Nanomaterials for Solid State Hydrogen Storage

Nanomaterials for Solid State Hydrogen Storage Fuel Cells and Hydrogen Energy Series Editor: Narottam P. Bansal NASA...

Author: Robert A. Varin | Tomasz Czujko | Zbigniew S. Wronski

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Nanomaterials for Solid State Hydrogen Storage

Fuel Cells and Hydrogen Energy Series Editor:

Narottam P. Bansal NASA Glenn Research Center Cleveland, OH 44135 [email protected]

Aims and Scope of the Series During the last couple of decades, notable developments have taken place in the science and technology of fuel cells and hydrogen energy. Most of the knowledge developed in this field is contained in individual journal articles, conference proceedings, research reports, etc. Our goal in developing this series is to organize this information and make it easily available to scientists, engineers, technologists, designers, technical managers and graduate students. The book series is focused to ensure that those who are interested in this subject can find the information quickly and easily without having to search through the whole literature. The series includes all aspects of the materials, science, engineering, manufacturing, modeling, and applications. Fuel reforming and processing; sensors for hydrogen, hydrocarbons and other gases will also be covered within the scope of this series. A number of volumes edited/authored by internationally respected researchers from various countries are planned for publication during the next few years. Titles in this series Nanomaterials for Solid State Hydrogen Storage R.A. Varin, T. Czujko, and Z. S. Wronski ISBN 978-0-387-77711-5, 2009 Modeling Solid Oxide Fuel Cells: Methods, Procedures and Techniques R. Bove and S. Ubertini, eds. ISBN 978-1-4020-6994-9, 2008

Robert A.Varin • Tomasz Czujko Zbigniew S. Wronski

Nanomaterials for Solid State Hydrogen Storage

Robert A. Varin University of Waterloo Department of Mechanical and Mechatronics Engineering 200 University Ave. W Waterloo, Ontario Canada N2L 3G1

Tomasz Czujko University of Waterloo Department of Mechanical and Mechatronics Engineering 200 University Ave. W Waterloo, Ontario Canada N2L 3G1

Zbigniew S. Wronski CANMET Energy Technology Centre Hydrogen Fuel Cells and Transportation Energy Natural Resources Canada 1 Haanel Drive Ottawa, Ontario Canada K1A 1M1

ISBN: 978-0-387-77711-5 e-ISBN: 978-0-387-77712-2 DOI: 10.1007/978-0-387-77712-2 Library of Congress Control Number: 2008929618 © Springer Science+Business Media, LLC 2009 All rights reserved. This work may not be translated or copied in whole or in part without the written permission of the publisher (Springer Science+Business Media, LLC, 233 Spring Street, New York, NY 10013, USA), except for brief excerpts in connection with reviews or scholarly analysis. Use in connection with any form of information storage and retrieval, electronic adaptation, computer software, or by similar or dissimilar methodology now known or hereafter developed is forbidden. The use in this publication of trade names, trademarks, service marks, and similar terms, even if they are not identified as such, is not to be taken as an expression of opinion as to whether or not they are subject to proprietary rights. Printed on acid-free paper springer.com

Preface

Although hydrogen as a chemical element has been known to the humankind and used in various capacities for a very long time, only in the past 15 years its importance to the world population as an energy vector has gradually emerged. A long-term reliance of humanity on the energy derived solely from fossil fuels, such as coal in the nineteenth and crude oil and natural gas in the twentieth century, has led to a number of new challenges facing all of us in the twenty-first century, such as sharp reduction in the world crude oil and eventually coal supply, global warming and following climate changes due to the release of growing amounts of greenhouse gas CO2, and poor urban air quality. Hydrogen is essentially the only viable remedy for the growing world energy problems. Hydrogen is a very attractive alternative energy vector for replacing fossil fuel-based economy. The future Hydrogen Economy offers a potential solution to satisfying the global energy requirements while reducing (and eventually eliminating) carbon dioxide and other greenhouse gas emissions and improving energy security. Hydrogen is ubiquitous, clean, efficient, and can be produced directly from sunlight and water by biological organisms and using semiconductor-based systems similar to photovoltaics. Hydrogen can also be produced indirectly via thermal processing of biomass or fossil fuels where the development of advanced technological processes combined with a CO2 sequestration is emerging. However, this rosy picture, as it usually happens in a real life, is marred by a number of obstacles which must be overcome before the Hydrogen Economy becomes a reality. One of these obstacles is safe and efficient storage of hydrogen particularly for mobile/automotive applications where hydrogen gas will be supplied to fuel cells that, in turn, will power the transport vehicles in a clean, inexpensive, safe, and efficient manner. From all possible solutions to hydrogen storage the one which relies upon storage in solid media (hydrides) is the most attractive one. The fast emerging nanoscience/nanotechnology will allow fabricating nanomaterials for solid-state hydrogen storage that can, in a long run, revolutionize hydrogen storage. This book is our modest contribution to this innovative area of hydrogen storage. Wherever possible we tried to illustrate the hydrogen storage behavior by our own results. In Chap. 1, we introduce the reader to the motivation for the transformation to the Hydrogen Economy. In a number of following sections/subsections, we v

vi

Preface

provide a comprehensive synchronic history of development of hydrides and nanomaterials including the existing fabrication methods with a special emphasis on ball (mechanical) milling in high-energy mills. Important hydride properties and experimental techniques for assessing hydrogen storage behavior are also discussed. In Chap. 2, we review hydrogen storage properties of selected simple metal and intermetallic hydrides with the most emphasis on magnesium hydride (MgH2) which now can be treated as a model hydride whose hydrogen storage properties in nanostructured form can be used as a benchmark for comparing the properties of other hydrides. Chapter 3 brings a thorough review of the properties of complex hydrides whose high volumetric and gravimetric capacities make them most attractive for the vehicular solid-state hydrogen storage in transportation. Chapter 4 provides information on carbons and nanocarbons as alternative means of hydrogen storage to solid hydrides. This includes diamond and nanodiamond, graphene, ordered graphites and nanographites, disordered and active carbons, fullerenes, carbon nanotubes, and other nanoshapes. Chapter 5 is a sort of an executive summary where we provide a critical assessment of the present state of knowledge and make predictions for the future developments. Waterloo, ON Waterloo, ON Ottawa, ON

Robert A. Varin Tomasz Czujko Zbigniew S. Wronski

Contents

1

Introduction ..............................................................................................

1

1.1 1.2

1

Motivation: The Hydrogen Economy ................................................ Brief, Synchronic History of Development of Hydrides and Nanomaterials......................................................... 1.2.1 Early Investigations of Metal–Hydrogen Systems and Hydrides .......................................................................... 1.2.2 Early Routes to Nanomaterials .............................................. 1.2.3 Historical Development of Classical Hydrogen Storage AB5 Alloys................................................................ 1.2.4 Historical Development of Interstitial Hydrides in Other Intermetallic Systems .............................................. 1.2.5 Historical Development of Nanophase AB2 Intermetallic Hydrides ........................................................... 1.2.6 New Routes to Nanomaterials: Mechanical Alloying and Mechanochemical Activation ......................................... 1.2.7 Historical Development of Lightweight Metal Hydrides and Hydride Complexes ........................................................ 1.2.8 Early Studies of Noninterstitial Transition Metal Ternary Hydrides ................................................................... 1.2.9 Toward Chemical/Complex Hydrides ................................... 1.2.10 Historical Development of Nanocarbons and Carbon Nanotubes .......................................................... 1.2.11 New Materials and Techniques.............................................. 1.3 Nanoprocessing in Solid State in High-Energy Ball Mills................ 1.3.1 Processes for the Synthesis of Nanostructured Materials ................................................................................ 1.3.2 Milling Processes and Equipment ......................................... 1.3.3 Nanoprocessing Methods and Mechanisms .......................... 1.3.3.1 Mechanical Milling ................................................ 1.3.3.2 Mechanical Alloying ..............................................

7 7 10 13 15 16 17 18 20 21 23 25 27 27 28 37 38 39

vii

viii

Contents

1.3.3.3 1.3.3.4

Mechanochemical Activation ................................. Mechanochemical Synthesis (Mechanosynthesis) of Nanohydrides .................... 1.3.3.5 Mechanical Amorphization .................................... 1.4 Important Hydride Properties and Experimental Techniques ........... 1.4.1 Thermodynamics ................................................................... 1.4.1.1 Pressure–Composition–Temperature (PCT) Properties ..................................................... 1.4.1.2 Calculation of Activation Energy ........................... 1.4.2 PCT and Kinetic Curves Determination by Volumetric Method in a Sieverts-Type Apparatus ................................... 1.4.3 Microstructural Characterization of Ball-Milled Hydrides ................................................................................ 1.4.4 Weight Percent of a Hydride Phase and Hydrogen by DSC Method ..................................................................... References .................................................................................................. 2

40 52 55 56 56 56 60 65 71 73 74

Simple Metal and Intermetallic Hydrides .............................................

83

2.1

83 83

Mg/MgH2 ........................................................................................... 2.1.1 Crystallographic and Material Characteristics ...................... 2.1.2 Hydrogen Storage Characteristics of Commercial Mg and MgH2 ............................................... 2.1.2.1 Absorption .............................................................. 2.1.2.2 Desorption .............................................................. 2.1.3 Hydrogen Storage Characteristics of Mechanically (Ball) Milled MgH2................................................................ 2.1.3.1 Microstructural Evolution During Milling and Subsequent Cycling of Commercial MgH2 Powders .................................................................. 2.1.3.2 Hydrogen Absorption of Ball-milled Commercial MgH2 Powders ................................... 2.1.3.3 Hydrogen Desorption of Ball-milled Commercial MgH2 Powders ................................... 2.1.4 Hydrogen Storage Characteristics of MgH2 Synthesized by Reactive Mechanical (Ball) Milling of Mg ...................... 2.1.5 Aging Effects in Stored MgH2 Powders ................................ 2.1.6 Other Methods of Synthesis of Nanostructured MgH2 than Ball Milling ......................................................... 2.2 MgH2 with Catalytic Additives ......................................................... 2.2.1 Mg/MgH2–Metals and Intermetallics .................................... 2.2.1.1 Desorption in Vacuum ................................................. 2.2.1.2 Desorption at Atmospheric Pressure of Hydrogen..... 2.2.2 Mg/MgH2–Metal Oxides ....................................................... 2.2.3 Mg/MgH2–Carbon / Graphite and Carbon Nanotubes ...........

87 87 93 102

103 112 115 129 146 147 151 152 152 153 165 169

Contents

3

4

ix

2.3 Other Metal Hydrides Containing Mg............................................... 2.4 AlH3 ................................................................................................... 2.5 Other Metal and Intermetallic-based Hydrides: New Developments ........................................................................... 2.5.1 Metal Hydrides ...................................................................... 2.5.2 Rare-Earth AB5 Compounds.................................................. 2.5.3 Titanium–Iron AB Compounds ............................................. 2.5.4 Titanium and Zirconium AB2 Compounds ............................ 2.5.5 Other Novel Intermetallic Hydrides ...................................... References ..................................................................................................

170 174

Complex Hydrides ...................................................................................

195

3.1

Ternary Transition Metal Complex Hydrides ................................... 3.1.1 Mg2NiH4 ................................................................................ 3.1.2 Mg2FeH6 ................................................................................ 3.1.3 Mg2CoH5 ................................................................................ 3.2 Alanates ............................................................................................. 3.2.1 NaAlH4 .................................................................................. 3.2.2 LiAlH4.................................................................................... 3.2.3 Mg(AlH4)2 and Ca(AlH4)2...................................................... 3.3 Amides .............................................................................................. 3.4 Metal Borohydrides........................................................................... 3.5 Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing ...................................................... 3.5.1 MgH2–LiAlH4 Composite System ......................................... 3.5.2 MgH2–NaAlH4 Composite System ........................................ 3.5.3 MgH2–NaBH4 Composite System ......................................... References ..................................................................................................

196 196 198 204 206 206 213 223 231 240

Carbons and Nanocarbons ......................................................................

291

4.1 4.2

291 294 294 295

Diamond and Nanodiamonds ............................................................ Graphene, Ordered Graphite, and Nanographites ............................. 4.2.1 Graphene................................................................................ 4.2.1.1 In-Plane σ and Out-of-Plane π Bonding................. 4.2.1.2 Van der Walls Interplanar and Intermolecular Interactions ............................................................ 4.2.1.3 Physisorption of Hydrogen on Carbons ................. 4.2.1.4 Chemisorption of Hydrogen on Carbons ................ 4.2.2 Graphitic Nanofibers, Whiskers, and Polyhedral Crystals .... 4.2.3 Graphite ................................................................................. 4.3 Disordered and Active Carbons......................................................... 4.3.1 Disordered Graphites and Mechanically-Activated Carbons ..................................................................................

177 179 181 182 183 183 183

253 255 265 270 281

296 297 298 299 299 301 301

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Contents

4.3.2 Active Carbons and Chemically Activated Carbons ............. 4.3.3 Amorphous Carbon ............................................................... 4.4 Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns...................................................................... 4.4.1 Fullerenes and Hydrofullerenes ............................................. 4.4.2 Carbon Nanotubes ................................................................. 4.4.3 Carbon Nanohorns ................................................................. 4.4.4 Nanostructured Carbon Shells and Carbon Onions ............... References ..................................................................................................

305 305 308 312 314 317

Summary...................................................................................................

321

5.1 5.2 5.3

Metal/Intermetallic Hydrides ............................................................ Complex Hydrides ............................................................................. Nanocarbons and Others ...................................................................

322 323 324

Index ................................................................................................................

327

5

303 304

Chapter 1

Introduction

1.1

Motivation: The Hydrogen Economy

The energy supply to the humankind in the last two centuries was solely based on fossil fuels such as coal in the nineteenth century and crude oil and natural gas in the twentieth century. Unfortunately, this fossil fuel-based economy has led to a number of new challenges facing all of us in the twenty-first century such as global warming and following climate changes due to the release of growing amounts of greenhouse gas CO2, poor urban air quality, and reduction in the world crude oil supply, which could reach so-called Hubbert’s Peak around year 2011–2020. It is also noted that no major oil field has been discovered since 1970 [1]. Since the mid-1970s a concept of the ecologically clean Hydrogen Economy has been gaining momentum as essentially the only viable remedy for the growing world energy problems. The Hydrogen Economy offers a potential solution to satisfying the global energy requirements while reducing (and eventually eliminating) carbon dioxide and other greenhouse gas emissions and improving energy security. Hydrogen is a very attractive alternative energy vector. It is ubiquitous, clean, efficient, and can be produced directly from sunlight and water by biological organisms and using semiconductor-based systems similar to photovoltaics. Hydrogen can also be produced indirectly via thermal processing of biomass or fossil fuels where the development of advanced technological processes combined with a CO2 sequestration [2] have the potential to produce essentially unlimited quantities of hydrogen in a sustainable manner. For example, electricity produced by wind turbines or nuclear power plants during off-peak periods [3] can be used for the electrolysis of water into hydrogen [4] and the latter stored for future distribution to places of use [1]. When hydrogen burns, it releases energy as heat and produces water: 2H2 + O2→2H2O. Since no carbon is involved then using hydrogen produced from renewable or nuclear energy as energy resource would eliminate carbon monoxide and CO2 emissions and reduce greenhouse warming. However, if air is used for flame combustion of hydrogen, small amounts of NOx may be produced. In a fuel cell, hydrogen is converted directly to electricity in a similar reaction to the earlier one for burning hydrogen, and in essence just electricity and water are produced. Pure hydrogen enters the anode channel in a fuel cell and diffuses through a porous R.A. Varin et al., Nanomaterials for Solid State Hydrogen Storage, DOI: 10.1007/978-0-387-77712-2_1 © Springer Science + Business Media, LLC 2009

1

2

1

Introduction

anode toward the catalyst (Pt) where the hydrogen molecules H2 are stripped of their electrons and become positively charged ions (protons) (H2 → 2H+ + 2e−). Protons then migrate through the proton-permeable polymeric (Nafion) membrane (Proton Exchange Membrane also called Polymer Electrolyte Membrane – PEM) and the electrons generated during oxidation pass through the external circuit to the cathode, thereby creating electric current. On the cathode side, humidified air enters the cathode channel and diffuses toward the cathode-side catalyst layer. At the catalyst Pt surface, hydrogen protons recombine with electrons and oxygen molecules in air to produce water and heat according to the overall reaction ½O2 + 2H+ + 2e− → H2O + heat. This waste heat gives a PEM fuel cell an operating temperature of 60–80°C [5]. And last but not least is the fact that a hydrogen fuel-cell car can convert hydrogen energy into motion about 2–3 times as efficiently as a normal car converts gasoline energy into motion: depending on how it is designed and run, a good fuel-cell system is about 50–70% efficient in converting hydrogen to electricity, while a typical car engine’s efficiency from gasoline to output shaft averages to only about 15–17% [6]. Scott [7] succinctly summarized the inevitability of hydrogen becoming the fuel of the future by two rationales: depletion-based rationale and climate-based rationale. Driven by depletion, civilization must move from fossil fuels to sustainable energy sources. Realistically, the only way sustainable sources can be harvested to make chemical fuels is via hydrogen. Otherwise, how else can we get energy from wind, solar, or nuclear power to fuel an airplane? On the other hand, atmospheric CO2 growth is such that the concentration of CO2 in the atmosphere has increased from 280 to 370 ppm over the past 150 years. CO2 emissions can only be slowed by the extensive use of hydrogen and can only be stopped with the supremacy of sustainable-derived H2 among chemical fuels. The realization of the enormous benefits of the Hydrogen Economy has triggered over the last 15 years intense activities in the development of hydrogenrelated technologies. There are three major technological obstacles to the full implementation of the Hydrogen Economy in the next few decades. The first is the cost of safe and efficient production of hydrogen gas. At present, 48% of hydrogen is produced from methane steam reforming, 30% from oil/naphtha reforming, 18% from coal gasification, and only 3.9% from the electrolysis of water [8]. Apparently, the bulk of hydrogen production still relies upon fossil fuels, a by-product of which is CO2. However, the fossil fuel-based processes are much cheaper than the electrolysis of water. The efforts are underway to reduce the price of electrolysisderived hydrogen to $2–3 per kg. The second obstacle is the further development of PEM fuel cell (PEMFC), which is the primary cell most suitable for transportation. Most importantly, the research is focused on the extension of the usable service life, water flooding, dynamics, and reliability. The present cost of energy derived from a PEMFC is around $200 per kW and this must be reduced to around $30 per kW before PEMFC could be fully commercialized. The third obstacle is hydrogen storage for supplying PEMFC. There are three major competing technologies for hydrogen storage: compressed gas cylinders, liquid hydrogen tanks, and metal hydrides [9, 10]. Their comparison is shown in Table 1.1. A major drawback

1.1

Motivation: The Hydrogen Economy

3

Table 1.1 Comparison of three major competing technologies for hydrogen storage (based on [9, 10])

Storage system

Volumetric hydrogen capacity Drawbacks (kgH2 m−3)

Compressed hydrogen gas under 80 MPa pressure

~40

Liquid hydrogen at cryogenic tank at −252°C (21 K) Solid state hydrides

~71 80–160

Safety problem since enormous pressures are required; cost of pressurization; large pressure drop during use; hydrogen embrittlement of storage tanks Large thermal losses (open system); safety; cost of liquefaction None of the above

of compressed hydrogen storage for transportation applications is the small amount of hydrogen that may be stored in a reasonable volume (volumetric capacity/density). As can be seen in Table 1.1 compressed hydrogen gas technologies, even at such enormous pressures as ~80 MPa, also suffer from low volumetric densities not exceeding ~40 kgH2 m−3. As pointed out by Sandi [10] even at such a high pressure as 70–80 MPa, the energy content of compressed hydrogen is significantly less than that for the same volume of gasoline: 4.4 MJ L−1 (at 70 MPa) for hydrogen compared with 31.6 MJ L−1 for gasoline. Even though considered to be quite simple and inexpensive, the high pressure of 80 MPa involved in hydrogen gas cylinders raises safety concern. There is also some cost involved with compression to such high pressures. Another consideration is the large pressure drop during use. In addition, because most of the system parts exposed to hydrogen will be metallic, there is a concern of well-known hydrogen embrittlement [10]. The liquid hydrogen tank for its part offers almost twice as high storage capacity by volume as pressurized hydrogen; however, this is still less than half that required by the Department of Energy (D.O.E. FreedomCAR goal) (the requirements will be discussed later). A major drawback of liquid storage is a big cost of liquefaction, which today can add as much as 50% to the cost of H2 [6, 9, 10]. There are also safety issues associated with the handling of cryogenic liquids and the problem of evaporative loss. Solid-state hydrides that include metal/intermetallic and complex (chemical) hydrides are characterized by the highest volumetric capacities, and they do not suffer drawbacks as those experienced by compressed and liquid hydrogen. Because of the low pressures involved in metal hydride technologies and the fact that the release of hydrogen takes place via an endothermic process, this method of hydrogen storage is, however, the safest of all. Moreover, the hydrogen released from a metal hydride is of very high purity and, therefore, can be used directly to feed a PEM fuel cell. The U.S. Department of Energy (DOE) introduced a number of targets for onboard hydrogen storage systems within the frame work of its FreedomCAR program for the years 2007, 2010, and 2015 [11, 12], which are listed in Table 1.2. It is now appropriate to discuss solid state hydrogen storage in hydrides in the context of the targets shown in Table 1.2.

4

1

Introduction

Table 1.2 US DOE FreedomCAR hydrogen storage system targets [11, 12] Targeted factor −1

Specific energy (MJ kg ) System gravimetric capacity (wt%) System volumetric capacity (kgH2 m−3) Energy density (MJ L−1) Storage system cost ($ per kgH2) System cost ($ per kg per system) Operating temperature (°C) Min/max delivery temperature (°C) Cycle life-time (absorption/desorption cycles) Flow rate (full throttle) (g s−1) Delivery pressure from tank to FC (bar) Transient response(s) (10–90% and 90–0%) Refueling rate (kg H2 min−1)

2007

2010

2015

– 4.5 36 – 200 – −20/50 −30/85 500 3 2.5 30 0.5

7.2 6 45 5.4 133 6 −30/50 −40/85 1,000 4 2.5 15 1.5

10.8 9 81 9.72 67 3 −40/60 −40/85 1,500 5 2.5 15 2.0

Table 1.3 Hydrogen storage properties of intermetallic compounds (from [12]) Maximum hydrogen capacity Type

Intermetallic

Hydride

wt%

Temperature for 1 atm Pdesorption (°C)

A 2B AB AB AB2 AB5/MmB5 AB2

Mg2Ni FeTi ZrNi ZrMn2 LaNi5 TiV2 (TiV0.62Mn1.5)

Mg2NiH4 FeTiH2 ZrNiH3 ZrMn2H3.6 LaNi5H6 TiV0.62Mn1.5H2.5

3.6 1.86 1.85 1.77 1.49 2.15

255 −8 292 167 12 −6

Conventional metal hydrides based on metals such as V, Nb, Pd, Li, Na, etc. have gravimetric capacities too low for any commercial consideration in mobile hydrogen storage with the exception of LiH, which has high capacity but extremely high desorption temperature [13]. A notable exception of a metal hydride is Mg, which has a relatively high gravimetric capacity and can desorb at 300°C and slightly below after nanostructuring treatment (it will be discussed later). In essence, none of the metal hydrides can meet the DOE targets [13, 14]. Similarly, hydrides based on intermetallic compounds AB (FeTi, ZrNi), AB2 (ZrMn2/TiMn2/TiCr2), AB5 (LaNi5 or MmNi5 where Mm – mischmetal), and A2B (Mg2Ni) have relatively low gravimetric storage capacities as shown in Table 1.3, which are unsuitable for transportation storage although a number of them desorb hydrogen within the temperature range targeted by DOE at the desorption pressure (Pdesorption) of 1 atm (which is more or less an operating pressure of a PEM FC). However, there exist many complex hydrides having high and very high gravimetric storage capacities some of which are shown in Table 1.4 [13–22]. Their theoretical capacity is calculated as the ratio of the atomic mass of hydrogen in

1.1

Motivation: The Hydrogen Economy

5

Table 1.4 Hydrogen storage properties of selected high-capacity hydrides [13–22] Metal– Theoretical maxi- Theoretical revers- Approx. desorpible gravimetric hydrogen tion temperature mum gravimetric system Hydride range (°C) H2 capacity) (wt%) capacity (wt%) Li–B–H Mg–B–H Fe–B–H Ca–B–H Na–B–H Li–Al–H Al–H Mg–Al–H Li–N–H Zn–B–H Ca–Al–H Mg–H Na–Al–H Mg–N–H Mg–Fe–H Na–N–H

LiBH4 Mg(BH4)2 Fe(BH4)3 Ca(BH4)2 NaBH4 LiAlH4 AlH3 Mg(AlH4)2 LiNH2(+LiH + TiCl3) Zn(BH4)2 Ca(AlH4)2 MgH2 NaAlH4 Mg(NH2)2(+LiH) Mg2FeH6 NaNH2

18.4 14.9 12.1 11.6 10.6 10.6 10.0 9.3 8.8 8.5 7.9 7.6 7.5 7.2 5.5 5.3

~13.8 ~11.2 Unknown Unknown 10.6 ~7.9 10.0 ~7.0 ~6.0 8.5(?) ~5.9 7.6 5.6 ~7.0 5.5 Unknown

~470 ~300 Unknown ~320(?) 400–600 110–260 ~150 110–160 150–280 85–140 80–180 300–400 229–247 140–250 300–400 <200(?)

the hydride formula to the molecular mass of hydride. Some of these hydrides called complex hydrides (like borohydrides) decompose in a multistage sequence such as, for example, LiBH4 that decomposes in the first stage into LiH + B + (3/2)H2 and in the second stage LiH decomposes into Li and H. However, only the first reaction that releases about 13.8 wt% H2 (1.5 mol of H per 1 mol of LiBH4) is potentially reversible [23]. Therefore, the fourth column in Table 1.4 includes so-called theoretical reversible gravimetric capacity, which is the amount of hydrogen that is potentially feasible to be reversibly desorbed/absorbed from such a complex hydride. Unfortunately, as shown in the last column in Table 1.4 the major problem of complex hydrides is that their desorption temperatures are not even close to the operating temperature range required by the DOE targets (Table 1.2). Some of the borohydrides such as Zn(BH4)2 that start desorbing around 80°C, close to the operating temperature of PEM FC, release a toxic borane gas B2H6 together with hydrogen [24, 25], which can quickly destroy a PEM FC membrane. Therefore, the major focus of research in the last decade is on finding the means to substantially reduce the desorption/absorption temperature of hightemperature complex hydrides and in addition to improve their absorption/desorption kinetics where applicable. As can be clearly seen in Table 1.4 only high-capacity hydrides are important for onboard hydrogen storage for vehicular applications. As pointed out in [11] depending on the storage material and on the system design, material capacities may need to be a factor of up to two times higher than system capacity targets. If such a rule is to be applied to the highcapacity hydrides in Table 1.4 than vis-à-vis extremely restrictive DOE targets

6

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Introduction

in Table 1.2 most of the high-capacity hydrides should not be even considered as on-board storage materials for PEMFC vehicles. This of course is very unreasonable and it seems that sooner or later the DOE targets must be updated to be in line with the reality. The second parameter of an utmost importance is desorption/absorption temperature. In reality, it may be extremely difficult to find a material desorbing in exactly the same temperature range as required by the DOE targets in Table 1.2. However, materials with higher desorption temperatures cannot be completely ruled out of hand as viable storage media. First, there is a recent trend to increase the working temperature of the PEM fuel cell by using an electrolyte membrane of polybenzimidazole (PBI) doped with phosphoric acid [26]. Second, a nanostructured MgH2 with a catalyst that reduces desorption temperature to below 300°C could be used on board in a recently proposed two-stage reservoir [27]. Another very restrictive and debatable DOE requirement for onboard hydrogen storage is the reversibility of solid state hydrides (e.g., onboard refueling with the target rate of 1.5 kgH2 min−1 for 2010 as shown in Table 1.2). However, as pointed out by Sandrock et al. [28] it is an immense problem to remove the exothermic heat at that charging rate. For example, if we charged H2 at 1.5 kg min−1 into a vehicular storage tank based on NaAlH4 (ΔH= 37 kJ mol−1 H2), we would have to remove heat at the rate of 450 kW! This would require very substantial and costly heat exchangers and in practice would completely eliminate the possibility of quick onboard recharging. Another point of importance is an enormous cost attached to building an entire infrastructure of a hydrogen gas refueling station such as steel piping, vessels, etc. Last but not least is the hydrogen embrittlement phenomenon in which various metals, such as high-strength steel, aluminum and titanium alloys become brittle and eventually crack under load following exposure to hydrogen. This process could be a problem for hydrogen steel piping and other components of a hydrogen gas refueling station. In our opinion much easier solution is off-board recharging in which the depleted solid state hydride container is repleted with nanostructured solid state storage materials manufactured in a dedicated plant as schematically shown in Fig. 1.1. According to this new vision, the refueling station is just a retail station for hydrogen storage containers where refueling could become eventually a fully automated process reduced to a quick replacement of containers without any complicated infrastructure. This could also lead to the growth of new businesses for the Hydrogen Economy. Also, an off-board recharging solves the problem of chemical reversibility, which creates a difficulty for a number of hydrides otherwise quite attractive as storage materials. Ball milling can be used to induce a mechanical reversibility carried out in a ball milling device instead of applying high pressure and temperature for the chemical synthesis. Such hydride storage technology based on mechanochemical reactions and nanostructuring processes conducted in high-energy ball mills constitutes the core of this book. Alternative, nonhydride storage, as exemplified in carbon-based nanomaterials, is discussed in the second part of this book.

1.2

Brief, Synchronic History of Development of Hydrides and Nanomaterials

Full container Fuel cell car (Neckar, Daimler Chrysler) or hydrogen internal combustion engine

Container with nanostructured hydrogen storage material * No infrastructure is needed to load H2 in a refuelling station

Hydrogen refuelling station*

7

Depleted container

Depleted container

H2 Off board recharging - Production plant for containers with nanostructured hydrogen storage material (e.g. by ball milling)

Fig. 1.1 A novel vision of refueling/retail station for fuel cell-powered vehicles based on the concept of the off-board recharging of spent hydride fuel and the mechanochemical synthesis of nanostructured hydrogen storage materials in ball-mill plants. Such is not the currently accepted vision of hydrogen economy but one which may turn out to be viable alternative to onboard recharging

1.2

1.2.1

Brief, Synchronic History of Development of Hydrides and Nanomaterials Early Investigations of Metal–Hydrogen Systems and Hydrides

Although not considered to be stored and delivered from metal-hydrogen systems, hydrogen has been used for energy generation since the 1800s. Being a major constituent, up to 50 vol.%, of so-called syngas, synthetic gas manufactured by gasification of coal, wood, or waste, it was widely used in homes in Eastern USA and Europe from circa 1850 until the Second World War. The water gas reaction utilized in production of syngas was the first in the man’s endeavor to generate energy from hydrogen, the most energetic and the most abundant ingredient of our geosphere and biosphere. C + H 2 O (steam ) → CO + H 2

(1.1)

8

1

Introduction

At about same time when the hydrogen-rich syngas had been combusted for electricity generation in gas-fueled electric plants, some studies driven by pure scientific curiosity were conducted, which led to observation of a direct conversion of hydrogen chemical energy into electrical energy. Sir Humphrey Davy (1802) experimented on a cell that utilized a carbon anode in aqueous nitric acid, and Sir William Grove built a gaseous voltaic battery, what is considered the first hydrogen fuel cell. Obviously, flameless burning of hydrogen in such a fuel cell as to generate electric current was not in the ranks with direct combustion of hydrogen in syngas for generation of electric energy in just launched electromagnetic generators. Obviously, this was also long before concerns have become raised about pollution of air and the change to the environment and the climate brought by burning coalbearing fuels for heat, electricity, and mass transportation. Simply, both the economical drivers and the technology to directly convert hydrogen energy into electric energy were not in place. Some 100 years after Cavendish’s discovery of hydrogen, and only 3 years after it was realized that hydrogen sorbed from chemical or electrochemical sources causes blistering and embrittlement to steel vessels, Graham [29] observed the ability of palladium to absorb hydrogen and wrote in Philosophical Transactions of the Royal Society of London: “Hence palladium has taken up a large volume of gas when the temperature of the metal never exceeded 245°C,” and again “1 vol. palladium held 643.3 vols. hydrogen. By the care of my zealous assistant, Mr. W.C. Roberts, the hydrogen employed in these experiments was purified to the highest degree by passing it in succession through alcohol, water, caustic potash, and tubes of 0.7 meter each, filled with broken glass impregnated with nitrate of lead, sulphate of silver, and oil of vitriol. The gas was inodorous, and burned with a barely visible flame.”

The reversible absorption was observed to proceed in presence of either metallic palladium or in palladium–silver alloys; much less hydrogen was absorbed in Cu sponge (1 vol. Cu: 0.6 vol. H.), and not at all in Os–Ir. The reaction of hydrogen with palladium, so being described by Graham, was: Pd + H 2 ↔ PdH x

(1.2)

Palladium hydride is not a stoichiometric chemical compound but simply a metal in which hydrogen is dissolved and stored in solid state, in space between Pd atoms of crystal lattice of the host metal. Relatively high solubility and mobility of H in the FCC (face-centered-cubic) Pd lattice made the Pd–H system one of the most transparent, and hence most studied from microstructural, thermodynamic, and kinetic points of view. Over the century that followed many metal-hydrogen systems were investigated while those studies were driven mostly by scientific curiosity. Researchers were interested in the interaction of hydrogen molecule with metal surfaces adsorption and diffusion into metals. Many reports on absorption of H2 in Ni, Fe, Ni, Co, Cu, Pd, Pt, Rh, Pd–Pt, Pd–Rh, Mo–Fe, Ag–Cu, Au–Cu, Cu–Ni, Cu–Pt, Cu–Sn, and lack of absorption in Ag, Au, Cd, Pb, Sn, Zn came from Sieverts et al. [30–33].

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9

At the time when Sieverts et al. reported first studies on absorption of hydrogen by metals and alloys, a series of seminal papers had been published by Avrami [34–36] on the kinetics of phase transformations driven by nucleation and growth of nuclei and leading to microstructural granulation, viz. grain size distribution in solid material. Therefore, foundations for hydrogen storage and for formation of fine and ultrafine-grained microstructures in solid materials had been developing concurrently. At that time when world emerged from the Second World War, and became falling into another Cold War with the Soviet Union, liquid hydrogen has been intensively studied as a rocket fuel. In contrast to the dangers of a new world deal, and in the climate of overwhelming optimism triggered by the end of war, hydrogen was also contemplated as the fuel of the future for a civil supersonic aircraft flying between Europe and North America, with Canadian Montreal’s supermodern Mirabel airport being considered for accommodating overseas flights from Europe. After World War II, investigations of hydrides were driven mostly by nuclear reactor applications as to understand hydride-caused embrittlement of reactor metals, such as Zr, or to take advantage of high populations of hydrogen atoms in hydrides to scatter, moderate, or shield from energetic neutrons in high-temperature, mobile nuclear reactors. Zirconium alloys in water-cooled reactors, in particular, were known to pick up hydrogen, and precipitate zirconium hydrides on cooling. Therefore, zirconium–hydrogen system was the one studied in close reference to development of nuclear reactors. The first intermetallic hydride reported was ZrNiH3 [37]. On historical milieu we would like to bring attention of our readers to a curious synchronic development of the discipline of hydrides and a new discipline of materials science and technology. Concurrently to the reporting of first intermetallic hydrides came postwar major advances in examination of microstructures in metal alloys and compounds. First, it was an appropriate interpretation of data from X-ray diffraction for crystalline and polycrystalline materials and realization that line broadening in the diffraction pattern can be caused by both small grain sizes and/ or internal strains, while the separation of these two contributions can be achieved using plots for Williamson-Hall method [38]. Then came the first edition of a classic book on X-ray diffraction procedures for polycrystalline and amorphous materials (term nanomaterials was still not coined) [39]. The late 1960s brought the Rietveld method for profile refinement of a whole diffraction pattern, composed of several phases distinct in both chemical and grain size characteristics [40]. A new dislocation theory and a transmission electron microscopy were applied in studies of dislocations and dilatational misfit of zirconium hydrides precipitated in Zr metal matrix [41]. These early metallographic investigations quite often were giving evidence to formation of nanometric precipitates, as metastable hydride phases in metals under nonequilibrium conditions imposed either by high dose of ionizing radiation or fairly rapid cooling (>10 °C s−1). Now, let us come back to our other historical thread: the development of hydrided materials. Large body of results was coming in the 1960s from research on apparently nanometric structures observed in fine metal powders and deposits,

10

1

Introduction

like dispersed Fe, and cathodically charged Ni layers prepared in electrochemical cells [42]. Yet, the thermodynamic information extracted from such electrochemically charged metals concerned inherently nonequilibrium processes, and was considered doubtful and less of value. In fact, an important experiment on direct synthesis of nickel hydride, NiH, at 18 kbar of H2 gas in a high-pressure cell, carried out by Baranowski in Poland’s Institute of Physical Chemistry in Warsaw in 1966 [43] was driven by the desire to achieve true equilibrium conditions between H2 gas and solid nickel hydride phase. At that time a common approach in investigations of metal–hydrogen systems was to conduct studies on well-ordered polycrystalline metals or, if possible, single crystals. Obviously, ordered metals used in experiments let investigators to construct the most transparent and instructive models for a metal-hydrogen system. Although by the end of 1960s many metals were reported to form hydrides, those were either too stable, like ZrH2, or too unstable, as NiH to attract interest in respect to reversible hydrogen sorption.

1.2.2

Early Routes to Nanomaterials

There is a strong interplay between science and technology. New science often creates new technological opportunity, and reciprocally, new technology development often triggers new opportunity to advance science. Such interplay is well illustrated in the tread in which the science of new materials for hydrogen storage has always been strongly intertwined with the parallel development of innovative materials processing techniques and technologies. Initially, those were techniques of highpressure physics and chemistry, and instrumental techniques in electrochemistry. Later development was fueled by an emerging discipline of materials science. The advent of nanomaterials, as one of a few founding blocks of future nanotechnology, came from materials science laboratories around the world. Nanotechnology is a technology that owes its name to the prefix nano, a Greek word for dwarf, as applied to objects that exhibit billionth (10−9) meter dimensions. The first reports on unusual new metallic phases, where observed structures exhibit partitioning into nanosized grains, or are simply grainless, have been coming from groups in the USA and other countries worldwide. A new process of rapid quenching of metals, with the cooling rate of 106 K s−1, first applied by Pol Duvez and his colleagues [44] in the California Institute of Technology in Pasadena in 1960, produced unexpected new fine-grained and grainless, metastable alloy phases, the first nanometals and amorphous metals. Using a quench technique capable of cooling metal melts to ambient temperatures with such extraordinary cooling rates, through just spraying and splashing of milligrams of melt alloys on a chill surface [sic!], the process of nucleation and growth, as described by Avrami 20 years earlier, was kinetically bypassed to yield a configuration of frozen liquid or amorphous metal. Duvez et al. realized this possibility by reporting complete solid solubility in Cu–Ag and Ga–Sb–Ge systems and formation of new nonequilibrium phases in Ag–Ge and

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11

Au–Si – these being the first nanocrystalline alloys and the first amorphous alloys. Cohen and Turnbull [45] were quick to point out that amorphous alloys can exhibit glass-to-crystal transition. Indeed, when processed by low-temperature annealing they formed first metal alloys with nanometric grains, hence, the first nanoalloys. The new science of amorphous and noncrystalline solids would not have developed so rapidly if Chen and Polk had not invented useful amorphous ferromagnetic materials and this technological breakthrough had not been vigorously pursued by the big US company, the Allied Chemical Corporation (later Allied Co). Indeed, in 1972, a new chapter in metallurgical technology has been opened, when a new rapid quenching process of melt spinning, viz., casting of a stream of molten metal on a copper drum, rotating with speeds ranging from 5 to 20 m s−1, was used by Chen and Polk that worked for the Allied Chemical Corporation in USA [46] to spin the first noncrystalline ferrous (and ferromagnetic) metal ribbon, Fe80B20. Further to this a new process to form alloy ribbons with structures ranging from nanometric (then called microcrystalline) to amorphous was invented by several researchers, among them Bedell, Narashiman, and Ray working for corporate research in companies in the USA. It came as a shock to many physicists that amorphous iron alloy can be ferromagnetic in spite of absolute lack of crystalline lattice (and indeed, absence of a theory of amorphous magnetism). With the invention of a new class of soft magnetic materials it became clear that applications could be found that will utilize the outstanding combination of physical and chemical properties as being the direct consequence to the lack of structural crystalline long-range order (LRO), and the presence of short-range order (SRO); the latter being limited to nanometric distances between atoms. Among those new materials were binary Fe–B, Fe–Zr, Cu–Zr, Ni–Zr, Pd–Si, Mg–Zn, and many ternary or even quaternary amorphous systems. Many patents were filed in the USA and Japan, and by many groups worldwide, on amorphous alloys and intermetallics that were termed metallic glasses or glassy metals [47] and had been manufactured by a new rapid solidification technology. This technology became intensively developed by MIT’s Center for Materials Science and Engineering in the USA [48] and Japan’s Tohoku University, among many other materials research laboratories worldwide. From the historical point of view it is important to realize that many if not all of these amorphous alloys, i.e., metallic glasses, when carefully annealed at low temperatures, change to devitrified, nanostructured alloy phases. Those were the first nanomaterials engineered on purpose. They were produced in form of thin ribbons and wires via rapid solidification processing of melt alloys. They allowed controlled exploration of chemical, mechanical, magnetic, and other properties as arising from nanostructures. Shortly, the most prominent benefits of new nanostructures were achieved in the outstanding improvement in the energy product of new neodymium supermagnets produced by crystallization of amorphous Fe–Nd–B alloys (as reviewed in [49]). Shortly after, among research stimulated by the outstanding combination of chemical and magnetic properties in metallic glasses were the first papers on the action of hydrogen on amorphous intermetallics and magnetic metallic glasses. Absorption and diffusion of hydrogen in amorphous alloys was studied in UK by Harris [50] and by Cantor with his colleagues [51], in Germany by

12

1

Introduction

Kirchheim et al. [52, 53], and in the USA by Bowman [54], among few others. In Canada, Wronski et al. [55] studied hydrogenation of F90Zr10 amorphous intermetallic via cathodic polarization, and concluded that hydrogen can be reversibly released by low-temperature annealing at less than 100°C. The hydrogenation was accompanied by a concurrent drastic change of magnetic properties, from the paramagnetic nonhydrogenated phase to the ferromagnetic hydrogenated amorphous solid solution; during those transformations, the F90Zr10 remained amorphous in its bulk, yet Fe magnetic nanocrystals precipitated on the surface layer of the hydrogenated amorphous F90Zr10 [56]. In the USA, Johnson and colleagues in Caltech observed quite the opposite: the first example of a hydrogen-induced crystal-toglass transformation in metallic alloys resulting in the formation of amorphous hydrides [57]: Zr0.75 Rh 0.25 + H 2 → am Zr0.75 Rh 0.25 H1.14

(1.3)

Reaction (1.3) proceeded only at a relatively low temperature just above 150°C and although amorphous alloy exhibited large H/M (hydrogen-to-metal) atom ratio, the charging and discharging of amorphous phase proceeded in a range of pressures, while a crystalline alloy could have been cycled at almost constant pressure [58]. Bowman [59] explained that the strong exothermic reactions between hydrogen and most amorphous intermetallics release excess energy, which causes crystallization of amorphous phase. He also reported that many metallic glasses tend to absorb less hydrogen then their nanocrystalline counterparts; he explained this behavior by the random order of atoms in amorphous phase, which might restrict the number of interstitial sites that are favorable for hydrogen occupancy. While interest in amorphous alloys was becoming vane over the time, the paradigm shift in materials science brought about by the discovery of amorphous metals and rapid solidification cannot be underestimated. The interest of physicists, chemists, materials scientists, and engineers became refocused from well-ordered crystalline materials to disordered and nanocrystalline phases. Shortly, substantial R&D programs were initiated in the USA, Japan, and worldwide both at universities and national government laboratories. Foundations have been established for the advance of the science of nanomaterials. However, the transition from nanoscience to nanotechnology had to come from yet a concurrent innovation in tools used by scientists. This was the invention of the first scanning tunneling microscope (STM) in 1981 [60], followed by the invention of the atomic force microscope (ATM) in 1986 [61]. While STM and ATM microscopes provided powerful tools to advance new science of solids at nanometric region materials scientists were also greatly inspired by papers coming from Gleiter and his group at Saarbrücken University in Germany. Those were the papers that redefined nanocrystalline materials as such where the fraction of atoms located at grain boundaries of polycrystalline material is comparable to the fraction of atoms at the core of grains when the grain size reaches nanometers [62]. A claim was put forward that such materials are to be fundamentally different from, and often superior to, those of the conventional

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13

polycrystals and the amorphous solids. Gleiter observed that nanometer-sized crystalline materials being polycrystals with very small crystallite sizes of about 2–10 nm in diameter are composed of randomly oriented high-angle grain boundaries. The first such nanocrystalline phases produced on purpose came from Saarbrücken about 1984 by evaporation of the material in a high-purity inert gas atmosphere followed by condensation and compaction in ultrahigh vacuum [63]. By this time chemists working in the field of metal catalysts were well aware of importance of nanometric features, such as edge sites and surface clusters on chemical catalysis. Molecular orbital calculations performed 10 years earlier predicted that the ionization potential of nanostructured metals should increase as their grain size decreases [64]. The high surface-to-volume ratio observed in nanostructured materials not only appeared important to the number of active sites in a catalyst, but also seemed to influence the defect chemistry (like oxygen and other anion defects). Siegel [65] computed that the percentage of metal atoms on the surface of grain increases from a few percent in a 100-nm particle to about 90% in a 1-nm crystallite. These simple experiments by Gleiter and equally simple calculation by Siegel were eye opening for others in research community in material science. Immediately, a new consensus reached among materials scientists was that extending our understanding of structure–property relationship in solid materials down to the nanometer regime should be attractive for development of engineering materials with an outstanding combination of properties or novel properties. Among materials that became studied were nanophases produced by mechanical alloying [66] and nanohydrides where the size effect determines hydrogen storage properties [67].

1.2.3

Historical Development of Classical Hydrogen Storage AB5 Alloys

Development of first practical hydrogen storage alloys, AB5-type intermetallics, had its beginning in a typical accidental laboratory experimentation, although it was predetermined by a vigorous development of a new discipline of materials science and engineering in late 1960s. The outstanding hydrogen sorption properties of rare-earth AB5 intermetallics were accidentally discovered in Philips Laboratories in Eindhoven, Netherlands at about 1969 in a program to develop a new permanent magnet alloy [68]. In the work on Sm–Co magnet alloy it was observed that a loss of coercivity occurs, when magnets were aged in humid air, which was related to two concurrent reactions: corrosion of SmCo5 intermetallic with release of hydrogen, and sorption of the released hydrogen by still uncorroded alloy: 2SmCo 5 + 3H 2 O → Sm 2 O3 + 10Co + 3H 2

(1.4)

SmCo 5 + 1.25H 2 ↔ SmCo5 H 2.5

(1.5)

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Introduction

Zijlstra et al. [69] demonstrated that the reaction of (1.5) is reversible at several bars of pressure at ambient temperature. In the following studies on the origin of magnetic coercivity in these new magnets, Phillips laboratories came with the hydride of LaNi5, lanthanum–nickel hydride, which reversibly bonded more than six atoms of H per one formula unit (H/M ratio > 1) [70]. LaNi 5 + 3.35H 2 ↔ LaNi 5 H 6.7

(1.6)

Reversibility of hydrogen charging was very good at room temperature. This is because the hexagonal lattice of the metal host, like this of LaNi5, does not undergo major transformation as hydrogen is inserted interstitially. This was the first AB5type interstitial hydride in which hydrogen is stored between metal atoms. Those rare-earth AB5-type hydrides were quickly utilized in rechargeable nickel metal hydride batteries where electrochemical hydrogen charging and discharging take place at ambient temperature. Such electrochemical hydrogen storage is reversible, when the negative hydride electrode (anode) is combined with the positive Ni electrode (cathode) in the battery cell.. AB5 H x + xOH - ↔ AB5 + xH 2 O + xe -

(1.7)

NiOOH + H 2 O + e - ↔ Ni (OH )2 + OH -

(1.8)

As a matter of fact, the first hydrides with practical hydrogen storage capacities were realized in rechargeable nickel metal hydride batteries. For more information on electrochemical hydrogen storage in rechargeable batteries a reader can be referred to several recent reviews on this subject [71–73]. Renewed interest in solid-state hydrogen storage came with first Oil Crisis in 1970s. By early 1973 a patent was granted to US Brookhaven National Laboratory, where Reilly had produced first MmNi5 alloy for hydrogen storage [74]. MmNi5, where Mm – mischmetal – is a unrefined (cerium free) mixture of rare-earth metals, mostly La and Nd, obtained from Bastanite from California (also found in large deposits in Inner Mongolia, China), was much less expensive than La and has already been used as a deoxidizer in steel industry. Soon after, the potential for energy-related application for nickel and nickel alloys attracted the major Ni metal and alloy producer in Canada and the USA, INCO. Garry Sandrock, then with INCO, achieved optimization of MmNi5 compositions for gas hydrogen storage through careful partial substitutions for A and B in AB5 formula [75]. The compositional changes have reduced the raw materials costs to about 30% of that for the LaNi5 alloy. LaNi5 has an attractive value of the equilibrium hydrogen desorption pressure dropping in the range between 1 and 2 atm at 25°C. In the MmNi5 this pressure rises to a much less attractive 30 atm, which can imply difficult engineering design for gas hydrogen storage, and was completely impractical in MmNi5 used for the anode in a NiMH cell; obviously, battery cells operate at near-atmospheric pressure. The substitutions of MmNi5 with Ca for Mm and Al for Ni were quite effective in lowering the equilibrium pressure of hydrogen back to about 1–2 atm at

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15

25°C, which was desired for easy hydrogen charging and discharging at ambient conditions. Approaches along the same line came in the same time from Japan [76]. The alloying of polycrystalline AB5 alloy was an excellent (and very practical) result of new materials design involving adjustment of the strength of chemical bonds between metallic elements in AB5 hexagonal lattice and its interstitial hydrogen atoms. The bond strength in the metallic host lattice was further optimized with cosubstitutions, which was effective in reducing the volume change on hydriding of the hexagonal AB5 unit [77]; the density of material can, indeed, be related to the bond lengths through an approximate formula: Ro = 0.0115Aav/a2c, where Aav is an average atomic weight for the AB5 formula unit, and a and c are the hexagonal lattice parameters in nanometers [78]. The lower dilatation of the AB5 cage on placing hydrogen in it was equivalent to better cohesiveness of atomic structure in this intermetallic, hence improved resistance to alloy decrepitation on repetitive (cyclic) charging and discharging, hence better corrosion resistance of MmNi5 alloys in an aggressive caustic battery electrolyte. Although substitution with Al came at the price of the lowered hydrogen capacity, another Mn substitution ameliorated this problem [79]. With compliance with the well-observed 10–15-year shift from new material development to its full commercialization, the Al and Mn substitutions have been combined in multielement MmNi(Co, Mn, Al)5 battery alloy and optimized for the use in anodes in NiMH commercial batteries. According to recent review a typical commercial AB5 alloy can consist of many elements, such as in hydroalloy F from GfE Metallen und Materialen GmbH: La0.64Nd0.36Ni0.95Cr0.19Mn0.41Co0.15 (in wt%) [71]. Again, one can observe how the development of materials for gas hydrogen storage for fuel cells has been intertwined with a parallel development of alloys for electrochemical hydrogen storage in rechargeable battery cells. AB5 alloys used in either hydride battery cell or in a hydrogen storage tank for fuel cell are not per se nanocrystalline but the microstructural design process, which led to their development in 1980s and 1990s, reached nanoscale phase partitioning.

1.2.4

Historical Development of Interstitial Hydrides in Other Intermetallic Systems

The first systematic research into gas hydrogen storage for practical applications began in the early 1970s at the Brookhaven National Laboratory (BNL) on New York’s Long Island. The investigations were conducted in a program that explored possibility of storing electric energy produced in off-peak hours by nuclear power stations through generation of hydrogen, and then the storage of hydrogen in metals instead of storing electric current. So, at the same time as the AB5-group hydrides were intensively investigated in the BNL, new AB-type hydrides came into view. Iron–Titanium, FeTi, hydride was discovered in 1974 by Reilly and Wiswall [80] in the same Brookhaven National Laboratory. This metal hydride formed from a 50–50 atom ratio of titanium and iron, the lowest cost and the most convenient

16

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Introduction

hydride, offered a reversible capacity of almost 1.5 wt% H2 operating at ambient temperatures. A desirable feature of this new hydride was that the hydriding– dehydriding reaction occurs near ambient temperature and during the charging pressure ranges between 500 and 15 psi. The prototype 400-kg FeTi hydride storage unit was built by a New Jersey’s electric utility. The charge–discharge reactions were observed to be very rapid but strongly limited by the rate of heat transfer. FeTi material was not losing hydrogen storage capacity even after thousands of cycles when high-purity hydrogen was used. Yet, the alloy was very sensitive to O2 and CO poisons, and even to traces of H2O. Obviously, while meeting the expectations of stationary storage of nuclear energy in the form of solid hydride it was too heavy for storage of hydrogen for transportation energy. Anyway, the needs for cars propelled by hydrogen fuel were not in place at the end of the 1970s.

1.2.5

Historical Development of Nanophase AB2 Intermetallic Hydrides

A new venue for hydride research has been evolving from the study of multielement Ti–Ni alloy thin films for hydrogen oxidation in fuel cells conducted by Stanford Ovshinsky in the USA in the early 1980s. The new disordered multielement, multiphase AB2-type alloys for hydrogen storage exhibited a large degree of structural disorder and submicron partitioning of phases. Ovshinky has been pioneering a view [81] that structural disorder at nanoscale results in outstanding combination of hydrogen storage in crystalline grains, while the disordered grain-boundary phase provides channels for hydrogen delivery and release from nanocrystalline grains. A first, quintessential tool of the coming nanotechnology, a scanning tunneling microscope, was then used to reveal that in the TiVZrCrNi multiphase alloy the hydrogen is preferentially sorbed in the Ti-rich phase, while V-rich phase was likely to corrode preferentially to expose Ni nanoparticles [82]. This shift of view was congruent with enormous intellectual potential, which was brought about by research on amorphous and rapidly quenched phases (amorphous selenium, chalcogenide glasses, and metallic glasses) in the 1960s and 1970s. US Ovonics patents came just in time when interest has been growing in Europe, Japan, and North America in chemical properties of amorphous metals and metallic glasses prepared by rapid solidification technology. The materials challenge was to optimize hydrogen capacity and corrosion properties in multiphase nanostructures for the anode in nickel metal hydride (NiMH) cell. Shortly, the outcome of international race began toward development of a second generation of multiphase and multielement NiVZrTiCr hydrides. The nanostructured battery materials were developed applying new principles of materials science established in previous decade and reflecting on the role of disorder and nonequilibrium processing in overcoming the conflicting requirements for materials properties. In the early 1990s sealed nickel metal hydride batteries became available in Japan and elsewhere for consumer goods. Shortly afterward, the new nanopowders, which were proven in electrochemical hydrogen storage, began to be investigated for gas hydrogen storage. This came when new

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drivers emerged for growth of hydrogen energy in the 1990s. The concern now was for clean environment, and mitigation of greenhouse gases, and appeared in the context of evidence for accelerating global climate change.

1.2.6

New Routes to Nanomaterials: Mechanical Alloying and Mechanochemical Activation

Like the discovery of nonequilibrium quenching of metal melts in the 1960s and 1970s led to the development of disordered multiphase AB2 hydrides in the 1980s, an interest in nonequilibrium milling of metal powders, in the 1980s, has been setting foundations for the development of the nanohydrides of the following decade. The nonequilibrium synthesis of new metal alloys by high-energy ball milling of powders was developed almost 20 years earlier by John Benjamin and his coworkers at the International Nickel Company (INCO) [83]. The development came from research on new alloys intended for gas turbines, and was an outcome of an effort to produce nickel superalloys, which were mechanically hardened through an inoculation with fine refractory metal oxides, Al2O3, Y2O3, and ThO2, in so-called oxide dispersion-strengthened (ODS) process. In his research Benjamin pointed out that mechanical alloying (MA), a new process induced by intensive milling of metal and compound powders, could make new materials with unique properties [84]. (In a historical note like this, the credit for coining the term mechanical alloying must be given to a Patent Attorney for INCO company.) However, it was not until the early 1980s that the discovery by Koch and coworkers [85] of amorphous alloys made by MA renewed interest in intermetallic compounds, which were difficult to synthesize by orthodox metallurgy. Those authors demonstrated that, during mechanical alloying of elemental metal powders of nickel and niobium in a special high-energy ball mill (SPEX™ model), the final product becomes an amorphous alloy. This was demonstrated by the complete disappearance of the sharp Bragg reflections in X-ray diffraction pattern, which are characteristic of any crystalline phase, and appearance of a broad maximum, characteristic of amorphous alloys; the latter was previously seen in metals that were rapidly quenched from melt. The amorphous structure was confirmed in the exothermic glass-to-crystal transformation peak revealed by differential scanning calorimetry (DSC). Two years before the publication by Koch, Yermakov et al. [86] reported on the amorphization of a number of Y–Co intermetallic compounds by grinding of ingots, which were prealloyed by casting. This process was later termed mechanical milling (MM) process. Further early investigations were carried out by Schwarz et al. in the US Los Alamos National Laboratories [87], as well as by groups in Germany and the Netherlands. Already, in these early stages of new research, it was well recognized that by ball milling the average size of grains in powder particulates was reduced to nanometric dimensions. Either MA or MM processes drop under more general class of solid-state amorphization reactions, SSAR. Amorphization by irradiation of solids was observed yet in the era of study of materials for nuclear reactors. In 1962, Bloch [88] amorphized U6Fe by exposing it to fluxes of nuclear fission fragments. Others observed amorphization

18

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Introduction

of intermetallic compounds by high-voltage electrons [89]. Schwarz and Johnson [90] demonstrated the phenomenon and a new process of amorphization by diffusion between thin-film sandwiches of crystalline La and Au. In the late 1970s, Johnson and coworkers [57] in Caltech demonstrated that in addition to techniques of amorphization by rapid solidification of molten alloys and amorphization by mechanical alloying of solid alloys, the process can be conducted by repetitive exposure of an alloy to hydrogen pressure and vacuum. This provided a new solid-state processing route to amorphization by hydrogenation. To this class also belong reactions of hydrogen decrepitation of metals and alloys, the method proposed by Harris [50] for preparation of fine powders for subsequent sintering into new generation of neodymium supermagnets. Through cycling of a cast metal ingot of Fe–Nd–B in hydrogen the ingot was effectively changed into fine powder; however, nanometric hydride phases that must have formed concurrently to the decrepitation were not considered for hydrogen storage. So it took another thread in materials chemistry history tracks to bring mechanical processing for advancing research on hydrogen sorption in solid media. This thread can be attributed to studies of effects of mechanical grinding on activation of inorganic solids for chemical reactions. As a matter of fact, the approach to apply mechanical energy, in form of grinding and milling, as to activate or bring about new paths for chemical reactions was much older. In a historical review on mechanochemistry, as a branch of solid state chemistry dealing with mechanically activated chemical reactions, pioneering role is attributed to photochemist M. Carey-Lea who as early as 1892 [91] began publishing reports of systematic studies on decomposition of silver and mercury halides, then dozens of other stable compounds, under applied mechanical pressure, using so simple a method as grinding compounds in a mortar. Extracting mercury from its sulfide while rubbing it in a copper mortar was known yet in ancient Greece [92]. However, in modern times it was a large body of work on mechanochemical activation of inorganic compounds that was coming from laboratories in Russia and the USA that demonstrated the role of defects in activation of materials for chemical reactions conducted in a solid state or at the interface of solids and gases. Researchers from the US Naval Civil Engineering Laboratory engineered mechanical alloys as composites where the second phase of Fe (or Ni, Ti, or Cu) has been dispersed at nanometric manner in the Mg phase; these alloys were intended as supercorroding alloys to react rapidly with seawater to produce heat and hydrogen gas in underwater welding of naval vessels [93]. Shortly after, in years 1984–1997, Ivanov et al. [94] from Russia were the first to explore mechanical alloying to obtain magnesium alloys for hydrogen storage.

1.2.7

Historical Development of Lightweight Metal Hydrides and Hydride Complexes

Formation of magnesium and calcium hydride compounds was first recognized by German chemists yet at the end of nineteenth century [95] but it took more than a half a century before a substantial yield of Mg hydride was obtained from direct

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19

synthesis of the elements and the first determination of values for the decomposition pressure on temperature, the enthalpy of formation, and the activation energy were obtained [96]. Since then, Mg metal with its hexagonal lattice was well researched for hydrogen storage. MgH 2 → Mg + H 2

(1.9)

Magnesium dihydride can store several times as much hydrogen per unit weight as AB-type TiFe hydride. However, the first and the most limiting problem with lightweight Mg alloys for hydrogen storage was always the very poor kinetics of hydrogenation. Indeed, metallic magnesium does not react with hydrogen at low temperatures, and reacts only very slowly at high temperatures. Because of the relative inertness with respect to hydrogen, Mg was even recommended as a structural metal suitable for contact with hydrogen up to the melting temperature [sic!]. Not surprisingly, many controversial and poor reproducing data have been published on the hydrogenation/dehydrogenation kinetics. The natural inertness of Mg to molecular hydrogen can be attributed to well-known factors: (1) thick surface oxides that cover magnesium metal do not provide sites catalytically active for dissociation of hydrogen molecule, which is known to occur on contact with a metallic surface (clean metallic surface often provides active sites for catalytic breaking of very strong H–H bond), (2) surface-adsorbed impurity gases further slow the kinetics of hydrogen absorption into the bulk grains (the process that may be related to the hydrolysis of Mg with water vapor and causing even more thick oxide/hydroxide surface barriers; Sect. 2.1.5). On the other hand hydrogen dissolved in magnesium metal is readily removed by heating it in vacuum up to 300–400°C. The stable Mg dihydride was prepared as early as 1951 by direct heating of Mg in gaseous hydrogen under high pressures; however, the reaction was slow and difficult to complete [97]. While the chemical reactivity of Mg metal to hydrogen is poor it was early realized in works by Dymova et al. [98] that its dihydride can be a strong chemical reductor, and its increased reactivity depends strongly on small grain size. So the effect of grain size was pointed on as being one of high importance. Mg2Ni alloys and magnesium–nickel A2B-type complex hydrides, Mg2NiH4 were well known, being first developed in the US Brookhaven National Research Laboratory [80], and studied in Nordic countries [99], where the interest in Mg has been stimulated by availability of hydropower required to produce Mg metal in the first place. (Even earlier, Pebler studied A2B-type hydride of Zr2Ni [100].) The new Mg2Ni alloy alleviated the problem inherent to Mg metal: very poor kinetics of hydrogenation reaction. Reilly and Wiswall [80] noticed that by the presence of intermetallic magnesium compounds, Mg2Ni and Mg2Cu, on the Mg surface this reaction can be greatly accelerated. The hydrogen release can proceed without major crystallographic transformation as the hydride is converted back to the Mg2Ni: Mg 2 NiH 4 ↔ Mg 2 Ni + 2H 2

(1.10)

20

1

Introduction

The next alkali earth in the periodic table is Ca and its ternary hydride was found to be less stable than the Mg ternary hydride. Indeed, the desired reaction for CaNi5 is: CaNi 5 + 3H 2 ↔ CaNi 5 H 6

(1.11)

and this reaction was observed to work repeatedly and reversibly at near room temperature. Yet, the following disproportionation reaction is a preferred reaction path as long as the temperature is slightly increased (to 200°C) and the diffusion of the metal atoms becomes significant in comparison to the diffusion of smaller H atoms. The reaction product observed in the disproportionation was the mixture of more stable dihydride and metallic nickel: CaNi 5 + H 2 → CaH 2 + 5Ni

(1.12)

Although in later studies most of the intermetallic compounds were found to be thermodynamically unstable and falling into the disproportionation reaction path easily their actual tendency to do so varied markedly from system to system, as later it was learnt from studies on transition metal ternary hydrides. Many of these hydrogen-storing hydrides became available by the end of 1970s, most of them under the trademark of commercial HY-STOR hydrogen storage alloys: 300 Series (Mg-base), 200 Series (Ni-base, including both CaNi5-type and MmNi5 type), and 100 Series (FeTi-based). The first and the most important alloy, for which the H storage properties clearly benefited from nanostructure and nanoprocessing, was Mg2Ni, an A2B-type intermetallic. Although yet in 1961 Dymova et al. [98] pointed to the small grain size as a primary factor in improving hydrogen sorption in Mg, and other Russians, Ivanov et al. [94], demonstrated the high-energy ball milling route to hydrogen storage in magnesium, it took another four decades and series of studies reported from the McGill University in Canada to bring again the benefits of fine grains to the attention of the community of researchers in hydrogen storage. The clear demonstration of the effect of nanostructuring came in works of Zaluskis’ [101]. A study of ball milled powders showed that when the smallest X-ray diffracting metal grains in milled particles reach the range 5–50 nm the kinetics of both absorption and desorption is improved by an order of magnitude. These were nanostructured hydrides or nanohydrides. It must, however, be pointed out that they completely ignored the effects of simultaneous decrease of particle size during milling (Sect. 2.1.3).

1.2.8

Early Studies of Noninterstitial Transition Metal Ternary Hydrides

With the beginning of the 1990s the interest in hydrides was somehow refocused. Firstly, interest evolved in other hydrides of Mg with transition metals, beyond already well-researched Mg2Ni stoichiometry. Mg and Fe do not alloy to form

1.2

Brief, Synchronic History of Development of Hydrides and Nanomaterials

21

Mg2Fe binary intermetallic by either ingot casting or metal powder sintering procedures; however, when sintering is done in hydrogen, a very stable Mg2FeH6, noninterstitial, ternary magnesium–iron hydride, forms [102]. The Fe and Co transition metal hydrides (in reality they are complex hydrides) with Mg were firstly synthesized in the University of Geneva in 1991 in the same group that observed formation of this hydride during sintering of powders. While presenting a formal derivative of A2B stoichiometry, the hydrides were, in essence, Fe–metal and Co–metal hydride anionic complexes [103]. The new anionic complexes of Fe and Co were stable rendering the desorption temperature too high. When desorbed, the hydrides could not change to the hydrogen-free A2B alloys because Mg2Fe and Mg2Co do not exist in the binary Mg–Fe and Mg–Co systems (e.g., iron and magnesium are immiscible in the solid state and form a large miscibility gap in the liquid state). Indeed, the following reaction (1.13) is in striking contrast to reactions (1.2–1.4) and (1.5–1.6) where the dehydrogenated phase was converted back to stable host alloy (LaNi5 or Mg2Ni), and indeed similar to the disproportionation reaction (1.12) Mg2 FeH 6 → 2MgH 2 + Fe + H 2

(1.13)

One can notice that in (1.13) the Fe3+ in the (FeH6) complex is reduced to metallic Fe, the process that requires a major rearrangement of atoms. Therefore it is not surprising that the kinetics of desorption was poor while being controlled by diffusion of the metal atoms. Although more research had been coming, particularly from groups in the University of Geneva and in Stockholm, research on Mg2FeH6 and similar transitionmetal complex hydrides was put aside for some time. This, perhaps, was unjustified as some of those ternary hydrides formed with alkali earths and transition metal complexes, viz., BaReH6 [104], exhibited interesting storage properties: desorption below 100°C and 2.7 wt% H2 capacity at very high volumetric density of 134 g H2 L−1. Although too expensive for storage in cars, such hydrides would find a niche in fuel cell power sources in portable electronics, if only the desorption temperature is lowered. However, the needs were not in place, at the time of great success of hydrides in rechargeable NiMH. A stable aluminum hydride, AlH3, has been known to exist since some time. It has been prepared by various techniques including the direct reaction of metallic Al with atomic hydrogen by Siegel [105]. AlH3 is soluble in ether when freshly synthesized but polymerizes while aged and forms insoluble precipitate whose structure depends on the degree of polymerization of –(AlH3)– monomer.

1.2.9

Toward Chemical/Complex Hydrides

With the coming of a New Millennium, and in the dramatic context of the September 11, 2001 terrorist attacks other drivers for hydride research became evident: energy security as to secure uninterrupted supply of energy for still unsustainable world

22

1

Introduction

economies, and ever-growing needs for fuels for mass transportation. Also, the rapid economic growth in China and Asia put big pressure on markets for fossil fuels. USA Congress and President issued a “Great Challenge” in respect to development of hydrogen storage for future fuel cell cars. The interest has been renewed for storage of hydrogen in chemical media other than metal and alloy powders. So far, there are about 170 hydrides listed on the website hydpark.ca.sandia.gov. Of particular interest became complex hydrides with participation of alkaline metal ions and metal complexes and in mixed ionic and covalent coordination compounds. This class of hydrides has become known as chemical or complex hydrides. Among them are nontransition metal sodium and lithium alanates and borohydrides. The two most investigated are sodium aluminum hydride (sodium alanate), NaAlH4, and sodium borohydride, NaBH4. The former, alanate, was synthesized almost two decades earlier in Russia by Dymova et al. [106]. Those complex hydrides were not thought to be suitable for solid-state hydrogen storage media as they release hydrogen irreversibly at conditions that are outside the temperature– pressure range desirable for fuel cell propelled cars. In particular, their hydriding and dehydriding characteristics have been creating significant challenges. This opinion has been changed with a breakthrough work by Bogdanovic´ and Schwickardi [107], who discovered that the addition of Ti compound (chloride) to the complex hydride NaAlH4 causes reversible hydrogen storage in the range up to 3.7 wt% H2 under moderate conditions of temperature and pressure. However, the reaction has long been known to proceed in two steps [108] NaA1H 4 ↔ 1 / 3Na 3 A1H 6 + 2 / 3H 2 → NaH + Al + 3 / 2H 2

(1.14)

and again, like in the reaction of (1.13), the completely discharged phase did not form an alloy, NaAl, from the original metals; instead alanate is reduced to metallic Al, in fact nanoaluminum, that does not easily recharge with H2. In 1997 a study conducted by Ford Corp. (USA) compared the weights and volumes of various hydrogen storage media in the context of their use for a lightweight four- to five-passenger fuel-cell car with a 500-km range per one hydrogen charge [109]. Similar studies were solicited by other major car manufacturers, and government organizations in several G7 countries. The interest was raised by the theoretical electrical conversion efficiency for an ideal hydrogen–oxygen fuel cell being an impressive 83%, and up to 60% of the hydrogen chemical energy converted to electric energy (electric current). This compared very favorably with circa 45% achieved in hydrogen internal combustion engines. Since then, search for new reversible lightweight complex hydrides was heating up. Interest aroused in other complexes of hydrogen with Al and B that have been known since 1950s. Alkali metal and magnesium tetrahydroborides were synthesized [110, 111] and investigated by Germans and Russians for thermal decomposition from the late 1950s to the early 1970s. Lithium tetrahydroboride, LiBH4, which has already been known as a strong reducing agent in organic synthesis, has received renewed attention after Züttel [112] reported the onset of hydrogen desorption at approximately 200°C promoted by SiO2 admixed to this borohydrides.

1.2

Brief, Synchronic History of Development of Hydrides and Nanomaterials

23

Among the alanates, the sodium alanate and the lithium alanates have been well studied. Magnesium aluminum hydride (magnesium alanate), Mg(AlH4)2, with the total 9.3 wt% H2, has been known since its synthesis by Wiberg in the 1950s; however, there have been difficulties with good reaction yield to produce this hydride until more than 50 years later Fichtner and colleagues [113] produced a sufficiently pure compound for structural investigations. Again, the reversibility of the reaction was questionable as the experiments so far demonstrated that after discharge of hydrogen the alanate yields Al phase, not unlike sodium alanate in the reaction of (1.14). Nevertheless, the intensive research on complex hydrides has established itself solidly in place in early years of the new Millennium.

1.2.10

Historical Development of Nanocarbons and Carbon Nanotubes

Concurrent stream of the development of nanomaterials for solid-state hydrogen storage comes from century-old studies of porous materials for absorption of gasses, among them porous carbon phases, better known as activated carbon. Absorption of gases in those materials follows different principles from just discussed absorption in metals. Instead of chemisorption of H2 gas into the crystalline structure of metals, it undergoes physisorption on crystalline surfaces and in the porous structure formed by crystals. The gases have also been known to be physisorbed on fine carbon fibers. Reports on growth of fine filaments of carbon have been coming since 1890, when filamentous carbon was obtained by passing cyanogens over red-hot porcelain. Hollow carbon fibers were reported in a Russian journal in 1952 [114]. An interesting view point about the discovery of carbon-nanotube-like materials can be found in Guest Editorial to the journal Carbon [115]. Fine fibers showing graphitic structures were frequently encountered as unwanted phases in both steel industry and in industrial catalysis, where they deposited on walls of metallurgical furnaces and chemical reactors. Since then development of carbon filaments was driven by their use as reinforcement phase in lightweight composite materials for space and aerospace industries. Research focused on fabrication of carbon fibers from polymerbased precursors such as rayon, polyacrylonitrile (PAN), or mesophase pitch. In the 1970s Oberlin et al. [116] published images of what can be now called single- or double-wall carbon nanotubes (SWNT, DWNT), although at this time, images obtained from transmission electron microscopes cannot clearly reveal the number of concentric hollow tubes in the fiber. In the 1980s he produced hollow carbon filaments by chemical vapor deposition (CVD) in a floating catalyst reactor where they were nucleated and grown on 10-nm metal catalyst nanoparticles. Such CVD method to produce fine hollow filaments became with time the most versatile production method for carbon nanotubes (CNT). At the time, it was mechanical property that was of primary interest, as it was realized early that well bulk carbon fibers resist crack propagation. Therefore, little attention was paid toward the

24

1

Introduction

Oberlin’s small and hollow objects, which only latter were termed carbon nanotubes, and were becoming studied around the world for other than mechanical, mostly electronic, and later hydrogen storage properties. The situation changed diametrically in 1991 when Iijima [117] published a breakthrough paper entitled “Helical microtubules of graphitic carbon.” Iijima’s rediscovery of CNTs came during the research conducted to follow on a breakthrough research on another carbon nanophase, fullerene, and spherical C60 carbon atom configuration, which was discovered 6 years earlier by Kroto et al. [118] and published in a seminal paper on C-60 buckminsterfullerene. The latter molecule consisted of 60 carbons that satisfied a mathematical theorem by Leonhard Euler requiring a spherical surface, as to be entirely built up from pentagons and hexagons, the molecule must have exactly 12 pentagons like the geodesic domes of the architect R. Buckminster Fuller, and not unlike a soccer ball. Shortly after, the now clear CNT images were observed by powerful high-resolution transmission electron microscope (HRTEM) in soot produced by the just published Kratschmer-Huffman method on the laboratory scale production of C60 [119]. CNTs can be seen as single, one atom-thick, graphitic planes, viz. graphene planes, rolled to form nanoscale cylinders, which are often closed at the end by half a sphere of the earlier discovered fullerene. These structures, while comprising of only one single-wall cylinder were termed single-walled nanotubes, SWNTs, whereas those consisting of two or more concentric graphene cylinders became known as multiwalled nanotubes (MWNTs). Depending of how the graphene sheet is rolled up, the three forms of CNTs, zigzag, armchair, and chiral, change their electronic properties from metallic, to superconducting, to insulating. In 1992 it was shown that MWNTs can be produced in DC arc discharge in He with use of two graphite electrodes [120], and in the following year the SWNTs were also produced in arc-discharge using alloy catalysts (like Fe–Co) [121–123]. Recently, unrolled single-layer graphene phase was also prepared by Novoselov et al. [124]. In 1997 Dillon and coworkers [125] reported the first experimental result of high hydrogen uptake by carbon nanotubes; their estimate was 5–10 wt% H2. This was followed by hot race to claim by many researchers ever-larger potency of carbon nanotubes to absorb and retain hydrogen. The most known became the claim by Baker et al. [126] that some nanofibers can absorb over 40–65 wt% H2. However, these experimental results and their interpretation were strongly criticized and have not yet been able to be reproduced. More careful studies, that follow, have shown that only 0.7–1.5 wt% of hydrogen gas is adsorbed in nanofibers under ambient temperature and pressures slightly above 100 bar [127]. Since then, the area of nanocarbons and carbon nanotubes for hydrogen storage has seen continuing interest from many research groups worldwide. Carbon nanotubes, as graphene and graphite, are highly ordered carbon phases. However, a separate line can be drawn for historical development of disordered carbon phases; among them is an amorphous carbon (am-C) . In it, strong bonding between carbons did not allow for completely chaotic distribution of carbon atoms in solid-state phase. Instead, amorphous carbon exhibits random distribution of three possible coordinations of carbon atoms in a planar sp2, tetrahedral sp3, and

1.2

Brief, Synchronic History of Development of Hydrides and Nanomaterials

25

even linear sp1 configurations of electronic orbitals. Some of these phases were shown to be greatly stabilized by hydrogenation. Hydrogenated amorphous carbon has been observed to contain as much as 40–60 at.% H2 in a structure consisting of more than 70% tetrahedrally bonded atoms in sp3 configuration [128]. At lower hydrogen content at 25–30 at.% H2 the sp3 fraction is high. At 70 at.% H2, the hydrogenated tetrahedral amorphous carbon, known better as diamond-like carbon (DLC), has been developed for many very useful applications [128]. These amorphous phases have only been produced as thin films in deposition from high-density plasma. Recently, a new spherical graphitic nanophase, termed carbon nanoshell, was found by Wronski and Carpenter [129] in carbon residues obtained by acid leaching of Ni from commercial carbonyl nickel powders produced in century-old CVD carbonyl nickel process in INCO refineries in Wales, UK and Canada. However, these and other new nanocarbons are still waiting to be investigated for hydrogen storage. As for the year 2007, the recent Viewpoint Paper by Chahine and Bénard [130] states clearly that these new nanocarbons and nanostructures that store hydrogen by physisorption will still require development to bring about qualitative changes; means are still waiting for looking at them from a new angle.

1.2.11

New Materials and Techniques

In the past several years, some old compounds were put into perspective from a new angle keeping still in mind that to meet hydrogen storage targets in hydrogen storage for vehicular applications such storage must be reversible in near-ambient temperatures (less than 100°C) and pressure range. One of the systems that gained great interest was lithium nitride. Li3N has been known since the 1910 study by Ruff in Germany to compound with hydrogen, and release it under strong heating. However, it was only in recent investigations by Ping Chen and coworkers in National University of Singapore [131] that the Li–N–H system was reassessed and demonstrated hydrogen absorption in as low temperature as 100°C, and desorption that is rapid, yielding as much as 9.3 wt% H2 in half an hour when temperature is raised to 255°C. What attracted wide interest in this research was that about 6.3 wt% H2 was desorbed below 200°C, and the desorption reaction follows the reverse of absorption: Li3 N + 2H 2 ↔ Li2 NH + LiH + H 2 ↔ LiNH 2 + 2LiH

(1.15)

However, the desorption proceeds under 10−5 bar vacuum, and traces of ammonia that is toxic for fuel cells were detected. While lithium nitride was an old material revisited, zeolites are well-researched materials for their catalytic and physisorption properties. These microporous inorganic frameworks were known to trap hydrogen; however, their potential for hydrogen storage was not great as they show only moderate storage capacities. On the other hand, a new microporous metal–organic frameworks (MOFs) ignited return of interest in hydrogen storage based on physisorption rather then chemisorption.

26

1

Introduction

In 2003, Rossi et al. [132] reported a MOF material with high hydrogen sorption capacity, 4.5 wt% H2 at cryogenic temperature of 78 K. The promising hydrogen sorption properties of MOF materials were discovered only 4 years from the time the first synthesis of this class of materials was performed [133]. Much more modest, 1 wt% H2 storage, yet at room temperature and 20-bar pressure was observed. This started another vigorous research in materials, in which scaffolding-like nature leads to extraordinarily high surface area of 2,500–4,000 m2 g−1, i.e., more than four times the specific surface area of zeolites. In the following research the capacity of systems such as IRMOF-6 and -8 in room temperature exceeded that of carbon nanotubes in cryogenic temperatures. Since then, the area of mesoporous hydrogen storage media, where hydrogen is sorbed in very-high porosity materials, has been experiencing great interest from the community of researchers in hydrogen storage. The intertwining of the development of nanomaterials and methods of nanoprocessing is still ongoing. While development of new MOF-type materials is an example of a quintessential bottom–up nanotechnology, an opposite route, top–down nanotechnology approach, realized through reduction in the grain and particle size has also been furthered. New techniques have become developed, which have potential to take materials research beyond metals and carbons, and beyond the limitations imposed by the choice between Scylla and Charybdis of a strong chemisorption in nanohydrides and a weak physisorption in nanocarbons. New explored materials are nanocomposites or hybrid hydrogen storage media where chemical bonds are optimized for reversible storage of hydrogen in solids. One of these new techniques having potential for preparation of such materials is the controlled reactive hydrogen ball milling/alloying (CRMM/MA) of metals, nonmetallic elements, and compounds in hydrogen-filled ball mills. This approach to preparation of new hydrides and hydride nanocomposites has been developed at the University of Waterloo and CANMET’s government laboratories in Ottawa, Canada. The milling is conducted in specialty, high-energy ball mills, where trajectories of balls are well controlled, and H2 gas can be supplied and absorbed during milling. The complex Mg2FeH6 hydride was prepared in this way by Varin et al. [134] in the direct, mechanically driven synthesis, viz. direct mechanosynthesis reaction by reacting inexpensive elemental Mg and Fe metals: 2Mg+Fe+3H 2 → Mg2 FeH 6

(1.16)

The reaction of (1.16) follows reverse path to the thermal decomposition reaction of (1.13) and proceeds at room temperature and only slight overpressure of hydrogen supply. This presents a new mechanical activation route to manufacturing of nanomaterials for hydrogen storage. This brief history of century-old investigations toward hydrogen interaction with solid materials and nanomaterials brings us to the current state of affairs when the hydrogen storage for fuel cell systems still remains to be solved. Indeed, in the first decade of the new Millennium, and at the advent of the Hydrogen Economy, fuel cell stacks for use in mass transportation, like those developed by Ballard Power Systems based in Canada, are ready for mass commercialization. Also, hydrogen

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

27

production systems, as those widely used in the chemical and fossil-fuel (oil)refining industries are ready to supply hydrogen fuel to fuel cell cars. Sources of hydrogen are numerous. More than 90% of hydrogen is today produced through thermochemical reforming of methane gas, CH4, at high-temperature (800–1,700°C) reaction with steam or oxygen [(1.17) and (1.18)] to make the syngas…not unlike the gas produced by gasification of coal in the nineteenth century (e.g., (1.1)). CH 4 + H 2 O → CO + H 2 (reforming reaction)

(1.17)

CO + H 2 O → CO2 + H 2 (gas shift reaction)

(1.18)

Alternative production of clean hydrogen is considered in the context of new nuclear reactor technologies, like Generation IV reactors, and as a complement to nuclear electric energy systems in which hydrogen is to be produced electrolytically from off-peak nuclear power. Research, development, and demonstration of production of clean hydrogen from water via innovative photochemical water splitting, and from renewable energy sources, such as wind and solar energy exhibit surge of activity. Yet, at the same end to first decade of Millennium, time-efficient and safe hydrogen storage for use in mass transportation is still unsolved. The hydrogen challenge is waiting solution. Can advances of nanotechnology and nanomaterials meet the hydrogen challenge?

1.3 1.3.1

Nanoprocessing in Solid State in High-Energy Ball Mills Processes for the Synthesis of Nanostructured Materials

Since the interest arouse in unusual chemical, physical, and mechanical properties of nanostructured materials, progress in the synthesis of such materials has only been accelerating. Already established wet chemistry methods, with use of aqueous and nonaqueous solutions, have been reexamined with respect to feasibility to conduct chemical precipitation of nanocrystalline phases under conditions that hinder growth of customary crystalline phases. These methods involve precipitation of metals from its salts with use of strong reductors as borohydrides. In general, methods have been explored that accelerate the rate of reaction and thus limit the time available for ions to diffuse and contribute to growth of the nucleating crystals. Among the advantages of the established wet chemistry methods, like sol–gel processing, are good control over microstructure and particle morphology, and no need for special equipment. Sol–gel approaches were delivering good results even when a synthesis of complex structures was required (e.g., in complex multimetal oxides). Generation of nanometer-sized phases through deposition of metals and metal oxides from gas phase was among the first dry-chemical methods [62]. Many of new processes required development of special equipment or methods as to conduct

28

1

Introduction

physical vapor deposition (PVD) or chemical vapor deposition (CVD) under thermodynamically nonequilibrium conditions. Among these old and new processes for preparation of nanocrystalline alloys and compounds are the following: • • • • • • • • • • • •

Solution-precipitation methods, e.g., those with use of strong borohydride reductors Sol–gel technologies Reverse-micelle synthesis Polymer-mediated synthesis Protein-templated methods Rapid solidification and devitrification of amorphous metals and metallic glasses Combustion-flame chemical vapor condensation processes (Kear) Induction-heating chemical vapor condensation processes DC and RF magnetron sputtering, inclusive of the method of thermalization Laser ablation methods Supercritical fluid processing Sonochemical synthesis and microwave hydrodynamic cavitation synthesis

Many if not all of these processes have already been developed and well established even before the term nanomaterials was coined; they remain further refined with focus on manufacturing materials for particular properties, electronic, etc. All of them are well covered in books and monographs, which are too many to make good recommendation to our reader. For recent progress the reader can consult Nanomaterials Handbook, where the editor, Yury Gogotsi, assembled impressive body of authors to write chapters on processing and properties of many nanomaterials, particularly well covered being carbon-based nanomaterials, metallic semiconductors, and ceramic nanomaterials [135]. All of these methods stem from a natural property of atoms to self-organize into ordered structures and crystallites as being driven by energy benefits when atoms join into structures. Short-range ordering in metal alloys may be formed during deposition of atoms or atom clusters from gas phases, and survive nonequilibrium processing of metals via solidification from rapidly quenched melts. The short-range order of atoms may even be preserved in amorphous metallic alloys. The approach to build up nanocrystals through such self-organization of atoms is well recognized in the preparation of nanomaterials and termed the bottom–up nanotechnology approach. Bottom–up processing of nanomaterials relies upon use of either liquid solvents and vapor gas phases, or solidification from molten metals and compounds. However, in this book we concentrate on nanoprocessing conducted entirely in a solid state. This is a class of methods based on processes conducted in high-energy ball mills. Such methods are based on high-energy grinding and milling of materials. They provide top–down approach toward manufacturing nanomaterials.

1.3.2

Milling Processes and Equipment

The milling, grinding, and pulverizing of materials have always been one of unit operations in chemical engineering. They are well described in chemical engineering

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

29

handbooks along with heat transfer, mass transfer, fluid flow, and thermodynamic processes. Each of these unit operations gather knowledge of physical laws and practice that is necessary for reliable engineering design in many industries, viz., mineral, ceramic, and powder metallurgy. The primary objectives of milling have always been mixing or blending, change in particle shape (morphology), and first of all: size reduction. The use of the milling and pulverizing for causing the change in fundamental mechanical, chemical, and physical properties of the materials itself was not the target of milling or grinding. Use of mechanical milling for synthesis of new alloys, compounds, and nanomaterials was not envisioned until recent decades. Equipment used for milling may be classified according to the way in which mechanical forces are applied: (1) between two solid surfaces (crushing, shearing), and (2) at one solid surface (impact). The milling in ball mills combines both crushing/shearing and impact forces combined in various proportions, depending on the equipment used. There are many different designs of ball mills, which can be used for processing of advanced materials. Among them are the following: • • • • • • •

Tumbler, jar, drum, or cannon ball mills Szegvari attritor vertical mills and other vertical stirred ball mills Planetary Fritsch and Retsch model mills Shaker (vibratory) SPEX model mills A.O.C. magnet-controlled mechanical model mill (Uni-Ball magnetomill) A.O.C. electric discharge-assisted mechanical mill ZOZ continuous-fed horizontal mill, and other horizontal high-energy bead mills

Ball or tube mills have a cylindrical or conical shell, rotating on a horizontal, vertical axis. The ball mill differs from the tube mill by being short, and having the crucible length and diameter almost identical. The typical ball mill, used as much in laboratory like in industry, has been the tumbler ball mill (in laboratory practice also known as the jar mill and in industry drum mill). The grinding balls (usually in large numbers) impact upon the powder charge when cylindrical container placed on rollers rotates along horizontal axis. The balls may roll down the inside wall surface, which produces shear forces on powder trapped between the wall and the ball, but mostly they fall freely accelerated only by gravitation force, then impacting the powders (and other balls) beneath them. To maximize the impact forces imposed on powder the criterion of critical speed may be applied, where Nc expressed in rotations per minute (RPM) is the theoretical speed at which the centrifugal force on a ball (at the height of its orbital path) becomes equal to the force on it due to gravity [136]: N c = 42.3 / D ½

(1.19)

where D is diameter of the mill in meters, and the ball diameter is kept small with respect to the mill diameter. The milling in ball mills can be described as kinetic processing when the contact of the balls with the material is the main event of kinetic energy transfer from the

30

1

Introduction

grinding media into the powder. A well-known fundamental equation describes the relation between the kinetic energy – Ekin, and the mass – m, and the velocity – v: Ekin = (1 / 2 )mv 2

(1.20)

The velocity v of the ball in free fall is gt, where g is the Earth’s gravitational (downward) acceleration of the ball, equal to 9.8 m s−2, and t is the time to cover a distance from the height of its path on the shell to the impacted powder at the bottom of the shell. Hence Ekin = (1 / 2 )m(gt)2 [kgm 2 s − 2 ] or 48.02mt 2 [Nm]

(1.21)

where the N in brackets represent force units in Newtons. Since the time-of-flight of a ball in free downward direction is greatly limited by diameter D of the shell, and the energy scales as square root of time, the kinetic energy, which can be achieved in a simple jar ball mill, is greatly limited. The contact, on impact, between the grinding (balls) and ground (powder) media is also limited. Therefore, laboratory jar and industrial drum (tumbler) mills are low-energy mills (although they can provide higher energy milling if sufficiently large diameter in order of meters and many balls are used, and the drum is operated just short of the critical speed (see 1.19). From the earlier equations it becomes clear that the maximum velocity of balls relative to milled material is the primary factor in determining the energy transferred to or deposited in the material. Larger acceleration than gravitational force, hence higher velocity and higher kinetic energy of balls can be achieved in Szegvari attritor mills. In such mills the milling is conducted in a cylinder filled with balls that are stirred by rotating a horizontal shaft (agitator). The impact of shaft causes differential velocities between the balls and the powder. This can be seen in Fig. 1.2. The collective impact of a group of balls, and the impact on powder particulates trapped between a ball and the wall, is suppressed. Attritor mills are vertical mills but the same mechanical milling principle is extended to variety of other vertical mills generically referred to as stirred ball mills. There are also horizontal stirred mills. In the latter the shell is stationary and filled with multitude of small balls or beads (diameter 0.6 cm; ¼ in. or smaller); hence they are also known as bead mills. Attritor-type bead mills stir the milled media at speeds from 100 to 1,500 RPM. These mills are available commercially in batch, continuous, and circulation types. The shear mechanical milling is a predominant mode [137] in these mills as the energy of accelerated balls is dispersed among the other balls in the mill (see Fig. 1.2) whose relative motion exerts mostly shearing forces on trapped powder particulates. An attritor-type ball mill delivers ten times more energy than comparable size drum ball mill, and some of its models can be considered as medium-energy ball mills. Although the efficiency of vertical attritor mills can be low (when the volume available for powder is small and the reactor cylinder is full with balls, and when the powder has tendency to fall by gravity to the bottom of the cylinder), the new horizontal bead mill attritors are quite efficient in preparation of fine-grained materials.

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

Fig. 1.2 Motions of balls in attritor mills; the dotted vertical line represents the axis of rotation of a vertical shaft on which a number of horizontal shafts (arms) are mounted, which strike balls and agitate their motion in a very energetic manner; powder particles trapped between the balls are subjected to high-rate deformation processes

Planetary mill

31

rotating shaft

Vibrational mill

ωp ωp

Strong shearing

a

b

Strong shearing Strong impact

Fig. 1.3 Motion of balls in a planetary (a) and a vibrational (b) mill

In planetary ball mills the force acting on balls is increased. Centrifugal force of 50 times gravitational force – m × g –can easily be produced. The motion of the shell and the balls in a planetary mill is shown in the schematic Fig. 1.3a. The size of a planetary ball mill will be smaller than the size of a drum (tumbler) ball mill of the same energy. Like attritor and horizontal ball mills they can be considered as medium- to highenergy ball mills, however, milling times needed to process submicron size and nanostructured powders may be long. The commercially available mills, popular in research laboratories, particularly in Europe are Fritsch Pulverisette™ planetary mills. The models where two or four batches of powders can be processed in the cylinders mounted on rotating plate were the workhorse instruments in many laboratories. The recent Pulverisette model 7 was designed with larger rotational speeds specially for processing of powders to nanometric-level size. The Fritsch model Pulverisette™ mills, models 4, 5, and 6, are compatible with an interesting reactor cylinder, in which the cover is equipped with pressure and temperature sensor and a small radiofrequency unit allows for wireless transfer of the sensors’ data to the receiver interfaced with computer. This GTP (temperature–pressure) system developed

32

1

Introduction

Fig. 1.4 Planetary mill vial with antenna mounted on cover equipped with temperature and pressure sensors for GTP wireless monitoring gas conditions during milling

by the Fraunhofer Institute in Dresden, Germany, allows for monitoring temperature and pressure of gases evolved (or absorbed) in the reactor vial in in-situ and real-time conditions. This vial is shown in Fig. 1.4. Since the energy transferred to the material under processing in solid state increases with the frequency of impacting as square of the ball velocity, the most energetic ball mills are such that the balls are shot at high speeds and at high frequencies. This takes place in vibratory ball mills. In those mills, known also as shaker mills, an eccentric motion is imparted to the cylindrical container (or rather armature on which it is mounted) at frequencies ranging from several impacts per minute, viz., several hertz, up to 1,800 Hz, and at small amplitudes of vibrations. The balls oscillate in three mutually perpendicular, or more, planes within a small vial (several tens of cm3) as illustrated in Fig. 1.3b. At average frequency 1,200 and the amplitude of ball vibration ca. 5 cm (diameter of the rotating vial) the balls achieve velocities ca. 5 m s−1 at the moment of impact. The kinetic energy transferred to the material can be very high, even with one 5-g ball, as one can realize substituting the square of this velocity to (1.20). Most of energy transfer in vibratory mills is conducted in the mechanical impact mode although substantial shear of powder particulates is also present, as the balls rotate on the shell and the particulates are trapped between the shell and the balls. Maurice and Courtney [138] have calculated that the shock impacts imparted on the powder in the high-energy SPEX model vibrational mill are typically 40 kbars with stainless steel balls of 6 mm in diameter vibrating in a vial 50 mm of diameter. The SPEX model mill has been used extensively in research laboratories to process small batches of nanomaterials (ca. 10 cm3 or 5 g). Although this mill has been developed for grinding hard materials, such as ZrSiO4, TiO2, SiO2, and Al2O3, in analytical spectroscopy laboratories, this has been the mill that brought the solid nanoprocessing and mechanical alloying

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

33

research to laboratories worldwide, and particularly in North America. Figure 1.5 shows the SPEX Model 8,000 Duo high-energy mill for simultaneous milling of two powder batches, as manufactured by the CertiPrep Company (Metuchen, NJ, USA). The vials, Fig. 1.5, are made of hardened ferritic steel, but vials made of zirconia, alumina, agate, and hard-metal tungsten carbide are also used to greatly limit contamination by grinding media (Fig. 1.6), which can be as high as several percent of Fe when milling in steel vials. The milling that yields nanometric structure takes place about 10 times faster in a SPEX ball mill than in a planetary ball mill. The milling time needed to produce fine powders and nanostructure depends obviously on the nature of material, but also on other milling parameters. The weight of the milling balls is one important parameter, the ratio of the weight

Fig. 1.5 Vibratory high-energy SPEX ball mill

Fig. 1.6 Tungsten carbide and agate vials for SPEX mill

34

1

Introduction

of balls to the weight of powder charge B/P, is the other. One can opt for balls having particular density, which can vary from 2.3 g cm−3 to 4.0, 5.7, 7.8, and 16.4 g cm−3 for alumina, zirconia, stainless steel, and tungsten carbide, respectively. The ball-to-powder weight ratio can vary too, from 1 to 100, and obviously, this is important to maintain the same ratio when one intends to compare results from different milling. As a first approximation one can expect the frequency of the impacts to be proportional to the number of balls used (usually two or three balls are sufficient) but also on the B/P ratio. The time needed for milling to the same fines of a powder usually decreases with the increase in the B/P ratio. Said this, we have to bring the reader’s attention to often unexpected behavior of powder during high-energy milling. For example, the charge of pure aluminum powder milled for as short as 30 min in a SPEX mill changes to a multiplicity of small aluminum balls with 2–4-mm diameter, instead grinding the coarse Al powder into fine Al powder [sic!]. Agglomeration and cold welding of metal particulates to the reactor wall and the milling balls or to form balls made of the milled powder itself are common phenomena that occur during milling of pure ductile metals in vials filled with inert gas, such as argon. On the other end, milling of graphite leads rapidly to amorphous carbons. Therefore, when formation of nanostructure is this intended task, one must design the process carefully so that neither cold-welded balls nor amorphous powders are produced. Other way to apply force to a milling ball, besides the gravitational and the centrifugal means discussed earlier, is to drive ball motion by magnets. On such a magnetic principle works the A.O.C. model Uni-Ball Mill magnetic mill developed in Australia (Wollongong, NSW, Australia). The ferritic steel balls in the reactor cylinder are attracted to the inside surface of the shell by strong Nd–Fe–B permanent magnets placed outside this shell (Figs. 1.7 and 1.8). The direction of the strong magnetic force is well defined by the lines of magnetic induction, which penetrate the shell made of nonmagnetic austenitic steel to end up in the ferritic (hence soft-magnetic) balls. The pull imparted by the external magnet on the magnetic balls inside is so strong that the centrifugal force acting on the balls becomes a secondary factor in milling. High-energy milling can be achieved at low rotations (30–200 rpm). At low rotation the mill design is greatly simplified and two cylinders rotated on the same horizontal pulley can be filled sequentially with a process gas (such as hydrogen). What is more important, the motion of the balls inside the shell is well controlled by the external magnets to the point that the fraction of the impact and the shear modes of mechanical milling can be controlled. The highenergy mechanical shear force is imparted on the powder trapped between the shell and the ball when the magnet is in the 6 o’clock position. In high-energy impact mode each ball travels from the 6 o’clock position, at the bottom of the cylinder, to the 3 o’clock position, and then falls to the bottom of the reactor. The point where each ball is detached from the wall is well determined by the position of the magnet; hence each ball imparts the same energy impact on the milled powder. Since each ball travels only quarter of the circumference of the shell it falls four times per one rotation. This results in 525 impacts per minute for a shell rotating at relatively low speed of 175 rpm.

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Nanoprocessing in Solid State in High-Energy Ball Mills

35

Fig. 1.7 Motion of balls in the magnetic ball mill (courtesy of A.O.C. Scientific Engineering, Australia)

Fig. 1.8 Milling cylinder and magnet mounted for mechanical nanoprocessing in the mechanical strong impact mode (IMP2)

36

1

Introduction

At this frequency of impacts the energy transferred to the powder would be comparable with that achieved in a planetary ball mill. The milling cylinder suitable for reactive milling is shown in Fig. 1.8. High-energy ball milling is a complex process, which requires optimization of many parameters to assure repeatability of nanostructure from batch to batch. To illustrate this complexity we can list the important parameters that must be decided when conducting the process in the magnetic A.O.C. model Uni-Ball Mill: • Milling mode depends on the magnet position. By changing the magnet position one can change milling mode from low-energy shearing, through high-energy shearing, to low-energy impact, and to strong impact. Two-magnet options are also possible: (very strong shearing mode and very strong impact mode). • Number of balls used for milling usually varies from two to four balls. Maximum five balls can by used in one cylinder but the optimal number is four. • Milling speed that can be controlled in a range of 0–200 rpm depends on milling mode. More energetic modes usually require fast rotation. • Milling time that is closely related to milling mode, working distance, and type of process. Ball milling requires usually shorter time than mechanical alloying or reactive ball milling • Milling atmosphere, which is a neutral protective gas (helium or argon) during mechanical milling or hydrogen under pressure up to 0.9 MPa under reactive mechanical alloying processes. • Ball-to-powder weight ratio: This parameter depends on the mass of milled powder and number of balls and usually is in a range 10–100. The maximum mass of powder in one cylinder is 25 g what allows for milling 50 g of powder at once in two cylinders. Ball-to-powder ratio affects efficiency of milling or synthesis process and can be controlled by changing mass of milled powder or number of balls. However, there is no visible influence of the ball-to-powder weight ratio on the particle size of the synthesized MgH2 after reactive milling for 30 h. • Working distance (WD) is the distance between magnet (magnets) and cylinder. This parameter affects mainly the attractive force between the magnet and balls inside the cylinder. It is shown that increasing the WD reduces substantially the attractive force between the magnet and the 25-mm steel ball. Neodymium (NdFeB) magnets used in the A.O.C. mills may vary in attraction force. For high-energy ball milling that requires two magnets (Fig. 1.9), they are usually applied in tandem such that stronger magnets are paired with weaker ones on both sides of the Uni Ball Mill 5 (left and right) and the weaker magnet is always above (top) the stronger magnet (bottom). Ideally, magnets should have opposite polarity to double magnetic induction on balls operating in magnetic air gap. There are other advantages of employing magnetic ball mills besides the control of mechanical milling modes. Since the centrifugal force becomes a secondary factor in milling, and the reactor shell rotates at low RPM, contamination from balls and shell wear is lower than in a vibrational or a planetary mill; there is less ball wear involved and contaminations with Fe from steel become less of the problem. Also lower rotations and uniaxial movement of reactors paced on horizontal axle allow

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

37

12 Right bottom

Left bottom Force (kgf )

10 8 6 4 2

Left top

Right top

0 0

2

4

6

8

10

12

14

16

Distance (mm)

Fig. 1.9 The attractive force between the 25-mm steel ball metal vs. distance from the cylinder (working distance – WD) and four various magnets

equipping reactors with valves for filling them with reactive or inert gases. Most of the experimental results in this book that we use to demonstrate the effects of ball milling in processing of nanomaterials for hydrogen storage come from such magnetic mills employed in our laboratories. In controlled reactive mechanical milling/mechanical alloying modes (CRMM/ CRMA), discussed latter in this chapter, powders can be nanoprocessed under sequential supply of reactive process gases, such as N2, O2, etc. When the process gas is H2 and the powder is a metal, we can consider this mechanical alloying with hydrogen and term this process hydrogen alloying (HA). This has been used for a direct mechanical synthesis (mechanosynthesis) (MAS) of interstitial metal hydrides. We will discuss these processes in the following chapters.

1.3.3

Nanoprocessing Methods and Mechanisms

Five major processing methods that are employed during milling of materials in high-energy ball mills are as follows: • • • • •

Mechanical milling (MM) and mechanical disordering (MD) Mechanical alloying (MA) Mechanochemical activation synthesis (MCAS) Mechanochemical synthesis (MCS), and hydrogen alloying (HA) Mechanical amorphization (MAM)

In the following paragraphs we will briefly describe all the aforementioned methods, while paying particular attention to MA, which is at the core of all benefits brought about by mechanical processing of nanomaterials. Yet, multiplicity of phenomena of mechanochemical coupling that occurs during processing of alloys and compounds in high-energy ball mills have not been well understood, yet well reviewed before.

38

1.3.3.1

1

Introduction

Mechanical Milling

High-energy ball milling is the only nanotechnology top–down approach for the synthesis of nanoparticles. In the process of mechanical milling (MM) brute force is applied to material, whether it is a metal, a pre-alloyed intermetallic, or a solid chemical (stoichiometric) compound: the force that is sufficient to disperse the material into fine nanometric particulates or agglomerates of such material. Since only reduction in particle size is required and not substantial rearrangement of atoms takes place in milled particles the time required for MM is short. Yet, even short milling times can break thin chemically passive surface coatings (e.g., surface oxides) and expose fresh, clean chemically active metallic surface. Such milling can also introduce defects into solid compounds. It results in an increased chemical activity of milled media toward both gasses and chemical reactions in solutions and electrolytes; the process is often termed mechanochemical activation synthesis (MCAS). Longer milling of intermetallics leads to changes in the long-range order in intermetallics. Such a process termed sometimes mechanical disordering (MD) leads to formation of disordered intermetallics, alloys, and compounds. A process wherein mixture of elemental metal powders, or powders of metal and nonmetal are milled long enough to trigger alloying of elemental powders in a solid-state process is termed mechanical alloying (MA). Since substantial rearrangement of chemical species must take place during alloying the latter requires long times of milling, usually more than twice the time needed for preparation of nanopowders by mechanical milling. Even longer milling results in what we can term mechanical amorphization (MAM) of a crystalline solid. The amorphous alloys produced in this way are analogous to amorphous metal phases (metallic glasses) produced by rapid quenching of metal melts as described in “Introduction,” Sect. 1.2.2. When the milling process initiates a solid-state reaction and yields a new stoichiometric or quasi-stoichiometric chemical compound, such as carbide, nitride, silicide, etc., such a process can be termed reactive mechanical alloying or reactive mechanical milling (RMA/RMM). When mechanical modes of milling (shear, impact) are controlled we have controlled reactive mechanical alloying or controlled reactive mechanical milling (CRMA/CRMM). Reactive mechanical milling can be realized by reacting two chemical compounds A and B in a solid-state synthesis that yields a distinct chemical compound C. This is termed mechanical synthesis or mechanosynthesis. Reactive milling can also be conducted by milling metal powders in ball mills filled with a reactive gas, such as N2. Hydrogen can also be used for reactive milling and leads to reduction of oxides during milling. Hydrogen can also alloy with metals to form interstitial metal-hydrogen solutions, hydrides, or chemical hydrides. Reactive mechanical milling in hydrogen-fed ball mills, we can term hydrogen alloying (HA), is a new efficient way to the discovery and development of new hydrogen storage materials. Obviously, hydrogen alloying can be used for mechanosynthesis of chemical hydrides. In the following several paragraphs we describe briefly mechanisms that may be at play during ball milling of powders in high-energy ball mills.

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

1.3.3.2

39

Mechanical Alloying

Mechanical alloying is conducted through milling of two elemental metal powders. Longer milling, order of hours, leads to atomic-level mixing of the metals and produces an alloy consisting of these metals. Until the advent of this new, nonequilibrium and low-temperature solid-state processing method, metal alloys could only be manufactured by melt casting and metal foundry practices. There is a sequence of consecutive mechanical events together with concurrent events of chemical nature that leads to the realization of the mechanochemical process described as mechanical alloying. Mechanical alloying can be regarded as a repeated stressing, deformation, fracture, and cold welding of powder particulates, as illustrated schematically in Fig. 1.10. These mechanical processes are interlaced with enhanced diffusion of chemical species and increase in structural disorder. The first step is development of high-rate stresses in the milled powders. The key step is repeated high-rate, severe plastic deformation followed by series of chemical and mechanical processes, which result in generation of structural disorder, extended solid solubility of chemical species, and formation of metastable phases and nanostructures. The sequence of concurrent mechanical and chemical events can be written as follows: Mechanical stressing ® severe plastic deformation ® formation of a submicron lamellar microstructure ® cold interdiffusion of metal atoms between lamellae or nanograins (cold welding) ® fracture ® formation of nanostructure ® extended solid solubility ® mechanical alloying with formation of thermodynamically stable and/or metastable phases ® amorphization. A great deal of research papers, also review articles, was published on the subject of mechanical alloying. Also, there are serial conferences and international meetings on the subject, where new developments in mechanical alloying for processing of nanomaterials, including those for

repetitive stressing

welding

fracturing

Fig. 1.10 Mechanical alloying through repetitive cold welding and fracturing

40

1

Introduction

use in hydrogen storage, are reported.. Since this field of materials science is growing rapidly at the time we write this book it would be difficult to select scholarly papers published on this complex subject. Instead, we will refer the reader to an extensive review, which could be of particular help [139].

1.3.3.3

Mechanochemical Activation

Stresses, Deformations, and Equilibrium Mechanical Processing The local pressure applied to solid by two colliding balls, or a ball hitting a solid trapped between it and the shell of the cylinder can reach several GPa [140]. The mechanical stress in the material under milling develops as particles are subjected to the effect of pressure created between two milling surfaces. If this stress on ball impact is normal to the surface of solid, tensile or compressive stress occurs. In tangential movement of the milling surfaces the stresses are of shear stress type (Fig. 1.11). Actually, both stresses occur simultaneously as the material is processed. In attritor (beads) or vibrational ball mills the particles move relative to milling balls and collide either with the milling surface or with each other. The material particles and the milling balls can not only move forward but can also rotate (Fig. 1.2), so more complex stresses develop in addition to the impact stresses. In more general continuum mechanics approach to three-dimensional solid bodies, one must identify several planes on which the forces act, and to calculate the stress on each plane in the milled material, as it is shown in Fig. 1.12. One must note that if the stresses are expanded in terms of the normal-impact and tangential-shear

tension-on ball impact

compression-on ball impact

shear, on tangential ball impact

Fig. 1.11 Deformation of an ideal material under normal (tension, compression) forces and shear forces exerted by balls. The positions of balls are shown for impact and shear mode of milling

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Nanoprocessing in Solid State in High-Energy Ball Mills

41

Fig. 1.12 Expansion of the ball-milling stress into normal (impact) and shear stresses

components, then the set of all components can be called a tensor σ in Euclidean space. Then, any stress state can be represented by its minimal and maximal components, known as principal components [141]. These stresses can be measured as the average stress in many milled particles, and reflected in the stress broadening in addition to the grain size broadening of X-ray diffraction peaks in milled powders (how, it will be shown latter in this section for the Williamson–Hall method, Fig. 1.19, and in Sect. 1.4.3 for the Cauchy/Gaussian method). The principal stresses can also be measured in larger single milled particles used in a stress diffractometer [142]. For example, during milling of cast Mg crash in SPEX mill, normal compressive stress, several times greater than the shear compressive stress was measured and exhibited gradual change to 0, while the shear negative (compressive) stress changed to positive (tensile). Such processing resulted in mechanical activation of Mg metal for absorption of hydrogen [142]. During milling conducted in the A.O.C. Uni-Ball Mill the occurrence of predominant stresses, shear or normal (impact), can be somehow controlled through special configuration of magnets that affects the trajectory of motion of milling balls. While action of forces and stresses during high-energy ball milling has not been well investigated it is apparent that they cause a high degree of elastic and plastic deformation. In simple terms elastic deformation is defined as the reversible deformation that occurs when a load is applied. Most materials deform in a linear elastic fashion i.e., the amount of reversible deformation is a linear function of the applied stress up to a certain stress level, i.e., the strain (tensile, compressive, or shear) divided by original length e, is proportional to the applied stress, s (force divided by cross-section area), according to the Hooke’s Law s = Ee

(1.22)

where E, Young’s modulus is the materials constant, which quantifies elastic behavior. Most nonmetallic materials, such as salts, oxides, and ceramics deform also in such a linear fashion, although in a very small range: if the applied force is further increased the compound fractures in a catastrophic manner. One can notice that the atoms in the stressed body change their relative positions (with respect to one another) more under the shear elastic stress than under tensile

42

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Introduction

or the compressive elastic stresses. Therefore, we can focus more on the elastic shear stress. If, further to the above, one assumes proportional relationship to reflect on lowering energy by the action of mechanical force then a term proportional to the elastic shear stress, a • s, must be subtracted from the activation energy EA [143]: k = k0 exp[ −

( EA − as ) ] RT

(1.23)

where a is a proportionality constant, and k is the rate at which covalent bonds in a nonmetallic material are ruptured. Now, one can substitute the stress for the strain, and the relation between applied elastic stress and the rate k can be proposed: k = k0 exp[ −

( EA − Eae ) ] RT

(1.24)

Therefore, the rate at which chemical bonds break increases with elastic shear stressing of the material. The rupture of chemical bonds, hence fracture of material, leads to its fragmentation into particles. This reduces the average particle size in powder as fractured particles multiply into even smaller particles. Equation (1.24) points to the importance of elastic shear strains in mechanical activation of chemical bonds for particle size refinement and production of nanoparticles. Most metals initially deform elastically and then begin to deform plastically. Plastic deformation allows energy of stresses to be dissipated rather than building up to the point where the material breaks. Such dissipated energy is used for grain size refinement in each of the deformed material particles. This results in decrease of the average grain size. Therefore, the energy of balls, which was transferred to the material, is utilized for both particle and grain size refinement. Amount of energy used for each of these two processes is largely determined by the value of Young’s modulus. We will come back to the issue of grain refinement in the following paragraphs.

Nonequilibrium Mechanical Processing and Excitations Although less to be seen this way, (1.24) points also to the importance of elastic shear strains in a possible mechanism of mechanical activation of materials for chemical reactions. Apparently, ruptured chemical bonds should lead to increased reactivity of the milled powder toward gases, and predominately toward hydrogenation reaction. However, for the sake of efficient mechanical activation the elastic strains should be isolated from the plastic strains. One way of doing this could be preventing the dissipation of energy injected into material by the impacting ball. Dissipation of energy in deformed materials needs time, and if there is no time sufficient for such dissipation between two stressing events (being these impact or shearing events) then the energy must be reabsorbed locally in near

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

43

adiabatic processes. This should demonstrate itself in formation of adiabatic shear bands, as caused by strain-localization phenomena that are generally attributed to a plastic instability arising when thermally activated softening exceeds work hardening in the material. Adiabatic shear bands are well known to play a key role in dynamic fracture and fragmentation, especially in hard (or work hardened) materials [144]; unfortunately, such processes have not been well described in the context of events taking place during high-energy milling, when powder particles are deformed severely at high rates of mechanical deformation. Nevertheless, it would be possible to view a picture where the energy of mechanical process is transformed locally and instantaneously into the elastic energy of either vibrational modes of thermally softened lattice of atoms or into higher oscillational modes of electrons in the atoms in such a lattice. The atoms, whose electrons achieved higher oscillation states, become ready for chemical reactions; thus, the material becomes mechanically activated for chemical reactions. The mechanism we are putting forward here is one of several possible mechanisms that may be at play during nonequilibrium processing and mechanochemical activation. The usual way to activate material for chemical reaction is through heating it to the average temperature that triggers the reaction. This constitutes thermal activation. It cannot be said firmly that milling process does not include local heating and local rises in temperature. Such a local and instantaneous rise in temperature is difficult to quantify, yet it was reported to reach 400°C in highly energetic milling of Ni–Zr powder mixtures [145]. However, the average temperature of the powder, measured with a thermocouple touching the wall of milling cylinder rarely exceeds 60°C in a SPEX mill. Theoretical analysis of ball milling process in a SPEX mill demonstrated the same ~60°C increase [140]. In thermal activation, mechanical energy transformed to kinetic energy of vibrating atoms can activate atoms for chemical reactions. Thermally activated reactions can be described in mechanistic (kinetic) terms where atoms or molecules with increased kinetic energy (hence temperature) exchange energy and interact with each other in collisions. The kinetic energy of the colliding atoms is on an average lower than the energy of chemical bond, but the local quasi-stresses at the moment of impact can be sufficient to cause extensive polarization (deformation of shape) of the molecule, or cause polarization of the chemical bond as in Fig. 1.13. For atoms in a crystal lattice, the change in the force that bonds atoms, from zero net force (point r0) to attraction (rt) or repulsion (rc), will result in increasing the population of higher oscillational states of electrons (not shown in this picture). Then, the relaxation of elastic energy with the excitation of interatomic bonds triggers formation of an unstable transition state or a short-life complex in which interatomic bonds differ from the initial molecules. Stabilization of such a complex and its dissociation along a new reaction path may lead to formation of new reaction products. Whichever be the mechanism, mechanochemical activation in high-energy ball mills is thus an overlooked method to activate solids for chemical reactions. We will discuss this in the following sections.

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Introduction

V(r/r0)

rt

attraction net force

repulsion

r0

r

rC

compressive stress

tensile stress

Fig. 1.13 The elasticity of the chemical bond, gray area regions points to the polarization of chemical bond between the atoms in position 0 and r0, rc or rt

Disordering, Faulting, and Grain Refinement The relaxation of elastic energy with the excitation of interatomic bonds may be one fundamental mechanism that drives mechanochemical activation process. The other mechanism can be breaking of passive oxide layers during milling of metal powders, hence creation of fresh metal surfaces. There is, yet another important aspect of milling processes that may lead to mechanochemical activation. This is disordering and faulting that occurs in crystalline lattice near the particle surface, as well as in the particle interior. As we stated before, original source of novel properties of mechanically alloyed powders is a stressing process. If there is enough time for dissipation of energy imparted on material by ball milling, such energy is dissipated by plastic deformation processes. Such processes generate highly disordered structures within the processed particles. In metal particles, plastic deformation requires the movement of dislocations within the crystal structure, as shown, although in a simplistic manner, in Fig. 1.14. Movement of dislocations can be hindered by foreign solute atoms, precipitations of second phases, and other dislocations. Pinned dislocations can contribute further to generation of new dislocations since they form sources that emit more dislocations, as in the famous FrankRead source. As dislocations become entangled, the work hardening begins.

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45

Fig. 1.14 One step in the movement of an edge dislocation through the crystal lattice. As a consequence of the extra-half plane of atoms regions are formed in the vicinity of dislocation where compressive and tensile stresses/strains are imposed on the neighboring atoms

The dislocations snag and pile up. Material loses its plasticity. It hardens. As the cohesive forces that keep atoms in the hardened material are overcome by the milling forces, which now cannot be plastically dissipated, the material structure becomes disjoined and rearranges into a polycrystalline structure consisting of many grains. Once polycrystalline structure is formed the Hall-Petch relation can relate the yield strength, YS, and the grain size, d YS = YS0 + Kd −1/ 2 ,

(1.25)

where YS0 and K are material constants. It must be noted that starting from grains so small that dislocations can not pack and pile up, the inverse Hall-Patch relationship can be at work and increased plasticity (superplasticity) has been reported in ceramic nanomaterials. However, until the grains achieve critical size the ball milling produces multigrain particulates. The atomic structure of a single grain is, in principle, across the grain. However, in the regions of grain boundaries, the average atomic density, interatomic spacing, and the coordination between nearest neighbor atoms deviates from that inside the grain, and differs from region to region. The presence of these two structural constituents (grains and grain boundaries) at comparable volume fractions and with typical crystal sizes of a few nanometers is crucial for the properties of nanocrystalline materials [62]. In consequence, many of the physical and mechanical properties of nanocrystalline solids, such as thermal expansion, elastic constants, fracture stress, and ductility, are widely different from those of the same material having conventional grain sizes. This is a direct consequence of the fraction of atoms in intercrystalline positions being significant. Consequently, interface structures in these materials are bound to play a major role in material properties. Once the crystal size and boundary dimensions become comparable within certain length scales new physical effects are to be expected. This grain boundary fraction is truly disordered, if not plainly amorphous. This is illustrated in Fig. 1.15. The fraction of atoms that counts into grain boundaries should increase with milling time; however, in powder particulates below some

46

1

Introduction

a grain A grain B b

c

Fig. 1.15 Enlarged view of hydrogen atoms (black circles) in the grain boundary between two grains A and B (gray circles). Note varying short-range order around atoms a and b, which belong to the grain boundary, and fixed order around atoms of type c, which belong to the grain core

critical size there is little or no grain boundaries. This may have significant consequences on chemical reactivity of milled powders because the disordered grainboundary phase is seen as the place when diffusion of chemical species is faster and chemical reactions are facilitated. Nanostructured materials are such materials where the density of grain boundaries is maximized. It must be noted that nanopowders do not necessarily consist of nanometric particles (nanoparticles). The powder particles may be actually of micrometric sizes, yet the particle interior may be divided into many nanograins, as this is depicted in Fig. 1.16. Such nanograins are the smallest diffracting domains in X-ray diffractometry, and being incoherent to X-ray diffraction they cause peak broadening in X-ray diffraction (XRD). The average grain size is determined from such broadening of peaks. We will discuss how this can be done later on.

Hydrogen Entry into Grain-Refined and Disordered Powder Structures Hydrogen molecule, H2, dissociates on contact with catalytic metal surface, and then enters the material as dissociated hydrogen atoms through grain boundaries. Subsequently, the hydrogen atoms diffuse preferentially through quasi-amorphous grain boundary network (regions a and b in Fig. 1.15) before they begin to diffuse into grain interior (region c). Therefore, the specific surface area (SSA) and

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Nanoprocessing in Solid State in High-Energy Ball Mills

47

H2

particle size

grain size

Fig. 1.16 Particle size and grain size in a nanostructured powder. Black points depict hydrogen molecules () in pores between particles, and hydrogen atoms () diffusing through a grain boundary network and gradually entering the grain interior; black triangles point to triple points that are preferential sites for hydrogen atom accumulation

the fraction of material that belongs to grain boundaries are equally important. The beauty of nanostructuring in ball mills comes in the observation made that both the particle size and the grain size are equally rapidly decreased with even short processing in a high-energy ball mill, as shown in Fig. 1.17. Yet, the picture in which the particle size and the grain refinement are sole reasons for mechanochemical activation of powders for reactions with hydrogen is not complete. Nanostructuring by ball milling introduces a variety of defects, vacancies, dislocations, and stacking faults besides the earlier described grains and grain boundaries. These defects raise the free energy of the system making it accessible to formation of thermodynamically metastable phases. Also, defects lower the activation energy of reactions limited by poor kinetics. One type of such defects that would be of particular interest is stacking faults. Heavy deformation via cold working, like that occurring in ball mills, introduces stacking faults on (111) planes in fcc metals, such as Ni, and in the basal (0001) or prismatic {1010} planes of the hexagonal close-packed metals, such as Mg. The former is a well-known catalyst for hydrogen storage, and the latter is a hydrogen storage metal itself, storing as much as 7 wt% H2 in a reversible process. There are few investigations toward the effect of stacking faults [147] – such as twins and deformation faults – on gas hydrogen storage in hexagonal metals and alloys.

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Introduction

Commercial MgH2-ball milled 100 Particle size ECD (mm)

350

Grain size (nm)

300 250 200 150

10

1

0.1

100

0

5

50

10 15 Milling time (h)

20

25

0 0

5

10 15 Milling time (h)

20

25

Fig. 1.17 Simultaneous reduction of the particle size and the grain size in fully hydrogenated magnesium powder, MgH2 [146]

Interplanar Hydrogen in Disordered Layered Compounds Formation of stacking faults has been observed in some inorganic compounds under even mild milling conditions. Some layered inorganic compounds of interest for sorption and desorption of gas hydrogen (H2) and electrochemical hydrogen (H+) exhibit disorder in stacking of atomic planes. Among them are graphites and some nanocarbons to be discussed in Chap. 4. Stacking faults were also observed in battery-grade nickel hydroxide, Ni(OH)2, used in the positive electrode of rechargeable alkaline nickel metal hydride (NiMH) battery [148]. We will turn to nickel hydroxide powder to demonstrate the effect of ball milling on improved reversible insertion and deinsertion of hydrogen ions between hexagonal basal planes of layered nanocrystals. Generic Ni(OH)2, having a Mg(OH)2 brucite-type structure, is a compound with well-stacked hexagonal basal planes of NiO2 composition, seen in TEM image, Fig. 1.18a, and H atoms bonded to O atoms in hydroxyl, –OH, groups that hang inside interplanar space. The planes are stacked through action of relatively weak van der Walls bonds. Battery-grade Ni hydroxides were found to have large degree of disorder in stacking of planes, Fig. 1.18b, in contrast to well-stacked planes in generic hydroxide (Fig. 1.18a). The existence of such stacking fault disorder and its role in improving insertion of hydrogen ions into electrodes were recently investigated [148, 149]. Mechanochemical activation of these materials in ball mills was proposed to be an efficient method for improvement of reversible insertion of electrochemical hydrogen, viz., H+, into these structures [148]. A concept of stacking fault disorder, which was postulated on the evidence from X-ray diffraction (Fig. 1.19)

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

49

a

b Fig. 1.18 Stacking of hexagonal basal planes in Ni(OH)2. (a) Generic and (b) ball milled; sets of chevrons underline the shift of planes due to shearing forces in the milling process, after [72]. Also, shown hydroxyl – OH – groups, light grey – O, and black – H atom dumbbells extending into the interplanar space

and direct imaging in high-resolution electron microscope (Fig. 1.18b) [148], has been since confirmed in computer simulation [149] and batch-scale experiments with use of ball mills [150]. Even a short ball-milling for less than half an hour in SPEX mill results in proliferation of stacking faults, followed by fragmentation of the structure into small domains of layered nanocrystals of this compound [151]. Therefore, insertion and deinsertion of hydrogen ions is facilitated in short-layered crystallites, i.e., layered nanocrystallites, and this leads to the observed improvement in the electrode capacity, lower swelling on charge/discharge cycles, and extended durability [152]. Stacking-fault-type disorder is rarely discussed in the context of milled metals, alloys, and hydrides. We shall not discuss it more here either. However, we will come back to this type of disorder later in Chap. 4 when we will discuss disordered graphites and nanocarbons.

50

1

Introduction

a

b Fig. 1.19 (a) X-ray diffraction pattern with nonuniform line broadening, sharp lines for h − k = 3n, and broad lines for h − k ≠ 3n, which gives evidence to presence of stacking fault-type disorder in the structure of hexagonal Ni(OH)2 for electrochemical hydrogen insertion; (b) WilliamsonHall plot construction is used to determine the degree of stacking fault disorder [38]; the plot (b* = bcosq/l) vs. (d* = 2sinq/l) exhibits large scatter of b* values for the material with stacking faults(●, ) as compared with the material without faults (▲, Δ); the points for the indices hk0 (●, ▲,) are not affected by the disorder, as well as are the points for the fcc Ni reference powder (▲); l and q are the wavelength and the half of two-theta angle at which Bragg peaks are centered (adopted from [148])

Interstitial Hydrogen in Disordered Alloys Well-ordered intermetallic compounds (alloys), when processed in high-energy ball mills exhibit atomic (chemical) disordering in the early stages of ball milling [153]. Let us take as an example ordered AlRu intermetallic crystals that are of B2 type and β-CuZn structure. This structure consists of two simple cubic interpenetrating

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

51

sublattices, one mostly occupied by Cu atoms and the other having mostly Zn in lattice nodes as represented by black and white atoms in Fig. 1.20a. At high temperatures, above the critical temperature of 1,015 K, the long-range ordering disappears and both kinds of atoms become randomly distributed in the two sublattices in equal fractions. Long-range chemical order is lost and two sublattices become indistinguishable in X-ray diffraction pattern, characteristic of BCC structure. Such disordering phenomena that occur in alloys at increased temperatures have long been known, and large body of knowledge has already been accumulated in the domain of physical metallurgy and material science. However, similar disordering, induced without use of thermal activation and proceeding in room temperature, has become to be apprehended only since the advent of mechanical alloying. Indeed, high-energy ball milling of AlRu in ambient temperature for 32 h in a vibratory ball mill under high vacuum generates stable disordered solid solutions [154]. Even more intensive milling can lead to unstable solid solutions with extended solubility limits of one type of atoms in the crystalline lattice formed by other atoms. Ultimately, atoms in an alloy or a compound can be mechanically mixed to form amorphous phase, where chemical disorder is convoluted with topological disorder. We will come back to these disordered phases in the following paragraph, when we will discuss mechanochemical synthesis routes that lead to formation of nonequilibrium phases. This could have far-reaching consequences, as mechanically induced disorder can be a factor in activation of well-ordered intermetallic compounds (alloys) for reversible reactions with molecular hydrogen. Figure 1.20b shows the PCT plots, which details hydrogen sorption properties for B2-type FeTi, an important hydrogen storage alloy with capability for reversible storage at room temperature applications. The plateau of equilibrium sorption is lowered in ball milled, disordered FeTi in comparison with the not-milled counterpart. Therefore, the alloy has been

1

b sublattice

In P (MPa)

0 −1

c b

−2

a

A atom −3

B atom

0.0

a sublattice

a

b

0.4

0.8 H/FeTi

1.2

Fig. 1.20 The B2-type CsCl structure of FeTi stoichiometric, ordered, compound (left), and room-temperature hydrogen PCT properties for B2-type FeTi hydrogen storage alloys: amorphous – a, nanocrystalline − b, and crystalline − c (adopted from [155])

52

1

Introduction

mechanically activated for sorption of hydrogen at lower pressures. Inspecting the plots (Fig. 1.20b) one can also notice that overmilling the alloy to the amorphous powder is detrimental, as sorption plateau disappears, and the powder capacity for hydrogen sorption becomes null. Said this, we can let the reader to recall Fig. 1.15, where we depicted amorphous-like phase regions at grain boundaries as the pathways open for preferential diffusion of hydrogen atoms. Apparently, an alloy can benefit from some fraction of amorphous phase to improve kinetics of hydrogen absorption, but complete amorphization of crystalline lattice lowers capacity for storing hydrogen [156]. Mechanochemical activation is therefore a complex process where kinetic and thermodynamic effects must be firstly well understood, and then optimized. Summing up, we identified several reasons and phenomena for mechanochemical activation of powders for chemical reactions with hydrogen at room temperatures in high-energy ball mills. Among these discussed here were the following: ●

● ● ●

●

●

●

Breaking of oxide layers and passivation coatings, hence exposure of fresh catalytic sites on metal surfaces High levels of elastic shear and other stresses induced by milling Particle size reduction, hence increase in the specific surface area Grain size reduction and maximizing the density of grain boundaries in nanostructured particles Stacking fault disorder and fragmentation of structures into multiplicity of layered nanocrystals (viz. in oxides, hydroxides, and graphites) Atomic (chemical) disorder that leads to formation of stable and unstable (extended) solid solutions Topological disorder that leads to formation of amorphous phases

In mechanochemical activation it would be the action of several of the aforementioned parameters, and their relative importance is often difficult to separate. There may be also other phenomena and synergetic effects too, which have not been elucidated yet. Whatever are the mechanisms, mechanochemical activation works amazingly well, and presents an attractive alternative to thermal activation. It also provides easy, top– bottom method for manufacturing nanostructured hydrides and nanocarbons.

1.3.3.4

Mechanochemical Synthesis (Mechanosynthesis) of Nanohydrides

A primary method of mechanochemical synthesis of nanostructured hydrides (nanohydrides) is processing by mechanical (ball) milling. Processes of manufacturing of nanocrystalline/nanostructured hydrides by ball milling are shown in Fig. 1.21. There are three major processes all of which start from raw metallic/nonmetallic elements [157]. (1) In the first procedure, called a two-step method, either premilled pure metallic elements or a mechanically milled (MM) precast intermetallic ingot are hydrogenated under gas hydrogen pressure to form a desired metallic or intermetallic-

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

53

Raw materials Mixture of chemical compounds or hydrides

Elemental powders

Pure elements

Inert gas Inert gas

Arc melting

Preformed hydride

Mechanical alloying (MA)

H2

Bulk material

Milling of powders (MM)

Alloy powders

Reactive mechanical alloying (RMA)/ Mechano-chemical Synthesis (MCS)

Activation and hydrogenation

Mechano-chemical activation synthesis (MCAS)

By-product removed

Nanostructured hydrides

Fig. 1.21 Flow chart showing the possible methods of manufacturing nanocrystalline/nanostructured hydrides (modified from [157])

based hydride such as, for example, MgH2 or Mg2NiH4. For hydrogen alloying of a powder mixture, elemental powders are mixed in desired proportion and then mechanically alloyed (MA) by ball milling under inert gas (e.g., argon) to form intermetallic compounds. The formation of intermetallics in milled metal–metal and metal–metalloid (e.g., B, Si, C) systems is the first step of mechanosynthesis [158]. The general reactions for a binary system A and B are as follows: nA + mB → AmBn via ingot / powder

(1.26a)

AmBn + (x/2)H2 → AmBnHx hydrogenation

(1.26b)

(2) In the second method, so-called reactive mechanical milling (RMM), elemental metallic powders are milled under hydrogen atmosphere. This procedure is a one-step method. The reactions are a direct alloying of metals and hydrogen, i.e., hydrogen alloying. When sufficient amount of the hydrogen reactant is supplied this reaction yields hydrides with well-defined stoichiometry: A0 + (x/2)H20 → Ax+Hx1–

(1.27)

As seen from the change in valency of metal and hydrogen, this is a redox reaction because the metal is oxidized (loses n electrons) by the hydrogen, while the valence of hydrogen becomes, formally, −1. Magnesium dihydride, MgH2, can be manufactured in ball mills in such a manner, with the reaction yield close to 90%. The reaction of

54

1 Introduction

(1.27) is accompanied by both the particle- and the grain-reduction, as discussed in Sect. 1.3.3.3. In a variant of the second method described earlier the premixed metallic powders (or pulverized ingots) are milled under hydrogen atmosphere to directly form an intermetallic hydride. It can be also viewed as hydrogen alloying of metal powders and powder mixtures in hydrogen alloying mills. This method is called a reactive mechanical alloying (RMA) or mechanochemical synthesis (MCS). nA + mB + (x/2)H2 → Am BnHx

(1.28)

MCS-type reactions can occur while milling a simple metal hydride with a metallic element M, 2AH + 2Al+ 3H2 → 2AAlH4

(1.29a)

for instance NaH with Al [159]: NaH + Al + 3/2H2 → NaAlH4

(1.29b)

Mechanosynthesis reaction can also be conducted with use of a second hydride, instead of hydrogen gas reactant. Most often the first hydride is a simple metal hydride, while the second hydride is a complex hydride. 2NaH + NaA1H4 → Na3A1H6

(1.30)

Such reactions have been observed to occur where an alanate such as NaAlH4 is milled with alkali metal hydride, such as NaH or LiH. (3) In the third method called mechanochemical activation synthesis (MCAS), a mixture of metal compound, viz. metal chloride, is ball-milled to induce a reaction to yield a high-hydrogen capacity hydride. If an n-valent metal chloride or fluoride (AXn) is used in the reaction with alkaline metal borohydride B(BH4) then in general terms the reaction can be written as AXn + nB(BH4 ) → A(BH4)n + nBX

(1.31)

where A(BH4)n is a newly synthesized complex metal borohydride and BX is a salt (X = Cl or F). Either Li or Na borohydrides, LiBH4 or NaBH4, are used as source of hydride complex. Similar reactions are used for preparation of alumohydrides, viz. alanates, using Li or Na alanates, LiAlH4 or NaAlH4. These reactions are mutual substitution reactions where the ions in two compounds change sites in the crystalline lattice. The valence of the metals and the hydrogen do not change from the reactants to the products. This is analogous to well-known metathesis reaction, which is a common method for preparation of borohydrides in donor nonaqueous solvents (tetrahydrofuran, ethers). The MCAS advantage is that solid-state process

1.3

Nanoprocessing in Solid State in High-Energy Ball Mills

55

does not require use of an organic solvent, whereby the borohydride product does not come with the organic solvent molecules bonded in its crystalline lattice. Said this, we must observe that the MCAS process yields a mixture with alkali metal chloride or fluoride, and to purify the complex hydride product one must call upon a difference in solubility between the borohydride and the metal chloride in the same THF or ether solvents. There are also disadvantages that can be mentioned when opting for preparation of borohydrides or alanates by solid-state mechanosynthesis. Certain hydride complexes are not stable, and become stable only when borohydride or alanate molecule is solvated. Also, certain structures based on hydride complexes may be not stable under mechanical ball milling. The mixture of Mg and B milled with hydrogen yields an amorphous phase, instead of the anticipated Mg(BH4)2, with release of hydrogen gas [160]. The MCAS product can also be a mixture of an amorphous phase and a boride, such as MgB2. This brings us to the subject of synthesis of amorphous phases in mechanochemical processing.

1.3.3.5

Mechanical Amorphization

When a mixture of elemental powders is milled, the formation of the amorphous phase is due mainly to an ultimate interdiffusion of atoms that occurs at fresh surfaces and interfaces created by mechanical milling [87]. This interdiffusion is promoted by defects and chemical disorder in the crystalline structure. However, no such interdiffusion is prominent while milling intermetallics, where constituents are already alloyed. In this case the accumulation of the chemical disorder (as discussed in previous paragraphs) leads to collapse of crystalline structure if the rate of dynamic recovery of crystalline order is less than the rate of deformation. This is a nonequilibrium thermodynamic process. The phase formed under nonequilibrium milling is an unstable solid solution. In a solution consisting of two kinds of atoms that differ in size (more than 15%) their packing on the crystalline sites becomes less dense due to the disorder and defects. This constitutes a high-energy phase that has an inherent tendency to transform to more stable phase. This is a metastable amorphous phase, in which atoms abandon the crystalline sites and rearrange into close-packed topologically random structure. Yet, there is an alternative rearrangement of atoms to even lower energy structure. Such is an equilibrium crystalline compound. Addition of the third kind of atoms (in respect to atomic radii) during milling with hydrogen seems to increase ability of disordered solid to form an amorphous structure [160]. Hence, it is possible that the topological disorder brought about by elastic transformation under dynamic shear and impact stress can be stabilized by hydrogen atoms [161]. However, the alternative is to rearrange atoms into an equilibrium crystalline hydride compound. This would require nucleation of hydride grains in the matrix of unstable crystalline solid solution that was formed on milling. Since nucleation proceeds en masse, yet crystal growth rate is slow, the process yields hydride nanophase. Therefore, there are two parallel routes to lower the energy of unstable mechanically disordered solid solution: (1) formation of an intermediate amorphous phase first, followed by crystallization an

56

1

Introduction

equilibrium compound and hydride nanophases [160, 161], and (2) transformation to equilibrium crystalline compound and hydride nanophases directly. What reaction path is chosen depends on both thermodynamic and kinetic conditions, and the products in mechanochemical reactions shown in the reaction of (1.31) may come with sometimes unavoidable fraction of an amorphous phase. Summing up, we discussed four main routes toward nanoprocessing of hydrides in ball mills. These are as follows: ● ● ●

●

Mechanical milling (MM) of cast ingots followed by hydrogenation Mechanical alloying (MA) of elemental powders followed by hydrogenation Mechanochemical synthesis (MCS) by reactive milling of metal powders or intermetallics in hydrogen gas Mechanochemical activation synthesis (MCAS) by reactive milling of compounds

All of these syntheses are preceded by mechanical activation of reagents in ball mills and all of them can be considered when manufacturing of nanohydride powders is anticipated. However, solid-state reaction routes are not always well investigated, and can be complex. More often than not, the yields of reactions leading to nanohydrides are not reaching 100% and an amorphous phase, which provides lower storage capacity, is formed.

1.4

Important Hydride Properties and Experimental Techniques

1.4.1

Thermodynamics

1.4.1.1

Pressure–Composition–Temperature (PCT) Properties

Pressure–composition–temperature (PCT) curve called also pressure–composition isotherm (PCI) curve can be a source of important fundamental information related to thermodynamic properties of solid hydrides. There are several methods of determining PCT properties ranging from thermogravimetric to precise volumetric measurements obtained by using a classical Sieverts-type apparatus. The thermogravimetric methods unfortunately are extremely limited in pressure applied during test to the maximum 2 MPa, which is usually around equilibrium pressure only for low-capacity hydrides. The plateau pressure at the decomposition temperature for high-capacity hydrides is much higher and reaches from 2.5 up to 15 MPa. A typical isotherm of a reversible hydride is shown in Fig. 1.22 By measuring the changes in hydrogen pressure and corresponding changes in hydrogen concentration in metal at a given temperature, PCT curves can be constructed that are expected to give a flat plateau. Most practical hydrides do not show perfectly flat plateau or zero hysteresis. Sloping behavior is observed possibly due to different equilibrium pressure, localized defects, and surface inhomogenities [21] (will be discussed in more detail in Sect. 2.1.4).

1.4

Important Hydride Properties and Experimental Techniques

57

H2 absorption

lnP

Hysteresis = lnPa /lnPd

H2 desorption

Slope = dlnPa / dHcap

H capacity wt.%

Fig. 1.22 Schematic isothermal pressure–composition hysteresis loop

Tc

T3

lnP = ΔH/RT - ΔS/T

↵ α+β

T2

↵

lnP

T1

a

↵ lnP

α

↵ β

−ΔH/R

↵

H capacity wt.%

↵

b

1/T

Fig. 1.23 (a) Pressure–concentration–temperature plot and (b) Van’t Hoff plot

The effect of temperature on the PCT curves is shown in Fig. 1.23a. The metal initially dissolves only small amount of hydrogen (<0.3 wt%), which creates a solid solution of hydrogen in a metal matrix α-phase. As the hydrogen pressure and hydrogen concentration in the metal are increasing, interactions between hydrogen and metal atoms become locally important and nucleation and growth of a new metal hydride β-phase is observed. In the plateau region there exists a mixture of solid solution α-phase and metal hydride β-phase. The length of plateau determines how much H2 can be stored reversibly with a small pressure variation. It can be seen in Fig. 1.23a that increasing temperature increases plateau pressure and beyond the critical temperature Tc, plateau region disappears and the α phase converts to the β phase continuously. The relation between midplateau pressure P and temperature T is given by the well known Van’t Hoff equation:

58

1

ln( P / P0 ) =

ΔH ΔS − RT R

Introduction

(1.32)

where P0 is atmospheric pressure, ΔH and ΔS are enthalpy and entropy changes of the hydriding/dehydriding reaction, respectively, T is the absolute temperature, and R is the gas constant (8.314472 J mol−1 K−1). For almost all hydrides (but a few exceptions exist) the enthalpy and entropy of hydriding reaction are negative, i.e., the hydriding reaction is exothermic and dehydriding reaction is endothermic. To avoid confusion, in this book both ΔH and ΔS are always expressed per 1 mol H2. The knowledge of ΔH is especially important to the heat management required for practical engineering devices and is a fundamental measure of the M–H bond strength. The enthalpy of absorption and desorption process, ΔH, can be determined from the slope (−ΔH/R) using the Van’t Hoff plot (logarithm of the midplateau pressure against the reciprocal temperature: lnP vs. 1/T (or more preferably 1,000/T) (where P is most conveniently given in atm) presented in Fig. 1.23b. The enthalpy term characterizes the stability of metal–hydrogen bond and the operating temperature of the metal hydride is fixed by the plateau pressure thermodynamically and by the overall reaction kinetics. The entropy term corresponds mostly to the change from molecular hydrogen gas to dissolved atomic hydrogen and is more or less constant for all hydrides. Hydrogen capacity HC in Fig. 1.24 can be expressed in either atomic H/M ratio (H – number of H atoms, M – number of metal atoms) or weight percent (wt%), both of which are commonly used [14]. It must be noticed that calculating wt% both mass of hydrogen mH and mass of metal mM (not only mass of metal) must be considered in the denominator. HC =

mH [wt%] mHydride

(1.33)

ln P

H2 absorption

H2 desorption Reversible capacity Δ(mH/mHydride)R

Max. Capacity Δ(mH/mHydride)max

Capacity HC wt.%

Fig. 1.24 Schematic presentation of hydrogen capacity defined by various methods

1.4

Important Hydride Properties and Experimental Techniques

59

where mHydride – mass of hydride = mM+ mH Capacity presentation in wt% is very useful from technological point of view, because it gives direct information on how much of hydrogen can be stored in a material. Regardless of the units, there are several ways to express hydrogen capacity. The reversible capacity Δ(mH /mHydride)R, is conservatively defined as the plateau width, which can be considerably less than the maximum capacity D(mH /mHydride)max (Fig. 1.24). In practice, depending on available pressure and temperature ranges, engineering capacity is usually between reversible and maximum capacity. Substituting P = 1bar (or 1 atm) in (1.32) one can find a simple relationship between the equilibrium temperature (Tplateau) required to give a midplateau pressure of 1 atm H2, ΔH and ΔS in the following form ΔH = ΔSTplateau

(1.34)

Equation 1.34 is plotted for a number of hydrides in Fig. 1.25. As can be seen all the data points fit very well in a simple straight line whose slope is equal to ΔS ≈ −130 J mol−1K−1 [162]. This clearly shows that the entropy term is, indeed, a nearly constant value for all the solid state hydrogen systems. Figure 1.25 also shows that a low desorption temperature at 1 atm of pressure (more or less an operating pressure of a PEMFC) can only be achieved with hydrides having the formation/decomposition enthalpies not larger than 50 kJ mol−1. For example, hydrides that desorb at room temperature such as LaNi5 and TiFe have ΔH ~ 30 and 33.3 kJ mol−1, respectively [163]. However, too small an enthalpy term would require Tplateau at 1 atm to be much below 0°C. From this point of view the enthalpy term is one of the most important factors characterizing any hydride.

Hydride formation enthalpy [kJ/mol H2]

0

TiCr1.8 FeTi LaNi5 CaNi5 PdH0.7 AlH3 Mg2Ni VH2 CsH ZrMn2

−50

MgH2

−100

NaH KH UH3

−150

−200

RbH

ZrNi

Lide Hydpark Mueller et al.

ZrH2 LiH TiH2

BaH2 CaH2

SrH2

0

500 1000 Plateau temperature [K]

1500

Fig. 1.25 Hydride formation enthalpy, ΔH, per mole H2 as a function of the plateau temperature at 1 bar. The plateau temperature is calculated from reported thermodynamic parameters using the Van’t Hoff equation [162]

60

1.4.1.2

1

Introduction

Calculation of Activation Energy

Thermodynamics can be used only to calculate the driving force for transformation but it cannot say how fast a transformation will proceed. The study of how fast process occurs belongs to kinetics. In Fig. 1.26 is shown the free energy curve of a single atom, which transforms from initially metastable state into a state of lower free energy. If GI and GF are the free energies of initial and final states, the driving force for the transformation will be ΔG = GF – GI. However, before the free energy of the atom can decrease from GI to GF the atom must pass trough a so-called activated state with a free energy ΔGA above GI. The energies shown in Fig. 1.26 are average energies associated with large number of atoms. As a result of random thermal motion of the atoms the energy of any particular atom will vary with time and occasionally it may be sufficient for the atom to reach the activated state. This process is known as thermal activation [164]. The two most popular methods of calculation of energy of activation will be presented in this chapter. First, the Kissinger method [165] is based on differential scanning calorimetry (DSC) analysis of decomposition or formation processes and related to these reactions endo- or exothermic peak positions are connected with heating rate. The second method is based on Arrhenius equation and determination of formation or decomposition rate from kinetic curves obtained at various temperatures. The critical point in this method is a selection of correct model to estimate the rate of reaction.

Calculation of Activation Energy by the Kissinger Method By performing the Kissinger analysis [165], i.e., an analysis of the sensitivity of the peak positions, in terms of Tmax the temperature of the peak maximum, to the applied heating rate, b, the apparent activation energy, EA, can be obtained from the following equation:

G

ΔGA GI

ΔG GF

Fig. 1.26 Transformation from initial to final state through an activated state of higher free energy

Initial state

Activated state

Final state

1.4

Important Hydride Properties and Experimental Techniques

61

⎛ b ⎞ d1n ⎜ 2 ⎟ ⎝ Tmax ⎠ E =− A R ⎛ 1 ⎞ d⎜ ⎟ T ⎝ ⎠

(1.35)

max

In Fig. 1.27 and Table 1.5 a typical DSC peak position shift related to various heating rates for thermal decomposition of LiAlH4 is presented. Thus, EA can be obtained as the slope in a plot of ln(b/Tm2) vs. 1,000/Tm (Fig. 1.28). Calculation of Activation Energy from the JMAK Model and the Arrhenius Equation The hydrogen absorption/desorption kinetics are usually analyzed by applying the JMAK (Johnson-Mehl-Avrami-Kolmogorov) theory of phase transformations, which is based on nucleation and growth events [166–168] where a is the fraction transformed at time t or alternatively for hydrides the fraction absorbed

I

20

DSC/(mW/mg)

15

4⬚C/min

6⬚C/min

10⬚C/min

exo III

10 5 0

IV V

II

−5 100

200

300

400

500

Temperature / ⬚C

Fig. 1.27 Dehydrogenation of as-received LiAlH4 investigated by DSC

Table 1.5 The temperatures of peaks connected with various processes during thermal dehydrogenation of as-received LiAlH4 investigated by DSC Temperature (°C)

Heating rate (°C min−1)

I

II

III

IV

V

4 6 10

150.76 150.32 153.69

166.04 167.33 166.66

174.83 180.90 186.67

237.35 238.30 246.98

447.09 449.26 466.59

1

ln(b /T 2)

62 −9.9 −10 −10.1 −10.2 −10.3 −10.4 −10.5 −10.6 −10.7 −10.8 −10.9 2.17

Introduction

y = −14.647x + 21.93 R2 = 0.9936

EA = 122 kJ/mol

2.18

2.19

2.20

2.21

2.22

2.23

2.24

2.25

1000 / T (K−1)

Fig. 1.28 Kissinger plot for dehydrogenation of as-received LiAlH4. Energy of activation for the 1 decomposition LiAlH4→ 3 Li3AlH6 + 2/3Al + H2 (peak III in Fig. 1.27)

h

α = 1 – e–(kt)

(1.36)

or desorbed at time t. It must be kept in mind that the JMAK model applies when growth of a new phase begins randomly in the bulk and at the surface (nucleation is spatially random), the sample size is much greater than any individual transformed region, growth proceeds homogeneously throughout the sample, and nucleation rate is constant [166–168]. The parameters describing the reaction kinetics, such as the nucleation and growth rates, are contained within an effective kinetic parameter, k, while the exponent,h, called the Avrami exponent or reaction order, provides some information about the dimensionality of the transformation, i.e., whether it is one-, two- or three-dimensional and whether it is interface-limited or diffusion-limited. Equation (1.36) can be rearranged to the following linear equation ln[–ln(1 – a)] = (hlnk) + hlnt

(1.37)

from which the values of the reaction order h and subsequently the rate constant k can be interpolated by plotting ln[-ln(1-a)] vs. ln(t). Such a plot for each constant temperature should give a straight line with slope h and intercept hln(k). From the latter, the rate constant k can be easily computed knowing the h value. It must be pointed out that only a nearly linear initial portion of the isothermal kinetic curve a vs. time (t) is to be taken into account for calculations. From our experience it is common that the h values can differ depending on the temperature for which they are being calculated. The different values of h suggest that different mechanisms are rate controlling of absorption/desorption at various temperature ranges. Therefore, we recommend the usage of free h values obtained from a double-logarithm fitting procedure as more true than the fixed h values [169].

1.4

Important Hydride Properties and Experimental Techniques

63

The apparent activation energy for the absorption/desorption process is usually evaluated from the Arrhenius plot of rate constant k values with temperature [166] by simply plotting a straight line lnk vs. 1/RT. k = koe–EA/RT

(1.38)

where EA is the apparent activation energy, R is the gas constant (8.314472 J mol−1 K−1), and T is the absolute temperature in K. The most popular and universal model describing absorption and desorption process in hydrides seems to be the JMAK and we will focus on calculation of energy of activation using this approach. However, it must be noticed that in some particular cases other mechanisms of hydride formation and decomposition can be applied. There are a few other kinetics models listed in Table 1.6. For example, Barkhordarian et al. [170, 171] and Liang et al. [172] have applied contracting volume model for the kinetic analysis of sorption process in MgH2 catalyzed by Nb2O5 and V, respectively. Nevertheless, in similar composites (Mg + Nb, MgH2 + V) [166, 173] the kinetics of sorption was analyzed using JMAK. We tested the three respective models from Table 1.6: JMAK, contracting volume 1 − [1 − a]1/2 = kt (two-dimensional growth with constant interface velocity) and surface reaction a = kt, for MgH2 hydride doped with 5 wt% of micrometric Ni catalyst. The results of these calculations are presented in Figs. 1.29 and 1.30. The obtained activation energy values equal 105, 101, and 105 kJ mol−1, respectively. This shows that in practical situation each of these three models gives almost identical activation energy.

2 375⬚C

1

350⬚C

ln[-ln(1-a)]

275⬚C

300⬚C

0

y = 1.22x − 4.08

−1

y = 1.01x − 3.86

−2 −3

y = 1.21x − 5.54 y = 1.64x − 8.89

−4 −5

325⬚C

0

1

2

3

4 ln(t)

y = 2.12x − 14.4

5

6

7

8

Fig. 1.29 Plot ln(−ln(1 − a)) vs. ln(t) for dehydrogenation of MgH2 (Tego Magnan®) milled for 20 h and catalyzed by 5 wt% Ni (tests were done in temperature range 275–375°C using nonactivated powder)

64

1

Introduction

−2 Surface controled y = -105260x + 16.149 R2 = 0.98

−3

lnk

−4

JMAK y = -105293x + 16.491 R2 = 0.98

−5

−6

CV (n = 2) y = -101264x + 14.728 R2 = 0.99

−7

-8 0.1800 0.1850 0.1900 0.1950 0.2000 0.2050 0.2100 0.2150 0.2200 0.2250 1000/RT

Fig. 1.30 Arrhenius plots for dehydrogenation of MgH2 (Tego Magnan®) milled for 20 h and catalyzed by 5 wt% Ni (tests were done in the temperature range 275–375°C using a nonactivated powder). Arrhenius plots were obtained using various models of decomposition process (Table 1.6)

Table 1.6 Kinetic equations used for fitting experimental hydrogen sorption data Model equation

Description

References

a = kt [−ln(1 − a)]1/n = kt

Surface controlled (chemisorption) Johnson-Mehl-Avrami-Kolmogorov (JMAK): n = 3 – three-dimensional growth of existing nuclei with constant interface velocity n = 2 – two-dimensional growth of existing nuclei with constant interface velocity Contracting volume (CV): n = 3 – three-dimensional growth with constant interface velocity n = 2 – two-dimensional growth with constant interface velocity

168, 170 166–168, 173

Contracting volume (CV): Three-dimensional growth diffusion controlled with decreasing interface velocity

168, 171

1 − 1n(1 − a)1/n = kt

⎛ 2a ⎞ 1 − ⎜ ⎟ − (1 − a )1/3 = kt ⎝ 3 ⎠

170–172

1.4

Important Hydride Properties and Experimental Techniques

1.4.2

65

PCT and Kinetic Curves Determination by Volumetric Method in a Sieverts-Type Apparatus

The hydrogen desorption or absorption PCT curves of various hydrides can be evaluated using a Sieverts-type apparatus. The scheme of typical Sieverts-type apparatus is shown in Fig. 1.31. The Sieverts-type apparatus consists of a calibrated volume determined physically, a reactor whose temperature is controlled by the temperature control system and the cooling system, a vacuum system, a pressure monitoring system, valves, and source of hydrogen and argon delivery. The quantity of desorbed hydrogen (number of molls) is calculated using ideal gas flow: PV = nRT

(1.39)

where P – gas pressure, V – gas volume, n – number of moles of gas, T – absolute temperature of gas, and R – the universal gas constant. The value and units of R depend on the units used in determining P, V, n, and T. ● ●

The quantity of gas, n, is normally expressed in moles. The units chosen for pressure and volume are typically atmospheres (atm) and liters (L), however, other units may be chosen.

Pressure Monitoring System Vv

T H2

vent

VH

VAr

VC

VR

Ar

VP

cooling water Vacuum System cooling water

VC

R

Cooling System Temperature Control System

Fig. 1.31 Scheme of Sieverts-type apparatus where T transducer, VH hydrogen cut-off valve, VAr argon cut off valve, VP vacuum system cut-off valve, VR reactor cut-off valve, VC calibrated volume and its cut-off valve, Vv vent valve, R reactor

66

1

Introduction

Therefore, R can be expressed, for example, in L atm mol−1 K−1 where R = 0.08206. For thermodynamic calculation the universal gas constant includes energy unit R = 8.314472 J mol−1 K−1. Let us assume that we can treat hydrogen as an ideal gas. Before beginning of absorption or desorption the relation between pressure of hydrogen in a system and number of moles of hydrogen at temperature T of the analyzed process can by described by: P1V = n1RT

(1.40)

After desorption or absorption we have: P2V = n2RT

(1.41)

where P1> P2 for absorption and P1< P2 for desorption. Rearranging (1.40) and (1.41) we obtain: n1 =

PV PV 1 n2 = 2 RT RT

(1.42)

Therefore, the difference between number of moles of hydrogen in the system resulting from absorption or desorption is as follows: Δ n = n1 – n2 = ΔP

V RT

(1.43)

where DP = P1 − P2. The mass of absorbed or desorbed hydrogen can be calculated using number of moles of gas and molecular mass of hydrogen: mH= 2.016Δn, which finally gives us: mH = 2.016ΔP

V RT

(1.44)

When change in hydrogen mass is known using (1.33) we can easily calculate hydrogen capacity in the investigated material. The ideal gas law should be corrected by the Van der Waals equation for the volume of gas molecules and molecular interactions at higher hydrogen gas pressures. ⎛ n2 a ⎞ ⎜⎝ p + V 2 ⎟⎠ (V − nb ) = nRT

(1.45)

where a is the measure of attraction between hydrogen particles (0.2476 L2 bar mol−2) and b is the volume excluded by a mole of hydrogen particles (0.02661 L mol−1). From our tests it seems that (1.45) should be used for pressures higher than about 5 MPa. From the scheme of Sieverts apparatus presented in Fig. 1.31, we can see that the volume that must be taken into account during calculation of quantity of desorbed or absorbed hydrogen does not consist of only a reactor volume. The volume V in (1.44) in a real system is made up of volume of reactor plus internal volume

1.4

Important Hydride Properties and Experimental Techniques

67

of connecting pipes, valves, as well as the transducer. The direct measurement of these parameters is impossible. Moreover, the transducer measures pressure changes in this part of apparatus where temperature is ambient, so for calculation using (1.44) we apply T = 297 K (ambient temperature). Nevertheless, increasing the reactor temperature to the temperature of process (usually in the range from 40 to 400°C) increases the temperature of hydrogen gas in the entire system. That is why an application of the ideal gas law would not be possible. Therefore, to evaluate the volume of gas that participates in sorption process the system volume calibration process is necessary. Calibration process consists of a few steps. At first we have to obtain advisable pressure P1 (usually about 10 atm) of argon in a calibrated volume VC and atmospheric pressure in the rest of the system (reactor and connections). Then by opening calibrated volume cut-off valve VC and by pressure reduction to the value P2 the total volume of the system for the apparatus with the relative transducer can be calculated using the formula: P1Vc = P2 (Vc + Vs )

(1.46)

where Vs is made up of volume of reactor plus internal volume of connecting pipes, valves, and the transducer. Rearranging (1.45) we obtain

Vs =

(P1 − P2 )V P2

c

(1.47)

To eliminate the temperature effect, the calibration curve must be obtained by repeating described process at various temperatures (RT to Tmax, where Tmax – maximum temperature of process). To determine PCT curve by volumetric method at first we have to know mass of analyzed powder (hydride or pure metal). The typical mass of powder used in volumetric method is in a range 50–500 mg and depends on VR (reactor volume with volume of connecting pipes, valves, and transducer). After the mass measurement, the powder is loaded into specimen holder and then it is placed in the Sieverts apparatus reactor. To prevent any oxidation and for safety reason the system must be purged a few times by argon and then evacuated. However, one must be careful how much powder is appropriate for the absorption/desorption volume of a Sieverts-type apparatus. Figure 1.32 shows the desorption experiments from two different masses of MgH2 doped with nano-Ni (n-Ni) powder. It is seen that the mass of 140 g gives about 3.0 wt% of desorbed hydrogen while the smaller mass of powder (50 mg) let more hydrogen to be desorbed. This behavior can be explained if we resort to the imaginary desorption PCT curve at 275°C in Fig. 1.33. The equilibrium plateau pressure at 275°C is higher than 0.1 MPa at which desorption is carried out and at this temperature MgH2 can desorb at atmospheric pressure. However, the kinetics of desorption will depend on the driving force (as shown in Fig. 1.26). Since a larger mass of hydride will produce

68

1

Introduction

Hydrogen desorbed [wt.%]

8.00 MgH2 + n-Ni milled for 10 min - 275⬚C

7.00 6.00

4.6wt.%H2 Mass = 50 mg

5.00 4.00 3.00

3.0wt.%H2

2.00 Mass = 140 mg

1.00 0.00 0

1000

2000 Time [s]

3000

4000

Fig. 1.32 The effect of the mass of desorbing hydride powder on the amount of hydrogen desorbed at 275°C at 0.1 MPa from MgH2 doped with 5 wt% of nanonickel (n-Ni)

lnP Very small driving force PCT at 275ⴗC Mass >100 mg Mass 50 -100 mg 1 atm (0.1MPa)

Reasonably large driving force

wt.%H2

Fig. 1.33 Schematic plot showing the effect of mass of MgH2 hydride on the driving force for desorption at 275°C at atmospheric pressure

a higher overpressure, much above 0.1 MPa, in a confined desorption volume of a Sieverts apparatus, which is closer to the plateau at 275°C, then kinetics could be slower and as a result a smaller amount of hydrogen could be desorbed as opposed to a small amount of hydride creating lower overpressure and higher driving force allowing more complete desorption. When we want to obtain a PCT curve for decomposition process the pressure of hydrogen must be set up above the expected equilibrium pressure to prevent decomposition during heating to this temperature. If we do not have any information about equilibrium pressure for investigated material we should apply maximum pressure

1.4

Important Hydride Properties and Experimental Techniques

69

that is allowable by our Sieverts apparatus. High hydrogen pressure should prevent decomposition of hydride during heating up to the temperature of the test. When temperature of the test is reached and the system is thermally stable we can start gradually decreasing the pressure and registering the amount of released hydrogen related to the equilibrated pressure. In a case of the PCT absorption curve after evacuation, the system must be heated up to a test temperature and thermally stabilized. Then by gradually increasing pressure we will observe hydrogen absorption related to a particular pressure of hydrogen. It is important that the time of visible pressure change at every step is closely connected with the kinetics of process at an applied temperature. An analysis of absorption and desorption kinetic is very similar to the PCT curve determination. Before starting the desorption test, the inner tubing of the apparatus is also evacuated and purged a few times with argon and then two times with hydrogen. Subsequently, hydrogen at a high pressure, preventing desorption, is admitted into the airtight cylindrical desorption chamber containing the hydride powder and the temperature of the chamber is gradually increased up to the desired desorption temperature. The high pressure barrier is applied to prevent any desorption during the period of temperature stabilization (15–30 min.) within the desorption chamber. Once the temperature is stabilized the hydrogen pressure is quickly reduced to 0.1 MPa (1 atm), and the hydrogen desorption process begins (pressure increases above initial 0.1 MPa). No elaborate activation procedure usually is applied to the powders. Absorption test starts with purging process as well as evacuation and then system is thermally stabilized under vacuum. Subsequently, hydrogen at desired absorption pressure is admitted into the system and by observing the pressure decreasing as a function of time, the kinetic curve is registered. The difference in pressure between the initial pressure and that at the instant of time is taken for calculation of the mass of absorbed/desorbed hydrogen from (1.44) and the wt% of absorbed/desorbed hydrogen (capacity) from (1.33). However, a question arises whether or not the hydride annealing under highpressure barrier that is applied to prevent any desorption during the period of temperature stabilization (15–30 min) can influence the sorption properties? Bouaricha et al. [174] measured PCT and kinetics of hydrogen sorption using activated samples. The samples were activated at 350 or 400°C, by exposure to hydrogen at a high pressure. This was done until the hydrogen capacity of the material reached a maximum and did not evolve anymore with time. Generally, this state was reached after 12 h. However, it is clearly visible that time applied for activation process using high-pressure annealing must be definitely longer than approximately 30 min. To be absolutely sure that thermal stabilization of measuring system under high preventing pressure of hydrogen does not affect hydrogenation/dehydrogenation properties of analyzed hydride we studied some properties, important from a decomposition point of view, for powder before and after annealing process. Magnesium hydride (type ABCR – see Chap. 2) ball milled for 100 h under HES 57 mode (Sect. 1.3.2) was heated up to various temperatures (300, 325, and 350°C) under preventing 4.5 MPa pressure of hydrogen in a Sieverts apparatus exactly in this same manner as during thermal stabilization. To eliminate any doubts the annealing time was increased up to 45 min. After annealing, the powder was cooled

70

1

Introduction

down up to ambient temperature and investigated by DSC and XRD analysis. Figure 1.34 shows the influence of annealing temperature during thermal system stabilization on DSC test results. A slight increase of both onset and peak temperature is observed, which means that tendency to the decomposition of annealed powder is even less than that for the powder directly after milling. The structural analysis by XRD method (Fig. 1.35 and Table 1.7) shows progressive decrease of the amount of γ-MgH2 (will be discussed in more detail in Sect. 2.1) in powder after annealing as compared with phase composition in a powder directly after milling. Moreover, slight increase of grain size of β-MgH2 during thermal stabilization of a Sieverts apparatus is observed. However, the crystallite (grain) size of magnesium hydride is still in the nanorange (35 nm for powder annealed up to 350°C). These observations clearly prove that the thermal stabilization of Sieverts apparatus under high pressure of hydrogen carried out under previously specified conditions can not be treated as an activation process and it causes negligible structural changes in the investigated hydride. Finally, a small but important technical detail related to the method of measuring pressure by a transducer attached to a Sieverts apparatus (T in Fig. 1.31) may be in place here. Some transducers measure pressure relative to the atmospheric pressure and display a 0 value when pressure equals atmospheric pressure. In this case, during constructing PCT curve 1 atm should be always added to each pressure reading from the transducer’s display. If a Sieverts’ system is equipped with a transducer measuring and displaying the value of 1 atm at an atmospheric pressure, this correction is redundant. Also, this correction is redundant during measuring in hydrogen absorption/desorption kinetic curve because in this case what is needed for calculation of the hydrogen quantity is just ΔP (1.44).

420

Temperature [ⴗC]

400 380 360 340

Peak Onset

DSC analysis

320 300 Milled

300

325

350

Annealing temperature [ⴗC]

Fig. 1.34 The influence of annealing temperature on onset and peak temperature of decomposition peak obtained by DSC test

1.4

Important Hydride Properties and Experimental Techniques

71

10000 b-MgH2 g-MgH2 MgO

8000

Counts

6000

Annealed 350ⴗC 4000 Annealed 325ⴗC 2000

Annealed 300ⴗC Milled

0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 1.35 XRD pattern of MgH2 hydride after milling and annealing in Sieverts apparatus during thermal stabilization of the measuring system for 45 min at 300, 325, and 350°C under high pressure of hydrogen (4.5 MPa)

Table 1.7 Phase composition and grain size of β and γ-MgH2 after milling and annealing in Sieverts apparatus during thermal stabilization of the measuring system, at 300, 325, and 350°C under high pressure of hydrogen (4.5 MPa) Sample

Grain size of β-MgH2 (nm)

Grain size of γ-MgH2 (nm)

Phases present in powder

Milled Annealed 300°C Annealed 325°C Annealed 350°C

13 ± 2 19 ± 1 28 ± 1 30 ± 2

13 ± 0.1 19 ± 0.3 29 ± 2 –

β-MgH2, γ-MgH2, MgO β-MgH2, γ-MgH2, MgO β-MgH2, γ-MgH2, MgO β-MgH2, MgO

1.4.3

Microstructural Characterization of Ball-Milled Hydrides

Two important morphological parameters characterizing ball-milled powders are the particle and grain size of constituent phases within the powders. In our laboratory, the size measurement of the powder particles is carried out by attaching loose powder to sticky carbon tape and taking pictures under secondary electron (SE) mode in the SEM. The images are then analyzed by an image analysis software. The size of the powders is calculated as the particle equivalent circle diameter, ECD = (4A/p)1/2, where A represents the projected particle area. Usually from ~300 to 700 particles are analyzed for each batch.

72

1

Introduction

The crystalline structure of hydride powders is characterized by powder diffraction. The nanograin (crystallite) size of phases residing in the milled powders is calculated from the broadening of their respective X-ray diffraction (XRD) peaks. Since the Bragg peak broadening in an XRD pattern is due to a combination of grain refinement (nanograin/crystallite) and lattice strains, it is customary to use computing techniques by means of which one can separate these two contributions. The separation of crystallite size and strain is obtained from a Cauchy/Gaussian approximation by a linear regression plot according to the following equation [39]: d 2 (2q ) Kl ⎛ d (2q ) ⎞ = + 16e2 tan 2 q L ⎜⎝ tan q sin q ⎟⎠

(1.48)

where the term Kl/L is the slope, the parameter L is the mean dimension of the nanograin (crystallite) composing the powder particle, K is a constant (≈1) and e is the so-called maximum microstrain (calculated from the intercept), l is the wavelength, and q is the position of the analyzed peak maximum. The term d(2q) = B[1−(b2/B2)] (rad) is the instrumental broadening-corrected pure XRD peak profile breadth [39], where B and b are the breadths in radians of the same Bragg peak from the XRD scans of the experimental and reference powder, respectively. The B value is approximated as the full width at half maximum, FWHM, and calculated by the diffractometer software. The b value is approximated as FWHM from the X-ray diffraction pattern of a compound LaB6, the National Institute of Standards and Technology (NIST) standard reference material (SRM) 660 for subtracting the instrumental broadening from the experimental FWHM and finding d(2q) as given earlier. It must be noted that when FWHMs of the instrumental line profiles are obtained in this manner, the Bragg peaks for the LaB6 SRM are occasionally at different 2q angles than those of the analyzed hydride in the milled powders. The interpolated FWHM values between angles for the SRM peaks are found using a calibration curve. A typical regression plot according to (1.48) is shown in Fig. 1.36. 0.00035

d2(2q)/tan2(q)

0.0003 0.00025 0.0002

kλ/L = 0.0039

L = 39nm

λCu = 0.1541838 16e2 = 1.0E-05 e = 9.0E-04

y = 0.0039x + 1E-05 R2 = 0.99

0.00015 0.0001 0.00005 0 0.01

0.02

0.03

0.04

0.05

0.06

0.07

0.08

d(2q) / [tan(q)sin(q)]

Fig. 1.36 Cauchy/Gaussian plot (1.48) for the grain size of dehydrogenation of MgH2 (ABCR) milled for 20 h and catalyzed by 5 wt% Ni

1.4

Important Hydride Properties and Experimental Techniques

73

It must be also noted that modern X-ray diffractometers are usually equipped with software that calculates grain size and lattice strain.

1.4.4

Weight Percent of a Hydride Phase and Hydrogen by DSC Method

In DSC measurements, the weight percent of a phase was calculated using the peak area of the DSC curve and its reported heat of formation. For example, weight percent of the β-MgH2 in a reactively milled powder can be estimated using the peak area of the DSC curves and the reported β-MgH2 heat of formation (−74 kJ mol−1 [175], which equals to −2,811 J g−1). The DSC curve was analyzed by the NETZSCH thermal analysis software. First, the onset and end temperature of the peak were determined. Then, the peak area was calculated using the linear approach from the onset temperature to the end temperature (Fig. 1.37) by the DSC software. The weight percent of β-MgH2 is given by: wt% of β-MgH2 = peak area (J g−1)/β-MgH2 heat of decomposition (J g−1), where heat of decomposition = −(heat of formation). From the decomposition of β-MgH2 (MgH2 → Mg + H2), the weight percent of desorbed hydrogen can be calculated by: wt% of H2 = wt% of MgH2 × (molecular weight of H2/molecular weight of MgH2).

DSC /mW/mg 2.5

[1] 332.8⬚C

exo

2.0

1.5

1.0

0.5 [1] 346.9⬚C

[1] 321.5⬚C

0

[1]

280

300

320 Temperature /ⴗC

340

360

Fig. 1.37 Schematic of calculating the weight percent of a hydride phase by DSC analysis

74

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Introduction

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Chapter 2

Simple Metal and Intermetallic Hydrides

2.1 2.1.1

Mg/MgH2 Crystallographic and Material Characteristics

Reaction of hydrogen with the elemental Mg is one of the most widely researched reactions in the field of solid-state hydrogen storage. The Mg–H system is quite simple as shown in the binary phase diagram in Fig. 2.1. At moderate hydrogen pressures the only hydride phase existing in equilibrium with Mg is magnesium dihydride, MgH2, more commonly referred to as “magnesium hydride.” Table 2.1 shows the crystal structure data of the phases existing in the Mg–H system. Pure Mg has a hexagonal crystal structure and its hydride has a tetragonal lattice unit cell (rutile type). The low-pressure MgH2 is commonly designated as β-MgH2 in order to differentiate it from its high-pressure polymorph, which will be discussed later. Figure 2.2 shows the crystal structure of β-MgH2 where the positions of Mg and H atoms are clearly discerned. Precise measurements of the lattice parameters of β-MgH2 by synchrotron X-ray diffraction yielded a = 0.45180(6) nm and c = 0.30211(4) nm [2]. The powder diffraction file JCPDS 12–0697 lists a = 0.4517 nm and c = 0.30205 nm. The density of MgH2 is 1.45 g/cm3 [3]. Noritake et al. [2, 4] investigated the bonding nature of MgH2 employing the maximum entropy method and synchrotron radiation powder data. They found that the bonding nature of hydrogen in MgH2 was quite complex, consisting of the mixture of ionic and covalent bonding. As stated by the authors, there are weak but significant covalent bonds between Mg and H as well as between H and H. The weak covalency of the Mg–H bond may be advantageous on hydrogenation/ dehydrogenation performance. Modeled charge density distribution revealed that the ionic charge of Mg and H can be represented as Mg1.91+ and H0.26−, respectively. It means that Mg is ionized almost as Mg2+, while hydrogen is very weakly ionized. They also found weak but significant bonds between Mg and H as well as between H and H. Under increasing hydrogen pressure, substantial changes occur in the Mg–H system. Bastide et al. [5] investigated the behavior of MgH2 phase under high pressures up to 80 kbar and found that at ambient temperature (20°C) and 80 kbar of R.A. Varin et al., Nanomaterials for Solid State Hydrogen Storage, DOI: 10.1007/978-0-387-77712-2_2 © Springer Science + Business Media, LLC 2009

83

84

2

0

1

675

Simple Metal and Intermetallic Hydrides

Weight Percent Hydrogen 2 3 4

650⬚C

6

7 8

P = 25 MPa

L + H2 643⬚C

L

5

625 (Mg) + H2

575

566⬚C

525 MgH2

Temperature, ⬚C

(Mg)

475

425

375 0 Mg

10

20

40 50 30 Atomic Percent Hydrogen

60

70

Fig. 2.1 Mg–H binary phase diagram. Reprinted with permission of ASM International. All rights reserved. www.asminternational.org [1]

t.1 t.2 t.3 t.4 t.5 t.1 t.2 t.3 t.4 t.5

Table 2.1 Crystallographic data of the phases in the Mg–H system [1] Phase

Composition at.% H

l

Space group

Strukturbericht designation

Prototype

(Mg) MgH2

0–11 66.7

hP2 tP6

P63/mmc P42/mnm

A3 C4

Mg TiO2 (rutile)

Fig. 2.2 The crystal structure of b-MgH2 [2]

2.1

Mg/MgH2

85

pressure the β-MgH2 phase (designated α-MgH2 in the original paper [5]) transformed partially into another polymorphic phase γ-MgH2 forming a mixture (β + γ)-MgH2. X-ray diffraction (XRD) studies showed that γ-MgH2 has an orthorhombic unit cell structure of α-PbO2 type with the lattice parameters a= 0.453 nm, b= 0.544 nm, and c = 0.493 nm (this phase is included in the powder diffraction file JCPDS 35–1184). They also found another metastable phase δ-MgH2 (designated β-MgH2 in the original paper [5]), which upon release of pressure transformed into γ-MgH2. They claimed that above 350°C, γ-MgH2 was supposedly transformed into the equilibrium β-MgH2 phase. However, as will be shown later, the γ-phase decomposes above 300°C. Hydrogen absorption and desorption characteristics of two different commercial MgH2 powders were investigated in our laboratories. The first powder was purchased from Degussa-Goldschmidt, sold under the trade name “Tego Magnan.” Its average purity claimed by the supplier is ~95% (the remaining Mg) but we found that the powder was slightly more inhomogeneous than the ABCR (a German company) powder. The theoretical purity-corrected hydrogen content is ~7.2 wt.%. The powder particle size was provided by the supplier as 325 mesh. Scanning electron micrograph of the morphology of the Tego Magnan powder is shown in Fig. 2.3a. The size of the powders was calculated as the particle equivalent circle diameter (ECD) as described in Sect. 1.4.3 (~260 to 650 individual particles were usually analyzed). The size distribution is plotted in Fig. 2.3b. It is observed that the ECD distribution is well fitted by a lognormal curve (marked by a solid line in Fig. 2.3b) as confirmed by a quantitative statistical analysis. It must be pointed out that a lognormal ECD distribution has been a common characteristic of all MgH2 powders either commercial or synthesized from the elemental Mg which we have investigated so far. The evidence will be presented throughout this book. From the statistical point of view, a lognormal distribution of the powder particle sizes seems to be an indicator that two populations of particle sizes are present in a powder, one small and the other large, whose closely spaced normal distributions partially overlap giving in effect a lognormal distribution curve. Hence, the calculated mean particle size (ECD) and its standard deviation (SD) should only be considered as an estimate since according to the statistical point of view the concept of the mean value is valid only for normal (Gaussian) distribution. Nevertheless, the mean ECD is still a useful estimate for comparing powders subjected to various processing, e.g., ball milling. The mean value of the particle size of as-received Tego Magnan MgH2 is 36 μm with SD = ±16 μm (Fig. 2.3b). The grain (crystallite) size and lattice strain of the as-received Tego Magnan MgH2 were determined from the broadening of the X-ray diffraction (XRD) peaks using a Cauchy/Gaussian approximation as described in Sect. 1.4.3. The estimated value of ~67 nm (no lattice strains) [6] is at the border of nanocrystallinity if one defines it as the grain size smaller than 100 nm [7]. Our result correlates very well with the grain size of 78 nm reported very recently by Kojima et al. [8] for their commercial MgH2. Both results suggest that certain commercial varieties of MgH2 could be subjected to either ball milling or other type of postdeformation in the proprietary manufacturing process, which results in the final nearly nanosize grains (crystallites).

86

2

Simple Metal and Intermetallic Hydrides

Fig. 2.3 (a) Scanning electron micrograph of the morphology of as-received Tego Magnan MgH2 powder and (b) powder particle size distribution (equivalent circle diameter, ECD)

The second type of MgH2 powder was purchased from ABCR GmbH & Co. KG, sold under the trade name MG-5026. Its average purity claimed by the supplier is ~98% (remaining Mg). For simplicity, it will be referred to hereafter as the “ABCR powder.” Volumetric desorption tests were very reproducible, giving the average value of desorbed hydrogen equal to ~7.5 wt.% which is nearly identical to the theoretical purity-corrected capacity of 7.51 wt.%H2 at 98% purity. This testifies to

2.1

Mg/MgH2

87

a high homogeneity of the powder. Since its morphology and particle size (ECD) distribution are identical to those shown in Fig. 2.3, they are not shown here. Its mean ECD value is 41 μm with SD = ±21 μm. Indeed, particle size is very close to that of Tego Magnan. However, the grain (crystallite) size determined from the XRD peak broadening is ~300 nm (no lattice strain). This is a major microstructural difference between ABCR and Tego Magnan MgH2 powders.

2.1.2

Hydrogen Storage Characteristics of Commercial Mg and MgH2

2.1.2.1

Absorption

Hydrogenation/dehydrogenation studies of commercial Mg date almost 50 years back to the pioneering studies of Stampfer et al. [9] and Kennelley et al. [10]. The former authors reported the enthalpy of MgH2 formation being in the range of –70.8 to –72.6 kJ/mol and corresponding entropy as –127.3 to −130.6 J/molK. The latter authors reported the heat (enthalpy) of formation of MgH2 as −74.1 ± 2.9 kJ/mol. Vigeholm et al. [11–16] carried out very thorough investigations of hydrogen absorption/desorption of commercially pure Mg. They concluded that any comminuted Mg product with characteristic dimensions of less than approximately 75 μm and with impurity concentrations and compositions typical of normal Mg products will absorb hydrogen readily above 300°C at pressures exceeding the equilibrium level (i.e., plateau pressure on a corresponding pressure–composition–temperature (PCT) curve). They claimed that activation was in general either unnecessary or required only a few cycles under normal absorption conditions. The highest absorption rate was observed at an optimal temperature of around 340–350°C. Increasing the temperature above this value had slightly adverse effect on absorption kinetics. At ~263°C and 1.2 MPa hydrogen pressure, the magnesium powder absorbed ~2.3 wt.%H2 in roughly 6,000 s. At ~340°C, the same powder absorbed ~5.4 wt.%H2 in roughly 3,600 s. At the range of absorption temperatures up to ~340°C, the increase of pressure above 2 MPa increased the initial rate of kinetics but simultaneously reduced the degree of conversion of Mg into MgH2. In general, in order to obtain the degree of reaction close to 100% (i.e., Mg completely converted to a nearly single phase MgH2), the absorption temperature had to be higher than 400°C at a moderate hydrogen pressure of about 2.0–2.5 MPa. For example, in the initial absorption process at 400°C (after initial activation procedure by annealing in helium for 1 h), it was observed that the lower the hydrogen pressure, the closer the ultimate reaction approached stoichiometric hydride: 97.6% at 2.0 MPa, 95.5% at 3.2 MPa, and 81.5% only at 4.8 MPa [16]. In absorption process, the reaction of magnesium with hydrogen is a nucleation and growth mechanism where the nucleation rate is pressure dependent. They estimated the enthalpy and corresponding entropy of MgH2 formation as −70.0 kJ/mol and −126 J/mol K, respectively.

88

2

Simple Metal and Intermetallic Hydrides

More recently, very detailed thermodynamic studies of the Mg–H system were carried out by Bogdanovic´ et al. [17]. They specifically focused their attention on the possible effect of particle size on thermodynamics. Fractions with particle sizes of 25–40 μm and 70–100 μm were chosen for PCT determinations. Prior to all measurements, Mg powder fractions were activated. However, the range of selected temperatures 400–520°C was rather high. The hysteresis between absorption and desorption isotherms was very small. The reaction enthalpy and entropy for different hydrogen loadings and temperatures in the plateau region were determined using the Van’t Hoff relationship. A summary of their results is listed in Table 2.2. Thermodynamic parameters calculated for a hydrogen content of 50% (average value) are marked (calc.). It is quite clear that within the range of investigated particle sizes and temperature there is no dependence between particle size and thermodynamic parameters. In our laboratory, absorption experiments were carried out on an unmilled but activated ABCR MgH2 powder. In the preliminary stage, the powder was heated up to 350°C in a volumetric Sieverts-type apparatus (Sect. 1.4.2) under a preventing pressure of 3.5 MPa (so as not to allow desorption), held for about 30–45 min to stabilize temperature, and then subjected to the following desorption/absorption cycles: 1. Desorption at 350°C; annealing at 350°C/pre-vacuum/15 min; cooling to 325°C under pre-vacuum; absorption at 325°C/1.2 MPa/4,750 s; 2. Desorption at 350°C; annealing at 350°C/pre-vacuum/15 min; cooling to 300°C under pre-vacuum; absorption at 300°C/1.2 MPa/4,750 s; 3. Desorption at 350°C; annealing at 350°C/pre-vacuum/15 min; cooling to 275°C under pre-vacuum; absorption at 275°C/1.2 MPa/4,750 s; 4. Desorption at 350°C; annealing at 350°C/pre-vacuum/15 min; cooling to 250°C under pre-vacuum; absorption at 250°C/1.2 MPa/4,750 s. Consequently, the powder which absorbed at 250°C went through four full cycles as described above. Figure 2.4a shows typical kinetic curves of absorption at various temperatures. The absorption curves shown in Fig. 2.4a are quite similar to the ones reported by Vigeholm et al. [12] under the same 1.2 MPa hydrogen pressure as the one used in our work. In our case, about 4.5 wt.%H2 is absorbed at 325°C in ~4,500 s vis-à-vis about 4.5 wt.%H2 at 312°C absorbed in 4,000 s as reported by Vigeholm et al. At 263°C, Vigeholm et al. reported ~2.3 wt.%H2 absorbed in ~6,000 s which compares favorably with ~2 wt.%H2 absorbed at 250°C in 4,500 s in our experiments (Fig. 2.4a).

Table 2.2 Dependence of formation enthalpy and entropy of MgH2 (average values) upon the particle size of Mg powders used [17] ΔHf (calc.) Particle size ΔSf(calc.) ΔSf(J/molK) Mg (μm) Temp. (°C) ΔHf (kJ/mol) (kJ/mol) (J/mol K) 25–40 70–100

400–520 400–520

−72.6 ± 1.7 −74.7 ± 1.7

−132.2 −134.7

−72.9 −72.9

−132.3 −132.3

2.1

Mg/MgH2

89

5.00

1-250ⴗC 2-275ⴗC 3-300ⴗC 4-325ⴗC

Absorbed H2 (wt.%)

4.50

4

4.00

2

3.50

3

3.00 2.50 2.00

1

1.50 1.00 0.50 0.00 0

1000

a

2000 Time (s)

3000

4000

−3 −4

y = -64703x + 7.0755 R2 = 0.9969

ln k

−5 −6 −7 −8 −9 0.00019

b

0.0002

0.00021

0.00022

0.00023

0.00024

1/RT

Fig. 2.4 (a) Absorption kinetic curves of unmilled but activated and cycled ABCR powder and (b) estimate of apparent activation energy of absorption from the Arrhenius plot of ln k vs. 1/RT using data for all four temperatures: 250, 275, 300, and 325°C (EA ~65 kJ/mol). Coefficient of fit R2 = 0.997

Fernández and Sánchez [18] investigated the kinetics of hydrogen absorption and desorption by activated magnesium powder (several cycles of hydrogen absorption/desorption at 375°C) using a volumetric technique. They pointed out that formation/decomposition of metal hydrides comprises a number of steps taking place in series: H2 transport to the surface, H2 dissociation, H chemisorption, surface– bulk migration, H diffusion and nucleation, and growth of hydride/metal (absorption/desorption) phases. A summary of the detailed mechanisms taking place during nucleation and growth of MgH2 upon absorption and their dependence on absorption pressure (temperature) was provided by Friedlmeier and Groll [19]. Within the framework of the nucleation and growth model, they divided an isothermal absorption event into two cases depending on the applied hydrogen pressure as visualized in Fig. 2.5.

90

2

Mg

b

Mg

1

0

a

b

a

2nd Phase

a

1st Phase

1

Mg

2nd Phase

b

1st Phase

Mg

Simple Metal and Intermetallic Hydrides

t

0

t

b

Fig. 2.5 Schematic of the cross-section of an Mg particle at two different times and typical α = f(t) curves for hydriding at (a) pappl ≈ ppl and (b) pappl >> ppl (pappl – applied hydrogen pressure at T = const; ppl – plateau pressure at the same T = const; α – fraction transformed; β – β-MgH2 hydride) (adapted from [19])

Figure 2.5a shows that for an applied hydrogen pressure close to the equilibrium plateau pressure, ppl, the rate-controlling step is the chemical reaction which is proportional to the Mg: b-MgH2 interface area. The absorption curve is then sigmoidal. In this case, only a few hydride grains nucleate and they subsequently grow to relatively large sizes until a completely closed hydride layer is formed within the particle (first phase in Fig. 2.5a). The situation is similar to the solidification from the melt with a very small undercooling (a limited number of nuclei). The transformation mechanism in the first phase results in a large hydrogenated volume of a Mg particle. Normally, hydrogen diffusion takes place either through Mg phase or, alternatively, through the Mg:β-MgH2 interface. However, once the massive hydride layer is already formed, the phase transformation of Mg into β-MgH2 essentially stops because of the virtual impermeability of the β-MgH2 layer to hydrogen diffusion. It must be recalled now that the hydrogen diffusion rate in β-MgH2 is a few orders of magnitude lower than that in metallic Mg [20]. A very slow diffusion of hydrogen is now the rate-controlling reaction step (second phase in Fig. 2.5a). As shown in Fig. 2.5b, for high applied hydrogen pressures (pappl >> ppl), a large amount of β-MgH2 nuclei is formed similar to the nucleation during solidification from the melt under a large undercooling. The growth rate is much smaller as compared to that in Fig. 2.5a and only a very thin layer of β-MgH2 can be formed near the surface of each particle until the reaction stops (first phase in Fig. 2.5b). The first phase is most probably surface controlled (chemisorption). The complete chemical reaction is excluded since no sigmoidal curve is observed. The reaction in

2.1

Mg/MgH2

91

the first phase results in a small hydrogenated volume of a Mg particle, although the initial kinetics (the slope of the linear portion of the curve in Fig. 2.5b) is faster than that in Fig. 2.5b. Furthermore, as pointed out by Friedlmeier and Groll [19], the mechanisms delineated in Fig. 2.5 mean that, in practice, if one allows a sufficiently long period of time, higher hydriding capacities will be achieved for lower pressures. The higher the applied pressure, the faster the start of the hydriding reaction, but the lower the hydriding capacities achieved. All this leads to a crossing of isothermal α = f(t) curves obtained under various applied hydrogen pressure, the effect which is not observed for most hydride materials. Karty et al. [21] pointed out that the value of the reaction order η and the dependence of k on pressure and temperature in the JMAK (Johnson-Mehl-AvramiKolmogorov) equation (Sect. 1.4.1.2), and perhaps on other variables such as particle size, are what define the rate-limiting process. Table 2.3 shows the summary of the dependence of η on growth dimensionality, rate-limiting process, and nucleation behavior as reported by Karty et al. [21]. The absorption curves in Fig. 2.4a were analyzed by a linear fitting to the JMAK equation from which the reaction rate constant, k, and the reaction order,η, can be determined. The values of the reaction order η are listed in Table 2.4. At the lowest absorption temperature of 250°C, the η parameter is close to 2, and then it decreases to about 1, remaining close to this value at all the other temperatures. The different values of the reaction order suggest that different mechanisms are controlling the rates at various temperature ranges of absorption. It can be seen from Table 2.3, that the value of ~2 (1.83 in Table 2.4) suggests that the transformation of Mg to Table 2.3 Dependence of η in the JMAK equation [21] Growth Rate limiting process dimensionality Diffusion 1 2 3 Interface transformation 1 2 3

η

Constant nuclei number 0.5 1 1.5 1 2 3

Table 2.4 The values of the reaction order η in the JMAK equation for the absorption experiments on the unmilled, activated, and cycled ABCR powder presented in Fig. 2.4 Absorption temperature (°C) Reaction order η 200 250 275 300 325

no absorption 1.83 0.91 0.74 0.84

Constant nucleation rate 1.5 2 2.5 2 3 4

92

2

Simple Metal and Intermetallic Hydrides

β-MgH2 upon absorption is either diffusion-rate limited occurring by two-dimensional growth at constant nucleation rate, or alternatively, it is interface-controlled transformation with one-dimensional growth at constant nucleation rate or two-dimensional growth at constant nuclei number. From the standpoint of the previous discussion of the transformation mechanisms shown in Fig. 2.5 and a nearly sigmoidal shape of the transformation curve in Fig. 2.4a, it is quite likely that at the lowest temperature of 250°C the Mg to β-MgH2 transformation during absorption is interface controlled. However, one-or two-dimensional growth is inconsistent with microscopic observations of three-dimensional hydride grain growth [19]. The η value close to 1 suggests either diffusion rate–limited transformation occurring by two-dimensional growth at constant nuclei number or, alternatively, interface-controlled transformation with one-dimensional growth at constant nuclei number. A word of caution is needed here. As pointed out by Karty et al. [21], the listing in Table 2.3 does not exhaust all possible η behavior, but it does illustrate that in general the determination of η does not provide a unique identification of the kinetics. However, the two extreme η values of 0.5 and 4 do more or less define unique processes. Karty et al. reported that the value of η for hydriding (absorption) of Mg2Cu-catalyzed Mg was equal to 0.5. In turn, η was equal to 1 for hydrogen absorption transformation of vapor-deposited Mg. The latter is in good agreement with the results shown in Table 2.4 for the temperature range higher than 250°C. Fernández and Sánchez [18] reported that a nucleation and growth (NG) mechanism as expressed by the JMAK equation with exponent values of η = 0.5–1 for absorption gave the best fit to the experimental data. The apparent activation energy of absorption in Fig. 2.4a was estimated through the Arrhenius plot of rate constant, k, with temperature (Fig. 2.4b). All the fitting procedures have already been described in Sect. 1.4.1.2. The most surprising result is a relatively low value of the apparent activation energy equal to ~65 kJ/mol. It must be pointed out that in view of the fact that the mechanisms of the transformation of Mg to β-MgH2 can depend on temperature range as clearly shown in Table 2.3, the estimated apparent activation energy of absorption reflects the occurrence of all those mechanisms. It must be mentioned that the elimination of the absorption curve at 250°C, which is characterized by the largest η = 1.83 (Table 2.3) from the estimation, does not change the final value of the apparent activation energy for absorption. Jensen et al. [22] compiled experimentally determined apparent activation energy values for hydrogenation/dehydrogenation of MgH2 from a number of published references. The apparent activation energy of the hydrogenation of cycle-activated Mg powder was reported as ~100 kJ/mol. For vapor-deposited Mg it was 138 kJ/mol. It must be pointed out that in our work absorption in Fig. 2.4a was conducted on a cycle-activated commercial MgH2 powder rather than on pure Mg. This factor could also contribute to the lowering of the apparent activation energy (Fig. 2.4b). The highest value of ~270–314 kJ/ mol was reported by Jensen et al. for deactivated (oxidized due to air exposure) Mg foil. According to Jensen et al., the major factor affecting the apparent activation energy and giving rise to a substantial scatter among values reported in the literature is the presence of MgO/Mg(OH)2 surface layers, retarding the diffusion

2.1

Mg/MgH2

93

of hydrogen into Mg (absorption) or out of MgH2/Mg (desorption). This factor will be discussed in more detail later on.

2.1.2.2

Desorption

As mentioned earlier, Kennelley et al. [10] carried out pioneering studies of dehydrogenation (desorption) of MgH2. They reported the equation for the equilibrium temperature–pressure relationship in the following form: log p (mm of Hg) = −3857/T + 9.78 where p is an equilibrium pressure and T is temperature in K for MgH2 decomposition. Later on, Stander [23] investigated the decomposition kinetics of MgH2 (92% purity) synthesized from the elemental Mg powder under 15 MPa hydrogen pressure at 400°C for 20 h. He reported that if the difference between the experimental dissociation pressures (e.g., 384 kPa) and the equilibrium dissociation pressure corresponding to T = const via PCT curve (e.g., 890 kPa at T = 368°C for MgH2 as determined by Stander) is large, the nuclei of Mg metal are formed virtually instantaneously at sites on the MgH2 particles and the rate of decomposition (kinetics) of MgH2 is controlled by movement of the cylindricalshaped Mg–MgH2 interface as given by the equation: 1 – (1 – α)0.5 = k1t

(2.1)

The apparent activation energy of decomposition estimated from the Arrhenius plot for k1 gave 120 kJ/mol. Conversely, Stander noticed that if the difference between the experimental dissociation pressures (e.g., 384 kPa) and the equilibrium (plateau) dissociation pressure corresponding to T = const (e.g., 404 kPa at T = 335°C for MgH2) is relatively small, then better fits were obtained with the model of random nucleation followed by one-dimensional growth or instantaneous nucleation followed by two-dimensional growth as given by the equation: (– ln (1 –α))0.5 = k2t

(2.2)

The Arrhenius plot for k2 yielded the apparent activation energy of 126 kJ/mol. Since a higher desorption temperature is usually related to a higher equilibrium pressure via Van’t Hoff equation, there is a tendency to attribute the kinetic effects to the difference in temperatures when desorption temperature is being changed from a relatively low value to a high value. However, one must remember that the underlying factor is still the difference between the experimental desorption pressure and the equilibrium dissociation pressure. Stander [23] concluded that slow movement of the phase boundary is the ratecontrolling step at pressures appreciably different from the equilibrium pressure. At pressures near the equilibrium pressure nucleation seems to play a role. The difference in activation energy for the two mechanisms cannot be regarded as being significant. Furthermore, as the reaction proceeds, nucleation will become less important. Eventually he concluded that nucleation is kinetically unimportant in the decomposition

94

2

Simple Metal and Intermetallic Hydrides

of MgH2. With regard to desorption, Vigeholm et al. [12] observed that the increase of desorption pressure from 0.15 MPa to over 1 MPa dramatically reduced the rate of desorption. Complete decomposition required high temperatures (in excess of 390°C to ensure completion in less than 10 min) and pressures below 100 kPa. In general, they found that during desorption the reversed nucleation took place. They also found that the reaction rates were not influenced significantly by the number of absorption/ desorption cycles within 150 cycles but the absorbed/desorbed amount of hydrogen dropped to 50–60% of the stoichiometric value up to 500 cycles. Figure 2.6a shows typical kinetic curves of first desorption carried out in a Sieverts-type apparatus at the initial hydrogen pressure of 0.1 MPa (atmospheric pressure of 1 bar) for the as-received, nonmilled, and nonactivated Tego Magnan powder. For each temperature, a fresh load of sample was desorbed. At each temperature in Fig. 2.6a, the desorption process is complete with 100% of MgH2 des-

Hydrogen desorbed [wt.%]

9

Desorption

8 7

4 3

6

1

2

5 4 1-350ⴗC 2-375ⴗC 3-400ⴗC 4-420ⴗC

3 2 1 0 0

1000

a

2000 Time [s]

3000

−3 −3.5

y = -120307x + 16.664 R2 = 0.9961

−4 ln k

−4.5 −5 −5.5 −6 −6.5 −7 0.00017

b

0.00018

0.00018 0.00019 1/RT

0.00019

0.0002

Fig. 2.6 (a) Desorption kinetic curves at various temperatures under initial hydrogen pressure of 0.1 MPa of the as-received, nonactivated, commercial MgH2 powder Tego Magnan and (b) the Arrhenius plot of the desorption rate for the estimate of the apparent activation energy, EA, using kinetics data for four temperatures: 350, 375, 400, and 420°C (EA ~120 kJ/mol). Coefficient of fit R2 = 0.996

2.1

Mg/MgH2

95

orbed (~7.6 wt.%). It must be pointed out that no desorption was observed at temperatures below 350°C after a testing time of a few hours. It is not to say that desorption would not occur after a much longer time at temperatures below 350°C. For instance, Vigeholm et al. [12] observed about 83% of hydrogen desorbed (~6.3 wt.%H2) after 15 h at ~320°C. As can be seen in Fig. 2.6a, at 350°C it takes 3600 s to completely desorb the powder (~7.6 wt.%H2 desorbed). However, an increase of temperature to 375°C and higher increases the rate of desorption enormously such that the powder is completely desorbed within 200–700 s range. The desorption kinetics observed in Fig. 2.6a is faster than that reported by Huot et al. [24] for an unmilled, commercial purity MgH2 (~95%). At 350 and 375°C, their unmilled MgH2 desorbed ~5.0 wt.%H2 in 2000 s and ~6.5 wt.%H2 in ~900 s, respectively. Figure 2.6b shows the Arrhenius plot for the estimate of the apparent activation energy of desorption, EA, using all the kinetic curves from Fig. 2.6a. The obtained value is ~120 kJ/mol with an excellent coefficient of fit R2 = 0.996. Another example of hydrogen desorption kinetic curves is presented in Fig. 2.7 for the second commercial powder ABCR. There is no desorption at 325°C but full desorption occurs relatively easily at 350°C and higher temperatures (~7.5 wt.%H2 desorbed). In this respect, the ABCR powder behaves in a manner similar to its Tego Magnan counterpart. At the first glance, the desorption rates at the same temperatures in Figs. 2.6 and 2.7 are not very different. Figure 2.7b shows the Arrhenius plot for the estimate of apparent activation of desorption, EA, using all the kinetic curves from Fig. 2.6a. The obtained value is ~168 kJ/mol with an excellent coefficient of fit R2 = 0.996. The activation energy value is much higher than the one obtained for the Tego Magnan powder. In order to shed more light on the noted discrepancy, we analyzed the reaction order parameters η in the standard JMAK equation. Table 2.5 compiles the obtained values of η for both analyzed powders. It is interesting that the highest value of ~3.6 for both powders corresponds to the lowest desorption temperature of 350°C. At higher temperatures, the η parameter substantially decreases to the average value of ~1.5. Apparently, the transformation mechanism is quite different at 350°C as compared to that at the higher temperature range. However, for both powders the values of η are remarkably close to each other in spite of the substantial differences in the apparent activation energy of desorption for Tego Magnan and ABCR powders. Unmilled, commercial Tego Magnan and ABCR powders in the as-received state were also subjected to activation by short cycling. In the first step, the powders were heated to 350°C under a hydrogen pressure of 2.7 MPa to prevent desorption during temperature stabilization. This step took about 15 min. Subsequently, the powders were subjected to activation, which consisted of three cycles and each cycle consisted of the following steps: (a) desorption at 350°C under initial atmospheric pressure of hydrogen (0.1 MPa) for approximately 60 min, (b) annealing under pre-vacuum at 350°C for 15 min, and (c) absorption at 350°C under hydrogen pressure of 2.7 MPa for 30 min. After activation, the powder was desorbed at a constant temperature in a volumetric Sieverts-type apparatus under atmospheric pressure of hydrogen. After each desorption, the same powder sample was re-absorbed at 350°C under a hydrogen pressure of 2.7–3.5 MPa for 15–30 min and this was followed by desorption at a

96

2

Hydrogen desorbed [wt.%]

8.00

Simple Metal and Intermetallic Hydrides

5

7.00 Desorption

4

6.00

3 5.00 2

4.00 3.00 2.00 1.00

1

1-325ⴗC 2-350ⴗC 3-375ⴗC 4-400ⴗC 5-420ⴗC

0.00 0

1000

a

2000

3000

4000

Time [s] −2 y = -167806x + 25.419 R2 = 0.9955

−3

ln k

−4 −5 -6 −7 −8 0.00017

b

0.000175

0.00018

0.000185

0.00019

0.000195

1/RT

Fig. 2.7 (a) Desorption kinetic curves at various temperatures under an initial hydrogen pressure of 0.1 MPa of the as-received, nonactivated commercial MgH2 powder ABCR and (b) the Arrhenius plot of the desorption rate for the estimate of the apparent activation energy, EA, using kinetics data for four temperatures: 350, 375, 400, and 420°C (EA ~168 kJ/mol). Coefficient of fit R2 = 0.996

Table 2.5 The values of the reaction order h in the JMAK equation for the desorption experiments on the unmilled and nonactivated powders presented in Fig. 2.6 and 2.7 Reaction order η Desorption temperature (°C) Tego Magnan® ABCR GmbH & Co. 350 3.60 3.65 375 1.76 1.45 400 1.39 1.73 420 1.56 1.31

2.1

Mg/MgH2

97

desired temperature. The kinetic curves of desorption are shown for Tego Magnan and ABCR powder in Figs. 2.8 and 2.9, respectively. A low rate desorption is already observed at 300°C, and a fair amount of desorption is observed at 325°C for both powders (Figs. 2.8a and 2.9a). This behavior is in a stark contrast to the lack of desorption at 325°C for a nonactivated powder (Fig. 2.7a). As mentioned previously, both nonactivated Tego Magnan and ABCR powders did not desorb hydrogen at temperatures lower than 350°C. It is quite obvious that the activation procedure allows desorption to occur at much lower temperatures. It must, however, be pointed out that after activation cycling, the maximum desorbed capacity of hydrogen, even

Hydrogen desorbed [wt.%]

7 Desorption

6

4

5

3 4

1-300ⴗC

3

2-325ⴗC 3-350ⴗC

2

2

4-375ⴗC 1 1 0 0

a

1000

2000 Time [s]

3000

−4 −4.5

y = -118468x + 16.746 R2 = 0.9964

−5

ln k

−5.5 −6 −6.5 −7 −7.5 −8 −8.5 0.00018 0.00019 0.00019 0.0002

b

0.0002 0.00021 0.00021 0.00022

1/RT

Fig. 2.8 (a) Desorption kinetic curves at various temperatures under an initial hydrogen pressure of 0.1 MPa of the as-received, activated commercial MgH2 powder Tego Magnan and (b) the Arrhenius plot of the desorption rate for the estimate of the apparent activation energy, EA, using kinetics data for four temperatures: 300, 325, 350, and 375°C (EA ~118 kJ/mol). Coefficient of fit R2 = 0.996

2 Hydrogen desorbed [wt.%]

98 7.00 Desorption

6.00 4

5.00

1 3

4.00

1-325ⴗC

2

3.00

ln k

2-350ⴗC

1

2.00

3-375ⴗC

1.00

4-400ⴗC

0.00

0

1000

a

b

Simple Metal and Intermetallic Hydrides

−3 −3.5 −4 −4.5 −5 −5.5 −6 −6.5 −7 −7.5 −8 0.00017

2000 3000 Time [s]

4000

y = -126254x + 18.056 R2 = 0.9749

0.00018

0.00019

0.0002

0.00021

1/RT

Fig. 2.9 (a) Desorption kinetic curves at various temperatures under an initial hydrogen pressure of 0.1 MPa of the as-received, activated commercial MgH2 powder ABCR and (b) the Arrhenius plot of the desorption rate for the estimate of the apparent activation energy, EA, using kinetics data for four temperatures: 325, 350, 375, and 400°C (EA ~126 kJ/mol). Coefficient of fit R2 = 0.975

at a high temperature range of 375–400°C in Figs. 2.8 and 2.9, seems to be lower than that of nonactivated MgH2 in Figs. 2.6 and 2.7. This drop in storage capacity after cycling was reported by Vigeholm et al. [13–15] and is definitely not related to the presence of residual impurity gases in the reacting hydrogen. They calculated that under the assumption that all available ppm impurities in high-purity H2 gas actually reacted with Mg, a maximum passivation of Mg per cycle would be about 0.002%, i.e., completely negligible. Most probably, the drop is related to the kinetics of the absorption stage of cycling due to applied hydrogen pressure being much higher than the equilibrium one as explained in Fig. 2.5, which will be discussed in more detail in connection to the PCT curve in Fig. 2.11. The values of the reaction order η in the JMAK equation for the desorption experiments on the unmilled and activated powders are presented in Table 2.6. On average, the values are slightly larger than those obtained for the nonactivated powders (Table 2.5), which may suggest slightly different transformation mechanisms in the nonactivated and activated powders (Table 2.3). Nonetheless, the η values are quite similar for both Tego Magnan and ABCR powders. As suggested by Dal Toè et al.

2.1

Mg/MgH2

99

Table 2.6 The values of the reaction order η in the JMAK equation for the desorption experiments on the unmilled and activated powders presented in Fig. 2.8 and 2.9 Reaction order η Desorption temperature (°C) Tego Magnan ABCR GmbH & Co. 300 3.52 – 325 2.44 2.63 350 2.04 2.44 375 1.84 1.86 400 – 1.57

[25], since MgH2 powder constitutes an agglomerate of particles with greatly varying sizes, the particles within a certain size range can decompose/compose involving slightly different rate-limiting processes and therefore different values of the reaction order η. Therefore, it seems that the use of free η values obtained from a double-logarithm fitting procedure as described in Sect. 1.4.1.2 is preferable to using fixed η values. Table 2.7 lists the estimated values of the apparent activation energy of desorption. Nonactivated and activated Tego Magnan as well as activated ABCR have almost the same average apparent activation energy of ~118–126 kJ/mol. The estimated average is quite close to that reported by Stander [23] (120 and 126 kJ/mol). Only for the nonactivated ABCR powder, the apparent activation energy of ~168 kJ/ mol is much higher than the others. Huot et al. [24] reported ~156 kJ/mol for the unmilled powder, which was estimated assuming a constant η = 3 in the JMAK equation. It was not specified whether the powder was activated or not. Jensen et al. [22] made an overview of experimentally determined apparent activation energies for desorption (dehydrogenation) of MgH2. They found that for activated MgH2 samples the activation energies of desorption were scattered within the range 120–160 kJ/mol. The samples intentionally exposed to air (called “deactivated” by the authors) exhibited very high values of the reaction rate parameter η, being on the order of 4–7, and extremely high apparent activation energies on the order of 230–300 kJ/mol. They attributed this behavior to the existence of magnesium oxide (MgO) surface layer formed on exposure to air. We also observed activation energies of this magnitude in so-called “aged” MgH2 powders, as will be discussed in Sect. 2.1.5. Friedrichs et al. [26] showed that a thin, amorphous magnesium hydroxide (Mg(OH)2) layer can form on the surface of nanocrystalline MgH2 powder after even a relatively short exposure to air. Apparently, moisture (H2O) in air can easily react with the surface of MgH2 according to the following hydrolysis reaction: MgH2 + 2H2O ⇒ Mg(OH)2 + 2H2

(2.3)

Furthermore, Varin et al. [27] reported that a long-term air exposure of nanocrystalline MgH2 for a few months leads to a massive transformation of a large fraction of MgH2 particles into the crystalline Mg(OH)2 phase. Such a reaction is especially detrimental to a nanocrystalline MgH2 synthesized by reactive ball milling as will be

100

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discussed in more detail in the following sections. Highly hydrolyzed samples of MgH2 may exhibit much higher values of activation energy, being within 230–300 kJ/mol range. Andreasen et al. [28] pointed out that the case of Mg variations in the reported apparent activation energies correlates with the presence of a MgO surface layer inhibiting diffusion of hydrogen. Thus, oxidized samples show large apparent activation energies and well-activated samples show smaller activation energies. However, the mere surface hydrolysis/oxidation layer model cannot explain these peculiar differences in the activation energy of nonactivated/activated Tego Magnan and ABCR powders in Table 2.7. We propose that it can be explained by a combined effect of hydroxide/oxide layer on the surface and, in addition, differences in grain size between both powders: ~67 and 300 nm for Tego Magnan and ABCR, respectively. A model explaining this behavior is shown in Fig. 2.10. It shows particles of Tego Magnan and ABCR powders before and after activation covered with a layer of hydroxide/oxide. It is logical to assume that the impurity layer is always broken/discontinuous along the intersections of grain boundary planes with the particle surface, which provide excellent paths for hydrogen penetration and diffusion into the bulk. Since Tego Magnan has much smaller, nearly nanosized grain, the layer on its particles will be heavily broken into small pieces and practically fully discontinuous owing to a large number of intersections of grain boundary planes with the surface (Fig. 2.10a). Activation, in principle, will not change this picture because the layer cannot be broken into even smaller pieces (Fig. 2.10b). Since hydrogen can easily penetrate and diffuse at many places, the activation energy of desorption remains nearly the same, at a relatively low level before and after thermal activation of Tego Magnan. In contrast, owing to much larger grain size, the ABCR powder has much less grain boundary plane intersections with the particle surface and the layer is semicontinuous (Fig. 2.10c), leading to more difficult penetration of hydrogen and high activation energy of desorption of nonactivated powder. After activation, the layer is heavily broken (Fig. 2.10d) allowing easy hydrogen penetration, which is associated with much lower activation energy of desorption. Figure 2.11a shows an example of PCT desorption curves for the activated commercial MgH2 Tego Magnan powder obtained at four different equilibrium temperatures. Figure 2.11b shows the Van’t Hoff plot of ln p vs. 1,000/T (where p is the pressure and T is the temperature), from which the enthalpy of decomposition of MgH2 of 71kJ/molH2can be obtained. This value is very close to all the values for the enthalpy of formation of MgH2 obtained during absorption as reported Table 2.7 Summary of apparent activation energies of desorption for unmilled Tego Magnan and ABCR powders estimated from the Arrhenius plot Apparent activation energy Coefficient of fit Kinetic curves at of desorption, EA R2 in the Arrhenius temperatures taken Powder (kJ/mol) equation for calculation (°C) Activation Tego Magnan 120 0.996 350, 375, 400, 420 No Tego Magnan 118 0.996 300, 325, 350, 375 Yes ABCR 168 0.996 350, 375, 400, 420 No ABCR 126 0.975 325, 350, 375, 400 Yes

2.1

Mg/MgH2

101

a

b

c

d

Fig. 2.10 (a) Particle of a Tego Magnan powder with the grain size of ~67 nm before activation and (b) after activation. (c) Particle of ABCR powder with the grain size of ~300 nm before activation and (d) after activation (dark envelope around particles represents Mg(OH)2/MgO)

previously in Sect. 2.1.2.1. It can be seen that the maximum hydrogen capacity of activated samples obtained in a PCT test (~5.5wt.%) is about 1.7 wt.% short of the purity-corrected hydrogen capacity of Tego Magnan powder (~7.2 wt.% at 95% purity). The same hydrogen capacity deficit is seen in Fig. 2.8a and 2.9a where the maximum amount of hydrogen desorbed from the activated MgH2, even at the highest desorption temperature, is always smaller than that desorbed from non-activated samples (Fig.2.6a and 2.7a). As mentioned previously, during activation the samples were subjected to absorption at 350° under 2.7 MPa pressure. In addition, intermediate re-absorption steps after each desorption were also conducted at the same temperature and similar pressure. One can notice in Fig. 2.11a that at 350°C the equilibrium (plateau) pressure is about 1MPa for desorption. Assuming a more or less similar plateau pressure for absorption, it means that all the re-absorption steps were conducted at the condition of applied pressure much higher than the plateau pressure (pappl >> ppl). As shown in Fig. 2.5b, under this condition a large amount of β-MgH2 nuclei is being formed, the growth rate is relatively small, and only a thin layer of β-MgH2 can be formed near the surface of each particle until the reaction stops, leading, in effect, to faster kinetics but lower hydriding capacities. Obviously, a smaller amount of absorbed hydrogen results in a smaller amount of desorbed hydrogen, as clearly seen in Fig. 2.11a. However, it must be pointed out that such a hydrogen deficit of practical maximum capacity does not affect the thermodynamic behavior of the powder during desorption, resulting in a very true enthalpy of decomposition/formation.

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Simple Metal and Intermetallic Hydrides

10.00 Desorption Pressure (MPa)

2.77 MPa 1.92 MPa 1.11 MPa

1.00 0.565 MPa

400C 375C 350C 325C 0.10

0

1

2

a

3 4 5 6 H2 capacity (wt.%)

7

8

4 DHdes = 71 kJ/molH2

3.5 3

ln p

2.5 2 y = -8.5824x + 16.135 R2 = 0.9909

1.5 1 0.5 0 1.45

b

1.5

1.55 1.6 1000/T (1/K)

1.65

1.7

Fig. 2.11 (a) PCT desorption curves at various temperatures for the activated commercial MgH2 Tego Magnan powder; numbers indicate the average mid-plateau pressure. (b) The Van’t Hoff plot for finding the enthalpy and entropy of decomposition, which is equal to ~71 kJ/mol and~134 J/ mol K, respectively. Note excellent coefficient of fit R2 = 0.991 (p – pressure)

Figure 2.12 shows an example of a curve obtained in a differential scanning calorimeter (DSC), which illustrates general features of hydrogen desorption from a nonactivated MgH2 sample. A single and quite symmetrical endothermic peak characterizes hydrogen desorption. The shape of the peak for nonactivated MgH2 is quite narrow but very strong. The peak maximum is centered at 428.3°C (at heating rate 4°C/min), which is the most common range of the peak maximum temperatures for nonactivated MgH2 powders. Usually, most of the peak temperatures fall within the range 410–440°C (depending on the DSC heating rate).

2.1.3

Hydrogen Storage Characteristics of Mechanically (Ball) Milled MgH2

As discussed in Sect. 1.3, the major objective of mechanical (ball) milling is to obtain a substantial nanostructuring of hydrides (nanohydrides), which in effect can

2.1

103

Mg/MgH2 DSC / mW/mg 14

+

exo

[1] 428.3⬚C

12 10 8 6 4 2 0 200

250

300

350 400 Temperature / ⬚C

450

500

Fig. 2.12 DSC curve of a commercial, as-received and nonactivated Tego Magnan powder obtained under argon flow and a heating rate 4°C/min

improve their hydrogen storage properties. The first pioneering experiments on the synthesis of nanohydrides by ball milling were conducted in the early and mid1990s by the group led by Prof. J.O. Ström-Olsen at the McGill University in Montréal, Canada, in which a prominent role was played by L. Zaluski and A. Zaluska [29–33]. A parallel line of work on the same subject was initiated by the group led by O.N. Srivastava in India [34]. However, it must be pointed out that in technical terms the hydrides in early research were not physically ball-milled in order to induce nanocrystallinity. In reality, either intermetallic compounds were first synthesized by mechanical alloying, or pure metals (e.g., Mg) were simply ballmilled under argon and subsequently subjected to hydrogenation at the appropriate temperature and hydrogen pressure. Such a method is called in the present text “a two-step” method. In addition, those early works, especially in the group led by Prof. J.O. Ström-Olsen, overemphasized the effect of nanograins (crystallites) formed within the heavily milled powder particles on the hydrogen absorption/desorption properties and somehow marginalized the role of the reduction of particle size that occurs simultaneously with the decrease of nanograin (crystallite) size. Especially, for milled Mg, as well as for the Mg-based and other intermetallic compounds having nanograin microstructure, subsequent hydrogenation/dehydrogenation cycles at elevated temperatures should give rise to nanograin growth. Hence, it is hard to understand why improved hydrogen storage properties would, indeed, still require the nanosized grains which in all practical terms do not exist any longer. This very important issue will be discussed in more detail in the following sections.

2.1.3.1 Microstructural Evolution During Milling and Subsequent Cycling of Commercial MgH2 Powders In the context of this book, the name “commercial” MgH2 will be used to describe any MgH2 powder purchased from a supplier of chemicals, the prime examples of

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Simple Metal and Intermetallic Hydrides

which are Tego Magnan from Degussa-Goldschmidt and ABCR from ABCR GmbH&Co.KG, which were already discussed in Sect. 2.1.1. The first attempts to ball-mill a commercial undoped MgH2 powder were undertaken by the group in the Hydro-Quebec Research Institute in Quebec, Canada [24, 35–36]. In contrast to the prior work in the Ström-Olsen’s group, the Hydro-Quebec authors at least pointed out that not only nanograins are being formed but also particle size reduction associated with an increase of the specific surface area (SSA) occurs simultaneously during ball milling. Nevertheless, no attempts to correlate the structural changes occurring during ball milling with hydrogen sorption properties were made. In our laboratories, we have carried out a thorough investigation of microstructural evolution occurring during ball milling of both commercial Tego Magnan and ABCR MgH2 powders. Figure 2.13 shows the evolution of XRD patterns as a function of controlled mechanical milling (CMM) of the commercial ABCR MgH2 powder under the HES57 mode in a hydrogen gas atmosphere in the magneto-mill Uni-Ball-Mill 5 (Sect. 1.3.2). It is clearly seen that, indeed, ball milling induces drastic changes in the microstructure of the milled powders. There are three important microstructural changes occurring upon milling of MgH2 that have already been described to a smaller or larger extent in a number of papers mentioned above. First, the breadth of the XRD peaks of β-MgH2 increases with increasing milling time. As discussed in Sect. 1.4.3, this is related to the formation of crystallites (nanograins) within the powder particles, which may be accompanied by the introduction of lattice strains. Table 2.8 lists nanograin size and lattice strain of β-MgH2 as a function of milling time estimated from the procedure described in Sect. 1.4.3. 30000

β-MgH2 γ-MgH2 Mg MgO s substrate

Counts

20000

s

s

s

As received 0.25 h

10000

1h 5h 10 h 20 h

0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 2.13 Evolution of XRD patterns as a function of ball milling time of ABCR powder under the HES57 mode (high energy shearing; two magnets at five and seven o’clock positions) in hydrogen atmosphere under ~600 kPa pressure in the magneto-mill Uni-Ball-Mill 5

2.1

Mg/MgH2

105

Table 2.8 Grain size variations of b-MgH2 as a function of milling time of ABCR powder under HES57 mode from Fig. 2.13 Number of XRD Milling time (h) Grain size (nm) Strain R2 peaks 0 (as received 299 0 0.9989 4 0 (as received) 303 0 0.9936 6 0.9904 4 0.25 (15 min) 52 9.31 × 10−4 0.9954 3 1 30 3.16 × 10−3 0.9901 4 5 12 1.17 × 10−3 10 10 0 0.9999 3 0.9926 3 20 11 5.26 × 10−3

1000

Grain size (nm)

ABCR ABCR cycled after milling Tego Magnan

100

Average grain size after cycling

10

1 0

5

10 15 Milling time (h)

20

25

Fig. 2.14 Grain size of both commercial powders Tego Magnan and ABCR as a function of milling time in the magneto-mill Uni-Ball-Mill 5 under shearing and impact modes. For comparison, the level of the grain size measured after activation by a short cycling for the powders milled for the same duration is also shown (cycling scheme: heating to 325°C for ~15 min under 3.4 MPa H2 to prevent desorption-first desorption at 325°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 3.4 MPa H2 for 30 min/second desorption at 350°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 2.7 MPa H2 for 30 min)

As can be seen, the nanograin size is being reduced pretty fast with milling time such that barely after 15 min of milling it is already reduced approximately sixfold as compared to the original grain size of the as-received ABCR powder. After milling for approximately 5 h, the nanograin size is already saturated within the ~10–12 nm range and further milling for 10–20 h does not bring about any more changes in the nanograin size. It must be pointed out that the time to reach the saturation level of the β-MgH2 nanograin size depends on the milling mode, such as low-energy shearing (LES), high-energy shearing (HES), and impact, however when the milling process is carried out in the magneto-mill Uni-Ball-Mill 5 the saturation time is more or less within a few hours range regardless of the mode. The situation could be slightly different when using different type of mill, e.g., Spex or Fritsch, but there is no systematic study of a saturation behavior in these types of mills. Figure 2.14 shows

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Simple Metal and Intermetallic Hydrides

a plot of grain size as a function of milling time for two commercial powders Tego Magnan and ABCR, which were ball-milled under various shearing and impact modes in the magneto-mill Uni-Ball-Mill 5. It is apparent that within the experimental scatter the grain size saturation level does not depend on the type of commercial MgH2 powder and the mode of milling. The strains of the β-MgH2 phase in Table 2.8 are minimal (<0.5%) as opposed to the values of over 1% reported by Huot et al. [24]. Taking into account that the β-MgH2 phase is tetragonal and its atomic bonding is mostly ionic with a little covalency [2, 4], i.e., ceramic-like, one can hardly envision substantial dislocation activities that would lead to dislocation accumulation in the lattice and, in turn, strains of such a large magnitude (>1%). It seems that the strains reported by Huot et al. are rather overestimated. For clarification, it is to be pointed out that the XRD peaks of MgO observed in Fig. 2.13 arise because of the exposure of residual Mg present in the as-received MgH2 powder to air during powder handling for XRD tests. After milling, the residual Mg becomes nanostructured and in this state exhibits a very strong affinity to oxygen in air. Second, even after a short milling time of 15 min (0.25 h), the XRD peaks of an orthorhombic γ-MgH2 phase already appear on the pattern (Fig. 2.13). As pointed out by Schulz et al. [35, 36], the pressure increase due to the mechanical action of the milling balls produces the structural transformation of a tetragonal β-MgH2 into an orthorhombic γ-MgH2 which normally would occur under enormous static pressures of around 8 GPa (Sect. 2.1.1). Interestingly, the γ-MgH2 XRD intensities do not seem to increase measurably with increasing milling time beyond 1 h, which suggests that their volume fraction may saturate after a relatively short milling time under this specific milling mode. Huot et al. [24] estimated using the Rietveld refinement of an XRD pattern that the γ-MgH2 phase abundance was around 18% and did not increase with increasing milling time. They explained this apparent dilemma by two possible mechanisms. The first mechanism relies upon the occurrence of two competitive processes, one promoting the formation of the γ−phase due to mechanically driven transformation, and the other enhancing the formation of the β−phase due to thermally driven γ− β transformation. The alternative mechanism is based on the assumption that the increase of strain in the γ−phase beyond a certain level induces transformation of γ into β. It is, however, hard to decide unambiguously which mechanism is indeed true. However, as will be discussed later, our results show that the volume fraction of the γ−phase seems to be higher after milling for 100 h than that after 20 h. Third, parallel to the decrease of grain size there is always a decrease of the particle size of milled powders, as shown in Fig. 2.15. The particle size reduction occurs within a very similar time frame as does the reduction of grain size. As can be seen in Fig. 2.15a, barely after about 15 min the particle size is reduced from the initial ~40 to ~1 μm. Further prolonged milling up to 100 h may bring about very incremental particle size reduction down to ~0.6 μm (Fig. 2.15b). This behavior is of a very general nature and practically does not depend on the mode of milling and the type of MgH2 commercial powder as can be seen in Fig. 2.15. Almost identical time frame for grain and particle size variations as a function of milling time, as

2.1

Mg/MgH2

107

Particle size ECD (mm)

100 ABCR IMP68

ABCR HES57

Tego HES57

Tego IMP68

10

1

0.1 0

20

a

40

60

80

100

Milling time (h)

Particle size ECD (mm)

100 ABCR IMP68

ABCR HES57

Tego HES57

Tego IMP68

10

1

0.1 0

b

1

2

3

4

5

Milling time (h)

Fig. 2.15 Powder particle size vs. milling time for two commercial MgH2 powders Tego Magnan (Tego) and ABCR which were milled in the magneto-mill Uni-Ball-Mill 5 under shearing and impact modes. (a) Milling for up to 100 h and (b) enlarged section up to 5 h (HES57 – high-energy shearing with two magnets at five and seven o’clock positions; IMP68 – strong impact with two magnets at six and eight o’clock positions)

shown in Figs. 2.14 and 2.15, makes it rather difficult to identify unambiguously which factor is, indeed, governing hydrogen storage characteristics. This difficulty has led to a common belief that the grain size is mostly responsible for the observed enhancement of hydrogen storage properties, which is not necessarily the case as will be discussed later. Also, it must be pointed out that even after milling up to 100 h, which was the longest milling duration in the Uni-Ball-Mill 5, and regardless of the mode of milling, we never observed a contamination of MgH2 with Fe coming from the steel milling tools. The presence of Fe was not detected either by energy dispersive spectroscopy (EDS) or XRD measurements. Apparently, the content of Fe, if any, was below the detectability level of both analytical techniques. Recently, Ares et al. [37] reported that the concentration of Fe in the MgH2 milled in a steel vial increased to 4 wt.% after 700 h of milling in a Fritsch P5 planetary mill. They found

108

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Simple Metal and Intermetallic Hydrides

that such a high concentration of Fe facilitated particle agglomeration and coldwelding. It is also interesting to note that Ares et al. found that the lattice strains of MgH2 milled in a steel vial were much lower than those observed in the material milled in a ceramic vial. In our work, a stainless steel vial was used, resulting in negligible lattice strains of MgH2 (Table 2.8). Thermal activation by desorption/absorption cycling at elevated temperatures of premilled MgH2 reverts some of the microstructural changes induced by milling. ABCR powders initially milled, as shown in Fig. 2.13, were subjected to the following cycling: heating to 325°C for ~15 min under 3.4 MPa H2 to prevent desorption/first desorption at 325°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 3.4 MPa H2 for 30 min/second desorption at 350°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 2.7 MPa H2 for 30 min. Figure 2.16 shows the XRD patterns of the premilled powders after the last absorption stage in a cycling scheme. By comparison with the XRD patterns of the milled powders in Fig. 2.13, noticeable changes occurred after cycling. First, the diffraction peaks of γ-MgH2 completely disappeared. Second, the diffraction peaks of the β-MgH2 phase have become much narrower. Table 2.9 shows the estimated grain sizes and lattice strains of β-MgH2. There has been substantial grain 40000

β-MgH2 Mg MgO s substrate

Counts

30000

20000

s

0.25 h

s

1h

10000

5h 10 h 20 h

0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 2.16 XRD patterns of initially milled ABCR powders from Fig. 2.13 after short cycling: heating to 325°C for ~15 min under 3.4 MPa H2 to prevent desorption/first desorption at 325°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 3.4 MPa H2 for 30 min/second desorption at 350°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 2.7 MPa H2 for 30 min

2.1

Mg/MgH2

109

Table 2.9 Grain size variations of premilled ABCR powder after short cycling Number of Milling Grain XRD peaks time (h) size (nm) Lattice strain R2 0 166 1.09 × 10-3 0.9865 3 0.25 47 0 0.8441 3 1 34 0 0.9171 3 5 40 0 0.9994 3 10 35 0 0.9992 3 20 43 0 0.9907 3

Particle size (ECD) (mm)

2.5 2 1.5 1 0.5 0 0

1

2

3

4

5

6

Number of cycles

Fig. 2.17 Particle size (ECD) of Tego Magnan powder milled for 100 h under the low-energy impact mode with one magnet and subsequently subjected to five cycles of desorption at 300°C under 0.1 MPa H2 pressure for ~4,700 s/absorption at 300°C under 2.5 MPa H2 for 60 min

growth, although the β-MgH2 grains can be still classified as nanosized (<50 nm). The lattice strains are almost completely eliminated. The grain sizes of premilled powders after cycling are plotted in Fig. 2.14. It is interesting to note that regardless of the initial milling time and different initial grain sizes after milling (Table 2.8), the average grain size after cycling is almost identical, being within a narrow range of ~40 nm. It should also be pointed out that the average grain size of as-received and cycled powder in Table 2.9 seems to be smaller (~166 nm) than that before cycling (~300 nm) (Table 2.8). This grain size reduction might be induced by the cyclic lattice expansion upon the hydrogenation/dehydrogenation process [11–16]. In contrast to quite a substantial nanograin growth of premilled β-MgH2 upon cycling, the particle size does not undergo any measurable changes upon cycling as shown in Fig. 2.17 for the Tego Magnan powder milled for 100 h under low-energy impact mode with one magnet and subsequently subjected to the following cycling: desorption at 300°C under 0.1 MPa H2 pressure for ~4,700 s/absorption at 300°C under 2.5 MPa H2 for 60 min. Within the experimental scatter, the particle size (ECD) after five cycles is almost the same as the one directly after milling. On the other hand, Fig. 2.18 clearly shows that after the first cycle the grain size rapidly increases from about as-milled 10 nm to about 50 nm and then saturates for the next

110

2

Simple Metal and Intermetallic Hydrides

70

Grain size (nm)

60 50 40 30 20 10 0 0

1

2

3

4

5

6

Number of cycles

Fig. 2.18 Grain size of Tego Magnan powder from Fig. 2.17 as a function of number of cycles (1 cycle: desorption at 300°C under 0.1 MPa H2 pressure for ~4,700 s/absorption at 300°C under 2.5 MPa H2 for 60 min)

a

b

Fig. 2.19 Scanning electron micrographs of Tego Magnan powder (a) milled for 100 h under the high-energy impact mode (IMP68 with two magnets at six and eight o’clock positions) and (b) after three cycles of desorption at 350°C under 0.1 MPa H2 pressure for 15 min/annealing at 350°C under pre-vacuum for 15 min/absorption at 350°C under 3.2 MPa H2 pressure for 15 min

four cycles. There is no lattice strain after the first cycle but after the third and fifth cycle it is 4.8×10−5 and 8×10−5, respectively. Apparently, a minimal strain appears during cycling. Another important behavior of powders upon cycling is a complete lack of sintering. Fig. 2.19 shows the scanning electron microscopy (SEM) micrographs of the Tego Magnan powder milled for 100 h under IMP68 high-energy mode (with two magnets at six and eight o’clock positions) and subsequently subjected to three cycles: desorption at 350°C under 0.1 MPa H2 pressure for 15 min/annealing at 350°C under pre-vacuum for 15 min/absorption at 350°C under 3.2 MPa H2 pressure for 15 min. Despite the relatively high cycling temperature, no evidence of sintering can be observed. Additional evidence that no sintering occurred upon cycling is provided

2.1

Mg/MgH2

111 45 Mean: 0.6 mm Std. Dev.: 0.4 mm

Relative frequency (%)

40 35 30 25 20 15 10 5 0 0.0

0.5

a

1.0 1.5 2.0 Particle size ECD (mm)

2.5

3.0

45 Mean: 0.6 mm Std. Dev.: 0.4 mm

Relative frequency (%)

40 35 30 25 20 15 10 5 0 0.0

b

0.5

1.0 1.5 2.0 Particle size ECD (mm)

2.5

3.0

Fig. 2.20 Particle size distribution of Tego Magnan powder (a) milled for 100 h under the highenergy impact mode (IMP68 with two magnets at six and eight o’clock positions) and (b) after three cycles of desorption at 350°C under 0.1 MPa H2 pressure for 15 min/annealing at 350°C under pre-vacuum for 15 min/absorption at 350°C under 3.2 MPa H2 pressure for 15 min

in Fig. 2.20, which shows the comparison of particle size (ECD) distributions of the powder from Fig. 2.19 after milling and after cycling, and in Fig. 2.21, which shows a plot of particle size ranges of the as-received, as-milled, and cycled powders. It is obvious that both the particle size distribution and the particle mean value (ECD) are not affected by cycling and remain essentially at the same level as those directly after milling. The characteristic lognormal distribution of particles is also retained after cycling. It can also be seen in Figs. 2.14, 2.18, and 2.21 that after cycling the grain size of β-MgH2 is ~40–50 nm, which is fourfold larger than that after milling. Our

112

2

Simple Metal and Intermetallic Hydrides

100

Particle size ECD (mm)

67 nm

10

As received

13 nm

54 nm

1

After milling

After cycling

0.1 Powder

Fig. 2.21 Particle size of Tego Magnan powder in the as-received state, after milling for 100 h under the high-energy impact mode (IMP68 with two magnets at six and eight o’clock positions) and after three cycles desorption at 350°C under 0.1 MPa H2 pressure for 15 min; annealing at 350°C under pre-vacuum for 15 min/absorption at 350°C under 3.2 MPa H2 pressure for 15 min. Grain size of β-MgH2 is shown above each data point

results are in good agreement with those of Walton et al. [38], who reported that after two cycles at 300°C the nanograin size of MgH2 powder premilled for 5 h grew from the initial ~25 nm to ~70 nm. Huhn et al. [39] reported that the Mg grains grew from an initial size of 11 nm to 200 nm during desorption of the MgH2 sample between 300 and 400°C. However, the average grain sizes after cycling obtained in our work and those reported in [38, 39] are generally larger than those reported by Fátay et al. [40].

2.1.3.2

Hydrogen Absorption of Ball-Milled Commercial MgH2 Powders

A profound effect of ball milling on the absorption kinetics of ABCR powder is indeed clearly seen in Fig. 2.22 for the milled, activated, and cycled sample. The as-received powder was milled for 5 h under 600 kPa of hydrogen gas using strong shearing mode with two magnets (designated HES57). Subsequently, the powder was heated up to 350°C in a volumetric Sieverts-type apparatus (Sect. 1.4.2) under an applied hydrogen pressure of 3.5 MPa (to prevent desorption), held for about 30–45 min to stabilize temperature, and then subjected to nine desorption/absorption cycles consisting of desorption at 350°C; annealing at 350°C/pre-vacuum/15 min; cooling under pre-vacuum to the required absorption temperature and finally absorption at T = 325–200°C (at every 25°C)/1.2 MPa/4,750 s. The procedure is exactly the same as the one described in Sect. 2.1.2.1 for absorption of the same unmilled powder. Consequently, the powder that absorbed at 200°C went through nine full cycles. The absorption curves in Fig. 2.22a are very different than those for the unmilled powder in Fig. 2.4a. First, the milled ABCR powder is able to

2.1

Mg/MgH2

113

7.00 1-200ⴗC 2-225ⴗC 3-250ⴗC 4-275ⴗC Absorbed H2 (wt.%)

6.00

5-300ⴗC 6-325ⴗC

6

5.00

5 4

4.00

Absorption 3

3.00 2

2.00

1

1.00 0.00 0

1000

a

2000 Time (s)

3000

4000

−3 y = -71301x + 10.276 R2 = 0.9685

−4

ln k

−5 −6 −7 −8 −9 0.00018

b

0.0002

0.00022

0.00024

0.00026

1/RT

Fig. 2.22 (a) Absorption kinetic curves of milled (HES57;5 h), activated, and cycled ABCR powder and (b) estimate of apparent activation energy of absorption from the Arrhenius plot of ln k vs. 1/RT using data for all six temperatures: 200, 225, 250, 275, 300, and 325°C (EA ~71 kJ/mol). Coefficient of fit R2 = 0.969 (activation and cycling: heating 350°C/3.5 MPa; 30–45 min/nine desorption/absorption cycles: desorption 350°C/annealing at 350°C/pre-vacuum/15 min/cooling under pre-vacuum to the required absorption temperature and finally absorption at T = 325–200°C (at every 25°C)/1.2 MPa/4,750 s)

absorb ~3 wt.%H2 at a low temperature of 200°C after about 4,000 s, which is in a great contrast to the unmilled powder, which did not absorb at all at such a low temperature. Second, at 250°C after about 1,500 s the milled powder absorbs 5 wt.%H2, which strongly contrasts with the barely 0.5 wt.%H2 absorbed by the unmilled powder under the same absorption conditions. Third, the highest absorbed capacity of the milled powder after 4000 s is about 5.5 wt.%H2 in the temperature range 250–325°C. Again this strongly contrasts with 4.5 wt.%H2 which the unmilled powder is able to absorb at the highest temperature of 325°C. Indeed, as can be seen, the absorption kinetics of MgH2 are greatly improved by ball milling.

114

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Simple Metal and Intermetallic Hydrides

Nevertheless, the maximum hydrogen capacity achieved after absorption at 325°C is still about 2 wt.% lower than the purity-corrected capacity of about 7.5 wt.% for the ABCR powder. The absorption curves in Fig. 2.22a were analyzed by a linear fitting to the JMAK equation from which the reaction rate constant, k, and the reaction order η can be determined. The values of the reaction order η are listed in Table 2.10. Quite remarkably, regardless of absorption temperature, the η parameter is always very close to 1. This value is similar to the η values calculated for the same but unmilled powder (Table 2.4) except that for the unmilled powder at the lowest desorption temperature of 250°C the η parameter was equal to ~2. According to Table 2.3, the η value close to 1 at the entire range of absorption temperatures in Table 2.10 suggests either diffusion rate–limited transformation occurring by twodimensional growth at constant nuclei number, or, alternatively, interface-controlled transformation with one-dimensional growth at constant nuclei number. More importantly, the η value, which seems to be independent of absorption temperature within the 200–325°C range (Fig. 2.22a), may indicate a very uniform mechanism of absorption after milling and activation cycling. The apparent activation energy of absorption in Fig. 2.22a was estimated through the Arrhenius plot of the rate constant, k, with temperature (Fig. 2.22b). Within the experimental error, the obtained value of ~71 kJ/mol is close to that obtained for the unmilled ABCR powder in Fig. 2.4b (~65 kJ/mol). In view of apparently faster desorption kinetics of the milled ABCR powder in Fig. 2.22a as compared to the unmilled powder in Fig. 2.4a, one might expect a lower value of the apparent activation energy of the former. This brings us to the point when a question should be asked as to what microstructural features are responsible for the better absorption kinetics of the milled powder. Since the powder after milling was subjected to high-temperature cycling for relatively long time durations, it is obvious that the γ-MgH2 phase already disappeared because of the cycling. In addition, the original nanograin size, which was within the 10 nm range after 5 h of milling (Fig. 2.14), should have grown to the size of ~40–50 nm characterizing cycled materials as discussed in Sect. 2.1.3.1. Therefore, it is hard to believe that a nanometric grain size is responsible for the improved absorption kinetics. In contrast, the mean particle size (ECD) after 5 h of milling under the HES57 mode could be slightly smaller than 1 μm (Fig. 2.15b). Since it remains unchanged after cycling (Figs. 2.17 and 2.21), it seems apparent that the reduced particle size is responsible for improvements in the absorption kinetics after milling and cycling. Similar to our results, Walton et al. [39] reported that hydroTable 2.10 The values of the reaction order η in the JMAK equation for the absorption experiments on the milled, activated, and cycled ABCR powder presented in Fig. 2.22 Absorption temperature (°C) Reaction order η 200 1.20 225 0.99 250 1.07 275 1.01 300 0.90 325 0.97

2.1

Mg/MgH2

115

gen cycling of milled MgH2 resulted in no change in the particle size but produced a marked increase in the grain size without any deterioration in the hydrogen absorption kinetics. They concluded that the changes in grain size observed on hydrogen cycling did not play a significant role in the hydrogen absorption kinetics. In view of all this evidence, it seems prudent to conclude that the major factor improving absorption properties of milled MgH2 is the reduced particle size rather than a nanometric grain size. More evidence will be provided in the following sections. 2.1.3.3

Hydrogen Desorption of Ball-milled Commercial MgH2 Powders

Figures 2.23 and 2.24 show the desorption curves and the Arrhenius plots for the calculation of the apparent activation energy of the Tego Magnan powder milled for 20 and 100 h, respectively. In both cases desorption tests were made under 0.1 MPa

Hydrogen desorbed [wt.%]

8 Desorption 7 5 6 4 5

1-275ⴗC 2-300ⴗC 3-325ⴗC 4-350ⴗC 5-375ⴗC

3

4 3

2

2 1 0

1 0

1000

2000

3000

4000

Time [s]

a −4 −4.5 −5 −5.5

ln k

−6 −6.5 −7 −7.5

y = −139834x + 21.653 R2 = 0.9877

−8 −8.5 0.00017

b

0.00018

0.00019

0.0002

0.00021

0.00022

1/RT

Fig. 2.23 (a) Desorption kinetic curves at various temperatures under initial hydrogen pressure of 0.1 MPa of the commercial Tego Magnan MgH2 powder milled continuously for 20 h under the IMP68 mode in argon and (b) the Arrhenius plot of the desorption rate for the estimation of the apparent activation energy, EA, using kinetics data for four temperatures: 300, 325, 350, and 375°C (EA ~140 kJ/mol). Coefficient of fit R2= 0.988. Desorption tests were carried out directly after milling

116

2

Simple Metal and Intermetallic Hydrides

Hydrogen desorbed [wt.%]

8 5

7

i 4

6 5

3

4

2

3 2 1

1

1-275ⴗC 2-300ⴗC 3-325ⴗC 4-350ⴗC 5-375ⴗC

0 0

1000

a

2000 Time [s]

3000

4000

−2 −3

y = -146449x + 23.08 R2 = 0.9743

ln k

−4 −5 −6 −7 −8 −9 0.00018

b

0.00019

0.0002

0.00021

0.00022

1/RT

Fig. 2.24 (a) Desorption kinetic curves at various temperatures under an initial hydrogen pressure of 0.1 MPa of the commercial Tego Magnan MgH2 powder milled continuously for 100 h under the IMP68 mode (in argon) and (b) the Arrhenius plot of the desorption rate for the estimation of the apparent activation energy, EA, using kinetics data for four temperatures: 300, 325, 350, and 375°C (EA ~146 kJ/mol). Coefficient of fit R2= 0.974. Desorption tests were carried out directly after milling

(atmospheric) hydrogen pressure, directly after milling without any activation. Tables 2.11 and 2.12 show the reaction orders η in the JMAK equation. In the first place, it is quite revealing to compare desorption kinetics after milling in Figs. 2.23 and 2.24 with the kinetics of the same Tego Magnan powder in the as-received state shown in Fig. 2.6. Unmilled powder did not desorb below 350°C. After 20 h of milling, ~4 wt.% H2 is already desorbed at 300°C in 4,000 s. It is interesting to note that milling for 100 h makes desorption rate even faster, since at 300°C 4 wt.% H2 is desorbed in ~2,500 s and 5 wt.% in ~3,200 s. This advantage of a shorter desorption time for the 100 h milled powder is preserved at each temperature in Fig. 2.24a. The grain size of the powders milled for 20 and 100 h is ~13 nm (R2 = 0.997; zero lattice strain) and ~14 nm (R2 = 0.996; zero lattice strain), respectively. Their average ECD particle size is essentially identical and equal to

2.1

117

Mg/MgH2 Table 2.11 The values of the reaction order η in the JMAK equation for the desorption experiments on the Tego Magnan powder milled for 20 h presented in Fig. 2.21 Absorption Reaction order η temperature (°C) 300 3.18 325 1.69 350 1.42 375 1.47

Table 2.12 The values of the reaction order h in the JMAK equation for the desorption experiments on the Tego Magnan powder milled for 100 h presented in Fig. 2.22 Absorption Reaction temperature (°C) order η 300 2.53 325 2.02 350 1.91 375 1.60

Fig. 2.25 DSC traces of Tego Magnan powder milled continuously for 20 (continuous line) and 100 h (broken line) under the IMP68 mode in argon (heating rate 4°C/min)

0.6 ± 0.4 μm. The question is then which factor is responsible for the faster desorption rate of the 100 h milled powder? In our opinion this factor is the higher volume fraction of the γ-MgH2 phase in the 100 h milled powder. Figure 2.25 shows a peculiar behavior of the DSC trace of the 20 and 100 h milled powders. The lowtemperature shoulder (at 363.7°C) in the DSC curve of the 20 h milled powder becomes a principal low-temperature peak (365.1°C) in the DSC curve of the 100 h milled powder. According to Gennari et al. [41], the low-temperature shoulder/

118

2

Simple Metal and Intermetallic Hydrides

peak is due to the total decomposition of γ-MgH2 and the partial decomposition of a fraction of β-MgH2, whereas the high-temperature DSC peak corresponds to the decomposition of the remaining β-MgH2. This strongly suggests that the amount of γ-MgH2 is much greater in the 100 h milled powder than that in the 20 h milled powder, which makes desorption kinetics more rapid in the former. The apparent activation energy of desorption in Fig. 2.23a was estimated through the Arrhenius plot of rate constant, k, with temperature (Fig. 2.23b). Using all five data points, we obtained 140 and 146 kJ/mol for powders milled for 20 and 100 h, respectively. Within the experimental error both values are the same. As mentioned previously for the absorption process, the apparent activation energy does not seem to be sensitive enough to reveal more subtle differences in desorption kinetics. At this point it must be stressed that deleting the highest η values in Tables 2.11 and 2.12 gives the activation energy 120 (R2 = 0.999) and 117 kJ/mol (R2 = 0.986) for the 20 and 100 h milled powder, respectively. This value is much smaller than those obtained taking into account 300°C in Figs. 2.23 and 2.24 but still almost identical for both milling times. Also, these results show that the selection of too narrow a range of temperatures for calculating the activation energy of desorption can substantially affect the calculated values. Finally, short cycling after milling does not change the activation energy. The Tego Magnan powder milled continuously for 20 h under IMP68 mode from Fig. 2.23 showed an activation energy EA = 138 kJ/mol (R2 = 0.987) after three cycles of desorption at 350°C/0.1 MPaH2 and absorption at 350°C/2.6 MPaH2. A careful analysis of desorption curves in Figs. 2.23a and 2.24a shows one commonly observed effect of milling on the practical maximum hydrogen capacity which is reduced after desorption. The maximum purity-corrected hydrogen capacity of the Tego Magnan MgH2 powder is around 7.2 wt.% at 95% purity. This capacity is in practical terms obtained by hydrogen desorption from the unmilled commercial Tego Magnan powder in the temperature range 350–420°C as evidenced in Fig. 2.6a. At the lowest desorption temperature of 350°C in Fig. 2.6a, the total purity-corrected theoretical capacity of the as-received Tego Magnan MgH2 powder can be easily obtained after about 3,600s. However, after milling for 20 and 100 h the maximum hydrogen capacity obtainable after desorption at 350°C is less than 7 wt%. Apparently, the maximum purity-corrected hydrogen capacity is impossible to be achieved for a milled Tego Magnan MgH2 powder even after 4,000 s of desorption as evidenced by Figs. 2.23a and 2.24a. In order to shed more light on this intriguing behavior, the Tego Magnan MgH2 powder milled for 100 h in argon was directly after milling desorbed in a Sievertstype apparatus at three temperatures: 350, 375, and 400°C. Desorption curves are shown in Fig. 2.26. It is clear that even at 375–400°C the milled Tego Magnan MgH2 powder does not completely desorb, showing ~6.2–7 wt.% desorbed hydrogen as compared to its 7.2 wt.% capacity. The microstructure of desorbed powders was investigated by XRD as shown in Fig. 2.27. After desorption at all three temperatures, there are high-intensity peaks of newly formed Mg and small but sharp peaks of retained MgH2 discernible on the XRD pattern. Apparently, for whatever reason, the decomposition process of ball-milled MgH2 hydride does not proceed to completion

2.1

Mg/MgH2

119

Hydrogen desorbed [wt.%]

8 2

7

3

6

1

5 4

1-350ⴗC

3

2-375ⴗC

2 3-400ⴗC

1 0

0

1000

2000 Time [s]

3000

4000

Fig. 2.26 Hydrogen desorption curves at three different temperatures of Tego Magnan powder milled continuously for 100 h under the IMP68 mode in argon

30000

retained b-MgH2

20000 Counts

Mg

400ⴗC 10000 375ⴗC

350ⴗC

0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 2.27 XRD patterns of MgH2 (Tego Magnan) powders milled continuously for 100 h under the IMP68 mode in argon and subsequently desorbed directly after milling in a Sieverts-type apparatus at various temperatures

even at a temperature as high as 400°C. Table 2.13 shows the grain size and lattice strains of β-MgH2 and Mg after milling and after desorption, which were estimated from the broadening of appropriate XRD peaks in Fig. 2.27. The ~14 nm nanograin size of β-MgH2 obtained after milling for 100 h (no XRD pattern shown here) increased to ~60–80 nm range for the retained β-MgH2 present after desorption.

120

2

Simple Metal and Intermetallic Hydrides

Table 2.13 Grain size of b-MgH2, Mg, and retained β-MgH2 in the Tego Magnan powder after milling and desorption at various temperatures Grain Number Powder size(nm) Strain (%) R2 of peaks 14 0 0.9964 7 MgH2 – milled 0.9838 3 350°C (Mg) 86 8.2 × 10−4 55 0 0.9947 4 350°C (MgH2) 0.9997 3 375°C (Mg) 99 6.6 × 10−4 80 0 0.9973 5 375°C (MgH2) 0.9948 6 400°C (Mg) 78 5.4 × 10−4 62 0 0.9948 6 400°C (MgH2)

The freshly formed Mg from the decomposed β-MgH2 has the same range of grain size. The lattice strains in β-MgH2 are nil after both milling and desorption. Some minimal strains are observed in the freshly formed Mg after desorption and they decrease with increasing desorption temperature. An exact explanation as to why milled nanocrystalline MgH2 becomes stabilized during desorption is not available at the present time. Nevertheless, a few relevant factors should be considered. First, as discussed previously, inevitably γ-MgH2 is always formed during milling because of the transformation of β-MgH2 (Fig. 2.13). However, during subsequent high temperature desorption or cycling this orthorhombic hydride phase quickly disappears (Fig. 2.16). According to Gennari et al. [41], the initial decomposition of the γ−phase produces synergic effects during hydrogen desorption, which stimulate β-MgH2 decomposition by creating a volume contraction, which, in turn, generates stresses acting on β-MgH2. Conversely, one may argue that if the γ−phase decomposes too quickly, then the β-phase may become too stable and small amounts of it may persist even up to high desorption temperatures. Second, during desorption of a milled MgH2 powder and decomposition of β-MgH2, there always occurs a simultaneous growth of nanograins of β-MgH2 (Table 2.13). Since particle size is not changed during desorption, one may hypothesize that the growth of nanograins within a β-MgH2 particle might somehow decelerate the decomposition of β-MgH2. This, however, should be more carefully investigated, since unmilled MgH2 having quite large grain size does not show a reduced desorption capacity (Figs. 2.6 and 2.7). The kinetic curves in Figs. 2.23, 2.24, and 2.26 were obtained after continuous milling in which the milling vial was opened only once and a powder sample was collected from the vial after completion of milling. However, on many occasions samples for various tests are collected periodically after a prescribed period of milling duration such that the milling process is discontinued a number of times during the entire milling process. This kind of processing by ball milling apparently degrades the desorption properties of the milled MgH2 powder even further. Figure 2.28 shows the desorption curves at 325°C of the ABCR powder that was discontinuously ballmilled under HES57 mode for various time durations described in the inset. During milling, the milling vial was periodically opened and a small sample was collected for microstructural studies; then the vial was closed and milling continued. As seen in Fig. 2.28 and described previously, the as-received powder does not desorb at

2.1

Mg/MgH2

121

Hydrogen desorbed [wt.%]

7.00 Desorption

6.00

325ⴗC 5

4 5.00

6

2

4.00 3 3.00 2.00 1.00 1

1-As received 2-0.25 h milling 3-1 h milling 4-5 h milling 5-10 h milling 6-20 h milling

0.00 0

1000

2000

3000

4000

Time [s]

Fig. 2.28 Desorption curves at 325°C of the ABCR powder premilled discontinuously for various times under the HES57 mode in hydrogen

325°C. However, milling for a relatively short time 0.25 h (15 min) makes the powder desorb about 5 wt.%H2 in about 4,000 s. The 5 h milling time seems to be optimal, as the desorbed amount of hydrogen from this powder increases to 6 wt.% after ~3,600 s. Longer milling for 10 and 20 h does not improve desorption kinetics any further. However, in comparison to continuously milled powders in Figs. 2.23, 2.24, and 2.26, the kinetics of discontinuously milled powders in Fig. 2.28 is worse, and after 4,000 s the desorbed amount of hydrogen is less than that desorbed from its continuously milled counterpart at 325°C. It seems that discontinuous milling and, in particular, periodical opening of milling vial, even under protective argon gas, introduces more impurities into the powder despite the fact that both the argon and hydrogen protective gases used in milling were of high purity with only a few ppm of other gas impurities. The hydrogen desorption process investigated using DSC is profoundly affected by milling. Figure 2.29 shows the DSC hydrogen desorption curves of the ABCR powder premilled for various time durations as compared to the curve of the as-received sample. All peaks are endothermic. First, the temperature of the principal desorption peak maximum is gradually shifted from around 408°C for the as-received MgH2 powder to lower temperatures with increasing milling time from 0.25 h to 5 h (Fig. 2.29a). For the 5 h milled powder, the DSC peak temperature is about 30°C lower. However, further milling for 10 and 20 h does not measurably shift the peak maximum to progressively lower temperatures (Fig. 2.29b). Second, a smooth and symmetrical DSC peak of the as-received powder transforms after milling into an asymmetrical peak having a pronounced low-temperature shoulder (Fig. 2.29a,b). The appearance of this shoulder on a DSC pattern can be assigned to the decomposition of mostly γ-MgH2 being formed. An important question now arises: what microstructural factor(s) is (are) responsible for the observed decrease of the DSC desorption peak temperature with increasing

122

2

DSC / mW/mg 7 ↓ exo 6 [1] —— ABCR as received [2] ------ ABCR milled for 15 min 5 [3] –– –– ABCR milled for 1 h [4] –– - –– ABCR milled for 5 h 4 3

Simple Metal and Intermetallic Hydrides

[1] 407.9⬚C

[2] 401.1⬚C

[4] 377.1⬚C

[3] 386.5⬚C

[4] 355.3⬚C

TON

2

[3] 366.6⬚C

1

[2] 378.5⬚C

0 250

300

350

400

450

Temperature / ⬚C

a ↓ exo 4

[2] 371.6⬚C

[1] 382.4⬚C

[1] —— ABCR milled for 10 h [2] ------ ABCR milled for 20 h

3 [1] 358.6⬚C

[2] 388.1⬚C

2 1 0 250

b

300

350

400

450

Temperature / ⬚C

Fig. 2.29 DSC hydrogen desorption curves at the heating rate of 4°C/min of the ABCR powder as received and premilled discontinuously for (a) 0.25 to 5 h and (b) 10 and 20 h under the HES57 mode in hydrogen

milling time? A common belief has been that this is related to the nanostructuring of β-MgH2 hydride, i.e., formation of nanograins within the milled powder particles. However, two factors change during milling, one of them is grain size and the other is particle size. More recently, Hanada et al. [42] attempted such a correlation and reported that the reduction of crystallite size associated with the increasing lattice strains led to the decrease in hydrogen storage capacity while the onset of dehydrogenation temperature was decreased by about 70 K. No reduction of particle size was taken into account in the analysis. Theoretical [43] and experimental [44] attempts trying to explain the role of the particle size of MgH2 on its hydrogen storage properties have been made by other researchers. Wagemans et al. [43] systematically investigated the effect of the size of MgH2 on its thermodynamic stability

2.1

Mg/MgH2

123

using ab initio Hartree–Fock and density functional theory calculations. They reported that the hydrogen desorption energy and the corresponding desorption temperature decrease significantly when the cluster size becomes smaller than ~1.3 nm. Obviously, such a small particle size is most likely unrealistic from any practical point of view but the study clearly indicates the general importance of particle size in the hydriding/dehydriding behavior of MgH2. Shao et al. [44] prepared ultrafine particles of magnesium by hydrogen plasma– metal reaction. The particle size of ultrafine magnesium ranged from 20 to 600 nm with an average size of about 300 nm. The absorption kinetics of hydrogen under 40 bar pressure were studied after activation by one absorption and desorption cycle at 400°C. At 200°C, about 1.9 wt.%H2 was absorbed in ~4,800 s; at 250°C, ~5.7 wt.% was absorbed in ~4,800 s; and at 300–350°C, ~7.6 wt.% was absorbed in ~1,800 s. The absorption kinetics of ultrafine Mg particulate are fast but not faster than the absorption kinetics of milled, activated, and cycled ABCR MgH2 powder as shown in Fig. 2.22. However, the activation of ultrafine Mg particulate involves only one activation cycle. As expected, the hydrogenation of ultrafine Mg results in the exclusive formation of β-MgH2 [44], which is also supported by the recent results of Friedrichs et al. [45] who synthesized a nanoparticulate Mg (particle size 30–50 nm with an average of 35 nm) by a gas-phase condensation method and obtained a nanoparticulate β-MgH2 after hydrogenation (particle size 70–150 nm with an average of 109 nm). This makes a big difference with respect to a milled MgH2, which usually exhibits both γ-and β-MgH2. Shao et al. [44] obtained from the Van’t Hoff plot the formation enthalpy of ultrafine MgH2 as –75kJ/mol. This result shows that even such a substantial particle refinement does not seem to change the enthalpy of formation as long as there is no γ-MgH2 present. This problem will be additionally discussed in the next section. Fátay et al. [46] reported a decreasing hydrogen onset desorption temperature of MgH2doped with Nb2O5with increasing milling time as measured by thermogravimetry analysis (TGA). Their results suggest that particle size is predominantly responsible for the lower onset desorption temperature. The MgH2 particle size was reduced to about 1 μm from the initial 40–50 μm and even more when the premilled 1-μm sized MgH2 powder was mixed with Nb2O5 and additionally milled for 15, 40, and 90 min. The grain size of MgH2 remained within 9–11 nm range regardless of milling time. However, none of the above study has taken into account all the factors that usually change during milling of MgH2. Table 2.8 already shows the decrease of the initial grain size after milling as a function of milling time under the HES57 mode for the ABCR powder, whose DSC curves are shown in Fig. 2.29. Table 2.14 lists the variations of the mean particle size of the same powder. As mentioned previously, there is a simultaneous reduction of the grain size and mean particle size of the hydride powder upon milling. Therefore, we plotted in Fig. 2.30a,b the temperature of DSC peak maximum (Tpeak) and the onset temperature (TON) from Fig. 2.29 as a function of particle and grain size, respectively. In Fig. 2.30b, the symbol LT denotes the temperature maximum of the low temperature shoulder/peak and HT denotes the temperature maximum of the high-temperature shoulder/peak (Fig. 2.29). It is quite clear that in the first stage the DSC peak temperatures only slightly

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2

Simple Metal and Intermetallic Hydrides

Table 2.14 The mean particle size (ECD) of ABCR powder after milling under the HES57 mode and after subsequent cycling as described previously in Sect. 2.1.3.1 After milling After cycling Mean Standard Mean Standard Milling particle size deviation particle size deviation time (h) (ECD) (nm) (SD)(nm) (ECD) (nm) (SD)(nm) As received 40,973 21,116 35,279 21,894 0.25 1,110 959 1,308 586 1 977 806 961 663 5 850 643 989 512 10 844 571 951 614 20 854 634 917 430

decrease with decreasing particle size, but after reaching some critical particle size around 1,000 nm they start decreasing quite dramatically with even small reductions in the particle size. It must be kept in mind that the particles after milling consist of two hydride phases, β-and γ-MgH2, as marked in Fig. 2.30. As discussed previously, thermal cycling eliminates γ-MgH2 from the microstructure of milled powders as evidenced by the XRD pattern in Fig. 2.16. Therefore, the milled powders were thermally cycled in hydrogen (Sect. 2.1.3.1: heating to 325°C for ~15 min under 3.4 MPa H2 to prevent desorption/first desorption at 325°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 3.4 MPa H2 for 30 min/second desorption at 350°C under 0.1 MPa H2 pressure for ~4,700 s/annealing under pre-vacuum at 350°C for 15 min/absorption at 350°C under 2.7 MPa H2 for 30 min) and subsequently tested in DSC (Fig. 2.31). The mean particles sizes (ECD) after cycling are also listed in Table 2.14. Within the experimental scatter, the mean particle sizes after cycling in Table 2.14 do not differ from those after milling. The DSC curves of the cycled samples are shifted back to the right (higher temperatures) as compared to their milled counterparts in Fig. 2.29. What is interesting is the fact that the desorption temperature of cycled (activated) ABCR is also higher than that of the noncycled one. Apparently, the cycled sample is somehow stabilized by thermal cycling as discussed with respect to Fig. 2.9. Nevertheless, the DSC peak temperatures of the cycled samples still show dependence on milling time since their peak temperatures in Fig. 2.31 decrease with increasing time of prior milling. The DSC curves of the cycled samples are smooth and symmetrical without any shoulders. The DSC peak temperatures of the cycled samples from Fig. 2.31 are also plotted in Fig. 2.30a,b as a function of their respective particle size (ECD). Interestingly, the DSC peak temperature curves for cycled samples are shifted up to higher-temperature range with respect to the milled samples. Nevertheless, their shape is very similar to the latter, i.e., gradual decrease in the first stage and subsequent fast decrease after reaching a critical value of the particle size. It must also be pointed out that the grain size of the β-MgH2 phase after milling indicated by the number beside each data point in Fig. 2.30 does not seem to affect the DSC hydrogen desorption temperature in any systematic manner. There are three important observations made in this study: (a) γ-MgH2 is virtually eliminated from the microstructure because of cycling; (b) the DSC peak temperatures of

2.1

Mg/MgH2

125

430 420

166 nm

HES57

b–MgH2

410

DSC TON [ⴗC]

400 390 380

34 nm 43 nm

40 nm

52 nm

b –and g – MgH2

30 nm

340

a

As-received

11 nm

350

330 100

299 nm

47 nm

35 nm

370 360

As-received-cycled

After milling After milling and cycling

12 nm 10 nm 1000

10000

100000

Powder particle size - ECD [nm] 430 420

b–MgH2

DSC Tpeak [ⴗC]

410

b – and g – MgH2

400 390 380 370 After milling-HT After milling-LT After cycling-HT

360 350 100

b

1000

10000

100000

Powder particle size - ECD [nm]

Fig. 2.30 Changes of DSC hydrogen desorption temperatures from Fig. 2.29 as a function of the particle size (ECD) of (a) milled (HES57 mode) and (b) cycled ABCR powder. Numbers beside each data point indicate the grain size of β-MgH2. (a) Onset temperature (TON) and (b) peak temperatures (Tpeak)). Standard deviations for the mean particle size (ECD) from Table 2.14 are omitted for clarity

the thermally cycled powder samples are shifted to a temperature range which is higher than that of the milled samples; and (c) the DSC peak temperatures of the cycled samples still exhibits dependence on the mean particle size within this higher range of temperatures. Accordingly, on the basis of this quantitative evidence it can be concluded that two microstructural factors, the γ-MgH2 phase residing within the

126

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Simple Metal and Intermetallic Hydrides

DSC / mW/mg [1] 427.5⬚C

↓ exo 8 6

All powders cycled [1] —— ABCR as received [2] ------ ABCR milled for 15 min [3] –– –– ABCR milled for 1 h [4] –– - –– ABCR milled for 5 h

4

[2] 418.2⬚C [4] 405.1⬚C [3] 415.4⬚C

2 0 350

300

a

400 Temperature / ⬚C

450

500

450

500

DSC / mW/mg 5 4 3

[2] 402.8⬚C

↓ exo Powders cycled after milling [1] —— ABCR milled for 10 h [2] ------ ABCR milled for 20 h

[1] 401.1⬚C

2 1 0 300

b

350

400 Temperature / ⬚C

Fig. 2.31 DSC hydrogen desorption curves at the heating rate of 4°C/min of the ABCR powder, premilled discontinuously under the HES57 mode in hydrogen and subsequently cycled as described in Sect. 2.1.3.1. (a) Milled for 0.25 to 5 h in hydrogen and cycled and (b) milled for 10 and 20 h in hydrogen and cycled

powder particles and refined powder particle size acting additively, are responsible for a substantial reduction of hydrogen desorption temperature of MgH2 hydride as observed in DSC. However, it seems from Fig. 2.30 that the presence of the γ-MgH2 phase is a predominant factor leading to the decrease of the desorption temperature. The range of the desorption temperature drop due to just the particle size seems to be smaller than that due to the presence of the γ-MgH2 phase. A very similar dependence of the hydrogen desorption temperatures measured in DSC on the particle size of powders has been found for the Tego Magnan MgH2 powder after milling under IMP2 mode with a single magnet [6]. Figure 2.32 shows

2.1

Mg/MgH2

127 DSC / mW/mg

DSC / mW/mg ↓ exo 8

4.0

[1] 414.6⬚C

3.5

As received commercial b-MgH2 (TEGO MAGNANTM)

6

[1] 404.2⬚C

↓ exo

15 min milling b-MgH2

3.0 2.5 2.0

4

1.5 1.0

2

0.5 0

0

a

300

320

340

360 380 400 Temperature / ⬚C

440

HT

DSC / mW/mg 4.0

420

b

3.5

1 h milling b-MgH2

2.5

440

420

440

400

420

440

400

420

440

[1] 361.7⬚C

1.0 [1] 371.5⬚C

0.5 0

0 300

320

340

360 380 400 Temperature / ⬚C

420

440

DSC / mW/mg

1.5

300

320

340

360 380 Temperature / ⬚C

400

[1] 379.2⬚C

3.0

50h milling b-MgH2+g-MgH2

2.5

25 h milling b-MgH2+g-MgH2

2.0

d

DSC / mW/mg ↓ exo

[1] 388.6⬚C

2.5 ↓ exo

2.0

[1] 406.9⬚C

[1] 394.5⬚C

1.5

1.0

1.0

0.5

0.5 0

0 300

320

340

360 380 Temperature / ⬚C

400

420

440

DSC / mW/mg 2.5

420

[1] 380.8⬚C

1.5

LT

TON

0.5

e

400

2.0

1.0

c

360 380 Temperature / ⬚C

10 h milling b-MgH2+g-MgH2

2.5

2.0 1.5

340

↓ exo

3.0

3.5 3.0

320

DSC / mW/mg

[1] 395.0⬚C

↓ exo

300

f

320

340

360 380 Temperature / ⬚C

DSC / mW/mg

↓ exo

2.0

−0.5 300

2.5

[1] 376.1⬚C [1] 385.7⬚C

75h milling b-MgH2+g-MgH2

2.0

[1] 380.0⬚C

↓ exo

100h milling b-MgH2+g-MgH2

[1] 365.2⬚C

1.5 1.5 1.0

1.0

0.5

0.5 0

g

300

0 320

340

360

380

Temperature / ⬚C

400

420

440

300

h

320

340

360 380 Temperature / ⬚C

Fig. 2.32 DSC curves for the as-received MgH2 (Tego Magnan) and ball-milled MgH2 powders (IMP2; single magnet). (a) As-received commercial MgH2. Milled for (b) 15 min, (c) 1 h, (d) 10 h, (e) 25 h, (f) 50 h, (g) 75 h, and (h) 100 h in hydrogen (heating rate 4°C/min) [6]

128

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Simple Metal and Intermetallic Hydrides

the DSC curves of the as-received Tego Magnan powder and the ball-milled powders. A single endothermic hydrogen desorption peak with the maximum at ~415°C characterizes hydrogen desorption from a coarse-particle, as-received Tego Magnan powder. MgH2 ball-milled for 15 min also shows a single hydrogen desorption peak at ~404°C although much wider than that for the as-received powder (Fig. 2.32b). With increasing milling time, the formation of a shoulder is observed on the DSC curves (Fig. 2.32c–h). This shoulder is most likely due to the decomposition of γ-MgH2 phase as discussed previously. Figure 2.33 shows the plots of DSC desorption temperatures as a function of powder particle size (ECD) of the milled Tego Magnan powder. It can be seen in 430 420 MgH2 - as-received - 67 nm

DSC TON [ⴗC]

410 400 390 380 370 360 350 340

73 nm 13 nm 14 nm 71 nm 7 nm 19 nm 6 nm

330 100

1000

a

10000

100000

Powder particle size - ECD [nm] 430 MgH2 - as-received - 67 nm

420

13 nm 73 nm

DSC Tpeak [ⴗC]

410 400

71 nm

390 380

14 nm

7 nm 19 nm

6 nm

370

HT

360 350 100

b

LT 1000

10000

100000

Powder particle size - ECD [nm]

Fig. 2.33 DSC hydrogen desorption temperatures vs. particle size for as-received and ball-milled Tego Magnan powder. (a) Onset temperature (TON) and (b) low-temperature (LT) and high-temperature (HT) DSC peaks. Numbers beside data points indicate grain (crystallite) size of the β-MgH2 phase. Standard deviation bars for the particle size (ECD) are omitted for clarity [6]

2.1

Mg/MgH2

129

Fig. 2.33a that the onset temperature (TON) initially slowly decreases with decreasing mean ECD down to ~2,000 nm (2 μm), and then starts decreasing more rapidly with further decrease of ECD. The total drop of TON from its initial value for the as-received MgH2 to the value attained for the milled MgH2 with the particle size of ~500–600 nm is about 60°C [6]. This is a very substantial reduction of the desorption onset temperature. In Fig. 2.33b the LT and HT peak temperatures behave in a similar manner as the onset temperature: first decreasing slowly and after achieving some critical (threshold) range of ECD, more rapidly. The total observed drop of LT and HT peak temperatures from the initial values for the as-received MgH2 value is about 40–50°C. It is to be noticed that the biggest change in particle size by a factor of 15 from ~36 μm for the as-received MgH2 down to ~2 μm for the 15 min milled MgH2 yields a smaller difference in the lowering of DSC temperature than further ECD reduction by a factor of about 4 from ~2 to ~0.47 μm (470 nm) after milling for 10 h [6]. This clearly shows that a critical (threshold) range of particle size is required to achieve a substantial reduction in the DSC hydrogen desorption temperature. In turn, the nanograin (crystallite) size corresponding to each milled powder is also shown beside each data point in Fig. 2.33a,b. It is very clear that within the investigated range of initial crystallite (grain) sizes from ~67 nm down to ~3 nm, there does not seem to be a correlation between the nanograin size of MgH2 and its corresponding DSC desorption temperature. Once again, this evidence reinforces the previous statement that the presence of the γ-MgH2 phase is a predominant factor, which, in combination with the reduction of particle size, leads to the decrease of the desorption temperature. The crystallite (grain) size seems to have a secondary role. This secondary role may (or may not) be reflected in the reduction of the hydrogen capacity of a premilled powder after desorption, which is usually lower than the purity-corrected capacity, owing to nanograin growth at the desorption temperatures (Table 2.13) when the average particle size remains unaffected by desorption (Table 2.14).

2.1.4

Hydrogen Storage Characteristics of MgH2 Synthesized by Reactive Mechanical (Ball) Milling of Mg

As pointed out in Sect. 2.1.3, in the early works on the synthesis of MgH2 [29–33] the elemental Mg metal was simply ball-milled under argon and subsequently subjected to hydrogenation at the appropriate temperature and under required hydrogen pressure in a separate step. Such a “two-step” method is rather cumbersome and time consuming. Mg is not the best material for ball milling, as it requires various chemical additives that can prevent cold welding of this very soft metal. Indeed, a heavy deformation during ball milling of Mg can reduce its grain size even to a nanosized level, but cold welding usually prevents a substantial particle refinement. Much more efficient method of the synthesis of metal hydrides is a “one-step” method of reactive ball milling of elemental metal directly under hydrogen gas. The method of reactive mechanical alloying/milling of either elemental metals or prealloyed

130

2

Simple Metal and Intermetallic Hydrides

intermetallic compounds under active gas atmospheres such as nitrogen [47–50] and hydrogen [50, 51–60] became popular in the 1990s. A major realization from this early research on reactive mechanical milling/alloying was that a substantial nanostructuring, i.e., formation of nanograins (nanocrystallites) within the synthesized powder particles of greatly reduced sizes, can be achieved quite easily and simultaneously with the formation of a hydride. An overview of nonreactive and reactive mechanical milling/ alloying processes for the synthesis of nanocrystalline intermetallic hydrides that were published in the open literature up to 2002 can be found in [61]. For the sake of clarity, it must be pointed out that in the literature various authors have used quite liberally the terms such as “mechanical alloying” and “mechanical milling” without paying too much attention to the exact physical nature of the milling process. For example, Chen and Williams [54] titled their pioneering paper “Formation of metal hydrides by mechanical alloying.” In reality, the technique they used was a simple milling of elemental metals (Mg, Ti, and Zr) under hydrogen. Therefore, it is probably more appropriate to name such a process “mechanical milling” rather than “alloying.” To avoid any confusion, in the present work we reserve the term “reactive mechanical alloying” to the process in which there is a simultaneous formation of an intermetallic compound (also amorphous) and its hydride, e.g., Mg2NiH4 or something similar (Sect. 1.3.3). Formation of a simple metal hydride by milling under hydrogen will be preferentially called “reactive mechanical milling.” The first successful synthesis of MgH2 by reactive mechanical milling was carried out by Chen and Williams [54], as mentioned above. Besides MgH2 they also synthesized such metal hydrides as TiH1.9 and δ-ZrH1.66. Only limited desorption studies by DSC were carried out. This first effort was followed several years later by a more thorough study of hydrogen desorption from a reactively synthesized magnesium hydride by Gennari et al. [41]. Their major finding, as quoted several times in the previous sections, was the discovery of a peak doublet associated with the hydrogen desorption from the γ-MgH2 + β-MgH2 mixture at a lower temperature (LT peak in the doublet), where a fraction of β-MgH2 is destabilized by desorbing γ-MgH2, and eventually desorption from the remaining β-MgH2 at a higher temperature (HT peak in the doublet). It is rather surprising that only a handful of papers have been published on the reactive mechanical milling synthesis of MgH2. In order to cover this gap, in the past few years we have been carrying a thorough research on various aspects of the reactive synthesis of MgH2. Some results have already been published [62]. A large number of samples of elemental Mg powder (325mesh, 99.8% purity) were milled under hydrogen atmosphere in the magneto-mill Uni-Ball-Mill 5 (Sect. 1.3.2) using a wide variety of milling parameters (Table 2.15). Particle size as well as grain size and lattice strain of the phases embedded within the particles were calculated and are listed in Tables 2.16 and 2.17. Thermal hydrogen desorption behavior of the milled powders was studied by DSC in a Netzsch Model 404 system at a heating rate of 4°C/min and an argon flow rate of 16 or 50 ml/min. Desorption kinetics were investigated with a Sieverts-type apparatus (Sect. 1.4.2). The results are compiled in Table 2.17. Typical evolution of particle size (ECD) as a function of reactive milling time of selected synthesized powders from Table 2.15 is shown in Fig. 2.34. SEM micrograph

Milling

Time (h)

Mg1H Hydrogen (750kPa) 100 30 40:1 2 2 HES S 5,20,50,100,150t Mg7H Hydrogen (880kPa) 100 30 40:1 4 2 HES S 5,20,50,100,150 Mg6H Hydrogen (880kPa) 100 30 40:1 4 2 HES C 150 Mg2H Hydrogen (700kPa) 170 55 40:1 2 2 IMP2 S 100 Mg4H Hydrogen (880kPa) 200 60 40:1 4 2 IMP2 C 150 Mg5H Hydrogen (880kPa) 200 60 80:1 4 2 IMP2 C 150 Mg9H Hydrogen (880kPa) 225 70 40:1 4 2 IMP710 C 60 Mg10H Hydrogen (880kPa) 225 70 40:1 5 2 IMP710 C 60 Mg11H Hydrogen (880kPa) 225 70 20:1 4 2 IMP710 C 100 Mg12H Hydrogen (880kPa) 230 75 40:1 5 2/10 IMP68 C 30 Mg13H Hydrogen (880kPa) 230 75 100:1 5 2/10 IMP68 C 30 Mg14H Hydrogen (880kPa) 230 75 20:1 5 2/10 IMP68 C 30 HES high-energy shearing with one magnet; IMP2 strong impact mode with one magnet; IMP710 impact mode with two magnets at seven and ten o’clock position; IMP68 milling mode with two magnets at six and eight o’clock position; S sequential milling; C continuous milling.

Table 2.15 Composition of powders and processing parameters for reactive mechanical milling of Mg-H powders [63] Powder Ball to powder No. WD description Gas RPM Speed (%) weight ratio of balls (mm) Mode

2.1 Mg/MgH2 131

132

2

Simple Metal and Intermetallic Hydrides

Table 2.16 Nanograin size and strain of b-MgH2 synthesized by reactive milling of Mg–H powders (calculated as shown in Sect. 1.4.3) [63] Nanograin size, L, and strain (e) of β-MgH2 Powder Mg1H

Mg7H

Mg6H Mg2H Mg4H Mg5H Mg9H Mg10H Mg11H Mg12H Mg13H Mg14H

Milling time (h)

L (nm)

e

R2

No of XRD peaks

5 20 50 100 150 5 20 50 100 150 150 100 200 150 150 60 60 100 30 30 30

39 13 9 7 6 42 15 9 8 7 8 8 8 8 8 8 7 8 9 9 9

0 0 0 0 0 3.5 × 10−4 1.4 × 10−3 0 0 0 0 0 0 0 0 0 0 0 0 0 0

0.763 0.909 0.991 0.988 0.99 0.949 0.961 0.991 0.993 0.981 0.979 0.979 0.978 0.989 0.990 0.968 0.990 0.983 0.994 0.988 0.986

4 5 6 6 6 7 6 7 7 7 9 4 4 5 5 5 5 8 10 10 9

insets in Fig. 2.34 show the morphology of powders. After 20 h of milling under the high-energy shearing (HES) mode using only one magnet (Sect. 1.3.2), the average ECD particle size of the powder is already reduced from the mean 9 μm (Fig. 2.34a) to just ~1.8 μm (Fig. 2.34b). An extension of the synthesis time either under HES to 50 h (Fig. 2.34c) or IMP2 (one magnet) to 150 h (Fig. 2.34d) further reduces the particle size to ~0.71 and ~0.34 μm, respectively. Such a continuous reduction of particle size during reactive synthesis of MgH2 differs from a saturation of the particle size at ~1 μm after a relatively short time during a simple mechanical milling of a commercial MgH2 (Fig. 2.15). If one compares the ECD values of ball-milled commercial MgH2 powder in Table 2.14 with the ECD values of reactively milled MgH2 powders in Table 2.17 and Fig. 2.15, it seems that during reactive milling of Mg the rate of particle size reduction is slower, but eventually much smaller final particle size can be achieved after a long milling duration of 100–150 h. Reactive milling for only a short time results in a profound agglomeration of soft Mg particles as can be seen for Mg7H sample milled for 5 h in Table 2.17. It is also clearly seen in Fig. 2.34 that the relative frequency distribution of powder particles remains log-normal starting from Mg through any powder regardless of the duration of reactive milling time. This is exactly the same behavior as already described for mechanically milled commercial MgH2 powders (Fig. 2.20). The experimental coefficient of variation, CV(ECD) = SD(ECD)/M(ECD) (where SD is the standard deviation

HES

HES

HES IMP2 IMP2 IMP2 IMP710 IMP710 IMP710 IMP68 IMP68 IMP68

Mg1H

Mg7H

Mg6H Mg2H Mg4H Mg5H Mg9H Mg10H Mg11H Mg12H Mg13H Mg14H

5 20 50 100 150 5 20 50 100 150 150 100 150 150 60 60 100 30 30 30

38 13 9 7 6 33 15 9 8 7 8 8 8 8 8 7 8 9 9 9

– 1,789 ± 1,432 712 ± 570 – – 39,671 ± 33,930 2,123 ± 1,783 444 ± 363 384 ± 281 374 ± 263 393 ± 255 374 ± 162 374 ± 233 338 ± 201 – – – 728 ± 456 643 ± 399 563 ± 347

– 0.80 0.80 – – 0.86 0.84 0.82 0.73 0.70 0.65 0.43 0.62 0.59 – – – 0.63 0.62 0.62

311.8 372.1 361.2 341.1 339.5 398.8 375.3 350.8 349.0 351.8 351.7 335.4 347.0 326.0 351.8 351.0 350.5 346.2 348.7 347.9

– 380.8 369.5 352.8 354.2 – 387.3 372.9 365.5 370.5 371.3 349.7 361.6 348.3 375.2 367.8 365.5 365.9 367.5 363.8

375.7 406.8 408.4 392.4 382.1 410.5 398.1 389.3 397.7 392.7 391.3 376.3 380.3 369.0 394.6 388.6 392.0 392.1 379.7 383.8

5.8 35.3 53.2 54.9 56.0 9.8 48.5 67.1 73.0 74.6 78.1 84.3 81.2 74.9 75.0 71.4 70.1 84.7 86.0 88.3

Table 2.17 Selected results of XRD analysis and DSC analysis of the reactively milled powders in the Mg–H system [63] DSC peak (°C) (at heating rate 4°C/ Mean particle size min) (ECD) ± SD* Nanograin Yield of CV (ECD) Sample of hydride size (nm) of for size MgH2 LT HT description Mode Time (h) TON (wt.%) powders (nm) β-MgH2 distribution

2.1 Mg/MgH2 133

2

Fig. 2.34 Particle size distributions and morphology of (a) initial Mg and selected synthesized MgH2 powders from Table 2.15 reactively milled for (b) 20 h under HES (Mg1H20), (c) 50 h under HES (Mg1H50), and (d)150 h under IMP2 (Mg5H) [62]

134 Simple Metal and Intermetallic Hydrides

2.1

Mg/MgH2

135 Mg

16000

b-MgH2 MgO g-MgH2

Counts

12000

5h

8000

20h 50h

4000

100h 150h

0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 2.35 A general example of the evolution of the XRD patterns as a function of reactive milling time [63]

and M is the mean), which is listed in Table 2.17, indicates the uniformity of the particle distribution such that relatively low values of CV(ECD) on the order of 0.5–0.6 and less indicate uniformity and much larger values point towards rather nonuniform distribution, i.e., the existence of small and large size populations of particles [62]. Figure 2.35 shows the general evolution of the XRD patterns as a function of reactive milling time, which can be observed regardless of the mode of milling in Table 2.15. Tetragonal β-MgH2 is formed in the initial period of reactive milling (5 h) and coexists with the still unreacted Mg. The latter is completely converted into MgH2 between 50 and 100 h of milling. A metastable orthorhombic γ-MgH2 appears between 20 and 50 h of milling under the HES, IMP2, IMP710, and IMP68 modes (Table 2.15). The only exception was the Mg1H powder milled under the HES mode with only two balls (Table 2.15), in which the only hydride phase was β-MgH2 after milling for more than 20 h. No further changes are observed between 100 h and 150 h of milling. Small peaks of MgO in Fig. 2.35 are due to the oxides always present on the elemental metal powders, specifically Mg, and additionally to the reaction of highly reactive, milled nanostructured retained Mg with atmospheric oxygen during XRD tests [41, 64–66]. XRD patterns allow the calculation of grain size/lattice strain of hydride phases from broadening of their respective Bragg peaks as described in Sect. 1.4.3. The values of grain sizes and strains are listed in Tables 2.16 and 2.17. It must be pointed out that the lattice strains of β-MgH2 are either nonexistent or very minimal. This is not very surprising since, as mentioned in the preceding section, the β-MgH2 phase is tetragonal and its atomic bonding is mostly ionic with a little covalency [2, 4], i.e., ceramic-like, and one can hardly envision substantial dislocation activities

136

2

Simple Metal and Intermetallic Hydrides

40 Mg1H/HES-S

Nanograin size of β-MgH2 (nm)

35

Mg7H/HES-S Mg6H/HES-C

30

Mg2H/IMP2-S Mg4H/IMP2-C

25

Mg5H/IMP2-C Mg9H/IMP710-C

20

Mg10H/IMP710-C Mg11H/IMP710-C

15

Mg12H-14H/IMP68-C

10 5 0 0

50

100

150

200

250

Milling time (h)

Fig. 2.36 Nanograin (crystalline) size of β-MgH2 phase synthesized in the Mg–H powders under different reactive milling modes (Table 2.16). S sequential milling; C continuous milling [63]

that would lead to dislocation accumulation in the lattice and, in turn, strains of any large magnitude. Figure 2.36 shows the nanograin (crystalline) size of β-MgH2 in the Mg–H powders milled under different milling modes (values from Table 2.16). It can be seen that in all the milled powders the β-MgH2 phase undergoes a very fine nanostructuring and becomes saturated (crystallite/grain size ~6–9 nm) after ~30–50 h of reactive milling time. Further refinement proceeds slowly, and the crystallite/ nanograin size finally reaches a steady-state value of ~6–9 nm rather than decreasing indefinitely. This behavior is quite similar to the one observed for the grain size of mechanically milled commercial MgH2 powders shown in Fig. 2.14. However, it seems that the saturation of the grain size of β-MgH2 occurs faster during the mechanical milling of a commercial product. On the basis of the heavy peak broadening, it can be reasonably concluded that the grain size of γ-MgH2 is most likely on the same level as that of β-MgH2, i.e., within the 10 nm range [6]. Figure 2.37 shows the yield of the MgH2 phase estimated by DSC analysis (Sect. 1.4.4) as a function of milling time in the Mg–H powders milled under various milling modes. The dependence of the yield of MgH2 on the milling time can be best represented using as examples the Mg1H and Mg7H powders milled under the HES mode (Tables 2.15 and 2.17). The yield of MgH2 first increases dramatically in the early stage of milling and attains a saturation level after milling for 100 h. The lowest yield of MgH2 (~56.0 wt.%) at the saturation level is found in the Mg1H powder after milling for 150 h under the HES mode with a single magnet, which in general is not a very energetic mode, while the highest value is ~88 wt.% in the Mg14H powder after milling for 30 h under the IMP68 (two magnets) mode, which is an extremely energetic mode of milling. This result also shows clearly that achieving

2.1

Mg/MgH2

137 100

88.3 (β+γ)

90

86 (β+γ) 84.7 (β+γ)

Yield of MgH2 from DSC (wt.%)

80

70

84.3 (β+γ) 81.2 (β+γ) 78.1 (β+γ) 74.6 (β+γ)

73.0 (β+γ)

75.0 (β+γ) 71.4 (β+γ)

70.1 (β+γ) 74.9 (β+γ) 67.1 (β+γ)

60 53.2 (β)

50

48.5 (β)

56.0 (β) 54.9 (β)

Mg1H/HES-S Mg7H/HES-S Mg6H/HES-C

40

Mg2H/IMP2-S Mg4H/IMP2-C

35.3 (β)

30

Mg5H/IMP2-C Mg9H/IMP710-C Mg10H/IMP710-C

20

Mg11H/IMP710-C Mg12H/IMP68-C

10

9.8 (β)

Mg13H/IMP68-C

5.8 (β)

Mg14H/IMP68-C

0 0

50

100

150

200

250

300

Milling time (h)

Fig. 2.37 DSC yield of MgH2 phase in the milled powders as a function of milling time. Numbers beside data points indicate DSC yield of MgH2 (wt.%). β or β + γ indicates that the milled powders contain only β-MgH2 or both of γ- and β-MgH2 phases, respectively. S sequential milling; C continuous milling [63]

the yield of synthesized MgH2 much higher than 90% during reactive mechanical milling even under very energetic mode is quite difficult and most probably impossible in practice. Using the yield of MgH2 and milling time as rough estimators of the efficiency of milling modes applied in the present work in the Uni-Ball-Mill 5, the weakest to the strongest mode is HES < IMP2 < IMP710 < IMP68. It would be interesting to compare the present results with the results of a similar study in which reactive milling was carried out in a different type of ball mill, e.g., planetary. However, a study of this type is unavailable in the open literature. In the Mg1H, Mg7H, Mg9H, and Mg10H powders, the ball-to-powder weight ratio was kept at 40:1, while the number of balls was increased from two balls in Mg1H to four balls in Mg7H powder, and from four balls in Mg9H to five balls in Mg10H powder. Yield of MgH2 in Mg7H powder (74.6 wt.%) after milling for 150 h is much higher than that in Mg1H powder (56.0 wt.%) after 150 h under the same

138

2

Simple Metal and Intermetallic Hydrides

7

7

6

6

5 4 3 2

Mg9H-IMP710-40:1

1

Absorbed H2 (wt.%)

Absorbed H2 (wt.%)

milling mode (Fig. 2.37). When number of balls increases from two to four under the HES mode with the same ball-to-powder weight ratio, the yield of β-MgH2 increases. This might be the result of increased contacts between balls and powders when the number of balls increases. Energy applied to the powder particles increases as the number of collisions (contacts) increases. As a result, the yield of the product increases. However, the dramatic increase of the yield of β-MgH2 is not observed while milling under IMP710 mode as the number of balls changes from four in Mg9H (~75 wt.%) to five in Mg10H (~71 wt.%) (Fig. 2.37). It seems that four balls are required to produce sufficient numbers of contacts between balls and powders in the Uni-Ball Mill 5, which results in a very efficient reactive milling. The effect of ball-to-powder weight ratio on the efficiency of reactive milling was also assessed (Fig. 2.38). As can be seen in Fig. 2.38a, longer time is required to reach the saturation level of absorbed hydrogen when the ball-to-powder weight ratio is reduced from 40:1 to 20:1 under the IMP710 mode. This indicates that milling time is dependent on the ball-to-powder weight ratio. As a general rule for mechanical alloying processes, Suryanarayana [67] and Lü et al. [68] reported that the higher the ball-to-powder weight ratio, the shorter is the milling time. However, under the IMP68 mode (Fig. 2.38b) the time to reach the saturation with a 40:1 ball-to-powder weight ratio is very close to that obtained with a 20:1 ratio. Since IMP68 is slightly more energetic than IMP710, although both are impact modes, it can be concluded that when the number of balls is fixed, the milling time is a function of the ball-to-powder weight ratio and milling mode, while the latter has the stronger effect. In addition, as shown in Fig. 2.39 the magnitude of ball-to-powder weight ratio does not have any measurable effect on the particle size of MgH2 synthesized by controlled reactive mechanical milling (CRMM) in the magneto-mill Uni-Ball-Mill 5. The corresponding DSC curves of Mg powders milled in hydrogen atmosphere after various milling times under the HES mode are shown in Fig. 2.40. The 5 h milled sample has only one small peak at 410.5°C, but the others show double peaks (peak doublets). Desorption peak temperatures of the doublet peak first decrease as

Mg11H-IMP710-20:1

4 3 2

Mg14H-IMP68-20:1

1

Mg12H-IMP68-40:1

0

0 0

a

5

20

40 60 80 Milling time (h)

100

120

0

b

5

10 15 20 25 Milling time (h)

30

35

Fig. 2.38 The effect of ball-to-powder weight ratio on the efficiency of milling under two impact modes (a) IMP710 and (b) IMP58 in the magneto-mill Uni-Ball-Mill 5 [63]

2.1

Mg/MgH2

139

Particle size ECD (mm)

1.4 1.2 1 0.8 0.6 0.4 0.2 0 0

20

40 60 80 Ball/powder ratio

100

120

Fig. 2.39 The effect of ball-to-powder weight ratio on the particle size of MgH2 synthesized by reactive ball milling for 30 h in the magneto-mill Uni-Ball-Mill 5 [63] DSC / mW/mg 3.0

exo

[5] 370.5⬚C [3] 389.3⬚C

5

2.5

1.0

[2] 387.3⬚C [2] 398.1⬚C

[4] 365.5⬚C

2.0 1.5

[5] 392.7⬚C [4] 397.7⬚C

1. 5h 2. 20h 3. 50h 4. 100h 5. 150h

4 [1] 410.5⬚C

3

0.5

2

1

0 −0.5 300

320

340

360 380 Temperature / ⬚C

400

420

Fig. 2.40 Evolution of representative DSC curves for synthesized MgH2 as a function of reactive milling time in the HES mode (heating rate 4°C/min) [63]

the milling time increases and then stabilize after 100 h. Regardless of the microstructure present in the milled powder, the DSC curve of a powder milled for less than 20 h shows a single peak (e.g., 5 h), while those milled for longer than 20 h show either a hidden peak (shoulder) or a doublet (Fig. 2.41). XRD analysis performed on the samples after DSC desorption only revealed Mg and MgO phases (XRD pattern not shown here). This indicates that the decomposition of γ- and β-MgH2 is completed after DSC heating. The origin of the shoulder or the doublet in the single-phase synthesized β-MgH2 is most likely due to the presence of fine and coarse particle fractions in the powders and possibly in association with MgO formed on smaller

140

2

Simple Metal and Intermetallic Hydrides

DSC / mW/mg HT

4.0 ↓ exo 3.5 3.0 2.5

LT: low-temperature peak

2.0

HT: high-temperature peak

1.5

TON: peak on-set temperature

LT

1.0 0.5

TON

(a)

0 200

(b) 250

300

350 Temperature / ⬚C

400

450

500

Fig. 2.41 Representative DSC curves of MgH2 powders synthesized by reactive milling for more than 20 h showing (a) peak with a shoulder (hidden peak) and a well-developed peak or (b) double peaks (doublet) (heating rate 4°C/min) [63]

particles. These different fractions may desorb at slightly different temperature ranges, leading to a shoulder/doublet [62, 69]. If the synthesized powders contain both γ- and β-MgH2 then the appearance of a doublet can be explained by the mechanism put forward by Gennari et al. [41]: i.e., first, the hydrogen desorption from the γ-MgH2 + β-MgH2 and then from the remaining β-MgH2 as discussed earlier in the text. Data from Table 2.17 were used to construct the dependence of DSC desorption temperature of synthesized MgH2 on its particle size (ECD) similar to Figs. 2.30 and 2.33 for the ball-milled commercial MgH2, as shown in Fig. 2.42. In essence, the behavior is very similar to that seen in Figs. 2.30 and 2.33. Both the onset (TON) and peak (LT and HT) desorption temperatures either slowly decrease or remain constant up to a certain critical (threshold) value/range of ECD and then dramatically decrease with even a slight decrease of ECD. An interesting observation is that for the synthesized MgH2 in Fig. 2.42 the critical mean ECD range lies roughly within 1,000–750 nm as opposed to ~1–2 μm (1,000–2000 nm) for the commercial MgH2 subjected to a nonreactive, conventional mechanical milling in Figs. 2.30 and 2.33. As discussed in Sect. 2.1.3.3, the presence of γ-MgH2 in the mixture with β-MgH2 is the primary factor and the reduced particle size is secondary factor, which combined lead to the reduction of the hydrogen desorption temperature of MgH2. From this view point, it can be hypothesized that during mechanical milling of the commercial MgH2 powder the γ-MgH2 phase is produced more easily in a fraction of particles with larger sizes than during reactive synthesis, which results in a desorption temperature drop at the greater particle size range for the former. Thermodynamic parameters of the Mg6H powder synthesized for 150 h (Tables 2.15 and 2.17) were investigated by PCT curves for absorption and desorption at various temperatures as shown in Fig. 2.43a. The powder after such a long reactive

2.1

Mg/MgH2

141

420 67nm

400 DSC TON (ⴗC)

33nm 15nm

380 14nm

360 340 320

Grain size in the range of 7-9nm

300 100

1000

a

10000

100000

Powder particle size - ECD (nm) 420 9 nm

410

33 nm 8 nm

DSC Tpeak (ⴗC)

400 390 380 370

67nm

14 nm

8 nm 9 nm 8 nm 8 nm

7 nm

15nm 9 nm

9 nm 9nm

8 nm

360 LT

350 340 100

b

1000

10000

HT

100000

Powder particle size - ECD (nm)

Fig. 2.42 DSC desorption temperature from Table 2.17 versus powder particle size (ECD) of synthesized MgH2 powders. (a) Onset temperature (TON) and (b) low-temperature (LT) and high-temperature (HT) peaks in a doublet. Numbers beside data points indicate the grain size of the MgH2 phases [63]

milling contains both γ-MgH2 and β-MgH2 (Fig. 2.35). In order to eliminate the presence of g-MgH2 and its effect on PCT curves, the powder was desorbed at 350°C under pre-vacuum for 60 min. Subsequently, absorption/desorption cycles for acquiring PCT curves were conducted as follows: (1) Absorption 350°C; desorption 350°C; annealing 325°C/pre-vacuum/30 min temperature stabilization; (2) Absorption 325°C; desorption 325°C; annealing 300°C/pre-vacuum/30 min temperature stabilization;

142

2

Simple Metal and Intermetallic Hydrides

10

Pressure (MPa)

375C abs 350C abs 325C abs 300C abs

375C des 350C des 325C des 300C des

1

0.1 0

1

a

2

3 4 5 H2 capacity (wt.%)

6

7

8

3.5

Absorption 3

DHabs = -72 kJ/mol

2.5

ln p

2

R2 = 0.9815

Desorption 1.5

DHdes = 83kJ/mol 1

R2 = 0.9934 0.5 0 1.5

b

1.55

1.6

1.65

1.7

1.75

1.8

1000/T (1/K)

Fig. 2.43 (a) PCT curves for absorption and desorption of the Mg6H powder reactively synthesized for 150 h (Tables 2.15 and 2.17). (b) Van’t Hoff plot (ΔHabs = –72 kJ/mol, ΔHdes = 83 kJ/mol, ΔSabs = 138 J/mol K and ΔSdes = 151 J/mol K)

(3) Absorption 300°C; desorption 300°C; annealing 375°C/pre-vacuum/30 min temperature stabilization; (4) Absorption 375°C; desorption 375°C. The plateau region on the PCT curves in Fig. 2.43a is very flat, showing only a minimal slope (Sect. 1.4.1). However, each absorption–desorption pair of the PCT curves clearly exhibits a pressure hysteresis. This means that the pressure needed for absorption (hydride formation), pabs, is always greater than that of hydride decomposition, pdes. The cause of pressure hysteresis in metal hydride systems is not fully understood. A number of models that attempted to explain this phenomenon

2.1

Mg/MgH2

143

Table 2.18 Ratios of Pabs/Pdes as a function of temperature for the Mg6H powder from Fig. 2.43a PCT temperature (°C)

Mid-plateau pabs (MPa)

Mid-plateau pdes (MPa)

pabs/pdes

Δp (pabs–pdes) (MPa)

375 350 325 300

2.192 1.632 0.782 0.402

1.57 0.887 0.492 0.205

1.396 1.84 1.589 1.976

0.622 0.612 0.358 0.197

have been proposed throughout the last 70 years. Models developed up to 1992 were reviewed by Qian and Northwood [70, 71] and Esayed and Northwood [72]. A general observation is that PCT pressure hysteresis seems to be an inherent feature of all metal–hydrogen systems. If hysteresis is measured by the ratio pabs/pdes, its degree usually decreases with increasing temperature (or increases if it is measured as Δp). As shown in Table 2.18, this rule is reasonably well obeyed by the PCT curves in Fig. 2.43a. However, at this point it is important to recall that Bogdanović et al. [17] reported very small and in reality almost nonexistent absorption/desorption hysteresis during absorption/desorption experiments with Mg particles in the temperature range 403–520°C. In a nutshell, most of these early models attribute the hysteresis to either the elastic or plastic accommodation of the volume change during hydride formation (absorption) or decomposition (desorption) since the hydride phase is less dense than the surrounding metal matrix it replaces during formation. Obviously, the opposite is true during decomposition. The accommodation process is frequently associated with the irreversible processes of plastic deformation and dislocation generation primarily in the hydride formation stage (absorption) when the hydride nucleus is growing in the hydrogen supersaturated metallic matrix, but this can also be true during hydride decomposition/desorption stage. Frequently, it is believed that hydride decomposition/desorption pressure is well defined, whereas formation/ absorption pressure is not. PCT data obtained during desorption are regarded as the best approximation to equilibrium between stable phases [70]. In 1993, Balasubramanian [73] expanded the elasto-plastic accommodation theory, outlined earlier by Qian and Northwood [71]. The major tenets of the elasto-plastic theory are such that elastic as well as plastic energies for both hydride formation and decomposition contribute to pressure hysteresis; the accommodation energy for hydride formation is, in general, not equal to that of hydride decomposition; furthermore, depending on whether the hydriding material is ductile or brittle either elastic or plastic accommodation terms will prevail in the hysteresis; and finally, for powder material in which the new phase is formed near the free surface or when it grows to a size comparable to the size of a powder particle, no plastic deformation occurs and the matrix remains in a purely elastic state. In 1995, Schwarz and Khachaturyan [74] took a slightly different approach and assigned the PCT absorption–desorption hysteresis to the coherency strain, which in open systems such as hydrides in equilibrium with a hydrogen gas reservoir at fixed pressure p and temperature T completely eliminates the two-phase equilibrium and as a result produces a large thermodynamically reversible hysteresis effect due to the presence of persistent coherency

144

2

Simple Metal and Intermetallic Hydrides

strain. More recently, in 2003, Rabkin and Skripnyuk [75] found that the Mg powder ball-milled for 8 h clearly exhibited a pressure hysteresis at 300°C, while no hysteresis was observed for the powder milled for 15 h. They argued that although the models of pressure hysteresis based on elastic, plastic, or elasto-plastic irreversible energy losses can be reasonable at low temperatures, they cannot be applied to the observed case of milled Mg hydrided/dehydrided at 300° because at such a high temperature a full recovery of deformed structure and the relaxation of elastic stresses should occur within a short time. Also, the model of Schwarz and Khachaturyan [74] is only relevant for the metal hydride transformations with a small relative volume change of less than a few percent. However, in the Mg–MgH2 transformation the volume change is about 20% and the loss of coherency should occur at a very early stage of transformation, which immediately removes the constraints required by the Schwarz and Khachaturyan theory. They proposed a new semiquantitative model which considers the fact that in a metallic powder during hydrogenation/absorption stage there exists a mixture of particles that are fully transformed (hydride) and nontransformed (supersaturated solid solution). The overall picture is of a pseudo-two-phase equilibrium during hydrogenation of metallic powders: a given transformation fraction does not mean that this is the fraction of hydride inside each individual metallic particle, but instead the fraction of fully transformed particles intermixed with the metallic particles where the transformation has not started at all. It follows from their model that the absorption PCT curve should exhibit a slope, while under assumption that elastic stresses are insignificant during the nucleation of the metallic solid solution from the hydride during desorption, the desorption curve should follow the true thermodynamic equilibrium and exhibit no slope in the two-phase region. The model also predicts that nanostructured powders produced by ball milling should not exhibit any hysteresis once their average particle size is smaller than the critical size of hydride nucleus (i.e., at which there is a sudden coherency loss with the matrix). Accordingly, in this case the whole particle can be transformed into hydride by thermal fluctuation, which eliminates the need for elastic accommodation of the expanding nucleus inside metallic matrix. The model of Rabkin and Skripnyuk [75] seems to be quite attractive in its explanation of pressure hysteresis observed in Fig. 2.43a. Our results always show a log-normal frequency distribution of powder particles (Fig. 2.34), which points towards the possibility of the existence of large and small fractions of particle sizes. This, in turn, is compatible with the model which states that fully transformed and nontransformed particles can coexist depending on their sizes. As shown in Fig. 2.43b, the enthalpy of absorption and desorption calculated from the Van’t Hoff plots using the mid-plateau pressures of PCT curves in Fig. 2.43a, which are listed in Table 2.18, is equal to −72 and 83 kJ/mol, respectively. The value of entropy is 138 and 151 J/mol K for absorption and desorption, respectively. The enthalpy value for absorption is very close to the values found in the literature for MgH2 as discussed in Sect. 2.1.2 and 2.1.3. Surprisingly, however, the enthalpy of desorption at 83 kJ/mol is much greater than the former and also greater than the enthalpy of desorption of the as-received and activated MgH2 as shown in Fig. 2.11. The coefficients of fit are excellent and give good credibility to the obtained values.

2.1

Mg/MgH2

145

The cause of greater enthalpy of desorption than that of absorption for the synthesized and desorption/absorption-cycled MgH2 is not well understood. The γ-MgH2 originally present after reactive milling synthesis was eliminated from the microstructure by desorbing the reactively synthesized powder at 350°C under pre-vacuum for 60 min, so this factor could not affect the results. Particle size is stable during any further thermal processing of ball-milled MgH2 without any sintering that could affect the average ECD (Figs. 2.17, 2.19, 2.20, and 2.21), so this factor can also be eliminated. Another factor, the grain size of MgH2, usually grows rather quickly in milled powders to about 60–80 nm range after desorption at 350–400°C (Table 2.13). Since the synthesized powder was desorbed at 350°C under pre-vacuum for 60 min, its grain size should be in this range. During further thermal cycling, the grain size becomes rather stable without measurable growth (Table 2.9, Figs. 2.14, 2.18, and 2.21). Also, this factor does not seem to be responsible for the greater desorption energy. However, in general, as has been shown in the preceding sections, the kinetics of the desorption process of ball-milled MgH2 powders are always more sluggish than those of absorption. Apparently, in view of our results, the desorption process is also more difficult thermodynamically, having greater enthalpy than its absorption counterpart. For the sake of clarity, it must be mentioned that Huot et al. [24, 35] reported that at 350°C the absorption/desorption PCT plateau pressure hysteresis of the unmilled MgH2 was quite substantial, while the hysteresis of the same material milled for 20 h was very small. They argued that the plateau pressure difference observed for the unmilled material was due to a very slow desorption kinetics which did not allow reaching equilibrium. In turn, ball milling increased desorption kinetics, which allowed reaching equilibrium and eliminated hysteresis. However, it must be noticed that the situation reported by Huot et al. is completely opposite to what is observed in Fig. 2.43, in which MgH2 synthesized by reactive mechanical milling shows a pressure hysteresis. In the context of the Rabkin and Skripnyuk model, which requires that desorption is better approximation of equilibrium then absorption, the high value of the enthalpy of desorption observed for a reactively synthesized MgH2 in Fig. 2.43b seems to be correct. If so, the hysteresis in Fig. 2.43a is due to higher enthalpy of desorption rather than to an inhomogeneity in the fraction transformed and nontransformed particles during the absorption process. A limited study of the desorption kinetics of reactively synthesized MgH2 was also carried out. The hydride was synthesized by reactive milling of Mg powder (Fig. 2.34a) in the magneto-mill Uni-Ball-Mill 5 and subsequently stored for 4 months in a sealed container under high-purity argon and removed only once for the desorption test. Figure 2.44 shows the kinetic curves and the corresponding Arrhenius plot for calculating the activation energy of desorption, which is ~140 kJ/mol. This value is close to those obtained for milled commercial MgH2 powders in Figs. 2.23 and 2.24. It is interesting to note that a relatively long and continuous storage of the synthesized powder under high-purity argon without its removal from a sealed container did not deteriorate the hydrogen desorption properties of the powder. Apparently, this type of continuous storage under a protective gas does not lead to aging of the powder as will be discussed in the following section.

146

2

Simple Metal and Intermetallic Hydrides

Hydrogen desorbed [wt.%]

7 Desorption EA~140 kJ/mol

6

4 5

3

4 2 3

1-300ⴗC 2-325ⴗC 3-350ⴗC 4-375ⴗC

2 1 1 0

0

1000

ln k

a

b

0 −1 −2 −3 −4 −5 −6 −7 −8 −9 0.00018

2000 Time [s]

3000

y = −139523x + 21.114 R2 = 0.9821

0.00019

0.0002

0.00021

0.00022

1/RT

Fig. 2.44 (a) Desorption kinetics curves at various temperatures under an initial hydrogen pressure of 0.1 MPa of MgH2 powder reactively synthesized for 100 h under the HES mode with one magnet in H2 at 850 kPa (ball-to-powder-weight ratio = 40; four balls). (b) The Arrhenius plot of the desorption rate for the estimation of the apparent activation energy, EA, using kinetics data for four temperatures: 300, 325, 350, and 375°C (EA ~140 kJ/mol). Coefficient of fit R2 = 0.982. Desorption tests were carried out after continuous storage in high-purity argon for 4 months

2.1.5

Aging Effects in Stored MgH2 Powders

In this book, by “aging” of MgH2 we understand a long-term storage of a MgH2 powder that has been subjected to a sequential removal from the storage container and exposed for longer or shorter periods of time to air or air-contaminated protective gas, which eventually leads to the formation of Mg(OH)2 (hydroxide) or MgO. As discussed in Sect. 2.1.2, formation of MgO/Mg(OH)2 layers on the surface of Mg/MgH2 powder is quite detrimental, as it affects its absorption/desorption properties [22]. In the already briefly mentioned paper, Friedrichs et al. [26] carried out very elegant HRTEM (high-resolution transmission electron microscopy) and XPS (X-ray photoelectron spectroscopy) studies of the initial stages of the formation of MgO/ Mg(OH)x on the ball-milled nanocrystalline MgH2 (BM-MgH2) and Mg (BM-Mg) powders, the latter obtained after hydrogen desorption of the former with particle

2.1

Mg/MgH2

147

sizes in the range 0.1–1 μm (having crystallites/grains ~10–20 nm in size), and additionally, on nanosized single-crystalline Mg produced by the gas-phase condensation technique (GPC-Mg)(average particle size ~40 nm). The analysis of BM-MgH2 under TEM was very difficult because it started releasing hydrogen as soon as the electrons interacted with the sample, resulting in a full decomposition to Mg in a few seconds. However, on the BM-and GPC-Mg they observed a thin passivation layer of MgO (possibly mixed with Mg(OH)x) with a thickness in the range of 3–4 nm, which prevented further oxidation. They concluded that when the handling of these samples is carefully done under inert-gas atmosphere (argon), vacuum, or toluene, the BM-and GPC-Mg are not oxidized any further. The presence of an oxide/hydroxide passivation layer on the BM-MgH2 was established by XPS, and it was found to be even thinner than in the case of nanocrystalline BM- and GPC-Mg samples. The authors concluded that the BM-MgH2 sample was the most resistant to oxidation. Exposure to air produced for all materials an increase of the thickness of the surface layer. On the BM-MgH2 sample, an amorphous hydroxide layer Mg(OH)x was formed. As also mentioned in Sect. 2.1.2, Varin et al. [27] showed that a long-term air exposure of nanocrystalline MgH2 for a few months led to a massive formation of crystalline Mg(OH)2 not only on the surface but also in the bulk by conversion of the entire MgH2 particles into Mg(OH)2. Apparently, the initial amorphous hydroxide layer Mg(OH)x grows and transforms into a crystalline-phase Mg(OH)2 (the reaction of (2.3)). Another problem arises when the sealed container containing argon gas with the stored MgH2 powder is frequently opened a number of times, even under protective gas, in order to remove powder samples for various purposes. We investigated the desorption behavior of the as-received ABCR powder that was stored almost a year under protective high-purity argon gas in a sealed container. However, the container was opened a large number of times in a plastic glove bag filled up with high purity argon. Even at the first look, the kinetics of desorption in Fig. 2.45a seem to be very sluggish when compared to the kinetics of the same fresh ABCR powder in Fig. 2.6a. The activation energy calculated from the Arrhenius plot in Fig. 2.45b yields ~217 kJ/ mol. This value is almost 50% larger than that obtained for the fresh powder (~168 kJ/mol; Fig. 2.6b). The activation energy value for the “aged” ABCR powder is close to the range of the activation energies reported by Jensen et al. [22], who intentionally exposed the samples to air as mentioned in Sect. 2.1.2. Apparently, even if MgH2 powder is not necessarily directly exposed to air but just exposed for a number of times to contaminated argon gas, as can happen in a glove bag, its hydrogen desorption properties can steadily deteriorate owing to the formation of a layer of Mg(OH)2 at the surface of the particles.

2.1.6

Other Methods of Synthesis of Nanostructured MgH2 than Ball Milling

Two methods of producing nanostructured MgH2 were already briefly mentioned in Sect. 2.1.3.3 and 2.1.5. Shao et al. [44] used the hydrogen plasma–metal reaction to obtain ultrafine magnesium with an average size of about 300 nm. The activation of

148

2

Simple Metal and Intermetallic Hydrides

Hydrogen desorbed [wt.%]

8.00 7.00 3

6.00

Desorption EA ~217kJ/mol

5.00 4.00 2

1

3.00 1-350ⴗC 2-375ⴗC 3-400ⴗC

2.00 1.00 0.00

0

1000

a

2000

3000

4000

Time [s]

−4 −4.5

y = -217170x + 34.198 R2 = 0.9894

−5

ln k

−5.5 −6 −6.5 −7 −7.5 −8 −8.5 0.000175

b

0.00018

0.000185 1/RT

0.00019

0.000195

Fig. 2.45 (a) Desorption curves at various temperatures under an initial hydrogen pressure of 0.1 MPa for MgH2 ABCR commercial powder which was stored under high-purity argon for almost a year but whose container was frequently opened in a plastic glove bag filled with argon. (b) The Arrhenius plot of the desorption rate for the estimate of the apparent activation energy, EA, using kinetics data for three temperatures: 350, 375, and 400°C (EA ~217 kJ/molH2). Coefficient of fit R2 = 0.989

such produced hydride involves only one cycle but its absorption kinetics is not much faster than the absorption of milled, activated, and cycled ABCR MgH2 powder. Friedrichs et al. [26, 45] used a “two-step” method for the preparation of nanocrystalline MgH2. Instead of ball milling for the refinement of Mg, they used the inert gas condensation technique developed a number of years ago by Granqvist and Buhrman [76] and subsequently modified and used by Gleiter and his coworkers in their pioneering work on the early synthesis of nanomaterials [77, 78]. The Mg nanoparticles were nearly spherical and single crystalline with the size range of 30–50 nm and an average of 35–40 nm. Hydrogenation of these particles was carried out in the same unit without air exposure (in situ hydrogenation). After hydrogenation,

2.1

Mg/MgH2

149

Fig. 2.46 SEM image of a net of nanowhiskers/nanofibers of MgH2 obtained after the hydrogeninduced disproportionation of Mg24Y5 at 20 bar and 370°C [79]

β-MgH2 had a particle size of 70–150 nm with an average of 109 nm. No absorption/ desorption tests were carried out. Zlotea et al. [79] used a method of disproportionation of Mg24Y5 bulk compound, which was fabricated by heating elemental metals sealed in a tantalum (Ta) tube filled with argon and heating it to 800°C in a high-frequency induction furnace under argon. Hydrogen absorption of the compound was carried out in an autoclave system at different hydrogen pressures and temperatures in the ranges of 1–30 bar and 280–380°C, respectively. With increasing hydrogen pressure at 370°C, the disproportionation reaction yielded at 5 bar metallic Mg + YH2, then at 10 bar Mg + MgH2 + YH2 + YH3, and finally at 20 bar a complete disproportionation reaction occurred yielding an equilibrium mixture of only two phases MgH2 + YH3. From this sequence of events, the authors concluded that the hydrogen absorption in Mg24Y5 occurs in a two-step mechanism. During the first step, only YH2 and pure metallic Mg are formed. The second step promotes the formation of YH3 and MgH2. The MgH2 phase synthesized in the second step has a very unusual, onedimensional structure of nanowhiskers or nanofibers forming a net as shown in Fig. 2.46. No absorption/desorption tests were carried out by the authors. Saita et al. [80] applied hydriding chemical vapor deposition (HCVD) for preparing MgH2. They used commercial Mg, which was heated to ~600°C and vaporized in a hydrogen atmosphere at a pressure of 4 MPa. The reaction product was deposited on a cooled Inconel substrate and subsequently collected for further investigation. Quite remarkably, the obtained morphology was nanofibrous, as shown in Fig. 2.47, and is very similar to the one fabricated by Zlotea et al. [79]. Each fiber was less than 1 μm in diameter and 10 or more micrometers in length.

150

2

Simple Metal and Intermetallic Hydrides

Fig. 2.47 SEM image of a nanofibrous structure of MgH2 synthesized by HCVD [80]

Saita et al. [80] investigated hydrogen storage properties of the nanofibrous MgH2 by carrying out PCT as well as absorption/desorption kinetics tests. In a PCT test at 305°C, the product reversibly absorbed and desorbed 7.6wt.%H2 without any elaborate activation treatment except preannealing at 305°C for 2 h in vacuum. The plateau pressure in PCT absorption/desorption was slightly above 1 MPa for absorption and ~0.5–0.6 MPa for desorption. At 287 and 270°C, the material was able to desorb only if the pressure decreased much below 0.1 MPa (1 atm.). Calculation of the thermodynamic parameters from the Van’t Hoff plot yielded ΔH ~73.4 kJ/mol and ΔS ~ 129 J/molK. During absorption test at 340°C, the nanofibrous MgH2 absorbed under 4 MPa H2 about 6wt.% in 600 s. During desorption test at 340°C, it desorbed the same quantity of H2 under 0.6 MPa in 300 s. By comparing these numbers with those obtained for the ball-milled MgH2 and presented in the previous sections, it must be concluded that the nanofibrous MgH2 does seem to be superior as far as its absorption/desorption properties are concerned. Li et al. [81] reported the synthesis of metallic Mg nanowires by a vaportransport approach using commercial Mg powders. The nanowires were deposited on stainless steel mesh, which were inserted into a stainless steel tube in a furnace at the evaporation temperature of 930°C and deposition temperature of 300°C. Three types of nanowires with various morphologies were produced depending on the flow rate of argon gas used in the synthesis. Type 1 nanowires had a uniform diameter of 30–50 nm, and type 2 had 80–100 nm diameter. Type 3 were rod-like with a diameter range of 150–170 nm. XRD studies showed the presence of β-MgH2 as the principal phase, mixed with some Mg in type 3. The hydrogen absorption and desorption kinetics depended on the type of nanowires. The fastest kinetics were observed for type 1, which could achieve absorption of about 2 wt.% within 2,000 s in the hydrogen pressure range of 0.4–2.0 MPa even at such a low

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151

temperature as 100°C. At 300°C, type 1 absorbed about 7 wt.%H2 after about 500 s. Desorption experiments were conducted at hydrogen pressures ranging from 0.02 to 0.6 MPa, and type 1 desorbed about 3 wt.% at 200°C and about 7 wt%. at 300°C % after about 900 s, although the authors did not specify under what hydrogen pressure the desorption process had occurred. Most probably it was vacuum (0.02 MPa). The activation energies of hydriding/dehydriding for nanowires/rods of type 1, 2, and 3 were found to be 33.5/38.8, 38.7/46.5, and 70.3/81.1 kJ/mol, respectively. These values seem to be indeed much lower than any other quoted in this chapter for ball-milled MgH2. The Van’t Hoff plot yielded the enthalpy of desorption for type 1, 2, and nanowires as 65.3, 65.9, and 67.2 kJ/mol, respectively, which can be compared to the standard range of ~70–74 kJ/mol for MgH2. The authors also found that the major disadvantage of nanowires was their collapse into nanoparticles after a relatively short cycling of no more than 10 cycles. However, after 50 cycles, the collapsed nanowires, i.e., nanoparticles, did not show any decrease in capacity. Finally, it must be noted that all the methods of fabricating nanostructured MgH2 described and discussed above do not yield the quantities of product that could be commercially viable and as such will probably remain a scientific curiosity.

2.2

MgH2 with Catalytic Additives

As shown in the preceding sections, nanostructuring of MgH2, either by ball milling or other more exotic processing techniques, has been unable to reduce the desorption temperature of MgH2 much below 300°C as it was originally hoped for when ball milling was introduced for nanostructuring of solid-state hydrides (Sect. 1.3). Therefore, the researchers worldwide started looking for catalytic additives to MgH2, which when combined with a nanostructure could further improve the hydrogen storage properties of this particular hydride. The catalytic additives that have been investigated in MgH2 can be roughly divided into three important groups: metals/intermetallics, oxides, and chemical compounds that include hydrides. Their effect on the desorption properties of MgH2 will be discussed in the following sections. We are not going to discuss absorption since, as mentioned before, absorption is usually much easier than desorption. For the sake of clarity, we have divided each of the following sections into two parts depending on whether the hydride was desorbed in vacuum or at atmospheric pressure of hydrogen (1 atm. or 0.1 MPa). This differentiation is very important from the technological view point. It is quite unfortunate that, so far, most of the desorption tests of MgH2 with additives have been conducted in vacuum. It must be stressed that the results obtained in such a way have essentially a very limited value. From the Van’t Hoff data in Figs. 2.11 and 2.43, one can estimate using (1.34) in Sect. 1.4.1.1 that for the pure MgH2 the equilibrium temperature at atmospheric pressure of hydrogen is within the range of 250–280°C. It simply means that, desorption at atmospheric pressure and at a temperature lower than its equilibrium

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temperature at atmospheric pressure of hydrogen is thermodynamically impossible. However, in vacuum, material can desorb at much lower temperatures and with better kinetics. The great pitfalls of conducting desorption experiments in vacuum are well illustrated by the results obtained by Song et al. [82, 83], who showed an increase of desorbed hydrogen with decreasing desorption pressure. Also, from the results of Li et al. [84] one can clearly see a substantial increase in desorption kinetics at 300°C under the pressure of 0.02 MPa as compared to the results obtained under 0.1 MPa. The results from both groups will be discussed in more detail in the following sections. Obviously, in practical situations one can hardly imagine a vacuum pump installed onboard of a fuel cell–powered vehicle. Even if so, the membrane of a Proton Exchange Membrane Fuel Cell (PEMFC) will be soon contaminated by the oil vapors released from the pump (dry pumps are possible but this would enormously complicate the entire design).

2.2.1

Mg/MgH2–Metals and Intermetallics

2.2.1.1

Desorption in Vacuum

A number of researchers investigated the effects of various metallic additives on the desorption properties of nanostructured, ball-milled MgH2 in vacuum. Liang et al. [85, 86], Dehouche et al. [87], Bouaricha et al. [88], Huot et al. [89], Shang et al. [90], and Au [91] investigated the effects of the addition of Ti, V, Nb, Mn, Fe, Al, Cu, La, Ni, and Pd to commercial MgH2 or Mg, which was after milling separately hydrogenated. The effects of the addition of intermetallics were investigated by Bobet et al. (YNi) [92], Hu et al. (TiMn1.5 and Ti37.5V25Cr37.5) [93, 94], Tran et al. (Mischmetal) [95], Skripnyuk et al. (Mg2Ni eutectic) [96], Yonkeu et al. (TiV1.1Mn0.9) [97], Fu et al. (LaNi5) [98], and Sai Raman et al. (Ce-free MischmetalNi5-nonmilled; synthesized by encapsulation method) [99]. For some of the metallic additives, desorption at a low 200–235°C temperature range [85, 86, 91] was reported and for some intermetallic additives the range of desorption temperatures was as low as 250–270°C (Hu et al. (TiMn1.5 and Ti37.5V25Cr37.5)[93, 94]) and ~250°C for the addition of LaNi5 by Fu et al. [98]) and Liang et al. [100]. It is interesting to note that Liang et al. [100] observed that the milled Mg–30 wt.% LaNi5 mixture transformed into MgH2 + LaH3 + Mg2NiH4 upon hydrogenation and a direct mechanical milling of MgH2–30 wt.% LaNi5 under argon led to the formation of the mixture of MgH2 + LaH3 + Mg2NiH4 + Mg. It means that the intermetallic LaNi5 is an unstable additive, which reacts with the other phases present in a mixture either upon hydrogenation or mechanical milling. Liang et al. concluded that LaNi5 does not have any direct catalytic effect on MgH2. Instead, LaH3 has a strong catalytic effect on absorption of Mg, but weak effects on desorption. Mg2Ni has better catalytic effect than LaH3 at temperatures above 100°C.

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153

Some researchers investigated multiphase additives. Kojima et al. [101] ballmilled MgH2 with a nano-Ni/Al2O3/C composite catalyst. The mixture decomposed at a really low temperature of ~200°C but in vacuum. A complex composite catalyst BCN/Ni/Pd/SWNT (where BCN is barium-calcium niobium high-temperature proton conducting compound and SWNT is single-wall-nanotubes) was also used by Yoo et al. [102] high-temperature proton conducting for ball-milled MgH2. It desorbed ~3 wt.%H2 in about 3,600 s at 230–250°C in vacuum. Instead of ball milling, Dufour and Huot [103] used cold rolling to produce Mg–Pd laminates, which after activation desorbed ~6 wt.%H2 at 300°C in 1,200 s under primary vacuum. The same authors produced a Mg6Pd compound by cold rolling and ball milling [104]. However, the desorption properties were rather mediocre since the compound desorbed, at 350°C in vacuum (0.01 MPa), only about 2.8 wt.% after 2,000 min. In conclusion, as mentioned at the beginning of this section, desorption of MgH2 in the range of 250–280°C with a good rate in vacuum is nothing unusual.

2.2.1.2

Desorption at Atmospheric Pressure of Hydrogen

Surprisingly, a rather small number of researchers tested various supposedly catalytic additives to MgH2, employing desorption at atmospheric pressure of hydrogen (0.1 MPa). Grigorova et al. [105] used a reactive mechanical milling for producing nanostructured MgH2 with Mg2Ni and Mg2Ni1–xMx (M = Fe and Co) additives. Desorption tests were conducted at 300°C under 0.15 MPa hydrogen pressure. The mixtures desorbed 5–6 wt.%H2 within 3,000–7,000 s. No desorption tests below 300°C were reported. Bobet et al. [92, 106] investigated the effect of Co, but the material desorbed only at 350°C. Yu et al. [107] also used reactive milling for producing MgH2 with Ni, Cu, and CrCl3 additives. The mixture desorbed a moderate 5 wt.% at 300°C within 2,400 s, but no desorption below 300°C was reported. Li et al. [108] combined 20 wt.%Ni and 1 wt.%TiO2 as a complex additive, which was reactively milled with Mg. Reported desorption at 305°C was very slow. Instead of ball milling, Løken et al. [109] resorted to equal channel angular pressing (ECAP) to process heavily deformed Mg-20 wt.%Ni-8 wt.%Mm (Mischmetal) alloy, which was subsequently activated, hydrogenated, and dehydrogenated. No desorption below 325°C at 0.1 MPa hydrogen was observed. Apparently, this kind of processing is not any better than ball milling, or may even be worse. Czujko et al. [110] investigated ball-milled Mg with 10 wt.%V, Zr, and Y additives, which was activated, separately hydrogenated, and then desorbed. At 300°C under 0.1 MPa hydrogen pressure, the V-doped material desorbed ~5 wt.%H2 within 3,600 s but the one with Zr and Y only 2.8 wt.%H2. Li et al. [111] added 20 wt.% YNi intermetallic to Mg and milled in 3 MPa of hydrogen pressure. The mixture absorbed at room temperature about 4 wt.%H2 after about 1,000–1,200 s; at 100°C it absorbed 4.5 wt.%H2 after about 100 s; and at 200°C it absorbed 5.5 wt.%H2 after about 200 s. The absorption process was indeed extremely fast. However, the mixture did not desorb much below 300°C at 0.1 MPa. At 300°C it desorbed ~5 wt.%H2 after 1,800 s, which indicated a rather modest desorption rate. It

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Simple Metal and Intermetallic Hydrides

is interesting to point out that the same mixture could desorb almost 1 wt.%H2 more, much faster, when the desorption pressure was reduced to 0.02 MPa at 300°C. Desorption at 250°C at 0.02 MPa yielded ~5 wt.%H2 after 3,000 s. This is another good example that shows how reduction of desorption pressure below 0.1 MPa eliminates the thermodynamic barrier for MgH2 desorption. The most interesting studies are those in which desorption tests were conducted at the atmospheric pressure of hydrogen and the catalyzed MgH2 was able to desorb a relatively large quantity of hydrogen at the temperature equal to or even lower than the equilibrium temperature at atmospheric pressure of hydrogen, i.e., below 250–280°C. Some interesting results were reported by Vijay et al. [112]. They synthesized by reactive mechanical alloying/milling (RMA/RMM) the Mg + 5 wt.%FeTi and 30 wt.%FeTi composites, which were subsequently hydrogenated to form the MgH2 matrix and the MgH2 + 40 wt.%FeTiMn composite by mechanical milling. The latter composite after ball milling could absorb 4 wt.%H2 at 80°C and desorb ~3.8 wt.%H2 at 300°C within 800 s and ~2 wt.%H2 at 240°C within 4,200 s under 0.3 MPa of hydrogen pressure. The temperature of 240°C is much lower than the equilibrium temperature at atmospheric pressure of hydrogen, which may be evidence for the lowering enthalpy of MgH2 by the addition of 40 wt.%FeTiMn. The beneficial mechanism of FeTiMn intermetallic additive was not explained, but it may have something to do with the fact that the FeTi intermetallic forms a lowtemperature hydride (Table 1.3 in Sect. 1.1) and Mn might somehow additionally improve its performance. A certain disadvantage of this type of approach is the large quantity of additive intermetallic required. Nevertheless, this composite system requires attention and more research. Wang et al. [113] produced a Mg–10.9 wt.%Ce alloy by induction melting, which after pulverizing and reactive mechanical milling in hydrogen was mixed with nano-Ni (average particle size ~10 nm) and additionally milled for 50 h under argon. The as-cast structure contained the Mg matrix with the precipitates of CeMg12 intermetallic, which after hydrogenation converted to the CeH2.53 hydride. The authors reported that this composite was able to absorb ~2.9 wt.%H2 at 120°C within ~1,800 s. Under 0.1 MPa of hydrogen pressure, the composite desorbed ~2 wt.%H2 at 180°C within ~1,200 s. The reported results look very interesting, although certain discrepancies exist in the presented data. For example, the authors showed a PCT curve for the Mg–Ce/nano-Ni composite that evidently exhibits a pressure of 1 atm. hydrogen at 280°C. With such an equilibrium temperature at 1 atm., the composite should not have been able to desorb at 180°C. Also, they claimed that ball-milled MgH2 began desorbing at 200°C, which is very unlikely. Nevertheless, the investigation of this kind of a composite should be repeated to check whether the reported results are indeed reproducible. The same group led by Wang [114] also investigated hydrogen storage properties of a composite Mg/ Mg2Ni0.8Cr0.2 containing nano-TiO2 (average particle size ~40 nm). The material containing 20 and 50 wt.% Mg2Ni0.8Cr0.2 (plus TiO2) desorbed 4.2 and 3.2 wt.%H2 at 240°C under 0.1 MPa pressure within 2,900 and 900 s, respectively. Very recently, Bystrzycki et al. [115] studied the possibility of destabilization of MgH2 by chemical reaction with Si and formation of an intermediate intermetallic

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155

as proposed by Vajo et al. [116](Vajo et al. approach will be discussed in more detail in Sect. 3.3). The commercial MgH2 and Si powder mixture corresponding to the stoichiometry of Mg2Si was ball-milled to obtain a nanocrystalline composite structure. The sluggish destabilization of MgH2 by solid-state reaction with Si forming the Mg2Si intermetallic compound was observed at 250°C. It was confirmed that the Mg2Si compound was formed after the dehydrogenation of the synthesized MgH2–Si mixture. The well-known catalytic additive to MgH2 is Ni. So far, it has been added as a coarse powder with the size range of a few tens of micrometers. A short review of these early experiments is presented in [117]. More recently, Hanada et al. [118] investigated the addition of nano-Ni to ball-milled MgH2. Unfortunately, he used continuous desorption under He gas and the material desorbed ~6 wt.%H2 at 160°C in 15,000 s. In our laboratory, we focused our efforts on the catalyzing of MgH2 with nano-Ni powders produced as experimental batches by Inco Specialty Services, Mississauga, Ontario. The effects of micrometer-sized nickel (m-Ni)(commercially produced by Inco Type 255 Ni) having specific surface area (SSA) = 0.7 m2/g, submicrometric Ni having SSA = 7.5 m2/g, and experimental batch of nano-Ni (n-Ni) having SSA= 30.3 m2/g on the rate of synthesis of MgH2 and its hydrogen desorption properties have been reported in [117]. It has been found that both micrometric and submicrometric nickels greatly improve the rate of hydrogen absorption during controlled reactive mechanical milling (CRMM) and conversion of Mg into the MgH2 hydride, and, subsequently, improve the hydrogen desorption properties of synthesized catalyzed MgH2. However, by comparison to the previous two, n-Ni is the most potent catalyst for the conversion of Mg into MgH2 during CRMM under hydrogen. Up to ~25 h of CRMM, the addition of barely 0.5 wt.% of n-Ni increases about twofold the rate of hydrogen absorption as compared to undoped Mg. Furthermore, the addition of 2 wt.% of n-Ni results in even faster absorption of hydrogen by Mg, resulting in ~6 wt.%H2 absorbed after ~15 h of CRMM. The hydrogen desorption kinetics at the engineering conditions of 0.1 MPa hydrogen pressure and no initial activation by cycling becomes very fast, which is reflected in the reduction of the activation energy for desorption by ~60 kJ/mol as compared to the reference synthesized MgH2. The effects of n-Ni and nano-oxide additives (Al2O3 and Y2O3) on the hydrogen storage properties have been reported in [119]. The addition of both oxides has a limited effect on improving the hydrogen storage properties. In contrast, the addition of specialty Inco m-and n-Ni substantially reduces hydrogen desorption temperatures, which is also accompanied by very fast desorption kinetics under 0.1 MPa H2 pressure. The activation energy of desorption is also substantially reduced. Prolonged milling for 100 h is detrimental for the hydrogen storage properties of the m- and n-Ni-doped MgH2. It is to be pointed out that a simple mechanical mixing of either as-received or premilled for 20 h Tego Magnan MgH2 powder with both mand n-Ni additives in a glass vial rotating in a steel milling cylinder of the magnetomill Uni-Ball-Mill 5 for 1 h, without using steel balls, does not affect the desorption properties. Figure 2.48 shows the DSC traces of the MgH2 as-received and mixed with the Inco m- and n-Ni (SSA= 30.3 m2/g) in a glass vial. The endothermic hydro-

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DSC / mW/mg [3] 415.4⬚C [2] 420.8⬚C

↓ exo 10

[1] 414.6⬚C

Mixing [1] —— MgH2 as received [2] – – – m-Ni [3] ------- n-Ni

8 6 4 2 0 250

300

a

350 400 Temperature / ⬚C

450

500

DSC / mW/mg 4.0

[1] 379.4⬚C

↓ exo

3.5

[2] 363.8⬚C [3] 372.7⬚C

Mixing [1] —— MgH2 after pre-milling (20h) [2] – – – m-Ni [3] ------- n-Ni

3.0 2.5 2.0 1.5 1.0 0.5 0 250

b

300

350 400 Temperature / ⬚C

450

500

Fig. 2.48 DSC traces of Tego Magnan MgH2 powder doped with m- and n-Ni, and (a) mixed for 1 h without milling (no steel balls) and (b) mixed for 1 h with Tego Magnan MgH2 which was premilled for 20 h (heating rate 4°C/min; argon flow rate 25 ml/min)

gen desorption peaks of MgH2 with Ni additives are not shifted with respect to the peaks of the as-received powder. The situation is changed dramatically when the powder containing the Ni additive is ball-milled even for a relatively short time. Very recently, we have investigated the effect of milling time in the magneto-mill Uni-Ball-Mill 5 on the hydrogen storage properties of ABCR MgH2 powder doped with micrometric and nanometric Ni [120]. Figure 2.49 shows the SEM micrographs of micrometric (m-Ni) (Type 255 produced by Inco) and nanometric Ni (n-Ni) used in this study. Micro-Ni has a very

2.2

MgH2with Catalytic Additives

157

Micro-Ni Micro-Ni

a

b Nano-Ni ECD = 42 ± 16nm Nano-Ni ECD = 42±16nm

c

d

Fig. 2.49 SEM micrographs (a,b,c) of micrometric Inco Ni (Type 255) having SSA = 0.7 m2/g, taken at various magnifications to reveal its morphology, and (d) nano-Ni (n-Ni) having SSA = 14.5 m2/g

unusual and complex shape. At relatively low magnification (5,000x), it exhibits a filamentary shape (Fig. 2.49a). However, at the magnification of 25,000x it is clearly seen that each filament is composed of flower-like corollas (Fig. 2.49b) joined closely together, each one resembling a rose. At very high magnification (50,000x) each rose-like corolla consists of small petal-like features whose thickness is on the order of or less than ~100 nm (Fig. 2.49c). Although in the macroscale the Type 255 Ni is micrometric in size, the thickness of the petal-like features falls within the nearly nanometric range. The Inco nano-Ni (n-Ni) has a filamentary shape but resembles a delicate “coral colony” and its dimensions are truly nano, i.e. below 100 nm (Fig. 2.49d). The measured mean diameter of a “coral filament” is 42 ± 16 nm. Figure 2.50 shows the microstructure of the powder after milling for 15 min with the addition of 5 wt.% Inco m-Ni and n-Ni. The distribution of m-Ni is not quite uniform (Fig. 2.50a) in contrast to that of n-Ni which shows a very uniform distribution of Ni particulate (bright specks) in the entire mass of the powder (Fig. 2.50b). The average ECD particle size of the powder milled with the m-Ni additive is slightly larger than that of the one milled with n-Ni, although both are below the 1 μm mark. The XRD pattern in Fig. 2.51a shows that the microstructure of both powders consists of hydride phases β- and γ-MgH2, Ni, and a small amount of retained Mg and MgO that is formed during XRD test from the retained Mg. Table 2.19 shows grain size of the phases in the ABCR MgH2 ball-milled with Inco m-

158

ECD = 0.93 ± 0.63 mm

a

2

Simple Metal and Intermetallic Hydrides

ECD = 0.74 ± 0.46 mm

b

Fig. 2.50 SEM micrographs of the ABCR powder after milling for 15 min (HES57 – two magnets mode; hydrogen 700 kPa) with the addition of 5 wt.% Inco (a) m-Ni and (b) n-Ni. Average ECD particle size with standard deviation is shown in the inset

and n-Ni additives and in the same powders cycled as will be discussed later. After milling for 15 min using HES57 – two magnets mode under 700 kPa hydrogen, the grain size of β-MgH2 and the Mg phase is on the order of ~40 nm. Nano-Ni retains its as-received grain size on the order of ~20–25 nm. The effect of a ball-milling time on the DSC desorption behavior of undoped and Ni-doped MgH2 is shown in Fig. 2.52a. Endothermic desorption peak for the MgH2 + 5 wt.% m-Ni powder milled for 15 min is only modestly shifted to lower temperatures, showing the onset temperature at ~350°C and the peak maximum at 392.8°C as compared to a pure MgH2 with the onset at ~380°C and the maximum at 418.2°C, respectively, also milled for 15 min. In contrast, the hydrogen desorption peak for the MgH2 + 5 wt.% n-Ni powder is substantially shifted to lower temperatures and shows the onset temperature at ~170°C and the peak maximum at 243.1°C. However, when the MgH2 + 5 wt.% m-Ni powder is milled for 20 h, its desorption properties are much improved such that the onset is at ~275°C and the peak maximum at 302.3°C (Fig. 2.52b). Apparently, longer milling time reduces desorption temperature, most likely due to a better dispersion of Ni particles within the MgH2 matrix. Nevertheless, the desorption temperature of the 20 h milled MgH2 + 5 wt.% m-Ni powder is still worse than that of the 15 min milled MgH2 + 5 wt.% n-Ni. This clearly shows that n-Ni exhibits superb catalytic properties. Figure 2.53 shows the desorption curves for the ball-milled powders. The mixture with n-Ni in Fig. 2.53a shows a very fast desorption in the range 300–350°C and moderate desorption at 275°C. No desorption has occurred at 200 and 250°C, which indicates that the basic thermodynamic property of the system (enthalpy) is not changed from that for pure MgH2. We calculated the apparent activation energy of desorption of the mixture with n-Ni using data points from two ranges of temperatures: 275–350°C and 300–350°C, and obtained 94 kJ/mol (R2 = 0.982) and 78 kJ/mol (R2 = 0.989), respectively. The mixture of Tego Magnan MgH2 with m-Ni ball-milled for 20 h in Fig. 2.53b shows slower desorption kinetics than that for the mixture with n-Ni milled for 15 min, which is also supported by greater apparent activation energies of 105 kJ/mol (R2 = 0.983) and 92 kJ/mol (R2 = 0.989) for the

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159

20000 b-MgH2 g-MgH2 Mg

16000

MgO Ni

Counts

12000

8000

MgH2+m-Ni

4000 MgH2+n-Ni 0 30

40

a

50 60 Degrees 2-Theta

70

80

90

20000 b-MgH2 g-MgH2 Mg

16000

MgO Ni Mg2NiH4

Counts

12000

Milled+cycled+EA tested

8000

Milled +cycled 4000

Milled 0 20

b

30

40 50 60 Degrees 2-Theta

70

80

90

Fig. 2.51 XRD patterns of ABCR powder after (a) milling for 15 min with the addition of 5 wt.% Inco m-Ni and n-Ni (HES57–two magnets mode; hydrogen 700 kPa), and (b) the ABCR MgH2 + n-Ni (SSA = 14.5 m2/g) mixture after milling and cycling as well as after testing for activation energy of desorption after cycling

275–375°C and 300–375°C range, respectively. Once again, the n-Ni clearly shows its superb catalytic properties. A cycling behavior of the MgH2 mixture with n-Ni was also studied. The cycling process consisted of five desorption/absorption cycles at 300°C. Desorption was carried out at the atmospheric pressure of hydrogen, and absorption was realized under 4.0 MPa pressure of hydrogen for 15 min. XRD after cycling showed formation of a

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Table 2.19 Grain size of the phases in the milled and cycled powder ABCR MgH2 containing Inco m- and n-Ni additive Grain size of Grain size Grain size Grain size of Sample MgH2 (nm) of Ni (nm) of Mg (nm) Mg2NiH4 (nm) ABCR + 5wt.% m-Ni 40.0 ± 2.9 25.5 ± 6.0 ABCR + 5wt.% n-Ni 34.9 ± 2.4 17.3 ± 1.3 90 ± 10 – ABCR + 5wt.% n-Nia 86 ± 9 – ABCR + 5wt.% n-Nib a Powder milled and cycled at 300°C b Powder milled and cycled at 300°C + test for EA

Exo

—— MgH2 ------ MgH2+m-Ni – – – MgH2+n-Ni

Heat Flow (W/g)

7

41.4 ± 9.6 36.0 ± 4.2 133 ± 17 132 ± 19

– – 27 ± 2 24 ± 2

418.2ⴗC 392.8ⴗC

243.1ⴗC

2

Milled for 15min −3 100

200

a

300 TemperatureⴗC

400

DSC/mW/mg 3.0 2.5

[1] 302.3⬚C

exo MgH2 + m-Ni milled for 20 h

2.0 1.5 TON~275⬚C

1.0 0.5 0 100

b

150

200

250

300 350 Temperature / ⬚C

400

450

500

Fig. 2.52 (a) DSC traces of ABCR MgH2 doped with 5 wt.% Inco m- and n-Ni (SSA = 14.5 m2/g) subjected to 15 min of ball milling (HES57 – two magnets mode; hydrogen 700kPa). (b) DSC desorption peak of Tego Magnan MgH2 + 5 wt.% Inco m-Ni after milling for 20 h (heating rate 4°C/min; flow rate of argon 25 ml/min)

2.2

MgH2with Catalytic Additives

161

Hydrogen desorbed [wt.%]

8.00 1-200⬚C 2-250⬚C 3-275⬚C 4-300⬚C 5-325⬚C 6-350⬚C

7.00

6

6.00

5

5.00

4

Ea = 94 kJ/mol (275÷350ⴗC) Ea = 78 kJ/mol (300÷350ⴗC)

3

4.00 3.00

MgH2 + n-Ni milled for 15 min

2.00 2

1.00

1

0.00 0

1000

2000

3000

4000

Time [s]

a

Hydrogen desorbed [wt.%]

8.00 1-275⬚C 2-300⬚C 3-325⬚C 4-350⬚C 5-375⬚C

7.00

5 6.00 4 3

5.00

2

Ea = 105 kJ/mol (275÷375ⴗC) Ea = 92 kJ/mol (300÷375ⴗC)

4.00 1

3.00 2.00

MgH2 + m-Ni milled for 20h

1.00 0.00 0

b

500

1000 1500 2000 2500 3000 3500 4000 Time [s]

Fig. 2.53 Desorption kinetic curves at various temperatures obtained in a Sieverts-type apparatus under 0.1 MPa of hydrogen pressure of (a) ABCR MgH2 + n-Ni (SSA = 14.5 m2/g) ball-milled for 15 min and (b) Tego Magnan MgH2 + m-Ni ball-milled for 20 h

small amount of Mg2NiH4 hydride (Fig. 2.51b). A profound grain growth of the constituent phases is observed upon cycling as shown in Table 2.19. Compared to the grain size of the phases directly after milling, the grain size of β-MgH2 (after the fifth absorption) increases 2–3-fold upon cycling and the grain size of Mg (after fifth desorption) increases almost fourfold. Desorption kinetics of the cycled mixture (after the fifth absorption) were again investigated, and the pertinent desorption curves are shown in Fig. 2.54. In the 275–350°C range the rate of desorption is slightly slower as compared to that in Fig. 2.53a, but this is not reflected in the apparent activation energy of desorption, EA, equal to 99 kJ/mol (R2 = 0.937) and 68 kJ/mol (R2 = 0.945) for the range 275–350°C and 300–350°C, respectively. Most importantly, one can see a substantial loss of final capacity after cycling at each desorption temperature as compared to the as-milled material in Fig. 2.53a. In general, this effect is very persistent after cycling of MgH2 hydride as discussed in Sect. 2.1.3, and, obviously, is unrelated to the presence of n-Ni in the mixture with MgH2.

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Hydrogen desorbed [wt.%]

7.00 EA = 99 kJ/mol (275-350ⴗC)

6.00

EA = 68 kJ/mol (300-350ⴗC) 325⬚C

5.00

350⬚C

4.00 300⬚C

275⬚C

3.00 2.00 1.00 0.00 0

1000

2000 Time [s]

3000

4000

Fig. 2.54 Desorption kinetic curves at various temperatures of the cycled ABCR MgH2 + n-Ni (SSA = 14.5 m2/g) mixture from which the activation energy EA is calculated

Table 2.20 Specific surface area (SSA), the content of carbon various experimental batches of Inco n-Ni Sample SSA (m2/g) Grain size (nm) C (wt.%) NI4 4.02 60 0.53 NI5 6.41 — 2.97 NI6 9.49 — 0.46 NI3 14.50 17 0.34 NI7 18.82 37 0.61 NI8 28.18 — 2.3 NI9 60.46 — 1.86 NI10 84.70 — 0.46

and oxygen and morphology of O (wt.%) 0.69 0.051 0.11 2.46 3.5 2.08 11.4 11.7

Morphology Spherical Spherical Spherical Filamentary Spherical Filamentary Filamentary Filamentary

The combined effect of the SSA and chemical composition of Inco n-Ni has also been investigated [121]. Table 2.20 lists the SSA, the content of oxygen and carbon, and morphology of various experimental batches of Inco n-Ni that have been studied as catalyst for MgH2. The SSA covers a wide range from about 4 to 85 m2/g. The ABCR MgH2 mixed with 5 wt.% n-Ni having increasing SSA was ball-milled under 700 kPa hydrogen for 15 min using the HES57 mode (two magnets). For comparison, pure MgH2 was also milled under the same conditions. Figure 2.55 shows the microstructure of the ABCR powder milled with the n-Ni having increasing SSA. It is seen that the spherical Ni particles are distributed in some localized places such as the contact areas between the MgH2 particles, while the filamentary n-Ni is more or less smeared off on the surface of the MgH2 particles although it does not form any continuous or semicontinuous film. The average ECD particle size is within the range 0.8–1.0 μm.

Filamentary - SSA = 28.18 m2/g

Spherical - SSA = 9.49 m2/g

MgH2with Catalytic Additives

Fig. 2.55 Scanning electron micrographs (SEM) in the back scattered electron (BSE) mode at 100,000× of the ABCR MgH2 + 5 wt.% n-Ni mixtures ballmilled powder under 700 kPa hydrogen with varying SSA (see the insets in the pictures)

Filamentary - SSA = 84.7 m2/g

Spherical - SSA = 18.82 m2/g

Filamentary - SSA = 14.5 m2/g

Filamentary - SSA = 60.46 m2/g

Spherical - SSA = 6.41 m2/g

Spherical - SSA = 4.02 m2/g

2.2 163

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XRD patterns in Fig. 2.56a,b show primarily the presence of β- and γ-MgH2 and some retained Mg. Bragg diffraction peaks of Ni can be recognized in all samples but their intensity gradually decreases with increasing SSA, which indicates that the n-Ni having a very large SSA is more intimately mixed with MgH2. Specifically, this is observed for the filamentary n-Ni in Fig. 2.56b. This observation correlates well with the SEM micrographs in Fig. 2.55. MgO peaks as usual are due to the oxidation of the retained Mg. No formation of Mg2NiH4 has occurred upon milling (Fig. 2.56), although there is a very high likelihood that it will be formed upon subsequent thermal cycling as evidenced in Fig. 2.51b. Grain size of β- and γ-MgH2 calculated from the broadening of XRD peaks is on the order of 20–30 nm with a minimal lattice strain (10−3–10−4). Grain size of the retained Mg is on the order of 40 nm. DSC tests show a substantial reduction of the hydrogen desorption onset (red circles) (Ton) and peak (Tpeak) temperatures due to the catalytic effects of n-Ni as compared to the hydrogen desorption from pure MgH2 also milled for 15 min. (Fig. 2.57). It is interesting to note that there is no measurable difference between spherical (Fig. 2.57a) and filamentary (Fig. 2.57b) n-Ni, although there seems to be some effect of SSA. We also conducted desorption tests in a Sieverts apparatus for each SSA and obtained kinetic curves (Fig. 2.58), from which the rate constant, k, in the JMAK equation was calculated. The enhancement of desorption rate by n-Ni is clearly seen. At the temperature of 275°C, which is close to the equilibrium at atmospheric pressure (0.1 MPa), all samples desorb from 4 to 5.5 wt.% H2 within 2,000 s. We have plotted in Fig. 2.59a the onset and peak temperatures from Fig. 2.57a,b as a function of SSA for each n-Ni additive. It is seen that a dramatic drop in the hydrogen desorption temperature of MgH2 occurs only up to a certain value of SSA of approximately ~15 m2/g. Greater values of SSA do not have better effect on the desorption temperature. The plot of k-values as a function of SSA is shown in Fig. 2.59b for two desorption temperatures 275 and 300°C. The k-dependence on SSA is very similar to that for the desorption temperature in Fig. 2.59a, meaning that k increases up to an SSA of ~15 m2/g and then there is no further dependence with increasing SSA. Another important finding, as can be seen in Fig. 2.59b, is that there is no apparent effect of carbon (0.34–2.97 wt.%) and oxygen (0.05–11.7 wt.%) content in the n-Ni on the hydrogen storage properties of MgH2. Finally, it is also obvious that there is no apparent effect of the n-Ni morphology on hydrogen storage properties. It is to be concluded that the observed effects of n-Ni on the hydrogen storage behavior of MgH2 strongly suggest that all the improvement of hydrogen storage properties is due to the kinetic effect of n-Ni acting as a powerful catalyst rather than to a modification of thermodynamic properties (enthalpy) of the MgH2 + n-Ni system. As can be seen in Fig. 2.53a, MgH2 doped with the n-Ni additive (SSA = 14.5 m2/g) does not desorb under 0.1 MPa hydrogen pressure at 250°C, which for MgH2 is lower than the equilibrium temperature at that pressure, while it easily desorbs over 4 wt.%H2 at 275°C, which is the equilibrium temperature at 0.1 MPa.

2.2

MgH2with Catalytic Additives 16000

165

n-Ni spherical

β-MgH2 γ-MgH2 Mg

MgO Ni ? unidentified

Counts

12000

NI7

8000 NI6

?

NI5

?

4000 NI4

MgH2

0 30

40

a

50 60 Degrees 2-Theta

70

80

90

n-Ni filamentary

16000

β-MgH2 γ-MgH2 Mg

MgO Ni

Counts

12000

NI10

8000 NI9 NI8

4000 NI3

MgH

0 30

b

40

50 60 Degrees 2-Theta

70

80

90

Fig. 2.56 XRD patterns of ABCR MgH2 + 5 wt.% n-Ni after milling for 15 min. (a) spherical and (b) filamentary morphology (see Table 2.20 for sample designation)

2.2.2

Mg/MgH2–Metal Oxides

Metal oxides seem to be one of the most potent catalytic additives to MgH2. A number of authors investigated only the absorption process of MgH2 catalyzed with a specific metal oxide without an insight into the desorption process. Obviously, the absorption

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DSC / mW/mg 8

MgH2 [1] 404.6⬚C

n-Ni spherical

exo

7 6

4

NI6 [4] 286.5⬚C NI7 [5] 280.8⬚C NI5

3

[2] 297.6⬚C

5

[3] 331.7⬚C

2

NI4 [2] 399.8⬚C

1 0 150

200

250

300

350

400

450

500

Temperature / ⬚C

a DSC / mW/mg 8

n-Ni filamentary [1] 404.6⬚C

exo

NI8

NI3

7 6

[2] 277.5⬚C [3] 290.5⬚C

5 4

[4] 276.3⬚C

[5] 280.2⬚C

NI9

NI10

3 2 1 0 −1 150

b

200

250

300

350

400

450

500

Temperature / ⬚C

Fig. 2.57 DSC traces of ABCR MgH2 + 5 wt.% n-Ni after milling for 15 min. (a) spherical and (b) filamentary morphology (heating rate 10°C/min; argon flow rate 50 ml/min). The range of onset temperatures of n-Ni containing mixtures is circled in red

kinetics were enhanced as compared to that of pure MgH2. However, as shown earlier in Figs. 2.4a and 2.22a, since for pure MgH2 the rate of hydrogen absorption is much faster than the rate of desorption, we will not review those papers, which were limited only to the absorption studies of MgH2 doped with a specific metal oxide. Therefore, this section will be restricted to review the studies of desorption. A very unfortunate aspect of the reported research results is that most of the investigators studied the effect of metal oxides on the desorption properties of MgH2

2.2

MgH2with Catalytic Additives

167

8.00 NI8

Hydrogen desorbed [wt.%]

7.00

NI3

NI7

NI9

T=275ⴗC

6.00 5.00 4.00

NI6

NI4, 5

NI10

3.00 2.00 1.00

MgH2

0.00 0

1000

2000 3000 Time [s]

a

4000

Hydrogen desorbed [wt.%]

8.00 NI7, 8, 9

7.00

NI10

T=300ⴗC

6.00 5.00 NI3, 6

4.00 NI4

3.00

NI5 MgH2

2.00 1.00 0.00 0

b

1000

2000 3000 Time [s]

4000

Fig. 2.58 Desorption kinetic curves at (a) 275°C and (b) 300°C obtained in a Sieverts-type apparatus under 0.1 MPa of hydrogen pressure of ball-milled ABCR MgH2 + 5 wt.% n-Ni having varying SSA (Table 2.20)

in vacuum. As mentioned earlier, the results obtained under such a condition have a very limited meaning. Oelerich et al. [122, 123] and Jung et al. [124] investigated the addition of V2O5. Under primary vacuum (0.1 kPa) the doped MgH2 was able to desorb 2–3 wt.%H2 at 250°C, while at 300°C about 6 wt.%H2 was desorbed within 120–360 s. The addition of Cr2O3 to MgH2 was studied by Jung et al. [124], Dehouche et al. [125], Barkhordarian et al. [126], Bobet et al. [127], and AgueyZinsou et al. [128]. Desorption of 4–7 wt.%H2 within 300–1,000 s was reported at 300°C, while some mixtures desorbed at 250–280°C in vacuum. Some studies were also made on the effect of oxides such as Fe2O3 and Fe3O4 [123, 126], Mn2O3 [123],

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Simple Metal and Intermetallic Hydrides

450 spherical ♦ filamentary

Temperature [ⴗC]

400 350 300

Tpeak 250 200

Ton

150 0

20

40

a

60

80

100

SSA [m2/g] 0.01

Rate constatnt k [s-1]

0.009 (0.34; 2.46)

(2.3; 2.08)

(1.86; 11.4)

0.008

spherical ♦ filamentary

(0.61; 3.5) (0.46; 0.11)

0.007 0.006

300ⴗC

0.005

(0.46; 11.7)

0.004

(2.97; 0.05) (0.53; 0.69)

0.003 0.002

275ⴗC

0.001 0

0

b

20

40

60

80

100

SSA [m2/g]

Fig. 2.59 (a) DSC peak (Tpeak) and onset (Ton) temperature and (b) rate constant, k, in the JMAK (6.4) as a function of the SSA of the Inco n-Ni additive. The first and second number in parentheses is the content of carbon and oxygen in wt.%, respectively

MgO [128], Al2O3 [124, 129], and La2O3 [130] without, however, any noticably better effect on the improvement of MgH2 desorption in vacuum than that brought about by V2O5, Cr2O3, or n-Ni for that matter. The most thoroughly researched oxide is Nb2O5, which, in particular, was extensively studied by a group from GKSS Geesthacht, Germany, led by Klassen and Bormann [126, 129, 131–136]. In vacuum, this oxide seems to make MgH2 desorb very fast at 300°C such that Barkhordarian et al., Friedrichs et al., and Aguey-Zinsou et al. [126, 129, 131, 135] reported that 5–7 wt.%H2 desorbed within 90–250 s from a catalyzed MgH2 ball-milled for 20–200 h. They also reported [126] desorption of about 6.6 wt.%H2 at 250°C within a very reasonable time of 600 s. However, in a recent paper from this group [135] they raised a question whether the Nb2O5 additive indeed acts as a typical catalyst or

2.2

MgH2with Catalytic Additives

169

whether the kinetic enhancement is mainly due to its effect on particle size reduction. They found that the desorption kinetics in vacuum of pure MgH2 milled for 700 h is the same as that of the mixture MgH2 + 17 wt.% Nb2O5 milled for 200 h, where the former is supposed to have nanometric particle size. However, their claimed particle size of pure MgH2 milled for 700 h being on the order of 0.5–1.5 nm is rather unlikely. Apparently, this interpretation seems to be somehow erroneous. So far, the lowest desorption temperature in vacuum of the ball-milled mixture MgH2 + Nb2O5 is 163 and 200°C, as reported by Hanada et al. [133] and Bhat et al. [136], respectively. Hanada et al. reported ~5 wt.%H2 desorbed within 6,000 s. They also reported that the activation energy for hydrogen desorption of the mixture of MgH2 + 1 mol.%Nb2O5 milled for 20 h was ~71 kJ/mol. This value is quite close to the activation energy that we obtained for MgH2 milled with Inco n-Ni (Figs. 2.52a and 2.54), which was obtained from the desorption experiments conducted in a Sieverts-type apparatus under atmospheric pressure of hydrogen (0.1 MPa). Apparently, Nb2O5 does not seem to be a more efficient catalyst than n-Ni. As mentioned earlier, the great pitfalls of conducting desorption experiments in vacuum are well illustrated by the results obtained by Song et al. [82, 83], who showed a profound effect of desorption pressure on desorption kinetics. In a Sieverts-type apparatus at 300°C, the ball milled mixture of MgH2 + 10 wt.%Cr2O3 desorbed ~4, 2.5, and 1.5 wt.%H2 under 1, 1.2, and 1.4 bar of hydrogen pressure, respectively. These results are in strong contrast to the amount of 4–7 wt.%H2 desorbed at 300°C and even at 250–280°C reported in [124–128] for the MgH2 + Cr2O3 mixture when the desorption process was conducted in vacuum. Song et al. also studied the desorption process of the mixture Mg–10 wt.%(Fe2O3,MnO,Ni) and found that at 320°C the material desorbed 2.2, 1.2, 0.9, 0.2, and 0 wt.%H2 under less than 1, 1.2, 1.4, 1.6, and 1.8 bar, respectively.

2.2.3

Mg/MgH2–Carbon/Graphite and Carbon Nanotubes

Imamura et al. [137] studied absorption of Mg ball-milled with graphite and benzene as milling additives. In a Sieverts-type apparatus, the mixture after 20 h milling was able to absorb at 180°C. This in itself is nothing outstanding because milled, activated, and cycled MgH2 can also absorb at ~200°C (Fig. 2.22a). Bouaricha et al. [138] also studied absorption of a Mg + graphite mixture at 300°C, which showed much better kinetics than that of just milled Mg. A number of researchers studied both absorption and desorption [25, 139–141]; but unfortunately, all desorption studies were conducted in vacuum. However, it must be pointed out that even under vacuum conditions, desorption kinetics were no better than those obtained with a number of other additives discussed earlier in the text. The lowest desorption temperature applied was 290°C. The addition of carbon nanotubes that were either reactively milled under hydrogen mixed with Mg powder [142] or simply mixed with MgH2 and subsequently milled [143, 144] was investigated. In vacuum, desorption at 200°C gave 3.6 wt.% within 1,800 s [142]. Another reference reports ~5 wt.%H2 desorbed at 300°C within

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Simple Metal and Intermetallic Hydrides

3,600 s and 6 wt.%H2 at 350°C in about 300 s from the ball-milled mixture of MgH2+ single-walled nanotubes (SWNTs) tested in vacuum [143]. The results indicate that there is no particular advantage of adding carbon nanotubes to Mg/MgH2. There are also a few papers reporting results on the hydrogen desorption of the ball-milled mixtures of MgH2 and carbon/graphite studied by temperature programmed desorption (TPD), DSC, and TGA [145–147]. However, they do not report any particularly exciting results that would warrant a more thorough discussion. In summary, there is no compelling evidence that carbon, graphite, or carbon nanotubes can act as potent catalytic additives for the enhancement of the absorption/desorption properties of Mg/MgH2. They are not really better than, for example, some elemental metals such as n-Ni, which can be treated as sort of a catalytic benchmark.

2.3

Other Metal Hydrides Containing Mg

In order to improve the hydrogen storage properties, a number of researchers have followed an approach where Mg was alloyed with the rare-earth (RE) elements to form intermetallic compounds of a stoichiometric composition. In the 1980s the research groups of Pezat et al. [148] and Khrussanova et al. [149–151] investigated the hydrogen storage properties of LaMg12, CeMg12, MmMg12 (Mm-mischmetal), La2Mg17, and La2Mg16Ni, where La can be substituted by Ca or Ce, which were fabricated by ingot metallurgy followed by grinding to a powder. It has been found that the first hydriding process produces in the binary system irreversible decomposition of the intermetallic in contact with hydrogen. For example, LaMg12 decomposes to LaH3 and MgH2, whereas La2Mg16Ni decomposes to LaH3, MgH2, and Mg2NiH4. In the subsequent reversible desorption–absorption cycling, the RE hydride reacts with MgH2 forming Mg, LaHx, and H2 (and eventually Mg2Ni) and upon absorption again returns to the original mixture of LaH3 and MgH2 (and eventually Mg2NiH4). More recent studies by XRD have confirmed this general scheme of reactions [152]. In these materials absorption may occur as low as 100–130°C, although the amount of absorbed hydrogen is only ~2.5 wt.% [148, 150]. At around 300°C and higher, these materials can absorb slightly more than ~4 wt.%H2. However, full desorption requires temperatures at least on the order of 300°C. More recently, it has been shown that ball milling of La2Mg16Ni in tetrahydrofuran and benzene improves the absorption/desorption kinetics [153–155]. In conclusion, one must say that the RE–magnesium alloys do not provide any superior hydrogen storage properties to catalyzed MgH2 (or Mg). In addition, their decomposition/disproportionation to the mixture of RE hydride and magnesium hydride on the first hydrogenation can be a drawback because it leads to a serious degradation if multicycling absorption/desorption is considered as an option. An offset of the research on the RE–Mg compounds for hydrogen storage resulted in an interest in the Mg3RE-type compounds, where RE = Y, La, Ce, Pr, Sm, Gd, Tb, and Dy. Using very high pressures on the order of a few gigapascals,

2.3

Other Metal Hydrides Containing Mg

171

Kamegawa et al. [156] synthesized Mg3LaH9, Mg3CeH8.1, and Mg3PrH9 from powders of elemental metals. These hydrides decomposed into Mg and RE-hydride at about 300°C with an endothermic reaction. Obviously, because of the high pressures involved in their synthesis, the hydrides are irreversible. Quyang et al. [157, 158] investigated hydrogen storage properties of Mg3La and Mg3LaNi0.1 (in reality a mixture of Mg3La and Mg2LaNi phases), as well as Mg3Mm and Mg3MmNi0.1, in which La was replaced with a cheaper mischmetal (Mm). The compounds were produced by conventional ingot metallurgy. These intermetallic compounds are able to absorb hydrogen even at room temperature and it takes only a few minutes to uptake hydrogen to 90% of their full hydrogen capacity, which is ~3 wt.%. Upon hydrogenation, the compounds do not disproportionate as it happens for the RE–Mg compounds, e.g., LaMg12. During the hydrogenation of Mg3La, the structure of the hydride that is formed is FCC and is totally different to that of the Mg3La intermetallic compound, but the exact crystallographic structure is still under investigation. The enthalpy (ΔH) and entropy ΔS of the Mg3La –H dehydriding reaction were determined from the PCT curves using the Van’t Hoff plot, giving values of 81 kJ/mol-H2 and 142 J/mol-H2 K, respectively. During the hydrogenation of Mg3LaNi0.1/Mg3MmNi0.1, a small amount of MgH2 appeared in addition to the major hydride phase, which was the same as the hydrogenated Mg3La. Unfortunately, desorption of about 2.5–3 wt.%H2 requires a temperature range of ~250–280°C. Zhang et al. [159] investigated a partial substitution of Mg with Al in Mg3La having a modified stoichiometry of (Mg1–xAlx)3La. For the x = 0.1–0.3 range, the alloy was a three-phase one consisting of Mg3La, Mg22AlLa6, and C15-Laves La(Mg,Al)2. Increasing x to 0.4 eliminated completely Mg3La. During hydrogenation, the phases disproportionated into a mixture of MgH2, LaH3, and La3Al11, similar to the RE-Mg compounds. Taking into account a relatively small hydrogen capacity and the desorption temperature range comparable to that of MgH2, the compounds based on Mg3La are not really competitive to catalyzed MgH2. Another intermetallic compound that attracted some attention is γ-Mg17Al12. Bouaricha et al. [160] investigated the structure and hydrogen absorption properties of Mg–Al alloys prepared by ball milling over a wide range of compositions. The major phases formed were γ-Mg17Al12 and Mg3Al2 with ratios depending on the composition (e.g., 100% of γ-Mg17Al12 at 58 Mg:42Al). Bouaricha et al. were the first to report that γ-Mg17Al12 disproportionated into MgH2 and Al during hydrogenation. Yabe and Kuji [161] and Crivello et al. [162] synthesized stoichiometric γ-Mg17Al12 by the bulk mechanical alloying (BMA) technique. BMA is based on repeated forging processes and is one of the novel methods to synthesize nanostructured materials, equivalent to mechanical milling. The intermetallic compound can theoretically absorb 4.44 wt.%H2 in a two-step reaction: Mg17Al12 + 9H2 ü9MgH2 + 4Mg3Al2 (2.40wt.%H2)

(2.4a)

4Mg3Al2 + 8H2 ü8MgH2 + 12Al (2.40wt.%H2)

(2.4b)

In practice, about 3.2 wt.%H2 is absorbed at 250°C under 5.3 MPa and slightly more at 350°C. Desorption requires at least 250°C.

172

2

Simple Metal and Intermetallic Hydrides

There is a large family of magnesium-based ternary metal hydrides with a perovskite-type crystal structure containing alkali and alkaline-earth elements, e.g., Na, K, Ca, Cs, Rb, Eu, Sr, Ba, Y, and a few more, which have recently been reviewed by Yvon and Bertheville [163]. Since most of them has hydrogen capacity within the 1–3 wt.% range, they are of no interest for automotive applications. Additionally, their equilibrium desorption temperature at atmospheric pressure is rather high, approaching 400°C. In this group, the highest hydrogen capacity of 6 wt.% is exhibited by NaMgH3. Very recently, Ikeda et al. [164] synthesized this hydride by ball milling the mixture of NaH and MgH2. At 380°C in a PCT test the alkaline ternary hydride desorbed 6 wt.% in a two-step reaction: NaMgH3 ® NaH + Mg + H2

(2.5a)

NaH + Mg + H2 ® Na + Mg + 3/2H2

(2.5b)

From the Van’t Hoff plot the decomposition enthalpy change ΔH) and entropy change (ΔS) for the first and second step were calculated to be ΔH(1) = 93.9 kJ/molH2 and ΔS(1) = 116.2 J/mol-H2K, and ΔH(2) = 102.2 kJ/mol-H2 and ΔS(2) = 125.9 J/ mol-H2K. Taking into account the very high desorption temperature of this compound, much higher than that of a catalyzed MgH2, and the fact that MgH2 is used in its synthesis by ball milling, the compound is not at all competitive to MgH2. Ball milling was extensively used for the synthesis of novel nanostructured BCC alloys Mg–Tm–V (Tm = Ni, Co, Cu) by Kuji et al. [164, 165]. Their approach was based on earlier findings that in the quasibinary TiCr2–V system, the Laves phase TiCr2 can dissolve in a BCC V at low temperatures. The same idea was applied to dissolve the Laves phases MgTm2 in BCC V. Alloys with the stoichiometric composition Mg1.0Tm1.0V1.0 (Tm = Ni, Co, Cu) formed a single-phase BCC structure having nanosized grains <10 nm after milling a stoichiometric mixture of elemental metal powders for 25 h. PCT tests at room temperature showed that they were able to absorb about 1–2 wt.% hydrogen under 3 MPa of hydrogen pressure. No desorption was investigated and no kinetic studies were carried out. In another attempt, CaMg2 was mixed with V, Fe, and Mo in a 50:50 ratio (at.%). After ball milling for 60 h, a single-phase nanostructured BCC phase having grain size on the order of 10 nm was synthesized. The BCC alloys were able to absorb hydrogen after hydrogenation for 100 h at 220°C under 15 atm. pressure. TDS (thermal desorption spectroscopy) tests showed the desorption peak temperature at ~388°C and the desorbed capacities were 3.0, 1.5, and 1.0 wt.%H2 for the V, Fe, and Mo alloying elements. A group of researchers led by Akiba [166, 167] synthesized by ball milling for 200 h nanostructured binary Mg–Co and ternary Mg–Co–X (X = Fe, Cu, Pd) BCC alloys. These BCC alloys are able to absorb 1–3 wt.%H2 at 100°C under ~6 MPa but they desorb only minute quantities of hydrogen at that temperature. Their hydrogen absorption and desorption enthalpies are within the 23–33 and 17–27 kJ/mol-H2 range, respectively. Such relatively small enthalpies in combination with the almost nonexistent desorption at 100°C suggest that this is a kinetic effect due to a very sluggish desorption kinetics in these types of alloys. Obviously, the BCC alloys are not any viable

2.3

Other Metal Hydrides Containing Mg

173

candidates for gaseous hydrogen storage in automotive applications, but they may be considered for electrochemical storage, for example, in batteries, although it is rather doubtful that they would perform better than the well-known and tried LaNi5 (MmNi5) alloys. Hanada et al. [168] synthesized by ball milling the ternary Laves compounds MgYNi4, MgCaNi4, and CaYNi4 having C15b(AuBe5)-type Laves phase structure. After hydrogenation at room temperature under 2.0 MPa hydrogen pressure, the MgYNi4-H sample desorbed ~1.0 wt.%H2 in a PCT test at 40°C. Guénée et al. [169] synthesized by ingot metallurgy ternary intermetallic compounds LaNi4Mg and NdNi4Mg having the cubic MgCu4Sn crystal structure. They can absorb reversibly up to four hydrogen atoms per formula unit at 7–8 bar and ~50°C. The hydrides are stable at room temperature but desorb quite rapidly at 80°C in vacuum. In air, they decompose by catalytic water formation. It must be mentioned for the sake of clarity that conventional ingot metallurgy (arc or induction melting) was employed to synthesize more exotic ternary RE intermetallic hydrides such as Ce6Ni2Si3, Sm{Co,Ni}2Si3, La2Ni0.8Si1.2H3.75, La2NiSiH3.9, La2Ni1.2Si0.8H4.4, Ce2Ni0.8Si1.2H3.7, Ce2NiSiH4.4, Ce2Ni1.2Si0.8H4.9, La6Ni2Si3H12, and Ce6Ni2Si3H10.9 by Morozkin et al. [170, 171] and Lushnikov et al. [172]. Hydrogenation was performed at hydrogen pressures up to 50 MPa but some hydrogen uptake was observed at room temperature under only 0.1–0.5 MPa hydrogen pressures. Desorption started at around 250°C but continued to rather high temperatures such that 40% of the hydrogen absorbed was released at ~600°C with a complete desorption and transformation to the starting intermetallic compound at ~980°C. No nanostructuring of these ternary hydrides by ball milling was attempted. Nevertheless, owing to their low hydrogen capacity which is slightly higher than 1 wt.% and to relatively high desorption temperature range, they do not constitute a viable storage medium and can be treated as rather a scientific curiosity. Very recently, Sahlberg and Andersson [173] synthesized Mg–Y–Zn ternary intermetallic compounds and investigated their hydrogen storage properties. The phase structure of the samples consisted of a mixture of ternary intermetallic phases Mg3Y2Zn3, Mg12YZn, and Mg. At temperatures above 300°C, Mg12YZn decomposed into Mg and Mg3Y2Zn3. During hydrogenation, Mg3Y2Zn3 decomposed into YH2 and MgZn at hydrogen pressures below 100 kPa and 400°C. Above 100 kPa, magnesium started to absorb hydrogen, which led to the formation of Mg4Zn7 and MgH2. At pressures above 300 kPa, YH2 transformed to YH3 and above 1 MPa Mg4Zn7 decomposed into MgH2 and MgZn2. No solid solution of hydrogen in Mg3Y2Zn3 was observed. There was no absorption of hydrogen below 400°C. DTA/ TGA measurements showed a big endothermic peak together with a decrease in mass at 407°C. XRD analysis of the desorbed sample indicated formation of Mg3YZn6 according to the following reaction: YH3 + MgZn2 ® Mg3YZn6 + YH2 + H2

(2.6)

No desorption occurred below 400°C. Automatically, these compounds can be eliminated from any consideration as viable storage materials.

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Zn was also used as an additive to Mg by Deledda et al. [174], who attempted to synthesize by ball milling two intermetallic compounds Mg7Zn and Mg7Zn3, which according to some earlier predictions were supposed to form hydrides. After milling up to 50 h in argon atmosphere, liquid nitrogen (cryomilling), and hydrogen (reactive milling), the final result was the formation of an amorphous Mg–Zn phase with a likely composition of about Mg45Zn55. A fraction of Mg was left unreacted and transformed to MgH2 if ball milling was carried out in a reactive H2 atmosphere. Hydrogen absorption in a Sieverts-type apparatus for the powders ball milled in argon resulted in the partial conversion of the Mg into MgH2. The amorphous Mg–Zn phase did not play any significant role in the hydrogen sorption behavior and eventually crystallized into Mg21Zn25. The latter partially decomposed into MgH2 and MgZn2, provided that the hydrogenation was extended for sufficiently long times. Goo and Hirscher [175] synthesized nanocrystalline MgS directly from MgH2 and S powders by high-energy ball milling in an argon atmosphere. The XRD peaks revealed the formation of MgS phase with NaCl-type structure. The average grain size estimated by the diffraction line profile analysis was about 11 nm, which was confirmed by TEM studies. SEM investigations showed that the average particle size of the powder was below 10 μm. At room temperature, no significant reaction with hydrogen was observed. Lomness et al. [176] investigated hydrogen absorption in a mechanically alloyed mixtures of Ti–-Mg–-Ni and claimed that the one of the mixtures Ti53 Mg47Ni20 was able to absorb (11 wt.% H2 for a sample ball- milled for 37 h at a ball/-powder mass ratio of 70:1 with a hydrogen absorption onset temperature of ~90°C. In an attempt to reproduce their results, Sheppard et al. [177] processed the same composition in accordance with that of Lomness et al. They showed that the powder mixture Ti53 Mg47Ni20 milled for 36 h could rapidly absorb 2.5 wt.H2 at room temperature under 1-–6 MPa pressure. This is significantly less than the capacity claimed by Lomness et al. Ball- milled samples comprised low- crystallinity elemental Ti, Mg, and Ni which converted to Ti2Ni, Ti, and Mg during temperature cycling. Examination of the hydrided sample revealed the presence of Ti2NiHx (x > 0.5), TiHy (0.7 < y < 2) and MgH2.

2.4

AlH3

The simple aluminum hydride, AlH3 (alane), is a covalently bonded, metastable solid at room temperature with a large gravimetric and volumetric hydrogen capacity of 10.1 wt.% and 148 kgH2/m3, respectively. Its density is 1.48 g/cm3. In reality, it has been known to the community of chemists for 60 years, but only recently its high values of hydrogen capacity attracted the attention of a number of researchers who perceive this hydride as one of the primary candidates that can meet the Department of Energy (DOE) targets for supplying fuel cell–powered cars (see Table 1.2 in Sect. 1.1). The solvated form of hydride was first synthesized by Finholt et al. [178] by the reaction of LiAlH4 with AlCl3 in ether, yielding an ethe-

2.4

AIH3

175

real solution of AlH3, mostly for various applications in organic and inorganic chemistry. A nonsolvated form of AlH3 was initially prepared by Chizinsky et al. [179] and subsequently prepared by Brower et al. [180] by organometallic route at the Dow Chemical Co. In general, the early research has shown that there exists a number of polymorphs of AlH3 which were identified as α, α′, β, δ, ε, γ, and ζ. The appearance of such a large number of polymorphs is attributed to the fact that the synthesis of AlH3 is extremely sensitive to the desolvating conditions, i.e., temperature and time. Freshly prepared, nonpassivated AlH3 is pyrophoric and reacts violently on contact with water. The first thermodynamic information on α-AlH3 was obtained by Sinke et al. [181]. They measured the enthalpy of formation and absolute entropy at 298.15 K as −11.4 ± 0.84 kJ/mol and 30.1 ± 0.4 J/mol K, respectively. Subsequently, Herley et al. [182–184] studied the thermal and photolytic (UV irradiated) decomposition kinetics of α-AlH3. In general, decomposition was quite slow. At 145°C, the length of time to obtain a full desorption was about 100 min and it was correspondingly longer at lower temperatures. Nearly no desorption was observed at 100°C. Slightly faster kinetics were obtained in the isothermal coirradiated decomposition, but still too slow for any practical application as a storage material. As pointed out by Sandrock et al. [185, 186], the decomposition/desorption reaction of α-AlH3 occurs in a single endothermic reaction a - AlH3 ® Al + 3/2H2

(2.7)

This reaction is in essence irreversible under reasonable pressure conditions, and H2 gas pressures in excess of 2.5 GPa are needed to rehydrogenate Al back to AlH3 [187]. Therefore, from the standpoint of perceived application as a hydrogen storage medium for fuel cell–powered cars, AlH3 would have to be recharged off board (see also Sect. 1.1). Despite its large dissociation pressure, AlH3 is metastable at room temperature, most likely due to the presence of an Al2O3 layer encapsulating each particle. Sandrock et al. attempted to improve the hydrogen desorption properties of α-AlH3 by simultaneous ball milling and doping to bring the hydride closer to the application as a storage material for hydrogen-fueled vehicles. They used the α-AlH3 (trigonal/rhombohedral crystal structure) produced in 1975 by Dow Chemical. It had somehow a lower capacity of 8.3 wt.% compared to 10 wt.% of pure AlH3. Using ball milling they achieved simultaneous particle size reduction and doping. TPD (temperature programmed desorption) showed substantial lowering of desorption temperature range of pure (undoped) α-AlH3 due to ball milling, which reduced the particle size from the as-received 100 μm to ~1 μm after 1 h, and to ~0.3 μm after 3 h of ball milling, respectively. Accordingly, desorption range was reduced from 175–200°C for as-received α-AlH3 to 125–175°C after 1 h of milling (for ~7 wt.%H2 desorbed). Doping with Ni and Ti was unsuccessful, but good results were obtained with alkali-metal hydrides such as LiH, NaH, KH, and LiAlH4. Addition of 10–20% LiH reduced desorption temperature range 40–50°C from the pure AlH3 milled for 1 h. During isothermal desorption at 100°C, 80 mol%AlH3–20 mol%LiH mixture desorbed 4 wt.%H2 within 300 min. They explained the enhancement of kinetics by invoking a “window model” for the trans-

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Fig. 2.60 Schematic representation of LiH + AlH3 → LiAlH4 mechanical alloying reaction during ball milling to produce H2-transparent LiAlH4 “windows” [186] Table 2.21 Thermodynamic and kinetic data for the decomposition of freshly made α-, β-, and γ-AlH3 polymorphs. after [188–190] Activation energy Total enthalpy of Surface area Particle desorption Ea (kJ/ formation, ΔHtot Polymorph (m2/g) diameter (nm) mol) (kJ/molAlH3) 10.9 204 102.2 ± 3.2 −9.9 ± 0.6 α-AlH3 14.6 152 92.3 ± 8.6 −8.0 ± 1.0 β-AlH3 16.4 136 79.3 ± 5.1 −7.1 ± 1.0 γ-AlH3

port of hydrogen from the AlH3 phase. According to this model, which is shown in Fig. 2.60, ball milling of LiH-doped AlH3 leads to the formation of LiAlH4 nanosized phase on the surface of a AlH3 particle originally encapsulated by an Al2O3 layer. These nano- LiAlH4 serve as “windows” through which H2 can easily escape during desorption. Addition of 5TiCl3 to 75AlH3 + 20LiH resulted in a slow desorption behavior during ball milling and afterwards in a Sieverts-type apparatus for several hours. Apparently, TiCl3, which is a well-known catalyst for LiAlH4 (it will be discussed in Chap.3), accelerates their “window” action even at room temperature. Graetz et al. [188–190] studied the decomposition of two other polymorphs β-and γ-AlH3 in comparison to the α-polymorph, all of which were freshly made by the organometallic method developed by Brower et al. [180]. Table 2.21 compiles thermodynamic and kinetic data extracted from their papers. It was found that at temperatures ≥100°C, decomposition of β- and γ-AlH3 occurs with the initial exothermic phase transformation of β- and γ-AlH3 → α-AlH3, which subsequently decomposes in the

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endothermic reaction of (2.7). Therefore, the total formation enthalpy for the β- and γ-AlH3 in Table 2.21 is the sum of both transformation enthalpy and heat of formation of Al + 3/2H2 (2.7). At low temperatures, <100°C, two decomposition pathways are possible for β- and γ-AlH3: (1) a phase transformation to α-AlH3 followed by decomposition to the elements (2.7), or (2) direct decomposition to the elements (2.7). In general, at temperatures <100°C, the isothermal decomposition of fresh α-AlH3 was observed to be faster than that of the “old” Dow Chemical hydride [185, 186], and those of β- and γ-AlH3 were slightly faster than α-AlH3. NMR studies generally confirmed the previous findings [191]. Also, similar results were obtained by a European group [192] who studied thermal decomposition of AlH3 by in situ synchrotron XRD and thermal desorption spectroscopy. They obtained the activation energy Ea of desorption as 136 and 92 kJ/mol for α-AlH3 and γ-AlH3, respectively. They also observed a significant anisotropic volume expansion of α-AlH3 during heating with the principal expansion along the a-axis. From the thermal decomposition of γ-AlH3 they deduced two parallel paths: (1) γ → Al and (2) γ → α → Al, although the latter seemed to be the prevailing one. Orimo et al. [193] studied the influence of ball milling on the thermal behavior of α-, β-, and γ-AlH3 polymorphs using thermogravimetric and differential thermal analysis (TG-DTA) supported by XRD. They observed a dehydriding reaction occurring during milling in a planetary ball mill (Fritsch P7) at 400 rpm and only Al was detected after 60 min of milling. This observation is in contrast with that of Sandrock et al. [185, 186], who did not report such a behavior upon milling. The dehydriding reaction during milling was also observed for the γ-AlH3 phase, but the α-phase which formed the γ-form was not decomposed after 60 min of milling. No change in the β-phase was observed after 60 min milling. In summary, the α-phase decomposes into metallic Al either upon milling or heating. The γ-phase initially undergoes transformation to the stabilized α-phase either upon milling or heating and subsequently decomposes to Al in an endothermic reaction. The β-phase undergoes exothermic dehydriding only upon heating, not upon milling. According to Orimo et al., this implies that β-phase is “chemically” stabilized by the chemical species (e.g., chlorides used in the synthesis of AlH3). The observed behavior of the β-phase seems to differ from the one reported by Graetz et al. [188–190], as described earlier, who observed transformation β → α before subsequent thermal decomposition of α. Finally, it must be mentioned that Graetz and Reilly [188] clearly pointed out that the conventional organometallic synthesis of AlH3 is a costly procedure. This factor, together with its on-board irreversibility, calls for novel, very cost-effective, and energetically efficient techniques to regenerate AlH3 from spent Al powder.

2.5

Other Metal and Intermetallic-based Hydrides: New Developments

A brief general overview of simple metal hydrides and intermetallic-based hydrides including their historical development has already been given in Sect. 1.1 and 1.2. Early research results on these hydrides have been summarized in a book by Mueller

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et al. [194]. Since then, a number of excellent reviews have appeared on classical metal and intermetallic hydrides, by Schlapbach [195], Zaluski et al. [31], Sandrock [196], and Grochala and Edwards [197]. As pointed out in Sect. 1.1, conventional pure metal hydrides and intermetallic-based hydrides listed in Table 1.3 (AB (FeTi, ZrNi), AB2 (ZrMn2/TiMn2/TiCr2 Laves phases), AB5 (LaNi5 or MmNi5 where Mm-mischmetal), and A2B (Mg2Ni)) have gravimetric capacities too low or desorption temperatures too high for any commercial consideration in mobile hydrogen storage. Therefore, none of the metal/intermetallic hydrides can meet the DOE targets (Sect. 1.1). However, because a number of them can reversibly store hydrogen at near ambient conditions of pressure and temperature, they can be potentially used in a variety of stationary or semistationary applications. Such near-ambient conditions enable the use of the heat of environment to be used as a heat source for providing the enthalpy of hydride desorption/decomposition or heat dissipation sink during absorption or charging a battery. In the near future, metal/intermetallic-based hydrides, owing to its advantage of high volumetric hydrogen density/capacity (Fig. 2.61), can find some niche market applications such as fuel cell–powered submarines, fuel cell–powered fork lifts, notebook computers, cellular phones, cordless

Fig. 2.61 Gravimetric and volumetric densities, and the corresponding specific energies and energy densities, of a variety of different hydrogen storage media, including two common fossil fuels, gasoline, and propane. The FreedomCar targets for 2005, 2010, and 2015 are indicated within the shaded regions. The gravimetric hydrogen densities for pressurized hydrogen gas and cryogenic liquid hydrogen (diamond symbols) include the mass of the storage container, whereas the values reported for the metal hydrides, complex hydrides, and other hydrogen storage materials are based on the absolute (theoretical) amount of hydrogen [198]

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tools [199], and hydrogen compression systems (based on alloyed TiMn2, TiV2, and TiCr2) [200]. In essence, metal/intermetallic hydrides of AB, AB2 (Laves phases), AB5, and A2B have not been improved to any remarkable extent since the end of the 1990s. Nevertheless, there have been some efforts directed to either improvement of synthesis or properties of metal/intermetallic hydrides particularly by application of ball milling associated with nanostructuring. Some of these more recent efforts will be briefly discussed in the following sections.

2.5.1 Metal Hydrides Elansari et al. [201] developed a novel method of synthesizing alkali metal hydrides Na, KH, RbH, and CsH by reactive mechanical milling of pure alkaline metals under hydrogen pressure up to 30 bars in a planetary mill (Retsch PM 400). The reaction proceeds in 16 h and gives 3–15 g of very pure alkali metal hydride with FCC crystal structure (space group Fm3m). Yukawa et al. [202] investigated the effect of alloying elements M = Ti, Cr, Fe, Co, Ni, Cu, Zr, Nb, Mo, Ru, Rh, Pd, Ag, Hf, Ta, W, Re, Os, Ir, Pt, Au, In, and Sn on the stability of vanadium hydrides γ-VH2 and β-VH or V2H (sometimes labeled β1-V2H and β2-VH; these are nominal compositions not stoichiometric and hydrogen content varies depending upon the solubility at various temperatures; the reader is encouraged to consult a V–H phase diagram, e.g., [1], p.437). In the V–H system, the β-phase is extremely stable and its hydrogen desorption never occurs under moderate conditions. In turn, the γ-phase is not so stable and can decompose according to VH2 → VH + ½H2 at moderate temperatures and pressures. They synthesized the V–1 mol%M and V–3 mol%M alloys by arc melting, cold rolling, cutting into plates, polishing mechanically and chemically, and finally activating the sample before hydrogenation. Subsequently, they acquired PCT curves at 40°C. They observed that the PCT plateau pressure (at the midpoint) changed in a systematic way following the order of elements in the periodic table. The pressure was high for the group 8 elements, Fe, Ru, and Os. Within the range 1–3 mol%, the logarithm of plateau pressure increased almost linearly with increasing amount of the element for Ni, Cr, Mo, Ru and decreased for Nb and Ti. Zr did not affect the plateau pressure. However, the authors were unable to explain the observed change unambiguously. Two peaks appeared in a DSC trace. XRD showed that the first peak at ~280°C was due to the transformation of the β-phase to the BCC (α)-phase containing hydrogen. The second peak at ~380°C was assigned to the transformation of the BCC (α)-phase to BCC vanadium metal. In an alloy with Mn, Mo, and W, the low-temperature peak disappeared. For Mn, Fe, Co, Ni, Rh, Pd, W, and Re, the high-temperature peak shifted to higher temperatures as compared to that for the pure V metal.

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Very recently, Yan et al. [203] developed a family of vanadium-based BCC solid solution alloys. They observed that for the Vx– (Ti-Cr-Fe)100–x (where Ti/(Cr + Fe) = 1.0, Cr/Fe = 2.5, x = 20–55) alloys, with increasing V content the hydrogen absorption capacity and desorption capacity both increased, and plateau pressure decreased. Two phases, the BCC and Laves phase, appeared when x < 30, and a single BCC phase appeared when x ≥ 0. For the Vx– (Ti-Cr-Fe)100–x (where Ti/(Cr + Fe) = 1.0, Cr/Fe = 2.5) alloys with V ≥ 30at.%, three alloys, V30Ti33Cr27Fe10, V42Ti28Cr21.7Fe8.3, and V55Ti21.5Cr17.1Fe6.4, were developed, which have hydrogen capacities 2.06, 2.12, and 2.23 wt.%H2 at 298 K, respectively. The maximum desorption capacity in this alloys system was obtained at a unit lattice dimension of 3.042 Å, where V40Ti28.4Cr23.6Fe8 alloy absorbed 3.82 wt.%H2 and desorbed 2.30 wt.%H2 at 298 K. The alloy with V/Fe ratio of 5:1 could be prepared from a lowcost FeV80 master alloy. Similarly, Song et al. [204] investigated the effects of vanadium content and alloying elements (Mn and Ni) on the microstructure and hydrogen absorption/desorption properties in vanadium-based solid solution alloys. The samples were prepared by conventional ingot metallurgy (arc melting) and subsequently activated at 400°C. They found that with increasing vanadium content from 5 to 10 to 35% in the Ti–xV–Cr alloys, the quantity of V-based solid solution phase increased and eventually became the majority phase. The microstructure of the 77.8 V–7.4Zr– 7.4Ti–7.4Ni alloy consisted mainly of V-based solid solution phase and a small amount of Laves phase. With increasing vanadium content from 5 to 10 to 35 to 77.8%, the hydrogen absorption capacity increased from 1.30 to 1.55 to 2.88 to 3.11 mass%. At the same time, the hydrogen desorption capacity increased from 0.52 to 0.70 to 1.79, but then decreased to 1.34. The sorption tests were carried out at room temperature. With the increasing vanadium content, the pressure plateau in the PCT curve became more and more distinctive and steadily shifted down to lower pressure levels. It was also found that a small quantity of Mn and Ni could effectively increase the plateau pressure in the 35% vanadium alloy and the hydrogen absorption capacity in the 5% vanadium content alloy, but had little effect on their hydrogen absorption kinetics. Chandra et al. [205] studied thermally cycled (up to 4,000 cycles between 398 and 298 K) and cold-worked V–0.5at.%C alloy from its standpoint as a potential candidate for use in metal hydride heat pumps (MHHPs) [206], hydrogen compressors, and closed-cycle cryogenic refrigerators [207]. A small amount of carbon was added to overcome the problem of microstrain development (excessive plastic deformation) during the γ-phase formation due to large volume changes during β2 ↔ γ transformation, and to determine whether the loss in absorption hydrogen storage capacity can be reduced. They found that C had minimal effect on the high pressure β2 + γ phase region, although stability in the low pressure α + β1 region increased. Thermal cycling between β2- and γ-phases increased hysteresis owing to the increased absorption pressures and unchanged desorption pressure. Hydriding of the cold-worked alloy increased the pressure hysteresis significantly owing to decreased desorption pressure and lowered the hydrogen capacity. In the β2-phase, microstrains decreased after 778 cycles, while crystallite sizes (called “domains”

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by the authors) increased. In contrast, in the γ-phase both microstrains and crystallite sizes decreased. After 778 cycles, the dehydrogenated α-phase showed residual microstrains similar to those in intermetallic hydrides.

2.5.2

Rare-Earth AB5 Compounds

Zhu et al. [208] prepared by ball milling nanocomposites of Mg and MmNi5–x(CoAlMn)x whose microstructure consisted of nanometer-sized Mg, MmM5, and MmMg17 phases (M = Co, Al, Mn, Ni). PCT measurements at 30°C showed increased hydrogen capacity of a nanocomposite. However, although initial absorption rate at 60°C and atmospheric pressure of a nanocomposite was higher than that of the melted/ascast sample, the desorbed hydrogen capacity of the former was much smaller. The authors explained this by the formation of a dense layer of hydride during absorption kinetic measurements under an applied constant pressure of 0.6 MPa, as opposed to a step-by-step pressure application in a PCT test. Analysis of the kinetic curves for absorption showed that the rate-controlling step of the hydriding process of melted/ as-cast MmM5 is different from that of the Mg–MmM5 nanophase composite. This is due to a profound difference between the coarse-grained structure of melted/as-cast material and nanometric-sized grains in a nanocomposite. Fujii et al. [209] studied the effect of mechanical milling under argon and hydrogen on the storage properties of LaNi5. Nanocrystalline material with grains on the order of ~20 and 10 nm was produced after 60 and 600 min of milling, respectively. Milling under argon resulted in a lowering of plateau pressure at room temperature PCT test and narrowing of the width of pressure plateau. Reactive milling under hydrogen resulted in a formation of nanocrystalline LaNi5H0.15 and an amorphous phase coexisting together when the milling time was <180 min. After longer reactive milling, LaNi5H0.15 disappeared and an amorphous phase dissociated into nanocrystalline Ni and amorphous LaNiyHx (y < 5). No plateau was observed on a PCT curve after reactive milling for more than 60 min. The authors concluded that the hydrogen storage properties cannot be improved by ball milling in systems containing metals with a strong affinity for hydrogen-like RE metals. Rozdzynska-Kielbik et al. [210] substituted a fraction of Ni in a LaNi5 compound by Zn. PCT tests from room temperature to 80°C showed decreasing plateau pressure, while XRD showed an increase of the unit cell volume. The hydrogen capacity slightly decreased with increasing Zn content. Laurencelle et al. [211] and Cheng et al. [212] investigated the effect of hydrogen absorption/desorption cycling on the storage properties of LaNi5, where Ni was partially substituted with Sn and Al, respectively. The alloys were prepared by conventional ingot metallurgy (arc and induction melting, respectively). The Sn-substituted compound was used in the first stage of a metal hydride compressor prototype. It was found that after 1,000 cycles of aging, the hydrogen storage capacity was reduced by 2.2%, the particle size decreased, but the crystalline structure and reaction kinetics remained unchanged. In contrast, when the Al-substituted

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LaNi5 was cycled 1,000 times, the plateau pressures at 90, 110, and 130°C were only slightly changed, while a significant degradation of absorption kinetics was observed. The pulverization of the bulk alloys occurred, and the volume mean particle diameter was reduced from ~35 μm before to ~11 μm after 1,000 cycles. The changes in properties were explained by disproportionation and pulverization phenomena and the effect of impurities in the hydrogen used for charging.

2.5.3

Titanium–Iron AB Compounds

Hotta et al. [213] focused on the synthesis of nanocrystalline TiFe intermetallic by mechanical alloying and its hydrogen storage properties. The time of ball milling to achieve a full synthesis was 20 h. Since TiFe intermetallic is notoriously difficult to activate, the ball-milled powder was annealed at 300°C in vacuum and subsequently under 15 MPa of hydrogen for five times. The PCT at room temperature showed two plateaus, the first at 0.5 MPa from 0.5 to 0.9mass%H2, and the second at a whopping 25 MPa. It has been reported that for the TiFe synthesized by conventional ingot metallurgy the first plateau usually occurs at ~0.4 MPa and is associated with the formation of β-TiFeH hydride having an orthorhombic crystal structure. The second plateau usually around 1 MPa corresponds to the formation of γ-TiFeH with a monoclinic structure. Apparently, for the nanocrystalline TiFe synthesized by ball milling the first plateau occurs at a similar pressure to that of the as-cast material but the second plateau is much higher than that for the conventionally synthesized intermetallic. The authors suggested that such a big difference could be related with atomic arrangement and/or coordination surrounding the octahedral sites in the γ-phase, but no quantitative explanation was given. In the follow-up paper from the same group, Abe et al. [214] combined mechanical alloying with postannealing. In some contradiction to the first paper, they claimed that after 20 h, there was also an amorphous phase in the structure although the XRD pattern shown in the paper is not convincing enough. They claimed that in a DSC test the intermetallics milled for more than 10 h recrystallized at ~480°C. An alloy ball-milled for only 5 h and postannealed at 500°C for 3 h in argon was nearly fully synthesized to the TiFe intermetallic as confirmed by XRD. It could absorb hydrogen in a PCT test without any activation and displayed only a single plateau in contrast to the previous report. The samples postannealed at 400°C did not exhibit a clear plateau. The maximum storage capacity in a PCT test was ~1.25 wt.% at 5 MPa H2. There are certain discrepancies in both reports that warrant some more studies of this interesting synthesis technique. Saita et al. [215] used hydriding combustion synthesis for a direct production of TiFe. In the experiments, an exothermic reaction of Ti with hydrogen (Ti + H2= TiH2 + 144 kJ) was utilized for HCS of TiFe because the adiabatic flame temperature of this reaction was estimated to be 2,000°C, which is sufficiently high for melting both iron and titanium. A 1:1 molar mixture of elemental Ti and Fe pow-

References

183

ders was covered with additional Ti powder and placed in a hydrogen atmosphere under 0.9 MPa pressure. After the ignition of the Ti powders by an electrically heated carbon filament, the TiFe synthesis reaction (Ti + Fe = TiFe + 40 kJ) occurred successfully. Before hydrogenation, the product was activated by evacuating and charging hydrogen at a temperature of 400°C; then, the product was cooled to 25°C and was able to store 1.7mass% H2. Unfortunately, no hydrogenation pressure was given in the original paper.

2.5.4

Titanium and Zirconium AB2 Compounds

Sometimes, phase transformations that occur upon milling can be detrimental to the hydrogen storage properties. Ball milling was used by Huot et al. [216] to process TiV0.9Mn1.1 alloy (TiMn2-type Laves phase). Arc-melted and cast alloy consisted of C14 Laves phase and a BCC phase. Upon ball milling for up to 80 h, the C14 phase disappeared while an FCC phase appeared. The same result was obtained when raw elemental metal powders were used for synthesis by mechanical alloying. The hydrogen storage properties were measured after activation by cycling between high hydrogen pressure of 5 MPa and vacuum, at 250°C. The absorption PCT test at 23°C showed that the as-cast alloy was able to absorb ~1.9 wt.%H2 under 7 MPa. The same alloy milled for 80 h did not absorb hydrogen. It seems that the presence of the FCC phase formed during milling somehow blocks the hydrogen intake but the exact explanation is lacking.

2.5.5

Other Novel Intermetallic Hydrides

Guénée and Yvon [217] synthesized LaNi2Mn3 intermetallic with a YNi2Al3-type crystal structure. However, its hydrogen storage properties are no better than those of LaNi5-type because it could barely absorb ~1.4 wt.%H2 forming LaNi2Mn3H6 and its equilibrium temperature at 1 bar H2 is about 60°C. No attempt of nanostructuring the compound by ball milling was made. The Aoki group [218] has been developing intermetallic alloys based on a CaSi compound that is alloyed with Si, Al, Ge, Mg, and Sr. However, the alloys cannot compete with the LaNi5-type or even TiFe because they absorb only slightly more than 2 wt.%H2 at 100°C and desorb at 200°C.

References 1. H. Okamoto, “Desk Handbook-Phase Diagrams for Binary Alloys,” ASM International, Materials Park, OH, (2000), p.430. 2. T. Noritake, S. Towata, M. Aoki, Y. Seno, Y. Hirose, E. Nishibori, M. Takata, M. Sakata, “Charge density measurement in MgH2 by synchrotron X-ray diffraction,” J. Alloys Compd. 356–357 (2003) 84–86.

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Chapter 3

Complex Hydrides

A brief description of the development of complex hydrides has been given in Sects. 1.2.8 and 1.2.9. Selected hydrides from this group have been thoroughly discussed in a number of review papers which appeared in the past few years [1–7] the most notable of them being the most recent one [7]. The term “complex” hydride is rather liberally applied to a rather large group of hydrides by various authors. In the broadest sense these are hydrides composed of an anionic metal–hydrogen complex or non metal–hydrogen complex bonded to a cationic alkali or transition metal [1]. Hence, the entire large group can be roughly subdivided into two categories: group I and II – salts of [AlH4]−, [NH2]−, [BH4]−, i.e., alanates, amides, and borohydrides [7], and transition metal (TM) complex hydrides that have anionic (TMHx)− complexes such as [FeH6]4− attached to a cationic light metal, e.g., Mg2+, in Mg2FeH6 [1]. Their bonding is usually an ionic– covalent mix. Similar transition metal ternary complex hydrides exist in the Mg–Co, Mg–Ni and Mg–Mn systems forming Mg2CoH5, Mg2NiH4 and Mg3MnH7 (the latter was synthesized under 20 kbarH2 at ~ 800°C [8]). Confusingly, some references also classify certain complex hydrides such as for e.g., NaBH4 as “chemical hydrides” due to the fact that they can easily react with water solution of KOH or NaOH (or water steam) releasing hydrogen and serving as potentially alternative irreversible source of hydrogen (e.g., so called Millennium Cell) [9–13]. According to the classification of Chandra et al. [5], and probably more correctly, chemical hydrides are such hydrogen-bearing compounds as methanol (CH3OH), methylcyclohexane (CH3C6H12), ammonia (NH3) and ammonia borane (NH3BH3), and other organic compounds. These compounds could also be used to generate hydrogen, e.g., by steam reforming, but still they are irreversible and may increase pollution. As already shown in Table 1.4 and in Fig. 2.61 a number of complex solid hydrides have high and very high theoretical gravimetric and volumetric hydrogen capacities combined with relatively low desorption temperature range due to quite favorable thermodynamics (enthalpies). This alone has made them a very hot topic of research in the past 10 years or so. Unfortunately, a number of them are still plagued by fundamental problems like high kinetic barriers to dehydrogenation and irreversibility. Nevertheless, in the all honesty, it must be said that if a very R.A. Varin et al., Nanomaterials for Solid State Hydrogen Storage, DOI: 10.1007/978-0-387-77712-2_3 © Springer Science + Business Media, LLC 2009

195

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3 Complex Hydrides

restrictive target of reversibility established by D.O.E. (see Sect. 1.1 and Table 1.2) could be moderated then at least a few complex hydrides would be quite close for vehicular applications in the near future. In the following sections we will discuss the most interesting of them with a special emphasis on their thermal decomposition from the engineering standpoint. We will not be discussing various aspects of their crystallographic structures. The crystallographic structures of complex hydrides as determined by X-ray or neutron diffraction data are rather complex and the interested reader is referred to [7] for more details on this subject.

3.1

Ternary Transition Metal Complex Hydrides

Three transition metal complex hydrides Mg2NiH4, Mg2CoH5 and Mg2FeH6 have been extensively studied, especially from the standpoint of their synthesis to nanostructured hydrides by ball milling.

3.1.1

Mg2NiH4

The Mg2Ni intermetallic compound with the hexagonal hP18 (Pearson) lattice structure exists in the equilibrium binary Mg–Ni system [14]. Its complex hydride Mg2NiH4 has a theoretical hydrogen capacity of ~3.6 wt%. In this hydride two Mg2+ ions donate two electrons each to stabilize the [NiH4]4− complex [15]. The crystallographic and hydrogen storage properties of Mg2NiH4 were first investigated and reported by Reilly and Wiswall [16]. From XRD they determined the lattice structure of Mg2NiH4 as tetragonal. The measured density of the product was 2.57 g/cm3. PCT tests at various temperatures showed a very flat plateau between the 3 and 10 atm. pressure range, depending on temperature. They established a full hydrogen reversibility of the compound according to a simple reaction Mg 2 Ni + 2H 2 ↔ Mg 2 NiH 4

(3.1)

They suggested that in the first stage of hydrogenation hydrogen dissolves in the Mg2Ni phase to the extent that the ratio of H/(Mg + Ni)≅0.1 according to the reaction 0.54 Mg2 NiH 0.3 + H2 ↔ 0.54Mg2 NiH 4

(3.2)

The standard enthalpy (ΔH) and entropy (ΔS) of formation at 298 K was −64.5 kJ/molH2 and −122.3 J/molH2K, respectively. It is to be noted that the enthalpy of formation/decomposition of Mg2NiH4 is slightly lower than that of MgH2 (Sect. 2.1) which allows the former to desorb at slightly lower temperatures than MgH2.

3.1

Ternary Transition Metal Complex Hydrides

197

Reilly and Wiswall reported that the polycrystalline form of Mg2Ni produced by a conventional ingot metallurgy still required activation by short cycling at ~ 300°C. Difficulty with activation, still relatively high temperature of absorption/desorption and rather sluggish kinetics, prompted Singh et al. [17] and immediately after that Zaluski et al. [18] to use mechanical alloying (MA) of elemental metal powders for the synthesis of nanocrystalline Mg2Ni with the hope of achieving substantial improvements in its hydrogen storage properties. Indeed, ball-milled nanocrystalline Mg2Ni showed better hydrogen sorption properties than the conventional polycrystal. A number of other researchers followed this path of producing nanocrystalline Mg2Ni by MA or direct synthesis of Mg2NiH4 by reactive mechanical milling (RMA). Varin and Czujko [19] critically reviewed the results reported in the literature up to 2001 with a special emphasis on the improvements of storage properties brought about by the formation of nanostructure in the processed powders. In the past few years there still have been some research efforts focused on Mg2NiH4. The Akiyama’s group in Japan [20–22] has developed a hydriding combustion synthesis technique for the direct fabrication of Mg2NiH4. The results showed that the product of Mg2NiH4 from the hydriding combustion synthesis absorbed the maximum (3.4–3.6 wt%) near the theoretical value just after synthesis without any activation process. Hampton et al. [23–24] activated the Mg2Ni and Mg2.35Ni alloys (Hy–Stor 301 by Ergenics; slightly enriched in Mg to prevent disproportionation of Mg2Ni to Mg and MgNi2) by a reaction with liquid water or water vapor for hydrogen uptake. Shao et al. [25] prepared Mg2Ni from magnesium and nickel nanoparticles produced by hydrogen plasma–metal reaction. Two preparation methods were developed to obtain the compound. One is heating the nanoparticles under 0.10 MPa argon pressure at 430°C and the other is under 3.00 MPa hydrogen pressure at 280°C. No hydrogen storage properties of this material were assessed. Varin et al. [26, 27] studied microstructural evolution in the nearly dual-phase Mg–Ni alloy containing various fractions of the Mg2Ni and MgNi2 intermetallic phases during mechanical milling under a controlled shearing mode in the magneto-mill Uni-Ball-Mill 5. It was observed that the Mg2Ni phase underwent a partial amorphization in the Mg–Ni alloys containing ~61 and ~79 vol.% of the MgNi2 phase while no amorphization of Mg2Ni was observed in the alloy containing only ~9 vol.% of MgNi2. The results were rationalized in terms of the enthalpy effects based on the application of Miedema’s semi-empirical model to the phase changes in ball-milled intermetallics and the critical nanograin size required to be formed in the Mg2Ni phase before triggering its amorphization, which is enhanced by the presence of hard MgNi2 phase during ball milling. Unfortunately, the milled powders of 27.9 ± 11.1 at.% Ni alloy, after long-term milling for 100 h, did not absorb hydrogen. Chen et al. [28] investigated the hydrogenation characteristics of the slurry composed of the NH4F solution treated Mg2Ni and liquid C6H6. At 210°C and under a hydrogen pressure of 4.0 MPa alloy absorbed hydrogen first, transformed into hydride and then the benzene was hydrogenated to cyclohexane with the hydride as

198

3 Complex Hydrides

the catalyst. The hydrogen absorption capacity of the slurry system composed of 20 wt% treated alloy and benzene reached ~ 6.4 wt% and the hydrogenation completed in ~ 20 min. More recently a number of research groups focused on various synthesis methods for unalloyed and alloyed Mg2Ni. Liu et al. [29] synthesized nanostructured Mg2Ni alloys by hydriding combustion synthesis and subsequent mechanical milling. The product was able to absorb ~ 3 wt%H2 at 100°C under initial hydrogen pressure of 3.0 MPa. The same amount of hydrogen was released during continuous heating under vacuum from ~100 to ~300°C. Xie et al. [30] used hydrogen plasma-metal reaction to synthesize nanostructured Mg2Ni1−xCox (x = 0.05, 0.1) alloys. However, no substantial improvement of storage properties was observed as the alloys could only absorb slightly over 2 wt%H2 at 200°C but desorption in primary vacuum still required about 300°C. Hara et al. [31] synthesized Mg2Ni alloys with the addition of Y and depleted Mg by casting and vacuum heating. The alloys could absorb ~2–3 wt%H2 at 200°C but only a minimal amount of ~0.5 wt%H2 at room temperature. Kumar et al. [32] prepared Mg2Ni by polyol reduction method which exhibited increased hydrogen absorption characteristics. The maximum absorption and desorption capacities were 3.2 and 2.8 wt%, respectively, at 275–330°C. As can be seen from this brief review, Mg2NiH4 does not have a real potential for applications in gaseous hydrogen storage, primarily, due to its relatively low capacity and high temperature of hydriding/dehydriding. Its capacity is much lower than that of MgH2 and sorption temperature range is almost the same as that of catalyzed MgH2. There is no advantage here over a simple MgH2. On the other hand, alloyed Mg2Ni could find some applications in electrochemical hydrogen storage as a potential replacement for MmNi5-type alloys in the Ni-MH batteries as reviewed by Wronski [33].

3.1.2

Mg2FeH6

Mg2NiH4 is rather an exception to a general situation of transition metal complex hydrides in the sense that it has a corresponding equilibrium intermetallic compound Mg2Ni in the binary system of Mg–Ni, as mentioned in the preceding section. In the binary Mg–Fe system there is no solubility of Fe in either solid or liquid Mg and obviously no intermetallic compounds are formed between both metallic elements (e.g., [14, p. 366]). However, when the stoichiometric 2:1 mixture of Mg and Fe (2Mg–Fe) is synthesized under hydrogen, the complex hydride Mg2FeH6 forms. This hydride was originally synthesized by Didisheim et al. [34] by sintering of elemental metal powders at ~ 500°C under around 60 bar of hydrogen pressure according to the reaction 2Mg + Fe + 3H 2 → Mg 2 FeH 6

(3.3)

3.1

Ternary Transition Metal Complex Hydrides

199

XRD measurements showed that the compound has a cubic lattice with the lattice parameter a = 0.6443 nm and metal atom arrangement of the fluorite type (space group Fm3m). Theoretical studies of its X-ray absorption spectra were conducted soon after its discovery by Orgaz and Gupta [35]. The compound attracted attention of a number of researchers due to its relatively high theoretical hydrogen capacity of 5.5 wt% (Table 1.4). The expectation was that its absorption/desorption temperature could be much lower than that of MgH2 owing to the presence of Fe in the compound and its possibly catalyzing effect on MgH2 which was discovered [36] a few years before the successful synthesis of Mg2FeH6. Ivanov et al. [37] and Konstanchuk et al. [38] were the first to apply ball milling to the mixture of Mg-25 wt%Fe and after successive hydrogenation/ dehydrogenation reported the formation of Mg2FeHx hydride where x was assumed to be between 4 and 6. The alloy had enhanced hydriding/dehydriding properties. Reiser et al. [39] and Bogdanovic´ et al. [40] studied the Mg2FeH6 compound from the standpoint of its potential application for thermochemical thermal energy storage, for example as the storage of solar or excess industrial heat. The compound was synthesized by hydrogenation/dehydrogenation cycles at 450/547°C, 82/95 bar H2 of the powder mixture of 2Mg–Fe for 1.5/1.5 h. They measured PCT curves at 350, 400, 450, 500 (desorption/absorption) and 400 and 525°C (desorption). All PCT curves showed horizontal plateaus with no hysteresis with the following mid-plateau pressures: 3.6 bar/350°C, 11.25 bar/400°C, 28.5 bar/450°C, 65.17 bar/500°C, and 93.7 bar/525°C. From the Van’t Hoff plot they calculated the enthalpy (ΔH) of formation as −77.4 kJ/molH2 which was much lower than the previous values of −98 ± 3 kJ/molH2 [34] or 86 ± 6 kJ/molH2 [38]. Bogdanovic´ et al. concluded that Mg2FeH6 and a mixture MgH2 + Mg2FeH6 are highly suitable materials for thermochemical thermal energy storage at around 500°C, due to their excellent stability in cycle tests and further advantages including low price of the starting materials, a free choice and constancy of the heat delivery temperature by controlling the applied hydrogen pressure, and the absence of heat losses with time. A search for improvements in the synthesis which would avoid sintering at high temperatures and pressures, and in the hydrogen storage properties of Mg2FeH6, particularly in lowering its desorption temperature, prompted a large number of researchers to apply mechanical alloying and reactive mechanical alloying for the synthesis of nanocrystalline Mg2FeH6 [41–47]. The reactive mechanical alloying synthesis of nanocrystalline Mg2FeH6 was extensively studied in our laboratory [48–52]. We tested various controlled conditions of reactive mechanical alloying in the magneto-mill Uni-Ball-Mill 5 (Fig. 1.7 in Sect. 1.3.2). It was found that a lowenergy milling under 400–450 kPa of hydrogen of the 2Mg–Fe elemental mixture using shearing mode led to the formation of β-MgH2 within the first 10–15 h. Afterwards, the XRD peaks due to β-MgH2 disappeared with increasing milling time up to 59 h and only unreacted Fe peaks remained in the XRD pattern (Fig. 3.1). Energy dispersive spectroscopy (EDS) analysis clearly showed that Mg was still present in the microstructure of powder particles after milling up to 59 h. This very unusual behavior was interpreted as the formation of amorphous Mg and its hydride although it must be clearly pointed out that the XRD pattern of the 59 h reactively

200

3 Complex Hydrides Mg MgH2

Intensity (cps)

Fe

59h 40h 20h 15h 10h 5h 30

40

50 60 Degrees 2θ

70

80

90

Fig. 3.1 XRD patterns of 2Mg–Fe powders reactively mechanically alloyed under low-energy shearing mode

milled powder in Fig. 3.1 does not resemble a typical XRD pattern for an amorphous material which usually has two very diffused peaks. So far, the best results were obtained in our laboratory when controlled reactive mechanical alloying (CRMA) was carried out under impact 2 mode with a single magnet (IMP2; Fig. 1.7). Figure 3.2a shows the evolution of XRD patterns with increasing milling time when milling was conducted in a sequential mode in which small samples were extracted at selected times after opening the milling vial (shown in Fig. 3.2a). For comparison, the XRD pattern of a powder milled continuously for 270 h (270 h-cont.) is also shown. After 270 h of reactive milling the microstructure of the processed powder is dual phased and consists of Mg2FeH6 and retained Fe. Figure 3.2b shows the morphology of the Mg2FeH6 powder synthesized continuously for 270 h. The powder structure is very homogeneous. The grain size of Mg2FeH6 calculated from the XRD peak broadening (Sect. 1.4.3) is on the order of 6–8 nm while the grain size and lattice strain of retained Fe decreases and increases, respectively, with increasing reactive milling time (Fig. 3.3). This behavior of grain size and strain is typical for ball-milled metals. Figure 3.4 shows DSC traces of 2Mg–Fe mixtures milled sequentially for 210 and 270 h and continuously for 270 h. A single strong endothermic peak corresponding to the decomposition of the Mg2FeH6 phase is observed. Interestingly, such a single peak is also observed for a mixture of β-MgH2 + Mg2FeH6 [49, 51] at milling

3.1

Ternary Transition Metal Complex Hydrides

201

6000

Mg2FeH6 MgO Fe β-MgH2 (220)

(111)

(311)

4000

(400) (331)

(422) (511)

Counts

270h-cont. 270h 210h

2000

188h 100h 18h 0 20

a

30

40

50 60 Degrees 2-Theta

70

80

90

b Fig. 3.2 (a) Evolution of the XRD patterns of 2Mg–Fe mixture reactively milled sequentially for various times under IMP2 mode in 880 kPa of hydrogen. For comparison the XRD pattern of the mixture milled continuously for 270 h is also shown. (b) Morphology of 2Mg–Fe mixture reactively milled for 270 h in a continuous manner

times shorter than 210 h (Fig. 3.2a). The onset temperature of desorption is rather low being ~200°C and this without any catalyst. The maximum peak tempe-ratures are all below 300°C. The abundance of the ternary complex hydride Mg2FeH6, produced during sequential milling, and estimated from DSC (Sect. 1.4.4) reached

202

3 Complex Hydrides

Fe

50 40

Strain (%)

Grain size (nm)

60

30 20 10 0 0

50

a

100 150 200 Milling time (h)

250

300

b

0.2 0.18 0.16 0.14 0.12 0.1 0.08 0.06 0.04 0.02 0

Fe

0

50

100 150 200 Milling time (h)

250

300

Fig. 3.3 (a) Grain size and (b) lattice strain changes of retained Fe with increasing reactive milling time of 2Mg–Fe mixture DSC / mW/mg 0

exo

−0.5 −1.0 −1.5 [2] 271.9⬚C −2.0

1-210h-seq. 2-270h-seq. 3-270h-cont.

[1] 295.8⬚C

−2.5

[3] 273.3⬚C 50

100

150

200

250 300 Temperature / ⬚C

350

400

450

500

Fig. 3.4 DSC traces of 2Mg–Fe mixtures milled sequentially for 210 and 270 h and continuously for 270 h (heating rate of 4°C/min and argon flow rate of 16 ml/min)

~ 44 wt% after 188 h, and afterwards it slightly decreased to 42 wt% after 210 and 270 h. In contrast, the DSC yield of Mg2FeH6 after continuous CRMA for 270 h was estimated from DSC as ~ 57 wt%. Much higher yield after continuous milling was attributed to the absence of MgO in this sample (Fig. 3.2a). Unfortunately, quite promising hydrogen desorption behavior in DSC as shown in Fig. 3.4 did not translate into desorption in a Sieverts-type apparatus as shown in Fig. 3.5. The powder milled sequentially for 270 h desorbed in a Sieverts-type apparatus at 250 and 290°C (Fig. 3.5) under primary vacuum only about 1.2 wt%H2 which is approximately a half of the hydrogen content obtained during DSC and TGA tests. No desorption of hydrogen was detected in a Sieverts-type apparatus at 250 and 290°C after 128 and 70 min, respectively, from the powder continuously milled for 270 h. The latter easily desorbed 3.13 and 2.83 wt%H2 in DSC and TGA

3.1

Ternary Transition Metal Complex Hydrides

203

Desorbed hydrogen (wt.%)

0 −0.2 −0.4 −0.6 −0.8 −1 −1.2 −1.4 0

10

20

30

40

50

60

Desorption time (min)

Fig. 3.5 Desorption curve at 290°C under primary vacuum obtained in a Sieverts-type apparatus from the 2Mg–Fe mixture reactively milled in a sequential mode for 270 h

tests, respectively. On the basis of the experimental results we proposed [52] that the following reaction is most likely responsible for the formation of Mg2FeH6 during mechano-chemical synthesis of 2Mg–Fe 2Mg(s) + 2H 2 (g) → 2MgH 2 (s)

(3.4a)

2MgH 2 (s) + Fe(s) + H 2 (g) ↔ Mg2 FeH 6 (s)

(3.4b)

where (s) refers to “solid” and (g) to “gas.” As can be seen a persistent problem in the mechano-chemical synthesis of Mg2FeH6 is the presence of retained Fe in all powders subjected to CRMA. That indicates that the synthesis reaction of Mg2FeH6 is still incomplete. As reviewed in Sect. 2.2.1, vanadium is one of a few catalytic metals which improve storage properties of Mg/MgH2. Therefore, we decided to add V to the reactively pre-milled 2Mg–Fe mixture in order to check if it improves the rate of the Mg2FeH6 synthesis upon further CRMA in a similar manner as the addition of micrometric-, submicrometric- and nano-Ni benefited the synthesis of MgH2 (Sect. 2.2.1). In the first step, the vanadium turnings were mechanically milled under the IMP2 mode in the protective He gas for 13 h with nine drops of toluene added as a lubricant which prevents the relatively ductile V powder from adhering and cold-welding to the balls and the milling vial wall. In the second step, the elemental mixture 2Mg–Fe was subjected to CRMA for 200 h. Finally, 4.5 g of this powder was mixed with 0.1 g of milled V powder (~ 2.2 wt%) and subjected to CRMA under IMP2 mode for 70 h. Crystallite (nanograin) size of Mg2FeH6 as estimated from the XRD peak breadth was about ~ 10 nm in V-bearing milled powder. The enthalpy of DSC desorption from V-bearing 2Mg–Fe mixture was lower than that for the reactively pre-milled 2Mg–Fe mixture resulting in a lower yield of Mg2FeH6 [52]. Apparently, addition of V did not improve the yield of Mg2FeH6.

204

3 Complex Hydrides

More recently Zhou et al. [53] investigated the energy and electronic structure of Mg2FeH6 by using a first-principles plane-wave pseudopotential method to calculate heats of formation and on that basis the formation mechanism of Mg2FeH6. They concluded that after formation of MgH2 according to the reaction of (3.4a), the iron atoms progressively dissolve in the MgH2 lattice forming an intermediate (MgFe)H2 solid solution which subsequently transforms into the Mg2FeH6 phase. In other new developments Saita et al. [54] studied parameters of the hydriding combustion synthesis of Mg2FeH6 and Mg2Ni1-xFex hydrides with and without pre-milling. Puszkiel et al. [55] carried out quite exhaustive thermodynamic and kinetic studies of Mg2FeH6 which was fabricated by mechanical milling in argon followed by sintering under hydrogen. Li et al. [56] used hydriding combustion synthesis under a magnetic field for the synthesis of Mg2FeH6. Unfortunately, none of these recent studies resulted in the product whose hydrogen storage properties would be superior over materials synthesized by a simple reactive mechanical alloying. In summary, it must be pointed out that the research on Mg2FeH6 brought about rather disappointing results. Nanostructured Mg2FeH6 has lower hydrogen capacity than nanostructured MgH2 and the desorption temperature of ternary complex hydride is not much lower than, for example, that of MgH2 catalyzed with nano-Ni (Sect. 2.2.1). The retained Fe usually present in the microstructure after reactive milling is very difficult to remove and is just a ballast reducing the practical hydrogen capacity of the compound. Hydrogen release rate from Mg2FeH6 seems to be slower than that of MgH2. On the other hand, Mg2FeH6 is so far the one of only two complex hydrides (the second is Mg2CoH5 which is discussed in the following section) which can be relatively easily synthesized with the yield ~ 60–70 wt% by reactive ball milling of elemental metal powders under hydrogen or mechanical alloying of pre-formed MgH2 and Fe and it is most likely fully mechanically reversible.

3.1.3

Mg2CoH5

The complex hydride Mg2CoH5 is very similar to Mg2FeH6. In the binary system of Mg–Co there is no solubility of Co in either solid or liquid Mg and no intermetallic compound, Mg2Co, exists in equilibrium with other phases. However, in contrast to the Mg–Fe system, the intermetallic compound MgCo2 exists in equili-brium in the Mg–Co binary system (e.g., [14, p. 251]). The theoretical hydrogen capacity of Mg2CoH5 is only 4.5 wt% which is obviously lower than that of Mg 2FeH 6 due to the presence of the heavier Co element and one less H atom in the hydride formula. The first successful synthesis of polycrystalline Mg2CoH5 and its deuteride Mg2CoD5 was reported by Zolliker et al. [57]. The 2Mg–Co elemental mixture was sintered under 4.0–6.0 MPa of hydrogen at temperatures between 350 and 500°C. The formation of the hydride was not complete and the main remaining phases, MgH2 and Co had to be removed by a separate purification process. X-ray and neutron diffraction data suggested a tetragonally distorted CaF2-type crystal structure

3.1

Ternary Transition Metal Complex Hydrides

205

of Mg2CoH5 with a = 0.4463 nm and c = 0.6593 nm (for the deuteride). The structure transformed at ~ 215°C into a disordered cubic structure with a = 0.6453 nm. They estimated the enthalpy of formation of Mg2CoH5 as 86 ± 5 kJ/molH2. Ivanov et al. [58] were the first to use ball milling of the 2Mg–Co elemental mixture to synthesize Mg2CoH5 either directly during milling or first, by pre-milling and subsequent pressing into pellets and hydriding. They observed three plateaus in a desorption PCT curve corresponding to Mg2CoH5, the hexagonal phase which was identified as Mg3CoH5 and MgH2. From the Van’t Hoff plot they calculated ΔH = −79 ± 4 kJ/molH2 and ΔS = −134 J/molH2K for Mg2CoH5, and ΔH = −70 ± 4 kJ/molH2 and ΔS = −118 J/molH2K for Mg3CoH5. Thermal decomposition of both phases led to the formation of an intermetallic compound MgxCo with an FCC crystal structure (a = 1.143 ± 0.001 nm). Selvam and Yvon [59] synthesized a nearly single-phase Mg2CoH5 hydride by sintering at 9.0 MPa and 450–500°C. Depending on the hydrogen pressure they observed the simultaneous formation of Mg2CoH5 and Mg6Co2H11. Huot et al. [60] employed ball milling of the 2Mg–Co elemental mixture, followed by sintering at 350°C for 1 day under 50 atm. of hydrogen pressure. The yield of the ternary hydride was only about 30 wt%. Chen et al. [61] directly synthesized Mg2CoH5 by mechanical alloying of 2MgH2–Co mixture under hydrogen. They observed two plateaus on absorption/desorption PCT curves (at 280 and 320°C) which they interpreted as corresponding to MgH2 and Mg2CoH5. At 380°C only a single plateau appeared on the PCT curve. Mg2CoH5 decomposed/formed according to the following reversible reaction Mg 2 CoH 5 ↔ MgcCo + Mg + Co + H 2

(3.5)

From the Van’t Hoff plot they calculated the enthalpy and entropy change of absorption as ΔH = −69.5 kJ/molH2 and ΔS = −129.6 J/molH2K, respectively. Correspondingly, for desorption they calculated ΔH = −83.2 kJ/molH2 and ΔS = −146.7 J/molH2K. It is interesting that the enthalpy of desorption of ball-milled, nanocrystalline Mg2CoH5 is higher than that for absorption. This is a very similar behavior to the one observed for the ball-milled, nanocrystalline MgH2 as discussed in Sect. 2.1.4 (Fig. 2.43). Shao et al. [62] used Mg and Co nanoparticles with the size of 200–300 and 5–60 nm, respectively, for the synthesis of nanoparticulate Mg2CoH5 (particle size 50–300 nm range) by pressing 2:1 powder mixtures of Mg and Co into pellets and then very complicated sintering under a hydrogen scheme. They observed the formation of two hydride phases: Mg2CoH5 and Mg3CoH5. From the Van’t Hoff plot the formation enthalpy and entropy were estimated as −82.3 kJ/molH2 and −138.8 kJ/molH2K for Mg2CoH5, and −73.2 kJ/molH2 and −123.0 kJ/molH2K for Mg3CoH5, respectively. The most exhaustive studies of the synthesis and hydrogen storage properties of Mg2CoH5 synthesized by RMA (reactive mechanical alloying) and a combination of RMA and sintering were carried out by the group led by Gennari [63–65]. They observed a transition from tetragonal to cubic structure of Mg2CoH5 at

206

3 Complex Hydrides

~ 200°C as first reported by Zolliker et al. [57]. The ternary hydride decomposed at 390°C with the formation of MgCo intermetallic. On a PCT curve they identified two plateaus corresponding to the formation of Mg2CoH5 and Mg6CoH11 hydrides. Hydriding/dehydriding reactions were reversible on cycling with particularly, the fast absorption rate in the range 250–400°C. However, the amount of absorbed/ desorbed hydrogen did not exceed ~ 3 wt%. In summary, it is rather obvious that Mg2CoH5, similar to Mg2FeH6, is not competitive in comparison with a catalyzed MgH2 and have no potential whatsoever as a hydrogen storage medium in mobile or stationary applications.

3.2

Alanates

As discussed earlier in this chapter alanates constitute a sub-category in the category of complex hydrides which roughly covers alanates, amides, and borohydrides. The theoretical maximum and reversible gravimetric hydrogen capacities of alanates are on an average rather high and most of them exceed or are at least equal to the theoretical gravimetric capacity of MgH2 which can be taken as a sort of a benchmark (Table 1.4 and Fig. 2.61). However, the situation does not look so rosy if one compares their theoretical reversible gravimetric capacities which are the amounts of hydrogen that are potentially feasible (but not necessarily practically available) to be reversibly absorbed/desorbed from an alanate. A quick look at the numbers in Table 1.4 (Sect. 1.1) shows that only LiAlH4 has theoretical reversible gravimetric capacity which is more or less equal to the capacity of MgH2. However, one can also see in Table 1.4 that the approximate desorption temperature range of most of the important alanates is much lower than that of MgH2. This factor combined with their very adequate volumetric capacities (Fig. 2.61) make them so attractive for being considered as prime candidates for mobile hydrogen storage in automotive to supply fuel cells or internal combustion engines with hydrogen. In this section we will focus on four alanates with high practical potential such as NaAlH4, LiAlH4 and Mg(AlH4)2/Ca(AlH4)2. Alanates with theoretical gravimetric capacities lower than that of NaAlH4, being impractical for mobile applications, will not be discussed.

3.2.1

NaAlH4

The sodium alanate hydride NaAlH4 is so far the most investigated and promising hydrogen storage material, at least on a short time scale (for reviews see [1–7]). According to the literature the thermal decomposition of NaAlH4 occurs in a three-step reaction:

3.2

Alanates

207

(R1a) NaAlH 4 (s) → NaAlH 4 (1)

(3.6)

(R1b) NaAlH 4 (l) → 1/3Na 3 AlH 6 (s) + 2/3Al(s)

(3.7)

+ H 2 (g) ( ΔH = 37 kJ/molH 2 ) (R2) 1/3Na 3 AlH 6 (s) → NaH + 1/3Al + 0.5H 2 ( ΔH = 47 kJ/molH 2 )

(3.8)

(R3) NaH → Na + 0.5H 2

(3.9)

where (s) refers to “solid” and (l) refers to “liquid.” (R1b)–(R3) proceeds with a theoretical hydrogen release of 3.7, 1.9 and 1.8 wt%, respectively (and obviously lower for a purity-corrected capacity). Figure 3.6 shows SEM pictures of as-received and ball-milled NaAlH4 powder (purity 90%) which we investigated in our laboratory [66]. The as-received powder is rather coarse having the particle size larger than 10 μm and seems to cluster to form much larger aggregates (Fig. 3.6a). Quite energetic milling for 5 h in the magneto-mill, Uni-Ball-Mill 5, does not seem to be very effective leading to some refinement of particle size but there is still a tendency for the clustering of smaller particles into larger aggregates (Fig. 3.6b). Figure 3.7a shows a typical DSC trace obtained from a pure, undoped NaAlH4 obtained in our laboratory. The exothermic peak at 175.4°C is most probably due to the presence of surface hydroxyl impurities in the powder as reported by Andreasen [67] for LiAlH4. For the endothermic 182.1°C peak (R1a), the corresponding TGA weight loss (TGA traces not shown here) is almost non-existent. As a result, it is most likely that (R1a) is due to the melting of NaAlH4 which is reported by Claudy et al. [68] and Gross et al. [69] to occur at ~ 180°C. Dilts and Ashby [70], who were the first to investigate the thermal behavior of NaAlH4 by DTA/TGA assigned this endo peak to an unspecified phase change. In turn, Zaluski et al. [71] interpreted it as the first-step decomposition of NaAlH4. The origin of the small peak at 258.6°C is not clear. TGA shows a very small but still

5h milled NaAlH4

As received NaAlH4

a

b

Fig. 3.6 Secondary electron micrographs showing morphology of (a) as-received NaAlH4 and (b) after milling for 5 h in the magneto mill Uni-Ball-Mill 5 under HES57 mode in argon (two magnets at 5 and 7 o’clock positions; ball to powder weight ratio 50:1)

208

3 Complex Hydrides

Fig. 3.7 DSC trace of as-received, undoped NaAlH4 (purity 90%) and (b) the same hydride after milling for 5 h in the magneto mill Uni-Ball-Mill 5 under HES57 mode (two magnets at 5 and 7 o’clock positions). Heating rate 10°C/min at argon flow 50 ml/min

recognizable weight loss at this temperature range. Claudy et al. [68] reported a phase transition of pseudocubic α-Na3AlH6 into face-centered cubic β-Na3AlH6 at about 252°C which is close enough to the peak centered at 258.6°C. However, such an explanation would require the existence of α-Na3AlH6 from at least partially decomposed NaAlH4. TGA does not provide any firm evidence of partial decomposition of NaAlH4 in this range of temperatures although it is possible that the decomposition of NaAlH4 starts at this temperature with a simultaneous transformation of α-Na3AlH6 into β-Na3AlH6. For the peak with the maximum at 297.9°C the corresponding TGA curve shows in the temperature range 250–320°C, a substantial weight loss of ~ 3.5 wt%, which is close to the theoretical purity-corrected value of ~ 3.3 wt%H2 of the first-step decomposition of NaAlH4 (R1b). As a result,

3.2

Alanates

209

the 297.9°C peak is most probably attributed to the decomposition of NaAlH4 into Na3AlH6, Al and hydrogen (R1b). For the peak at 379°C, the corresponding TGA weight loss (350–400°C) is ~ 1.6 wt%, which is close to the theoretical value ~ 1.7 wt%H2 of the second-step decomposition of NaAlH4 (R2). As a result, the 379°C peak is most probably due to the decomposition of Na3AlH6 into NaH, Al and hydrogen (R2). NaH might decompose at a temperature higher than 400°C, which gives a weight drop in a TGA curve (not shown here). Further XRD study of powders heated to certain DSC peak temperatures is needed to reaffirm the nature of DSC-TGA thermal events in NaAlH4. It is to be pointed out that severe ball milling of NaAlH4 for 5 h does not change, within an experimental scatter, the positions of the DSC peaks as shown in Fig. 3.6b. According to the literature the reaction (R3) of (3.9), corresponding to the decomposition of NaH, can only occur at the range 425–500°C and possibly even higher [4, 6, 72] which is considered too high for any practical application. Therefore, the maximum theoretical hydrogen capacity available from NaAlH4 is about 5.6 wt% from the reactions (R1b) and (R2) of (3.7) and (3.8). As will be shown in the following paragraphs, doping can eliminate melting of highly purified NaAlH4. Also, Gross et al. [69] clearly showed that if a pure uncatalyzed NaAlH4 was isothermally annealed at 150°C for 4 h under vacuum it showed no desorption. Apparently, melting is first needed to desorb a pure NaAlH4. This picture is dramatically changed if catalysts are added as will be discussed further. For a long time NaAlH4 was considered in all practical terms as an irreversible hydride although it could be synthesized under hydrogen pressure from solvents [7]. Dymova et al. [73] achieved the first direct synthesis of NaAlH4 from the elemental metals Na and Al at the temperature range 270–280°C under hydrogen pressure of 17.5 MPa and higher. The clue of this synthesis route was that Na was in a molten state. Nevertheless, this route was still far away from any practicality. In the 1990s Bogdanovic´ and Schwickardi [74] discovered that the doping of NaAlH4 with small amount of Ti compounds through wet chemistry rendered the hydride reversible in the solid state under moderate conditions. NaAlH4 doped with 2 mol%Ti(O-n-C4H9)4 (or Ti(OBu)4) desorbed ~ 4.5 wt%H2 at 160°C within ~ 10 h and accordingly shorter time was needed at higher temperatures. The rehydrogenation of Ti-doped samples dehydrided to the NaH + Al stage occurred at 170°C under 15.2 MPa of hydrogen pressure. Higher initial rates of hydrogenation were obtained by using Ti(OBu)4 instead of β-TiCl3. Soon after this seminal discovery Zaluska et al. [71, 72] used mechanical milling of NaAlH4 to enhance its desorption/absorption properties. They reported that after 2 h of ball milling undoped NaAlH4 could desorb ~ 3 wt%H2 within ~ 2 h at 160°C. Even faster desorption rates could be obtained if carbon was used as an activator milled together with the hydride. Such a milled mixture with carbon could also easily absorb 3–4 wt%H2 at ~ 130°C under ~ 9 MPa of hydrogen pressure. It must, however, be pointed out that our own studies of the desorption of ball-milled undoped NaAlH4 do not confirm such fast desorption kinetics as claimed by Zaluska et al. [71, 72]. Figure 3.8a shows desorption curves of NaAH4 ball milled for 5 h as described earlier which was subjected to desorption during continuous

210

3 Complex Hydrides

Desorbed hydrogen (wt.%)

7 425ⴗC

6 5

400ⴗC

4

300ⴗC

3 2 225ⴗC 1 0

0

1000

a

Ball milled 100wt.% NaAlH4

2000 3000 Time (sec)

4000

450 400 400ⴗC

Temperature (ⴗC)

350 300

300ⴗC

250 200

225ⴗC

150 100 50 0 0

b

1000

2000 3000 Time (sec)

4000

Fig. 3.8 (a) Desorbed hydrogen (wt%) from the NaAlH4 milled for 5 h as a function of heating time while heating to 225, 300, 400 and 425°C. Quasi-Temperature Programmed Desorption (TPD) carried out in a Sieverts-type apparatus. (b) Temperature profile change as a function of time while heating to 225, 300 and 400°C in a Sieverts apparatus

heating from room temperature to the desired temperature. Figure 3.8b shows the profile of temperature change with time during heating. As a result, the sample is heated under dynamic heating for a few hundred seconds followed by an isothermal (static) heating (quasi-TPD test). As can be seen in Fig. 3.8a desorption rate at 225°C, which is the lowest temperature used in our tests, allows desorption of approximately 4 wt%H2 in more or less 2 h of continuous heating. This is apparently a slower desorption rate than that reported by Zaluska et al. It must be pointed out that the same purity of NaAlH4 was used in our studies as that used by Zaluska et al. Therefore, differences in the desorption kinetics between our work and in Zaluska et al. are not due to different purity of the material. This problem requires more thorough investigations. In further developments Sandrock et al. [1, 75] optimized the amount of Ti catalyst needed to effectively increase the rates of absorption/desorption and introduced

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211

a dry preparation technique of dispersing the chloride catalyst precursor by the ball milling of NaAlH4 + TiCl3 mixtures (0–9 mol%TiCl3). It must be pointed out that the NaAlH4 was not technically a pure compound but was a highly purified material made by cryopumping tetrahydrofuran (THF) from a solution of NaAlH4 in THF and then vacuum drying. After 3 h of processing the mixture of alanate plus catalyst in a high-energy SPEX® mill they observed a remarkable improvement in desorption/absorption kinetics. With the addition of 2–4 mol%TiCl3 the alanate was able to desorb ~ 3 wt%H2 within ~ 1 h at 125°C under 0.13 MPa pressure (atmospheric). Rehydrogenation of dehydrided NaAlH4 (NaH + Al) at 125°C under 8.8–7.9 MPa containing 2 and 4 mol%TiCl3 resulted in the absorption of ~ 4–4.5 wt%H2 within ~ 0.75–1.5 h. In general, the authors observed a linear decrease of total desorption/ absorption capacity with increasing mol% of TiCl3. This effect can be explained by invoking the mechanism by means of which each TiCl3 provides metallic Ti as a catalyst. As argued by Sandrock et al. various Ti-bearing compounds such as Ti(OBu)4 or β-TiCl3 are called “catalyst precursors.” X-ray diffraction shows that after ball milling (or wet chemical processing) the TiCl3 is completely reduced by the Na in NaAlH4 to form NaCl and most likely zero-valent (metallic) Ti [75, 76]. The general solid state reaction proposed by the authors is based on the partial destruction of a small amount of NaAlH4 in the process in the following way (1−x ) NaAlH 4 + xTiCl3 → (1 −4x ) NaAlH 4 + 3xNaCl + xTi + 3xAl + 6xH 2 (3.10) where x is the mole fraction of TiCl3 added to the initial mixture before mechanical milling. This reaction clearly indicates that the metallic Ti is a catalyst and TiCl3 is only the precursor. The authors pointed out that (3.10) is not necessarily exact and it may not be the elemental Ti which is the final catalyst. It could be TiH2, a Ti-subhydride, a Ti-alloy or an intermetallic compound/hydride. The catalyst size is very small being either nanometric or even amorphous and impossible for detection by XRD. Obviously, according to the reaction of (3.10) an increase of the content of the TiCl3 precursor in the initial mixture NaAlH4 + TiCl3 will produce more “ballast” after ball milling in the form of NaCl and Al which are useless from the standpoint of hydrogenation as they do not form hydrides. As such, the total hydrogen capacity will decrease proportionally to the amount of catalyst precursor added as described above. Sandrock et al. also found that all kinetics of undoped and doped NaAlH4 followed the Arrhenius equation (Sect. 1.4.1) in a wide temperature range from 20 to 225°C. For comparison, they found that the activation energy of desorption for an undoped NaAlH4 was equal to ~ 118 kJ/molH2 (and correspondingly ~ 121 kJ/molH2 for Na3AlH6) in contrast to the doped NaAlH4 exhibiting ~ 73–80 kJ/molH2 (and correspondingly ~ 97 kJ/molH2 for Na3AlH6) for TiCl3 precursor range from 0.9 to 6 mol%. The activation energy seems to be stabilized after initial drop for 0.9 mol%TiCl3 when more TiCl3 is added to NaAlH4. Fichtner et al. [77, 78] added small Ti clusters (Ti13·6THF; the synthesis of these clusters is described in [77]) to highly purified NaAlH4 by ball milling for 30 min and observed a substantial improvement of desorption/absorption kinetics for an alanate doped with 2 mol% of Ti clusters. A doped alanate desorbed on

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3 Complex Hydrides

the average ~ 4 wt%H2 under 0.5–0.6 bar (slightly below atmospheric pressure) at 150°C. The first step of absorption from NaH + Al to Na3AlH6 (the reverse reaction R2 of (3.8) but in solid state) was able to occur at 100°C under 20 bar H2 pressure within about ~ 3 h. The second step to NaAlH4 ((R1b) of (3.7)) usually occurred at 100°C under 100 bar H2 pressure within another 3–4 h duration. It is to be pointed out that these results are not much superior to those already reported by Sandrock et al. [1, 75]. Wang and Jensen [79] used a commercial Ti powder which they added to a highly purified NaAlH4 via ball milling for 1–10 h. At 150°C under <1 Torr H2 pressure (vacuum < 0.001 atm.), the mixture desorbed in the first desorption ~ 3.5 wt%H2 within 1–3 h depending on the milling time. They were able to recharge this material at 120°C under ~ 12 MPa H2 within 8–10 h. After cycling, the Ti-doped NaAlH4 lost capacity to ~ 2.8 wt%H2 after eight cycles. Quite recently, Pukazhselvan et al. [80] added mischmetal (Mm) as a catalyst to highly purified NaAlH4 by ball milling. The alanate doped with 2 mol% Mm was continually desorbing up to 200 min at ~ 150°C (rapid heating from room temperature) under ~ 1 atm. pressure reaching finally about 5 wt%H2. The same alanate doped with 2 mol%Ti desorbed under the same conditions exhibiting a pronounced plateau which according to the authors corresponded to the first stage of desorption, R1b, (in solid state) of (3.7). However, after about 200 min the final amount of hydrogen was almost identical to the one desorbed by Mm doped alanate (~ 5 wt%). It is to be pointed out that Wang and Jensen [79] who used a commercial Ti powder as a dopant did not observed any unusual shape of the desorption curve (plateau). Jun Wang et al. [81, 82] investigated addition of chlorides of Ti, Zr and Fe as well as Ti and graphite as a co-dopant to a highly purified NaAlH4 by ball milling. Samples doped with metal chlorides were able to desorb at 90–130°C in TGA (thermogravimetric analyzer) under atmospheric pressure of nitrogen but the desorbed amounts of H2 did not exceed 2 wt% after 40 and 10 min at 110 and 130°C, respectively. Samples doped with graphite as a co-dopant to Ti desorbed under atmospheric pressure no more than 2–2.5 wt%H2 at 90 and 110°C. The first report on rehydrogenation to NaAlH4 from NaH and Al by ball milling and subsequent annealing under hydrogen pressure was made by Zaluska et al. [83]. After ball milling and nanostructuring the mixture was subjected to 90 bar H2 at 130°C resulting in the formation of NaAlH4. In addition, Zaluski et al. [71] reported the synthesis of Na3AlH6 using ball milling of the mixture NaH + NaAlH4. During subsequent continuous heating in a DSC experiment only two endothermic peaks were observed corresponding to decomposition of Na3AlH6 (centered at ~ 300°C) and NaH (centered at ~ 400°C). Very recently, Eigen et al. [84] synthesized NaAlH4 doped with TiCl4 (less expensive than TiCl3) by ball milling of NaH and Al (plus TiCl4) under approximately 6 bar of H2 pressure. The material synthesized for 240 h desorbed about 4 wt%H2 at 125°C under vacuum. It seems that this methodology is not commercially viable due to the very long milling times required. In summary, one can say that NaAlH4 is one of the most extensively researched hydride. Although the desorption/absorption temperatures and time durations for a

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213

catalyzed compound are quite close to the D.O.E. targets (Table 1.2 in Sect. 1.1) the hydrogen capacities are far too low for any mobile application. In stationary applications most probably MmNi5-type is still by comparison far superior to NaAlH4 for any applications, for example, in small electronic devices. It must also be pointed out that the observed temperatures/times of desorption/absorption for catalyzed NaAlH4 discussed above were obtained for the highly purified alanate and mostly under vacuum conditions for desorption. It is not clear if such low desorption temperatures and relatively short desorption/absorption times can be also obtained for catalyzed but unpurified, commercial NaAlH4 under atmospheric pressure for desorption.

3.2.2

LiAlH4

The results of the research on LiAlH4 have been reviewed in a number of review articles (for reviews see [1–7]). This is the easily commercially available alanate with the highest theoretical storage capacity which equals 10.6 wt%H2 (Table 1.4 in Sect. 1.1). Such a high capacity makes it a potentially very attractive hydride for solid state hydrogen storage. It is well established in the literature [67, 70, 85–89] that hydrogen desorbs from a pure, uncatalyzed LiAlH4 hydride in a three-step decomposition very similar to that of NaAlH4, the first of which goes through the melting of LiAlH4 (R1a) LiAlH 4 (s) → LiAlH 4 (l)

(3.11)

(R1b) LiAlH 4 (l) → 1/3Li3 AlH 6 (s) + 2/3Al(s) + H 2 (g)

(3.12)

(R2) 1/3Li3 AlH 6 (s) → LiH + 1/3Al + 0.5H 2

(3.13)

(R3) LiH → Li + 0.5H 2

(3.14)

where s-solid, l-liquid and g-gas. R1a is endothermic, R1b is exothermic, R2 and R3 are both endothermic reactions. (R1b), (R2) and (R3) proceed with a theoretical hydrogen release of 5.3, 2.6 and 2.6 wt%, respectively. One must keep in mind that these numbers will be lower for purity-corrected capacity. In thermal analysis the reaction (R1a,b) of (3.11) and (3.12) occurs around 112– 220°C, (R2) takes place around 127–260°C and (R3) occurs at too high temperatures (400–450°C) to be practical for hydrogen storage applications [67]. Hence, the useful theoretical capacity accessible below, say, ~ 250°C is about 7.9 wt%H2 from the reactions (R1b) and (R2) of (3.12) and (3.13). This capacity is slightly larger than MgH2 and accessible at much lower desorption temperature range than that of MgH2 (Chap. 2). Figure 3.9a shows a typical DSC trace at the heating rate of 10°C/min obtained in our laboratory from as-received LiAlH4 (purity 97%). It is clearly seen that (R1a) is proceeded by an exothermic peak centered at 152.5°C. In the literature, this first

214

3 Complex Hydrides R1a

DSC / mW/mg ↓ exo

[1] 172.0⬚C

2

R3 [1] 438.1⬚C

R2

0

[1] 242.7⬚C

hydroxyl

−2

[1] 152.5⬚C

−4

R1b −6

[1] 193.9⬚C

50

100

150

250 300 Temperature / ⬚C

200

a DSC / mW/mg ↓ exo 2

350

400

450

[1] 171.4⬚C

1 [1] 447.0⬚C

0

[1] 222.9⬚C

−1 −2

[1] 137.2⬚C

−3 [1] 182.1⬚C

50

b

100

150

200

250 300 Temperature / ⬚C

350

400

450

Fig. 3.9 (a) DSC trace of as-received, undoped LiAlH4 (97% purity) and (b) the same hydride after milling in the magneto mill Uni-Ball-Mill 5 under HES57 mode (two magnets at 5 and 7 o’clock positions) for 20 h in argon. DSC heating rate 10°C/min at argon flow 50 ml/min

exotherm is usually assigned to the interaction of LiAlH4 with hydroxyl impurities [67]. In the next stage LiAlH4 melts showing an endothermic melting peak (R1a) centered at 172°C, and subsequently exothermally decomposes and crystallizes to form solid Li3AlH6 which gives the (R1b) peak centered at 193.9°C. This is immediately followed by another endothermic reaction (R2) centered at 242.7°C. The decomposition of LiH (R3) is endothermic with the maximum at 438.1°C. This temperature is much lower than about 680°C, claimed to be the decomposition temperature of LiH by Zaluski et al. [71]. The thermal behavior of pure LiAlH4 is almost exactly the same as has been observed for pure NaAlH4 in Fig. 3.7a with the exception that the reaction (R1b) for pure NaAlH4 is endothermic and both (R1b)

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215

and (R2) occur at much higher peak temperatures than those for LiAlH4 if one compares Figs. 3.7a and 3.9a. There is some limited experimental evidence that upon continuous heating at very low heating rates, on the order of 0.5°C/min, the melting (R1a) of the deuterided LiAlD4 does not occur [87]. Wiench et al. [90] used variable temperature (VT) in situ nuclear magnetic resonance (NMR) to study an isothermal decomposition of LiAlH4. After 30 min of exposure to 100°C, only trace amounts of metallic Al and Li3AlH3 (reaction (R1b)) were detected which means that practically no decomposition occurred. At 150°C the decomposition process was accelerated and after 2 h of exposure no LiAlH4 could be detected. Since the temperature of 150°C is much lower than the melting point in Fig. 3.9a, the decomposition had to occur without melting. Strangely enough, Wiench et al. also found that the concentration of Li3AlH6 in the hydride sample after exposure to 150°C for ~ 2 h was equal to ~ 18 mol% instead of an equilibrium value of 33.3 mol% if the reaction proceeded according to the reaction (R1b) of (3.12). On that basis they concluded that the alternative reaction path during isothermal heating was LiAlH 4 → LiH + Al + 3/2H 2

(3.15)

Also, Andreasen et al. [86] reported that as-received LiAlH4 (~ 97% purity) desorbed isothermally about 5 wt%H2 even at such a low temperature as 115°C after 20 h. At higher temperatures in the range of 130–150°C, alanate desorbed ~ 6 wt% within approximately 1.2–2.5 h. From the shape of the desorption curves it was clear that the decomposition of LiAlH4 in solid state occurs according to a twostep mechanism given by the reaction (R1b) and (R2). It must be noted that the isothermal decomposition of LiAlH4 without melting as reported in [86] and [90] is in striking contrast with the reported lack of isothermal decomposition of NaAlH4 even at 150°C [69]. From the kinetic measurements under isothermal conditions Andreasen et al. [86] calculated the activation energies of desorption arriving at 82 ± 4 kJ/mol for the reaction (R1b) in solid state, and 90 ± 4 kJ/mol for the reaction (R2). From the Kissinger analysis at the heating rates 2, 3, 4 and 5°C/min, Andreasen [67] obtained 81 ± 4 kJ/mol for the reaction (R1b) from the liquid and 108 ± 8 kJ/mol for the reaction (R2). Blanchard et al. [91] reported the activation energy of 102 kJ/mol for the main desorption stage of the deuteride LiAlD4. On the other hand, from the Kissinger analysis shown in Fig. 1.28 in Sect. 1.4.1, we calculated that the activation energy of decomposition for the reaction (R1b) from the liquid equals −122 kJ/mol. This is slightly higher than the value obtained by Andreasen. Figure 3.10a and b show the Kissinger analysis for the reaction (R2) and (R3) of pure LiAlH4 obtained in our laboratory which yields 153 and 158 kJ/mol, respectively. Again, our value of activation energy for (R2) is higher than that obtained by Andreasen. Another factor which suppresses melting of LiAlH4 is catalysts. Andreasen [67] used 1 min ball milling to disperse 2 mol%TiCl3·1/3AlCl3 in LiAlH4. DSC experiments with the heating rates 3–5°C showed only two endothermic reactions, the

3 Complex Hydrides

ln(b/T2)

216 −10.1 −10.2 −10.3 −10.4 −10.5 −10.6 −10.7 −10.8 −10.9 −11 −11.1 −11.2 1.9

y = −18.382x + 24.943 R2 = 0.9451

EA = 153 kJ/mol

1.91

1.92

ln(b/T2)

a

b

−10.8 −10.9 −11 −11.1 −11.2 −11.3 −11.4 −11.5 −11.6 −11.7 −11.8 −11.9 1.34

1.93 1000/T (K-1)

1.94

1.95

1.96

1.39

1.4

y = −19.01x + 14.719 R2 = 0.9719

EA = 158 kJ/mol

1.35

1.36

1.37

1.38

1000/T (K-1)

Fig. 3.10 Kissinger analysis of the activation energy, EA, of decomposition for the reaction (a) (R2) and (b) (R3) for pure LiAlH4

first one with the maximum around 150°C assigned to the dehydrogenation of solid LiAlH4 into Li3AlH6 according to modified reaction (R1b) where l-liquid should be replaced by s-solid, and the second one with the maximum around 200°C assigned to the decomposition of solid Li3AlH6 according to the reaction (R2) of (3.13). In a DSC experiment the metal chloride catalytic doping apparently lowers the first decomposition reaction of LiAlH4 below its melting point practically eliminating melting. Also the second decomposition reaction of Li3AlH6 seems to be slightly shifted to lower temperatures by comparison with (R2) in Fig. 3.9a. In our laboratory, we have studied the effect of ball milling on the thermal decomposition of LiAlH4 in DSC experiments. The morphologies of LiAlH4 powder as received and after milling in the magneto mill Uni-Ball-Mill 5 under HES57 mode (two magnets at 5 and 7 o’clock positions) for 20 h in argon are shown in Fig. 3.11a, b, respectively. The average ECD particle size with standard deviation (measured as described in Sect. 1.4.3) of the as-received powder equals 10.5 ± 4.8 μm and the particles have very sharp edges. After ball milling the average ECD is reduced to 5.2 ± 4.3 μm but smaller particles have a tendency to agglomerate into larger aggregates in a very similar fashion as observed for NaAlH4

3.2

Alanates

a

217

b

Fig. 3.11 (a) SEM picture (secondary electrons) of as-received LiAlH4 powder and (b) the same powder after milling in the magneto mill Uni-Ball-Mill 5 under HES57 mode (two magnets at 5 and 7 o’clock positions) for 20 h in argon

(Fig. 3.6b). It must be mentioned that both alanates NaAlH4 and LiAlH4 are extremely difficult to be effectively refined during ball milling. As can be seen the particle size reduction of LiAlH4 is only twofold after 20 h of milling. This can be compared with almost 40-fold particle size reduction of MgH2 milled for a much shorter time (Fig. 2.15 in Sect. 2.1.3). Chen et al. [85] reported a substantial grain size refinement of LiAlH4 doped with TiCl3·1/3AlCl3 after milling for 0.5–1 h. However, their result is in contradiction to the value of the average grain size of ~ 58 nm, which we measured in our experiments for LiAlH4 after ball milling. In turn, our result agrees well with Andreasen et al. [86] who reported the average grain size of ~ 50 nm for LiAlH4 after 6–10 h of milling in a Retsch PM 100 planetary mill. It is rather hard to explain why the addition of a small amount of TiCl3·1/3AlCl3 would enhance the reduction of grain size. This problem needs further studies. As shown in Fig. 3.9b severe ball milling of LiAlH4 for 20 h carried out in our laboratory does not change in any substantive manner, within an experimental scatter, the positions of DSC peaks as compared to the as-received alanate in Fig. 3.9a. Only the hydroxyl reaction peak and both (R1b) and (R2) peaks are slightly shifted to lower temperatures. In contrast, Resan et al. [92] reported that milling of undoped LiAlH4 in SPEX 8000M resulted in a decrease in the temperature of each step of desorption (R1b and R2) and a decrease in the amount of hydrogen released at each desorption step. During a continuous heating in a Temperature Programmed Desorption (TPD) experiment the total amount of hydrogen released decreased from 6.3 to 5.8 wt% upon milling for 5 min and to 5.0 wt% after milling for 60 min. The temperature of each hydrogen release step was decreased by approximately 20°C after 5 min of milling and by 30°C after 60 min of milling. On a DSC trace the temperature of exothermic and endothermic peaks were also lowered compared to the unmilled samples. These observations are in contradiction to the DSC trace (in Fig. 3.9) from our laboratory which shows that if there is an effect of milling it is really quite minor. Resan et al. concluded that undoped/pure LiAlH4 decomposes

218

3 Complex Hydrides

during milling in contrast to undoped NaAlH4 which after milling under the same conditions did not show any loss in capacity. The effect of ball milling on the isothermal desorption kinetics of pure LiAlH4 was investigated by Andreasen et al. [86]. After 6 h of milling they observed the appearance of additional peak on XRD pattern due to the formation of either monoclinic or rhombohedral Li3AlH6 which suggests a partial decomposition of pure LiAlH4 during milling. However, interestingly enough no peaks due to metallic Al were observed as required by the reaction (R1b) in solid state. They suggested that the lack of Al peaks is due to its amorphous structure. The most likely cause of partial decomposition of pure/undoped LiAlH4 during milling which was put forward by Andreasen et al. is the increase of the milling vial temperature upon milling. This effect is similar to the one observed by Resan et al. [92] and discussed above. It is quite possible that during milling in a high-energy SPEX 8000 mill by Resan et al. there was a substantial increase in temperature unnoticed by the authors which led to a partial decomposition of pure/undoped LiAlH4. It seems that controlling the milling conditions is very crucial to a full understanding of the behavior of LiAlH4 (and probably NaAlH4) during milling. Largely controlled milling conditions can be only achieved in the magneto-mill Uni-Ball-Mill 5 described in Sect. 1.3.2. Andreasen et al. [86] also found that ball milling increased the rate constant, k, in the JMAK equation (Sect. 1.4.1), of reaction (R1b) in solid state but virtually had no effect on the rate constant of reaction (R2). They also showed that the reaction constant, k, of reaction (R1b) in solid state increases with decreasing grain size of ball-milled LiAlH4 within the range 150–50 nm. Andreasen et al. concluded that the reaction (R1b) in solid state is limited by a mass transfer process, e.g., long range atomic diffusion of Al while the reaction (R2) is limited by the intrinsic kinetics (too low a temperature of decomposition). In conclusion, one must say that ball milling alone is not sufficient to improve the kinetics of reaction (R2). A solution to improvement of the kinetics of reaction (R2) could be a suitable catalytic additive. A number of metal chlorides such as TiCl3, TiCl3·1/3AlCl3, VCl3, TiCl4 were added to LiAlH4 as catalysts by ball milling with the goal to improve dehydrogenation/hydrogenation properties. However, mechanical milling of metal chlorides with LiAlH4 is rather tricky because dopants facilitate decomposition of LiAlH4 and its deuteride LiAlD4 during ball milling even if milling takes only a few minutes [67, 88, 89, 91, 93]. However, there is one paper by Chen et al. [85] where they reported that doping LiAlH4 with 2 mol%TiCl3·1/3AlCl3 by agitating in a Fuji Auto mixer for 1 h did not cause it to decompose. Once again, uncontrolled milling conditions in most of ball mills used in various experiments can lead to such contradicting final results. Some light was recently shed on this important problem by Blanchard et al. [94] who attempted to control milling energy by changing gyration rates in a Pulverisette 6 ball mill during milling of LiAlD4 doped with VCl3. They showed that at any milling energy applied VCl3 was reduced to Li–V–Cl metastable phases, LiCl and free Al and V or Al–V phases. Even under mild milling conditions, at or close to room temperature, the two first decomposition steps of

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Alanates

219

LiAlD4 enhanced by VCl3 occurred during the first week after milling at temperature in the range of 20–50°C. According to Chen et al. [85], Andreasen [67] and Blanchard et al. [91] milling reduces Ti3+ in TiCl3 according to the following reaction 49LiAlH 4 + TiCl3 ⋅ 1/3AlCl 3 → 46LiAlH 4 + Ti + 3Al + 3LiCl + 1/3AlCl 3 + 6H 2

(3.16)

The reduced Ti species are possibly distributed as metallic Ti across the surface of the powder particles or they may even be present as TixAly intermetallic compound [89, 92]. The above reaction is in essence almost identical to the reaction of (3.10) for NaAlH4 catalyzed with TiCl3. The effect of catalytic metal chloride additives on the kinetics of isothermal decomposition of LiAlH4 in a Sieverts-type apparatus has been studied by a few research groups and the results seem to be rather contradictory. On the one hand, Chen et al. [85] found from the Arrhenius plot that the activation energy of desorption of LiAlH4 catalyzed with 2 mol%TiCl3·1/3AlCl3 was 42.6 kJ/mol for the reaction (R1b) in solid state and 54.8 kJ/mol for the reaction (R2). Indeed, the desorption curves of TiCl3·1/3AlCl3 doped LiAlH4 showed quite fast kinetics in two stages (R1b) and (R2) of desorption. Undoped LiAlH4 did not desorb at 100°C whereas doped alanate desorbed ~ 3 wt%H2 within 100 min. At 125°C the desorbed amount of H2 was ~ 4 wt%H2 and eventually ~ 5.5 wt%H2 at 175°C within 50 min. It must be pointed out that the addition of metallic Ti and TiH2 hydride instead of TiCl3·1/3AlCl3 did not improve dehydrogenation/hydrogenation properties. The authors argued that some multiple core composite nanoparticles containing Ti ↔ Ti3+ defect sites are formed due to the reaction of (3.16) and then they react with each other and with hydrogen which is enhanced by the existing defects sites. However, this model still does not explain the lack of catalytic effect of metallic Ti. Also, Wang et al. [95] observed substantial enhancement of desorption kinetics of LiAlH4 doped with 0.5 mol%TiCl3 and ball milled for 20 min. At a really low temperature of 90°C the mixture released 3 wt%H2 within 30 min and 4 wt%H2 within 150 min (note that the second stage of isothermal desorption governed by the reaction (R2) is always slower than that of the reaction (R1b) in a solid state). For comparison, the NaAlH2 doped with 4 mol%TiCl3 and ball milled for 120 min released only 0.5 wt%H2 within 150 min and Mg(AlH4)2 doped with 1 mol%TiCl3 and ball milled for 15 min released less than 1.5 wt%H2. On the other hand, Andreasen [67] found from the Kissinger analysis that the apparent activation energies of decomposition of LiAlH4 doped with 2 mol%TiCl3·1/3AlCl3 are equal to 89 ± 9 and 103 ± 1 kJ/mol for the reaction (R1b) in a solid-state and (R2), respectively. The apparent activation energies found are very close to those for the pure/undoped sample equal to 81 ± 4 kJ/mol ((R1b) from the liquid) and 108 ± 8 kJ/mol (R2) as discussed earlier in this section. Furthermore, the apparent activation energies of doped LiAlH4 found by Andreasen are almost twice as high as those found for a doped LiAlH4 by Chen et al. [85]. It is possible that the observed discrepancies in the apparent activation energy of desorption of

220

3 Complex Hydrides

doped LiAlH4 are owing to various milling conditions used by different research groups which in effect can lead to various uncontrolled degrees of decomposition of doped LiAlH4 during milling. As mentioned earlier, controlled milling conditions should be applied to investigate this problem. In order to establish the catalytic activity of various compounds Kojima et al. [96] investigated the amount of hydrogen released during ball milling for 5 min to 24 h of doped LiAlH4. They found that the amount of desorbed of hydrogen decreased and hence, the activity increases in the following order of chloride catalysts TiCl3 > ZrCl4 > VCl3 > NiCl2 > ZnCl2. The most active seem to be TiCl3 and ZrCl4. They also found that nano-Ni is an excellent catalyst which substantially enhances desorption during milling at room temperature. They proposed the following reactions for the investigated catalysts with LiAlH4: 4LiAlH 4 + ZrCl 4 → Al3 Zr + Al + 4LiCl + 8H 2

(3.17)

3LiAlH 4 + VCl3 → Al3 V + 3LiCl + 6H 2

(3.18)

3LiAlH 4 + NiCl 2 → Al3 Ni + LiH + 2LiCl + 5.5H 2

(3.19)

Apparently, according to Kojima et al. intermetallic trialuminide phases are formed which act as effective catalysts rather than pure zero-valent Ti as argued by Chen et al. [85] but in accord with the suggestions in [89, 92]. Brinks et al. [97] ball-milled LiAlD4 with the TiF3 additive and concluded that there were no signs of any Ti-containing phases, and the unit-cell of LiAlD4 and Al gave no indication of any solid solutions. Hence, they eventually concluded that the Ti was in an amorphous state directly after ball milling. In addition, they observed that samples stored in a glove box were slowly desorbed, and after 6 months for a LiAlD4 + TiF3 sample, the reaction to LiD + Al was nearly finished. Finally, an important issue of the reversibility of LiAlH4 should be briefly discussed. In contrast to NaAlH4 there is no report of any successful rehydrogenation of dehydrogenated LiAlH4 (a mixture of LiH and Al) back to Li3AlH6 and LiAlH4 (reverse of (R2) and (R1b)). Wang et al. [95] tried to rehydrogenate under ~ 8 MPa of hydrogen pressure at 125°C the LiAlH4 doped with TiCl3 which was discharged of hydrogen at 125°C for 16 h under ~ 0.35 MPa. No successful hydrogenation occurred. In contrast, NaAlH4 was easily rehydrogenated under same conditions. Therefore, it is important to understand whether the difficulty in a rehydrogenation of LiAlH4 arises from a thermodynamic reason or from a kinetics reason. The early information on the thermodynamic properties of LiAlH4 and Li3AlH6 comes from Claudy et al. [98] who reported experimentally measured values of the heat of formation of LiAlH4 and Li3AlH6 as −107.2 ± 0.84 and −311.1 ± 1.26 kJ/mol, respectively. These values were later confirmed by a theoretical work of Løvvik et al. [99] who used density-functional theory (DFT) for predicting the enthalpies of formation at 298 K for LiAlH4 and Li3AlH6 to be equal to −113.42 and −310.9 kJ/mol, respectively. They also predicted the enthalpies of decomposition of LiAlH4 to form Li3AlH6 (the reaction (R1b)) and

3.2

Alanates

221

the decomposition of Li3AlD6 to form LiD (the reaction (R2)) to be endothermic and equal to 9.79 and 15.72 kJ/mol at 298 K, respectively. They concluded that their results suggest that LiAlH4 is after all a stable phase. Brinks et al. [97] showed using Sieverts-type measurements that the plateau pressure for a sample of LiAlD4 + 2 mol%TiF3 at 80°C was higher than 85 bar. Jang et al. [100] used the CALPHAD method to compute phase equilibria between Li3AlH6 and LiAlH4 as a function of temperature and hydrogen partial pressure. The aim of the computation was to check whether the hydrogen absorption/desorption reaction between Li3AlH6 and LiAlH4 would be thermodynamically possible in a practically accessible temperature and hydrogen pressure range. Figure 3.12 shows the stability diagram of LiH, Li3AlH6 and LiAlH4 computed using two different sets of thermodynamic parameters, in a form of the Van’t Hoff plot. Two sets of data adopted for computations from two references are represented by either solid curves or by dotted curves. The two lines of each calculation represent temperature-hydrogen pressure equilibrium conditions between two pairs of hydrides LiH/Li3AlH6 and Li3AlH6/LiAlH4. It is very clear that the reversible reaction Li3AlH6 → LiAlH4 requires a partial hydrogen pressure well above 103 bar at room temperature (298 K). In other words, in practical terms such a reaction is inaccessible. Although it seems that LiAlH 4 does not exhibit a chemical reversibility (at temperature and under hydrogen pressure) still it appears to be prone to a

106 10

Chen et al. (2001) Brinks et al. (2005)

5

LiAIH4

Hydrogen Pressure, bar

104 10

3

Li3AIH6

102 101 100 LiH −1

10

10−2 2.0

2.5 1000/T, K−1

3.0

3.5

Fig. 3.12 Computed stability diagram of LiH, Li3AlH6 and LiAlH4, in comparison with experimental data [85, 97] for the LiH/Li3AlH6 equilibrium. The solid curves are those calculated using the parameter Set 1 obtained by fitting to Chen’s et al. data [85] while the dotted curves are using the parameter Set 2 obtained by fitting to Brinks et al.’s data [97] (after [100])

222

3 Complex Hydrides

mechanical reversibility where rehydrogenation occurs during mechanical milling at the hydrogen atmosphere in the presence of catalysts. A general concept of mecha-nical reversibility in hydrogen alloying mills was recently discussed by Wronski et al. [101]. Wang et al. [102] developed a mechanochemical method to rehydrogenate LiAlH4 mixed with the catalytic precursor TiCl3. This sample was milled in a vial containing hydrogen at the pressures ranging from 4.5 to 97.5 bar using a SPEX 8000 shaker mill for 20 min. The ball-milled sample was subjected to dehydrogenation at 90°C for 5 h and afterwards again ball milled for 2 h at the hydrogen atmosphere at the pressures ranging from 4.5 to 97.5 bar. Then tetrahydrofuran (THF) was added to the material and the mixture was additionally milled for 2 h at the same pressure ranges. The resulting mixture was vacuum filtered through 0.7 μm filter paper and vacuum dried to collect the rehydrogenated LiAlH4. The remaining residue could be used as a dopant for the next cycle. In essence, the method consists of five general cycles: catalyst dispersion by ball milling, dehydrogenation, rehydrogenation, vacuum filtration, vacuum-drying and then catalyst redispersion. The last step is the first step of the next cycle. Wang et al. pointed out that the Ti catalyst was ineffective in the absence of THF, suggesting that the Ti was somehow acting on the LiAlH4·4THF adduct which was formed during milling with THF. Also, the largest rehydrogenation conversion yield (~ 90%) was observed when the hydrogen pressure was at least equal to 60 bar but no further gain was observed when the pressure was much above 60 bar. Nevertheless, at 4.5 bar the rehydrogenation yield was still ~ 44%. In another recent development Kojima et al. [103] mechanically milled LiH and Al without and with the TiCl3 additive for 24 h in a H2 gas atmosphere at a pressure of 1 MPa at room temperature. They found that a small amount of LiAlH4 could be directly synthesized by the mechanochemical reaction with concomitant formation of Li3AlH6. The latter can be relatively easily formed by mechanochemical synthesis of LiAlH4 and LiH as originally reported by Zaluski et al. [71] and later by Balema et al. [104]. LiAlH4 is a very interesting complex hydride which deserves more research efforts with the aim of making its storage properties even more compatible with the D.O.E. targets (Sect. 1.1). Doping with metal chloride additives which are dispersed by ball milling substantially improves its desorption properties to the extent that 4–5 wt%H2 could be obtained at the range of temperature between 90 and 100°C. This temperature range is very close to the temperature generated by the waste heat of a PEM fuel cell. Further research efforts focused, in particular, on the ball milling aspects and search for more advanced catalytic additives could most likely improve its hydrogen desorption properties even more. The only drawback remains its chemical irreversibility. However if on board recharging requirement can be relaxed then research efforts could also be directed towards the studies of its off board rechargeability by using the concept of “mechanical reversibility” by ball milling.

3.2

Alanates

3.2.3

223

Mg(AlH4)2 and Ca(AlH4)2

The theoretical gravimetric hydrogen capacity of Mg(AlH4)2 is 9.3 wt% (Table 1.4 in Sect. 1.1) and the hydride is based on two inexpensive light metals Mg and Al. Its theoretical reversible gravimetric hydrogen capacity is ~ 7.0 wt% which is lower than that of LiAlH4. The preparation of Mg(AlH4)2 was first reported by Wiberg and Bauer [105] who used solvent mediated metathesis reactions of MgH2 and AlCl3, MgH2 and AlH3, and LiAlH4 and MgBr2. Ashby et al. [106] investigated the reactions of lithium and sodium aluminum hydrides with magnesium halides in ether solvents to prepare Mg(AlH4)2. Soon after, the thermal decomposition of Mg(AlH4)2 was studied by Dilts and Ashby [70] and Claudy et al. [107]. More recently, the chemical synthesis of Mg(AlH4)2 was replicated by Fichtner et al. [108–111] who used a chemical reaction of solvent-free sodium alanate (NaAlH4) with magnesium chloride (MgCl2). Fichtner et al. [108–111], Schwarz et al. [112] and Fossdal et al. [113, 114] investigated the dehydrogenation of Mg(AlH4)2 synthesized by solvent mediated metathesis reaction and found that it decomposed in two major steps with the first step around 160°C. This relatively low dehydrogenation temperature of the first step is attractive enough to warrant further studies. According to the observations the decomposition of Mg(AlH4)2 occurs in two steps, at two different temperature ranges Mg(AlH 4 )2 → MgH 2 + 2Al + 3H 2 (110 − 200° C)

(3.20a)

MgH 2 → Mg + H 2 (240 − 280° C)

(3.20b)

In the first step of the reaction of (3.20a) the theoretical amount of desorbed hydrogen is 7.0 wt% and the remaining 2.3 wt%H2 is desorbed in the second step (3.20b). In the first step of desorption Fichtner et al. [109] reported the amount of ~ 6.6 wt%H2 (i.e., about 73% of the total 9 wt%H2) desorbed in a thermogravimetric analysis experiment at around 160–200°C range. Since the second step given by the reaction of (3.20b) is simply related to the hydrogen desorption from MgH2 (Chap. 2), its desorption temperature range is too high to be accessible for any practical applications in supplying a PEM fuel cell. An alternative indirect method of the synthesis of Mg(AlH4)2 is the mechanochemically activated metathesis reaction which was first reported by Dymova et al. [115, 116] according to the following reactions between MgH2, AlH3 and AlCl3 during ball milling

or

3MgH 2 + 2AlCl3 → 2AlH 3 + 3MgCl 2

(3.21a)

MgH 2 + 2AlH3 → Mg(AlH 4 )2

(3.21b)

2MgH 2 + AlCl3 → 0.5Mg(AlH 4 )2 + 1.5MgCl 2

(3.22)

224

3 Complex Hydrides

This concept was followed by Mamatha et al. [117, 118] who used a slightly different path of reaction using a mixture of NaAlH4 and MgCl2 which was ball milled for 3 h (in a Retsch ball mill) and obtained a mixture of synthesized Mg(AlH4)2 and NaCl formed according to the following reaction occurring during ball milling 2NaAlH 4 + MgCl 2 → Mg(AlH 4 )2 + 2NaCl

(3.23)

Subsequent DSC analysis of the mechano-chemically prepared powders showed three endothermic peaks with the maxima around 140, 270 and 450°C, which were interpreted as corresponding to the decomposition of Mg(AlH4)2 into MgH2, Al and H2, decomposition of MgH2 into Mg and formation of Al3Mg2, and eventually melting of Al3Mg2 and Al mixture, respectively. Kim et al. [119] synthesized a mixture of Mg(AlH4)2 and NaCl after 1 h of milling in a SPEX 8000 ball mill. DSC test was carried out only up to slightly over 300°C and showed the first exothermic peak at the range of 115–150°C and the second endothermic peak at the range of 240–290°C. They concluded that the first exothermic reaction corresponds to the decomposition of Mg(AlH4)2 into MgH2, Al and H2 and the second endothermic peak to the decomposition of MgH2. Apparently, there is a clear controversy between the results of Mamatha et al. [117, 118] who report endothermic and those of Kim et al. [119] who report exothermic reaction for the decomposition of Mg(AlH4)2. In our laboratory we also investigated the mechanochemical activation synthesis (MCAS) of Mg(AlH4)2 [120] with the first goal of finding the effect of the milling time longer than 3 h on the microstructural development of the initial mixture of NaAlH4 and MgCl2 and the second goal of clarifying the thermal events occurring during DSC analysis of the synthesized mixture of Mg(AlH4)2 and NaCl with a special emphasis on the heat flow event in the range of 115–150°C. Figure 3.13 shows the evolution of XRD patterns as a function of milling time during MCAS. It is seen that milling for 5 and 10 h results in the formation of the [Mg(AlH4)2 + 2NaCl] mixture according to the reaction of (3.23). The calculated grain size of Mg(AlH4)2 hydride in the mixture is on the order of 18 nm. Nevertheless, after relatively short milling for just 10 h a partial decomposition of the initially formed Mg(AlH4)2 into the nanocrystalline β-MgH2, elemental Al (grain size ~ 26 nm) and hydrogen gas is observed. Prolonged milling for 40 h results in a complete decomposition of the Mg(AlH4)2 into two nanocrystalline solid phases such as β-MgH2 (calculated grain size ~ 11 nm) and the elemental Al (grain size ~ 20 nm), and hydrogen gas. This indicates that Mg(AlH4)2 is unstable under prolonged milling. Figure 3.14 shows a DSC trace obtained at the scan rate of 4°C/min from the mixtures milled for 5, 10, and 40 h. A very small thermal flow effect of either exothermic or endothermic nature is seen around 140°C at the DSC traces of mixtures milled for 5 and 10 h only. At temperatures >180°C, mixtures milled for 5 and 10 h exhibit three strong endothermic effects centered at around 271, 315 and 452°C but the one milled for 40 h shows only two endo effects at around 292 and 452°C.

3.2

Alanates

225 Mg(AlH4)2 NaCl Al b-MgH2

12000

Counts

8000

40h 4000 10h

5h 0 40

30

50 60 Degrees 2-Theta

70

80

90

Fig. 3.13 Evolution of XRD patterns as a function of milling time during MCAS of Mg(AlH4)2 from the mixture of NaAlH4 and MgCl2 DSC / mW/mg ↓ exo

[1] 452.2⬚C

0.8

[2] 451.9⬚C

0.6

0.4

[3] 451.0⬚C

1. 5h 2. 10h 3. 40h

[2] 270.7⬚C [3] 291.6⬚C [1] 270.0⬚C [1] 314.8⬚C [2] 315.5⬚C

0.2 2

3 1

0 100

150

200

250

300 350 Temperature / ⬚C

400

450

500

Fig. 3.14 DSC curves of powders after milling for 5, 10 and 40 h upon heating to 500°C registered at the scan rate of 4°C/min

We analyzed the phase composition of the Mg(AlH4)2 + 2NaCl mixture during heating in a DSC apparatus by stopping the test at various temperatures and taking powder for XRD tests. Figure 3.15 shows XRD patterns of powder milled for 10 h and subsequently heated in a DSC apparatus to the selected temperatures shown in Fig. 3.15. The XRD pattern from the sample heated up to 180°C contains only Bragg

226

3 Complex Hydrides NaCl Al(Mg) Al b-MgH2 Al3Mg2

10000

8000 425ⴗC Counts

6000

350ⴗC 4000

295ⴗC 2000 180ⴗC 0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 3.15 XRD patterns of powder milled for 10 h and then heated to 180, 295, 350 and 425°C in a DSC test in Fig. 3.14

peaks of NaCl which does not undergo any changes during heating, Al and small amount of β-MgH2. The absence of Bragg peaks of Mg(AlH4)2 shows that the thermal events seen in Fig. 3.14 around 140°C correspond to a decomposition of Mg(AlH4)2, synthesized after 10 h of milling (Fig. 3.13), into MgH2 and Al according to the reaction of (3.20a). The XRD pattern of sample heated up to 295°C contains the Bragg peaks of NaCl, the solid solution of Mg in Al designated Al(Mg), weak peaks of Al3Mg2 and very weak peaks of β-MgH2. Apparently, around the 271°C peak the majority of β-MgH2 decomposes in the endothermic reaction into free Mg and H2. Simultaneously, reactions of Mg and Al to form Al3Mg2 and Al(Mg) solid solution must have occurred. The reactions can be written as follows: β-MgH 2 + 2Al (+2NaCl) → Mg + H 2 + 2Al (+2NaCl) (225 − 340° C at 4° C/min) Mg + 2Al → 0.5Al3 Mg2 + 0.5Al(Mg) (225 − 340° C at 4° C/min)

(3.24a) (3.24b)

After DSC run up to the temperature of 350°C in Fig. 3.15, which is higher than the temperature of the second DSC peak maximum at ~ 315°C, the decomposition

3.2

Alanates

227

of the remnant MgH2 is completed and the microstructure consists of residual NaCl, large amount of Al3Mg2 intermetallic compound (high intensity XRD peaks) and large amount of Al(Mg) solid solution (high intensity XRD peaks). As such, two endothermic peaks at ~ 271 and ~ 315°C might arise either due to the decomposition of β-MgH2 and the formation of the Al3Mg2 intermetallic compound or alternatively, due to the two-step decomposition of β-MgH2. Mamatha et al. [117, 118] and Kim et al. [119] reported only a single endothermic peak with the maximum around 280–290°C for the mixture Mg(AlH4)2 + 2NaCl. Subsequently, the lattice parameters of Al(Mg) were calculated as shown in Table 3.1. It is clear that the lattice parameter of Al in the powders milled for 10 h is much larger after heating in DSC up 295, 350 and 425°C than that of the sample heated up to 180°C. The lattice expansion is consistent with the magnitude of the atomic radius of Al and Mg which is equal to 0.1432 and 0.1604 nm, respectively [121]. Fossdal et al. [114] claimed the formation of the Al(Mg) solid solution during low-temperature desorption of Mg(AlH4)2 below 180°C which is not confirmed here. Mamatha et al. [117, 118] never reported the formation of the Al(Mg) solid solution although Kim et al. [119] mentioned about increase in the lattice parameter of Al most probably due to formation of the Al(Mg) solid solution. The endo peak centered at 292°C for the mixture milled for 40 h in Fig. 3.14 is due to the decomposition of β-MgH2 and formation of Al3Mg2 and Al(Mg). It is to be pointed out that the DSC run of the 40 h milled powder at the scan rate of 4°C/min up to 350°C which is above the peak maximum at ~ 316°C in Fig. 3.14, resulted in the formation of the mixture of Al3Mg2 and Al(Mg) solid solution whose lattice parameter is also listed in Table 3.1 [120]. This clearly confirms that the formation of Al3Mg2 and Al(Mg) in the present work occurs during DSC heating from the elemental Mg and Al formed according to the reactions of (3.20a) and (3.20b). Finally, the high-temperature peak at ~ 452°C in Fig. 3.14 is due to the eutectic melting of the Al3Mg2 and Al(Mg) mixture according to the binary Mg–Al phase diagram [122]. Thermogravimetric analysis (TGA) showed [120] that the weight losses in the first and second step decomposition for the 5 h milled powder were 2.37 and 0.08 wt%, respectively. Corresponding weight losses for the 10 h milled powder were 2.14 and 0.02 wt%. Assuming the MCAS is completed (i.e., products have a 1 mol of Mg(AlH4)2 and 2 mol of NaCl), the theoretical hydrogen capacity in the mixture Table 3.1 Lattice parameters of Al and Al(Mg) solid solution in the as-milled and DSC tested samples Milling time (h)

Temperature of DSC test (°C)

10 As-milled 10 180 10 295 10 350 40 350 10 425 a Al(Mg) solid solution

Lattice parameter of Al (FCC) (nm) 0.4037 ± 0.0015 0.4039 ± 0.0015 0.4059 ± 0.0018a 0.4062 ± 0.0016a 0.4065 ± 0.0012a 0.4063 ± 0.0013a

228

3 Complex Hydrides

of Mg(AlH4)2 + 2NaCl would be ~ 3.97 wt% rather than 9.3 wt% as in a pure Mg(AlH4)2. From TGA analysis, the largest total weight loss ~ 2.45 wt%, which was observed for the 5 h milled powder, is much less than the theoretical value of ~ 3.97 wt%. A deficiency of ~ 1.5 wt% might be a combined effect of the partial decomposition of Mg(AlH4)2 during milling for 5 h and the underestimation of the weight loss in the second step (~ 270–300°C) due to the oxidation of Mg during TGA run under flowing nitrogen. The total TGA hydrogen desorption of only ~ 2.16 wt% for the 10 h milled powder is most probably due to the prior partial decomposition of Mg(AlH4)2 during milling. No weight loss observed in TGA in the 100–180°C range for the 40 h milled powder further confirms that Mg(AlH4)2 is fully decomposed after a prolonged milling time. In order to shed more light on the nature of thermal events close to 140°C in Fig. 3.14, we carried out DSC experiments at the scan rate of 4 and 20°C/min. Figure 3.16a shows the enlargement of the low temperature region up to 200°C of the DSC curve of powder ball milled for 10 h. Taking the flat DSC trace portion for the 40 h sample as the baseline (at this temperature range), one can invoke the existence of either two exothermic peaks or one endothermic peak due to the decomposition of Mg(AlH4)2. DSC traces at the scan rate of 20°C/min in Fig. 3.16b of one sample milled for 5 and one milled for 10 h, analyzed against the baseline, show the DSC humps which could be considered as endothermic events. In contrast, the other samples milled for 5 and 10 h show a diffuse trough that could be considered as an exothermic reaction. XRD analysis showed that samples heated up to 125 and 140°C still contained Mg(AlH4)2, residual NaCl, traces of β-MgH2 and the elemental Al rather than the Al(Mg) solid solution as mentioned earlier [120]. This in essence is almost the same phase content as that after ball milling for 10 h (Fig. 3.13). The presence of the traces of β-MgH2 indicates that the beginning of the decomposition of Mg(AlH4)2 starts around 125°C (at a scan rate 4°C/min). After a DSC run up to 150°C the microstructure contained only a negligible amount of Mg(AlH4)2. The absence of Mg(AlH4)2 is an indication of its complete decomposition. This analysis has clearly shown that approximately up to 180°C, Mg(AlH4)2 is already decomposed. However, the nature of this transformation as being either endothermic or exothermic could not be unambiguously established. Most probably, this situation arises owing to a very small enthalpy of decomposition of Mg(AlH4)2 to MgH2, Al and hydrogen (20a) which is cited in the literature as being equal to ~ 1.7 kJ/molH [118]. It was also predicted theoretically by Zhou [123] as being equal to 2.2 kJ/mol which agrees very well with the value obtained in [118]. It is quite possible that at such a small value of the enthalpy of decomposition of Mg(AlH4)2, which in all practical terms is almost close to zero, even small microstructural fluctuations always present in the powders synthesized by mechanochemical reactions may change the character of the decomposition reaction from endothermic to exothermic. In addition, as pointed out by Mamatha et al. [118] owing to such a small enthalpy of decomposition of Mg(AlH4)2 the thermodynamic stability of this alanate is far below that suitable for reversible hydrogen storage, roughly in the range of 20–40 kJ/mol (Sect. 1.4.1). Wang et al. [95] dehydrogenated Mg(AlH4)2 doped with 2

3.2

Alanates

229

DSC *10−2 / mW/mg 2 ↓ exo 1 0 −1 −2

[1] 142.9⬚C

baseline [1] 136.6⬚C

−3 −4

4ⴗC/min

[1] 157.2⬚C

−5 120

100

160 140 Temperature / ⬚C

180

200

DSC / mW/mg ↓ exo 0.5 0.4 0.3

[1] 162.7⬚C [3] 169.5⬚C

10h (endo)

0.2 0.1

(baseline)

20ⴗC/min

0 100

b

10h (exo)

5h (exo)

5h (endo) 120

140

160

180

200

Temperature / ⬚C

Fig. 3.16 DSC traces up to 200°C of NaAlH4 + MgCl2 mixtures ball milled for 5 and 10 h (the baseline is the trace of the mixtures milled for 40 h where Mg(AlH4)2 decomposed during milling)

mol%TiCl3 at 125°C for 16 h under 3.5 atm. hydrogen and attempted to rehydrogenate the sample at 125°C under ~ 8 MPa hydrogen pressure. No hydrogen uptake was observed. Apparently, this experiment confirms that Mg(AlH4)2 is not reversible under reasonable conditions of temperature and pressure. We tried to overcome this problem by investigating if mechanical reversibility would be possible by controlled reactive mechanical alloying (CRMA) of elemental Mg and Al powders under hydrogen in the magneto-mill Uni-Ball-Mill 5 [124]. We used three stoichiometric Mg–2Al mixtures such as (a) elemental Mg and Al powders, (b) elemental Al powder and commercial AZ91 alloy (Mg–Al–Zn alloy) and (c) powder of as-cast Mg–2Al alloy. Unfortunately, no successful synthesis of Mg(AlH4)2 has been achieved. The only nanocrystalline hydride formed up to 270 h of CRMA was β-MgH2. It has been found that there is a strong competition

230

3 Complex Hydrides

between formation of Al(Mg) solid solution and the β-MgH2 hydride occurring to a various extent up to ~ 10 h of CRMA in all three Mg–2Al mixtures. A hypothesis has been put forward that the presence of Al(Mg) solid solution inhibits the reaction of β-MgH2, Al and H2 to form Mg(AlH4)2. Furthermore, despite the fact that after prolonged milling the Al(Mg) solution eventually decomposes into secondary Al(s) (derived from solid solution), the latter retains its physico-chemical characteristics of the former solid solution which still inhibits the reaction to form Mg(AlH4)2. It is to be noted that the accelerated formation of the Al(Mg) solid solution has been observed during reactive milling of the elemental powders or pre-alloyed Mg–2Al ingots and during thermal decomposition of Mg(AlH4)2 as discussed above. This behavior has its roots in a very substantial solubility of Mg in Al vs. temperature as can be seen in the binary Mg–Al system [122]. It is to be pointed out that the major problem in the MCAS technology is the formation of a large amount of waste by product NaCl. With 2 mol of NaCl per only 1 mol of Mg(AlH4)2 (the reaction of (3.23)) the excellent theoretical hydrogen capacity ~ 9.3 wt% of pure Mg(AlH4)2 is reduced to the very inferior theoretical ~ 3.97 wt%H2 in the mixture of Mg(AlH4)2 + 2NaCl. As reported by Mamatha et al. [118] and Sterlin Leo Hudson et al. [125], it is possible to separate Mg(AlH4)2 from NaCl by the Soxhlet extraction method which is based on the suspension of the Mg(AlH4)2 + 2NaCl mixture in the diethyl ether (Et2O) solvent. However, there are a couple of serious disadvantages of this extraction method. First, the extracted Mg(AlH4)2 is usually contaminated with the solvent adduct [119]. Second, the extraction process needs several days to be completed and then the product requires some additional long-time annealing in vacuum [118]. Such a long production time may not be well suited for industrial environment. On the other hand, a solvent mediated metathesis reaction results in the formation of Mg(AlH4)2 with adducts which can substantially affect its hydrogen desorption behavior [108–111, 117–119]. A very similar alanate to Mg(AlH4)2 is its counterpart Ca(AlH4)2. However, the theoretical maximum hydrogen gravimetric capacity of Ca(AlH4)2 is 7.9 wt% and the theoretical reversible hydrogen gravimetric capacity is only 5.9 wt%. These numbers are smaller than those of Mg(AlH4)2 (Table 1.4 in Sect. 1.1). The MCAS synthesis of Ca(AlH4)2 has been investigated by Mamatha [117, 118] and very recently by Komiya et al. [126]. In this synthesis instead of MgCl2 as in the reactions of (3.23) a CaCl2 is used. The thermal dissociation of the mixture of Ca(AlH4)2 + 2NaCl obtained by the MCAS route and that of a single-phase Ca(AlH4)2 was observed upon heating from room temperature to 250°C at 4°C/min and yielded ~ 2.5 wt%H2 within ~ 80 min and ~ 5 wt%H2 within ~ 40 min. The dissociation took place in two steps (similar to the reactions of (3.20a) and (3.20b) where MgH2 is substituted by CaH2) with approximately equal amounts of hydrogen released in both steps. The expected thermolysis products, CaH2 and Al, apart from NaCl, could be identified by the XRD patterns. Addition of 2 mol%TiCl3 to the Ca(AlH4)2 + 2NaCl led to the accelerated desorption during ball milling resulting in about 2 wt%H2 desorbed within ~ 200 min of milling [117, 118]. The Ca(AlH4)2·THFx synthesized by a metathesis reaction of CaCl2 and NaAlH4, in a tetrahydrofuran (THF) solution, appeared to decompose in three steps at ~ 170, 241 and 277°C as observed by using

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a thermal desorption mass spectroscopy (TDS) but seemed to be more stable than Mg(AlH4)2 [126]. In summary, neither Mg(AlH4)2 nor Ca(AlH4)2 seem to have a potential for further development as solid state hydrogen storage materials for mobile applications. Their storage properties seem not to be better than those of LiAlH4 or NaAlH4. Nevertheless, they still deserve more research to arrive at the final verdict.

3.3

Amides

The extensive studies of hydrogen storage properties of lithium amide were activated by the pioneering efforts of Chen et al. [127]. They showed that Li3N compound can absorb and desorb hydrogen at a very reasonable pressure range according to the following reversible reaction path Li3 N + 2H 2 ↔ Li 2 NH + LiH + H 2 ↔ LiNH 2 + 2LiH

(3.25)

where Li2NH is “lithium imide” and LiNH2 is “lithium amide.” The theoretical amount of reversible hydrogen in this reaction is ~ 10.4 wt% (2H2/(Li3N + 2H2)) or 11.6 wt% if expressed per a molecule of Li3N. In fact, as far back in the past as 1910, Dafert and Miklauz reported erroneously that the reaction of (3.25) produced Li3NH4 which in reality was a mixture of LiNH2 + 2LiH [127]. Chen et al. reported that the overall heat of the reaction of (3.25) was −161 kJ/mol. They observed two plateaus on an absorption PCT curve (under 20 bar), the first one with a very low equilibrium pressure <0.07 bar which means that hydrogen absorbed in this region would not be easily desorbed unless under a high vacuum. It corresponded to a high-temperature portion of TGA curve which they also examined. The second absorption plateau was sloped and the overall equilibrium pressure was below 0.2 bar at 195°C, 0.5 bar at 230°C and 1.5 bar at 255°C. That simply means that relatively high temperatures are required for desorption at atmospheric pressure and in addition, at temperatures below 250°C a desorption of reasonable amount of H2 requires at least a primary vacuum. During absorption/desorption under 20 and 0.04 bar, respectively, they observed that Li2NH absorbed and desorbed reversibly ~ 6.5 wt%H2 according to the following reaction Li 2 NH + H 2 ↔ LiNH 2 + LiH

(3.26)

having the enthalpyΔH = −45 kJ/mol. That means that this amount of hydrogen can be reversibly stored in Li2NH. Using Van’t Hoff plots Chen et al. found that the standard enthalpy of absorption for Li2NH, ΔHab, is about −66.1 kJ/mol. Since after a few cycles absorption and desorption isotherms coincide, the standard enthalpy of desorption should be equal to ΔHab. Chen et al. observed a similar hydrogen reversibility for Ca2NH but in this case temperatures were much higher being in the range of 350–600°C and the amount of hydrogen stored reversibly was only ~ 1.9 wt%.

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3 Complex Hydrides

The above paper published in Nature by Chen et al. [127] triggered up an avalanche of research papers on various aspects of the hydrogen storage properties of Li–N–H system and its various modifications in the last 5 years. A number of these papers published up to 2006 have already been reviewed [2–5, 7]. Further research [128–140] firmly established that the general reaction of (3.25) occurs in a two-step path: Li3 N + H 2 ↔ Li 2 NH + LiH

(3.27a)

which is not really reversible requiring ~ 0.01 bar at 255°C with ΔH = −165 kJ/ molH2 [131] Δ H = −148 kJ/molH2 [134, 138], yielding hydrogen capacity ~ 5.5 wt% and Li 2 NH + H 2 ↔ LiNH 2 + LiH

(3.27b)

with Δ H = −44.5 kJ/molH2 [131]. Because of the smaller enthalpy change, the reaction of (3.27b) can more easily absorb/desorb ~ 6.5 wt%H2. Unfortunately, the utility of the reaction of (3.27b) is limited by two major drawbacks which will be briefly discussed. The first drawback is the evolution of ammonia (NH3) as a transient gas during the reaction of (3.27b) [128, 130, 131, 133, 135, 137, 138, 140, 141]. Hu and Ruckenstein [141] were the first to clearly point out that NH3 is formed through the decomposition of LiNH2 but is quickly captured by LiH. After careful analysis of all available experimental results Ichikawa et al. [130, 131, 138] proposed that the reaction of (3.27b) progresses with the involvement of the following two elementary sub-reactions which are essential for its completion 2LiNH 2 → Li 2 NH + NH 3

(3.27c)

which releases 37 wt%NH3, and NH 3 + LiH → LiNH 2 + H 2

(3.27d)

which releases 5.8 wt%H2. They calculated that the enthalpy change, ΔH, of the reaction of (3.27c) and (3.27d) is +84 kJ/molNH3 and −42 kJ/molH2, respectively. Obviously, the presence of NH3 in the hydrogen desorbed from LiNH2 is poisonous for the polyelectrolyte membrane of a conventional PEM fuel cell, even at trace levels, so at present even the smallest release of ammonia in the hydrogen gas cannot be tolerated in the system. In order to avoid the release of NH3, LiH must be present in the mixture and the capture of NH3 by LiH in the reaction of (3.27d) must be as fast as possible. The effects of various catalyzing agents in combination with high energy ball milling on the reaction of (3.27d) were studied. Ichikawa et al. [131, 138] used thermal decomposition mass spectroscopy for examining the basic properties of the 1:1 mixture of LiNH2 and LiH. They studied ball-milled mixture without any catalyst and the ball-milled mixture doped with 1 mol% of dry metals Ni, Fe, Co with particle size of several tens of nanometers and metal chlo-

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233

rides VCl3 and TiCl3. The undoped ball-milled mixture released hydrogen in the temperature range from 180 to 400°C while also emitting some amount of NH3, however, much smaller than the sample which was just mixed by hand using mortar and pestle. For nanoparticulate Ni, Co and Fe additives, a little emission of NH3 was observed above ~ 300°C. According to the interpretation provided by the authors the reaction rate is so low that a small amount of LiNH2 remained unreacted up to 300°C and started emitting NH3 at the temperature >300°C (the reaction of (3.27c)). In contrast, the sample doped with 1 mol%TiCl3/VCl3 easily desorbed ~ 5.5 wt%H2 in the 150–250°C range at the heating rate of 5°C/min. In addition, for the TiCl3 catalyzed sample no NH3 emission was detected up to 400–450°C. This results indicates that all LiNH2 was consumed as a result of complete reaction with LiH and transformed into Li2NH (the reaction of (3.27d)). The activation energy for hydrogen desorption calculated by the Kissinger method yielded 110 kJ/mol. Short cycling at 220°C desorption/180°C absorption showed that the cycle retention of TiCl3 bearing sample was good. In further studies from this group, Isobe et al. [137] reported that Tinano, TiCl3 and TiO2nano doped 1LiNH2:1LiH composites processed by ball milling revealed a superior catalytic effect on the TDS properties, while for the Timicro and TiO2micro additives catalytic effect was similar to the sample without any additives. Pinkerton [135] investigated the kinetic behavior of decomposition of ball-milled LiNH2 by ammonia release using thermogravimetric analysis (TGA). They noticed slow but significant decomposition by NH3 release in an H2 atmosphere (the reaction of (3.27d)) even under moderate pressure and temperature conditions (0.13 MPa, 175°C). The release of NH3 continued unabated even after a full conversion of the sample to LiH. The activation energy for the decomposition reaction was determined to be about 128 kJ/mole, virtually independent of the source and purity of LiNH2, its stoichiometry, ball milling time, and TGA sample size. Yao et al. [140] investigated the hydrogen storage behavior of pure LiNH2 and 1 LiNH2:1.2 LiH mixtures (overdose of LiH) with and without additives which were ball milled up to 4 h in a SPEX 8000 mill. They found that the decomposition mainly generated NH3. The only major endothermic peak at 380°C in the DSC curve corresponded to the melting of LiNH2. They noted that NH3 was released relatively slowly from the solid-state LiNH2, but rapidly from the liquid LiNH2, due to rapid diffusion of NH3 in the liquid at a high temperature. Milling only enhanced NH3 release in the solid-state LiNH2, but not in the liquid LiNH2 and NH3 started evolving from the milled LiNH2 around 50°C. The addition of Mn, MnO2, V or V2O5 to LiNH2 did not affect LiNH2 melting temperature, but enhanced NH3 release before the melting. With regard to the desorption from the 1 LiNH2:1.2 LiH mixture, 4 h of milling was needed to avoid escape of NH3. After 4 h milling the weight loss in TGA experiments was 5 mass% and the onset hydrogen desorption temperature was lowered from 160 to 120°C. No new phases were formed due to milling. They concluded that the effect of mechanical milling on the properties of (LiNH2 + LiH) should be attributed to improved contact between LiNH2 and LiH particles, as a result of homogeneous mixing and an increase in surface area of LiNH2 and LiH particles. Such an intimate contact ensures that the released NH3 is immediately captured by LiH particles and reacts according to the reaction of

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(3.27d), generating H2. Slightly higher content of LiH in the mixture (ratio 1:1.2) enhances this behavior. The addition to Mn, V and their oxides, had little influence on the production of H2 from the (LiNH2 + LiH) mixture. According to the authors it indicates that the rate-limiting step for dehydrogenation from (LiNH2 + LiH) is the reaction between LiH and NH3 and its kinetics could be enhanced by nanosized catalytic additives. Shaw et al. [142] investigated in detail the effect of ball milling in a modified Szegvari attritor on the hydrogen storage behavior of the mixture 1 LiNH2:1.1 LiH without any additives. Ball milling for 90 min reduced the crystallite size of LiNH2 and LiH to below 10 nm and slightly higher than 10 nm, respectively. Further increases in ball milling time beyond 90 min resulted in additional decreases in the crystallite size, but the reduction rate of the crystallite size was dramatically reduced. The equivalent particle diameters were reduced to 0.12, 0.11, and 0.09 μm for the powder mixture with the ball milling time of 90, 180, and 1,440 min, respectively. Apparently, high-energy ball milling reduced most of LiNH2 and LiH particles to submicrometer sizes and increased their specific surface area by at least one order of magnitude and created nanograins inside of particles. Dehydriding behavior was also greatly enhanced such that TGA curve for the powder milled for 180 min was shifted by ~ 70°C to the lower temperature range. DSC desorption peak of (1LiNH2:1.1LiH) (the reaction of (3.27b)) was reduced from 308 to 215°C by ball milling for 90 min. This peak was further reduced to 200°C if the ball milling time was increased to 180 min. Shaw et al. also observed a substantial effect of milling on the escape of NiH3. After ball milling for 180 min, the level of the NH3 intensity in the LiNH2 and LiH mixture dropped dramatically (by about 36 times). In fact, the NH3 intensity became lower than the detection limit, indicating that the intermediate reaction of (3.27d) between LiH and NH3 has been enhanced so much through high-energy ball milling that escaping of NH3 has been prevented. They have also found that escaping of NH3 is a kinetic issue. The escaping of NH3 depends not only on the powder characteristics and thus ball milling conditions, but also on the testing condition. For a given LiNH2 and LiH powder mixture, slow heating rates (e.g., no more than 5°C/min) can prevent escaping of NH3, whereas fast heating rates (e.g., 10°C/min) can promote escaping of NH3. The activation energy of desorption of (1LiNH2:1.1LiH) according to the reaction of (3.27b) gradually decreased with increasing milling time from ~ 164 kJ/mol for an “as-received” sample to ~ 63 kJ/mol after 1,440 min (24 h) of milling. Shaw et al. pointed out that the reduced particle size and increased surface area lead to an increased reaction rate, whereas a change in the activation energy is related to a change in the reaction mechanism or in the energy state of the reactants. Therefore, it was proposed that the reduction in the activation energy observed by the authors was due to an increase in the energy state of the LiNH2 and LiH mixture induced by high energy ball milling. The structural defects and the large area of grain boundaries within mechanically activated, nanostructured particles can contribute to this high-energy state. They also found the following trends: (1) the activation energy of decomposition of LiNH2 into Li2NH and NH3 according the reaction of (3.27c) exhibits the same trend as that for the reaction of (3.27b), i.e., decreases with increasing milling time, (2) the activation energy of the decomposition

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235

of LiNH2 into Li2NH and NH3 (the reaction of (3.27c) is always higher than that of decomposition of a mixture (LiNH2 + LiH) (the reaction of (3.27b)) and (3) the activation energy of the decomposition of the LiNH2 and LiH mixture with milling and then mixing is higher than that for the decomposition of LiNH2 as well as for the LiNH2 and LiH mixture with direct milling. The conclusion is that the role of LiH is not simply reacting with NH3 as defined by the reaction of (3.27d). In effect, LiH not only reacts with NH3, but also has catalytic effects on the reaction of (3.27c). Thus, the presence of LiH reduces the activation energy of the reaction of (3.27c) so that the “apparent” activation energy of the reaction of (3.27b) can be lower than the activation energy of the reaction of (3.27c) in the absence of LiH. This also explains why the LiNH2 and LiH mixture with direct milling has a lower activation energy than the mixture with milling and then mixing. Direct milling provides better mixing for the LiNH2 and LiH mixture than the mixture with milling and then mixing. Because of the better mixing of LiNH2 and LiH and the catalytic effect of LiH, the LiNH2 and LiH mixture with direct milling has a lower activation energy than the mixture with milling and then mixing. The second drawback of the reaction of (3.27b) is too high a temperature of reversible hydrogen absorption/desorption. For example, Kojima and Kawai [133] reported from PCT experiments that a single phase Li2NH could reversibly absorb about 6wt%H2 at 300°C under 9 MPa H2 pressure and desorb the same amount under vacuum (0.001 MPa) (PCT desorption plateau ~ 0.15 MPa). Their calculations using the Van’t Hoff equation gave ΔHab = −64.5 kJ/molH2 and ΔSab = −118 J/molH2K (absorption), and ΔHdes = −66.6 kJ/molH2 and ΔSdes = −120 J/molH2K (desorption). This is a larger enthalpy change than that reported in [131]. Unfortunately, even more thorough understanding of various mechanisms of the reaction of (3.27b) cannot change its unfavorable thermodynamics, specifically, because its enthalpy change, Δ H, seems to be closer to ~ 60 kJ/molH2 rather than the previously estimated ~ 45 kJ/molH2. A few approaches have been adopted to improve the storage conditions given by the reaction of (3.27b). Searching for the means to destabilize Li2NH (lithium imide) a number of authors studied a partial substitution of Li by metallic elements with larger electronegativity such as Mg [129, 132, 139]. Xiong et al. [129] observed that, for example, the Li–Mg–N–H system had a desorption peak of around ~ 166°C as compared to the peak desorption temperature of ~ 272°C for the undoped Li2NH. After hydrogenation the structure consisted of Mg(NH2)2 and LiH. This system had a plateau at around 10 bar for absorption and above that value for desorption at 180°C. Orimo et al. [132] and Nakamori et al. [139] prepared M(NH2)y where M = Li−x at.%Mg (x = 0–100 and y = 1–2). For LiNH2 with partial Mg substitution at x = 30 they observed, using thermal desorption spectroscopy, that the start of the thermal desorption was at around 100°C as compared to 280°C for LiNH2 without partial Mg substitution (a mixture of LiNH2 and LiH). It must however be pointed out that a full desorption was not completed even at ~ 300°C. These developments evolved into a search for novel hydride systems substantially differing from the original mixture (LiNH2 + LiH). Some sort of a breakthrough was achieved by Luo, and Luo and Rönnebro [143–145]. They replaced LiH with

236

3 Complex Hydrides

MgH2 in the original mixture LiNH2 + LiH and used ball milling for 2 h in a SPEX 8000 mill to prepare a mixture 2LiNH2:1.1MgH2 or 2LiNH2:1MgH2 (10% excess of MgH2 to prevent LiNH2 to decompose with the release of NH3 during cycling). The PCT experiments for absorption/desorption showed the plateau pressure at 220°C at ~ 50 and ~ 40 bar, respectively. That for desorption at 240°C was close to 70 bar. Isothermal desorption of a lightly cycled sample at 220°C in a Sieverts-type apparatus yielded ~ 4.5 wt%H2 at the pressure increasing up to 1.6 bar within about 1 h. The fresh sample (uncycled) required more than 15 h for desorption at 240°C. From the Van’t Hoff plot the decomposition enthalpy change was calculated as equal to ~ 34–39 kJ/molH2. XRD studies showed that the fresh sample contained mainly LiNH2 and MgH2 but the rehydrided sample contained LiH and Mg(NH2)2 phases. On that basis they proposed the following equation for the process 2LiNH 2 + MgH 2 → Li 2 Mg(NH)2 + 2H 2 ↔ Mg(NH 2 )2 + 2LiH

(3.28)

One should note that the reversible reaction in this equation takes place between (Mg(NH2)2 + 2LiH) and (Li2Mg(NH)2 + 2H2). Further to these original papers this new system has been more thoroughly investigated by a number of researchers [146–154]. It has been found [146] that a PCT curve for hydrogen desorption at 220°C from the mixture of Mg(NH2)2 and 2LiH exhibits two pressure-dependent segments at a given temperature and about 2/3 of the hydrogen is released at relatively higher pressures and forms a pressure plateau and the other 1/3 is desorbed at lower pressures (~ 5 wt%H2 total desorbed). The overall reaction heat measured in a differential scanning calorimeter is 44.1 kJ/ molH2, while the heat of desorption of H2 in the higher pressure plateau is about 38.9 kJ/molH2. These thermodynamic values are very favorable for PEM fuel cell application. An extension of the Van’t Hoff plot to 1 bar yields ~ 90°C but TPD measurements show that desorption starts at a temperature >100°C. The calculated activation energy of desorption Ea = 102 kJ/mol is high and sets a kinetic barrier. The authors suggested that such high activation energy may be related to the energy required for the reconstruction of reactants and dissociation of chemical bonds. Chen et al. [147] reported the results of cycling absorption at 200°C at 5 MPa hydrogen pressure and desorption in a primary vacuum (pressure<0.1 MPa) of ball-milled (2LiNH2:1MgH2) mixture. The mixture was able to desorb about 4 wt%H2 from the second to the fourth cycle within about 2 h. Desorption rate at 180 and 160°C under primary vacuum was much lower (e.g., 12 h to desorb ~ 3 wt%H2 and barely ~ 1.5 wt%H2 at 160°C). Luo and Sickafoose [148] investigated more carefully the progress of the desorption/absorption reaction of (3.28) and concluded that the reactions at the higher pressure plateau of a PCT curve observed by Xiong et al. [146] can be described by the following reaction (for wt%H2 > 1.5) Li 2 MgN 2 H 3.2 + 1.4H 2 ↔ Mg(NH 2 )2 + 2LiH

(3.29a)

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237

and for the lower pressure plateau and sloping region (for 0 < wt%H2<1.5) with Li 2 MgN 2 H 2 + 0.6H 2 ↔ Li 2 MgN 2 H 3.2

(3.29b)

Yang et al. [149] investigated two routes of activating the mixture of LiNH2 and MgH2. They noted that direct dehydriding of the mixture LiNH2 and MgH2 under vacuum at high temperature may result in the formation of NH3 from the decomposition of LiNH2 (3.27c) which is competitive to the first stage of the reaction of (3.28). They subjected the ball-milled sample (2LiNH2:1.1MgH2) to vacuum and heat (235–300°C for 1–5 h). Upon evacuating the milled sample (with nanograins 7–25 nm) for 5 h the XRD phase analysis showed the presence of LiNH2, Li2NH, MgH2, Li2Mg(NH)2 and LiH. After evacuation at 300°C for 5 h the majority phase was Li2NH with a small amount of LiH and LiMgN. This clearly shows that the decomposition of LiNH2 into Li2NH and NH3 (3.27c) proceeds slowly at 235°C but dominates at 300°C. PCT for the sample activated by this method had reduced capacity to 2.1 wt%H2 due to the partial decomposition Li2NH and irreversible release of NH3. The second method of activating was heating a sample under 110 bar pressure at 235°C for 1–5 h. The XRD phase analysis revealed the presence of Mg(NH2)2 and LiH confirming the reaction of (3.28). After dehydrogenation at 240°C to 1 bar the presence of Li2Mg(NH)2 was confirmed as required by the reaction of (3.28). They also calculated from the Van’t Hoff plot that the enthalpy change was 41.6 kJ/molH2. Absorption process under 110 bar was slow at 180°C (<2 wt%H2 in 8 h). Luo et al. [150] found that exposure of the ball-milled mixture (2LiNH2:1.1MgH2) to water-saturated air had a small positive effect on the rates and final wt%H2. Lohstroh and Fichtner [151] studied the sorption properties of ball-milled 2LiNH2–MgH2 mixtures. They found that during the first desorption cycle or intense ball milling Mg(NH2)2 was formed which then was a part of the reversible reaction. Their findings support an ammonia mediated reaction mechanism as was discussed earlier for the LiNH2–LiH system. Accordingly, the desorption involves various steps: most likely, initially Mg(NH2)2 emits NH3 which then reacts with the binary hydride LiH to form LiNH2. In a second step, MgNH and either LiNH2 or LiH react to form a mixed LiMg-imide phase. The addition of TiCl3 enhanced the milling intensity and hence lowered the desorption temperature but it did not have a catalytic effect on the reversible reaction. The escape of NH3 was also investigated by Luo et al. [152, 153]. They reported that the concentration of NH3 in desorbed H2 increased with the desorption temperature. For the (2LiNH2 + MgH2) system the NH3 concentration was found to be 180 ppm at 180°C and 720 ppm at 240°C. The capacity loss after 270 cycles at a temperature of 200°C was 25%, with 1/3 of the loss due to NH3-formation. They concluded that more research is needed to determine the cause for the remaining capacity loss. Most recently, Barison et al. [154] investigated the effect of intensive ball milling up to 48 h on the absorption/desorption properties of a mixture (2LiNH2:1.1MgH2). The milled samples were activated by means of heat treating

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3 Complex Hydrides

for 18–20 h at 220–230°C under 11–12 MPa of H2 as suggested in [149]. The XRD calculations of the ball-milled mixture yielded the grain size of less than 25 nm for MgH2 and less than 35 nm for LiNH2. The fastest absorption kinetics were observed for the mixture milled for 48 h. Under 9 MPa of H2 the mixture was able to absorb ~ 4.5 wt%H2 at 220°C. Desorption was less affected by milling and the differences were not great for the mixtures milled for 2, 12, 24 and 48 h. Roughly ~ 4.5 wt%H2 could be desorbed within 20–30 min at 220°C under 0.25 MPa of H2 pressure. PCT curves for desorption were investigated from which the enthalpy change for desorption was calculated from the Van’t Hoff plot as equal to 40.4 kJ/molH2, in excellent agreement with previous estimates. A parallel route taken by a number of researchers was based on the approach of replacing lithium amide (LiNH2) at the right hand side of the reaction of (3.27b) with an equivalent metal amide, M–N–H (M-metal). The selected amide has been magnesium amide (Mg(NH2)2). This approach although similar to the one described above (the reaction of (3.28)) differs in one important aspect, namely, that the selected amide, Mg(NH2)2, is pre-synthesized. In a sense this is a composite of two hydrides (two dissimilar hydrides which were processed together). Mg(NH2)2 can be synthesized relatively easily either by annealing a commercial MgH2 powder under gaseous NH3 in a reactor at the temperature range 330–380°C for 1 week, or by ball milling of MgH2 in a gaseous NH3 at room temperature [155, 156]. Magnesium amide is not produced commercially yet [157]. The reaction to form Mg(NH2)2 is as follows: MgH 2 + 2NH 3 → Mg(NH 2 )2 + 2H 2

(3.30)

A single-phase (unmixed) Mg(NH2)2 decomposes during continuous heating up to 500–600°C in a two-step process [155, 156, 158] Mg(NH 2 )2 → MgNH + NH3 (upto ~ 400° C)

(3.31a)

MgNH → 1/3Mg3 N 2 + 1/3NH 3 (above ~ 400° C)

(3.31b)

where MgNH is magnesium imide. An important question now is: “what is the ratio of hydrides in the mixture of Mg(NH2)2 and LiH?” The most popular choice [134, 156, 159–163] is the original ratio 3Mg(NH2)2 to 8LiH proposed by Leng et al. [156]. Thermal decomposition mass spectrum (TDMS) and TGA techniques show that H2 starts desorbing from this mixture processed by ball milling around 140°C and the desorption peaks around 190–200°C (heating rate 5°C/min) with the total desorbed H2 about 7 wt%. The ball-milled mixture dehydrogenated at 170°C in vacuum could be rehydrogenated at 200°C in 3 MPa of hydrogen. It seems that the 3Mg(NH2)2 + 8LiH system is reversible. The general reaction is given by 3Mg(NH 2 )2 + 8LiH ↔ Mg3 N 2 + 4Li 2 NH + 8H 2

(3.32)

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239

It has been suggested that the reaction of (3.32) is NH3 mediated according to the following partial reactions scheme [156] 3Mg(NH 2 )2 + (8 LiH ) → 3MgNH + 3NH 3

(3.33a)

3MgNH → Mg3 N 2 + NH 3

(3.33b)

4NH 3 + 4LiH → 4LiNH 2 + 4H 2

(3.33c)

4LiNH 2 → 2Li 2 NH + 2NH 3

(3.33d)

2NH 3 + 2LiH → 2LiNH 2 + 2H 2

(3.33e)

2LiNH 2 → Li 2 NH + NH 3

(3.33f)

2NH 3 + 2LiH → 2LiNH 2 + 2H 2

(3.33g)

2LiNH 2 → Li 2 NH + NH 3

(3.33h)

After summing up the phases on the right hand side of these partial reactions, one obtains on the right hand side of the overall reaction of (3.32) the following: Mg3N2, 4Li2NH, 8H2 and 8NH3 escaped. However, for the (3Mg(NH2)2 + 8LiH) mixture PCT absorption curves at 150 and 200°C after 1 h of waiting time are steep and they lack well-developed plateaus [160]. No desorption PCT curves are available. A small amount of NH3 was observed to be desorbed at the peak temperature of H2 desorption (~ 200°C) [161]. Another systems investigated in some detail are 3Mg(NH2)2 + 6LiH and 3Mg(NH2)2 + 12LiH [139, 164, 165]. Aoki et al. [164, 165] made a comparative study of 3Mg(NH2)2 + nLiH (n = 6, 8 and 12) using PCT curves and XRD. The reversible reactions for n = 6 and 12 are supposed to be expressed as follows: 3Mg(NH 2 )2 + 6LiH → 3Li 2 Mg(NH)2 + 6H 2

(3.34)

3Mg(NH 2 )2 + 12LiH → Mg3 N 2 + 4Li3 N +12H 2

(3.35)

They found that the amounts of the hydrogen desorbed from the mixtures with n = 6, 8, and 12 on a unit mass basis slightly decreased with increasing n (~ 5.4, 5.1 and 4.5 wt%, respectively). However, the molar ratios of the desorbed hydrogen to the mixtures were almost equal and the PCT isotherms were similar to each other. The plateau pressure for desorption of the (3Mg(NH2)2 + 12LiH) mixture was equal to 8–10 MPa, ~ 3.5 MPa and ~ 2 MPa at 250°C, 225 and 200°C, respectively. The desorption/absorption PCT curve at 250°C exhibited only very small hysteresis which means that the plateau pressures at this temperature are nearly identical. The Li2Mg(NH)2 and LiH phases were observed in XRD profiles of all the mixtures after PCT measurements. These results suggest that the dehydriding reaction of the

240

3 Complex Hydrides

mixtures of 3Mg(NH2)2 + nLiH (n = 6, 8, and 12) under hydrogen pressure is not dominantly affected by the value of n. Therefore, Aoki et al. [165] eventually proposed a general equation for the desorption reaction in the following form 3Mg(NH 2 )2 + nLiH → 3Li 2 Mg(NH)2 + ( n-6)LiH + 6H 2

(3.36)

They also concluded that the reactions given by (3.32) and (3.35) occur above 250°C. Interestingly, the signal intensity of ammonia desorbed from the mixtures detected by thermal desorption mass spectroscopy decreased with increasing n; this implies that the increase of n has an effect on decreasing the amount of ammonia desorbed from the mixture of 3Mg(NH2)2 + nLiH. The enthalpy and entropy of hydrogenation (Mg3N2 + 4Li3N mixture) calculated using Van’t Hoff plot obtained from the plateau pressures was equal to -46 kJ/molH2 and -104 J/molH2K, respectively. For the sake of clarity it must be mentioned that some more exotic amide combinations (composites) have also been studied: (Mg(NH2)2 + MgH2; desorbed ~ 2 wt%H2 and NH3) [158], (Mg(NH2)2 + 1 or 1.5 or 2NaH; reversible capacity ~ 2 wt%H2) [166], (LiNH2 + LiAlH4; possible capacity ~ 8 wt%H2?) [167], (2LiNH2 + LiBH4 + MgH2; max. ~ 8 wt%H2 desorbed at 260°C under 1 bar H2; reversible capacity ~ 2.8 wt%H2 at 180°C under 100 bar; NH3 suppressed) [168], (Ca(NH2)2 + NH; capacity ~ 1 wt%H2) [169], (Ca(NH2)2 + 2LiH and Ca(NH2)2 + 2LiNH2; desorbed ~ 4.5 wt%H2 between 100 and 250°C; NH3 suppressed) [170], (Mg(NH2)2 + CaH2; desorbed ~ 3 wt%H2 under 80 bar H2 pressure (!) but rehydriding unsuccessful) [171]. It must be concluded that these more exotic combinations of amides and other hydrides are not more promising than the original (LiNH2 + LiH) mixture. In summary, it must be concluded that amides as a category do not seem to have a substantial advantage over LiAlH4 system doped with catalytic precursors. There is no evidence that the amides could reversibly desorb/absorb at temperatures lower than 200°C. At the temperatures higher than 200°C amides and related materials barely desorb ~ 5 wt%H2. Their kinetics of isothermal desorption at the atmospheric pressure is unknown. Still relatively high pressures are needed for rehydrogenation. The unsolved problem yet is the escape of NH3 whose minute quantities also will destroy the membrane of presently used PEM fuel cell. After such a substantial amount of research work on amides the future looks rather bleak and it seems doubtful that any further quantum leap improvements could be achieved. Nevertheless, this is one of a very limited group of hydrides which are able to desorb a decent amount of hydrogen close to 200°C. Nanostructuring through ball milling and catalyst precursors such as metal chlorides provide some limited improvement of the storage properties.

3.4

Metal Borohydrides

Metal borohydrides are very well known chemicals and some of them have been used commercially in chemical industry for a number of years. The first metal borohydride, LiBH4, was synthesized almost 70 years ago by the reaction of ethyl

3.4 Metal Borohydrides

241

lithium with diborane (B2H6). More information on the synthesis by wet chemical methods of various borohydrides the reader can find in [7, 157 – paper by Soloveichik]. It is also to be pointed out that there is some experimental evidence of the direct synthesis from the elements under high pressure of hydrogen of such borohydrides as LiBH4, NaBH4, Mg(BH4)2 and Ba(BH4)2, reported in a 50-year-old German patent by Goerrig [172]. The synthesis temperatures were around 620– 650°C and pressures around 150 atm. The yield of magnesium borohydride was not too large but the mere fact that the hydride could be synthesized from the elements is very encouraging. An alternative synthesis to the exchange (metathesis) reaction in a wet chemistry is the same reaction in solid state using ball milling (mechanochemical activation synthesis-MCAS) as was first reported by Dymova et al. [115, 116] for the synthesis of magnesium alanate (Sect. 3.2.3). This method of solid state synthesis by ball milling was also applied to the synthesis of various borohydrides in [173–175]. In this synthesis the mixture of MgCln and either LiBH4 [173, 174] or NaBH4 [173– 175] is mechanically milled under argon and the synthesis reaction occurs according to the following: MCln + nLiBH4 → M(BH4)n + nLiCl

(3.37a)

MCln + nNaBH4 → M(BH4)n + nNaCl

(3.37b)

As mentioned in Sect. 3.2.3 a drawback of this method is the formation of a byproduct, which is salt, which reduces the overall hydrogen content of the synthesized mixture. Another peculiarity of this synthesis method is a complete lack of XRD peaks corresponding to a crystalline synthesized borohydride in the mixture in the right-hand side of the reaction of (3.37). Only the peaks of either LiCl or NaCl are present. On the other hand, the Raman spectra seem to indicate the presence of borohydride. Nakamori et al. [173, 174] interpreted this peculiar behavior as arising most likely due to “disordering” of the crystal structure of a synthesized borohydride, whatever it means for a solid state hydride. Application of MCAS to the synthesis of Mg(BH4)2 will be discussed later in the text. Very recently, ball milling was also used to synthesize sodium borohydride, NaBH4, by mechano-chemical reaction of dehydrated borax (Na2B4O7), MgH2 and Na2CO3 (at room temperature) according to the following reaction: [176] 8MgH2 + Na2B4O7 + Na2CO3 = 4NaBH4 + 8MgO + CO2

(3.38)

The attractiveness of metal borohydrides for solid state hydrogen storage stems from their very high hydrogen capacities as can be seen in Table 1.4 in Sect. 1.1. As a category they constitute the highest hydrogen capacity hydrides. Unfortunately, as will be shown further, they are, most likely, the most difficult of all other hydrides for implementing modifications which would bring their hydrogen storage properties closer to the requirements of D.O.E. The Russian scientists carried out a detailed study of the thermal stability of sodium borohydride (NaBH4), lithium borohydride (LiBH4), potassium borohydride (KBH4) and magnesium borohydride

242

3 Complex Hydrides

(Mg(BH4)2) at the temperatures up to ~ 700°C and hydrogen pressures up 10 atm. [177, 178]. The thermal decomposition of NaBH4 (and KBH4 as well) is very simple. The microstructure of “as-received” NaBH4 and its DSC trace obtained in our laboratory are shown in Fig. 3.17. The first endothermic peak centered at 498°C is due to melting. Its peak temperature is in an excellent agreement with ~ 505°C reported in [177, 178]. The second endothermic peak is due to the decomposition in a molten state of NaBH4 according to the reaction put forward by Stasinevich and Egorenko [177, 178] NaBH4 → Na + B + 2H2

(3.39)

The same decomposition path was observed for KBH4 by Stasinevich and Egorenko [178] except that the melting and decomposition temperatures were much higher (by about 100°C). We performed ball milling for 5 h on the “asreceived” NaBH4 but without any effect on its decomposition behavior. This is most likely due to the fact that NaBH4 exhibits a high structural stability during

a Decomposition

Melting

b

Fig. 3.17 (a) SEM picture of as-received NaBH4 powder of 98% purity and (b) its DSC trace up to 600°C (heating rate 10°C/min, argon flow ~ 50 ml/min)

3.4 Metal Borohydrides

243

ball milling, with a high resistance to form nanostructure (i.e., nanograins and fine particulate) [179]. LiBH4 undergoes first a polymorphic transformation around 110°C, then melts with an endothermic peak centered around 280°C, and eventually at the atmospheric pressure it decomposes with an endothermic peak centered at ~ 470°C through the following reaction path LiBH4 → LiH + B + 3/2H2

(3.40)

The reactions of (3.39) and (3.40) are the two types of reactions by means of which more stable metal borohydrides can decompose, i.e., either to the elements or to a metal hydride and boron. Unstable metal borohydrides decompose through the formation of boranes which will be discussed later. More recently, Züttel et al. [180] investigated the thermal decomposition of undoped and SiO2 doped LiBH4. They quoted the value of enthalpy and entropy change for the reaction of (3.40) as equal to −177 kJ/ mol and 238 J/molK. Apparently this is a very high enthalpy change which determines high temperature for desorption. The undoped hydride showed the first thermal peak at ~ 100°C which corresponded to the structural transition from the orthorhombic phase at low temperature to a tetragonal phase which was also accompanied by a small amount of desorption of ~ 0.3 wt%H2. The melting peak was observed around 270°C without any desorption and the first small desorption peak appeared at ~ 320°C. The second dominant hydrogen desorption started at ~ 400°C and reached the maximum ~ 500°C. The SiO2 additive resulted in the first peak to become the dominant hydrogen desorption peak which started at ~ 200°C and reached the maximum at ~ 350°C. The second peak became very sharp and slightly shifted to a lower temperature. Apparently, the SiO2 additive acted in a sort of catalytic way but still the improvement was insufficient. They also mentioned that any attempt to synthesize LiBH4 from the elements at temperatures up to 650°C and hydrogen pressure of 150 bar as claimed in [172] failed. However, Orimo et al. [181] managed to successfully rehydrogenate the LiH and B mixture of powders to LiBH4 under 35 MPa of hydrogen at 600°C. Very recently Mosegaard et al. [182] used synchrotron radiation to investigate the structural transitions during decomposition of LiBH4. They found three new phases observed during dehydrogenation of LiBH4. Phase I is hexagonal, and is observed in the temperature range ~ 200–300°C, and phase II is orthorhombic, and is observed in the temperature range ~ 300–400°C applying a constant heating rate of 5°C/min. Phase I transforms into II, e.g., at a constant temperature of 265°C after 5 h. Furthermore, a third phase, III, is observed in the temperature range RT to 70°C, and is caused by a reaction between LiBH4 and water vapor from the atmosphere. Hydrogen release is associated with the decomposition of III at ~ 65°C. However, as can be seen from the above discussion, a pure LiBH4 is a very stable hydride and has no potential for commercial application. Some efforts to destabilize LiBH4 thermodynamically have been made and will be discussed in Sect. 3.5. LiBH4 is also very volatile when in contact with water/moisture. It must also be mentioned that LiBH4 is a very expensive hydride. Taking into account all these factors its application as a potential hydrogen storage material is more than problematic.

244

3 Complex Hydrides

Thermal decomposition of Mg(BH4)2 was first studied thoroughly by Stasinevich and Egorenko [178] and a few years later by Konoplev and Bakulina [183]. The hydride was prepared by the wet exchange reaction (metathesis) of sodium borohydride (NaBH4) with anhydrous magnesium chloride (MgCl2) in diethyl ether. The density of the compound was determined as 0.989 g/cm3. Both groups found four main endothermic events as a function of temperature. The first one around 180°C corresponded to the structural transition of low-temperature tetragonal polymorph α-Mg(BH4)2 into high-temperature face centered cubic (FCC) polymorph β-Mg(BH4)2. The second endo effect around 300–320°C corresponded to the simultaneous melting and decomposition of β-Mg(BH4)2 according to the following reaction β – Mg(BH4)2 → MgH2 + 2B + 3H2 (300 – 320°C)

(3.41a)

The third endothermic peak occurred around 400–420°C and could be assigned to the decomposition of MgH2 according to the well known reaction MgH2 → Mg + H2 (400 – 420°C)

(3.41b)

The fourth peak at 660°C was due to the melting of Mg. Apparently, Mg(BH4)2 decomposes by the second type of decomposition reaction of borohydrides accompanied by the formation of a simple metal hydride (similarly to LiBH4 in the reaction of (3.40)). According to Table 1.4 in Sect. 1.1 the theoretical amount of hydrogen (theoretical reversible gravimetric capacity) which should hypothetically be desorbed from the reaction of (3.41a) is equal to ~ 11 wt%. Since this first decomposition reaction occurs around 300°C which is similar to that of the ballmilled MgH2 (Chap. 2) and the hypothetical amount of desorbed hydrogen is much larger than the one desorbed from MgH2, the Mg(BH4)2 hydride is a good candidate for further studies with the aim of substantially reducing its first decomposition temperature to much below 300°C. Vajeeston et al. [184] predicted on the basis of density-functional studies that synthesizing Mg(BH4)2 from the pure elements is energetically more favorable than synthesis via binary intermediates. In our laboratory, we have made an attempt to synthesize Mg(BH4)2 from the elements [101, 185]. A stoichiometric Mg–2B powder mixture of elemental Mg (−325 mesh, 99.8% purity) and amorphous B, both oxidized (containing B2O3) and oxide-free (annealed at 370°C for 20 h to remove B2O3), were subjected to controlled reactive mechanical alloying under 880 kPa pressure of hydrogen in the magneto-mill Uni-Ball-Mill 5. Figure 3.18 shows the morphologies of Mg–2B powder mixtures after reactive mechanical alloying for 50 and 200 h. No substantial difference is observed regardless of whether oxidized or oxidefree boron was used. Large quantities of powder particles are refined to 100–200 nm size (nearly nanometric) although some larger aggregates in the 1–2 μm range can also be seen after 200 h of milling which shows some tendency for clustering after long-time milling. XRD analysis showed that the as-milled Mg–2B mixture made with oxidized amorphous boron after 200 h consisted of β-MgH2 (calculated nanograin size ~ 6 nm and no lattice strains) and γ-MgH2 hydrides. No ternary Mg(BH4)2 hydride or MgB2 was formed. DSC test showed only endothermic peaks

3.4 Metal Borohydrides

245

a

b

c

Fig. 3.18 Secondary electron micrographs showing the morphology of Mg–2B powders reactively mechanically alloyed. (a) Made with oxidized amorphous boron after 200 h, (b) made with oxide-free boron after 50 h and (c) made with oxide-free boron after 200 h of milling

due to the decomposition of MgH2 and corresponding XRD pattern showed the presence of free Mg (nanograin size ~ 11 nm) from the decomposition of MgH2 hydride, MgB2 (calculated nanograin size ~ 13 nm and no lattice strains) which was formed due to the reaction of free Mg with the retained amorphous boron and MgO formed due to the exposure of free Mg to air during handling for XRD tests. However, a striking difference is observed when the oxide-free boron is used. Figure 3.19a shows that the Mg–2B powder made with the oxide-free boron after 50 h of reactive synthesis still consists of β- and γ-MgH2 hydrides, retained Mg and retained amorphous boron. MgO in the XRD pattern is most probably from the oxidation of retained Mg during the XRD test. After 50 h of milling the estimated nanograin size of β-MgH2 is equal to about 6 nm (no lattice strains) which is the same as that of the powder made with oxidized boron milled for 200 h, indicating a saturation of grain size at longer milling times which is a common observation for milled hydrides. Quite a remarkable transformation is observed as the milling time reaches 200 h. The XRD pattern (Fig. 3.19b) demonstrates a complete disappearance of crystalline MgH2 phase and the formation of MgB2. The prominent feature is the broad background, which rises between 25 and 45° and is roughly centered at 2Q ~ 35°, and rises again at low angles. This is a strong evidence of the formation of an amorphous phase coexisting with MgB2. The DSC curves of powders after

246

3 Complex Hydrides 1600

b-MgH2 Mg MgO g -MgH2 amo-Boron ? substrate

Milled 50h ?

Counts

1200

800

? ? ?

400

?

0 20

30

40

a

50 60 Degrees 2-Theta

70

80

90

1600

Milled 200h MgB2

Counts

1200

800

400

0

b

20

30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 3.19 XRD patterns of Mg–2B powder mixtures made with non-oxidized boron after milling for (a) 50 h and (b) 200 h

milling for 50 and 200 h are shown in Fig. 3.20. A single endothermic peak at 383.3°C characterizes hydrogen desorption from the γ- and β-MgH2 phases in the 50 h milled powder (Fig. 3.20a). A striking result is that the DSC curve of the powder milled for 200 h shows an endothermic peak at 357.7°C (Fig. 3.20b) although there are no crystalline hydrides in the powder (Fig. 3.19b). Since it is unlikely that hydrogen is released from the crystalline MgB2, the only explanation is hydrogen desorption from an amorphous hydride phase, which gives broad backgrounds in the XRD pattern after a 200 h milling (Fig. 3.19b). XRD pattern of the 200 h milled powder

3.4 Metal Borohydrides

247

DSC / mW/mg 2.5 ↓ exo

[1] 383.3⬚C

2.0 [2] 357.7⬚C

1.5

b

a. Milled 50h b. Milled 200h

1.0

a

0.5 0 [1] 329.5⬚C −0.5 100

150

200

250

300 350 Temperature / ⬚C

400

450

500

Fig. 3.20 DSC curves of the Mg–2B powder mixture made with oxide-free amorphous boron after milling for (a) 50 h and (b) 200 h

MgB2 MgO Mg

,

4000

Counts

3000

2000

1000 After DSC Milled 200h

0 20

30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 3.21 XRD pattern of the 200 h milled Mg–2B powder mixture made with oxide-free amorphous B after DSC test. For comparison the XRD pattern after 200 h milling is also shown

(made with oxide-free boron) after DSC in Fig. 3.21 shows the presence of nanometric free Mg (calculated nanograin size ~ 11 nm and no lattice strains), nanometric MgB2 (calculated nanograin size ~ 13 nm and no lattice strains) and MgO. The latter is most likely formed due to the fast oxidation of free-active Mg after exposure

248

3 Complex Hydrides

to air during XRD test because it is absent after milling for 200 h (Fig. 3.19b). The XRD pattern of the 200 h milled powder after DSC test in Fig. 3.21 indicates that the decomposition of an amorphous hydride phase occurred during a DSC test as also observed by the appearance of an endothermic peak in Fig. 3.20b. The desorption test of the 200 h milled powder in a Sieverts-type apparatus at 350°C yielded ~ 1.4 wt% of desorbed hydrogen which correlates well with the endothermic peak of hydrogen desorption in DSC test in Fig. 3.20b. Since enthalpy of formation of MgB2 is −92 ± 8.4 kJ/mol [186] vis-à-vis −74 kJ/mol for the MgH2 hydride, the former will have a tendency to form first in the Mg–B–H mixture during mechanical milling. However, if one compares the synthesized phases in the 200 h milled powders made with the oxidized and oxide-free boron it can be seen that the boron oxide prevents reaction of boron with Mg to form MgB2 in the 200 h milled powder made with oxidized boron. It is most likely that amorphous boron particles are covered by oxide layer and even with a prolonged milling time (200 h) the layer is not cracked to induce the reaction. From results of XRD in Fig. 3.19 and DSC analysis in Fig. 3.20 for the Mg–2B mixture made with the oxide-free boron, it can be proposed that as milling time passes 50 h and reaches 200 h, the γ- and β-MgH2 hydrides transform into an amorphous hydride phase and free-active Mg. Free-active Mg reacts with the oxide-free a-B to form MgB2. The phase transformations during milling in the Mg–2B mixture made with the oxide-free boron can be postulated by the following reactions: γ- and β-MgH2 → amorphous hydride + Mg (free-active) + H2 and subsequently Mg (free-active) + 2 amorphous-B (oxide-free) → MgB2. Unfortunately, in this study we were unable to reveal the nature of the amorphous hydride phase although as will be shown further, it is most likely a disordered Mg(BH4)2. The second attempt to synthesize Mg(BH4)2 has been made in our laboratory by the mechano-chemical activation synthesis (MCAS) (or solid state metathesis reaction) [175] according to the reaction of (3.37b). The starting powder mixture of “as-received” NaBH4 (98% purity) and anhydrous MgCl2 (99% purity) were mixed together in the 2:1 stoichiometric ratio with the objective of initiating the following metathesis reaction 2NaBH4 + MgCl2 → c-Mg(BH4)2 + 2NaCl

(3.42)

where c refers to “crystalline.” The mixture was subsequently, subjected to controlled mechanical milling under 600 kPa pressure of high purity argon (99.999% purity) in the magneto-mill Uni-Ball-Mill 5 with two Nd–Fe–B super-strong magnets at 6 and 8 o’clock positions (IMP68). The samples were milled continuously for 1, 3, 10, 20, 50 and 100 h at ~ 225 rpm. Figure 3.22 shows the phase evolution as observed by XRD vs. milling time. The presence of NaCl indicates that the reaction of (3.42) occurred during milling whereas the presence of retained MgCl2, MgCl2– 4H2O (most likely formed due to the exposure of retained MgCl2 to moist air during XRD tests), and retained NaBH4-type phase, show that the reaction of (3.42) has not been completed regardless of the milling duration even as long as 100 h. It is clear that the peaks belonging to the retained NaBH4-type phase in the milled powder mixtures are clearly shifted to the higher 2Q angles with respect to the peak positions of the reference “as-received” NaBH4. The shifting may indicate a decrease of

3.4 Metal Borohydrides

249 NaBH4 (Na,Mg)BH4 NaCl MgCl2 MgCl2-4H2O NaBH4 (Reference)

8000

1h

Counts

3h 10h

20h

4000

50h

100h

0 20

30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 3.22 XRD patterns of powder mixtures of 2NaBH4 and MgCL2 after milling for 1, 3, 10, 20, 50 and 100 h under IMP68 mode

the lattice parameter and shrinkage of the unit cell volume of NaBH4. The lattice parameters and unit cell volume were calculated and indeed, as shown in Fig. 3.23, there is a clear trend of unit cell volume decreasing with increasing milling time. We excluded any structural transition of NaBH4 [175] and concluded that the observed lattice shrinkage is owing to the dissolution of Mg which has a much smaller atomic radius of 0.1604 nm than that of Na (0.1858 nm) [187] in the original NaBH4 lattice. In other words, as milling time increases, more and more Mg from the decomposing MgCl2 diffuses into NaBH4 lattice, substitutes Na in the lattice forming (Na,Mg)BH4 solid solution and causes the shrinkage in the original unit cell volume of NaBH4. Based on DSC and TGA experiments in combination with XRD examination and comparison with thermal events observed by other authors [173, 174, 178, 183] who also investigated the thermal behavior of Mg(BH4)2 we proposed the following scheme of phase transformations taking place during milling and subsequent thermal experiments [175]: (1) Milling with MCAS • Partial reaction of 2NaBH4 + MgCl2 → d-Mg(BH4)2 + 2NaCl (where “d-” is “disordered”); the presence of d-Mg(BH4)2 we invoke just on the basis of the results presented by Nakamori et al. [173, 174] based on their Raman spectroscopy measurements (we are unable to confirm this)

250

3 Complex Hydrides

Lattice parameter (Na,Mg)BH4 (nm)

0.618

Unmilled

0.616 0.614 0.612 0.61 0.608 0.606 0.604 0.602

0

20

40

60

80

100

120

Milling time (h)

Fig. 3.23 Lattice parameter of (Na,Mg)BH4 solid solution as a function of milling time

• Formation of a solid solution (Na,Mg)BH4 from the unreacted portion of NaBH4 by Mg substitution into the lattice of NaBH4 where the solutionizing effect increases with increasing milling time (2) T~100–240°C; weight loss ~ 0.4–1.3 wt% • Decomposition of magnesium chloride hydrate (MgCl2–xH2O) (observed in TGA) (3) T~275–335°C; weight loss ~ 1.0–1.5 wt% • Melting and a simultaneous full or partial decomposition of d-Mg(BH4)2 into β-MgH2, B and H2 which gives an endothermic DSC peak with the maximum at ~ 298°C (4) T~335–420°C; weight loss ~ 1.0–1.8 wt% • A small DSC heat flow peak at ~ 350–360°C in Fig. 5a–d in Ref. [175] remains unexplained • Decomposition of β-MgH2, with the peak maximum around 410–420°C • Possible further decomposition of the remnants of d-Mg(BH4)2 in the above range of temperatures Regardless of the crystallographic nature of the synthesized d-Mg(BH4)2 hydride phase the milled powders desorb hydrogen within the 275–420°C temperature range and the total amount desorbed is the largest for the powder milled for 10 h (~ 3.2 wt%) and gradually decreases with increasing milling time reaching ~ 2.1 wt% for the powder milled for 100 h. This is still smaller than the theoretical capacity ~ 4.7 wt%H2 of the initial (2NaBH4 + MgCl2) mixture. It is also highly likely that the “amorphous” hydride phase synthesized from the elements by ball milling under hydrogen

3.4 Metal Borohydrides

251

[101, 185] as described earlier is in reality a disordered Mg(BH4)2 phase. We have also hypothesized that the presence of (Na,Mg)BH4 solid solution in the MCAS process inhibits the formation of crystalline Mg(BH4)2. Once a solid solution is formed, the amount of Mg is insufficient to form a large amount of perfectly crystalline Mg(BH4)2 hydride and more disordered Mg(BH4)2 (d-Mg(BH4)2) hydride forms. Despite a general agreement of the thermal behavior of disordered Mg(BH4)2 with the early works on the synthesis and thermal decomposition of crystalline β-Mg(BH4)2, a few outstanding issues remain puzzling. Why did the “wet” exchange reaction between NaBH4 and MgCl2 in diethyl ether as carried out by the Russian researchers over 30 years ago [178, 183] result in a perfectly crystalline Mg(BH4)2 whereas MCAS gives a disordered Mg(BH4)2? Is it possible that a disordered Mg(BH4)2 forms only if ball milling is applied to its synthesis? Interestingly enough, very recently Matsunaga et al. [188] synthesized a disordered Mg(BH4)2 by the reaction of LiBH4 with MgCl2 by first pressing a 2:1 mole ratio mixture into a pellet and subsequently annealing the pellet at temperatures greater than 250°C under 10 MPa of hydrogen pressure (in equilibrium with LiCl). This somehow suggests that ball milling is not a requisite for the synthesis of disordered Mg(BH4)2. Nevertheless, because a pressure was applied to the powder to make it into a pellet it cannot be ruled out of hand that some sort of external deformation still promotes the formation of disordered Mg(BH4)2. In addition, they established that the enthalpy change of the first decomposition reaction of Mg(BH4)2 into MgH2, B and H2 (the reaction of (3.41a) where β-Mg(BH4)2 is now d-Mg(BH4)2) is equal to 39.3 kJ/molH2. However, only the second step reaction (3.41b) is reversible at the condition up to 350°C at 10 MPa of hydrogen. At this point it should also be mentioned that recently Jeon and Cho [189] apparently synthesized a crystalline Zn(BH4)2 by using MCAS of a mixture of ZnCl2 and NaBH4 as opposed to Nakamori et al. [173, 174] who obtained a disordered Zn(BH4)2 as a result of MCAS. A few papers on the thermodynamical stability of metal borohydrides have recently been published by Nakamori et al. [173, 174, 176, 190, 191]. Using the first principles calculations they have found that there is a strong linear correlation between the heats of formation of metal borohydrides and the Pauling electronegativities of cations, cP, as shown in Fig. 3.24a. There is also a correlation between the desorption temperatures of metal borohydrides and the Pauling electronegativities of cations as shown in Fig. 3.24b. However, it must be pointed out that the desorption temperatures in Fig. 3.24b were found experimentally for the disordered metal borohydrides synthesized by MCAS rather than for the crystalline metal borohydrides. As can be seen, desorption temperature decreases with increasing Pauling electronegativity which could be very good news for storage applications. Unfortunately, the bad news is that the metal borohydrides with large Pauling electronegativities, greater than about 1.3, desorb with the release of toxic borane BnHm (e.g., B2H6) gas that can ruin the membrane of PEM fuel cell, despite their desorption temperature ranges looking very attractive (approximately 80°C) [189, 192, 193]. It is reported that the addition of nano Ni catalysts can lower the melting and thermal decomposition temperatures (at least 20–40°C) of Zn(BH4)2 as evidenced from the calorimetric analysis. At these low temperatures, the nanocatalyzed Zn(BH4)2 exhibits reduction

252

3 Complex Hydrides

Fig. 3.24 (a) Relation between the heat of formation, D H, of metal borohydride in the unit of kJ/ molBH4 and the Pauling electronegativity of the cation, cP [190]. (b) The desorption temperature, Td, as a function of the Pauling electronegativity of the cation, cP. Inset shows the correlation between Td and estimatedD Hdes for the desorption reaction [173]

in the amount of borane gases released by a factor of 20 as compared to the undoped sample. Unfortunately, the release of boranes cannot be suppressed completely and even minute quantities of this gas could easily destroy the PEM FC membrane. In summary, on the one hand, metal borohydrides with low Pauling electronegativities (<1.3) exhibit high decomposition temperatures. Unfortunately, at the moment there is no viable approach on the horizon by means of which high decomposition temperatures of metal borohydrides could be reduced to a reasonable level, say, 200°C. On the other hand, metal borohydrides with the Pauling electronegativities

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

253

greater than 1.3 have quite low decomposition temperatures, essentially compatible with the D.O.E. targets, but they desorb some quantities of toxic borane gases. It remains to be seen whether or not the release of boranes can be suppressed completely; e.g., by appropriate catalysts.

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

The concept of compositing various hydrides in order to destabilize one through the other in order to reduce the desorption temperature of one of the hydrides has been floating around for some time. Ball milling was used by Zaluska et al. [194] to produce a nanostructured mixture of 65 wt%MgH2 + 35 wt%Mg2NiH4 which showed a lowered desorption temperature around 240–280°C. Johnson et al. [195] produced a mixture MgH2 + 0.1 mol.%LiBH4 by annealing it in a quartz tube for 12 h but no desorption at the temperature lower than 300°C was reported. Vajo et al. in a series of papers [196–198] introduced the idea of creating a sort of a composite in which the principal thermally stable hydride is mixed, usually by ball milling, with the second constituent – either hydride or an intermetallic compound. This second component reacts during dehydrogenation of the principal hydride and forms another intermetallic compound, lowering the enthalpy change. This concept is illustrated schematically in Fig. 3.25a. In this case, in isolation, the pure hydride AH2 undergoes dehydrogenation to form A + H2 with a relatively high enthalpy. Consequently, the equilibrium hydrogen pressure at room temperature will be low. Alternatively, the temperature required for an equilibrium pressure of 1 bar, will be high. However, if the chemical environment of AH2 is altered by adding a second component, B, that form an intermetallic or other compound with A, then dehydrogenation can proceed to ABx + H2. This reaction occurs with a reduced enthalpy and, therefore, an increased equilibrium hydrogen pressure. Thus, AH2 is effectively thermodynamically destabilized, even though the bonding of AH2 is not altered. If the additive, B, does not form a hydride, destabilization occurs with a gravimetric penalty that depends on the relative atomic weights of A and B and the stoichiometry of the ABx compound. Using an additive that is itself a hydride can minimize this penalty. An example is the destabilization of LiBH4 with MgH2 according to this concept which is shown in Fig. 3.25b. Normally, the mixed ionic/ covalent bonding in LiBH4 is rather strong which results in high thermodynamic stability of the hydride. As reported by Vajo et al. equilibrium calculations based on known enthalpies, entropies and heat capacities predict Δ H = 67 kJ/molH2 and T = 410°C at 1 bar hydrogen pressure (atmospheric) for dehydrogenation to LiH + B + 3/2H2 (the reaction of (3.40)). However, when MgH2 is present as a destabilizing agent, dehydrogenation can proceed according to: 2LiBH4 + MgH2 « 2LiH + MgB2 + 4H2

(3.43)

254

3 Complex Hydrides

Dehydrogenated state DH large, T(1 bar) high

Stabilized (alloy) state ABX + H2 DH smaller, T(1 bar) lower

ENTHALPY (kJ)

ENTHALPY

A + H2

200

2LiH + 2B + MgH2 + 3H2

183

2LiH + MgB2 + 4H2

67 kJ/H2 T = 410⬚C 46 kJ/H2 T = 170⬚C 2LiBH4 + Mg + H2

76

76 kJ/H2 T = 285⬚C

AH2 + xB

Hydrogenated state 0

a

2LiBH4 + MgH2

b

Fig. 3.25 (a) Generalized enthalpy diagram illustrating destabilization through alloy formation upon dehydrogenation. Including the alloying additive, B, reduces the enthalpy for dehydrogenation through the formation of ABx and effectively destabilizes the hydride AH2. (b) Enthalpy diagram for the destabilization of LiBH4 by MgH2. Addition of MgH2 reduces the enthalpy for dehydrogenation of LiBH4 through the formation of MgB2. Dehydrogenation of MgH2 without LiBH4 decomposition is shown as a possible intermediate step [198]

accompanied by the formation of MgB2 boride compound which effectively destabilizes LiBH4 because now this reaction has a reduced enthalpy change Δ H = 46 kJ/molH2 and a calculated T = 170°C at 1 bar of hydrogen pressure. Vajo et al. developed several systems based on LiBH4 destabilized by compositing with MgH2, MgF2, MgS and MgSe. In each case the destabilizing intermetallic compound formed is MgB2 [196–198]. They demonstrated the reversibility of the LiBH4 + MgH2 system up to 10 wt%H2 absorbed/desorbed but only at very high temperatures ~ 350–400°C. Equilibrium PCT pressures varied from 4.5 bar at 315°C to 19 bar at 450°C. The kinetics were so slow that up to 100 h were necessary to attain equilibrium. Extrapolation of the equilibrium pressures at 315– 400°C to lower temperatures gave a pressure of 1 bar at 225°C. This value was significantly higher than 170°C calculated for this system. Even with metal chloride catalytic precursors and additional nanostructuring (nanoporous carbon support with a 13 nm pore size) the kinetics were too slow to obtain a meaningful amount of desorbed H2 below 200°C. Similar problems with the system based on Ca(AlH4)2 and destabilized by various additives has recently been reported by Hanada et al. [199]. Very recently [200, 201] we have formulated a hypothesis that by compositing hydride constituents having high and low temperatures of desorption, the desorption temperature of the constituent having high desorption temperature can be linearly reduced with an increasing volume fraction of the low-desorption temperature constituent according to the composite Rule-of-Mixtures (ROM). According to this hypothesis, the compositing of high decomposition temperature (Thigh) hydride with a low decomposition temperature (Tlow) hydride (metal or complex) would reduce

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

255

the decomposition temperature of the high temperature hydride according to the well-known Rule-of-Mixtures (ROM) Tdesorption = ThighVhigh + TlowVlow

(3.44a)

Tdesorption = T 0high – bVlow

(3.44b)

and after rearranging

where Vhigh and Vlow is the volume fraction of high and low temperature hydride, respectively, and T0high is the initial temperature at Vlow = 0. Since volume fraction (V in %) is directly proportional to weight fraction (W in %) (3.44b) requires a linear dependence of decomposition temperature with a negative slope b versus either the volume or weight fraction (Vlow or Wlow) of a constituent hydride having lower desorption temperature. Therefore, after compositing of high and low desorption temperature hydrides by ball milling which, in effect, results in the formation of intimately intermixed nanocomposite hydrides, one would expect that the desorption temperature of a hydride which has the higher desorption temperature decreases linearly with increasing volume fraction of a constituent hydride having much lower desorption temperature in accord with (3.44b). This hypothesis has been experimentally verified for the following composite systems: MgH2–LiAlH4, MgH2–NaAlH4 and NaBH4–MgH2 where the second hydride in a pair is supposed to have much lower desorption temperature than the first one. The composites have been synthesized by either controlled reactive mechanical milling (CRMM) or controlled mechanical milling (CMM) in a magneto-mill which in effect produced nanocomposites with the nanometric grain sizes of the constituent phases and substantial reduction of particle size.

3.5.1

MgH2–LiAlH4 Composite System

Since major decomposition events of LiAlH4 occur at the low temperature range of 150–250°C as discussed in Sect. 3.2.2, this complex hydride is an excellent constituent for testing the ROM behavior as given by (3.44b) of a composite system with MgH2 which has higher desorption temperature. “As-received” commercial MgH2 powder (sold under the trade name MG-5026® from ABCR GmbH&Co.KG; ~ 98 wt% purity; the remaining Mg; see Sect. 2.1) and LiAlH4 (97% purity) from Alfa Aesar were mixed to (MgH2 + Xwt%LiAlH4) compositions where X = 10, 20, 30, 50 and 70. As a reference, MgH2 powder without additives was used. (3.44b) is only a mathematical expression without invoking any physical mechanism. To determine if the catalytic effect of free Al, decomposed from LiAlH4, could be the underlying physical mechanism of the ROM behavior in this system, the mixture of MgH2 and Al (Al = 8, 15, 23, 42 and 62 wt%) was also prepared (the Al amount is equivalent to the amount of aluminum after decomposition of 10, 20, 30, 50 and 70 wt%LiAlH4). Controlled Mechanical Milling (CMM) was carried out for 20 h in

256

3 Complex Hydrides

the argon gas atmosphere under 700 kPa pressure in the magneto-mill Uni-BallMill 5 using strong shearing (HES57–two magnets) mode. Figure 3.26 shows the scanning electron micrographs of the microstructure of composites after milling for 20 h. The microstructure of the composites after milling can be compared with the unmilled and milled LiAlH4 powder in Fig. 3.11. It is clear that up to 50 wt%LiAlH4 the milling refines the particle size of the composite quite effectively. However, at 70 wt%LiAlH4 the particle size becomes much coarser. This is confirmed in Fig. 3.27 which shows the dependence of average particle size (ECD) on the MgH2 content. It is clear that at a small fraction of MgH2 milling becomes less effective. As mentioned in Sect. 3.2.2 ball milling for 20 h of a pure LiAlH4 reduced its particle size from 10.5 ± 4.8 to 5.2 ± 4.3 μm which is marked in Fig. 3.27 at 0 wt%MgH2. At 50 wt%LiAlH4 the composite particle size is reduced to ~1 μm. Figure 3.28 shows the XRD patterns of milled composites. The most interesting feature is the presence of diffraction peaks of Al for all compositions from 10 to 100 wt%LiAlH4. It means that a fraction of LiAlH4 decomposed during milling most likely to Li3AlH6 and Al (the reaction (R1b) of (3.12) in solid state) despite that the milling vials were continuously cooled off during milling by forced air from a fan. As mentioned in Sect. 3.2.2, Resan et al. [92] also observed that undoped/pure LiAlH4 decomposes during milling in contrast to undoped NaAlH4 which after milling under the same conditions did not show any loss in capacity. The presence of diffraction peaks of Li(OH)·H2O is most likely due to the oxidation (or hydrolysis) of a fraction of LiAlH4/Li3AlH6 in the milled powders exposed to air during XRD tests. Figure 3.29 shows the grain size of MgH2 in the composites (calculated form the breadth of diffraction peaks in XRD patterns as shown in Sect. 1.4.3) as a function of LiAlH4 content. At small contents of LiAlH4 the grain size is stabilized at ~ 12 nm. However, when the content of LiAlH4 increases to 30 wt% and more, the grain size of MgH2 jumps almost threefold to ~ 30 nm. This is another indicator of a low efficiency of milling of composites with a high content of LiAlH4. The grain size of Al whose Bragg peaks are present in Fig. 3.28 was estimated to be within 23–61 nm. Figure 3.30 shows the DSC traces for the (MgH2 + 20, 30, 50 and 70 wt%LiAlH4) composites. Only single endothermic peak centered at ~ 350°C is visible in DSC traces for the (MgH2 + 20 wt%LiAlH4) composite (Fig. 3.30a). This peak corresponds to the decomposition of MgH2. The first low temperature exothermic effect observed in Fig. 3.9 for a pure LiAlH4 (both unmilled and milled), which is usually assigned to the interaction of LiAlH4 with hydroxyl impurities [67], is not observed in Fig. 3.30a–c but it appears in Fig. 3.30d for (MgH2 + 70 wt%LiAlH4). Four endothermic events occur for (MgH2 + 30, 50 and 70 wt%LiAlH4) (Fig. 3.30b–d). The first endothermic peak at ~ 174–182°C has almost exactly the same temperature range as (R1a) in Fig. 3.9. No exothermic peak (R1b) of melting from Fig. 3.9 is seen in Fig. 3.30a–d. It seems that the addition of just 30 wt%MgH2 suppresses melting of LiAlH4 and its first decomposition into Li3AlH6 and Al ((R1b) in Fig. 3.9) occurs from a solid phase and is endothermic. This is supported by the observation of partial decomposition of LiAlH4 into (Li3AlH6 + Al) during milling as discussed before. The second endo peak in Fig. 3.30b–d at ~ 198, 193 and 223°C, respectively, corresponds to the decomposition

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

a

b

c

d

257

e Fig. 3.26 Scanning electron micrographs of (MgH2 + Xwt%LiAlH4) composites synthesized by ball milling in the magneto mill Uni-Ball-Mill 5 using HES57 mode for 20 h under 700 kPa argon. (a) X = 10, (b) x = 20, (c) X = 30, (d) x = 50 and (e) X = 70. All micrographs taken at the same magnification. Secondary electrons contrast

of Li3AlH6 into (LiH + Al) ((R2) in Fig. 3.9). The third peak in Fig. 3.30a–d at ~ 353, 335, 313 and 325°C, respectively, corresponds to the decomposition of MgH2 in the composite. Finally, the fourth endo peak at the ~ 425–450°C range in Fig. 3.30a–d is most likely due to the decomposition of LiH although some eutectic melting of Mg–Al cannot be ruled out of hand [120, 124]. The desorption peak temperature maxima for the MgH2 constituent in a composite are plotted in Fig. 3.31 as a function of the content of LiAlH4. It is seen that the ROM for the MgH2 desorption temperature in the (MgH2 + LiAlH4) nanocomposite

258

3 Complex Hydrides 10 9 8 ECD [mm]

7 6 5 4 3 2 1 0 0

20

40

60

80

100

MgH2 [wt.%]

Fig. 3.27 Dependence of the average particle size (ECD) of the composites on the content of MgH2 after ball milling in the magneto mill Uni-Ball-Mill 5 using HES57 mode for 20 h under 700 kPa argon

plotted vs. wt% of LiAlH4 is obeyed up to about 50 wt.% LiAlH4 with an excellent coefficient of the fit R2 = 0.95. At the higher content of LiAlH4 the ROM breaks down most likely due to a very ineffective milling as discussed earlier resulting in the increase of particle size of the composite and grain size of MgH2. As mentioned before in order to determine whether or not the free Al formed upon decomposition of LiAlH4/Li3AlH6 in the composite could act as a catalyst, we also prepared composites with the content of Al equivalent to the content of Al in the Xwt%LiAlH4. Their DSC desorption peak temperature maxima are also plotted in Fig. 3.31. The composites with the equivalent content of Al do not seem to follow the ROM behavior. Therefore, one can tentatively conclude that the underlying physical mechanism for the ROM behavior is not related to the catalytic effect of free Al. However, this possibility, however remote, cannot be completely ruled out of hand because the particle size of free Al formed upon decomposition might be much smaller than that obtained by ball milling of Al metal powder added to MgH2 powder. Nanosized free Al could aquire catalytic behavior. However, at the moment we do not have any evidence for that. Figure 3.32 shows XRD patterns of (MgH2 + LiAlH4) composites after DSC testing up to 500°C. The primary phases present are Mg and Al. Peaks of MgO and (LiOH)·H2O arise from the exposure of Mg and Li (or possibly even some retained LiH) to the environment during XRD tests. Apparently, XRD phase analysis indicates that a nearly full decomposition of original MgH2 and LiAlH4 hydride phases has occurred to the elements during a DSC experiment. In addition, no diffraction peaks of any intermetallic compound are observed in those XRD patterns. That means that no intermetallic compound was formed upon thermal decomposition of composites in DSC. Therefore, the mechanism of destabilization through the formation of an intermediate intermetallic phases proposed by Vajo et al. [196–198] and discussed in the beginning of this section seems to be ruled out of hand.

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing β-MgH2 γ-MgH2 Al

2000

259

MgO Li(OH)•H2O

MgH2 + 30wt.%LiAlH4

1600

MgH2 + 20wt.%LiAlH4

Counts

1200

800

MgH2 + 10wt.%LiAlH4 400

MgH2 0 20

30

a

40

50 60 Degrees 2-Theta

70

80

90

β-MgH2 LiAlH4 Al Li(OH)•H2O

1200

800 Counts

LiAlH4

MgH2 + 70wt.%LiAlH4 400

MgH2 + 50wt.%LiAlH4 0 20

b

30

40 50 60 Degrees 2-Theta

70

80

90

Fig. 3.28 XRD patterns of (MgH2 + Xwt%LiAlH4) composites synthesized by ball milling in the magneto mill Uni-Ball-Mill 5 using HES57 mode for 20 h under 700 kPa argon. (a) X = 10, 20 and 30 and (b) X = 50, 70 and 100 (pure LiAlH4)

Nonetheless, we have calculated the lattice parameters and the unit cell volume of free Mg which is formed from the decomposition of MgH2 and found that it decreases with increasing content of LiAlH4 in a composite up to 30 wt% and then more or less saturates as shown in Fig. 3.33. Such lattice shrinkage is most likely

260

3 Complex Hydrides 60

Grain size of MgH2

50 40 30 20 10 0 0

10

20 30 LiAlH4 [wt.%]

50

70

Fig. 3.29 Grain size of MgH2 in the ball-milled MgH2 + Xwt%LiAlH4 composites as a function of the LiAlH4 content

owing to the formation of a solid solution of Al in Mg, i.e., Mg(Al), during continuous decomposition of LiAlH4 providing free Al and decomposition of MgH2 providing free Mg. The formation of Mg(Al) solid solution is consistent with the magnitude of the atomic radius of Al equal to 0.1432 nm which is smaller than that of Mg equal to 0.1604 nm [121]. Therefore, an important question arises as to whether or not the formation of Mg(Al) solid solution in lieu of an intermetallic or other compound could reduce the enthalpy change of the overall decomposition reaction according to the model proposed by Vajo et al. There is no clear answer to this question. An alternative explanation is based on the reduction of the enthalpy of hydride formation of MgH2 by the elastic interactions with the hydride phases forming from the decomposition of LiAlH4 as proposed by Makihara et al. [202] for the composite system Mg-50 wt%ZrMn2 during its hydrogenation/dehydrogenation. They argued that the lattice contraction due to dehydriding of the composite constituent which dehydrogenates first are transmitted to the second higher desorption temperature constituent by the distortion stress through the interface area with nanocomposite structure. That’s why nanocompositing by ball milling is so important for this mechanism. Then, the lattice contraction is also induced in the second constituent resulting in its instability. Eventually, the hydrogen in the second constituent is discharged cooperatively as soon as that in the first constituent is being discharged. However, in the case reported by Makihara et al. [202] they observed only a single desorption peak from the composite, which was interpreted as a cooperative dehydrogenation. In our case of the (MgH2 + LiAlH4) composites, separate decomposition peaks are observed for the LiAlH4/Li3AlH6 and MgH2 constituents. Therefore, it is not clear if their model also could be applied here. The synthesized nanostructured composites (nanocomposites) (MgH2 + LiAlH4) were subjected to desorption experiments after ball milling for 20 h. Figure 3.34 shows desorption curves obtained under continuous heating up to 300°C (quasi-TPD)

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

Fig. 3.30 Examples of DSC curves for ball milled (a) (MgH2 + 20 wt%LiAlH4), (b) (MgH2 + 30 wt%LiAlH4), (c) (MgH2 + 50 wt%LiAlH4) and (d) (MgH2 + 50 wt%LiAlH4) composite (heating rate 10°C/min, argon flow ~ 50 ml/min)

3.5 261

262

3 Complex Hydrides 400 390 Al

Tpeak MgH2 (ⴗC)

380 370 360 350 340 330

y = −1.2209x + 387.86 R2 = 0.9509

320 310 300 0

20

40 60 LiAlH4 or Al (wt.%)

80

100

Fig. 3.31 Desorption temperature of the MgH2constituent in the (MgH2 + Xwt%LiAlH4)composites as a function of the LiAlH4 and equivalent Al content

4000

Mg(Al) Al MgO Li(OH)•H2O

3000

Counts

LiAlH4

2000 MgH2 + 70wt.%LiAlH4 MgH2 + 50wt.%LiAlH4 MgH2 + 30wt.%LiAlH4

1000

MgH2 + 20wt.%LiAlH4 MgH2 + 10wt.%LiAlH4

0 20

30

40 50 60 Degrees 2-Theta

70

80

90

Fig. 3.32 XRD patterns of (MgH2 + LiAlH4) composites after DSC testing up to 500°C.

under 0.1 MPa hydrogen pressure for composites containing 30, 50 and 70 wt%LiAlH4 and for the reference sample of pure LiAlH4 milled for 20 h. The initial, almost vertical, segments visible in all curves are due to the hydrogen desorption during the first transformation of LiAlH4 into Li3AlH6 and Al ((R1b) of (3.12)). The second segment with a small slope is due to the hydrogen desorption during

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing Unit cell volume of Mg [nm3]

3.5

263

0.0466 0.0464 0.0462 0.046 0.0458 0.0456 0.0454 0

20

40 LiAlH4 [wt.%]

60

80

Fig. 3.33 Unit cell volume of Mg formed upon decomposition of MgH2 in the composites during DSC experiments as a function of LiAlH4 content

8.00 4

Hydrogen desorbed [wt.%]

7.00 3 2

6.00 1

5.00 4.00 3.00

1 - MgH2 + 30wt.%LiAlH4 2 - MgH2 + 50wt.%LiAlH4

2.00

3 - MgH2 + 70wt.%LiAlH4

1.00

4 - LiAlH4

0.00 0

1000

2000

3000

4000

Time [s]

Fig. 3.34 Hydrogen desorption curves under continuous heating up to 300°C (quasi-TPD) under 0.1 MPa hydrogen pressure of composites (MgH2 + LiAlH4) synthesized by ball milling for 20 h (corrected for increasing pressure due to hydrogen temperature change)

the second transformation of Li3AlH6 into LiH and Al ((R2) of (3.13)). The third long segment discernible for the composites, which gradually saturates with time, is due to much slower desorption from the MgH2. Based on the DSC curve in Fig. 3.9b a pure LiAlH4 milled for 20 h should decompose at 300°C through the reactions (R1b) and (R2) of (3.12) and (3.13), respectively, to LiH and Al and in doing so, desorb a theoretical purity-corrected (97%) amount of H2 equal to ~ 7.66 wt%. However, it is seen in Fig. 3.34 that a pure LiAlH4 desorbs about 7 wt%H2 at 300°C. This amount is ~ 0.7 wt% deficient with respect to the theoretical value. This missing amount of H2 is lost, most likely, owing to the partial decomposition of LiAlH4 occurring during milling as evidenced by the presence of the

264

3 Complex Hydrides

diffraction peaks of Al in the XRD pattern of pure LiAlH4 milled for 20 h in Fig. 3.28b. The amount of hydrogen desorbed at 300°C for the composite containing 30, 50 and 70 wt%LiAlH4 is ~ 5.5, 6.0 and 6.5 wt%H2, respectively. The theoretical puritycorrected capacity for the desorption of the LiAlH4 constituent to LiH and Al in the 30, 50 and 70 wt%LiAlH4 composite is 2.3, 3.8 and 5.4 wt%H2, respectively. Apparently, the rest of H2 desorbed by the composites is due to the desorption from the MgH2 constituent. Although, in the strict sense, the desorption curves in Fig. 3.34 are not kinetic curves, since they are registered during continuous heating to a constant temperature of 300°C, it seems that the rate of hydrogen desorption is rather fast. The 30, 50 and 70 wt%LiAlH4 composite desorbs the corresponding maximum amount of H2 within ~ 1,200 s. Furthermore, Fig. 3.35 shows the desorption curves of the ball-milled (MgH2 + 50 wt%LiAlH4) composite during continuous heating up to 250, 260, 275 and 300°C under 0.1 MPa hydrogen atmosphere. Under these conditions the composite desorbs ~ 4.3 and 4.9 wt%H2 at 250 and 260°C, respectively. The purity-corrected amount of hydrogen which could be desorbed from the 50 wt%LiAlH4 constituent in a composite at these temperatures which are higher than the temperatures of the solid state reaction (R1b) of (3.12) and (R2) of (3.13) in Fig. 3.30b, is ~3.8 wt%. Experimentally observed values are larger by about 0.5–1.0 wt% than the theoretical one. This excess H2 could only be desorbed from MgH2. That means that MgH2 is able to desorb at temperatures 250 and 260°C, which are lower than its equilibrium temperature of desorption under 0.1 MPa equal to ~ 275°C. Apparently, MgH2 is thermodynamically destabilized by the second composite constituent LiAlH4. Additional evidence that MgH2 is, indeed, destabilized is provided by the shape of the desorption curves at 250 and 260°C in Fig. 3.35 in which one can see a clearly discernible third 8.00

Hydrogen desorbed [wt.%]

7.00 300ⴗC

6.00

275ⴗC

5.00 260ⴗC 4.00

250ⴗC

3.00 2.00 1.00 0.00

0

1000

2000

3000

4000

Time [s]

Fig. 3.35 Hydrogen desorption curves of ball milled (MgH2 + 50 wt%LiAlH4) during continuous heating up to 250, 260, 275 and 300°C under 0.1 MPa hydrogen atmosphere (corrected for increasing pressure due to hydrogen temperature change)

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

265

long segment due to the slowly desorbing MgH2 constituent in the composite. In other words MgH2 is able to desorb below ~ 275°C at the atmospheric pressure. However, the amount of H2 desorbed from the MgH2 constituent in a composite is still small. Most likely this is the effect of slow desorption kinetics. Finally, Fig. 3.36a shows the desorption curves for the (MgH2 + Xwt%Al) composites with the content of Al metal additive equivalent to the content of free Al formed after decomposition of Xwt%LiAlH4 in the (MgH2 + Xwt%LiAlH4) composites. The (MgH2 + Xwt%Al) composites were ball milled for 20 h. Figure 3.36b shows the dependence of an effective kinetic parameter, k, in the JMAK equation on the equivalent content of Al metal additive. The parameter k was calculated from the kinetic curves in Fig. 3.36a. Apparently, up to 23 wt%Al additive (curves 2, 3 and 4 in Fig. 3.36b) there is gradual improvement in the kinetics of desorption of MgH2. However, at the content of Al metal additive larger than ~ 23 wt% the kinetics dramatically worsens and the k parameter dramatically drops. This effect is in reality owing to a very ineffective milling of the MgH2 constituent in the Al-rich powders (cold welding). At the moment the question whether or not the free Al present in the (MgH2 + LiAlH4) composites exerts any catalytic effect on the desorption kinetics of the MgH2 constituent cannot be conclusively answered.

3.5.2

MgH2–NaAlH4 Composite System

“As-received” NaAlH4 (90% purity; Sigma-Aldrich) and MgH2 powder (~ 98% purity; ABCR Germany) were mixed in the (MgH2 + Xwt%NaAlH4) compositions where X = 0, 10, 30, 50 and 70. Controlled Mechanical Milling (CMM) was carried out for 5 h under 500 kPa pressure of argon atmosphere (99.999% purity) in the magneto-mill Uni-Ball-Mill 5 using strong shearing (HES57–two magnets) mode. Figure 3.37 shows scanning electron micrographs of the “as-received” NaAlH4 powder and the milled (MgH2 + Xwt%NaAlH4) composites (where X = 10, 30, 50, 70). The “as-received” NaAlH4 powder contains large aggregates (Fig. 3.37a) (also see Fig. 3.6a,b). After milling for 5 h, some NaAlH4 particles are refined but large aggregates still exist (Fig. 3.37b). The particle size in composites containing the majority of MgH2 (10 and 30 wt%NaAlH4) are relatively well refined. In the micrographs of milled composites of MgH2 with 50 and 70 wt%NaAlH4, refined particles of the MgH2 constituent and large aggregates of NaAlH4 are both observed (Fig. 3.37e, f). This clearly indicates that the milling of NaAlH4 even in the presence of MgH2 is not very effective for 5 h of duration. Nevertheless, EDS analysis revealed that the levels of Na, Mg and Al are similar in both the small particles and large aggregates. This indicates that NaAlH4 and MgH2 are blended well after 5 h of milling. Hence, it seems that the large aggregates are formed by cold welded small particles containing NaAlH4 and MgH2. Thermal behavior of the milled (MgH2 + Xwt%NaAlH4) composites with X = 0, 10, 30, 50 and 70 was investigated by DSC and TGA. DSC traces of MgH2 powder after milling for 5 h (not shown here) showed an endothermic peak doublet at 373.5 and

266

3 Complex Hydrides

Hydrogen desorbed [wt.%]

7.00 300ⴗC; 0.1MPa H2

6.00

4 5.00 3

4.00

1 - 0 wt.% Al 2 - 8 wt.% Al 3 - 15 wt.% Al 4 - 23 wt.% Al 5 - 42 wt.% Al 6 - 62 wt.% Al

3.00

1

2

2.00

5, 6

1.00 0.00 0

1000

2000

a

3000

4000

Time [s]

0.0012 0.001

k

0.0008 0.0006 0.0004 0.0002 0 0

b

10

20

30

40

50

60

70

Al (wt.%)

Fig. 3.36 (a) Desorption curves for the (MgH2 + Xwt%Al) composites with the content of Al additive equivalent to the content of free Al formed after decomposition of Xwt%LiAlH4 in (MgH2 + Xwt%LiAlH4) composites (300°C; 0.1 MPa H2). The (MgH2 + Xwt%Al) composites were ball milled for 20 h. (b) Dependence of an effective kinetic parameter, k, in the JMAK equation on the equivalent content of Al metal additive

389.0°C while that of “as-received” MgH2 gave one peak at 441.8°C. Figure 3.38 shows DSC traces of composites upon heating to 500°C. Composite with 10 wt%NaAlH4 exhibits only one symmetrical endothermic peak which is mostly due to the decomposition of MgH2. Its temperature of peak maximum is lowered by about 40°C with respect to MgH2 milled for 5 h. Addition of 30–70 wt%NaAlH4 to MgH2 results in the appearance of desorption peaks due to the decomposition of NaAlH4. The first exothermic peak at ~ 170°C possibly is due to the presence of surface hydroxyl impurities in the powder as discussed in Sect. 3.2.1. Its temperature increases very slightly with increasing content of NaAlH4 from 30 to 70 wt%. The next two endothermic peaks at the ~ 182–185 and 236–273°C range correspond to the melting reaction (R1a) of (3.6) and

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

a

b

c

d

e

f

267

Fig. 3.37 Scanning electron micrographs of (a) as-received NaAlH4 (90% purity), and powders milled for 5 h, (b) NaAlH4, (c) MgH2 + 10 wt%NaAlH4, (d) MgH2 + 30 wt%NaAlH4, (e) MgH2 + 50 wt%NaAlH4, and (f) MgH2 + 70 wt%NaAlH4

first decomposition reaction (R1b) of (3.7) for pure NaAlH4 in Fig. 3.7a, b, respectively. It looks like the addition of MgH2 does not suppress the melting of NaAlH4 as opposed to LiAlH4 as discussed in the preceding section (Fig. 3.30a, b). The most interesting is the third endo peak at the temperature range 330–380°C which for pure NaAlH4 corresponds to the reaction (R2) of (3.8) (Fig. 3.7a, b), i.e., decomposition of Na3AlH6 to NaH and Al. By comparing the temperature ranges of the decomposition peak of MgH2 in Fig. 3.38a with that for pure NaAlH4 and the third endo peak for the composites in Fig. 3.38b–d, it is obvious that the desorption range of MgH2 overlaps with the reaction (R2) peak due to the decomposition of Na3AlH6. This is a completely different situation

Fig. 3.38 DSC traces of composites (a) MgH2 + 10 wt%NaAlH4, (b) MgH2 + 30 wt%NaAlH4, (c) MgH2 + 50 wt%NaAlH4, and (d) MgH2 + 70 wt%NaAlH4 ball milled for 5 h. Heating rate 10°C/min at argon flow 50 ml/min

268 3 Complex Hydrides

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Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

269

than that for (MgH2 + LiAlH4) composites where the thermal events for both composite constituents were completely separated. It is interesting to observe that due to this overlap the third endo event for the composites in Fig. 3.38c, d, evolves into peak doublets or even triplets. Most probably, the first peak in a doublet or triplet (Fig. 3.38c, d) is due to decomposition of Na3AlH6 in a composite and the remaining ones are due to the decomposition of MgH2. The last endo peak at the ~ 430–450°C range observed only in DSC curves for the (MgH2 + 30, 50 and 70 wt%NaAlH4) is most likely due to the decomposition of NaH (some eutectic melting of Mg–Al cannot be ruled out of hand [120, 124] as in the MgH2 + LiAlH4 system). The peak temperature of the first endo peak does not change as the amount of NaAlH4 increases in a composite which is expected as this peak is due to the melting of NaAlH4. The position of the second endo peak shifts slightly to lower temperatures with decreasing amount of MgH2 in a composite getting closer to the temperature of this peak for pure NaAlH4 (Fig. 3.7a, b). This may indicate that MgH2 slightly destabilizes the reaction (R1b) of (3.7) (decomposition of NaAlH4 into Na3AlH6 and Al). However, an enormous amount of MgH2 (70 wt%) is required to reduce the temperature of the reaction (R1b) of (3.7) by ~ 60°C from that corresponding to pure NaAlH4 (e.g., compare Figs. 3.38b and 3.7b). In order to find out if the decomposition of NaAlH4 destabilizes the MgH2 constituent in a composite, the peak temperatures of the last peak in a doublet/triplet in Fig. 3.38b–d which supposedly corresponds to the decomposition of MgH2 are plotted as a function of vol.%NaAlH4 in Fig. 3.39. It is quite clear that the ROM behavior for the MgH2 temperature is not obeyed for this composite system in contrast to the (MgH2 + LiAlH4) composite system. There are two factors which may be responsible for this behavior. First, the decomposition temperature of Na3AlH6 overlaps with the decomposition temperature of MgH2. Second, the particle size of the composite powders is still coarse (Fig. 3.37e, f). 400 100%MgH2 Tpeak (ⴗC)(MgH2)

380 100%NaAlH 360 340 320 300

0

20

40 60 NaAlH4 (vol.%)

80

100

Fig. 3.39 Desorption temperature of the MgH2 constituent in the (MgH2 + NaAlH4) composites as a function of the NaAlH4 content

270

3.5.3

3 Complex Hydrides

MgH2–NaBH4 Composite System

In the (NaBH4 + MgH2) composite system the MgH2 constituent has much lower desorption temperature than NaBH4 and decomposes first which may have a destabilizing effect on the NaBH2 constituent. This was the underlying factor why we decided to examine the ROM for this specific system. As shown in Fig. 3.17b pure NaBH4 undergoes melting first at ~ 500°C and finally decomposes at ~ 580°C to the elemental Na, B and H2 as shown by (3.39). “As-received” commercial MgH2 powder (trade name MG-5026® from ABCR GmbH&Co.KG; ~ 98 wt% purity; the remaining Mg) and NaBH4 (98% purity) were mixed to (NaBH4 + Xwt%MgH2) compositions where X = 10, 20, 50, 70, 80 and 90. As a reference, NaBH4 and MgH2 powders without additives were also examined. Controlled Mechanical Milling (CMM) was carried out for 5 h for (NaBH4 + 10 wt%MgH2) and for 20 h for the remaining composites, under 700 kPa pressure hydrogen gas in the magneto-mill Uni-Ball-Mill 5 using strong shearing (HES57–two magnets) mode. To determine if the catalytic effect of free Mg, decomposed from MgH2, could be the underlying physical mechanism of the ROM behavior in this system, the samples of NaBH4 with the content of 12, 55, 73 and 93 wt%Mg were milled under similar conditions like the (NaBH4 + MgH2) composites. To avoid hydrogenation of magnesium during the milling process, hydrogen gas was replaced by argon protective gas. Scanning electron micrographs showing the particle morphology of “asreceived” constituents and milled (NaBH4 + MgH2) composites with low (10–50 wt%) content of MgH2 are presented in Fig. 3.40. The particles of “as-received” NaBH4 and MgH2 powders have irregular shape and average size ~ 145 and ~ 36 μm, respectively (Fig. 3.40a, b). After milling composites with 10, 20 and 50 wt%MgH2 (Fig. 3.40c–e), the NaBH4 constituent in these composites forms sort of “ameba-like” interconnected morphology rather than individual particles. The particles of MgH2 (brighter specs in the micrographs) seem to be slightly refined and uniformly distributed and embedded in a semi-continuous volume of the NaBH4 constituent. Due to this strange morphology of NaBH4 it was hard to calculate any meaningful ECD particle size. An increase in the content of MgH2 to 70, 80 and 90 wt% in the composites makes the ball milling process much more effective. It is seen now in Fig. 3.41 that the composites have very uniformly refined and globular particle size. With increasing MgH2 content, the refinement of the particulate increases (Fig. 3.41a, b) and finally the ECD reaches almost 0.7 ± 0.4 μm for the composite containing 90 wt%MgH2 (Fig. 3.41c). Phase analysis by XRD of the microstructure of composites after ball milling for 20 h is shown in Fig. 3.42. The peaks of NaBH4 are relatively narrow in the XRD patterns. This strongly suggests that the grain size of NaBH4 is not substantially reduced due to an ineffective milling. The peaks of γ-MgH2 appear only in composites with the content of MgH2 higher than 20 wt%. These XRD observations verify SEM observations (Fig. 3.40) that the milling process becomes effective only for

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

a

b

c

d

271

e Fig. 3.40 Scanning electron micrographs showing morphology of as-received powder particles: (a) NaBH4 and (b) MgH2, and the morphology of composites: (c) (NaBH4 + 10 wt%MgH2) (5 h milled), (d) (NaBH4 + 20 wt%MgH2) (20 h milled) and (e) (NaBH4 + 50 wt%MgH2) (20 milled)

composites with high content of MgH2. This behavior is also visible from the analysis of grain size calculated from the peak broadening and plotted as a function of composite composition in Fig. 3.43. Both composite constituents NaBH4 and MgH2 exhibit a substantial grain size reduction only in the composites with the MgH2 content around 50 wt% and more. It is also quite clear that the nanostructuring of NaBH4

272

3 Complex Hydrides

ECD = 1.4 ± 0.9 mm

ECD = 1.0 ± 0.6 mm

a

b ECD = 0.7 ± 0.4 mm

c Fig. 3.41 Scanning electron micrographs showing morphology of powder particles after ball milling for 20 h for (a) (NaBH4 + 70 wt%MgH2), (b) (NaBH4 + 80 wt%MgH2) and (c) (NaBH4 + 90 wt%MgH2) composites. Particle size ECD shown 16000

β-MgH2 NaBH4

γ-MgH2 MgO

12000

Counts

NaBH4 + 90 wt.%MgH2 NaBH4 + 80 wt.%MgH2 8000

NaBH4 + 70 wt.%MgH2 NaBH4 + 50 wt.%MgH2 4000

NaBH4 + 20 wt.%MgH2 NaBH4 + 10 wt.%MgH2 0 30

40

50 60 Degrees 2-Theta

70

80

90

Fig. 3.42 XRD patterns of composites after the synthesis by ball milling for 20 h

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

273

is not as effective as that of MgH2. The smallest grain size of NaBH4 in the composite containing 70 wt%MgH2 is within 40–50 nm range. The coarsest grain size of MgH2 in the composite containing 10 wt%MgH2 is within 35–40 nm range. Present results confirm our previous findings [179] that the formation of true nanostructure in NaBH4 is difficult to achieve and hence, clearly demonstrate a high structural stability of the compound under heavy deformation conditions imposed by milling. Figure 3.44 shows DSC traces of all (NaBH4 + MgH2) composites ball milled for 20 h. The DSC traces of composites containing 10, 20 and 50 wt%MgH2 exhibit three endothermic peaks from which the first is due to the decomposition of MgH2, the second is due to the melting of NaBH4 and the third is due to the decomposition of NaBH4 (Fig. 3.44a–c). The DSC traces of composites containing 70, 80 and 90 wt%MgH2 exhibit only two endothermic peaks, the first due to the decomposition of MgH2 and the second due to the decomposition of NaBH4 (Fig. 3.44d–f). Apparently, the addition of more than 50 wt%MgH2 suppressed melting of NaBH4, i.e., the decomposition temperature of NaBH4 was reduced below its melting point. This is a similar effect to the one observed for LiAlH4 when MgH2 was added (Sect. 3.5.1). One can note in Fig. 3.44a–f, a continuous reduction of melting and decomposition or just decomposition temperature of NaBH4 (when its melting is suppressed by MgH2) with increasing content of MgH2. It is also interesting to note that the decomposition temperature of MgH2 in the 10 and 20 wt%MgH2 composites (Fig. 3.44a, b) is very high being close to 452–458°C most likely due to an ineffective milling as discussed earlier. For comparison, the decomposition temperature of “as-received” MgH2 (ABCR) tested independently is ~ 442°C. In order to confirm that a large content of NaBH4 in a composite hampers effective ball milling of MgH2 and raises its decomposition temperature we also prepared composites in which the MgH2 and NaBH4 powder constituent was pre-milled for 20 and 5 h, respectively, and subsequently, the pre-milled constituents were mixed at the composition of (NaBH4 + 10, 20 and 35 wt%MgH2) by short ball milling for 15 min. These composites were subjected to DSC testing as shown in Fig. 3.45. Table 3.2 compiles the DSC desorption peak temperatures of the MgH2 constituent in composites as observed in Figs. 3.44 and 3.45. For comparison, the DSC desorption peak temperatures for pure MgH2 as received and after milling for 15 min and 20 h (tested separately in DSC; heating rate 10°C/min) are also listed in Table 3.2. It is seen that the decomposition temperature for the MgH2 constituent in the 10, 20 and 35 wt%MgH2 composites which were synthesized from the pre-milled powders is within the range of 418–420°C which is much lower than the 452–458°C range observed in Fig. 3.44a, b for the MgH2 constituent milled together with NaBH4 in the 10 and 20 wt%MgH2 composites. Since the pre-milling of a single-phase MgH2 for 20 h is more effective and results in particle size less than 1 μm, its decomposition temperature in a composite with NaBH4 is obviously lower (see Sect. 2.1 on the properties of ball-milled MgH2). Nevertheless, as can be seen in Table 3.2, still the decomposition peak temperatures of MgH2 in the 10 and 20 wt%MgH2 composites (~ 418–420°C) are higher than the decomposition peak temperature of pure

274

3 Complex Hydrides

Grain size of NaBH4 (nm)

100 90 80 70 60 50 40 0

10

10

20

a

20 50 MgH2 (wt.%)

70

80

45

Grain size of MgH2

40 35 30 25 20 15 10 5 0

b

50 70 MgH2 (wt.%)

80

90

Fig. 3.43 Grain size of (a) NaBH4 and (b) MgH2 as a function of MgH2 content in the (NaBH4 + MgH2) composites ball milled for 20 h (grain size calculated by the Rigaku Rotaflex diffractometer software)

MgH2 ball milled for 20 h equal to ~ 394°C (at the heating rate 10°C/min). This strongly suggests that a large content of NaBH4 in a composite exerts some stabilizing effect on the MgH2 constituent during desorption. However, gradual reduction of the NaBH4 constituent to 50, 30, 20 and 10 wt% also reduces the decomposition temperature of the MgH2 constituent to ~ 413, 363, 358 and 354°C. For the 30, 20 and 10 wt%NaBH4 compositions the decomposition temperature of MgH2 is below the decomposition temperature of MgH2 ball milled for 20 h which can be taken as a reference (394.3°C). Apparently, in the NaBH4-lean composites, the NaBH4 destabilizes the desorption of the MgH2 constituent. However, whether the nature of this destabilization is thermodynamic or kinetic is unclear. It is relevant to point out that we have recently reported [203] a similar destabilization of MgH2 in a nanocomposite with NaBH4 synthesized by reactive mechanical milling of (Mg + Xwt%NaBH4) where X = 2, 10 and 20. In this nanocomposite

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

DSC / (mW/mg) ↓ exo

502.9⬚C

591.6⬚C [1.1]

5

NaBH4+10wt.%MgH2

DSC / (mW/mg) ↓ exo 7 NaBH

275

589.0⬚C 4+20wt.%MgH2

6

4

5 3

4

488.6⬚C

3

2 1

458.0⬚C

[1.1]

2

451.2⬚C

1 0

0

a

100

200

300 400 Temperature / ⴗC

DSC / (mW/mg) ↓ exo 4

500

600

100

200

b

300 400 Temperature / ⴗC

DSC / (mW/mg) 7 ↓ exo

505.7⬚C

NaBH4+50wt.%MgH2 412.9⬚C 467.8⬚C

6

3

500

600

363.2⬚C

NaBH4+70wt.%MgH2

5

2

[1.1]

4 454.1⬚C

3

1

2 0

[1.1]

1

−1

0

c

100

200

DSC / (mW/mg) ↓ exo 4

300 400 Temperature / ⴗC

500

600

100

200

d

300 400 Temperature / ⴗC

DSC / (mW/mg) ↓ exo 3.0

357.9⬚C

NaBH4+80wt.%MgH2

2.5

500

600

354.5⬚C

NaBH4+90wt.%MgH2

2.0

3

1.5 2

1.0

435.7⬚C

421.7⬚C

0.5

1

0.0 −1.5

0 [1.1]

−1 100

e

200

300 400 Temperature / ⴗC

500

[1.1]

−1.0

600

f

100

200

300 400 Temperature / ⴗC

500

600

Fig. 3.44 DSC traces of composites: (a) (NaBH4 + 10 wt%MgH2), (b) (NaBH4 + 20 wt%MgH2), (c) (NaBH4 + 50 wt%MgH2), (d) NaBH4 + 70 wt%MgH2), (e) NaBH4 + 80 wt%MgH2) and (f) NaBH4 + 90 wt%MgH2) (with heating rate 10°C/min and argon flow rate 50 ml/min)

having the particle size of ~ 1 μm, the nanometric MgH2 constituent (12–14 nm grain size) was formed in situ during reactive milling under hydrogen. We observed that the DSC curve for the (MgH2 + 20 wt%NaBH4) nanocomposite showed an endothermic peak triplet with the first small peak at ~ 305°C and the second large peak at ~ 330°C most likely from the decomposition of β-MgH2, and the third smaller peak at ~ 405°C possibly from the decomposition of NaBH4. We concluded that the presence of 10 and 20 wt%NaBH4 in the nanocomposite with β-MgH2 lowers the decomposition temperature of the MgH2 constituent by as large as 100°C. The temperature of melting and desorption peaks of NaBH4 for all composites from Figs. 3.44 and 3.45 as a function of the content of MgH2 constituent is presented in Fig. 3.46a. It is seen that the melting temperature of NaBH4 decreases linearly with increasing content of MgH2. The coefficient of fit is R2 = 0.85. This dependence is rather weak. The decomposition temperature seems not to depend on the MgH2 content up to 20 wt%MgH2 but at larger contents of MgH2 it decreases

100

200

c

0

1

2

3

4

500

600

[1.1]

100

200

417.7⬚C

b

0

2

4

6

8

200

500

519.1⬚C

469.7⬚C

100

600

[1.1]

300 400 Temperature / ⴗC

479.0⬚C 420.1⬚C

Pre-milled and mixed 15min NaBH4+20wt.% MgH2

DSC / (mW/mg) ↓ exo 10

300 400 Temperature / ⴗC

Pre-milled and mixed 15min NaBH4+35wt.% MgH2

DSC / (mW/mg) ↓ exo

300 400 Temperature / ⴗC

392.0⬚C

418.1⬚C

490.8⬚C

578.9⬚C

500

586.0⬚C

600

[1.1]

Fig. 3.45 DSC traces of composites prepared with NaBH4 and MgH2 pre-milled for 5 and 20 h, respectively, and subsequently mixed for 15 min by ball milling. (a) (NaBH4 + 10 wt%MgH2), (b) (NaBH4 + 20 wt%MgH2) and (c) (NaBH4 + 35 wt%MgH2) (heating rate 10°C/min and argon flow rate 50 ml/min)

a

0

1

2

3

4

Pre-milled and mixed 15min NaBH4+10wt.% MgH2

DSC / (mW/mg) ↓ exo

276 3 Complex Hydrides

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Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

277

Table 3.2 DSC peak temperatures of the MgH2 constituent in composites from Figs. 3.44 and 3.45 and for pure MgH2 as received and after milling for 15 min and 20 h

Wt%MgH2 in composite

Wt%NaBH4 in composite

100 100 100 10 10

0 0 0 90 90

20 20

80 80

35

65

50 70 80 90

50 30 20 10

Processing As received Milled 15 min Milled 20 h Milled 5 h Pre-milled 20/5 h; mixed 15 min Milled 20 h Pre-milled 20/5 h; mixed 15 min Pre-milled 20/5 h; mixed 15 min Milled 20 h Milled 20 h Milled 20 h Milled 20 h

DSC desorption peak temperature of MgH2 (°C) (heating rate 10°C/min) 441.8 404.6 394.3 451.2 418.1 458.0 420.1 417.7 412.9 363.2 357.9 354.5

linearly with increasing content of MgH2. The coefficient of fit R2 = 0.98 is excellent. For the 80 wt%MgH2 the decomposition temperature of NaBH4 is reduced by ~ 150°C. It is clear that the presence of the MgH2 constituent in a composite destabilizes the NaBH4 such that its decomposition temperature at the content of MgH2 > 20 wt% obeys the ROM of (3.44b) very well. Since the ROM of (3.44b) is just a mathematical expression, of course, a question arises as to what the underlying physical mechanism is that could be responsible for the ROM dependence of decomposition temperature of the NaBH4 constituent on the content of MgH2? A similar question was raised for the case of the (MgH2 + LiAlH4) composite. The first physical model of interest is the one proposed by Vajo et al. [196–198]. According to this model the enthalpy change of LiBH4 during decomposition is reduced by the formation of an intermediate compound MgB2 from the free Mg obtained due to the decomposition of the MgH2 constituent. This model is presented graphically for the thermodynamic destabilization of LiBH4 by MgH2 in Fig. 3.25b. By analogy, the reaction of (3.43) can be adapted to the NaBH4 and MgH2 system in the following form 2NaBH4 + MgH2 ® 2NaH + MgB2 + 4H2

(3.45)

It requires that the strongest destabilization of NaBH4 by MgH2 should by observed for composite where the yield of MgB2 formed is higher than for the stoichiometric composition (NaBH4 + 25.8 wt%MgH2). In Fig. 3.46b the yield of MgB2 formed versus composition and theoretically predicted products of decomposition are presented. The vertical line shows the stoichiometric composition (NaBH4 + 25.8 wt%MgH2). Decomposition of stoichiometric composite is characterized by

278

3 Complex Hydrides 650 600 R2 = 0.98

Tpeak NaBH4 (ⴗC)

550

(3)

(2)

500 (5)

450

R2 = 0.85

400 decomposition with high content of MgH2 decomposition with low content of MgH2 decomposition with Mg melting with MgH2 melting with Mg

350 300 250 200 0

20

40

60

80

100

MgH2 or Mg (wt.%)

a

Yield of MgB2 (wt.%)

(NaBH4+MgH2) system 40 NaH + MgB2 NaH + MgB2 + B

20

NaH + MgB2 + Mg

0 0

b

20

40

60

80

100

MgH2 (wt.%)

Fig. 3.46 (a) The temperature of melting and decomposition peaks of NaBH4 for all composites from Figs. 3.44 and 3.45 as a function of the MgH2 content. Numbers in parenthesis show the number of multiple data points for this specific composition. Data for the (NaBH4 + Xwt%Mg) mixtures where Xwt% is an equivalent amount of Mg corresponding to the Mg decomposed from Xwt%MgH2 are also included. (b) Analysis of the yield of MgB2 based on the Vajo et al. [196– 198] model adopted for the (NaBH4 + MgH2) system

the presence of only NaH and MgB2. Hypostoichiometric composites should decompose to solid products such as NaH, MgB2 and B. Hyperstoichiometric composites decompose to NaH, MgB2 and Mg. According to the plot in Fig. 3.46b we should observe the strongest destabilization of NaBH4, reflected in the substantial decrease of its decomposition temperature, for hypostoichiometric compositions up to 25.8 wt%MgH2 where the yield of formed MgB2 increases with increasing content of MgH2. In contrast to the Vajo et al. model, we do not observe in Fig. 3.46a any decrease of the decomposition temperature of NaBH4, i.e., its destabilization, in the range up to 20 wt%MgH2. Moreover, we observe a strong destabilization of

3.5

Destabilization of High Desorption Temperature Hydrides by (Nano)Compositing

a

b

c

d

279

Fig. 3.47 Scanning electron micrographs of the morphology of powder particles after ball milling for (a) NaBH4 + 12 wt%Mg, (b) NaBH4 + 55 wt%Mg, (c) NaBH4 + 73 wt%Mg, (d) NaBH4 + 92 wt%Mg composites

NaBH4 in hyperstoichiometric range where the content of MgH2 increases and that of MgB2 decreases (Fig. 3.46a, b). It is seen by comparing Fig. 3.46a with 3.46b that the decrease of the decomposition temperature of NaBH4 coincides with the appearance and increase of Mg in the hyperstoichiometric range of compositions. This suggests that Mg may act as a sort of catalyst destabilizing the decomposition reaction of NaBH4. As mentioned earlier, we decided to probe this hypothesis by preparing mixtures of NaBH4 with 12, 55, 73 and 92 wt%Mg which were ball milled for 20 h. The content of magnesium in these composites is equivalent to the amount of magnesium as a product of decomposition of the equivalent amount of MgH2. The scanning electron micrographs of the morphology of the ball-milled mixtures with Mg are shown in Fig. 3.47. XRD phase analysis in Fig. 3.48 shows that these ball-milled mixtures consist of NaBH4, Mg, and surprisingly, some amount of β-MgH2. Since ball milling was carried out under argon the only explanation of the presence of β-MgH2 is that elemental Mg reacted with NaBH4 during milling. Two possible reactions could be suggested as follows: NaBH4 + Mg = MgH2 + NaH + a – B + 1/2H2

(3.46a)

NaBH4 + Mg = MgH2 + Na + a – B + H2

(3.46b)

280

3 Complex Hydrides 40000

β-MgH2 NaBH4

Mg(OH)2 Mg

Counts

30000

20000 NaBH4 + 92 wt.% Mg NaBH4 + 73 wt.% Mg 10000 NaBH4 + 55 wt.% Mg NaBH4 + 12 wt.% Mg 0 20

30

40 50 60 Degrees 2-Theta

70

80

90

Fig. 3.48 XRD patterns of (NaBH4 + Xwt%Mg) mixtures ball milled for 20 h where X = 12, 55, 73, 92. The content of magnesium in these composites is equivalent to the amount of magnesium as a product of decomposition of 10, 50, 70 and 90 wt%MgH2

where a-B is an amorphous boron. First reaction would require the presence of NaH in the XRD pattern which is not observed. Second reaction produces Na which usually should oxidize to Na2O having crystalline antifluorite structure (CaF2) [204]. Unfortunately, no diffraction peaks of crystalline Na2O can be discerned in the XRD patterns in Fig. 3.47 unless one assumes that Na2O became amorphous. Hence, it is hard to give any conclusive answer as to what kind of reaction could be responsible for the presence of β-MgH2 in the microstructure of (NaBH4 + Mg) composites after ball milling for 20 h. The argument for the existence of β-MgH2 is additionally reinforced by the presence of Mg(OH)2 which, most likely, was formed due to the hydrolysis of a part of β-MgH2 as a result of the exposure of powder to moist air during XRD tests. The grain sizes/strains of Mg were estimated from the peak broadening analysis and are equal to 103 nm/0 strain, 104 nm/2.12 × 10−4 and 88 nm/1.5 × 10−4 for the 55, 73 and 92 wt%Mg, respectively. DSC experiments were conducted up to 500 and 600°C for the ball-milled mixtures of NaBH4 with 12, 55, 73 and 92 wt%Mg (DSC traces not shown here). The temperatures of melting and decomposition peaks of NaBH4 are plotted in Fig. 3.46a. It is seen that the data points for the equivalent content of 12, 55 and 73 wt%Mg follow quite closely the data points for MgH2. Only data points for the equivalent content of 92 wt%Mg show slightly higher decomposition temperature than their MgH2 counterpart and lie slightly above the ROM line for MgH2. This behavior is probably related to different mechanical properties of Mg and MgH2 which results in different abilities to particle and grain size reduction during the milling process as can be seen in Fig. 3.47. The particle size reduction of ductile

References

281

magnesium is more pronounced for the powders with the low or medium content of Mg. The morphology of these powders is characterized by the presence of clusters formed with smaller particles (Fig. 3.47a–c). For the composite with the highest content of 92 wt%Mg in Fig. 3.47d, the powder particles (ECD ~ 9.2 μm) are practically identical to the “as-received” powder (ECD ~ 9.5 ± 5.5 μm), i.e., milling was rather ineffective. Assuming that Mg exerts some destabilizing effect on NaBH4, in the case of (NaBH4 + 92 wt%Mg) mixture the intimate contact between Mg and NaBH4 was insufficient and the destabilizing effect is less pronounced than that for the (NaBH4 + 90 wt%MgH2) counterpart. The evidence seems to indicate that the destabilization of (NaBH4 + MgH2) composites is somehow related to the presence of free Mg decomposed from MgH2. It is quite likely that Mg has a catalytic effect on the decomposition process of NaBH4. More research is needed to understand the fundamental mechanisms responsible for this peculiar behavior. In summary, the composite approach to complex hydrides renders some interesting results. However, although the observed absolute magnitude of the reduction of desorption temperatures in the investigated (MgH2 + LiAlH4) and (NaBH4 + MgH2) composite systems is quite large, they are still inadequate for any practical storage systems. More vigorous research is needed in this area to identify other systems and possibly catalysts which could further improve the observed composite behavior to an even larger extent.

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Chapter 4

Carbons and Nanocarbons

Carbon atoms can bond to one another in chains, rings, and branching networks to form a variety of structures, including minerals, solid carbon materials, fossil fuels, organic compounds and the large molecules essential to life. In solid carbon materials, atoms form single C–C bonds with each others, like in diamond, or bonds with the rate higher than one, like in graphite. In respect to crystalline structure there are three carbon polymorphs: • Diamond, with sp3 hybridized bonding • Graphite, with sp2 bonding • Amorphous carbon, with a mixture of bonding. There are other two ordered allotropes, which constitute the backbone of new carbon chemistry and nanotechnology: • Fullerenes, with perturbed sp2 bonding • Carbon nanotubes (CNTs), with sp2 bonding. To complete the list of forms in which elemental carbon was found, we must also list two other forms of carbon: • Graphene, exfoliated single sheets of atoms with sp2 bonds • Carbyne, with sp hybridized triple bonds.

4.1

Diamond and Nanodiamonds

In general, C–C bonding involves only the valence electrons. A brief recap of chemistry courses would be useful to the further discussion. The valence electrons of carbon atoms fill in the 2s and 2p orbitals, which represent electron distribution in 3D space as the square of amplitude of their wave function (Fig. 4.1). As the valence electrons s and p are not equivalent to one another, a molecule that is nonsymmetrical and energetically disadvantaged would result if the electrons fill their nonequivalent orbitals. Hence, one of the 2s electrons is “promoted” to the 2p level, which results in four half-filled atomic orbitals. The s orbital combines R.A. Varin et al., Nanomaterials for Solid State Hydrogen Storage, DOI: 10.1007/978-0-387-77712-2_4 © Springer Science + Business Media, LLC 2009

291

292

4 z

z y

–

x

+

Carbons and Nanocarbons z

+

+

y

–

y x

x

– py

px 2s

pz

2p

Fig. 4.1 Directional 2p orbitals, which determine direction and strength of carbon bonds

C

C

C

C C

Fig. 4.2 The sp3 hybridized orbital as an origin of tetrahedrally directed chemical s-type bonds in diamond

with the p orbital to make four equal sp3 hybridized, i.e., mixed, orbitals each with radially-symmetrical s-type character and some directional p-type character, as in Fig. 4.2. The geometry of the sp3-hybridized orbital is tetrahedral with 109.5° angles and high symmetry. Each hybrid orbital will contain one electron and it will overlap head-on with other carbon one-electron orbital forming two-electron s-type bonds with other carbon atoms. The bonding with two electrons per bond gives the bond rate 1 for diamond. This type of bonding takes place in diamond and diamondoid nanophase carbons. To follow space filling, and maintain directions in space forced upon electrons by arms (lobes) of sp3 orbitals, the diamond structure, Fig. 4.3a, has four nearneighbor carbon atoms bonded by s bonds, separated by a distance of 0.15445 nm. The high strength of the s bonds between carbon atoms determines the hardness of diamond. The resultant crystalline lattice of interlocking covalent bonds of the strong s-type makes the diamond very rigid, chemically not reactive, and the hardest material on earth. Since the valence electrons are strongly localized in sp3 orbitals, they cannot move under applied potential, and diamond has very poor electrical conductivity. Since the electrons are tightly bonded in s bonds between two carbon atoms, such bonds absorb light in the ultraviolet region and not in the visible region; hence pure diamond appears clear to the human eye.

4.1

Diamond and Nanodiamonds

293 C C C

C C

b

a

Fig. 4.3 (a) Crystal structure of diamond and (b) the smallest nanodiamond: adamantine

H

H

H

H

Fig. 4.4 The sp3 hybridized orbital on single carbon atoms terminated by four hydrogen atoms forms molecule of methane – CH4

Each lobe of the sp3 orbital containing one-electron overlap head-on carbon atoms and translational repetition of this tetragonal C–C4 cluster makes diamond. The three tetrahedral arms of orbital can also be terminated by hydrogen atoms that overlap head-on. If terminated by hydrogen, the tetrahedral C–H4 cluster makes the CH4 molecule of gas methane (Fig. 4.4). Like a single carbon atom capped with tetrahedrically coordinated hydrogen atoms, Fig. 4.4, a cluster of sp3-bonded carbon atoms can also be capped with hydrogen to form hydrogenated fragments of a diamond structure: diamondoids. Diamondoids are hydrogen-terminated nanodiamonds with sp3 hybrid bonding between the carbon atoms. The smallest of these hydrogenated nanodiamonds is adamantane (Figs. 4.3b and 4.5). Diamondoids can also be seen as cage-like saturated hydrocarbon molecules that possess a rigid carbon framework of diamond-like sp3-bonded carbon atoms. Lower-order

294

4

Carbons and Nanocarbons

H H

H

H

H

H H H

H

H H

H

Fig. 4.5 The sp3 hybridized orbitals on a cluster of ten carbon atoms terminated by hydrogen atoms form molecule of the smallest of nanodiamonds: adamantane C10–H16

diamondoids, e.g., adamantane (C10H16), diamantane (C14H20), triamantane (C18H24), and anti-tetramantane, (C22H28) have been synthesized in the laboratory. The homologous polymantane series has the general molecular formula C4n-6H4n+12, where n = 1, 2, 3, … (n = 1 for adamantane). The interest in those carbon nanophases has recently been renewed with the report by Dahl et al. [1] of Chevron-Texaco on the isolation of new higher diamondoids from petroleum oil. They provide a new class of carbon nanostructures that are markedly different from fullerenes and CNTs, which are discussed later in this chapter. Given the rigid structures and high hydrogen content of diamondoids, they might be potential building blocks for various applications in nanotechnology and of possible interest for hydrogen storage. Although the hydrogen content in solid diamondoid is as high as in petroleum hydrocarbons, the storage is through chemisorption, with the formation of strong s bonds, and the reversibility of the process is problematic. Little has been done so far to investigate desorption of hydrogen from diamondoid nanophases. It has been shown that through rather drastic graphitization treatment at 1,600°C, nanodiamond particles are converted to nanographite [2], but so far there has not been a report on hydrogen sorption properties of this very new nanophase.

4.2 4.2.1

Graphene, Ordered Graphite, and Nanographites Graphene

Single planar sheet of sp2-bonded carbon atoms corresponds to one hexagonal basal plane of graphite and is termed graphene. Recently, a successful attempt to isolate such a graphene layer has been reported [3], and apparently, it becomes possible to produce

4.2

Graphene, Ordered Graphite, and Nanographites

295

a new graphene phase by exfoliation or mechanical peeling of layers from bulk graphite. If separated completely, graphene layer should provide the specific surface area (SSA) of 2,622 g/m2 for two sides of the graphene [4]. Since such layers have finite and small dimensions, filling the empty sp2 orbitals with electrons from hydrogen atoms can stabilize dangling bonds along the periphery of graphene particles. The high degree of hydrogenation to the formula C62H20 stabilizes the graphene. 4.2.1.1

In-Plane s and Out-of-Plane p Bonding

Each graphene sheet is composed of rings of carbon atoms arranged on a hexagonal 2D tiling. This form of carbon has p bonds in addition to the s bonds, as shown in Fig. 4.6, leading to a bond order of 1.33. To form two distinct types of bonds, the s orbital and two of the p orbitals on each carbon have been mixed; thus the hybridization for each carbon is sp2. The three sp2 hybrid orbitals arrange themselves in three-dimensional space to get as far apart as possible. This can only be realized in trigonal planar geometry, where the bond angle between the hybrid orbitals is 120°. While two p orbitals, px and py, along x and y Cartesian axis mix to form three lobes of the electron density in the xy plane, the third unmixed pz orbital will be perpendicular to xy plane. Following the directions of the electron density lobes in 2D planes, carbon atoms arrange themselves into a structure where there are only three covalent bonds per carbon atom. This leads to a hexagonal structure, with the hexagons being arranged in a

b

π C

H

σ

C

a

C

π

a x b plane C

C

Fig. 4.6 sp2 hybridized orbitals on two C atoms that determine in-plane s bonding of carbon into hexagonal rings, viz., graphene plane; p electron charge distribution is shown above and under the plane; such electronic structure determines the crystallographic structure of graphitic materials. Possible scenarios for bonding of hydrogen atoms are via (a) strong covalent bond with participation of half-empty in-plane s orbitals, viz., dangling bonds, (b) more weak bond with participation of interplanar p orbitals, and (c) weak van der Walls bond, as discussed in Sect. 4.2.1.2

296

4

Carbons and Nanocarbons

flat a × b plane. Each carbon atom in Fig. 4.6 contributes one electron to each of three equivalent s bonds. The fourth electron is forced out of plane to occupy unmixed pz orbital, and its electron density shifts to above and below the hexagonal rings of atoms. Since the lobes of pz orbitals centered on neighboring atoms overlap side-by-side, above and below the hexagonal a × b planes, such overlapping forms another set of weak p bonds, in addition to strong in-plane s bonds. The p electrons are delocalized in the interplanar space and are free to move collectively in the ring (like electrons in benzene ring) and between the rings, hence determining good electric conductivity of graphite. Short 2D graphene sheets when stacked along c-axis normal to hexagonal basal plane produce 3D nanographitic structures. We will discuss such structures later in this chapter. When the hexagonal ordering of C atoms extends in a × b basal planes, and the stacking of planes extends along c axis normal to basal plane, then such a massive stacking makes graphite.

4.2.1.2

Van der Walls Interplanar and Intermolecular Interactions

However, this is not the p bonding between sheets that keep graphene sheets together. The lobes of p orbitals that belong to two consecutive sheets of carbon atoms do not overlap. The delocalized p electrons repeal each other, keeping the sheets at a fair distance. The sheets of atoms are arranged parallel to each other and kept together with dominant Van der Waals (vdW) weak forces [5]. These are attractive forces, which keep graphene layers well-stacked, and are the same intermolecular forces that play an important role in the physisorption mechanism of gasses and hydrogen on carbon surfaces. Van der Walls bonds stem from intermolecular forces between atoms (and/or molecules) that result from instantaneous charge distribution in atoms and molecules when they approach each other. Under mutual interaction, an asymmetric polarization of electron charge is induced in molecules that create temporarily dipole moments, and atoms or molecules become attracted by electrostatic forces (Fig. 4.7). They are always present between adjacent molecules, but they are usually rendered insignificant by stronger chemical (covalent) bonds. The vdW interactions are important because they not only determine bonding between layers of graphite but also physisorption of hydrogen on carbon surfaces. Their energy is of the order of a few kilo joules per mole. If no stronger bond exists between C or C and H atoms, then these atoms attract each other in the manner described in quantum chemistry as dispersive interactions with a potential VC–C or VC–H proportional to r−6. However, when they approach each other on the distance smaller than the sum of their radii (so-called van der Walls radii) the atoms are repulsed rapidly and proportionally to r−12. The “r6” potential for dispersive forces was proposed by London, and the overall “6–12”potential for dispersive van der Walls interaction is known as Lennard-Jones (LJ) potential: ⎡⎛ r0 ⎞ 12 ⎛ r0 ⎞ 6 ⎤ VLJ (C - X) = e ⎢⎜ ⎟ − ⎜ ⎟ ⎥ ⎝ r ⎠ ⎥⎦ ⎢⎣⎝ r ⎠

(4.1)

4.2

Graphene, Ordered Graphite, and Nanographites

EC−H2 ≈

297

aC • aH2 r6

αC

δ+

αH2

δ−

δ+

δ−

attraction qC

qH2

Fig. 4.7 Van der Walls bonding as originating from the polarizibility of electron density on hydrogen molecule and carbon atom

where e is the energy of attraction between C and X atoms. Therefore, there is a certain equilibrium distance r0, at which the dispersive–attractive and the repulsive forces balance and the system achieves minimum energy at the minimum of potential curve V(r0). The van der Walls radius, r0, for the C–H interaction can be assumed to be about 0.16 nm.

4.2.1.3

Physisorption of Hydrogen on Carbons

Many carbon-based systems such as graphitic plates, fibers, and wires aggregate with the use of van der Waals bonds. This weak bonding determines packing in bundles of CNTs, as discussed later in this section. Since surfaces of solid carbons interact with each other, and also form bonds with gasses, such weak interactions provide a challenge for quantum-chemical theoretical calculations. Indeed, the dispersion contribution to energy of C–H system is an effect of correlation between moving electrons, and methods that do not deal with such correlation, such as the Hartree-Fock (HF) method, have no chance to grasp the problem. Very popular density functional theory (DFT) calculations so far have hardly been able to properly account for hydrogen physisorption on nanocarbons, since weak dispersion interaction are missed in present DFT theory. Calculations of such interactions were recently tackled by other theories that include both weakly overlapping orbitals and dispersion interaction. It was determined that the energy of hydrogen molecule (H2) on flat graphitic platelets was between 4 and 7 kJ mol−1, depending on the orientation of H2 and on the particle size [6]. As for all electrostatic interactions, the strength of vdW bond should depend on charges, which in turn depend on the number of electrons per atom or molecule; hence the vdW interaction should be stronger between carbon atoms than carbon and hydrogen atoms and even stronger for interactions between atoms of multielectron

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noble gasses. This is not included in LJ potential and, in case of two charges, let us say qC centered on C atom and qH on H2 molecule, a term of type qCqH /r should be 2 2 added. The dispersion energy between hydrogen molecule and the carbon substrate, responsible for physisorption, is determined by the polarizibilities aC, aH (Fig. 4.7) 2

EC - H 2 ≈

a C ·a H2

(4.2)

r6

Since polarizibilities of hydrogen molecules do not vary the only way to increase physisorption is to increase the polarizibility of carbon p orbitals (Fig. 4.6). Unfortunately, the dispersive energy –e that corresponds to the minimum in the VLJ potential curve is low, on average 5 kJ/mol. In consequence the physisorption occurs by very weak bonding, whereby the hydrogen desorbs from graphitic surfaces at temperatures below room temperature. The 5 kJ/mol should be more than doubled to achieve ca. 10–50 kJ/mol range, which is now considered optimal value for reversible hydrogen storage at a temperature ca. 100°C. However, this is not easy to break unfavorable rules imposed by thermodynamics and weak bond strength that renders the enthalpy too low for efficient hydrogen storage by physisorption at ambient temperatures. There is also another topological limit on C:H ratio. Physisorption proceeds only in the monolayer above the boiling point of hydrogen, and the adsorption must follow the Langmuir isotherm. Hence, the storage capacity depends on the pressure of H2 gas above the carbon surface at a fixed temperature. It is also greatly limited by available SSA of carbon phase as provided by porous structure.

4.2.1.4

Chemisorption of Hydrogen on Carbons

Excess sorption of hydrogen, above the limits set by Langmuir isotherm and physisorption, is possible by chemisorption, i.e., absorption of H2 gas with the formation of strong covalent bonds between C and H atoms, as those seen in the diamondoids (Sect. 4.1). Highly defective structure of nanostructured graphite features a large fraction of broken bonds, i.e., carbon dangling bonds. They are good trapping sites for hydrogen storage. When 1s H orbital overlaps empty 2sp2 C orbital, a strong bond is formed, like those that were formed by the involvement of 2sp3 C orbital in the diamondoids, as discussed in the previous paragraph. Chemisorption is at play on the periphery of graphene plates (Fig. 4.6, position a). The major chemisorption mechanism of hydrogen bonding in hydrofullerenes is discussed in Sect. 4.4.1, and occurs as well in hydrogenated CNTs, Sect. 4.4.2, which exhibit a peak of desorption above 400°K [7]. As well as saturating the broken s bonds in graphene sheets, the electron from 1s H orbital can also interact with p electron density adjacent to the surface of graphene plate (Fig. 4.6, position b). In fact, if p bonding were to be fully utilized, every carbon atom could be a site for chemisorption. Yet, the deuterium labeling, neutron scattering, nuclear magnetic resonance (NMR), and electron energy loss spectroscopy (EELS) studies show that

4.2

Graphene, Ordered Graphite, and Nanographites

299

it is only a fraction of sites where D or H atoms are localized between the graphene layers of graphite, with the formation of p bonds [8–11]. However, because of the strength of chemical s + p bonds, which are much stronger than vdW interactions, the sorbed hydrogen can only be released at higher temperatures. Therefore, chemisorption of hydrogen in graphitic nanocarbons, as it is now understood and difficult to control, is not to be utilized in reversible hydrogen storage.

4.2.2

Graphitic Nanofibers, Whiskers, and Polyhedral Crystals

Graphitic polyhedral crystals have been reported with a high degree of order and small interplanar spacing, but not flat layers. The graphite layers wrap consistently around a polyhedral core to form a crystal-shaped structure. Graphite fibers and whiskers have a tree-ring concentric cylinder morphology, which is similar to the polyhedral graphitic particles but having very high aspect ratio and a fibrous shape. They consist of tubular graphene planes. Activated carbon fibers have been prepared by controlled pyrolysis of various structures, e.g., the synthetic polymer polyacrylonitrile (PAN) or coal-tar pitch. Graphitic carbon nanofibers (GNFs), or shortly carbon nanofibers (CNFs), have been catalytically synthesized. They consist of graphene planes arranged in platelet stacks, either in parallel or in angled arrangements that result in a conical fishbone (herringbone) structure. Very high hydrogen uptake values were claimed for these materials, up to H:C ratio equal to 24 (what gives 68 wt.% H2) in GNF herringbone fibers and 55 wt.% H2 in platelet stacks at 278 K and about 12 MPa [12]. These results have never been confirmed in other laboratories, and the group who made the claim have lowered the expectations for research breakthrough that was raised by the original paper, and reported much lower capacities at about 4 wt.% H2 at 6.5 MPa and room temperature [13]. Other studies came up with much lower storage capacities. Hydrogen storage properties of CNFs are determined by the SSA. Nanofibers synthesized via the decomposition of hydrocarbons from gas phase over the Ni70-Cu30 catalyst, exhibited SBET = 143 m2/g and small adsorption of 0.02 wt.% at 273 K and 750 torr H2 [14]. Even long exposure to 10.5 MPa high pressure of hydrogen provided only half-a-percent hydrogen adsorbed on ex-ethylene nanofibers (T decomposition = 450°C and catalyst Ni80-Cu20). Conclusion from a recent 2007 viewpoint paper by Chahine and Bernard is that “currently, physisorption of hydrogen undoped carbon nanostructures and MOFs falls short of the DOE targets for the use of hydrogen as transportation fuel” [15].

4.2.3

Graphite

Graphite is a polymorph of carbon consisting of stacks of graphene sheets of carbon atoms. The stacking of hexagonal basal a × b planes along c-axis changes 2D graphene to 3D graphite. The crystallographic unit cell for graphite is hexagonal (Fig. 4.8).

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H

Van der Walls bonds c-axis

covalent bonds σ +π

Fig. 4.8 Structure of graphite, pointing to 3D hexagonal unit and covalent bonds between carbon atoms, and van der Walls bonds between graphene planes and hydrogen

The theoretical density of graphite is 2.25 g/cm3. Mechanical milling can reduce this density to as low as 1.85 g/cm3 [16]. Highly oriented pyrolytic graphite (HOPG) is a synthetic material with nearly perfect planar structure that is very ordered with minimum interlaminar spacing. Bulk synthetic graphite exhibits BET surface area, SBET = 1–7 m2/g, but this increases to more than 100 m2/g in its chemically activated form. The BET surface area (SBET) is the surface area of the sorbent according to the BET model formulated by Brunauer and Emmett-Teller for planar surfaces, and the equation is formulated to assess multilayer adsorption of small inert molecules on substrate surface, including this that belongs to macropores (>50 nm). Added should be the surface area, St, which is derived from the mesopores (>2 nm), i.e., the amount of surface area excluding the macropores. With the total area of 204 m2/g the measured H2 adsorption on activated graphite was 14 mL H2/g in 77 K [17], which comes to only 0.125 wt.% H. The interplanar distance in regular graphite, 0.3354 nm, is too small for the insertion of gas hydrogen molecules, with a diameter of 0.406 nm. Even if the bulk graphite is exfoliated to multitude of isolated graphene sheets with the surface area of 1,315 m2/g, the amount of hydrogen stored by physisorption at cryogenic temperatures should not exceed 3 wt.% H [18]. Therefore, if higher value of storage capacities were reported, they must be attributed to a concurrent chemisorption. Such chemisorption usually occurs through the formation of very strong sigma bonds, which break and release hydrogen gas only at high temperatures (as discussed in Sect. 4.2.1.4). In view of its electronic structure, carbon atoms in graphite are strongly bonded in planes, and only weak van der Waals forces hold the planes together. This leads

4.3

Disordered and Active Carbons

301

to a material that can shear easily. It is also very brittle, because the s + p bonds hold the atoms of carbon very strongly in the lattice sites, and if strained, these bonds break and the crystal fractures. In fact, the graphite stands at the opposite side of scale of material hardness to the diamond. In view of such an easy sliding on hexagonal basal planes, and susceptibility to brittle fracture, graphite presents a great potential for production of disordered carbon phases and nanophases.

4.3

Disordered and Active Carbons

4.3.1

Disordered Graphites and Mechanically-Activated Carbons

Disordered graphites and nanographitic carbons feature low crystallinity as can be inferred from X-ray diffractograms. Usually, XRD patterns exhibit only strong peak for 00l Bragg reflections from well-stacked and equidistant hexagonal basal planes. The other interferences, caused by in-plane periodic structure, are weak. A number of models were proposed to explain the peculiar shape of XRD reflections. They have attempted to interpret the diffraction peak positions and shapes through partitioning a disordered graphite into a series of stacks of N equidistant graphene layers, where N is two, three, or more layers (Fig. 4.9). Certain assumptions have to be made as to statistical distribution of interlayer distances d between parallel basal planes, as to fit XRD patterns. Some models also reflected on the fact that graphene stacks can be tilted in respect to each other by the angle fi in Fig. 4.9; c-axis

ϕi c1

[001]

B

d1

N1

A

d2

B A N2

B

c2

A B

Ni

A

ϕi

ci

La La (i)

Fig. 4.9 Stacking of hexagonal basal planes in graphite (left). Mechanical activation results in a structure with proliferation of faults in the plane stacking (right). The change of stacking sequences in nearby subgrains is marked by arrows. The stacking fault disorder is corroborated from the nonuniform peak broadening or absence of hk indexed peaks in XRD pattern (see Sect. 1.3.3.3)

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consequently the basal c-axis changes direction from one stack to another [19]. This is illustrated in Fig. 4.9. The origin of the partitioning of massive graphitic structure into a mosaic of nanocrystalline stacks comes from easy shearing of planes, bonded only by weak vdW interactions. Forward shearing of hexagonal basal planes results in generation of stacking fault disorder, as seen in Fig. 4.9. The stacking fault disorder was recently postulated to explain improved insertion of H-ion into the structure of other layered compound: Hexagonal layered Ni hydroxide used in rechargeable alkaline batteries [20]. We have discussed it in Sect. 1.3.3.3. Stacking faults that originate in a mechanical shear of weak bonded graphene planes are deformation faults. Disordered graphites may also exhibit distinct growth faults created during the growth of a layered stack of planes. This view of disordered compounds based on deformed hexagonal crystal, whose planes are bonded only by week interactions, has been proposed in respect to insertion of Li+ and H+ ions in graphitic carbons and layered nickel hydroxide compounds, respectively [21, 22]. It was demonstrated that simple mechano-chemical activation of layered compounds in high-energy ball mills, such as those discussed in Sect. 1.3.2, is a simple yet unusually effective method for improvement of insertion of hydrogen ions into the structure (also seen in Fig. 4.9) for reversible electrochemical hydrogen storage [23, 24]. Another kind of faults are disclinations, and plane rotations, marked in Fig. 4.9 by angles fi and øi. The graphene layers are shown here as being planar but, in fact, these are rather wavy, curled planes. When this material has curled, twisted, and rotated graphene planes, it is called tubostratic graphite. The already weak bonding in graphene stacks become even weaker when the graphene layers are short. With weak attraction between short planes, the average interplanar distance along c-axis increases from 0.3354 nm for ideal graphite to about 0.36 in turbostratic graphite. The loosely bonded planes can rotate. This is where the term turbo (means: rotation) comes from. In turbostratic graphite, the atoms are arranged in layers, as in graphite, but stacked randomly instead of the ABABABA… sequence of graphite. The interplanar spacing, d, changes randomly, although in thin stacks it does not vary much along a short distance of c-axis (Fig. 4.9, right). Since such varying spacing cannot be resolved by X-ray diffraction, the diffractogram exhibits only (00l) peaks that are an average of these spacing. Such disordered graphites and nanographites can easily be produced by ball milling in high-energy mills, as discussed in Sect. 1.3. Hydrogen absorption in disordered graphite produced by ball-milling and shearing of layered structure of graphite was studied in the low temperatures from 35 to 110 K and at pressures up to 20 MPa. Between 5.9 and 6 wt% H was reported to be absorbed on cooling to 75 K [25]. Others reported that hydrogen absorption in mechanically activated nanostructured graphite reaches up to 7.4 wt.% H, where graphite is ball-milled in H2 at 1 MPa pressure in room temperature. Also, for the deuterided mechanically-activated graphite, prepared under 1 MPa D2 by ball milling for 80 h, two desorption peaks were observed, starting at about 600 and 950 K, indicating that there were two types of C–D type of bonding that differ in their strength [26]. A laser desorption time-of-flight mass spectroscopy experiment used

4.3

Disordered and Active Carbons

303

for characterization of a sample of graphite ball-milled with H2 for 100 and more hours suggest the formation of same hydrogenated carbon clusters, with C:H ratios 1.25 and 3.5 in addition to the majority nanostructured graphite phase [27]. Apparently, the effects of mechano-chemical activation of graphite in ball mills are not limited to promoting reversible hydrogen storage via physisorption only but also induce irreversible formation of carbon–hydrogen molecules via chemisorption. The fraction of hydrogen bonded in CHx molecules is not accessible for reversible storage under ambient conditions. The proliferation of stacking faults, edge dislocations, and other defects in disordered graphites result in many dangling bonds determined by unoccupied sp2 orbitals. If the carbon sp2 orbitals are overlapping with hydrogen s orbital, as in Figs. 4.4–4.6 (site a), the hydrogen storage would occur by chemisorption, instead of physisorption. So it is tempting to increase hydrogen storage in ambient temperatures by providing substrates with large density of empty orbitals. This can occur at the periphery of sp2-bonded graphene stacks as obtained by mechanical milling of graphite. The role of defects in changing physisorption to chemisorption can not only be limited to disordered carbons and nanocarbons but also discussed in respect to highly ordered nanocarbons, such as fullerenes and CNTs discussed in the following paragraphs. A view that the nonphysisorbed hydrogen observed in carbon nanostructures would be connected with the formation of structural defects was expressed recently by Zuttel and Orimo [28]. 4.3.2

Active Carbons and Chemically Activated Carbons

More conventional process to prepare active carbons is by chemical activation. So-called activated carbons consist of multitude of stacks of graphene planes of various sizes and degree of disorder. Such carbons are highly macro- and mesoporous materials. According to IUPAC definitions, three groups of pores are distinguished: • Macropores (above 50-nm diameter) • Mesopores (2–50-nm diameter) • Micropores (under 2-nm diameter). The microporosity is often reported in recent research papers as nanoporosity. Commercial activated carbon grades have an internal surface area of 500 up to 1,500 m2/g. Powdered activated carbon comes with particle size 1–150 µm. There are also granulated or extruded materials with granule size in the 0.5–4-mm range. Activation is often conducted by processing with steam or chemical agents. Carbons activated by steam can be prepared from raw materials such as coal, peat, or lignite, which are carbonized and reacted with high-temperature water steam, in the process where fraction of carbon atoms are gasified, leaving beside porous structure. Chemically, carbon can also be activated with phosphoric acid. So-called mesocarbon microbeads (MCMBs) were produced from coal tar pitch in the Osaka

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Carbons and Nanocarbons

Gas Cherry T process; they were used in Li-intercalated rechargeable batteries. Recently, this active carbon was further developed with manufacturing so-called superactivated carbon through reaction of coke or MCMBs with potassium hydroxide. This new form of carbon materials have a open cage-like type of porosity and SSA as high as 3,220 m2/g. Reversible hydrogen adsorption, following closely the characteristic Langmuir isotherm, was reported providing the storage of 5 wt.% H at 77 K and 1.3 wt.% H2 at 296 K [29] (Table 4.1). The heat of adsorption ΔH can be determined from Clausius–Clapeyron equation: ln P =

−ΔH +C RT

(4.3)

where R is gas constant, T is absolute temperature, and C is a constant. The reported values about 6 kJ/mol H2 are still short of the desirable range above 10 kJ/mol that targets bond strength required for reversible hydrogen storage and release at ambient temperatures. This desirable range would be at 10–60 kJ/mol, the range when the energy of physical adsorption is sufficient enough to form the optimal bond strength, and storage does not require the use of cryogenic temperatures. Higher energy than 50–60 kJ/mol results in chemical absorption through the formation of strong s bonds as discussed in Sect. 4.2.1.4, and demands high desorption temperature. This is shown on the “energy line” (Fig. 4.10). The hydrogen storage capacities for disordered graphites, nanographites, and activated carbons are collected in Table 4.1. One can conclude that activated carbons are better storage materials than CNTs and most experimentally investigated carbon nanophases (like GNFs). Yet, if one applies a broader definition of nanomaterials, the activated carbon phases are, indeed, the disordered and nanostructured carbons. 4.3.3

Amorphous Carbon

When the size of sp2-bonded clusters of carbons decreases below La = 1–2 nm, nanocrystalline graphitic carbon becomes amorphous carbon (am-C). In fact, am-C begin to exhibit any mixture of sp2, sp3, and even sp1 sites. This area of research is connected with investigation of thin carbon films for electronics, particularly ultrahigh density data storage in optical discs and MEMS. Amorphous carbon with high fraction of sp3 bonding like in diamond and diamondoids is known as diamond-like carbon (DLC). The amorphous carbon requires hydrogen to stabilize its amorphous structure. Amorphous carbon thin films can contain as much as 40–60 at.% H. About 70% of carbon atoms can be sp3 bonded and most of these sp3 orbitals, if not connected to carbon, are terminated with hydrogen. Like in diamondoids we discussed in Sect. 4.1, Fig. 4.5, hydrogen stabilizes the carbon structure. Lower hydrogen content of about 20 wt.% H2 results in increased fraction of sp2 hybridized bonds in a phase known as graphite-like hydrogenated amorphous carbon (GLCH). It is relatively easily produced by physical- or chemical-vapor deposition or magnetron sputtering [38]. Raman spectroscopy demonstrated that along with hydrogen covalently bonded to carbon, and hence hardly reversible, there exists a quasi-free

4.4

Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

physisorption

H-C 0

305

chemisorption

H-Mg

H-H 1

2

3

4

H-OH 5 x102 kJ/mol

MgH2 activated carbons carbon nanotubes graphitic nanofibers

Fig. 4.10 Current and targeted range (horizontal bar) of bond strengths for nanocarbons (red), shown against the ranges for physisorption and chemisorption, and compared with the bond strength for high-temperature MgH2, and energetic molecules of H2 and H2O

hydrogen (in a form of highly stretched or polarized H2 molecules) adsorbed by graphite-like structural fragments in otherwise amorphous structure. The amount of hydrogen released by breaking of weak C–H bonds in low-temperature annealing process (at less than 500°C) was determined to be 10 at.%. Another 12 at.% was released by breaking covalent bonds at higher temperature, 750°C. Thermallyprogrammed desorption indicated that the quasi-free hydrogen began to effuse from material by annealing at 100°C and its release peaked at 400°C [39].

4.4

4.4.1

Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns Fullerenes and Hydrofullerenes

In C60 fullerene-type carbon allotrope, there is only one structure in which all the pentagons are nonadjacent and this is icosohedral symmetry-Ih (Fig. 4.11). This structure is often referred to as backyball to reflect on its full name buckminsterfullerene (after Buckminster Fuller who popularized the geodesic dome as an architectural form). Fullerenes can be made by either combustion or pyrolysis of aromatic hydrocarbons but are mostly prepared in arc-discharge processes. They have been studied extensively, and in fact, their studies developed into a new branch of organic chemistry. In textbooks of fullerene chemistry, hydrogenation is the simplest reaction and fullerene hydrides are the simplest derivatives of fullerenes. Hydrogenation can be conducted under H2 pressure and elevated temperatures. Heating C60 at 400°C and 80 atm H2 yields red solid, C60H18 [40]. Up to 48 H atoms can be added to C60 under more forcing conditions. Catalytic hydrogenation at 280°C and 160 atm with use of Ru/C catalyst produce hydrofulerenes up to C60H50 [41]. Hydrogenation is quite

77 600–1,066 (desorp-tion) 77 77 77 298 77 300 77–300 Ambient 300 77 273 273 77 77 273 77

725

988 (+173) 70 2,800 2,800 922 (+206) 3,220 3,220

119 (85) 49 (26) 120 (+120) – – – –

77

T (K)

7 (+7)

Surface area SBET(+St) (m2/g)

1 bar 1 bar 1 bar 12 12 11 1.5 0.9 11 1 bar 3 3 3 1 bar

Ambient

1 bar

p (MPa)

0.125 0.05 0.10 1.4 12.4 5.7 1–1.8 0.5 < 0.1 1.27 0.10 0.37 4.7 1.5 1.3 5

1.5 1.5

0

H2 (wt.%)

Nijkamp et al. Nijkamp et al. Nijkamp et al. Hwang et al. Fan et al. Cheng et al. Strobel et al. Hirscher et al. Hirscher and Becher Nijkamp et al. Chahine and Benard Chahine and Benard Chahine and Benard Nijkamp et al. Kojima, et al. Kojima, et al.

Hentsche et al. Hirscher et al.

Nijkamp et al.

Reported by

[17] [17] [17] [31] [32] [33] [34] [35] [36] [37] [14] [14] [14] [37] [29] [29]

[25] [30]

[17]

References

4

CNFs carbon nanofibers; GNFs graphitic nanofibers; AC activated carbon

Graphite – synthetic Graphite – ball milled Graphite – mechanically charged with H2 in ball mill Graphite – activated 100 CNF – large CNF – median diameter CNF CNF CNF GNF – herringbone GNF – mechanically Activated GNF GNF – ACF 500 CNF – Ni70Cu30-ethylene AC – AX 21 AC – AX 21 AC – Norit SX1 AC – M30 superactivated AC – M30 superactivated

Material

Table 4.1 Several reported hydrogen storage capacities in graphitic and activated carbons and nanocarbons

306 Carbons and Nanocarbons

4.4

Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

307

Fig. 4.11 C60 fullerene. Image credit: http://www.xdray.uu.se/cluster/images/c60

effective with use of hydrogen in statu nascendi generated from reaction of Zn with HCl conducted in toluene solution; fullerene C70H38 in mixture with some C60hydrofulerenes can be prepared by this method. Also, C60H36 is easily prepared in reduction with hydrogen from Zn/conc. HCl in toluene. The reduction is very rapid if stirring can be complete in 1 h. However, synthesis of C60H36 is often accompanied by the formation of C60H18. Another method is transfer hydrogenation conducted with the use of organic compounds that have a tendency to loose hydrogen with the formation of aromatic compounds. For example, heating C60 fullerene in 9,10-dihydroanthracene in a sealed tube in 350°C for 30 min yields C60H36 in mixture with C60H18 after 24 h of heating [41]. The reaction is reversible and hydrofullerene can be back converted to C60 fullerene with DDQ reactant dissolved in toluene. Hydrogenation of fullerenes, being a reduction process, proceeds rapidly, and this means that fullerenes are quite strong oxidizers. They are able even to reduce H2S to elemental S; the product is C60H36 with some admixture of C60H18. Oxidation back to C60 can be conducted but requires wet chemical reactions. Release of hydrogen gas in solid state reaction is difficult as the hydrogen forms strong chemical bonds with fullerenes. The hydrogen bonds to the outer surface of such buckyballs. Formation of endohedral complexes of C60 with hydrogen was also shown to be possible, both, in first-principle calculations and in experiments; this fills the buckyballs with only few percent of hydrogen to about C60H composition. Apparently, C60 and other higher fullerenes show no hydrogen storage capacity in respect to reversible hydrogen solid-state storage. Molecular crystals can also be synthesized using C60 buckyballs as building blocks. The interstitial sites in molecular crystal [C60]n can be filled with K atoms. The atoms of potassium occupy tetrahedral or octahedral sites in the cubic close packed

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lattice made of n fullerene nodes. Molecular H2 can replace K in octahedral sites resulting in the maximum H/C ratio of 1:20. However, the ratio is low and precludes any application for solid-state hydrogen storage. Fullerenes in combination with metals form metallo-organic compounds, fullerids. C60 fullerene in metallo-organic combination with scandium can bind as many as 11 hydrogen atoms giving theoretical capacity ca. 9 wt.% H [42], and in combination with titanium ca. 8 wt.% H2 [43].

4.4.2

Carbon Nanotubes

A single-walled carbon nanotube (SWNT) is a single graphene sheet rolled up in a seamless cylinder, whose diameter is of the order of few nanometers (Fig. 4.12). A double-walled carbon nanotube (DWNT) consists of rolled two graphene layers, and a multiwall carbon nanotube (MWNT) exhibits several co-axial rolls of graphene sheets, one sitting in each other and separated by about 0.35 nm. CNTs are prepared by either laser ablation from graphite target, arc discharge, or chemical vapor deposition. In either case, to grow they need the presence of Co or Ni catalyst. Typically, outer diameter of CNTs prepared by a carbon arc process ranges between 20 and 200 Å, and inner diameter ranges between 10 and 30 Å. An aspect ratio (length-to-diameter ratio) is typically of 102–103. When SWNTs are prepared, they come out as mostly capped at each end with fullerene-like half-sphere- or polyhedral cups. A CNT presents a highly-ordered, hexagonal-packed structure, not unlike the C60 buckyball and other high-ordered fullerenes. Highly ordered structure is well confirmed by direct imaging in HRTEM and in Raman spectra, the two most commonly used tools for microstructural characterization of CNTs. Analysis of Raman spectra indicates that DWNT samples come as a combination of two or more kinds of double-wall tubes; for example, one reported sample exhibits inner/outer diameters 0.72/1.48 nm and 0.91/1.61 nm, respectively [44]. SWNTs have a tendency to self-organize and form more thick nanoropes. Nanoropes consist of tenths and hundreds of SWNTs, which are intertwined and aligned. Cutting

roll

2 0 1 3

Fig. 4.12 The structural relationship between graphene sheet and single-walled carbon nanotube; arrows point to two alternative directions of rolling; circles inside rings reflect on aromatic character of carbon rings and delocalization of resonant electrons inside them

4.4

Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

309

a rope in the plane normal to the rope alignment direction, one can see a triangular lattice of circles, as illustrated in Fig. 4.13, with the intertube spacing g of 0.3–0.4 nm, as measured from the center of the tube walls. If so, such a structure should manifest itself in X-ray diffractograms. Indeed, while SWNT material exhibits no strong XRD peaks, three broadened maxima at 2q = 1.95°, 4.1°, and 11.8° were reported for the DWNT nanorope material [44]. The presence of three peaks, although they are broadened, indicates that the well-ordered structure in bundles of DWNTs forms a 2D-triangular lattice. This intertube spacing is determined by van der Walls interactions. This is a van der Walls gap: g=d – 2r, where d is the spacing between tubes in triangular lattice of tube bundles, and r is the radius of the nanotubes (see Fig. 4.13). It is well-established that the hydrogen does not enter the interior of CNTs in any larger amounts. Instead, it prefers adsorption on the exterior wall of the tube, i.e., in the van der Walls gap. Elastic and quasielastic neutron scattering results imply that at 80 K, under H2 pressure of 110 atm, H2 molecules gradually condense in the SWNT sample while being physisorbed in the interstitial tunnels of the SWNT bundles [45]. Apparently, hydrogen is stored mainly in bundles of SWNTs, with H2 molecules physisorbed at the exterior surface of the wall, in either groove (a) or interstitial (b) channels. It is only at high pressures that hydrogen can penetrate the interior of tubes to occupy endohedral sites (c); in such a case, tube diameter must be ca. 0.7 nm [46]. Since CNTs are capped with semispherical fullerene caps at their end, the cups further obstruct hydrogen molecules to enter interior channels of CNTs. Chemical methods are devised to remove these end-caps to open nanotubes for access of gas molecules. Anticipated is linear increase in hydrogen storage capacity with larger tube diameters [7]. The low values of adsorption seems to be a direct consequence of a great fraction of the surface area that is excluded from adsorption, this being either the interior of a r

c

d g

b

Fig. 4.13 Triangular network of d-spaced nanotubes in a bundle; marked are van der Walls gap (g), and positions of H atoms on the exterior wall (a, b) and in the interior of tube (c)

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CNT or the regions of nanoropes where the van der Walls gap is too wide, 0.4 nm or more. The reported gap of about 0.32 nm [47] seems also be not optimal for efficient bonding of H2 molecule; the distance between tubes is large enough for H2 molecule to pass through without bonding to carbon. Compressing the bundle can reduce this distance to 0.26 nm, at which case chemical bonding between external wall and hydrogen can become possible [48]. Optimization of CNT arrays and ropes for storage of hydrogen was computer simulated using grand-canonical system and Monte Carlo algorithms, and such simulations suggest that self-organized triangular pseudo-lattice of tubes is unlikely to result in intermolecular potentials such that hydrogen is bonded in near-ambient temperatures, as is requested by its use in vehicular fuel cell systems [49]. A stronger intermolecular potential can be realized in van der Walls gap in the bundle of DWNTs. The DWNT bundle with wide g value can absorb two times more H2 than the close-packed SWNT bundle [44]. CNT storage medium does not consist of just straight SWCT, DWNT, or MWNTs. The tubes can be bended and deformed. There is a question how such tube deformation affects hydrogen storage. A rolled up graphene sheet that forms a CNT consists of hexagons of carbon atoms, which are sp2, hybridized like in graphite. We have already discussed that sp2 hybridization directs C–C bonds to lie in the plain of hexagonal rings formed by carbon atoms. However, graphene sheet is highly curved in a CNT. One can ponder how this high curvature of single graphene sheet affects the reactivity of carbon atoms toward hydrogen. Further to this, one could calculate an excess energy that is required to pull one C atom out of its equilibrium position in the graphene plane. Such an out-of-plane bond stretching results in the strain energy, EC–C (strain), and is composed of excess energy of four atoms: the pulled-up atom 0, and its three nearest neighbors, in the in-plane positions 1, 2, and 3 (Fig. 4.12). Then, of interest would be the binding energy EC–H (strain) of H on locally strained atom C. Apparently, such energy would require energy of breaking of in-plane bonds and moving electron into empty out-of-plane p orbital. Calculations for such a scenario, conducted by DFT computer simulations, were recently reported and the model predicted enhanced hydrogenation energy at the kink site of bent CNTs [50]. In another computer model it is found that charge transfer occurs from a low curvature region to a high curvature region of the deformed CNT bundle at 80 K and 10 MPa. It was suggested that carbon atoms develop charge polarization only on the deformed structure, like an oval-shaped CNT. The long-range electrostatic interactions of polarized charges on the deformed bundle increase the ordering of H2 molecules that interact with this surface. This causes condensation of hydrogen gas on deformed CNTs, although only in bundles and not on single SWNTs [51]. The condensed hydrogen may extract electron density from the carbon on the highly curved surface of a nanotube (hence, acts as a “hydrogen acid”) and forms covalent bond, as this is shown in Fig. 4.6, position a. Therefore, enhanced hydrogenation may benefit chemisorption instead of physisorption. Through chemisorption the hydrogen atoms will exothermally bond to the C atom on the tube wall. Again, this can be enhanced by the out-of-plain pullout, hence sp3 hybridization and out-of-wall charge polarization induced by the effect of high curvature. Bonds

4.4

Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

311

formed via chemisorption are much more stronger, and H2 is released only upon heating to temperatures above 400°C [52]. Unfortunately, such high temperatures are still not acceptable for vehicular hydrogen storage. The intercalation of graphite with alkali metals, particularly with Li+ ions, was investigated in reference to rechargeable Li-ion batteries; however, intercalation of carbonic materials with molecular hydrogen has not been successful [4]. Molecular H2 does not intercalate into graphites and carbonic materials but physisorbs on the carbon surfaces. A graphite intercalation compound, KC8, consists of alternating graphene planes and potassium, and lithiated graphites used in Li-ion batteries constitute the lithium analogue of this compound, LiC8. Electrochemical insertion of Li into SWNTs gives one Li+ per two atoms of C; however, the dynamic diameter of H2 molecule is much larger than Li+ ion. Intercalation of graphitic material with alkali metal causes the graphene layers to swell and thus facilitate the hydrogen insertion into graphitic nanocarbons. The early results of high H2 uptake of H2 by alkali-doped CNTs under ambient pressures and moderate temperatures [53] were not confirmed. Intercalated lithium have complex chemistry and high chemical affinity to water and carbon dioxide present in air. When re-examined, the weight gain attributed to hydrogen uptake appeared to be caused rather by absorption of water than by hydrogen. Lithiated SWNTs seems to store more than 4 wt.% H2, albeit the hydrogen desorbs only at higher temperatures above 600 K [54]. Summing up, considerable time and effort has been invested worldwide into utilizing CNTs as a hydrogen storage media; however, their use as viable hydrogen storage remains questionable. The hydrogen storage capacity of SWNTs has been studied by means of gravimetric and volumetric measurements. The result reported differs greatly (Table 4.2), to the extent not often met in reports from research. Hydrogen gas charged into CNTs at cryogenic temperatures does not bind firmly Table 4.2 2 Hydrogen storage in carbon nanotubes Material

T (K)

p (MPa)

H2 (wt.%)

Reported by

References

SWNT – SSA = 610 m2/g SWNT – 75% purity

77

0.05

~0.54

Miyamato et al.

[44]

<123 77 77

25 bar 0.1 0.05

2.4 ~0.7 ~0.76

Tarasov et al Miyamato et al. Miyamato et al.

[66] [44] [44]

77 300 473

0.1 7 0.1

~1.16 0.7–0.8 0.7

Miyamato et al. Badzian et al. Pinkerton et al.

[44] [67] [54]

663

0.1

4.2

Pinkerton et al.

[54]

Ambient

0.08

< 0.1

Hirscher et al.

[35]

823

Ambient

4.5

Chen and Huang

[68]

DWNT – SSA = 330 m2/g MWNT – high purity CNT-Li doped (purity <40%) CNT-Li doped (purity <40%) SWNT – ball milled in Ar MWNT-KOH modified

SWNTs single-walled nanotubes; DWNTs double-walled nanotubes; MWNTs multi-walled nanotubes; CNTs carbon nanotubes not purified

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to the carbon and most of it evaporates as soon as it is brought back to room temperatures. Currently, this is common feeling that CNTs themselves will not provide the solution to high-capacity, room-temperature hydrogen storage as needed for automotive applications. Under ambient conditions the amount of hydrogen stored in SWNTs is less than 1 wt.% [4]. The excess volume of hydrogen stored, which is the difference between the volume stored in nanocarbons and the volume taken by the nanomaterials itself, is not great and the storage in nanocarbons needs either cryogenic temperatures or high pressures in order to keep hydrogen bonded. Indeed, while adsorption in nanocarbons may reduce the pressure necessary for gas storage in high-pressure tank, the volume of such a material also reduces the total tank volume available for hydrogen. The argument was put forward that while about 1 wt.% H2 can be stored in nanocarbons per each 100 atm, an optimistic linear extrapolation gives for the US DOE target of 6.5 wt.% at as much as 700 atm pressure needed to keep this hydrogen at operational pressure plateau in the tank; however, at ambient temperatures this gravimetric density can be realized simply by pressurizing gas hydrogen with no help from nanocarbons [4]. A highly optimistic beginning of the hydrogen storage in CNTs is turning out quite disappointing these days. This provides good example of “overhype” brought about by nanotechnology, and in general by science when a new paradigm is born. The processes to prepare CNTs require high-vacuum systems to generate plasma either in laser-ablation, glow-plasma or arc-plasma discharges. The cost of scale-up and manufacturing are high, yield of CNTs synthesis low, and product purity questionable. Because hydrogen bonds rather to highly-curved exterior walls, and resists entering interior of the CNTs, the question arises: why do we need tubes at all, and why not storage in inexpensive disordered carbons rather than in costly CNTs?

4.4.3

Carbon Nanohorns

The disappointing results for H storage in highly ordered CNTs may be saved with the emergence of carbon nanohorns. Single-walled carbon nanohorns (CNHs) consist of single-walled graphitic structures formed out of a single graphene sheet rolled up to form conical (hornlike) shapes, which are rounded at the tip (Fig. 4.14a). Typically, their average size is of 2–3 nm. These horn-like shapes aggregate to form globular rosette structures with sizes of about 80–100 nm (Fig. 4.14b). CNHs exhibit very large surface areas approaching 1,500 m2/g. These phases are prepared by arc-discharge synthesis or laser ablation [55]. In contrast to CNTs, their growth does not require use of metal catalysts. Low-cost and high-purity CNH material can be produced in scaled-up process. CNHs samples can be obtained from MER Corporation (Tuscon, AZ) and from small batches processed by several university groups worldwide. Because of low cost, high purity and high surface area CNHs become attractive candidates for gas and/or liquid hydrogen storage. Indeed, recently it was reported isosteric heats of H2 adsorption corresponding to 100–120 meV energy for binding

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Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

313

a

CNH rosette nanoparticles

b Fig. 4.14 Carbon nanohorns with hydrogen condensed at the end of conical tip (a) agglomerate to form nanocarbon particles that exhibit rosette shape (b); TEM image credit: http://www.physics.siu.edu/migone/lab/carbon_nanohorns.htm

of H2 molecule to carbon [56]. It is much more than for SWNTs. This increase in the energy of hydrogen bonding has been attributed to enhanced interaction of H2 molecules at the conical tip of nanohorn, where gas-to-liquid transition, and even solid-like H2 was suggested to exist as a consequence of a quantum effect [56]. Inelastic and quasielastic neutron scattering experiments showed unambiguous signature of strong interaction between H2 molecule and CNHs, and the character of this interaction is quantitatively different to that in CNTs [57]. In this context, recent 13C NMR studies have identified two distinct processes with quite disparate relaxation times [58], one of them being attributed to slow motions of 13C, which takes place inside conical tips where hydrogen undergoes likely condensation process. We can expect that similar effect can be at play when hydrogen molecules concentrate

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at the tip of crack propagating through graphite crystal and in boundaries between graphitic nanocrystallites in disordered graphites, as discussed in Sect. 4.3. However, the density of crack tips in highly disordered graphites should be lower than the density of conical tips in carbon nanohorn phase.

4.4.4

Nanostructured Carbon Shells and Carbon Onions

Carbon nanoshells and carbon onions (Fig. 4.15) look like fullerene-related material, which they are not. These are sphere-shaped fine particles of carbon, which exhibit concentric graphitic shells. The carbon onions consist of nesting graphitic shells. These carbon shapes resemble Russian Matryoshka dolls with one shell of carbon atoms inside the other one, and so on. It was reported that intense irradiation of amorphous carbon in electron microscope results first in graphitization then in curling of graphene planes to the point that they close, forming perfectly spherical concentric set of graphene shells [59, 61]. Nucleation and growth of carbon onions was observed in polycrystalline metals implanted with carbon atoms [62], in some stages of carbonyl nickel chemical vapor deposition (Ni CVD) process, and in nanodiamond clusters heated above 800°C with or without the presence of ferrous catalyst [63]. They can also be produced in relatively inexpensive process (no vacuum required) by arc discharge conducted in water. It is interesting to note that good yield for production of such nanocarbons requires extremely sharp gradient

Fig. 4.15 Carbon onion. After [59]. Image credit: http://www.staff.uni-mainz.de/banhart/ c-nano-structures/onions. Note compression of interlayer distances – from 0.335 nm for ordered graphite to less than 0.3 nm; this is in contrast to the swelling of planes in turbostatic graphite and in carbon nanoshells [60]

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Highly Ordered Fullerenes, Carbon Nanotubes, and Carbon Nanohorns

315

between the gas bubble from the hot-plasma region and the water coolant [64] – this being reminiscent of the rapid quenching rate criterion applied to the formation of nanostructured intermetallics for hydrogen storage (see Sect. 1.2.2). Carbon nanoshells were extracted from some commercial nickel powders produced by INCO of Canada [60]. Commercial Ni metal powders are manufactured by Ni CVD carbonyl process from nickel tetracarbonyl. These commercially available Ni powders were found to be coated with a nanometric layer of defected graphitic carbon. Through a simple process of acid leaching of Ni metal, a dry filtrate is obtained. When imaged in a transition electron microscope, the dry powder filtrate consists of nanostructured hollow carbon shells that are positive replicas of the nickel particulates. These carbon nanoshells exhibited thickness that range from a few nanometers to tens of nanometers. They are composed of layered nanocrystals of turbostratic graphite, the average thickness of which is 9–10 nm, as determined by X-ray diffraction. In all cases, the graphene basal planes follow the curvature of the original Ni metal core [60]. A high-resolution transmission electron microscopy (HREM) picture of the planes of carbon atoms is shown in Fig. 4.16, and the model of the shell, as consisting of short stacks of turbostratic graphite, is depicted in Fig. 4.17. Carbon nanoshells of this type can be produced cheaply and in large quantities by recycling the carbonyl Ni powders. One can observe that carbon nanoshells, as consisting of small domains of graphitic sp2 sheets, must exhibit multitude of dangling bonds at their peripheries. These domains of stacked graphene sheets can also be seen as layered graphitic nanocrystals. The dimensions of these nanocrystals (few tenths nanometers thick by few hundred nanometers length) provide efficient constrains for mobility of p

Fig. 4.16 HREM image of carbon nanoshells showing that the basal planes roughly follow the curvature of the shell owing to the presence of a high density of dislocations. The image contrast of the basal planes was enhanced by use of fast Fourier transform (FFT) processing. After [60]

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Fig. 4.17 Model of a carbon nanoshell pointing to the stacks of 5–10 aligned layers within the shell. Hydrogen can possibly bond to the dangling bonds on the periphery of layered graphitic nanocrystals. After [60]

Fig. 4.18 The different degree to which electrons move collectively in various forms of carbon material as evidenced by distinct intensity of the plasmon peak located about 6 eV in EELS spectra (arrow). Hydrogen atoms can make less strong covalent bonds with participation of p electrons if the interplanar distance is increased in layered graphitic nanocrystals as seen in carbon nanoshells (frame in Fig. 4.17) and in disordered graphitic carbons (Sect. 4.3.1). After [60]

References

317

electrons, which otherwise are delocalized over all graphitic material. The localized p electrons and dangling bonds at the periphery of these nanocrystals make them a new interesting form of carbon, distinct from amorphous carbon, and graphite, as deduced from EELS spectra in Fig. 4.18. HREM and Raman studies by Tomita et al. [65] seem to suggest that the same fragmentation of graphite into graphitic nanocrystallites takes place also in carbon onions. Hydrogen storage properties of sp2-bonded carbon nanoshells and carbon onions have still not been reported in PCT experiments, as they were not reported for carbon nanohorns, and sp3-bonded nanodiamond phases. However, the reason we discuss these new nanocarbons here is that they bring the promise: that with the amazing propensity of C atoms to form bonds and structures, and more and more evidence coming that hydrogen storage properties can be controlled not only by physisorption but by various degree of chemisorption, new nanocarbons can be designed with just the right strength of C–H bonds. This should help current worldwide search, and the Holy Grail of reversible hydrogen storage for future hydrogen economy will be found one day, in one breakthrough research.

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59. F. Banhart, T. Fuller, P.H. Redlich, P.M. Ajayan, “The formation and self-compression of carbon onions.” Chem. Phys. Lett., 269 (1997) 349–355. 60. Z.S. Wronski, G.J.C. Carpenter, “Carbon nanoshells obtained from leaching carbonyl nickel metal powders.” Carbon, 44 (2006) 1799–1789. 61. D. Ugarte, “Curling and closure of graphitic networks under electron-beam irradiation.” Nature, 359 (1992) 700–709. 62. H. Abe, “Nucleation of carbon onions and nanocapsules under ion implantation at high temperature.” Diamond Relat. Mater., 10 (2001) 1201. 63. A. Hirata, M. Igarashi, T. Kaito, “Study on solid lubricant properties of carbon onions produced by heat treatment of diamond clusters and particles.” Tribology Int., 37(11–12) (2004) 899–905. 64. N. Sano, H. Wang, I. Alexandrou, M. Chhowalla, K.B.K. Teo, G.A. Amaratunga, “Properties of carbon onions produced by arc discharge in water.” J. Appl. Phys., 92(5) (2002) 2783–2788. 65. S. Tomita, M. Fujii, S. Hayashi, K. Yamamoto, “Electron energy-loss spectroscopy of carbon onions.” Chem. Phys. Lett., 305(3–4) (1999) 225–229. 66. B.P. Tarasow, J.P. Maehlen, M.V. Lototsky, V.E. Maradyan, V.A. Yartys, “Hydrogen sorption properties of arc generated single-wall carbon nanotubes.” J. Alloys Compd., 356–357 (2003) 510–514. 67. A. Badzian, T. Badzian, E. Brevel, A. Piotrowski, “Nanostructured nitrogen -doped carbon materials for hydrogen storage.” Thin Solid Films, 398–399 (2001) 170–174. 68. C-H. Chen, C-C. Huang, “Hydrogen storage by KOH-modified multi-walled carbon nanotubes.” Int. J. Hydrogen Energy, 32 (2007) 237–246.

Chapter 5

Summary

In this chapter, we will try to summarize the critical issues to be resolved and suggest some future research directions in the area of solid-state hydrogen storage. We will try our skills in predicting future trends, and come out with our personal assessment about the prospects of nanostructured hydrides and other nanomaterials to be used in solid-state hydrogen storage. The following discussion will be guided by observation that, as for the day we finish this book, no solution is found for onboard storage of hydrogen for future mass transportation, which we strongly believe will rely on a small, lightweight zero-emission car, being it either fuel-cell (FC) car or an FC-EV hybrid car. Subsequent is the assumption that the alternative and overlooked method of supplying hydrogen to an onboard FC should utilize an offboard recharging of hydride reservoirs or containers rather than an onboard recharging by direct pumping gaseous or liquid hydrogen at a refueling station. We introduced this idea in Sect. 1.1 (Fig. 1.1). We reiterate that in our opinion engineering problems with onboard direct hydrogen refueling are of such a magnitude that the efforts to overcome them would, at the end, make an offboard refueling viable alternative. Even at a pressure of 35 MPa (let alone at 70 MPa), how could gaseous hydrogen be safely transported hundreds of kilometers to the tens of thousands of stations in a country, then stored at a particular station for a reasonably long time, and eventually pumped into a tank of an FC powered car at that pressure? Undoubtedly, many attempts to build up conventional infrastructures, based either on pressurized or liquid hydrogen, will be undertaken. In fact many demonstrations are under way in many countries. To alleviate the problems of centralized distribution of high pressure hydrogen gas, Honda proposes individual home-located either solar-powered water electrolyzing or natural gas hydrogen generation stations (more details can be found on Hydrogen Station-the Honda FCX, http://world. honda.com/FuelCell/FCX/station/). Hydrogen generated at such a small station would be compressed to a required pressure by an on-site compressor and stored in a high-pressure underground tank. However, the economy and the safety of implementing this kind of a solution on a mass scale is a different story. Similarly, transporting, storing at the station, and pumping liquid hydrogen at 21K is an even more formidable task. Everyone knows it well what kind of safety precautions are undertaken when a space shuttle lifting main rocket is fuelled with liquid hydrogen and oxygen. R.A. Varin et al., Nanomaterials for Solid State Hydrogen Storage, DOI: 10.1007/978-0-387-77712-2_5, © Springer Science + Business Media, LLC 2009

321

322

5

Summary

Turning to solid-state storage, as already mentioned in Sect.1.1, onboard recharging a discharged hydride (assuming that such a hydride becomes available at some point in time) would generate a large amount of waste heat that could only be usefully utilized by applying a cumbersome onboard heat exchanger. In view of all these formidable engineering problems with onboard hydrogen refueling or recharging, an offboard method utilizing a neat recyclable reservoir or container (Fig. 1.1) seems to be a lot easier engineering solution. This offboard recharging concept is our guideline in the following discussion of suitability of various hydrides for solid-state hydrogen storage. We will first discuss metal/intermetallic hydrides, then complex hydrides, and finally other nanostructures such as nanocarbons and others.

5.1

Metal/Intermetallic Hydrides

The most attractive simple metal hydride is still magnesium hydride, MgH2. At the present stage, a catalyzed MgH2 (e.g., with a metal nano-Ni catalyst; Sect. 2.2.1) subjected to nanostructuring by ball milling has, at around 280–300°C and atmospheric pressure of H2, sufficient kinetics of desorption close to that targeted by the US DOE (Sect. 1.1; Table 1.2). It can easily provide 6–7 wt% H2 on desorption. After full decomposition to an elemental Mg, its recharging is easy by a simple hydrogenation under gaseous hydrogen. Even better, its recharging can be combined with a simultaneous nanostructuring by applying reactive ball milling in hydrogen alloying mills (Sect. 2.1). Obviously, the temperature range of around 300°C required for desorption (less for absorption) is far away from the 60–80°C range targeted by the US DOE. However, this problem can be relatively easily overcome by some smart engineering design of hydrogen reservoir or container, which utilizes either heating of MgH2 by flameless (catalytic) burning of hydrogen released from a low-temperature desorbing hydride held in an auxiliary container ([27] in Sect. 1.1) or even heating by a resistance heating element connected to a battery. A minor disadvantage of nanostructured MgH2 is that after decomposition to nanostructured Mg it becomes extremely pyrophoric and starts burning in contact with air. In our long-term research on MgH2 (Sect. 2.1), we have also noticed on a number of occasions that a nanostructured MgH2 powder obtained by ball milling becomes far more moisture sensitive than a coarse particulate, nonmilled MgH2. These two phenomena are most likely related to a substantial reduction of particle size upon ball milling. However, we have noticed that the addition of nanoNi (Sect. 2.2.1) almost completely suppresses pyrophoricity, and the nanostructured Mg with nano-Ni behaves in a benign manner in contact with air. In our opinion, at present, MgH2 is in essence ready for commercial application in prototype hydrogen storage reservoirs or containers. However, the major thrust should be shifted from quite useless and never-ending testing of various additives to MgH2 to pursuing more ingenious engineering designs of hydrogen storage reservoirs that could easily be accommodated into FC powered vehicles and then easily

5.2

Complex Hydrides

323

recharged offboard after being depleted. This could be an alternative hydrogen storage and delivery solution for mobile applications. The second simple metal hydride of some good potential for further development is AlH3. As reviewed in Sect. 2.4, the major advantages of this hydride include high gravimetric capacity, low temperature desorption range of 100–150°C, accompanied by reasonable kinetics for a catalyzed AlH3, and a very simple decomposition reaction to just purify Al and H2 in exactly the same manner as for MgH2. Disadvantages of AlH3 are a very costly conventional organometallic synthesis and the presence of a large number of polymorphs that may affect the conditions of desorption. In reality, such a large number of polymorphs relates to the less than perfect organometallic synthesis. It seems that the most important research focus would be on the development of inexpensive and fast bulk synthesis method of AlH3. Also, finding more effective catalysts would be an important research task. Regarding intermetallic-based hydrides, there is no hydride in this group that would be even remotely close to any commercial vehicular application. An interesting hydride is Mg2FeH6, whose great advantages are relative ease of synthesis in hydrogen alloying mills, better aging properties and its mechano-chemical reversibility by reactive mechanical alloying (Sect. 3.1.2). Although providing the storage capacity that is inferior to MgH2 and thus not considered for any serious storage for vehicular application, it is most likely the most ideal candidate for thermochemical thermal energy storage devices ([40] in Sect. 3.1.2).

5.2

Complex Hydrides

In the alanate group, NaAlH4 is one of the most extensively researched hydride. Its desorption parameters are now quite close to the US DOE targets (Sect. 3.2.1) and it is reversible under a reasonable temperature–pressure range, which is a great asset of a hydride, as considered for onboard storage. However, its hydrogen capacities are far too low for any mobile application. In our opinion, if a good engineering design could result in a reservoir based on nanostructured, catalyzed MgH2, this would be a far better solution than using NaAlH4 as a storage medium. Other hydride in this group that has some potential for a good storage medium is LiAlH4. Nanostructured and catalyzed LiAlH4 can desorb quite a large quantity of H2 with reasonable kinetics within a temperature range 150–200°C (Sect. 3.2.2). It is quite possible that more research could improve its properties even further. However, the fact that this hydride is based on Li might be a substantial drawback. Li is a relatively rare element and relatively expensive by comparison with more common metallic elements. Alanates of Mg and Ca such as Mg(AlH4)2 and Ca(AlH4)2, respectively, are based on much more abundant and rather inexpensive elements. However, in comparison to NaAlH4 and LiAlH4, the knowledge about their efficient synthesis and hydrogen storage properties is miniscule. They still deserve more intensive research efforts.

324

5

Summary

In the group of amides (Sect. 3.3), we do not see any single-phase amide that would give any hope for future vehicular storage. Only mixtures of amides with other hydrides (e.g., MgH2) show slightly improved storage properties. However, a complexity of these reactions between two dissimilar hydrides, a relatively small amount of hydrogen desorbed and a continuous presence of larger or smaller quantities of NH3, which is detrimental to the PEM membranes, make this system less promising than the others. In addition, LiNH2 and Li2NH used in these reactions have the same drawback as LiAlH4, i.e., they contain relatively expensive Li, as discussed earlier. In the metal borohydrides group (Sect. 3.4), there is no apparent hydride that would be of the same interest for hydrogen storage as the few alanates described earlier. A success was achieved in developing a simple rule that shows that metal borohydrides with low Pauling electronegativities (<1.3) exhibit high decomposition temperatures and those with the Pauling electronegativities greater than 1.3 have quite low decomposition temperatures, essentially compatible with the US DOE targets (Sect. 3.4). This would be an excellent guideline for selecting promising metal borohydrides for vehicular storage application if not for the fact that the metal borohydrides with high Pauling electronegativities desorb some quantities of toxic borane gases together with hydrogen. Boranes, in essence, kill the membrane of a PEM FC even at minute quantities. A prominent example here is Zn(BH4)2, which decomposes within an ideal temperature range of ~80°C releasing boranes mixed with hydrogen. It is apparent to us that in the group of metal borohydrides the research should be mainly focused on finding methods to completely suppress the release of boranes from the metal borohydrides having the Pauling electronegativities greater than 1.3. Destabilization of high-temperature hydrides by nanostructuring and compositing of nanohydrides (Sect. 3.5) is in its infancy. At the first look, it seems to be very promising but needs much more vigorous research.

5.3

Nanocarbons and Others

Nanocarbons and other carbonaceous materials such as carbon nanotubes with high specific surface area join hydrogen (adsorb) by physisorption, which requires cryogenic temperatures (77K). Relatively weak Van der Waals interactions (Chap. 4) is beneficial for debonding of hydrogen from porous carbonaceous materials but requires low temperatures to keep the hydrogen bonded in the first place. This type of bonding is completely different than the chemical covalent–ionic bonding observed for hydrides. In essence, the hydrogen capacity of these structures at cryogenic temperatures is proportional to the applied absorption pressure and the specific surface area of a carbon nanostructure. The rather small gravimetric and volumetric hydrogen density on carbon nanostructures, together with the need for cryogenic temperatures and very high pressures necessary for a reasonable absorption, are significant drawbacks for the development of commercial storage reservoirs.

5.3

Nanocarbons and Others

325

The research field still not well explored is the hybrid materials where hydrogen is bonded through physisorption mixed with chemisorption. Finally, there is the all new emerging field of metallo-organic frameworks, which we did not attempt to cover in this book, for the sake of maintaining conciseness and competency. A materials breakthrough coming from these areas cannot be excluded, and in contrary are quite anticipated.

Index

A ABCR MgH2 powder cycled samples, DSC curves of, 124 desorption behavior of, 147 desorption kinetic curves, 167 DSC hydrogen desorption curves, 121, 126 DSC hydrogen desorption temperatures, 125 DSC traces of, 160, 166 mean particle size (ECD) of, 124 milling and cycling, phases, 160 n-Ni, desorption kinetic curves, 162 particle of, 101 SEM micrographs of, 158 under HES57 mode, 104 XRD patterns of, 159, 165 Absorption curve, 90 Absorption kinetics, 88, 89 of ABCR powder, 113 of ultrafine Mg particulate, 123 Activated carbons constituents of, 303 hydrogen storage capacities for, 306 Activation energy calculation of, 60–64 value, 95 Adamantine, 293 Adiabatic shear bands, 43 Aguey-Zinsou, K.-F., 167, 168 Alanates, sodium alanate hydride. See Sodium alanate hydride AlH3 (Alane) ball milling and doping, 175 decomposition/desorption reaction, 175 Department of Energy (DOE), 174 schematic representation of, 176 thermal and photolytic, 175

thermodynamic and kinetic data for, 176 thermogravimetric and differential thermal analysis (TG-DTA), 177 Alloy formation upon dehydrogenation, destabilization through, 254 Al(Mg) solid solution lattice parameters of, 227 reactions of Mg and Al to form, 226 Al3Mg2 solid solution high intensity XRD peaks, 227 reactions of Mg and Al to form, 226 Aluminum hydride (AlH3), 21 Amides. See Lithium amide (LiNH2) Amorphization, by diffusion, 18 Amorphous alloys, 11, 17 Amorphous carbon (am-C), 24, 291 hydrogenated, 25 hydrogen released by breaking of weak C–H bonds in, 305 stabilization with hydrogen, 304 Amorphous ferromagnetic materials, 11 Amorphous magnesium hydroxide, 99 Amorphous metals, 10 Andersson, Y., 173 Andreasen, A., 207, 215, 217–219 Ares, J. R., 107 Arrhenius equation, 60 Atomic force microscope (ATM), 12 Atomic H / M ratio, 58 Attritor mills, 30, 31 Avrami exponent, 62 Avrami, M., 9

B Backyball, 305 Bakulina, V. M., 244

327

328 Ball mill, 17 Fritsch model, 31 HES57 mode, 121 high-energy, 37 Low-energy shearing (LES), 105 multigrain particulates, 45 planetary, 31, 36 vibrational, 40 Ball-milled hydrides, microstructural characterization of, 71–73 Ball-milled MgH2 powders, kinetics of, 145 Ball milling high-energy, 36 nanostructured mixture synthesis by, 253 novel nanostructured bcc alloys Mg–Tm–V, 172 Ball-to-powder weight ratio, 36 milling, efficiency of, 138 Barium-calcium niobium compound, 153 Barkhordarian, G., 63, 167, 168 Battery-grade Ni hydroxides, 48 Benjamin, J., 17 Biomass, thermal processing of, 1 Bloch, J., 17 Borohydrides, synthesis of, 241 Bouaricha, S., 69, 152, 169, 171 Bowman, R. C., 12 Brinks, H. W., 220, 221 Brower, F. M., 174, 176 Bulk alloys, pulverization of, 181 Bystrzycki, J., 154

C Ca(AlH4)2 hydrogen gravimetric capacity of, 230 hydrogen storage properties, 231 MCAS synthesis of, 230 Ca(AlH4)2·THFx, 230–231 Ca2NH, hydrogen reversibility for, 231 Cantor, A. M., 11 Carbon atoms in graphite, 300–301 valence electrons of, 291–292 Carbon-based nanomaterials, 6 Carbon nanofibers (CNFs), hydrogen storage properties of, 299 Carbon nanoshells, 25 constituents of domains of stacked graphene sheets, 315–316 sphere-shaped fine particles of carbon, 314

Index HREM image of, 315 hydrogen storage properties of, 317 Carbon nanotubes (CNTs), 169, 291 development of, 23–25 with high specific surface area, 324 hydrogen storage in, 311–312 storage medium, 310 synthesis of, 308 van der Walls gap in, 309 Carbon onions constituents of, 314 hydrogen storage properties of, 317 Carbon polymorphs, 291 Carbyne, 291 Carey-Lea, M., 18 Carpenter, G. J. C., 25 Cast metal ingot, 18 Cauchy/Gaussian approximation, 72, 85 Cauchy/Gaussian method, 41 Caustic battery electrolyte, 15 C–C bonding, 291 Chandra, D., 180 Chemical hydrides, 195 Chemical vapor deposition (CVD), 23, 28 Chemisorption, in nanohydrides, 26 Chen, C., 197 Cheng, H. M., 181 Chen, J., 205, 218–220 Chen, P., 25, 231, 232, 236 Chen, Y., 130, 181 Chizinsky, G., 174 Cho, Y. W., 251 C60H36, synthesis of, 307 Claudy, P., 207, 208, 220, 223 Cold interdiffusion, of metal atoms, 39 Cold-welded balls, 34 Cold welding, 34 Complex hydrides, 323–324 classification of, 195 composition of, 195 Composite, and thermally stable hydride, 253 Compressive elastic stresses, 42 Controlled reactive mechanical alloying (CRMA) composites synthesis by, 255 of elemental Mg and Al powders, 229–230 nanocrystalline Mg2FeH6, 200, 202, 203 Controlled reactive mechanical alloying modes (CRMA), 37, 38 Controlled reactive mechanical milling (CRMM), 37, 38, 138, 155 Corrosion resistance, 15 CO2 sequestration, 1 Courtney, T. H., 32

Index Covalent–ionic bonding, 324 CRMM. See Controlled reactive mechanical milling (CRMM) Cryogenic liquids, 3 Cryogenic temperatures, 324 Cryomilling, 173 Crystallographic transformation, 19 Crystal structure, dislocation, 44 Czujko, T., 153, 197

D Dahl, J. E., 294 Dal Toè, S., 98 Davy, H., 8 Deformation faults, 301 Dehouche, Z., 152, 167 Deledda, S., 173 Desorption kinetics, 121 curves at various temperatures, 96–98, 115, 116, 161 of cycled mixture, 161 Diamond, 291 crystal structure of, 292–293 Diamond-like carbon (DLC), 25, 304 Diamondoids as cage-like saturated hydrocarbon molecules, 293 hydrogen content in, 294 sp3 hybridized orbitals, 293, 294 Didisheim, J. J., 198 Differential scanning calorimetry (DSC), 17, 60, 101 desorption peak temperature, 121 hydrogen desorption temperature, 124 Dillon, A., 24 Disordered graphites crystallinity, 301 deformation faults in based on deformed hexagonal crystal, 301 disclinations and plane rotations, 302 effect of proliferation of, 303 hydrogen absorption in, 302 hydrogen storage capacities for, 304, 306 Double-walled carbon nanotube (DWNT), 23 hydrogen absorption capacity, 310, 311 Raman spectra of, 308 XRD peaks, 309 DSC. See Differential scanning calorimetry (DSC) Duvez, P., 10 Dymova, T. N., 20, 22, 209, 223, 241

329 E Edwards, P. P., 178 Egorenko, G. A., 242, 244 Eigen, N., 212 Elansari, L., 179 Elastic shear stress, 42 Elasto-plastic accommodation theory, 143 Elasto-plastic irreversible energy losses, 144 Electrical conversion efficiency, 22 Energy dispersive spectroscopy (EDS), 107 Energy security, 1 Enthalpy of formation/decomposition, of Mg2NiH4, 196 Euler, L., 24

F Fátay, D., 123 FeTi, 16 FC-EV hybrid car, 321 Fernández, J. F., 89, 92 Fichtner, M., 23, 211, 223, 237 Finholt, A. E., 174 Fossdal, A., 223, 227 Fossil fuel, 27 Fossil fuel-based economy, 1 Frank-Read source, 44 Friedlmeier, G., 89, 91 Friedrichs, O., 99, 123, 146, 148, 168 Fuel-cell (FC) car, 321 Fuel cell, hydrogen storage tank for, 15 Fuel cell-powered vehicles, refueling/retail station for, 7 Fujii, H., 181 Fullerenes, 24, 291 C60H36, 307 hydrogenation of, 305–306 molecular crystals, 307–308 preparation of, 305 structure of, 307 Fu, Y., 152

G Gaseous voltaic battery, 8 Gasoline, 2 Gas turbines, 17 Gaussian distribution, 85 Gennari, F. C., 140 Glass-to-crystal transformation, 11, 17 Gleiter, H., 13 Goo, N. H., 174

330 Graetz, J., 176, 177 Grain boundaries, 13, 45 Grain boundary planes, 100 Grain size distribution, 9 powders Tego Magnan and ABCR, 105 refinement, 42 variations, of premilled ABCR powder, 109 Graphene, 291 chemisorption of hydrogen, 298–299 definition of, 294 in-plane σ and out-of-plane π bonding, 295–296 physisorption of hydrogen, 297–298 van der Walls interplanar and intermolecular interactions, 296–297 Graphite, 291 crystallographic unit cell for, 299–300 density of, 300 disordered (see Disordered graphites) intercalation compound, 311 stacking of hexagonal basal planes in, 302 turbostratic (see Turbostratic graphite) Graphite-like hydrogenated amorphous carbon (GLCH), 304 Graphitic carbon nanofibers (GNFs), 299 Graphitic polyhedral crystals, 299 Greenhouse gas emissions, 1 mitigation of, 17 Grochala, W., 178 Groll, M., 89, 91 Gross, K. J., 207, 209 Grove, W., 8 GTP (temperature–pressure) system, 31 Guénée, L., 171, 183 Gupta, M., 199

H Hall-Petch relation, 45 Hanada, N., 122, 155, 169, 172 Harris, I. R., 11, 18 HCVD. See Hydriding chemical vapor deposition High-energy milling, 34 High-energy shearing (HES), 105 mode, 132 Highly disordered graphites, density of crack tips in, 314

Index Highly ordered nanocarbons, role of defects in changing physisorption to chemisorption, 303 Highly oriented pyrolytic graphite (HOPG), 300 High-resolution transmission electron microscope (HRTEM), 24, 146 High temperature hydride, 254–255 Hirscher, M., 174 H/M (hydrogen-to-metal) atom ratio, 12 Hooke’s Law, 41 Hotta, H., 182 Hubbert’s Peak, 1 Hudson, M. S. L., 230 Huot, J., 95, 99, 106, 145, 152, 153, 183, 205 Hu, Y. Q., 152 Hybrid materials, 325 Hydride anionic complexes, 21 Hydriding capacities, 91 Hydriding chemical vapor deposition, 149 nanofibrous structure, SEM image of, 150 Hydriding combustion synthesis technique, 197 Hydrogen absorption into bulk grains, 19 kinetics, 115 in Sieverts-type apparatus, 174 alloying mills, 323 atmospheric pressure of, 95 capacity deficit, 101 decrepitation, 18 desorption DSC tests, 164 energy, 123 kinetic curves, 95 process, 121 diffusion in amorphous alloys, 11 embrittlement, 3 equilibrium plateau pressure, 90 fuel cell, 8 internal combustion engines, 22 mechanical alloying with, 37 plasma–metal reaction, 123, 147 protons, 2 refueling station, 6 reversible storage of, 26 weight percent of hydride phase and, 73 Hydrogen alloying (HA), 37, 38 Hydrogenation/dehydrogenation kinetics, 19 Hydrogen Economy, 1 implementation of, 2 Hydrogen–oxygen fuel cell, 22

Index

331

Hydrogen storage competing technologies for, 3 development of AB5 alloys for, 13–15 electrochemical, 14 for fuel cell, 15 gravimetric and volumetric densities, 178 nanomaterials for, 26 properties of intermetallic compounds, 4 selected high-capacity hydrides, 5 reservoirs, 322 solid-state, 321 for transportation applications, 3 Hydrogen storage properties ball milling MgH2 controlled mechanical milling (CMM), 104 microstructural evolution, 103 specific surface area (SSA), 104 Hypostoichiometric composites, 278 Hysteresis, absorption and desorption isotherms, 88

Johnson, W. L., 18 Jung, K. S., 166

I Ideal gas law, application of, 67 Iijima, S., 24 Inco nano-Ni, filamentary shape, 157 Inert gas, 53, 147 Intermetallic-based hydrides, 177–178 Intermetallic compounds, hydrides based on, 4 Intermetallic hydride, 9 development of nanophase AB2, 16–17 Intermetallic systems, development of interstitial hydrides in, 15–16 Intermetallic trialuminide phases, 220 Interstitial hydrogen, in disordered alloys, 50–52 Ivanov, E. J./Ivanov, E., 20, 199, 205

L Lanthanum–nickel hydride, 14 Laurencelle, F., 181 Lennard-Jones (LJ) potential, 296 LiAlH4 DSC trace of, 213–214 catalyzed with TiCl3·1/3AlCl3 activation energies of decomposition of, 219–220 activation energy of desorption of, 219 hydrogen released during ball milling of, 220 chemical reversibility of, 220–221 DSC trace of as-received, undoped, 213–214 grain size refinement of, 217 hydrogen desorption, 213 isothermal desorption, 215 isothermal decomposition of, 215 effect of ball milling on, 218 effect of catalytic metal chloride additives on, 219 Kissinger analysis of activation energy of, 216 mechanical milling with TiCl3 metastable phases, 218 reduced Ti species, 219 mechanical reversibility of, 222

J Jang, J. W., 221 Jensen, C. M., 212 Jensen, T. R., 92, 99, 147 Jeon, E., 251 Johnson-Mehl-Avrami-Kolmogorov (JMAK) absorption curves, 91 DSC peak, 168 η, dependence of, 91 η reaction order, values of, 91, 96, 98, 99, 114 theory of phase transformations, 61

K Karty, A., 91, 92 Kawai, Y., 235 Kennelley, A., 87 Kennelley, J. A., 93 Khachaturyan, A. G., 143, 144 Kim, Y., 224, 227 Kinetic energy transfer, 29 Kirchheim, R., 12 Kissinger method, 60 Koch, C. C., 17 Kojima, Y., 85, 152, 220, 222, 235 Komiya, K., 230 Konoplev, V. N., 244 Konstanchuk, I. G., 199 Kratschmer-Huffman method, 24 Kroto, H. W., 24 Kumar, L. H., 198

332 LiAlH4 (cont.) melting of deuterided, 215 rehydrogenation of difficulty in, 220 mechanochemical method for, 222 Sieverts-type measurements, 221 SEM picture of, 217 stability diagram of, 221 thermal behavior of, 214 thermal decomposition of, 216–217 Liang, G., 63, 152 LiBH4 destabilization by compositing with MgH2, MgF2, MgS and MgSe, 254 with MgH2, 253–254, 277–278 synthesis of, 240–241 Li2NH. See Lithium imide LiNH2 + LiH mixture activation energy of decomposition of, 235 basic properties examination of, 232 dehydriding behavior, 234 reduction in activation energy, 234–235 release of hydrogen and ammonia, 233–234 effect of mechanical milling on properties of, 233 NH3 intensity in, 234 rate-limiting step for dehydrogenation from, 234 LiNH2 + MgH2 mixture absorption/desorption properties of, 236 routes of activating, 237 2LiNH2–MgH2 mixtures absorption/desorption properties of effect of intensive ball milling on, 237–238 Li3N, hydrogen storage properties, 231, 232 Li, Q., 204 Li, S., 150, 152, 153 Lithium amide (LiNH2) crystallite size of, 234 decomposition of activation energy of, 234–235 ammonia formation during, 232, 233 kinetic behavior of, 233 into Li2 NH and NH3, 237 formation of, 231 grain size of, 238 hydrogen storage behavior of, 231, 233 Lithium borohydride (LiBH4) polymorphic transformation, 242 synthesis of, 242–243

Index thermal decomposition of, 241–242 structural transitions during, 243 Lithium imide absorption/desorption kinetics, 231 hydrogen capacity evolution of ammonia, 232 reversible hydrogen absorption/ desorption, 235 Lithium nitride (Li3N), 25 Lithium tetrahydroboride (LiBH4), 22 Liu, X., 198 Li, W., 150 Li, Z., 152, 153 Lomness, J. K., 174 Long-range order (LRO), 11 Løvvik, O. M., 220 Low desorption temperature hydride, 255 Lü, L., 138

M Macropores, 303 Magnesium aluminum hydride (magnesium alanate) (Mg(AlH4)2), 23 Magnesium-based ternary metal hydrides, 171 Magnesium borohydride (Mg(BH4)2) mechano-chemical activation synthesis of lattice parameters and unit cell volume, 249 powder mixture, 248 phase transformations during milling, 249–250 synthesis from elements, 244 Mg–B–H mixture, 247 Mg-2B powder, 245–247 thermal decomposition of, 244 Magnesium dihydride, 19 Magnetic induction, 36 Magneto-mill Uni-Ball-Mill, 156 milling process, 105 Makihara, Y., 260 Mamatha, M., 224, 227, 228, 230 Maurice, D. R., 32 MCAS technology, problem in, 230 Mechanical alloying (MA), 17, 32, 38 for Mg2Ni synthesis, 197 through repetitive cold welding and fracturing, 39 Mechanical amorphization (MAM), 38 Mechanical disordering (MD), 38 Mechanical milling (MM), 17 Mechanical stressing, 39 Mechanical synthesis (mechanosynthesis) (MAS), 37

Index Mechanochemical activation, 17, 44, 48 of inorganic compounds, 18 Mechanochemical activation synthesis (MCAS), 38, 54 Mechanochemical coupling, phenomena of, 37 Mechanochemistry, 18 Melt spinning, process of, 11 Mesocarbon microbeads (MCMBs). See also Activated carbons synthesis of, 303–304 Mesopores, 303 Mesoporous hydrogen storage media, 26 Metal borohydrides, 324 with high Pauling electronegativities, 253 LiBH4 (see LiBH4) with low Pauling electronegativities, 252 thermodynamical stability of, 251 Metal hydrides, 178 alkali metal hydrides, 179 development of, 18–20 heat pumps, 180 metal hydride heat pumps (MHHPs), 180 transformations, 144 Metal–hydrogen bond, 58 Metal–hydrogen systems and hydrides, 7–10 Metal/intermetallic hydrides, 322–323 Metallic alloys, 12 Metallic glasses, 11 Metallic Mg nanowires, synthesis of, 150 Metallic powders, hydrogenation of, 144 Metal–organic frameworks (MOFs), 25 Metal powder sintering, 21 Methane (CH4), 27 sp3-bonded carbon atoms, 293 Mg–Al alloys, hydrogen absorption properties, 171 Mg(AlH4)2 dehydrogenation of, 223 desorption of, 227 doped with TiCl3, dehydrogenation of, 228–229 gravimetric hydrogen capacity of, 223 hydrogen storage properties, 231 mechanical reversibility investigation, 229–230 mechanochemical activation synthesis of, 224–225 and NaCl mixture exothermic reactions, 224 phase composition of, 225 XRD patterns of powder milled, 225–226

333 preparation of mechanochemically activated metathesis reaction, 223–224 solvent mediated metathesis reactions, 223 TGA and MCAS of, 227–228 thermal decomposition of, 223 exothermic/endothermic peak due to, 228 microstructural fluctuations, 228 weight losses in, 227–228 thermodynamic stability of, 228 d-Mg(BH4)2, 249 partial decomposition of, 250 synthesis of, 251 thermal behavior of, 251 Mg–2B powder mixtures morphology of, 244–245 synthesis with oxide-free boron, 245 DSC curves of, 247 XRD patterns of, 246, 248 Mg2CoH5 crystal structure of, 204–205 hydrogen storage, 206 synthesis of by ball milling, 205 by RMA, 205–206 sintering in, 204 Mg2FeH6 crystallite (nanograin) size of, 203 desorption in Sieverts-type apparatus, 202–203 DSC traces of, 200–202 energy and electronic structure of, 204 hydrogen capacity, 204 hydrogen desorption behavior in, 202 mechano-chemical synthesis of, problems in, 203 synthesis of, 198–199 XRD measurements, 199–201 β-MgH2 atomic bonding, 106 crystal structure of, 83, 84 diffraction peaks of, 108 grain size variations of, 105 hydrogen pressures, 90 Mg, phase transformation of, 90 nanograins, growth of, 120 nanograin size and strain of, 132 nanostructuring of, 122 strains of, 106 substantial nanograin growth of, 109 XRD peaks of, 104

334 γ-MgH2 diffraction peaks of, 108 DSC hydrogen desorption curves, 122 DSC peak, 121 Ni, Bragg diffraction peaks of, 162 X-ray diffraction (XRD), 85, 106 XRD intensities, 106 MgH2–LiAlH4 composite system ball milled, desorption experiments, 260 hydrogen desorption curves, 262–264 during transformation of LiAlH4, 262 MgH2 content of decomposition of, 260 dependence of ECD on, 256, 258 desorption at atmospheric pressure, 265 desorption peak temperature maxima for, 257–258 grain size of, 256, 260 ROM behavior, 255 synthesis by ball milling desorption experiments, 260, 262 DSC curves for, 256–257, 261 scanning electron micrographs of, 257 XRD patterns of, 258–259 MgH2–NaAlH4 composite system decomposition of, 268 milled, scanning electron micrographs of, 265, 267 milled, thermal behavior of endothermic peak doublet, 265–266 exothermic peaks, 266–267 MgH2–NaBH4 composite system ball milled DSC traces of, 272, 275 particle morphology of, 270–271 scanning electron micrographs of, 279–280 XRD patterns of, 273 MgH2 constituent catalytic effect of free Mg, 270 decomposition temperature of, 273–274 desorption temperature of, 268 destabilization in nanocomposite, 274–275 DSC peak temperatures of, 277 grain size of, 272, 274 particle size, 273 refinement of particulate, 270 ROM of, 277 Mg powder content in decomposition temperature of NaBH4 and, 279

Index destabilizing effect on NaBH4, 281 XRD patterns of, 280 NaBH4 constituent decomposition temperature of, 272, 278, 279 destabilization, 278–279 temperature of melting and desorption peaks of, 275, 280 Mg-H powders ball-to-powder weight ratio, 138 β-MgH2 phase synthesized, 136 DSC desorption temperature vs. powder particle size, 141 particle size distributions and morphology, 134 powders and processing parameters, 131 reactive milling time, 135 thermodynamic parameters of, 140 XRD analysis and DSC analysis of, 133 Mg-H (Mg6H) powder synthesis absorption and desorption, ratios of, 143 PCT curves for absorption and desorption of, 141, 142 thermodynamic parameters of, 140 Mg–H system binary phase diagram, 84 crystallographic data of phases in, 84 crystal structure data, 83 hydrogen pressure, 83 MgH2 Tego Magnan powder, 100 Mg/MgH2 ABCR commercial powder, 148 ball milling hydrogen absorption of, 112 hydrogen desorption of, 115 ball-to-powder weight ratio, effect of, 139 with catalytic additives carbon/graphite and carbon nanotubes, 169–170 desorption in vacuum, 152–153 hydrogen, desorption at atmospheric pressure, 153–164 important groups of, 151 metal hydrides, 170–174 metal oxides, 164–168 crystallographic data, of phases in the Mg–H system, 84 density of, 83 desorption of Arrhenius plot for, 93 dehydrogenation of, 99 η reaction order, values of, 98, 99 kinetic curves, 94, 145, 146 moisture, 99

Index stoichiometric value, 94 thermodynamic barrier, 154 in vacuum, 167 destabilization of, 154 DSC desorption temperature vs. powder particle size, 141 enthalpy and entropy, 88 hydrogenation/dehydrogenation of, 92 hydrogen desorption temperature of, 126 hydrogen storage characteristics of absorption/desorption of, 87 ball-milled, synthesis of, 129 nanocrystalline intermetallic hydrides, synthesis of, 130 nanocrystallites, 130 hydrogen storage properties, by ball milling, 103 lattice strains of, 107 material characteristics ECD distribution, 85 hydrogen absorption and desorption characteristic, 85 hydrogen storage, 103 nanofibrous structure, SEM image of, 150 nanostructured, synthesis methods, 147–151 nanowhiskers/nanofibers, SEM image of, 149 PCT plateau pressure hysteresis, 145 powder ABCR GmbH & Co. KG, 86 DSC yield of, 137 milled, DSC curve of, 139, 140 particle size distributions and morphology, 134 stored, aging effects in, 146 Tego Magnan, 87 scanning electron micrograph of, 86 synthesized nanocrystalline MgS, 174 Tego Magnan powder DSC traces/curve of, 102, 156 PCT desorption curves at various temperatures, 102 powder particle size vs. milling time, 107 XRD patterns of, 119 thermal activation, desorption/absorption cycling, 108 thermodynamic stability, 122 Mg(NH2)2 desorption, ammonia release, 237 and LiH mixture ammonia release, 237 desorption reaction in, 240 PCT absorption curves for, 239

335 ratio of hydrides in mixture of, 238 reversible reaction in, 236 XRD phase analysis, 237 reaction to form, 238 Mg2Ni stoichiometry, 20 synthesis of, 197 Mg–Ni alloy microstructural evolution in, 197 nanostructured, synthesis of, 198 Mg2NiH4 applications in gaseous hydrogen storage, 198 crystallographic and hydrogen storage properties of, 196 technique for direct fabrication of, 197 Mg particle, cross-section, 90 Micropores, 303 Microstructural evolution, evolution in nearly dual-phase Mg–Ni alloy, 197 Molecular crystals, synthesis of, 307 Molecular orbital, 13 Mueller, W. M., 177 Multiwalled nanotubes (MWNTs), 24, 308

N NaAlH4. See Sodium alanate hydride Na3AlH6, synthesis of, 212 Nakamori, Y., 235, 241, 249, 251 Nanoalloys, 11 Nanocarbons, 324–325 development of, 23–25 Nanocrystalline alloys, 11, 28 Nanocrystalline material, 181 Nanocrystalline Mg2FeH6 controlled reactive mechanical alloying of, 200, 202, 203 reactive mechanical alloying synthesis of, 199–200 Nanocrystalline phases, 13 Nanographites, hydrogen storage capacities for, 306 Nanographitic carbons, 301 Nanohydrides, 13 Nanomaterials, 10, 17 Nanometals, 10 Nanoprocessing, methods and mechanisms for, 37 mechanical alloying, 39–40 mechanical amorphization, 55–56 mechanical milling, 38 mechanochemical activation, 40–52 mechanochemical synthesis, 52–55

336 Nanoropes, constituents of, 308 Nanostructured battery materials, 16 Nanostructured hydrides, 321 Nanostructured materials, synthesis of, 27–28 Nanotechnology, 10 bottom-up and top-down approach for, 26 Natural gas hydrogen generation stations, 321 Nickel metal hydride (NiMH) battery/cell, 14, 16, 48 Ni-doped MgH2, 158, 160 Nonactivated Tego Magnan powder, 94 Noninterstitial transition metal ternary hydrides, 20–21 Nontransformed particles, 144 Northwood, D. O., 143 Novel intermetallic hydrides, 183

O Oelerich, W., 166 Organometallic synthesis, 323 Orgaz, E., 199 Orimo, S., 177, 235, 242, 303 Oxide dispersion-strengthened (ODS) process, 17

P Particle size refinement, 42 Pauling electronegativities, 324 PCT curve, 87 PEM fuel cell (PEMFC), 2 Permanent magnet alloy, 13 γ-Phase formation, thermal cycling, 180 Photovoltaics, 1 Physical vapor deposition (PVD), 28 Plastic deformation, 39, 42 Polyacrylonitrile (PAN), 23 Polybenzimidazole (PBI), 6 Polycrystalline materials, 9 Polycrystalline structure, 45 Polymer electrolyte membrane (PEM), 2 Powder metallurgy, 29 Powder particle size distribution, 86 Pressure–composition–temperature (PCT) curve, 56 Proton-permeable polymeric (Nafion) membrane, 2 Pulverisette™ mills, 31 Puszkiel, J. A., 204

Q Qian, S., 143

Index R Rabkin, E., 144, 145 Rare-earth AB5 compounds, 181 Reactive mechanical alloying (RMA), 38, 54, 199–200 Reactive mechanical milling (RMM), 38, 53 Rechargeable battery cells, 15 Redox reaction, 53 Reilly, J. J./Reilly, J., 15, 19, 176, 177, 196–197 Reiser, A., 199 Renewable energy sources, 27 Resan, M., 217, 218, 256 Rietveld method, for profile refinement, 9 Rossi, N. L., 26 Rotations per minute (RPM), 29 Rozdzynska-Kielbik, B., 181 Ruckenstein, E., 232 Rule-of-Mixtures (ROM), 254, 255

S Sahlberg, M., 173 Saita, I., 149, 150, 182 Sánchez, C. R., 89, 92 Sandrock, G., 6, 175, 177, 178, 210–212 Scanning electron micrographs (SEM) back scattered electron (BSE) mode, 163 of Tego Magnan powder, 110 Scanning tunneling microscope (STM), 12 Schwarz, M., 223 Schwarz, R. B., 18, 143, 144 Schwickardi, M., 209 Secondary electron (SE), 71 Selvam, P., 205 SEM micrographs, micrometric Inco Ni, 157 Shao, H., 123, 147, 197, 205 Shaw, L. L., 234 Shear elastic stress, 41 Short-range order (SRO), 11 Siegel, B., 13 Sieverts, A., 9 Sieverts-type apparatus, 65, 67, 94 Singh, A. K., 197 Single-walled carbon nanohorns (CNHs) hydrogen storage, 312–313 rosette nanoparticles, 312–313 Single-walled carbon nanotubes (SWNTs), 23, 153, 169 hydrogen storage capability of, 309–311 structural relationship between graphene sheet and, 308

Index Skripnyuk, V. M., 144, 145 Slurry system, hydrogen absorption capacity of, 197–198 Sodium alanate hydride desorption/absorption properties enhanced by mechanical milling, 209 as function of heating time, 210 Ti catalyst and, 210–211 Ti clusters role in, 211–212 Ti powder role in, 212 direct synthesis of, 209 doped with Ti compounds desorption in, 212 rehydrogenation in, 209 DSC trace of as-received, undoped, 207, 208 hydrogen capacities, 213 kinetics of undoped and doped, 211 phase transition of, 208 SEM pictures of as-received and ballmilled, 207 solid state reaction, 211 thermal decomposition of TGA curve, 208–209 three-step reaction, 206–207 Sodium borohydride (NaBH4) synthesis by ball milling, 241 thermal stability of, 241–242 Soft magnetic materials, 11 Sol–gel processing, 27 Solid-state amorphization reactions (SSAR), 17 Solid-state hydrogen storage, 22 Solid-state mechanosynthesis, 55 Solid state synthesis, by ball milling, 241 Song, M. Y., 152, 169, 180 Specific surface area (SSA), 46 SPEX mill, tungsten carbide and agate vials for, 33 SPEX™ model, 17, 32 sp3-hybridized orbital, geometry of, 292 Srivastava, O. N., 103 Stacking faults, 47 disorder, 48, 301 Stampfer, J. F., Jr., 87 Standard reference material (SRM), 72 Stander, C. M., 93, 99 Stasinevich, D. S., 242, 244 Ström-Olsen, J. O., 103 Superactivated carbon, 304 Surface coatings, 38 Surface-to-volume ratio, 13 Suryanarayana, C., 138 SWNTs. See Single-walled carbon nanotubes (SWNTs) Syngas, 7, 8

337 T TDS. See Thermal desorption spectroscopy Tego Magnan powder ball-milled powders DSC curves of, 128 DSC hydrogen desorption temperatures vs. particle size, 128 DSC traces of, 117 grain size of, 110 β-MgH2, Mg, 120 hydrogen desorption, 118 hydrogen desorption curves, 119 temperature of, 126 JMAK equation, 117 MgH2 powder, 118 desorption kinetic curves, 115, 116 DSC curves, 127 particle size of, 101, 109, 111, 112 scanning electron micrographs of, 110 thermal activation of, 100 Temperature programmed desorption (TPD), 170 Ternary magnesium–iron hydride, 21 Ternary transition metal complex hydrides Mg2FeH6 (see Mg2FeH6) Mg2Ni intermetallic compound activation, 197 Mg2CoH5 (see Mg2CoH5) Mg2NiH 4 (see Mg2NiH 4) nanostructured Mg2Ni alloy, 198 Thermal activation, 43 Thermal cycling, 145 Thermal desorption spectroscopy, 172 Thermodynamic equilibrium, 144 Thermodynamic parameters, 88 TiFe synthesis reaction, 182 Titanium, and zirconium AB2 compounds, 183 Titanium–iron AB compounds, nanocrystalline TiFe, 182 Tran, N. E., 152 Turbostratic graphite, 302

U Ultrafine-grained microstructures, 9 Ultrafine magnesium absorption kinetics of, 123 particle size of, 123 Uni-Ball Mill magnetic mill/ Uni-Ball-Mill 5, 34 milling modes, 137 milling time, 137

338 V Vajeeston, P., 244 Vajo, J. J., 154, 253, 254, 258, 260, 277, 278 Van der Waals equation, 66 Van der Waals interactions, 324 Van der Walls gap, 309 Van’t Hoff equation, 57, 93 Van’t Hoff plots, 100, 123, 142, 144 Van’t Hoff relationship, 88 Varin, R. A., 99, 197 Vigeholm, B., 87, 88, 93, 95, 98 Vigeholm, V., 87 Vijay, R., 154 Volumetric technique, 89

W Wagemans, R. W. P., 122 Walton, A., 114 Wang, J., 219, 222 Wang, P., 212 Wang, X. L., 154 Waste heat, 2 Water gas reaction, 7 Williams, J. S., 130 Williamson-Hall method, 9, 41 Wiswall, R. H., 15, 19, 196–197 Wronski, Z. S./Wronski, Z., 25, 198, 222

Index X Xiong, Z., 235, 236 XPS. See X-ray photoelectron spectroscopy X-ray photoelectron spectroscopy, 146 XRD patterns, of initially milled ABCR powders, 108

Y Yan, Y., 179 Y–Co intermetallic compounds, 17 Young’s modulus, 41 Yukawa, Y., 179 Yvon, K., 171, 183, 205

Z Zaluska, A., 103, 209, 210, 212, 253 Zaluski, L., 103, 178, 197, 207, 212, 214, 222 Zeolites, 25 Zhou, D. W., 204 Zhu, M., 181 Zijlstra, H., 14 Zirconium hydrides, 9 Zlotea, C., 149 Zn(BH4)2, nanocatalyzed, 251–252 Zolliker, P., 204, 206 Züttel, A., 22, 242, 303