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Synthesis of metal hydride from water
International Journal of Hydrogen Energy 32 (2007) 2501 – 2503 www.elsevier.com/locate/ijhydene
Synthesis of metal hydride from water Surendra K. Saxena ∗ , Vadym Drozd, Andriy Durygin Center for the Study of Matter at Extreme Conditions, College of Engineering, Florida International University, VH-140, University Park, Miami, FL 33199, USA Received 1 September 2006; received in revised form 18 September 2006; accepted 18 September 2006 Available online 8 December 2006
Abstract A number of hydrides are considered good candidates for hydrogen storage material for various applications in particular for automobile use. A metal hydride is synthesized through the reaction of a metal with hydrogen which is formed on industrial scale either by the electrolysis of water, by heating coke with steam in the water gas shift reaction or using hydrocarbons with steam. This study demonstrates that under certain conditions, it is possible to synthesize a metal hydride by the reaction of a metal with water or with a hydroxide. Such a synthesis route dispenses with the need for separately forming hydrogen by an expensive process and then to synthesize a hydride by metal–hydrogen reaction. If adopted in many of the hydrogen storage projects which plan to use a hydride for producing hydrogen through a chemical reaction or by a reversible dissociation for automobile use, this method could make a significant difference in making them cost-effective. 䉷 2006 International Association for Hydrogen Energy. Published by Elsevier Ltd. All rights reserved. Keywords: Metal hydrides; Synthesis
1. Introduction A binary metal hydride is usually prepared from a reaction between metal and hydrogen. Presence of oxygen in any form is considered detrimental to the synthesis and all water is avoided. In this study, we have shown that a hydride can be synthesized directly by a reaction between a metal and water or a water-bearing solid. Normal method of synthesizing a hydride is by reacting it with hydrogen, which requires that hydrogen be formed from an expensive process such as the electrolytic method or from fossil fuel. Many metals when brought in contact with water form hydrogen and cases of any hydride formation are rarely recorded [1]. We have shown both by thermodynamic calculations for a number of reactions and by experimenting with the Mg–H2 O system that a hydride can be synthesized by an exothermic reaction if a suitable adjustment is made in the amount of the reactants. The hydrides are important solids as hydrogen storage materials. There is currently a DOE supported project from Safe Hydrogen LLC [2] who use MgH2 slurry and water for ∗ Corresponding author.
E-mail address: saxenas@fiu.edu (S.K. Saxena).
vehicular use. Such projects will benefit immensely if MgH2 and other hydrides can be synthesized in a cost-effective way. The reaction studied here produces oxide or hydroxide along with the hydride. The total cost of winning the metal back must be compared with the cost of separately producing the hydrogen. The Safe Hydrogen LLC is working on a process for reduction of hydroxide to metal which they believe will be cost-effective [2]. 2. Hydride forming reaction To illustrate this process, the following reactions are considered: 2M+H2 O=MH2 +MO, and 2M+M(OH)2 =MH2 +2MO, where M could be a divalent cation such as Mg. Several other reactions using different metals are listed in Table 1. Thermodynamically both the sets of reactions proceed forward at room temperature (kinetics permitting). The availability of thermodynamic data restricts the list to the following but others could be added as more data becomes available: binary NaH, LiH, MgH2 and ternary LiAlH4 , NaAlH4 , LiBH4 and NaBH4 . Thermodynamic calculations showing the formation of these compounds are summarized in Table 1. The calculations have been performed using the FACTSAGE databases (mostly JANAF tables). The mixture of water and the solids must be carefully
0360-3199/$ - see front matter 䉷 2006 International Association for Hydrogen Energy. Published by Elsevier Ltd. All rights reserved. doi:10.1016/j.ijhydene.2006.09.032
2502
S.K. Saxena et al. / International Journal of Hydrogen Energy 32 (2007) 2501 – 2503
Table 1 Formation reactions of hydrides at 300 K No.
Reactants
Products
H , kJ
1. 2. 3. 4. 5. 6. 7. 8. 9.
Li + B + H2 O Li + B + Mg(OH)2 Na + B + H2 O 4 Li + H2 O 2 Mg + Mg(OH)2 2 Mg + H2 O Na + Al + H2 O Li + Al + H2 O at 225 bar Be + H2 O at 375 bar
0.5 LiBH4 + 0.5 LiBO2 MgO + 0.5 LiBH4 + 0.5 LiBO2 0.5 NaBO2 + 0.5 NaBH4 2 LiH + Li2 O 2 MgO + MgH2 MgO + MgH2 0.57 NaH + 0.36 NaAlH4 + 0.071NaAl9 O14 0.667 LiH + 0.333 LiAlH4 + 0.333 Al2 O3 0.5 BeO + 0.5 BeH2
−319 −282 −297 −494 −397 −392 −344 −365 −167
weighed for the hydride formation; otherwise the reaction to produce hydrogen sets in. It is possible to use hydroxide instead of water (Table 1). In each case, the product is a mixture of hydride and oxide, which must be separated from each other. This can be achieved with density separation methods because the hydrides are significantly lighter than the oxides. Data on Li–H2 O (reaction 4, Table 1) have been reported by Klanchar et al. [1] and others [3,4]. Klanchar et al. [1] used a thermoelectric calorimeter to study the lithium-water reaction and reported the final product as consisting of a mixture of Li2 O and LiH with small amounts of LiOH and H2 .
3. Experimental work Several experiments were conducted to test the thermodynamic predictions in the Mg–H2 O system. This system is important for hydrogen storage for automobile use because MgH2 may be used in oil-based slurry for filling the tanks [2]. In four different experiments, the formation of MgH2 from water proceeded via the reactions: Mg + 2H2 O = Mg(OH)2 + H2 and 2Mg + Mg(OH)2 = MgH2 + 2MgO. Therefore, only the data on the hydroxide reaction with metal is presented. The experiments were carried out in several different types of sealed containers. A mixture of micron sized magnesium metal and magnesium hydroxide was pressed into pellets. The pellets were placed in a quartz tube which was evacuated and sealed. The pellets were also put directly in a vacuum chamber. The samples were heated in the temperature range of 150–400 ◦ C. X-ray powder diffraction is done using Bruker GADDS/D8 X-ray system with Apex Smart CCD Detector and direct-drive rotating anode. The MacSci rotating anode (Molybdenum) is used with a 50 kV generator and 20 mA current. X-ray beam size can vary from 50 to 300 m. The high intensity X-ray beam results from the use of AXS optical device. The usual collection time is 1800 s. Fig. 1 shows the results obtained in vacuum heating. There are two small but distinct reflections for MgH2 and good development of the MgO reflections which show the validity of the thermodynamic reaction between the hydroxide and metal. The temperature for the dissociation of brucite to water and MgO is 263 ◦ C at 1 atm. It is likely that the
Fig. 1. X-ray results on a ball milled mixture of magnesium and brucite, Mg(OH)2 (1:1 molar), which was pressed in a disc and heated in vacuum for 24 h: the diffraction at 150 ◦ C shows a small peak due to MgH2 at around 2 value of 12.8 and shoulder on the Mg peak at 16.5; this development is more clearly established at 200 ◦ C.
dissociation temperature is lowered due to vacuum and water is released to combine with the metal. However, since any such water will be continuously removed, it is evident that the metal reacted directly with the hydroxide. The patterns in the evacuated and the vacuum heated samples at 200 ◦ C are compared again at 200 ◦ C to validate this argument (Fig. 2). While temperature may not be an important consideration in manufacturing a hydride, a low temperature production will of course be cost-effective. Fig. 2 show the X-ray diffraction of a sample of 1:2 molar mixture of Mg(OH)2 and Mg. At low temperatures there is a small recognizable shoulder peak of MgH2 at 2 of 16.3◦ . The final result of this heating at 400 ◦ C is a complete reaction of brucite and metal to form MgH2 and MgO.
S.K. Saxena et al. / International Journal of Hydrogen Energy 32 (2007) 2501 – 2503
Fig. 2. Comparison of X-ray results on a ball milled mixture of magnesium and brucite, Mg(OH)2 (2:1 molar). The diffraction from the sample heated in vacuum at 200 ◦ C clearly shows the formation of MgH2 (small peak at 2 of 12.8 and the shoulder peak on the Mg peak at 16.3◦ ). The other results are for a similar mixture, which was pressed in a disc and heated in a sealed evacuated quartz tube for 24 h: the diffraction at 150 ◦ C shows a small peak of MgH2 at around 2 value of 16.3◦ as a shoulder on the Mg peak; this development continues with heating and finally results in a fully developed pattern at 300 and 400 ◦ C. In the last heating practically all brucite and Mg is consumed resulting in MgO and MgH2 . The MgH2 peak intensities are half the intensity of the MgO peaks. A separately run mixture of 1:2 molar mix of MgH2 and MgO confirmed that the MgH2 produced is half in moles to that of MgO.
4. Discussion Since the purpose of this work was to establish that hydrides can be synthesized from a reaction between water or a hydroxide, no attempt is made here to study the kinetics which will depend on temperature and particle size. It appears that the reaction would be sufficiently rapid above 300 ◦ C. The oxides can be recycled to obtain the hydroxide. The metals will have to be recovered at some cost which needs to be taken into account for any application and must be compared with the cost of producing the hydride from a metal and hydrogen, when the latter is produced separately by an electrolytic reduction of water or from fossil fuel. McClaine et al. [2] have demonstrated a successful use of MgH2 slurry and water in producing hydrogen. In their process MgH2 is hydrolyzed resulting in Mg(OH)2 and hydrogen. It has been shown here
2503
that we can obtain MgH2 from Mg(OH)2 , thus making their process more economic by reducing the cost of making the hydride from Mg and water. Although we have used only the experimental information on the Mg–H–O system, several other hydrides, as listed in Table 1 and others not listed due to lack of thermodynamic data, can be synthesized by this method. Currently proposed methods that involve the hydrogen production in the fuel tank by chemical reactions under study by DOE (Department of Energy, US) supported researchers [5–7] are probably not cost-effective because of the need to remove the spent products and reduce the oxide to metal. There is some development in searching for a solution for recovering Mg from the hydroxide [2,8]. The study here does not solve the basic problem of the cost associated with the regeneration of the metal from the reaction products but in absence of any viable solution to a reversible dissociation of a hydride at ambient temperatures, the method suggested here should bring down the cost associated with hydrogen production in the fuel tank when using the chemical hydrides. Acknowledgment The authors’ work is supported through a grant from National Science Foundation (DMR-0231291) and a grant from Air Force (212600548). References [1] Klanchar M, Wintrode BD, Phillips JA. Lithium-water reaction chemistry at elevated temperature. Energy and Fuels 1997;11(4):931–5. [2] McClaine AW, Tullman S, Brown K. Chemical hydrogen slurry for hydrogen production and storage. FY 2005 Progress report: DOE hydrogen program. [3] Phillips J, Bradford MC, Klanchar M. A calorimetric study of the mechanism and thermodynamics of the lithium hydride-water reaction at elevated temperatures. Energy and Fuels 1995;9(4):569–73. [4] Bradford MC, Phillips J, Klanchar M. Novel high-temperature calvettype calorimeter for investigating metal-water reactions. Rev Sci Instrum 1995;66(1):171–5. [5] U.S. Department of Energy. Advanced Reciprocating Engines Program. Distributed Energy Resources.Washington, DC (www.eren.doe.gov/der). [6] U.S. Department of Energy. Federal Energy Technology Center. Distributed Generation—Securing America’s Future with Reliable, Flexible Power. Morgantown, WV. October (www.fetc.doe.gov). [7] Report of the Basic Energy Sciences Workshop on Hydrogen, Production, Storages and use, May 13–15, 2003. [8] FY 2005 Progress Report: DOE hydrogen program http://www. hydrogen.energy.gov/annual_progress05.html.
Synthesis of metal hydride from water Surendra K. Saxena ∗ , Vadym Drozd, Andriy Durygin Center for the Study of Matter at Extreme Conditions, College of Engineering, Florida International University, VH-140, University Park, Miami, FL 33199, USA Received 1 September 2006; received in revised form 18 September 2006; accepted 18 September 2006 Available online 8 December 2006
Abstract A number of hydrides are considered good candidates for hydrogen storage material for various applications in particular for automobile use. A metal hydride is synthesized through the reaction of a metal with hydrogen which is formed on industrial scale either by the electrolysis of water, by heating coke with steam in the water gas shift reaction or using hydrocarbons with steam. This study demonstrates that under certain conditions, it is possible to synthesize a metal hydride by the reaction of a metal with water or with a hydroxide. Such a synthesis route dispenses with the need for separately forming hydrogen by an expensive process and then to synthesize a hydride by metal–hydrogen reaction. If adopted in many of the hydrogen storage projects which plan to use a hydride for producing hydrogen through a chemical reaction or by a reversible dissociation for automobile use, this method could make a significant difference in making them cost-effective. 䉷 2006 International Association for Hydrogen Energy. Published by Elsevier Ltd. All rights reserved. Keywords: Metal hydrides; Synthesis
1. Introduction A binary metal hydride is usually prepared from a reaction between metal and hydrogen. Presence of oxygen in any form is considered detrimental to the synthesis and all water is avoided. In this study, we have shown that a hydride can be synthesized directly by a reaction between a metal and water or a water-bearing solid. Normal method of synthesizing a hydride is by reacting it with hydrogen, which requires that hydrogen be formed from an expensive process such as the electrolytic method or from fossil fuel. Many metals when brought in contact with water form hydrogen and cases of any hydride formation are rarely recorded [1]. We have shown both by thermodynamic calculations for a number of reactions and by experimenting with the Mg–H2 O system that a hydride can be synthesized by an exothermic reaction if a suitable adjustment is made in the amount of the reactants. The hydrides are important solids as hydrogen storage materials. There is currently a DOE supported project from Safe Hydrogen LLC [2] who use MgH2 slurry and water for ∗ Corresponding author.
E-mail address: saxenas@fiu.edu (S.K. Saxena).
vehicular use. Such projects will benefit immensely if MgH2 and other hydrides can be synthesized in a cost-effective way. The reaction studied here produces oxide or hydroxide along with the hydride. The total cost of winning the metal back must be compared with the cost of separately producing the hydrogen. The Safe Hydrogen LLC is working on a process for reduction of hydroxide to metal which they believe will be cost-effective [2]. 2. Hydride forming reaction To illustrate this process, the following reactions are considered: 2M+H2 O=MH2 +MO, and 2M+M(OH)2 =MH2 +2MO, where M could be a divalent cation such as Mg. Several other reactions using different metals are listed in Table 1. Thermodynamically both the sets of reactions proceed forward at room temperature (kinetics permitting). The availability of thermodynamic data restricts the list to the following but others could be added as more data becomes available: binary NaH, LiH, MgH2 and ternary LiAlH4 , NaAlH4 , LiBH4 and NaBH4 . Thermodynamic calculations showing the formation of these compounds are summarized in Table 1. The calculations have been performed using the FACTSAGE databases (mostly JANAF tables). The mixture of water and the solids must be carefully
0360-3199/$ - see front matter 䉷 2006 International Association for Hydrogen Energy. Published by Elsevier Ltd. All rights reserved. doi:10.1016/j.ijhydene.2006.09.032
2502
S.K. Saxena et al. / International Journal of Hydrogen Energy 32 (2007) 2501 – 2503
Table 1 Formation reactions of hydrides at 300 K No.
Reactants
Products
H , kJ
1. 2. 3. 4. 5. 6. 7. 8. 9.
Li + B + H2 O Li + B + Mg(OH)2 Na + B + H2 O 4 Li + H2 O 2 Mg + Mg(OH)2 2 Mg + H2 O Na + Al + H2 O Li + Al + H2 O at 225 bar Be + H2 O at 375 bar
0.5 LiBH4 + 0.5 LiBO2 MgO + 0.5 LiBH4 + 0.5 LiBO2 0.5 NaBO2 + 0.5 NaBH4 2 LiH + Li2 O 2 MgO + MgH2 MgO + MgH2 0.57 NaH + 0.36 NaAlH4 + 0.071NaAl9 O14 0.667 LiH + 0.333 LiAlH4 + 0.333 Al2 O3 0.5 BeO + 0.5 BeH2
−319 −282 −297 −494 −397 −392 −344 −365 −167
weighed for the hydride formation; otherwise the reaction to produce hydrogen sets in. It is possible to use hydroxide instead of water (Table 1). In each case, the product is a mixture of hydride and oxide, which must be separated from each other. This can be achieved with density separation methods because the hydrides are significantly lighter than the oxides. Data on Li–H2 O (reaction 4, Table 1) have been reported by Klanchar et al. [1] and others [3,4]. Klanchar et al. [1] used a thermoelectric calorimeter to study the lithium-water reaction and reported the final product as consisting of a mixture of Li2 O and LiH with small amounts of LiOH and H2 .
3. Experimental work Several experiments were conducted to test the thermodynamic predictions in the Mg–H2 O system. This system is important for hydrogen storage for automobile use because MgH2 may be used in oil-based slurry for filling the tanks [2]. In four different experiments, the formation of MgH2 from water proceeded via the reactions: Mg + 2H2 O = Mg(OH)2 + H2 and 2Mg + Mg(OH)2 = MgH2 + 2MgO. Therefore, only the data on the hydroxide reaction with metal is presented. The experiments were carried out in several different types of sealed containers. A mixture of micron sized magnesium metal and magnesium hydroxide was pressed into pellets. The pellets were placed in a quartz tube which was evacuated and sealed. The pellets were also put directly in a vacuum chamber. The samples were heated in the temperature range of 150–400 ◦ C. X-ray powder diffraction is done using Bruker GADDS/D8 X-ray system with Apex Smart CCD Detector and direct-drive rotating anode. The MacSci rotating anode (Molybdenum) is used with a 50 kV generator and 20 mA current. X-ray beam size can vary from 50 to 300 m. The high intensity X-ray beam results from the use of AXS optical device. The usual collection time is 1800 s. Fig. 1 shows the results obtained in vacuum heating. There are two small but distinct reflections for MgH2 and good development of the MgO reflections which show the validity of the thermodynamic reaction between the hydroxide and metal. The temperature for the dissociation of brucite to water and MgO is 263 ◦ C at 1 atm. It is likely that the
Fig. 1. X-ray results on a ball milled mixture of magnesium and brucite, Mg(OH)2 (1:1 molar), which was pressed in a disc and heated in vacuum for 24 h: the diffraction at 150 ◦ C shows a small peak due to MgH2 at around 2 value of 12.8 and shoulder on the Mg peak at 16.5; this development is more clearly established at 200 ◦ C.
dissociation temperature is lowered due to vacuum and water is released to combine with the metal. However, since any such water will be continuously removed, it is evident that the metal reacted directly with the hydroxide. The patterns in the evacuated and the vacuum heated samples at 200 ◦ C are compared again at 200 ◦ C to validate this argument (Fig. 2). While temperature may not be an important consideration in manufacturing a hydride, a low temperature production will of course be cost-effective. Fig. 2 show the X-ray diffraction of a sample of 1:2 molar mixture of Mg(OH)2 and Mg. At low temperatures there is a small recognizable shoulder peak of MgH2 at 2 of 16.3◦ . The final result of this heating at 400 ◦ C is a complete reaction of brucite and metal to form MgH2 and MgO.
S.K. Saxena et al. / International Journal of Hydrogen Energy 32 (2007) 2501 – 2503
Fig. 2. Comparison of X-ray results on a ball milled mixture of magnesium and brucite, Mg(OH)2 (2:1 molar). The diffraction from the sample heated in vacuum at 200 ◦ C clearly shows the formation of MgH2 (small peak at 2 of 12.8 and the shoulder peak on the Mg peak at 16.3◦ ). The other results are for a similar mixture, which was pressed in a disc and heated in a sealed evacuated quartz tube for 24 h: the diffraction at 150 ◦ C shows a small peak of MgH2 at around 2 value of 16.3◦ as a shoulder on the Mg peak; this development continues with heating and finally results in a fully developed pattern at 300 and 400 ◦ C. In the last heating practically all brucite and Mg is consumed resulting in MgO and MgH2 . The MgH2 peak intensities are half the intensity of the MgO peaks. A separately run mixture of 1:2 molar mix of MgH2 and MgO confirmed that the MgH2 produced is half in moles to that of MgO.
4. Discussion Since the purpose of this work was to establish that hydrides can be synthesized from a reaction between water or a hydroxide, no attempt is made here to study the kinetics which will depend on temperature and particle size. It appears that the reaction would be sufficiently rapid above 300 ◦ C. The oxides can be recycled to obtain the hydroxide. The metals will have to be recovered at some cost which needs to be taken into account for any application and must be compared with the cost of producing the hydride from a metal and hydrogen, when the latter is produced separately by an electrolytic reduction of water or from fossil fuel. McClaine et al. [2] have demonstrated a successful use of MgH2 slurry and water in producing hydrogen. In their process MgH2 is hydrolyzed resulting in Mg(OH)2 and hydrogen. It has been shown here
2503
that we can obtain MgH2 from Mg(OH)2 , thus making their process more economic by reducing the cost of making the hydride from Mg and water. Although we have used only the experimental information on the Mg–H–O system, several other hydrides, as listed in Table 1 and others not listed due to lack of thermodynamic data, can be synthesized by this method. Currently proposed methods that involve the hydrogen production in the fuel tank by chemical reactions under study by DOE (Department of Energy, US) supported researchers [5–7] are probably not cost-effective because of the need to remove the spent products and reduce the oxide to metal. There is some development in searching for a solution for recovering Mg from the hydroxide [2,8]. The study here does not solve the basic problem of the cost associated with the regeneration of the metal from the reaction products but in absence of any viable solution to a reversible dissociation of a hydride at ambient temperatures, the method suggested here should bring down the cost associated with hydrogen production in the fuel tank when using the chemical hydrides. Acknowledgment The authors’ work is supported through a grant from National Science Foundation (DMR-0231291) and a grant from Air Force (212600548). References [1] Klanchar M, Wintrode BD, Phillips JA. Lithium-water reaction chemistry at elevated temperature. Energy and Fuels 1997;11(4):931–5. [2] McClaine AW, Tullman S, Brown K. Chemical hydrogen slurry for hydrogen production and storage. FY 2005 Progress report: DOE hydrogen program. [3] Phillips J, Bradford MC, Klanchar M. A calorimetric study of the mechanism and thermodynamics of the lithium hydride-water reaction at elevated temperatures. Energy and Fuels 1995;9(4):569–73. [4] Bradford MC, Phillips J, Klanchar M. Novel high-temperature calvettype calorimeter for investigating metal-water reactions. Rev Sci Instrum 1995;66(1):171–5. [5] U.S. Department of Energy. Advanced Reciprocating Engines Program. Distributed Energy Resources.Washington, DC (www.eren.doe.gov/der). [6] U.S. Department of Energy. Federal Energy Technology Center. Distributed Generation—Securing America’s Future with Reliable, Flexible Power. Morgantown, WV. October (www.fetc.doe.gov). [7] Report of the Basic Energy Sciences Workshop on Hydrogen, Production, Storages and use, May 13–15, 2003. [8] FY 2005 Progress Report: DOE hydrogen program http://www. hydrogen.energy.gov/annual_progress05.html.