarada³⁴ reported complete fixation of CO₂ with Ca(OH)₂ for a Ca/molar ratio of 0.6 (stoichiometry dictates the ratio to be 1) along with enhanced decomposition of tar and char. In addition, the overall conversion rate of CO to CO₂ can be enhanced by the inclusion of an oxygen donor in the reaction zone.

Absorption enhanced steam reforming (AER) is another technology that has been explored recently for the continuous gasification of biomass and the production of a H₂-rich gas stream.^35,36 In the AER process, the CO₂ produced during steam gasification is separated from the reactor by an adsorbent, so that the resulting product gas contains a high H₂ concentration and low concentration of carbon oxides. CO₂ absorption not only shifts the equilibrium towards the desired product, but also delivers heat for the endothermic reactions. A gas with increased CO₂ concentration is produced along the regeneration step, which simplifies CO₂ separation. The concept of the AER process is outlined in Fig. 3. The two fluidised bed reactors are coupled in such a manner that the sorbent bed material circulates between the AER gasifier (CO₂ absorption) and the combustor (CO₂ desorption). A nearly nitrogen-free product gas with a caloric value of 12–14 MJ Nm⁻³ (dry) is produced in the AER gasifier.

Fig. 3 Coupling of two fluidized-bed reactors for the continuous production of an H₂-rich gas flow from biomass. The adsorbent bed material circulates between the AER reactor (absorption) and the regenerator (desorption).

AER is being used for the unpressurised steam gasification of biomass (eqn 5). Through simultaneous CO₂ absorption (in the example with CaO) (eqn 7), the equilibrium of the homogeneous WGS reaction (eqn 6) is shifted towards H₂ and CO₂ and all the parallel gasification/reforming reactions are also influenced in favour of the desired products. Accordingly, a H₂-rich product gas results with reduced CO and CO₂ concentration. Eqn 4 represents the idealized sum reaction for AER gasification. For the sake of simplicity, the formation of methane coke and tars is ignored.

CH_xO_y + (1−y)H₂O → CO + (0.5x +1−y)H₂   (ΔH > 0) (5)

CO + H₂O → CO₂ + H₂   (ΔH < 0) (6)

CaO + CO₂ → CaCO₃   (ΔH > 0) (7)

CH_xO_y + (1−y)H₂O + CaO → CO + (0.5x + 2−y)H₂ (8)

Regarding experiments conducted in dual fluidised-beds with continuous absorption and regeneration, a H₂ content of almost 70% could be achieved in most cases.³⁷ A typical gas composition of the 100 kW gasifier operated under AER-mode is shown in Fig. 4.³⁸ The experimental run at the 8 MW biomass CHP Güssing (Austria) plant showed that the AER-gas concept can also be realised in an existing dual fluidised bed steam gasification plant without major modifications. Product gas with increased H₂ content (>50%), low CO₂ content (<12%) and low tar content by about 1g Nm⁻³ (as in conventional gasification) has been obtained. It is clear that AER technology reduces the need for downstream processing. In addition, the removal of CO₂ by calcium oxide makes the reaction occur at a lower temperature (670–770 K vs. 1070–1270 K), reducing heat losses and material costs.

Fig. 4 Typical gas composition of a 100 kW gasifier operated in the Absorption Enhanced Reforming (AER) mode.

Sorbent enhanced reforming technology is still at the experimental stage, and shows promise for low cost H₂ production. Critical issues in this methodology are sorbent lifetime and system design.

2.2.3 Costs. Hydrogen prices vary considerably depending on the volume and form of delivery. For gaseous hydrogen produced from natural gas at a large-scale, central production facility at a pressure of around 30 bar, the plant gate price is about $5–$8/GJ. Generally, as the feedstock goes from natural gas to light hydrocarbons to heavy hydrocarbons and then to solid feedstocks, there is an increase in processing difficulty and capital costs. Thus, for hydrogen produced from biomass gasification the plant gate price is $12–13/MJ.³⁹ The relative costs (tax-free) and greenhouse gas emission levels (fuel supply and use) per MJ fuel for various fuel chains are plotted in Fig. 5. Fuels included are compressed and liquid hydrogen, produced via different paths for the transport sector.⁴⁰ The H₂ supply cost produced via gasification of wood is around three times higher that of gasoline/diesel, but on the contrary greenhouse gas emissions from supply and use is 7–8 times higher for gasoline/diesel fuel.

Fig. 5 Specific greenhouse gas emissions of the supply chain and use of fuel with respect to fuel costs (without taxes) from ref. 68.

Hydrogen costs delivered to the end user will generally be higher than the costs of current fossil fuel options. Especially hydrogen based on renewable sources will increase fuel costs compared with conventional fuels but at a very low specific CO₂ level. Higher efficiencies of fuel cell systems, costs for CO₂ capture and storage – necessary for fossil-based hydrogen – as well as bonus points for less specific carbon dioxide emissions will modify Fig. 5.

2.3 Biomass gasification in supercritical water

Supercritical water (SCW) is obtained at pressures above 221 bar and temperatures above 647 K. In the absence of oxygen, biomass is converted under supercritical water conditions into fuel gases (eqn 9), which are easily separated from the water phase by cooling to ambient temperature.^41,42

C_xH_yO_z + (2x−z)H₂O → xCO₂ + (2x−z + y/2)H₂ (9)

Thermodynamic calculations indicate that at temperatures above 873 K only a gas rich in H₂, CH₄, CO and CO₂, with no solid carbon, is formed. Unfortunately biomass does not react directly with steam at atmospheric pressure to yield gas products. Modell⁴¹ reported that the immersion of maple wood sawdust in supercritical water quickly decomposed to tars and some gas without the formation of char. Cellulose is the most stable component of biomass but suffers rapid decomposition at temperatures somewhat below the critical temperature of water. At temperatures above 463 K, a fraction of lignin and hemicellulose reacts via solvolysis after only a few minutes of exposure to hot water. The initial products generated in solvolysis undergo several reactions, such as dehydration, isomerization, fragmentation and condensation, finally forming gas and tars. The cellulose degradation kinetics began to have a much higher reaction rate as the reaction temperature approached and entered the supercritical water regime above 623 K and 25 MPa.⁴³

2.3.1 Reactors. Biomass gasification is a process under development and hence no commercial plants exist. The largest pilot installation was built in 2003 at the Forschungszentrum Karlsruhe (Germany). This plant has a capacity for 100 L h⁻¹, and was constructed to test the supercritical gasification of wet residues. Several smaller reactors are used at laboratory scale at Hiroshima University, University of Hawaii, Osaka Gas, Pacific Northwest National Laboratory and the University of Twente. Steel batch autoclaves are used in most cases to analyse product distribution and yield for different feedstocks, operating conditions and catalyst formulations. Steel autoclaves have the disadvantage of heating slowly and thus some time is required to reach reaction temperature. The possible catalytic effect of the walls has been suggested.⁴⁴

Other reactor types include capillaries and tubular steel reactors. Quartz capillaries have also been used as batch microreactors. This reactor configuration allows for inexpensive and high-speed testing, with the further possibility of visual observation. A drawback is that the pressure inside the capillary cannot be directly measured; it is calculated from the temperature and the sample plus reactor volume. For continuous operation, tubular steel reactors are often used. Other types of reactor, such as the stirred tank reactor can be used in principle, but to date this configuration has not yet been applied. To maintain the same capacity, the volume of a stirred tank reactor should be larger than the tubular reactor. The biomass concentration is lower in a stirred tank reactor due to the fast dilution to reactor outlet concentration.

2.3.2 Low temperature SCW gasification.
2.3.2.1 Chemistry. The mechanism of hydrothermal degradation of biomass is complex. Reactions involving cellulose and glucose were studied by Minowa's group in either the absence or presence of catalysts.⁴⁵ They used an autoclave as a reactor and ran the reaction in hot water within the 473 to 623 K temperature range. In all cases, the reaction products were gases, aqueous phase, oily material and solid residue. In the absence of catalyst, cellulose was slightly decomposed over 473 K to yield sugars, which are water-soluble products. As no gases, oil or char formation was produced, it could be inferred that hydrolysis is the primary step for the gasification reaction.⁴⁵ At 523 K, cellulose was decomposed to form gases, oil, char, sugars and other non-sugar compounds. Over 573 K, cellulose conversion was complete; sugars and oil decomposed while char production increased. Finally, char is mainly obtained with a yield ca. 60% on a carbon basis, with 15% of non-sugar water-soluble products and 10% of gas, mainly CO₂, with very small amounts of CO. From these results, a simple reaction scheme was proposed (Fig. 6). Additional experiments using glucose as starting feedstock concluded that hydrolysis is the first step in cellulose conversion.⁴⁶ For this feedstock, the product distribution of gas, oil and char at different reaction temperatures was essentially the same as that for cellulose. The degradation scheme for glucose and cellulose is therefore the same.

Fig. 6 Strategies for production of fuels from lignocellulosic biomass from ref. 4.

The onset temperature of cellulose degradation in the presence of nickel catalysts is similar to that found for catalyst-free operation, but the nickel phase gasifies the water-soluble products into a CO-free mixture containing CO₂, H₂ and CH₄. Oil and char are also produced, but their yields were very low. The change of gas composition revealed that CO₂ and H₂ are primary products, although a minor proportion of CH₄ is later formed via methanation.

2.3.2.2 Gasification under subcritical conditions. Catalytic gasification at lower temperatures (ca. 720 K) has been studied by several groups with the objective to produce either H₂ or a methane-rich gas with a medium calorific value. At temperatures below 723 K, the catalytic effects of the reactor walls are minimized, and the product distribution depends on the catalyst incorporated. For a slurry of 10% wood sawdust, Waldner⁴⁷ reported conversion to gases of about 21% at 409 °C in the absence of a catalyst. For cellulose, Minowa et al.⁴⁵ found only about 10% carbon conversion to gases at 623 K and 60 min residence time. In addition, 67% of the feed carbon was recovered as char, 5% as bio-oil and 13% dissolved in the aqueous phase.

In USA, Modell et al.⁴¹ were the first to demonstrate that wood could be gasified in supercritical water without the formation of char and tars at low conversions. Elliott et al.^48–51 at the Pacific Northwest National Laboratory developed a process to gasify biomass under subcritical conditions (623 K, 20 MPa) using a variety of catalysts. Rhodium, ruthenium and nickel phases deposited on ZrO₂ (monoclinic), α-Al₂O₃, TiO₂ (rutile) and carbon. In general, the product gas consisted of more than 50% vol% CH₄, 40–50 vol% CO₂, less than 10 vol% H₂ and light hydrocarbons at trace levels. Sealock et al.⁵² reported CH₄ yield of 0.22 g CH₄/g wood with 33 vol% CH₄ using a stirred batch autoclave provided with a stainless steel liner operated at 623 K and 34 MPa for 150 min with a Ni catalyst. The Ni phase sintered rapidly although it could be stabilized by adding a second metal.⁵³ The presence of inorganic salts and/or when N- and S-compounds were present resulted in catalyst deactivation.

In Japan, Osada et al.⁵⁴ investigated several supported noble metal catalysts for the gasification of lignin and propyl phenols.^54–57 They suggested that the metal catalyses the decomposition of lignin to lower molecular weight products, i.e. alkylated phenols, and also causes the gasification of these phenolics. Only the first decomposition step was affected by the water density, however higher water densities enhanced the decomposition of lignin but did not influence the gasification of 4-propyl phenol. The catalyst became deactivated in the presence of sulfur compounds, although the rate of deactivation did not depend significantly on the type of the sulfur compound. They also pointed out that sulfur most likely blocks the sites responsible for the C–C bond scission and for the methanation but not for the water gas-shift reaction nor for the decomposition of C₁ compounds such as formaldehyde. Minowa et al.⁵⁸ carried out a study on cellulose gasification at 473–673 K using nickel catalysts in a stirred autoclave. The gas yield was found to be a strong function of the amount of nickel catalyst. These authors proposed a simplified mechanism involving water-soluble intermediates that can either polymerize to oil and further to char or react to form gases.^59–61 Park and Tomiyasu⁶² gasified cellulose, several polymers and model compounds on unsupported RuO₂ catalyst at 723 K and 44 MPa. Experiments with deuterated water enabled them to propose a redox mechanism including formation of CO and H₂O in a first step (eqn 10), and H₂ production in a second (eqn 11):

(−CH₂O)_n + nRuO₂ → nH₂O + nCO + nRuO (10)

nRuO + nH₂O → nRuO₂ + nH₂ (11)

the CO and H₂ produced then react together to yield CH₄ and CO₂. Thus, the net balance of biomass gasification is the formation of CH₄ and CO₂, without appearance of H₂. Later on, the same authors reported the gasification of pulp, waste paper, lignin and paper sludge.⁶³ Watanabe et al.^64–66 employed metal oxides (ZrO₂, CeO₂, TiO₂, MoO₃), instead of metals, as catalysts in the decomposition of glucose, cellulose, lignin and formaldehyde at 673–713 K in a batch autoclave. Zirconia did not formed methane and it doubled the gasification efficiency for cellulose and glucose.⁶⁶ In addition, the more basic oxides ZrO₂ and CeO₂ yielded, as expected, more methanol than MoO₃ and TiO₂.^64,65

In Europe, Waldner and Vogel⁶⁷ gasified spruce sawdust slurries with feed concentrations up to 30 wt% around 673 K in a batch reactor. A number of catalysts were screened for their activity and selectivity. From the catalysts, Raney nickel, 1%Ru/TiO₂ and 2%Ru/C were then tested for their long-term stability in a continuous test setup. A mixture of five organic compounds representing hydrolyzed wood was used as a feed. The Raney nickel catalyst sintered after a short time whereas 1%Ru/TiO₂ was not active enough. On the contrary, the 2%Ru/TiO₂ catalyst was hydrothermally stable for more than 200 h on-stream at 673 K and 30 MPa. Dosing 8 ppm Na₂SO₄ in the feed resulted in slow catalyst deactivation. This deactivation could be explained by an irreversible bonding of sulfate anion to the surface ruthenium, although sulfate might not be the actual catalyst poison under reaction conditions as it may be reduced to a sulfide species.⁶⁷

2.3.3 High temperature SCW gasification. High temperature SCW gasification is conducted in the 773–1073 K range. Due to the high reactivity of biomass at these temperatures, high gasification efficiency is achieved when the concentration of the precursor is low, although efficiency falls at higher concentrations. Organic feedstock such as glycerol and glucose can be gasified in the absence of catalysts. The gasification of glycerol under SCW and temperatures below 873 K is very low, but it steadily increases to asymptotic values around 973 K, and complete gasification can only be achieved at concentrations below ca. 3%. Under these operating conditions, the yields of H₂ and CO₂ increase sharply, while that of CO follows the opposite trend. These results indicate a stronger water gas-shift activity at temperatures above 873 K. The pressure of the process has little effect on either product gas composition or gasification efficiency within a wide range of pressures, including supercritical as well as subcritical pressures (60–400 bar). Additionally, the concentration of the feedstock has a major influence on the gas yield, as illustrated by the drop in H₂ yield and carbon gasification efficiency when the concentration of the organic feedstock exceeds 5–10%.

Most studies in the field of high temperature SCW biomass gasification to hydrogen have been conducted in USA, Europe and Asia. A brief account of the activities carried out by all these groups in SCW has been given recently by Peterson et al.⁶⁸

In USA, Antal et al.⁴⁴ investigated supercritical gasification of wet biomass and glucose. They reported that complete gasification of glucose can occur at 873 K 34.5 MPa and a 30 s residence time, and also that Inconel walls of the reactor strongly catalyse the water gas-shift reaction. It was shown that wood sawdust, dry sewage sludge or other particulate biomass could be mixed with a corn starch to form a viscous gel.⁶⁹ At the critical pressure of water (22 MPa) this paste vaporizes without the formation of char. A packed bed of carbon catalyst at 923 K causes the tarry vapors to react with water to produce H₂, CO₂, some methane and only traces of CO, but this catalyst deactivated after several hours on-stream. More recently, Hong and Spritzer⁷⁰ at General Atomics designed a continuous reactor with a thermal sleeve inside the pressure vessel to study SCW oxidation under 23.5 MPa pressure for several feedstocks such as wood, corn starch, coal and solid wastes. According to this concept, external heating was avoided because the heat required for running the reaction was released by the partial oxidation. In their preliminary experiments, ethanol was added to the feed with the objective of enhancing the heating. They reported yields in a directly heated system similar to Antal's yields in an indirectly heated system, after compensating for heating differences. Based on these works, General Atomics completed pilot-scale testing and concluded that by using negative cost waste streams they will be able to produce H₂ at cost of about $3/GJ.

In Europe, the earliest work carried out by Schneider et al.⁷¹ and Kruse et al.⁷² on SCW gasification of a range of feedstocks demonstrated complete gasification to H₂ at 898 K and 25 MPa in both batch and continuous tubular-flow reactors. They found that potassium compounds such as KOH and K₂CO₃ drastically increased the yield of hydrogen. Based on these and other studies a pilot plant was built at Forschungszentrum, Karlsruhe. This pilot, which was designed for a continuous flow capacity of 100 kg h⁻¹, maximum temperature of 930 K and 28 MPa, has been in operation since 2003. Using feedstocks of 9–25 wt% ethanol, pyroligneous acid and corn silage, very high yields of H₂ were obtained. Notwithstanding, the plant showed some plugging of the preheaters even at low reaction times (3.5 h). As indicated, most of the SCW works conducted at FZK employed alkali salts as catalysts because they catalyse the water gas-shift reaction.^71–74 While the mechanism by which K-salts increase H₂ yield is still unclear, there is an important body of current agreement that the interaction of alkali salt solutions with the reactor walls under SCW conditions is a potential source of hydrogen. First, the chromium layer of stainless steel is solubilised by alkali, then the Cr (and/or Fe) atoms located in sub-surface layers react with water molecules to yield additional H₂. However, recent SCW experiments conducted in sealed quartz tubes examined the influence of KOH and NaOH without the complication of the metal reactor walls. Using this experimental setup Kersten et al.⁷⁵ reported that NaOH increased the H₂ yield from 9.9% to 17–21% during the gasification of 17% glucose at 873 K under 30 MPa and 60 s residence time. Indeed, detailed studies are required to get a clearer understanding of the catalytic effect of alkali salts during SCW gasification.

At the University of Twente, Potic et al.⁷⁶ have constructed small sealed quartz capillaries (id = 1 mm) with an objective of avoiding the catalytic effect of metal surfaces. Kersten et al.⁷⁵ investigated in detail the SCW gasification of glucose, glycerol and pine-wood in these reactors. They found that, in the absence of catalyst, complete gasification was only achieved at concentrations 2 wt%, however incorporation of a Ru/TiO₂ catalyst allowed glucose solutions of up to 17 wt% to be gasified. To demonstrate the catalytic effect of the reactor walls when using conventional autoclave reactors, Inconel (a typical alloy of SCW metal reactors) powder was added and found to have a dramatic effect, increasing the gasification of a 5% glucose solution.

In China, the work by Guo et al.⁷⁷ was basically focused on the use of different reactor configurations including a miniature plant with a movable-piston feed pump to handle slurry feeds. They used a great variety of model compounds such as glucose, cellulose, lignin and xylan as well as feedstocks like sawdust, rice straw, rice shell, peanut shell, corn stalk, corn cob and wheat stalk under typical reaction conditions of 873 K, 25 MPa pressure and 10 wt% feed. The gasification efficiency achieved was very high, often above 80% conversion, together with high hydrogen yields for every feedstock, except lignin which was the most difficult compound to gasify.

2.3.4 Costs. SCW gasification may become an important technology for converting wet biomass to a pressurized and clean medium caloric value gas with high hydrogen content. Technical hurdles are not yet solved completely but significant progress has been made with reactors of different sizes. Chars and tars formation are the most significant technological problem. However, catalysis should be a solution to obtain higher yields of hydrogen and to decrease the amounts of chars and tars. As conversion yield reaches 98% with a proportion of hydrogen higher than 50% in the gaseous phase, the next step to consider in a successful transfer to the industrial scale is how to evaluate the energy cost of the process. Calzavara et al.⁷⁸ carried out an evaluation of the energy efficiency of biomass gasification The results showed that the energy efficiency from thermodynamic calculations reached 60% when considering hydrogen, carbon monoxide and methane as valuable species in an ideal case. Including energy recovery from water at 280 bar and 1010 K, the overall energy yield reached 90%. These calculations are overestimated because no heat losses have been considered. Notwithstanding, this analysis shows that the key point of the process is energy recovery as the chemical reaction is endothermic and requires high temperature and a rather large water/biomass ratio.

The cost of hydrogen production from supercritical water gasification of wet biomass has been analyzed by Demirbas.⁷⁹ Cost analysis of hydrogen produced via gasification of biomass in SCW has been made at a series of temperatures, pressures and different resident times. For all explored gasification conditions, analysis revealed that the cost H₂ produced via SCW gasification is several times higher than the current price of hydrogen from steam methane reforming. In addition, an evaluation of SCW gasification and biomethanation in terms of process cost, energy efficiency, and CO₂ emissions has been made by Matsumura.⁸⁰ Gasification of 1 t/d (dry) of water hyacinth was selected as a model case. Assumptions were made that the system should be energetically independent, that no environmentally harmful material should be released, and that carbon dioxide should be removed from the product gas. Energy efficiency, carbon dioxide payback time, and price of the product gas were chosen as indices for energy, environmental, and economic evaluation, respectively. Under these assumptions, supercritical water gasification appeared to be more advantageous over biomethanation, but the cost of the product gas was still 1.86 times more expensive than city gas (in Tokyo), although some improvement in efficiency of supercritical water gasification could be obtained by proper design of the heat exchanger. Finally, utilization of fermentation sludge made biomethanation much more advantageous.

2.4 Reforming of biomass-derived products

2.4.1 Ethanol. The fermentation of carbohydrates is the primary technology for the generation of liquid fuels from renewable biomass resources, among which ethanol is the major one (eqn 10). According to this equation, 1 mol CO₂ is produced per mol of C₂H₅OH, which gives a maximum carbon selectivity of 50%. Ethanol is particularly suitable for hydrogen production because of: (i) its low toxicity; (ii) its low production costs; (iii) the fact that it is a relatively clean fuel in terms of composition; (iv) its relatively high hydrogen content; and (v) its availability and ease of handling.

C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂ (12)

Hydrogen can be obtained directly from ethanol via two main processes: steam reforming (SRE, eqn 13, and partial oxidation POE, eqn 14). The overall processes are a complex convolution of elementary steps that involve several organic intermediates. Whereas POE offers exothermicity and a rapid response, SER is endothermic and produces greater amounts of hydrogen, resulting in higher system efficiencies. A third option combines the advantages of both approaches by co-feeding oxygen, steam and ethanol simultaneously via the oxidative reforming process (ORE, eqn 15).

CH₃CH₂OH + 3H₂O → 6H₂ + 2CO₂   (ΔH⁰ = + 173.3 kJ mol⁻¹) (13)

CH₃CH₂OH + 1.5O₂ → 3H₂ + 2CO₂   (ΔH⁰ = − 552.0 kJ mol⁻¹) (14)

CH₃CH₂OH + xO₂ + (3−2x)H₂O → (6−2x)H₂ + 2CO₂ (0 < x < 0.5)   (ΔH⁰ = ((3−2x)/3)*207.7 – (x/1.5)* 545.2 kJ mol⁻¹) (15)

2.4.1.1 Steam reforming. The overall steam reforming reaction of ethanol can be represented by eqn 13. This reaction is usually conducted at temperatures in the 823–1073 K range in the presence of a catalyst. Ethanol-reforming reactions involve several reaction pathways (dehydration, decomposition, dehydrogenation, coking) depending on the catalysts and reaction conditions⁸¹ (Fig. 7). Therefore, the choice of catalyst is essential in the reforming process. Reactions to avoid are those that lead to C and C₂H₄ products that are precursors of carbon deposition on catalyst surfaces. Consistent with this, catalysts for the steam reforming of ethanol to produce H₂ selectively must be able to: (i) dehydrogenate ethanol; (ii) break the carbon–carbon bonds of surface intermediates to produce CO and CH₄; and (iii) reform these C₁ products to generate hydrogen. On the basis of the influence of the nature of both the metal and support on the catalytic characteristics of supported metals, the choice of these elements is a key factor in developing supported catalysts that will fulfill the above requirements. Different oxide catalysts,⁸² metal-based catalysts (Ni, Co, Ni/Cu,)^83–85 and noble metal-based catalysts (Pt, Pd, Rh)^86–89 have proven to be active in the ethanol-reforming reaction.

Fig. 7 Reaction mechanism for steam reforming of ethanol from ref. 4.

Both the metallic function and the acid–base property of the catalyst play major roles in the reforming reaction of ethanol. This is illustrated by the Cu/Ni/K/γ-Al₂O₃ catalyst, which exhibits acceptable activity, stability and hydrogen selectivity at relatively low temperature (573 K) and atmospheric pressure.⁸³ In this catalyst, the copper is the active phase and nickel promotes C–C bond rupture, increasing hydrogen selectivity, while the potassium neutralizes the acidic sites of the γ-alumina substrate and improves the general performance of the catalyst. The nature of the support influences the catalytic performance of the supported catalyst for the steam reforming of ethanol, since it affects the dispersion and stability of the metal and may participate in the reaction. Lanthanum oxide is a particularly suitable support for the metallic function of ethanol-reforming catalysts. The Ni/La₂O₃ or Ni–La₂O₃/Al₂O₃ catalysts have high activity and long-term stability for hydrogen production.^90,91 A 20% Ni/La₂O₃/Al₂O₃ catalyst has good stability at 1023 K for reaction times over 150 h, with only a small reduction in ethanol conversion from 95% to 90%, while hydrogen selectivity remains essentially unchanged. These results indicate the uniqueness of the Ni–La₂O₃ system in terms of its protracted stability. The unusual stability of the Ni–La₂O₃ catalyst has been attributed to the scavenging of coke deposition on the Ni surface by lanthanum oxycarbonate species existing on top of the Ni particles.⁸⁷ The effects of basic additives (K, Mg, Ca, Ce) that favour water adsorption and OH surface mobility in Al₂O₃ supports, to lower the rate of coke deposition on catalyst surfaces have also been investigated on nickel-based catalysts.^92,93 Coke formation on bare and Ce-, K- or Mg-modified catalysts does occur, but at orders of magnitude lower than that in alumina-supported Ni catalysts.

Cobalt-based catalysts have also been studied in the ethanol steam reforming reaction. Llorca et al.⁸⁵ performed the reaction between ethanol and water in the 573–723 K temperature range at atmospheric pressure over supported cobalt catalysts. The ZnO-supported cobalt catalyst had a very high catalytic performance. Using an EtOH)/H₂O = 1 : 13 (molar ratio) mixture, total conversion of ethanol and high values of H₂ and CO₂ were obtained, in the absence of deactivation. Complete EtOH conversion was also reached on the Co-free ZnO substrate, but the yields of H₂ and CO₂ alone were found to be substantially lower. The decomposition of EtOH into acetone occurs to a large extent on Co/ZnO catalysts. Since this reaction results from consecutive reactions, such as dehydrogenation and aldol condensation, activity tests conducted at low contact times have indicated that the reforming reaction is relatively fast, while EtOH decomposition to acetone via aldol condensation of acetaldehyde is supressed. In addition, the Co/ZnO catalyst accumulates a considerable amount of carbon throughout the reaction, which causes the deactivation of the cobalt catalyst.

Noble metals supported on porous oxide substrates (A₂O₃, SiO₂, CeO₂, TiO₂ and MgO)^{86,89,94–100} are highly active in the steam reforming of ethanol to CO_x and H₂. The support plays a significant role in the steam reforming of ethanol over noble metals. When CeO₂/ZrO₂, which has oxygen storage capability, is used as the support for noble metals, ethylene formation is not observed and the order of activity at higher temperature is Pt ≈ Rh > Pd.⁸⁶ Alumina-supported catalysts are very active at low temperatures in the dehydration of ethanol to ethylene. At higher temperatures, ethanol is converted into H₂, CO, CO₂ and CH₄, with an activity order of metals as follows: Rh > Pd > Ni = Pt.⁸⁵ Auprêtre et al.⁹⁹ studied the effect of both the metal and the support in the steam reforming of ethanol. They found that the hydrogen yield on alumina-supported metal catalysts at 973 K decreased in the following order: Rh > Pd > Pt > Ru. They concluded that the high activity of the metals in ethanol steam reforming and their poor efficiency in the water gas-shift reaction would provide active and selective catalysts for ethanol reforming. Auprêtre et al.⁹⁹ also reported that the H₂ yield on Rh/CeO₂ was higher than that on Rh/Al₂O₃ at 873 K. It was concluded that a metal–ceria interaction affects the absorption–decomposition of ethanol to CH₄ and CO products and their subsequent reforming reactions with steam. On the Rh/Al₂O₃ catalyst, Cavallaro et al.¹⁰⁰ reported that ethanol is firstly converted to ethylene by dehydration on the Al₂O₃ surface, or to acetaldehyde by dehydrogenation on Rh particles. The acetaldehyde undergoes decarbonylation on the rhodium surfaces to form methane and CO, while ethylene is also steam reformed on metal particles to C₁ (very fast reactions). Liguras et al.⁸⁷ also found that among the low-loaded catalysts, Rh was significantly more active and selective towards hydrogen formation than Ru, Pt and Pd. The catalytic performance of Rh was greatly improved by the increase of metal loading. In addition, the 5 wt% Rh/Al₂O₃ recorded good stability at 923 K and high H₂ selectivity (up to 95%) without carbon formation as demonstrated by long-term activity tests.¹⁰⁰

2.4.1.2 Catalytic partial oxidation. The partial oxidation of ethanol (POE, eqn 14) has been investigated with less intensity than in the case of steam reforming. Partial oxidation is a very interesting process for H₂ production because the catalysts used can be run autothermally, thereby eliminating the need for external heat. POE is much faster than the catalytic steam reforming that allows fast start-up and short response times to variations in H₂ production. It is a slightly endothermic reaction, so alone it will not produce the heat necessary for autothermal operation. Therefore, a small part of ethanol must be combusted to generate the heat for the operation in the 973–1273 K range necessary for the reaction to be sustained. It is emphasised that the pure POE process is not indicated for bio-ethanol reforming since bio-ethanol is an ethanol–water mixture in which the removal of all the water has a significant cost. Therefore, the processes for bio-ethanol partial oxidation are usually combined with steam reforming in autothermal schemes with the stoichiometry shown in eqn 16. Additionally, adding water to the reaction stream is very useful since catalyst stability is improved while coke formation is minimized. In addition, as the H atoms from H₂O can also be converted into H₂, complete conversion of the ethanol and water from this reaction could generate 5H₂ per C₂H₅OH, which gives a maximum H₂ selectivity of 5/3 or 167%.

C₂H₅OH + 2 H₂O + ½O₂ → 2CO₂ + 5H₂   (ΔH⁰_298K = −68.2 kJ mol⁻¹) (16)

The generation of hydrogen from ethanol via catalytic autothermal partial oxidation has been performed at temperatures of 700–1000 K using noble metal-based catalysts.^101,102 Ethanol reforming follows a very complex pathway, including several reaction intermediates formed and decomposed on both the metal and the support.^103,104 In the light of the above studies, it has been argued that the ethoxy species generated on the metal and on the support can be decomposed on the metal sites, forming CH₄, H₂ and CO, while part of the ethoxy species generated on the supports is further oxidized to acetate species, which decompose to CH₄ and/or oxidize to CO₂via carbonate species.¹⁰¹ Thus, supports with redox properties that help the oxidation of ethoxy species and metals with a high capacity to break C–C bonds and to activate C–H bonds are suitable for use in catalysts applied to the POE. Salge et al.¹⁰² studied the effect of the nature of the metal (Rh-, Pd-, Pt-) on the performance of catalysts supported on Al₂O₃ and CeO₂. The yield of H₂ production for catalysts supported on Al₂O₃ followed the order: Rh–Ru > Rh > Pd > Pt. Rh supported on CeO₂ was the most stable and gave greater H₂ selectivity than noble metals supported on Al₂O₃. The better activity and stability associated with the presence of CeO₂ can be related to the capacity of CeO₂ to store oxygen and make it available for reaction via a redox reaction.¹⁰⁵

2.4.2 Sugars. The aqueous-phase reforming (APR) process for the conversion of sugars and polyols into H₂ and C₁–C₁₅ alkanes has been developed by Dumesic et al.^106–109 Hydrogen, CO₂, CO, and light alkanes are produced by the APR of the aqueous sugar feeds under pressures ranging from 10 to 50 bar (eqn 17). One major advantage of APR is that it produces a hydrogen-rich stream with low levels of CO (100–1000 ppm), which makes it particularly suitable for feeding polymer electrolyte membrane fuel cells.

C₆H₁₄O₆ + 6H₂O → 6CO₂ + 13H₂   (ΔH⁰_298K = +443.5 kJ mol⁻¹) (17)

The reaction pathway of APR involves cleavage of C–H, C–C, and O–H bonds of sugar molecules to form adsorbed species on the catalyst surface. Adsorbed CO must be removed by the WGS reaction to form CO₂ and additional H₂. Undesired parallel reactions also occur and proceed via C–O bond splitting followed by hydrogenation to yield alcohols or even acids. Thus, good catalysts for the production of H₂ by APR reactions must be highly active for C–C bond cleavage and also capable of removing adsorbed CO by the WGS reaction, but must not facilitate C–O bond cleavage and the hydrogenation of carbon oxides. H₂ selectivity depends on the feed sugar, catalyst, and reaction conditions. As a general trend, H₂ selectivity decreases upon increasing molecular size of the feed molecule.

Kinetic studies were conducted for the APR of ethylene glycol (a probe molecule for sorbitol) over silica-supported Pd, Ni, Pt, Ir, Ru, and Rh catalysts at moderate temperatures (483–498 K) and moderate pressure (22 bar). The catalytic activity for APR of ethylene glycol, as measured by the rate of CO₂ formation per surface atom at 483 K followed the order: Pt ∼ Ni > Ru > Rh ∼ Pd > Ir.¹⁰⁹ Silica-supported Ni, Ru, and Rh catalysts recorded low selectivity for H₂ production and high selectivity for alkane production. In addition, the Ni/SiO₂ catalyst became rapidly deactivated at 498 K. On the other hand, Pt/SiO₂ and Pd/SiO₂ catalysts exhibited higher selectivity for the production of H₂, with lower rates of alkane production. It was also found that both the activity and selectivity of Pt-based monometallic catalysts can be enhanced by depositing Pt phase on TiO₂, carbon, and Al₂O₃ substrates¹¹⁰ or by adding Ni, Co, or Fe to a monometallic Pt/Al₂O₃ catalyst.¹¹¹ Alumina-supported PtNi and PtCo catalysts with Pt/Co or Pt/Ni atomic ratios ranging from 1 : 1 to 1 : 9 had the highest turnover frequencies for H₂ production (moles of H₂ per mole of surface site measured by CO adsorption) with values of 2.8–5.2 min⁻¹ for APR of ethylene glycol solutions at 483 K, compared to a value of 1.9 min⁻¹ for the monometallic Pt/Al₂O₃ under similar reaction conditions.

Nickel catalysts are also active for APR reactions; however, they have low selectivity and stability. The H₂ selectivity of Ni-based catalysts can be enhanced by adding Sn to the Ni catalyst, whereas its stability can be improved by using bulk Ni catalysts, for example, Raney Ni.¹¹² The rates of H₂ production by APR of ethylene glycol over a SnNi catalyst with Ni/Sn atomic ratios up to 14 : 1 are comparable to a 3 wt% Pt/Al₂O₃ catalyst, based on reactor volume. The incorporation of Sn to Raney Ni catalysts markedly decreases the rate of methane formation from reactions of CO_x with H₂, while maintaining the high rates of C–C cleavage necessary for the production of H₂. Nonetheless, the reactor must operate at near the bubble-point pressure of the feed and moderate space times to achieve high H₂ selectivities over Raney SnNi catalysts. Remarkably, these Raney SnNi catalysts are stable for more than 250 h time-on-stream.¹¹²

2.5 Biological hydrogen production

Biological H₂ production processes are becoming important because they can use renewable energy resources, and they usually operate at ambient temperature and atmospheric pressure. A number of microorganisms possess enzymes that can produce H₂ from water when they are exposed to an outside energy source, such as sunlight. Specific ways in which microorganisms can produce H₂ have recently been summarized by Das et al.¹¹³ There are three main processes for biological H₂ production: (i) biophotolysis of water using green algae and blue-green algae (cyanobacteria); (ii) photofermentation; and (iii) dark fermentation.

The direct biophotolysis (type (i) process) of H₂ production is a biological process that uses solar energy and photosynthetic algae systems to convert water into chemical energy:

2H₂O + photons → 2H₂ + O₂ (18)

The two photosynthetic systems responsible for the photosynthesis process are photosystem I (PSI), which produces reductant for CO₂, and photosystem II (PSII) which splits water to evolve O₂. The two photons obtained from the splitting of water can either reduce CO₂ by PSI or form H₂ in the presence of hydrogenase. Due to the lack of hydrogenase in plants, only CO₂ reduction takes place. By contrast, green algae and cyanobacteria (blue-green algae) contain hydrogenase and thus have the ability to produce H₂.¹¹⁴ In these organisms, electrons are generated when PSII absorbs light energy, which is then transferred to ferredoxin. A reversible hydrogenase accepts electrons directly from the reduced ferredoxin to generate H₂ in the presence of hydrogenase:

However, it is not yet clear how to manage such a process for efficient H₂ production. Indeed, when dark adapted (hydrogenase induced) cells are exposed to light, photosynthetic O₂ production commences inhibiting H₂ evolution.¹¹⁵

Type (ii) biological H₂ production process is photofermentation. Purple non-sulfur bacteria produce hydrogen mainly due to the presence of nitrogenase under oxygen-deficient conditions using light energy and reduced compounds (organic acids). The reaction is as follows:

C₆H₁₂O₆ + 12H₂O + photons → 12H₂ + 6CO₂ (20)

Major process bottlenecks involve low photochemical efficiencies (3–10%). Moreover, the lack of uniformity of the light distribution in the reactor also contributes to lowering the overall light conversion efficiency.

Dark fermentation (type (iii) process) is a ubiquitous phenomenon under anoxic or anaerobic conditions. The oxidation of the substrate by bacteria generates electrons, which need to be disposed of in order to maintain electrical neutrality. Under aerobic conditions, O₂ serves as the electron acceptor, whereas under anaerobic or anoxic conditions other compounds, such as protons, act as the electron acceptor and are reduced to molecular H₂.¹¹⁶ Carbohydrates, mainly glucose, are the preferred carbon sources for this process, which predominantly give rise to acetic and butyric acids together with H₂ evolution:¹¹⁷

C₆H₁₂O₆ + 2H₂O → 2CH₃COOH + 2CO₂ + 4H₂ (21)

C₆H₁₂O₆ + 2H₂O → CH₃CH₂COOH + 2CO₂ + 2H₂ (22)

Despite having a higher evolution rate, the yield of H₂ from the fermentation process is lower than that of other chemical–electrochemical processes. The theoretical H₂ yield is 4 mol of H₂/mol of glucose when the end product is acetic acid, while 2 mol of H₂/mol of glucose will be obtained if the metabolic end product is butyric acid. In practice, the yields are low since the end products contain both acetate and butyrate.¹¹⁷ Besides, as yields increase the reaction becomes thermodynamically unstable. Another constraint of the process is the low conversion efficiency of the substrate used.

A challenging problem in establishing biohydrogen as a source of energy is the renewable and environmentally-friendly generation of large quantities of H₂ gas. However, two major aspects need vital attention, viz., a suitable renewable biomass/wastewater and ideal microbial consortia that can convert this biomass efficiently to H₂. Comparative studies on the available processes indicate that biohydrogen production requires greater improvement on the process, mainly with respect to H₂ yield from the cheaper raw materials. The future of biological H₂ production depends not only on research developments, i.e., the improvement in efficiency through genetically engineered microorganisms and/or the development of bioreactors, but also on economic considerations.

3. The solar option

Of the various renewable energy sources, by far the largest resource is provided by the sun. More energy from sunlight strikes the earth in 1 h (4.3 × 10²⁰ J) than all of the energy currently consumed on the planet in 1 yr (4.1 × 10²⁰ J in 2001).¹¹⁸ Yet, in 2001, only <0.1% of electricity and <1.5% of fuels (mostly from biomass) were provided by a solar source.^118,119 Against the backdrop of the daunting carbon-neutral energy needs of our global future, the large gap between our present use of solar energy and its enormous undeveloped potential defines a compelling imperative for science and technology in the 21st century. Indeed, the solar energy striking the earth exceeds what can possibly be consumed by any technologically-advanced society. However, to make a material contribution to the primary energy supply, solar energy must be captured, converted, and stored to overcome the diurnal cycle and the intermittency of the terrestrial solar resource. Undoubtedly, the most attractive method for this energy conversion and storage is in the form of H–H bonds.

Within this scenario, the conversion of solar energy into hydrogen via the water splitting process assisted by photo-semiconductor catalysts is one of the most interesting ways of achieving clean and renewable energy systems.^120,121

3.1 Basic principles of water splitting

Among the methods for H₂ generation outside the C-cycle, photocatalytic hydrogen production via water splitting under visible light irradiation attracts the greatest attention for its potential to use the abundance of solar energy (the maximum direct insolation frequently reaches ca. 700 W m⁻² in the sunbelt regions) and water. Thermodynamically, the overall water splitting reaction is an uphill reaction with a highly positive change in Gibbs free energy (ΔG⁰ = + 237.2 kJ mol⁻¹):

H₂O(l) → H₂(g) + ½O₂(g)   (ΔG⁰ = + 237.2 kJ mol⁻¹) (23)

Fig. 8 shows a sketch diagram of the basic principle of overall water splitting on a solid photocatalyst. Under irradiation with an energy equivalent to or greater than the bandgap (E_g) of the semiconductor photocatalyst, the electrons (e⁻) of the valence band are excited into the conduction band (CB) while the holes (h⁺) are left in VB. Electrons and holes that migrate to the surface of the semiconductor without recombination can respectively reduce and oxidize the water molecules adsorbed on the semiconductor surface. To achieve overall water splitting, the bottom of the CB must be located at a more negative potential than the reduction potential of H⁺/H₂ (0 V vs. NHE at pH = 0), while the top of the VB must be positioned more positively than the oxidation potential of H₂O/O₂ (1.23 V vs. NHE). Therefore, according to this theoretical value, it can drive water splitting only if the photon energy is equal or superior to 1.23 eV. This energy is equivalent to the energy of a photon with a wavelength of around 1010 nm, indicating that visible light is energetically sufficient for the decomposition of water. However, the activation energy barrier in the charge transfer reaction between the water molecules and the semiconductor surface requires photon energy greater than the bandgap of the semiconductor to drive overall water splitting at a measurable reaction rate. The potential of the band structure is precisely the thermodynamic requirement. Other factors such as charge separation, mobility and the lifetime of photogenerated electrons and holes also have an effect on the photocatalytic properties of semiconductors. The generation and separation of photoexcited carriers (electrons and holes) strongly depend on both the presence of co-catalysts on the surface of the photocatalysts and the latter's structural and electronic properties. In general, high crystallinity has a positive effect on activity since the density of defects, which act as recombination centres of electrons and holes, decreases when crystallinity increases. Co-catalysts are usually loaded onto the photocatalysts to assist the redox reactions taking place on their surfaces.^122,123 The co-catalysts are typically a noble metal (e.g. Pt, Rh) or metal oxide (e.g. NiO, RuO₂) loaded on the surface as nanoparticles to reduce the electron–hole recombination and to produce active sites that reduce the activation energy for gas evolution.

Fig. 8 Sketch diagram of the basic principle of overall water splitting on a solid photocatalyst. Irradiation with an energy equivalent to or greater than the bandgap (E_g) of the semiconductor photocatalyst, the electrons (e⁻) of the valence band are excited into the conduction band (CB) while the holes (h⁺) are left in VB.

3.2 Photocatalysts for water splitting

Several types of semiconductors, over 130 materials including oxides, nitrides, sulfides⋯^124–141 have been reported to act as efficient photocatalysts for hydrogen evolution via water splitting. Among these photocatalysts, the higher quantum yields (QY) are reported for Ba-doped Sr₂Nb2O₇ (50% QY, pure water, UV)¹³⁵ and NiO/NaTaO₃ (56% QY, pure water at 270 nm).¹²⁹ Unfortunately, these exciting developments have only a limited value for practical, large-scale hydrogen manufacture because UV light accounts for only about 3–4% of solar radiation energy. Therefore, regarding the use of solar energy, it is essential to develop photocatalysts that split water efficiently under visible light (λ > 400 nm). The materials for visible light-driven photocatalysts had been quite limited. However, many oxides, sulfides, oxynitrides and oxysulfides have recently been found to be active for H₂ and O₂ evolution under visible light irradiation.^136–140 So far, the maximum quantum efficiency over visible light-driven photocatalysts achieves only a few percent at wavelengths as long as 500 nm (Cr/Rh–GaN/ZnO 2.5% QY, pure water, visible light).¹⁴¹ This value is still far from the QE of 30% marked as the initial starting point for practical application.¹⁴² Hence the development of new photocatalyst materials is still a major issue. There are several common approaches adopted in the search for photocatalyst for splitting water under visible light irradiation: (i) find new materials, (ii) tune the band gap energy by modifying cations or anions of UV-active photocatalysts with substitutional doping and, (iii) fabricate multi-component photocatalyst by forming solid solutions. The following sections will review the above approaches for the development of active visible light photocatalysts for water splitting.

3.2.1 Ti oxide-based photocatalysts. Ti-based oxide materials were found to be efficient photocatalysts for various photocatalytic reactions. However, most Ti oxide photocatalysts are unable to split water under visible light. This is mainly due to the large band gap of TiO₂ (3.2 eV), which allows solely for the use of a small fraction of solar spectrum (UV fraction, 3–4% of total solar spectrum). Intensive studies have been carried out to improve the visible light sensitivity of Ti oxide-based catalysts. One of the strategies for inducing visible light response in TiO₂ was the chemical doping of TiO₂ with metal ions with partially filled d-orbitals (V, Cr, Fe, Co, Ni...).^127,143,144 The metal ions may be incorporated into the TiO₂ by chemical (impregnation, precipitation...) and physical methods. Although TiO₂ chemically doped with metal ions could induce a visible light response, these catalysts showed limitations in reactivity for practical applications because dopants in the photocatalyst act not only as visible light adsorption centres but also as recombination sites between photogenerated electrons and holes.¹⁴⁵ In contrast to chemical doping, physical doping by the advanced ion-implantation technique has proven to be an effective method to improve the TiO₂ visible light response.^146,147 In this case, the absorption band of implanted TiO₂ shifts to visible light regions (up to 600 nm) due to the substitution of the Ti ions in the TiO₂ lattice with the metal ions. Thin TiO₂ films implanted with metal ions such as Cr and V have recorded photoactivity under visible light for H₂ evolution from aqueous solution involving methanol as sacrificial reagent with a quantum yield of 1.25 (λ = 420 nm).¹⁴⁸ Although the ion-implantation method provides a way of modifying the optical properties of TiO₂, it was not practical for mass production due to the high cost of the ion-implantation apparatus used to develop these TiO₂ modified photocatalysts.

The second main strategy followed to improve the visible light response of TiO₂ is related to the doping of anions such as N,¹⁴⁹ S¹⁵⁰ or C¹⁵¹ as substitutes for oxygen in the TiO₂ lattice. For these anion doped TiO₂ photocatalysts, the mixing of the p states of doped anion (S, N or C) with the O 2p states was reported to shift the valence band edge upwards to narrow the band gap energy of TiO₂. However, these materials do not record activity for pure water splitting, although the conduction and valence bands have enough potential for the reduction and oxidation of water, due to the large over-potential for H₂ and O₂ evolution on the surface of these photocatalysts.

The development of visible light-driven photocatalysts by doping metal ions into SrTiO₃ is also reported.¹⁵² A survey of dopants for SrTiO₃ revealed that the doping of Rh or the co-doping of Cr³⁺–Ta⁵⁺ or Cr³⁺–Sb⁵⁺ were effective in making SrTiO₃ visible light-responsive. Rh-doped SrTiO₃ produced H₂ from an aqueous methanol solution, while Cr–Ta-codoped-SrTiO₃ was effective for O₂ evolution from aqueous AgNO₃ solutions. These doping photocatalysts that show activities for only half reactions can be used to construct Z-scheme photocatalysts for overall water splitting.¹²⁷

3.2.2 Other transition metal oxide-based photocatalysts. In the doped photocatalysts commented on in the previous section, the formation of recombination sites by the dopant atoms is more or less inevitable. Moreover, the local energy levels formed by doping are usually discrete and thus unsuitable for the migration of holes formed in this level. Kudo et al. employed Bi³⁺ and Ag⁺ to modify the valence band of oxide photocatalysts.^130,153 BiVO₄ with scheelite structure and AgNbO₃ with perovskite structure were found to have photocatalytic activities for O₂ evolution from aqueous silver nitrate solution under visible light irradiation. The photocatalytic activity of BiVO₄ was higher than that of the WO₃, which is a well-known photocatalyst for the O₂ evolution reaction under visible light irradiation. The valence bands of BiVO₄ and AgNbO₃ consist of Bi and Ag orbitals mixed with O 2p states resulting in the increase of the valence band potentials and the decrease in the band gap. Although these catalysts did not have potential for H₂ production, they can be used as O₂ generators in Z-scheme photocatalysts for overall water splitting.

3.2.3 Nitrides and oxysulfide photocatalysts. Oxynitrides and oxysulfides are new types of visible light-driven photocatalysts.^{139,141,154–157} The valence bands of these photocatalysts consist of N 2p or S 3p orbitals in addition to O 2p, whereby band gap energy decreases without affecting the conduction band level, thus producing a visible light-driven photocatalyst with band edge potentials suitable for overall water splitting. For example, promising results have been reported for Sm₂Ti₂S₂O₅, which has proven to be responsive to excitation at wavelengths up to ca 650 nm.^139,157 The Sm₂Ti₂S₂O₅ works as a stable photocatalyst for the reduction of H⁺ to H₂ and oxidation of H₂O to O₂ in the presence of a sacrificial electron donor and acceptor under visible light irradiation. Other Ln₂Ti₂S₂O₅ (Ln, Pr, Nd, Gd, Tb, Dy, Ho and Er) oxysulfides with layered perovskite structure are also active for H₂ or O₂ evolution from aqueous solutions containing sacrificial reagents under visible irradiation.¹³⁹

Oxynitrides such as TaON, LaTiO₂N and (Ga_1−xZn_x)(N_1−xO_x) have evolved as promising water-splitting catalysts that can operate under visible light and without sacrificial reagents. Among these, the most promising formulation was the solid solution of GaN and ZnO ((Ga_1−xZn_x)(N_1−xO_x)). Density functional calculations indicated that the bottom of the conduction band of the GaN–ZnO solid solution was mainly composed of 4s and 4p orbitals of Ga, while the top of the valence band consisted of N 2p orbitals, followed by Zn 3d orbitals. The presence of Zn 3d and N 2p electrons in the upper valence band might provide p–d repulsion for the valence band resulting in a narrowing of the band gap. H₂ and O₂ were found to evolve steadily and stoichiometrically from solid solution photocatalysts with an average quantum efficiency of 0.14% in the range of 300–480 nm.¹⁴¹ Different transition metal oxides have been examined as cocatalysts to promote activity of (Ga_1−xZn_x)(N_1−xO_x) solid solutions. Among the various materials examined, the largest improvement in activity was obtained when (Ga_1−xZn_x)(N_1−xO_x) was loaded with a mixed oxide of Rh and Cr. The quantum efficiency of the Rh–Cr loaded (Ga_1−xZn_x)(N_1−xO_x) photocatalyst for overall water splitting reaches ca. 2.5% at 420–440 nm.¹⁵⁴

3.2.4 Metal sulfide-based photocatalysts. Sulfide photocatalysts, which have narrow band gaps and valence bands at relatively negative potentials compared to oxides, are good candidates for visible light-driven photocatalysts. However, semiconductor sulfides, which cause photocorrosion under visible light irradiation, are not suitable for water-splitting unless appropriate strategies are designed to minimize this problem. There are strategies for reducing photocorrosion by increasing the photocatalyst efficiency of semiconductor sulfides. A common method consists in the use of suitable sacrificial reagents. A Na₂S/Na₂SO₃ mixture added to the water/semiconductor suspension is considered a very effective way of reducing the photocorrosion of sulfide-based catalysts.¹⁵⁸ Although the use of sacrificial reagents makes the process unsustainable, the production of H₂ using sulfide sacrificial reagents might have practical application using by-products from the desulfurization process in petrochemical plants.
3.2.4.1 CdS-based photocatalysts. Among the available sulfide semiconductors, nanosized CdS is an interesting photocatalyst material, since it has a narrow band gap (2.4 eV) and a suitable conduction band potential to effectively reduce H⁺.^159–165 However, the photocatalytic properties of CdS are limited due to its photocorrosion under visible light irradiation.¹⁶⁶ Although photocorrosion can be significantly reduced by adding sacrificial reagents, life cycle assessment studies associated with the processes of water splitting using Cd-based photocatalysts must be performed in order to evaluate the potential environmental impacts associated with Cd. In spite of the drawbacks associated with CdS, considerable efforts are still being made to improve its photocatalytic properties. The steps reported in the literature to improve the activity of CdS include: (i) changes in the structural characteristics of CdS, (ii) combination of CdS with different elements or semiconductors to form mixed photocatalysts with different band gap size and, (iii) addition of small amounts of metals or transition metal oxides (cocatalysts) to CdS.
Combination of CdS with CdO and ZnO. Changes in the photoactivity of CdS can be achieved by combining the CdS semiconductor with other semiconductors with different energy levels (ZnS,¹³¹ ZnO,¹⁶⁷ TiO₂,¹⁶⁸ AgGaS₂¹⁶⁹…). It is claimed that photogenerated electrons in these composite systems move from CdS to the attached semiconductors, while photogenerated holes remain in CdS. This charge-carrier separation stops charge recombination, therefore improving the photocatalytic activity of CdS. The optoelectronic properties of ZnO and CdO make these materials interesting as semiconductors for combining with CdS.^152,171 CdO has a band gap in the interval 2.2–2.4 eV with high transmittance and very low resistance, while ZnO has a wide band gap (3.2 eV) that may improve the photoabsorption ability of CdS. In spite of the absence of studies in the literature on the effect of the presence of CdO and ZnO on the activity of CdS under visible light, there are interesting results showing the increase in CdS activity for samples mixed with CdO and ZnO.¹⁷² The CdS–CdO–ZnO catalyst prepared by sequential precipitation of CdS and CdO and ZnO showed that the activity of CdS in the hydrogen produced under visible light from aqueous solutions containing SO₃²⁻ + S²⁻ as sacrificial reagents increases with the addition of CdO and ZnO to CdS (Fig. 9). Taking into account the absence of significant changes in the structure of CdS with the addition of both oxides, the observed differences in activity of CdS–CdO–ZnO mixtures with respect to bare CdS are related to a better charge separation associated with the diffusion of photoelectrons generated in CdS toward surrounding CdO and ZnO.

Fig. 9 Hydrogen evolution rate from aqueous solution containing Na₂S + Na₂SO₃ under visible light irradiation over CdS, CdS–CdO and CdS–CdO–ZnO catalysts (catalyst: 0.1 g, reactant solution 150 ml (0.1M Na₂S +0.04M Na₂SO₃, 150 W Xe lamp).

The effect of thermal treatments under inert atmosphere on crystallinity and the phase transformations of precipitated CdS–CdO–ZnO catalyst was also studied.¹⁷²Fig. 10 shows hydrogen evolution under visible light irradiation over CdS–CdO–ZnO catalysts treated at 573, 773 and 923 K. Hydrogen production was found to increase with annealing temperature, with a decrease in activity for the sample treated at temperatures higher than 773 K. The physicochemical characterization of thermally treated CdS–CdO–ZnO shows that heating induces significant changes at structural and chemical levels. Fig. 11 shows that the catalyst has crystal structures of CdS that gradually change from a cubic structure to a more crystalline hexagonal phase with the increase in annealing temperature. The evolution of CdS phase in the CdS–CdO–ZnO mixture during thermal treatment is similar to the evolution observed in the bare CdS sample, suggesting that neither the presence of CdO nor that of ZnO modifies the thermal evolution of the CdS phase. Changes in structure after thermal treatment have a bearing on the photophysical properties of CdS–CdO–ZnO samples, recording an increase in band gap size from 2.24 to 2.33 eV for the sample annealed at temperatures higher than 773 K, associated with changes in its order degree. When comparing the photophysical properties of thermally treated CdS–CdO–ZnO samples with photocatalytic hydrogen production results, no direct correlation was found between visible light absorption and photoactivity. From this observation, it is inferred that the effective photo-utilization of electrons and holes, rather than photoabsorption ability, plays a key role in the catalytic activity of CdS–CdO–ZnO mixtures. The differences in activity are possibly related to the reduction in CdS surface exposure and contact between CdS and CdO/ZnO, as a consequence of the thermal crystalline growth that decreases the generation of photoelectrons in CdS and their diffusion toward surrounding CdO/ZnO nanoparticles, thereby increasing the probability of electron–hole recombination.

Fig. 10 Hydrogen evolution from aqueous solution containing Na₂S + Na₂SO₃ under visible light irradiation over CdS–CdO–ZnO catalysts annealed at different temperatures: (■) 573 K, (●) 773 K and (▲) 973 K (catalyst: 0.1 g, reactant solution 150 ml (0.1M Na₂S + 0.04 M Na₂SO₃, 150 W Xe lamp).

Fig. 11 XRD patterns of CdS–CdO–ZnO catalysts treated under inert flow at different temperatures; (a) 573 K, (b) 773 K and (c) 973 K ((*) hexagonal CdS, (o) cubic CdS (+) CdO).

As mentioned previously, the third way to enhance the activity of CdS photocatalyst is through the addition of small amounts of metals (cocatalyst) such as Pt, Rh and Ru.^170,171 Within this scenario, we have investigated the structural changes in CdS–ZnO–CdO mixtures associated with the photodeposition of metal cocatalysts (Pt and Ru).¹⁷² The photodeposition of Pt or Ru on the CdS–ZnO–CdO mixture changes the crystalline size of CdS with hexagonal phase as consequence of the photoetching phenomena on the CdS structures caused by the light irradiation used for photodeposition.¹⁷³ The absorption properties of CdS–ZnO–CdO mixture after the photodeposition of Pt or Ru showed only a slight blue-shift in the adsorption edge with respect to unloaded photocatalyst, consistent with the observed reduction in the crystalline degree of CdS after the photodeposition of noble metals. Hydrogen production on Ru- and Pt/CdS–ZnO–CdO catalysts markedly depends on the type of noble metal incorporated, with the Ru-based catalyst recording a higher level of hydrogen production than that obtained over its Pt counterpart (Fig. 12).

Fig. 12 Hydrogen evolution from aqueous solution containing Na₂S + Na₂SO₃ under visible light irradiation over CdS–CdO–ZnO with Pt and Ru cocatalysts: (■) bare CdS–CdO–ZnO, (●) Pt/CdS–CdO–ZnO and (▲) Ru/CdS–CdO–ZnO (catalyst: 0.1 g, reactant solution 150 ml (0.1M Na₂S + 0.04 M Na₂SO₃, 150 W Xe lamp).

The levels of hydrogen production over the noble metal catalysts contrast with the low rate obtained over the noble-metal-free CdS–ZnO–CdO counterpart. In particular, ruthenium strongly enhances the rate of hydrogen production by a factor of almost 50 times with respect to the CdS–ZnO–CdO substrate. The enhancement of photocatalytic activity has been explained in terms of a photoelectrochemical mechanism in which the electrons generated by irradiation of CdS are transferred to the loaded metal particles, decreasing electron–hole recombination. The XRD results showed similar changes in the CdS–ZnO–CdO substrate after Pt and Ru deposition, thereby indicating that the observed differences in activity are related to the state and surface concentration of noble metal entities after photodeposition onto the CdS–ZnO–CdO substrate. Accordingly, the better activity observed for Ru catalysts was associated with the presence of Ru oxide entities of high intrinsic activity with good surface coordination with cadmium sulfide particles that enhance the electronic transfer between both phases, thereby increasing the efficiency of water splitting by decreasing the probability of electron–hole recombination.

CdS–ZnS solid solutions. The incorporation of elements in the structure of CdS making solid solutions is a powerful strategy for improving the photocatalytic properties of CdS.¹⁷⁴ Solid solution photocatalysts have advantages over doped photocatalysts as solid solution allows for controlling the potentials of the conduction and valence bands, and the photogenerated electrons and holes are able to move in the continuous valence and conduction bands instead of the discrete donor levels seen in the doped photocatalysts. Zn is interesting as a semiconductor for combining with CdS. CdS and ZnS form a continuous series of solid solutions (Cd_1−xZn_xS) where metal atoms are mutually substituted in the same crystal lattice.^175–177 The Cd_1−xZn_xS solid-solution has a higher band gap than CdS, making the material attractive for solar applications. Photophysical and photocatalytic properties of Cd_1−xZn_xS solid solutions with different Zn concentration (0.2 < x < 0.35) prepared by co-precipitation were investigated.¹⁷⁸

X-ray diffraction patterns of Cd_1−xZn_xS samples (Fig. 13) displayed reflections corresponding to hexagonal CdS with higher hexagonal lattice parameters than those observed for pure CdS hexagonal phase, indicating the formation of solid Cd_1−xZn_xS solutions in all samples. The successive shift in lattice spacing indicated the higher substitution degree of Zn into CdS structure with the increase in Zn concentration. X-ray line broadening analysis revealed that the average crystalline size of Cd_1−xZn_xS solid solution varies with Zn concentration, with the highest crystalline size being obtained in the sample with a concentration of Zn equal to 0.3. UV–Vis absorption spectra of Cd_1−xZn_xS samples showed a blue shift of the absorption edge with the increase in Zn concentration. The energy band gap, determined from the adsorption onset, increases from 2.49 to 2.68 when Zn concentration in solid solution increases from 0.2 to 0.3. The changes in optical and structural characteristics of Cd_1−xZn_xS solid solutions have had an influence on their photocatalytic activities (Fig. 14). The photocatalytic activity of samples increases gradually when the Zn concentration increases from 0.2 to 0.3. The change in activity for H₂ production for these samples arises mainly from the modification of the energy level of the conduction band as the concentration of Zn increased. The activity of the solid solution decreases when the Zn concentration increases from 0.3 to 0.35 because it may be affected by a decrease in the number of available photons with the widening of the band gap and/or the decrease in the particle size of the solid solution.

Fig. 13 XRD patterns of Cd_1−xZn_xS with different Zn concentration (x = 0.2, 0.25, 0.30 and 0.35).

Fig. 14 Hydrogen evolution rate from aqueous solution containing Na₂S + Na₂SO₃ under visible light irradiation over Cd_1−xZn_xS with different Zn concentration (x = 0.2, 0.25, 0.30 and 0.35) (catalyst 0.1 g, reactant solution 150 ml (0.1M Na₂S + 0.04M Na₂SO₃, 150 W Xe lamp).

Taking into account the studies performed on CdS^165,179 that revealed the importance of the structural characteristics (crystalline phase, crystalline size and geometrical surface area) in the control of band structure and in the concentration and mobility of photocatalyst charges, studies have been conducted on the influence of structural changes induced by thermal treatments on the photophysical properties of Cd₁₋ xZn_xS solid solutions (x = 0.2).¹⁸⁰ The Cd_0.8Zn_0.2S solid solution subjected to thermal treatment under inert flow from 873 to 1023 K increases its crystalline size and slightly increases the substitution degree of Zn into CdS structure after heating at 1023 K. The structural changes induced by thermal treatments influence the UV–Vis absorption spectra of solid solution, with an increase being observed in the energy band gap of the solid solution (from 2.41 to 2.53 eV) with the increase in annealing temperature. The photocatalytic activity of Cd_0.8Zn_0.2S solid solution increases with treatment temperature, with a remarkable improvement in activity for the sample treated at 1023 K (Fig. 15). The combination of photoactivity results with structural characterization indicated that the improvement in photoactivity of the solid solution mainly arises from the increase in its crystalline size and absorbance derived from thermal treatments. The good crystallinity of phases with few crystal defects decreases the possibility of photoelectron–hole recombination, thereby leading to higher activity. In addition to these factors, the increase in the substitution degree of Zn into CdS structure observed for the sample treated at 1023 K is indicated as being responsible for the remarkable improvement in activity observed for this sample.

Fig. 15 Hydrogen evolution rate from aqueous solution containing Na₂S + Na₂SO₃ under visible light irradiation over Cd_0.8Zn_0.2S treated under inert flow at different temperature (873, 923, 973 and 1023 K) (catalyst: 0.1 g, reactant solution 150 ml (0.1M Na₂S + 0.04M Na₂SO₃, 150 W Xe lamp).

3.2.4.2 Other metal sulfide-based photocatalysts. ZnS is the other major metal sulfide investigated for photochemical water splitting. Combinations of AgInS₂ and ZnS produce a series of (AgIn)_xZn_2(1−x)S₂ solid solutions that crystallize in the cubic zinc blende or hexagonal wurtzite structure.¹⁴⁰ The optical adsorption of these materials can be adjusted between 400 and 800 nm depending on composition. (AgIn)_xZn_2(1−x)S₂ solid solutions showed photocatalytic activities for H₂ evolution from aqueous solutions containing sacrificial reagents, SO₃²⁻ and S²⁻, under visible light irradiation. Pt loaded (AgIn)_0.22Zn_1.56S₂ with a 2.3 eV band gap, showed the highest activity for H₂ evolution and the apparent quantum yield at 420 nm was 20%. Several ternary indium sulfides also show photochemical activity for water splitting under visible light. However these systems involve smaller quantities of H₂ under visible light. For example, a quantum efficiency of 3.7% at 420 nm was reported for the Na₁₄In₁₇Cu₃S₃₅ photocatalyst.¹⁸¹

Other metal sulfides (In₂S₃, SnS₂, HgS, Tl₂S, PdS, CoS, Fe₂S₃…) have also been tested without showing photoactivity under visible light because of their small band gaps (>2 eV).

4. Conclusions

The major source of the world's energy supply currently comes from fossil fuels, but they will decline in availability over the coming decades. The large-scale production of hydrogen from natural gas and other available hydrocarbons through catalytic reforming processes remains the cheapest source of hydrogen. Even when the cheapest production method is used (steam methane reforming), some authors estimate that the cost of hydrogen production is still four times that of gasoline for the equivalent amount of energy.¹⁸² However, H₂ production involving reforming technologies produces massive amounts of CO₂, which has an impact on global warming. One way to reduce CO₂ emissions is to apply reforming methods to alternative renewable precursors. Biomass precursors derived from plant crops, agricultural residues, woody biomass, etc., are being used for generating heat, electricity, and liquid transportation fuels (ethanol, sugars). Clean biomass precursors are converted into a gas mixture from which hydrogen is extracted. Virtually no net greenhouse gas emissions result because a natural cycle is maintained, in which carbon is extracted from the atmosphere during plant growth and is released during hydrogen production, although the time constants of the carbon cycle are different.

There is a need to develop non-conventional processes for H₂ production outside the C-cycle. Photonic energy technology offers the most promising outlook for the future because the water splitting reaction on semiconductor surfaces can generate potentially large quantities of clean, concentrated energy in H–H chemical bonds. Since the pioneering work by Fujishima and Honda,¹²⁴ numerous attempts have been made to develop efficient photocatalysts under visible light. Consequently, more than 130 photocatalytic materials have been described as catalysts for photochemical water splitting under visible light. In spite of these developments, current results still record low efficiencies for light-to-hydrogen conversion. So far, the maximum quantum efficiency over visible light-driven photocatalysts achieves only a few percent at wavelengths as long as 500 nm (Cr/Rh–GaN/ZnO, 2.5% QE pure water, visible light.¹⁵⁹ This value is still far from the QE of 30% marked as the initial starting point for practical application.¹⁴² Therefore, more efficient photocatalytic materials with a band gap as narrow as 2 eV (corresponding to 600 nm) need to be developed. This goal can quite possibly be achieved if sufficient knowledge is accumulated about the factors that determine the photoactivity of materials: molecular reaction mechanism, composition, structure, particle size, defect density, surface structure and charge transfer between semiconductor and cocatalysts. The molecular mechanisms and reaction kinetics of water reduction and oxidation on the semiconductor surface have yet to be elucidated in sufficient detail and should be investigated as a way of refining the materials to maximize efficiency. Finding new photocatalytic materials with a unique structure and phase is still likely to be the key to success. High throughput screening or combinational chemistry approaches, as well as a more rational search based on fundamental calculations/predictions, would be useful. The control of syntheses of materials for customising the crystallinity, electronic structure and morphology of catalysts at nanometric scale presents significant opportunities for improving water splitting photocatalysts, as these properties have a major impact on photoactivity. Taking into account the advances made in UV photocatalysts from the pioneering work of Fujisima and Honda in 1972 through to the present today, technically and economically viable visible light photocatalysts for water splitting could become available in the near future.

Acknowledgements

We are grateful to many of our colleagues for stimulating discussions and to our research sponsors CICyT and CAM (Spain) under grants ENE2007-07345-C03-01/ALT and S-0505/EN/0404, respectively. One of the authors (MCA) would also like to acknowledge financial support from the Ministry of Science and Education (Spain) through the R&C Program.

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