Ind. Eng. Chem. Res. 2009, 48, 3703–3712
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Hydrogen Generation from Chemical Hydrides Eyma Y. Marrero-Alfonso, Amy M. Beaird, Thomas A. Davis, and Michael A. Matthews* Department of Chemical Engineering, UniVersity of South Carolina, 301 Main Street, Columbia, South Carolina 29208
Complex chemical hydrides can be used to store and deliver hydrogen gas to fuel cells, and thus are one of several candidate materials to be used in storage systems for the hydrogen economy. These hydrides have high native hydrogen content, and hydrogen can be released via several chemical pathways. This review summarizes the extensive literature on the kinetic and thermodynamic properties of the various reactions of chemical hydrides, with an emphasis on hydrolysis. These properties are significant because they affect all aspects of system design, as well as the recovery and recycle of the byproducts. Hydrolysis of chemical hydrides takes place at relatively low temperatures and gives promising theoretical hydrogen storage efficiencies. Complications include metastable kinetic pathways as well as inefficient utilization of water in the byproducts. I. Introduction The term “hydrogen economy” has made its way into public consciousness in recent years. In general, the hydrogen economy refers to a vision of the future in which electric power will be generated not by direct combustion of fossil fuels or by nuclear fission, but by electrochemical oxidation of hydrogen in a fuel cell device. There are substantial opportunities to replace conventional power plants, whether on the grid system or in distributed power applications, with hydrogen-powered fuel cells. In addition, there is intense interest in finding replacements for fossil fuels in transportation applications. For transportation, hydrogen-fueled internal combustion engines (ICEs) are also an option,1 and so this approach to hydrogen utilization is also considered part of the hydrogen economy vision. Finally, a number of companies are investigating small fuel cells as an improvement over batteries in portable/mobile applications, such as laptop computers and mobile phones. Regardless of the application, hydrogen storage and delivery is a very challenging technical and economic issue to be resolved before the hydrogen economy can be realized on any national or global scale. Other key issues are low-cost and sustainable and efficient methods for hydrogen production and delivery2-4 as well as improving the efficiency and reliability and lowering the cost of the fuel cell itself.5 Research on these other key issues is being undertaken internationally, and a number of excellent science and technology reviews are available.6-8 While the production, storage, transmission, and utilization of hydrogen for petrochemical applications is mature and welldeveloped, hydrogen storage and generation for the hydrogen economy as defined above is still under development. There are several approaches to hydrogen storage and delivery presently under active research and development, and these have been described in both scientific9 and application-oriented reviews.10 Complex chemical hydrides are a means of storing hydrogen in the solid state at near ambient temperatures and pressures. * To whom correspondence should be addressed. E-mail: matthews@ cec.sc.edu.
Hydrogen can be released by a variety of chemical reactions for subsequent use in a fuel cell or ICE. The particular objective of this review is to summarize knowledge on the thermophysical and kinetic properties of chemical hydride reacting systems, with a focus on sodium tetrahydroborate (NaBH4, more commonly called sodium borohydride), which is the most-studied of the complex chemical hydrides. The thermophysical properties of chemical hydrides and the reaction products (e.g., density, crystal structure, water content, stability, and thermochemistry) are especially significant in determining the size of the storage and heat management equipment. The kinetic properties, along with the thermochemistry, ultimately dictate selection of catalysts and design of the reactor for hydrogen delivery to the power conversion device. Of course economics, safety, and infrastructure requirements are also important considerations for the hydrogen economy, but these issues have been summarized in other publications. Finally, it should be noted that this review is not intended to champion the adoption of chemical hydride technology over other approaches, such as compressed gas storage or reversible adsorptive storage. Indeed, the scenario that is already unfolding is that the full spectrum of hydrogen storage and delivery technologies will be utilized in the future hydrogen economy. II. H2 Storage System Technologies Four main approaches have been developed to store and deliver pure H2 gas at the point of use: compressed gas, cryogenic liquid, metal hydrides, and chemical hydrides. (A fifth approach is to store hydrocarbon or alcohol liquids, and to produce hydrogen by reformation or cracking. The reforming process generates CO, CO2, and other gaseous products at the point of use and is not further considered in this article.) Comparisons of these storage methods have been given in detail elsewhere.9,11 We briefly review the technologies to put chemical hydrogen storage in context of the current state of the art and to emphasize the importance of system design, not just material hydrogen capacity, for the success of any given hydrogen storage technology.
10.1021/ie8016225 CCC: $40.75 2009 American Chemical Society Published on Web 03/20/2009
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Figure 1. Volumetric and gravimetric capacities for several H2 storage technologies. Capacities for both materials-only basis and total system are shown.12,16
Two primary criteria used in the screening of hydrogen storage systems for vehicles are gravimetric and volumetric capacity.12 The gravimetric capacity (specific energy) is a measure of the usable energy per unit of system mass (e.g., kWh/kg). Volumetric capacity (energy density) is the useful energy per unit of system volume (e.g., kWh/L). Fuel cellpowered automobiles pose a significant challenge to hydrogen storage technology, and the U.S. Department of Energy through its FreedomCAR program has devoted considerable effort in evaluating current state-of-the-art technologies for the four main hydrogen storage routes. Figure 1, based on the U.S. Department of Energy (DOE) publications,12 presents the best known gravimetric and volumetric capacities for automotive transportation systems (ovals). To the DOE information we have added the native gravimetric materials-only efficiencies (stars) for gaseous and liquid hydrogen, and for NaBH4 hydrolysis. (There are many possible reversible hydride materials under consideration and, while the DOE reports a range of system efficiencies, we have not attempted to represent the materials-only storage capacity of any particular metal hydride). A comparison of the materials-only storage efficiency to the system efficiency reveals the penalties incurred by the tankage, valves, insulation, sensors, and other ancillary items necessary in a functioning system, regardless of the material basis. For reference, Figure 1 also shows the gravimetric and volumetric system storage efficiencies desired for passenger cars in the year 2015. This particular target, established by the U.S. Department of Energy FreedomCAR program,12 is intended to stimulate breakthroughs in both material storage capacity as well as efficient system design. The thermodynamic and kinetic properties of the native storage material influence system design; the penalties are different for each storage material. For example, pure liquefied hydrogen, at about 20 K and atmospheric pressure, has a volumetric storage capacity of 2.35 kWh/L13 and a gravimetric capacity of 33.3 kWh/kg11,13,14 (LHV). By comparison, pure gasoline has a storage efficiency equivalent of about 8.88 kWh/ L13 and 11.8 kWh/kg,11 and other liquid hydrocarbons are also extremely efficient even compared to pure liquid hydrogen. Whereas liquid hydrocarbon fuels can be stored and delivered from lightweight tanks near ambient conditions, pure H2 requires either high pressure tanks or highly insulated cryogenic storage. These system requirements lower the gravimetric and volumetric capacities significantly, as shown in Figure 1.
Hydrogen poses a unique challenge for storage because it is the lightest and least dense gas at only 0.0898 kg/m3 at standard temperature and pressure (STP).7 Compressed hydrogen gas is the most developed of the storage technologies and has already been used in early prototype fuel cell vehicles. The method involves pressurizing H2 to 5000-10000 psi (35-70 MPa) in specialized reinforced tanks. Despite increasing the density to 30 kg/m3 for 10000 psi gas tanks, the volumetric capacity is the lowest of the technologies, making it inconvenient for portable or automotive applications.15 Cryogenic liquid hydrogen further increases the density to 70.8 kg/m3 by condensing the hydrogen at 21.2 K.11 However, the liquefaction process requires large amounts of energy, approximately 30% of the overall energy content of the H2 to be liquefied.12 A combination of the two techniques called cryo-compressed hydrogen11 is also currently being investigated. Solid state hydrogen storage, encompassing both reversible metal hydride and irreversible chemical hydride approaches, are of particular interest because several of the materials offer gravimetric and volumetric capacities in excess of the DOE targets, leaving room for incorporation of system components without falling below the targets. Recent detailed reviews16-19 on metal hydride advances are available. Reversible metal hydrides comprise a broad range of materials including conventional/interstitial (e.g., LaNi5H6, Mg2NiH4) and complex hydrides (alanates, borohydrides, amides).20 Generally, hydrogen is stored by chemisorption of hydrogen onto the metal at elevated pressures. Hydrogen is then released by a combination of lowering pressure and increasing temperature.18 Because many of the materials allow reversibility (cycling) of this process, on-board H2 refueling is possible, which is potentially a major advantage especially for automobiles. A variety of reversible hydrides have been considered for vehicles by the DOE Metal Hydride Center of Excellence (MHCoE) and research is ongoing.21 Their potential is primarily based on their gravimetric density and reversibility. Each of the classes of metal hydrides has intrinsic pros and cons. For example, conventional hydrides are highly reversible and have high volumetric energy densities, but low gravimetric capacity.18 Complex hydrides of light elements such as Mg(BH4)2, Al(BH4)3, and AlH3 have high theoretical hydrogen capacities (10-17 wt %) but their reversibility is limited at reasonable temperatures and pressures by kinetics and/or thermodynamics.16 The metal hydride automotive system efficiencies shown in Figure 1 are projections from the Department of Energy based on a very limited number of full-scale systems or laboratory subscale prototypes.22 Reversible metal hydrides have the advantage of delivering pure hydrogen gas. System components typically include heat exchangers for heating (during hydrogen release) and cooling (during rehydrogenation), as well as a pressure vessel for containment (because rehydrogenation is done under pressure). III. Chemical Hydrides A current chemical hydride system evaluated by DOE in Figure 1 is based on aqueous-phase, catalyzed hydrolysis of NaBH4. Other pathways for reacting and utilizing chemical hydrides will be given subsequently. The materials-only H2 storage capacities shown for NaBH4 (stars) are based on the reaction below (eq 1) and include the mass and volume of both hydride and water: NaBH4(s) + (2 + x)H2O(g) f 4H2(g) + NaBO2·xH2O(s) + heat (1)
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Ideally, only two moles of water are required to release four moles of H2, but in fact (as will be discussed in detail later) the product sodium metaborate can exist in several degrees of hydration, with the generic stoichiometry x. On a materialsonly basis, this reaction gives 10.9 kg H2/kg reactants if x ) 0. If x is 2 or even 4 the materials-only storage capacity is still promising for some applications. However, in practice to date the best systems require that the NaBH4 be predissolved in a large excess of water to keep the solids in solution. The excess water, as well as the mass of the system hardware, significantly reduces the storage efficiency as shown. Details of stoichiometry, solubility, and other basic properties for additional chemical hydrides will be given below. Additional complex chemical hydrides such as NaAlH4, LiBH4, and LiAlH4 have good stoichiometric potential for gravimetric and volumetric efficiency. However, nearly ideal hydrolysis must occur with the minimum required water and 100% conversion of the hydride. Only under these conditions will there be leeway for component weight and volume to be designed to meet application-specific system requirements. IV. Stoichiometry of Hydrolysis and the Role of Water Chemical hydrides contain hydrogen in a reduced or electronrich state.23 Hydrides can be categorized as simple binary compounds or complex hydrides. In the simple hydrides (MHn, for example CaH2) the negative hydrogen is bonded covalently or ionically to a metal. In covalent hydrides the hydrogen-metal bond is created by a common electron pair. The ionic hydrides contain metal cations and negatively charged hydrogen ions. Simple hydrides react with water according to MHn+nH2O ) M(OH)n+nH2
(2)
where M is a metal of valence n. In simple chemical hydrides, the reaction product at near ambient conditions is always the hydroxide, not the oxide. Thus simple hydrides inherently have some inefficiency in that they retain n of the H atoms available in eq 2 as hydroxide. If the stable reaction product were the oxide, the gravimetric efficiency of hydrogen production would increase. In complex hydrides the hydrogen is combined with two other components, usually metallic elements. The complex hydrides (M(M′H4)n) react according to M(M′H4)n + (2n + x)H2O ) M(M′O2)n·xH2O + 4nH2 (3)
Table 1. Hydrogen Storage Capacity (H2 wt %) for Hydrolysis Products with Varying Degrees of Hydration hydration state of oxide product hydride
x)0
x ) 0.5
x)2
x)4
NaBH4 LiAlH4 LiBH4 NaAlH4
10.92% 10.90% 13.95% 8.96%
9.73% 9.72% 12.07% 8.14%
7.34% 7.33% 8.59% 6.40%
5.53% 5.52% 6.21% 4.97%
Table 2. Comparison of Gravimetric Efficiencies of Various Reactant Preparations
reactant disposition
molar ratio NaBH4:H2O
x
H2 storage capacity (wt %)
ideal hydrolysis (x ) 0, eq 1) saturated NaBH4 solution (25 °C) NaBO2 solubility limit (25 °C) 8.5 wt % NaBH4 + 5 wt % NaOH29
1:2 1:3.81 1:13.04 1:21.37
0 1.81 11.04 19.37
10.92% 7.57% 2.96% 1.81%
theoretical gravimetric efficiency decreases with increasing x for several complex chemical hydrides. Most of the systems that have been engineered to date utilize a route wherein both the hydride reactant, as well as the product, are dissolved in excess water. Because the solubility of the chemical hydride (55 g NaBH4/100 g H2O at 25 °C)27 is relatively low, and the solubility of the oxides is even lower (28 g NaBO2/100 g H2O at 25 °C),28 such systems inherently have a reduced gravimetric efficiency compared to the ideal. Table 2 shows the effective value of x for different amounts of water, and the corresponding hydrogen storage capacity on a materials-only basis. Ideal hydrolysis requires no excess water. Complete dissolution of NaBH4 at 25 °C requires extra water corresponding to x ) 1.81. If sufficient water is added to keep the NaBO2 in solution, the effective value of x is 11.04. Furthermore, the addition of NaOH or other base to stabilize the solutions also affects the hydrogen storage density. Shang et al.29 used thermodynamic modeling to find the optimum NaBH4 concentration as a function of alkali content and temperature. On the basis of their findings,29 the optimum concentration is 8.5 wt % NaBH4/5 wt % NaOH resulting in a 1.81 wt % H2 (x ) 19.37). Clearly, addition of water to achieve solubility of the reactants and/or products decreases the storage efficiency significantly. Several different hydrides have been considered for hydrogen storage and release via the hydrolysis route.20,30 Figure 2 shows the maximum hydrogen yield possible, assuming ideal hydrolysis (x ) 0) and 100% conversion of the hydride. LiBH4, NaBH4,
where n is the valence of the metal M, M′ is an element from group IIIA of the periodic table, such as boron or aluminum, and x is structural water of crystallization. Equation 3 is written with the assumption that the reaction has been carried out under conditions that produce a dry solid and dry gas; in this case the solid phase product is best described as a hydrated mixed metal oxide. In many practical designs, the reaction is carried out in excess water, in which case the (M′H4) anion and the metal oxide product are in the hydrated form in solution:24,25 (M′H4)- + 4H2O f M′(OH)4 + 4H2
(4)
In the context of hydrogen storage efficiency, Marrero et al.26 has also interpreted x as the “excess hydration factor”. Because the most efficient hydrolysis stoichiometry would lead to the anhydrous oxide, x ) 0, any water in excess reduces the gravimetric efficiency proportionally. Table 1 shows how the
Figure 2. Maximum deliverable hydrogen from hydrolysis or thermolysis of ionic hydrides.
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and NaAlH4, LiAlH4, which are the most commercially important complex hydrides, are high in hydrogen content. Hydrolysis is typically conducted in the aqueous phase at lower temperatures, and large quantities of excess water are required because of the low solubility of both NaBH4 and the borate byproducts in water.29 There are additional considerations that limit the kinetics and yield of hydrogen production. In pure liquid water the reaction generates species that are kinetically metastable depending on pH. The reaction slows or ceases at low yields of hydrogen even though the reaction is strongly favored thermodynamically.24,31-33 The metastable species require the addition of mineral acid or use of a heterogeneous catalyst to overcome kinetic limitations that exist in the aqueous phase at low temperatures.34 Furthermore, the solid products of the reaction are an unknown mixture of hydrated sodium metaborates. Hydrated products capture water that is not reduced to hydrogen, further decreasing the efficiency of the reaction. V. Other Pathways for Utilizing Chemical Hydrides In addition to the hydrolysis chemistry shown above, there are other chemical reactions involving complex hydrides that have been investigated,35 such as thermolysis and more recently, a combustion-assisted hydrolysis approach that combines thermolysis with hydrolysis. However, to date, none of these approaches has been implemented into a practical hydrogen generation system. In thermolysis (referred to as pyrolysis by some authors) the substance decomposes upon heating to generate hydrogen. Most simple hydrides decompose via a reversible thermolysis reaction35 shown below: 2MHn + heat S 2M + nH2
(5)
where M is a metal or an alloy. This reaction can be reversed by hydriding the metal under high pressure. Complex hydrides also decompose to produce hydrogen, but it has been reported35 that reversibility is not always possible. Decomposition and reversibility are possible for complex hydrides such as LiAlH4 and NaAlH4, making these compounds the most feasible for hydrogen storage. Thermolysis reactions use a stable starting material; no reactant is needed other than the hydride. However, in some examples the hydride has been mixed with a binder or other inert components to form a consolidated pellet, and the mass of these components reduce the net hydrogen storage efficiency. Thermolysis offers high energy density compared to hydrolysis routes. Disadvantages of this approach are the high temperature necessary to initiate reaction and the high sensible heat necessary to bring the reactants to the reaction temperature. Also, the thermolysis itself is highly exothermic and difficult to control.35,36 High temperatures and reaction control are safety concerns for a hydrogen delivery system. Figure 2 compares the ideal hydrolysis (x ) 0) and thermolysis (i.e., formula hydrogen content) routes for releasing hydrogen from several chemical hydrides. For all compounds except LiBH4 and LiH, ideal hydrolysis has a greater gravimetric H2 storage potential, because hydrolysis also generates hydrogen by reduction of water. For LiBH4 and LiH, ideal thermolysis is potentially more efficient because of the very low atomic mass of lithium. In the case of thermolysis, complete decomposition is assumed and Figure 2 does not account for any binder or inert material that might be used to control the temperatures. For the hydrolysis case, complete reaction is assumed using stoichiometric amounts of water and no water of hydration (x ) 0).
Hydrogen can also be released from complex hydrides via Beckert-Dengel chemistry, in which an ammonium salt reacts with a chemical hydride to produce hydrogen and a metal nitride. m/n(NH4)nX + Y(ZH4)m f YXm/nN + 4mH2
(6)
This reaction occurs according to the stoichiometry in eq 6 where X is a halogen/sulfate, Y is an alkali/alkaline earth metal, and Z is a trivalent metal.37 These are exothermic reactions that, like thermolysis, require high initiation temperatures and a binder to consolidate the hydride so that the reaction will propagate and attain complete conversion. The Beckert-Dengel pathway is also difficult to control and the substantial heat of reaction must be dissipated. However, all of the reactions take place between solids, and the only gas released is hydrogen. An interesting combination of thermolysis and hydrolysis has recently been demonstrated by Varma et al.38 In their approach, an elemental metal such as aluminum or magnesium was suspended in a polyacrylamide gel and added to an alkaline chemical hydride solution. When ignited, these mixtures release hydrogen by both hydrolysis of the chemical hydride, and oxidation of the elemental metal. This approach resulted in higher yield of hydrogen than would be possible with either pathway alone and did not require an additional catalyst. The temperatures required for this process are in excess of 1000 °C and performed best with nanoscale metal powders. One of the main byproducts of this pathway is methane.39 Therefore, condensation of the unwanted gaseous products is required to purify the H2 upstream of the fuel cell. Rather than generating hydrogen and then consuming it in a fuel cell or ICE, sodium borohydride can also be oxidized directly in the so-called direct borohydride fuel cell (DBFC). This is a relatively new approach with advantages and disadvantages related to thermophysical and kinetic properties of the electrochemical reaction. Several comprehensive reviews have recently been published on the topic.40-43 In general, this approach may avoid the need to handle gaseous H2 by oxidizing alkali-stabilized NaBH4 solutions at the anode to, ideally, release eight electrons (eq 7).44 When paired with oxygen reduction at the cathode (eq 8), the overall cell reaction is as shown in eq 9.43 anode:
BH4 + 8OH f B(OH)4 + 4H2O + 8e
cathode: 2O2 + 4H2O + 8e- f 8OHoverall:
BH4 + 2O2 f B(OH)4
(7) (8) (9)
The DBFC could give greater energy efficiency than the indirect method.42 Likewise, non-noble metals may be employed for the anode catalyst resulting in a cost savings over PEM fuel cells. One of the main setbacks in this approach is that the anode catalyst materials that are active for oxidation are also active toward hydrolysis of the NaBH4, resulting in an undesired release of H2.41 Furthermore, current membrane materials are not resistant to the highly alkaline conditions required and allow crossover of BH4- to the cathode. This is a significant issue for DBFCs as it leads to reduced efficiency, possible fouling of the cathode catalyst and safety concerns for handling the evolved hydrogen gas.41 However, significant progress has been made on specialized anodic catalysts, membrane materials, and other optimizations in a relatively short time frame. VI. Thermophysical Properties of Hydrated Byproducts The hydrolysis of NaBH4 is shown in eq 1. On the basis of the chemical equation, the amount of hydrogen produced per
Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3707 Table 3. Thermal Data for Selected Compounds Relevant to Chemical Hydrides compound (data source)
heat of formation (25 °C)
Gibbs free energy (25 °C)
heat capacity (25 °C)
NaBH4a LiBH4a LiAlH4a NaAlH4a H2O (g)a H2O (l)a H2O structuralb H2a Na+ (aq)a Li+ (aq)a B(OH)4-b NaBO2a H3BO3(c)a HBO2(c)a NaOHa B4Na2O7a
∆H°f (kJ/mol) -188.6 -190.8 -116.3 -115.5 -241.826 -285.83 -290.42 0 -240.1 -278.5 -1345.5 -977 -1094.3 -794.3 -425.8 -3291.1
∆G°f (kJ/mol) -3.9 -5 -44.7 ? -228.6 -237.18 -237.18 0 -261.9 -293.3 -1159.9 -920.7 -968.9 -723.4 -379.7 -3096
Cp (J mol-1 K-1) 86.8 82.6 83.2 ? 33.6 75.3 ? 28.8 46.4 68.6 ? 65.9 86.1 ? 59.5 186.8
a
CRC handbook. b Reference 49.
unit of chemical hydride is independent of the excess water in the reaction because the excess appears as hydration of the borate. Different studies26,45 have observed that the hydration state of the borate product depends on the conditions under which the reaction occurs. Species that have been detected include NaBO2 · 2H2O,26 NaBO2 · 4H2O, and Na2B4O7 · 5H2O.45 Matthews et al.26 determined that the NaBO2 product of steam hydrolysis at 110 °C is the dihydrate (NaBO2 · 2H2O). Steam/ solid reactions were conducted with various ratios of water to hydride, but regardless of the ratio the preferred product was sodium metaborate dihydrate. Studies on the stability of borates indicate that sodium borate dihydrate is stable at room temperature, while at temperatures above 115 °C the hemihydrate (x ) 0.5) is the stable form.47,48 More information about the reaction equilibrium is necessary to predict the preferred product. The Gibbs free energy and heat of formation determine whether the reaction is thermodynamically favored and how much energy must be added or removed during the reaction. Thermodynamic properties of the compounds involved in the reaction are necessary to estimate the heat of the reaction. There is limited information in the literature on the thermal properties (∆G°, f ∆H°) f of sodium borates in their different degrees of hydration. It is necessary to use correlations to estimate the Gibbs free energy and the heat of formation. Li et al.49 developed general equations (eq 10 and 11), based on the group contribution method, to correlate and predict the thermodynamic properties of hydrated borates based on structure. The Gibbs free energy of formation (∆G°) f and the enthalpy of formation (∆H°) f of a hydrated borate phase (Ma[BxOy(OH)z] · nH2O) can be estimated by considering the contributions of the cations in aqueous solution (M+v(aq)), the borate polyanions (BxOy(OH)z) and the structural water. Properties of the cations, polyanions, and water molecules are well-known and can be found in the literature. Some of the data for the calculations related to different hydrides are presented in Table 3. ∆Hf◦(borate) ) a∆Hf◦(M+v(aq)) + Hf◦(BxOy(OH)z) + nHf0(H2O) (10) ∆Gf◦(borate) ) a∆Gf◦(M+v(aq)) + Gf◦(BxOy(OH)z) + n∆Gf◦(H2O) (11) Equations 10 and 11 can be easily evaluated from the structure and stoichiometry (a and n) for a given compound. Each group
is easy to identify when the chemical formula is written in the following specific form: Ma[BxOy(OH)z] · nH2O. Using equations 10 and 11 and the properties obtained from literature (Table 3), the heat of formation and Gibbs free energy can be calculated for borates with different degrees of hydration. Marrero26 used this method to calculate the Gibbs free energy and the heat of formation of different hydrated sodium borates. It was found that the Gibbs free energy and the heat of formation decrease with the increase of the excess hydration factor. Thus the formation of more highly hydrated borates is more thermodynamically favored, which would reduce the gravimetric efficiency. From the heats of formation, the heats of reaction were also estimated.26 The reaction becomes more exothermic with increasing excess hydration factor. From the Gibbs free energy of reaction, hydrolysis is thermodynamically favored and is very exothermic. The removal of heat is a challenge for system design because of the mass and volume of heat exchange equipment, as well as safety. VII. Kinetics of Chemical Hydride Hydrolysis The hydrolysis of chemical hydrides has been a significant area of research since the mid 20th century. Initially, hydrolysis was investigated as a means to generate H2 for weather balloons and military applications during World War II.50 In addition, an interest in the reducing properties of sodium borohydride for organic and inorganic syntheses prompted studies on the solubility and reactivity of NaBH4 in various solvents.51,52 A variety of fundamental thermodynamic properties were also reported.53-57 Renewed interest in compact hydrogen storage in the late 1990s prompted new investigations based on Schlesinger’s pioneering work.34 Many authors have investigated the effects of various catalysts, temperature, pH, and NaBH4 concentration on reaction kinetics.32,34 Researchers have optimized these independent variables in an attempt to improve the gravimetric and volumetric properties described earlier. Complete conversions are achieved only by adjusting the solution pH to be acidic, by using a heterogeneous catalyst or by hydrolysis with steam.32,34,58 Most recent work has focused on developing more efficient catalysts for the reaction.28,59-73 Furthermore, mechanisms for the various acid additives and metal catalysts have been proposed, although the pathways are not completely understood. Understanding and controlling the release of hydrogen during the hydrolysis of NaBH4 is of crucial importance to the success of the technology. The reaction is known to be pH sensitive, reacting to completion under acidic conditions and stabilizing at low H2 yields under strongly basic conditions. Because of this, acids have been used as catalysts and bases are utilized as stabilizers to prevent premature reaction in some reactor designs. To promote reaction in stabilized basic solutions, metal catalysts are often employed and result in fast kinetics and high hydrogen yields that vary with catalyst preparation. Finally, the reaction rate also varies with temperature and interestingly, with the phase of the water used, for example steam. i. Uncatalyzed Hydrolysis. Almost all research to date is based on conducting the reaction in an aqueous mixture, with large amounts of excess water. Despite favorable thermodynamics, aqueous phase reactions with pure water give conversions of less than 10%, because the reaction mixture becomes basic and the reaction intermediates are stabilized at elevated pH. This stabilization under high pH conditions has also been exploited to prevent premature reaction. Fundamental investigations of the hydrolysis of chemical hydrides, particularly lithium and sodium borohydrides, were
3708 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009
performed by Schlesinger et al.34 in the early 1950s. They observed that the rate of the aqueous hydrolysis reaction decreased with the formation of the basic sodium metaborate B(OH)4- product and concluded that the reaction rate is dependent on both temperature and pH. Kreevoy et al.74 described the stability of NaBH4 in aqueous solution by an empirical rate law given in eq 12, where t1/2 is the half-life of NaBH4 in minutes and T is temperature in Kelvin. log10(t1/2) ) pH - (0.034T - 1.92)
(12)
This rate equation describes reaction in dilute solutions of NaBH4 with large concentrations of NaOH (up to 40% by weight) which prevents the liberation of hydrogen during storage. The Kreevoy equation describes kinetics in a region that is too dilute to be practical for hydrogen generating systems. Recently Moon et al.75 investigated the pH and temperature effects on the rate of decomposition of more concentrated NaBH4 solutions. As previous studies have shown, it was found that the uncatalyzed aqueous hydrolysis reaction of NaBH4 proceeds very slowly. Solutions ranging from 5 wt % NaBH4 to 25 wt % NaBH4 were studied, and they found that the rate of hydrogen evolution increases with increasing concentration up to a maximum of 20 wt % NaBH4 and then diminishes at higher concentrations. With more concentrated solutions, they assert that the rate was diminished probably due to solubility and mobility inhibition by the NaBO2 product and lower availability of water.75 For all solutions, initially linear (zeroorder) hydrogen generation was observed, but nonzero-order behavior became evident after several hours. It was found that the addition of alkali stabilizers such as NaBO2, KOH, and NaOH depressed the hydrogen evolution, a property utilized in many practical chemical hydride systems. The maximum H2 yield, that is the total amount of H2 released before stabilization occurs, also depends heavily on the pH and temperature. Yield is also conveniently expressed as percentage conversion, that is, the percentage of the maximum possible amount of hydrogen that is actually observed experimentally. Kojima et al.28 reported that at room temperature (20-23 °C) the overall reaction conversion stabilizes at 7% as the basic borate byproduct is formed. To create a basic environment in which NaBH4 hydrolysis is stifled, simple hydroxides such as NaOH and KOH are added. However, the resulting alkaline solution continues to undergo slow reaction to release H2. Several studies have looked at the decomposition rate of these solutions at various NaOH and NaBH4 concentrations and temperatures to gauge their shelf life. It is apparent that the rate of hydrolysis for dilute solutions is much different than that of concentrated solutions and that temperature has a dramatic effect on the stability of the solutions. The concentration of NaBH4 in solution must be maximized to improve energy density but low enough to keep the NaBO2 byproduct in solution.29 Minkina et al. investigated the effects of temperature and NaOH concentration on the degradation of concentrated NaBH4 solutions.76 They suggest that above 30 °C, at least 5 wt % NaOH must be added to decelerate the hydrolysis, consistent with Shang’s suggested 8.5 wt % NaBH4/5 wt % NaOH solution at room temperature.29 To release the hydrogen from these base-stabilized solutions, metal or acid catalysts are required. More applied studies have been conducted to minimize the amount of water in the reaction. In 1999, two separate papers regarding the interaction of NaBH4 with water vapor were published. Kong et al.30 exposed sodium borohydride to humid air at 25 °C and 1 atm. The hygroscopic NaBH4 powder
absorbed water and formed a solution, but no reaction was observed and the study was not pursued further. In contrast, Matthews et al.58,78 contacted sodium borohydride powder with saturated steam, varying the temperature from 110 to 140 °C. They observed high reaction rates and conversions at 110 °C, without a catalyst. However, at 140 °C the rate and overall yield diminished significantly. From a kinetic standpoint, the lower rate at higher temperature is unexpected. They also observed agglomerated products, which were hard and dry at 140 °C. Therefore, Matthews et al. postulated that the formation of an impermeable byproduct shell during reaction at high temperatures limits the mass transfer of steam to the reactants and reduces the yield. The reaction product was soft and wet when the hydrolysis was performed at 110 °C with pure steam. ii. Acid and Metal Catalysis. To improve rates and yields of the reaction for hydrogen generation, Schlesinger34 investigated two catalytic approaches. First, acid additives were used to overcome pH stabilization and allow the reaction to proceed to completion. This technique requires large amounts of acid, making the method unsafe, heavy and bulky. He also discovered that adding a few weight percent of certain metal salts, such as cobalt, nickel, and platinum, has a catalytic effect, independent of the pH of the solution. This improves the reaction rate without adding much additional weight. However, these catalysts were reported to cause foaming during hydrogen generation, limiting their application. Addition of acid to aqueous NaBH4 solutions causes the hydrolysis to proceed at an appreciable rate until complete conversion is achieved. Several studies have reported that the rate is independent of the specific acid used and depends only on the pH, referred to as general acid catalysis. Wang and Jolly24 proposed that borohydride hydrolysis is subject to an acid-base equilibrium between borate and boric acid. Under strongly basic conditions (pH > 9) the hydrolysis reaction (eq 13) yields the tetrahydroxyborate anion B(OH)4- which precipitates from saturated solution at room temperature as NaB(OH)4, also written NaBO2 · 2H2O. Under more acidic conditions (pH < 9) the hydrolysis yields boric acid (eq 14). BH4- + 4H2O f B(OH)4- + 4H2
pH > 9 (13)
+ BH4 +H + 3H2O f B(OH)3 + 4H2
pH < 9 (14)
The overall equilibrium was represented by Corti et al.79 as shown in Equation 15. B(OH)3+OH- f B(OH)4
(15)
While several mechanisms and intermediates have been proposed, the two most well regarded both assert that the BH4ion undergoes a stepwise hydrolysis, with the major discrepancy being the identification of the intermediate. The original mechanism, proposed by Davis and Swain80 and supported by Gardiner81 suggests a stepwise mechanism involving the intermediate BH3OH-. A later mechanism, proposed by Kreevoy and Hutchins33 accounts for kinetic data under strongly basic conditions, which the previous mechanism did not address. They suggest that rather than the BH3OH- intermediate, H2BH3 forms and reacts according to eqs 16-20. H+ + BH4 S H2BH3
(16)
Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009 3709
H2O +
BH4
-
S H2BH3 + OH
(17)
H2BH3 ⇒ H2 + BH3
(18)
BH3 + 3H2O ⇒ 3H2 + B(OH)3
(19)
OH- + B(OH)3 S B(OH)4
(20)
Furthermore, they proposed that the formation of the H2BH3 intermediate is the rate-limiting step in acidic solution. In general, the reaction is modeled as first order kinetics in NaBH4 and H+ across a range of pH, but does not take water into account in the rate equation. A new and valuable summary of the complex hydrolysis reaction chemistries (Figure 3) was developed by D’Ulivo et al.82 Their motivation was not hydrogen production, but the generation of boranes. This figure depicts the effect of pH and equilibriums of the intermediate species involved in the reaction. Catalysis by metals such as cobalt, nickel, copper, palladium, ruthenium, and their salts have also been widely investigated and found to follow zero-order kinetics32 and the rate is independent of the pH of the solution. Metal catalysts are also advantageous over acids, because they can be recovered and reused and can function in an on/off manner.83 The subject of low-temperature kinetics and metal catalysis in the aqueous phase has been reviewed recently by Liu et al.84 Some more applied studies attempted to reduce the kinetic limitations and the overall amount of water needed. The approach was to react solid NaBH4 by adding liquid water rather than predissolving, thereby increasing the gravimetric efficiency.45,75 Liu et al.45found that the reaction could achieve fast kinetics and more than 90% conversion. They observed that a ratio of water to borohydride of more than 4:1 was necessary in order to obtain high yields. This amount of water corresponds to an excess hydration factor (x) of 2. In this study cobalt chloride was found to be more effective than cobalt powder due to the initial acidic catalysis mechanism. The final solid product was analyzed by X-ray diffraction, and NaBO2 · 2H2O, NaBO2 · 4H2O, Na2B4O7 · 5H2O, and Na2CO3 · H2O were detected. VIII. Recycling and Production The hydrolysis reaction for production of hydrogen from sodium borohydride is nonreversible. Sodium borohydride utilization on a large scale will require a method of recycling
and reconstituting the solid byproducts that is less energy intensive than the current methods. The current state of chemical hydride manufacture is reviewed here. Sarkar and Banerjee85 have recently published a net energy analysis for the other three options of hydrogen storage, namely compressed gas, cryogenic liquid, and metal hydrides. Sodium borohydride production is currently the largest consumer of metallic sodium in the U.S., accounting for an estimated 30-35% of domestic sodium demand.86 Sodium borohydride is widely used in wood pulp bleaching, brightening recycled newsprint, and as a reducing agent in organic synthesis. Other applications include the control of odors in alcohols, electroless plating of nickel, gold, and other metals for electronic and aerospace applications, decolorization of ketones and the removal of toxic or heavy metals from process waste streams.87,88 The industrial production of sodium borohydride is based on the Schlesinger process,89-91 (eq 21) also known as the Rohm & Haas process.92 Sodium hydride (produced from the reaction of sodium with hydrogen) and trimethyl borate undergo a rapid reaction at 225-275 °C to produce sodium borohydride of high purity (90-96%) and high yields (94%). 4NaH + B(OCH3)3 f NaBH4 + 3NaOCH3
(21)
In the commercial process, trimethyl borate is added to sodium hydride dispersed in hydrocarbon oil at 250 °C. The components in the hydrocarbon oil phase are separated with addition of water, and sodium methylate (NaOCH3) hydrolyzes to methanol and sodium hydroxide. The methanol is recovered by distillation and utilized to make trimethyl borate. Sodium borohydride is in the remaining alkaline aqueous solution with a typical concentration of 12% NaBH4 and 40% NaOH. The solid sodium borohydride is recovered by extraction with isopropylamine and precipitation.93 Sodium borohydride is relatively expensive due to the multiple steps in the Brown-Schlesinger process and the use of an expensive reducing agent.94 The present cost of sodium borohydride is approximately $55/kg.95 The production of sodium borohydride is a highly energetic process; this represents a significant problem if sodium borohydride is to be used as an energy carrier on a large scale. Large processing plants will reduce production cost. Different alternatives94-102 for the synthesis of sodium borohydride have been recently published. Sodium borohydride has also been synthesized on a commercial scale by Bayer, by the reaction described in eq 22.94,102 This synthesis occurs in a batch process where a mixture of dehydrated borax, quartz, and sodium metal under hydrogen gas is heated in the presence of silica. 1 7 7 4Na + 2H2 + Na2B4O7 + SiO2 f NaBH4 + Na2SiO3 4 4 4 (22)
Figure 3. Reaction pathways for sodium borohydride hydrolysis across a range of pH. Reprinted with permission from ref 82. Copyright 2004 Elsevier.
This process has some complications that make it inappropriate for commercial production. There is some build-up of sodium silicate that requires disposal. The reaction has some explosion risks because it operates above the decomposition temperature of sodium borohydride. Also, the batch process is not ideal for economical commercial production. High production rates obtained from continuous flow processes are a key factor in the reduction of the final product cost. Previous work89 in the area of sodium borohydride production has been done with high temperature reactions. Suda et al.96 proposed a mechano-chemical reaction at room temperature by ball milling the reactants, a mixture of magnesium hydride, dehydrated borax, and sodium. Kojima et al.101 proposed a
3710 Ind. Eng. Chem. Res., Vol. 48, No. 8, 2009
process to recycle NaBO2 back to NaBH4 by annealing NaBO2 and MgH2 or NaBO2 and MgSi under high H2 pressure and high temperatures (350-750 °C). Kojima et al. also described the process of NaBH4 synthesis from coke, water, and NaBO2. More recently Suda et al.95 studied with more detail the mechanism of the anhydrous sodium metaborate with magnesium and hydrogen to form sodium borohydride. Suda et al.97 also explored the use of aluminum for the borohydride synthesis. It was found that NaBH4 cannot be formed through reaction of NaBO2 with Al and H2 directly, but Na4B2O5 reacts with Al and H2 to form NaBH4. The addition of NaBO2 to this reaction helps to achieve higher NaBH4 yields. Millennium Cell Inc.100 examined the electrolytic reduction of boron oxide and hydrogen gas leading to borohydride. It was concluded that there is potential for this technology to lower the costs of sodium borohydride production. The process is carried out in a batch system with no intermediates additions or separations. Park et al.101 also investigated electrolytic reduction at milder temperatures and pressures. They were able to convert the sodium metaborate to borax hydrate, which is more reactive toward NaBH4 production. They suggest that the combination of an electrochemical step, with a thermal treatment step can result in a more energy efficient recycle scheme. The borate byproduct of the hydrolysis of sodium borohydride potentially can be separated from the sodium borohydride solution, thereby increasing the gravimetric H2 storage density of the solution and facilitating the recycling of NaBO2 for production of NaBH4. Commercial nanofiltration membranes have been tested at different operating conditions to separate sodium borate from sodium borohydride.99 The selective separation of borates using membranes could prove useful in the production of H2 from solutions that contain more than 16 wt % of NaBH4. Researchers keep looking for an efficient approach that will decrease the cost and energy consumption of the recycling of NaBO2, such as the conversion of NaBO2 to borax. A more energy efficient process is possible if borax is used as the starting material for NaBH4 production. This will close the “hydrogen cycle”92 that begins with the production of sodium borohydride from borax, then the dehydrogenation of sodium borohydride that produces NaBO2 and finally the recycling of NaBO2 back to NaBH4. IX. Fundamental Properties and Relevance to Application Applications for hydrogen-powered fuel cells range from microwatts to megawatts. Regardless of the energy conversion device, the demands on the hydrogen storage system are severe. The Darnell Group102 reviewed markets for fuel cells for commercial portable power (
