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Studies of hydrogen storage in the sodium borohydride and magnesium nickel hydride composite system Afonso, Louis Greg 2013

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Studies of Hydrogen Storage in the 4NaBH4/5Mg2NiH4 Composite System by Louis Greg Afonso BASc, The University of British Columbia, 2008 A THESIS SUBMITTED IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF APPLIED SCIENCE in The Faculty of Graduate Studies (Chemical and Biological Engineering) THE UNIVERSITY OF BRITISH COLUMBIA (Vancouver) January 2013 © Louis Greg Afonso 2013 Abstract Complex metal hydrides typically have high enthalpies which lead to desorption tempera- tures that are too high for practical use. Thermodynamic destabilization is one method used to lower the enthalpy of decomposition and hence lower the temperature of desorption of a complex metal hydride. A lower temperature (less than 100ºC) would enable waste heat from a PEM (Polymer Electrolyte Membrane) fuel cell to drive the hydrogen desorption re- action. NaBH4 was destabilized by ball milling NaBH4 and Mg2NiH4 in a 4:5 molar ratio, respectively. Ball milling periods of up to 2 hours did not have an effect on the thermody- namics or the kinetics of the system. Grain sizes of the two phases, NaBH4 and Mg2NiH4, were reduced during the first 30 minutes of ball milling. The decomposition enthalpy of the system was measured and found to be 67 ± 4 kJ mol−1H2 for the decomposition of Mg2NiH4 in the composite, 76 ± 5 kJ mol−1H2 for the decomposition of NaBH4 and 95 ± 7 kJ mol −1 H2 for the decomposition of NaH, which corresponds to measured desorptions at 275, 360 and 420 ºC respectively. The enthalpy of absorption corresponding to Mg2NiH4 in the com- posite was 59 ± 4 kJ mol−1H2 . During dehydrogenation of the NaBH4 phase, the ternary boride phase MgNi2.5B2 is formed under a hydrogen back pressure of vacuum, 1 bar and 5 bar. The total capacity of the system is 5.1 wt%, and a capacity loss of 2.25 wt% hydrogen was noted during cycling studies partially due to the formation of MgNi2, which is a non- hydriding phase, loss of Na from the sample holder, and the formation of large crystals of Mg that could not be hydrogenated easily. Kinetic analysis was conducted and an activation energy of 131 ± 24 kJ mol−1 was determined for the decomposition of the Mg2NiH4 phase ii of the composite. XRD phase analysis showed that the Mg2NiH4 decomposed first starting at about 275 ºC, followed by the decomposition of NaBH4 at around 360 ºC. By 400 ºC, XRD analysis showed that the MgNi2.5B2 phase had formed. The effect of cycling on the crystallographic phases showed a change from monoclinic to cubic for the Mg2NiH4 phase of the composite as well as the formation of MgNi2. iii Preface The work in this thesis culminated in the production of a manuscript “Studies of Hydrogen Storage in the 4NaBH4/5Mg2NiH4 Composite System”. This Masters thesis work has been done by Louis Greg Afonso under the direct supervision of Professor David Wilkinson and Arman Bonakdarpour in the Department of Chemical and Biological Engineering at the University of British Columbia. Identification and design of the research program was conducted by Louis Greg Afonso with guidance by Professor David Wilkinson and Arman Bonakdarpour. Performance of the various parts of the research was conducted by Louis Greg Afonso with guidance by Professor David Wilkinson and Arman Bonakdarpour. Analysis of the research data was conducted by Louis Greg Afonso with guidance by Professor David Wilkinson and Arman Bonakdarpour. Preparation of manuscripts is currently underway and is being written by Louis Greg Afonso with guidance by Professor David Wilkinson and Arman Bonakdarpour. iv Table of Contents Abstract . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ii Preface . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . iv Table of Contents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . v List of Tables . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ix List of Figures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xi Nomenclature . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xviii Acknowledgments . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xxi 1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1 1.1 Thesis Overview . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2 1.2 Thesis Layout . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3 2 Background, Literature Review and Thesis Objectives . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4 2.0.1 US Department of Energy (DOE) Targets for Hydrogen Storage . . 4 2.1 Hydrogen Storage Technologies . . . . . . . . . . . . . . . . . . . . . . . 5 2.1.1 Compressed Hydrogen Gas . . . . . . . . . . . . . . . . . . . . . 5 2.1.2 Liquid Hydrogen . . . . . . . . . . . . . . . . . . . . . . . . . . . 6 v 2.1.3 Liquid Organic Hydrogen Carriers . . . . . . . . . . . . . . . . . 7 2.1.4 Solid-State Storage . . . . . . . . . . . . . . . . . . . . . . . . . 8 2.1.5 Applications of Metal Hydride Storage Systems . . . . . . . . . . 12 2.2 Thermodynamics of Metal Hydrides . . . . . . . . . . . . . . . . . . . . 13 2.2.1 Thermodynamic Destabilization . . . . . . . . . . . . . . . . . . . 17 2.3 Kinetics . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 2.3.1 Kissinger Analysis . . . . . . . . . . . . . . . . . . . . . . . . . . 19 2.4 Previously Studied Complex Hydride Systems . . . . . . . . . . . . . . . 23 2.4.1 LiBH4 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 23 2.4.2 LiBH4 / MgH2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25 2.4.3 NaBH4 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 26 2.4.4 NaBH4 / MgH2 . . . . . . . . . . . . . . . . . . . . . . . . . . . 27 2.4.5 Mg2NiH4 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28 2.4.6 LiBH4 / Mg2NiH4 . . . . . . . . . . . . . . . . . . . . . . . . . . 29 2.5 Thesis Objectives . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 31 3 Experimental Methods . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 32 3.1 Materials Synthesis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 33 3.1.1 Mg2NiH4 synthesis . . . . . . . . . . . . . . . . . . . . . . . . . 33 3.1.2 4NaBH4 / 5Mg2NiH4 synthesis . . . . . . . . . . . . . . . . . . . 33 3.2 Experimental Methodologies . . . . . . . . . . . . . . . . . . . . . . . . . 33 3.2.1 Ball Milling Experiments . . . . . . . . . . . . . . . . . . . . . . 33 3.2.2 Thermodynamic Studies (Pressure Composition Temperature) . . . 34 3.2.3 Hydrogen Back Pressure Experiments . . . . . . . . . . . . . . . . 35 3.2.4 Reservoir Volume Effects . . . . . . . . . . . . . . . . . . . . . . 35 3.2.5 Hydrogen Absorption and Desorption Cycling Studies . . . . . . . 36 3.2.6 Hydrogen Desorption Kinetics . . . . . . . . . . . . . . . . . . . 37 vi 3.2.7 Crystallographic Phase and Reaction Path Analysis . . . . . . . . 37 3.3 Equipment Used . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 37 3.3.1 PCTPro 2000 Pressure Composition Temperature (PCT) machine . 37 3.3.2 Argon Glove Box . . . . . . . . . . . . . . . . . . . . . . . . . . 40 3.3.3 Ball Mill . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41 3.3.4 X-Ray Diffractometer . . . . . . . . . . . . . . . . . . . . . . . . 41 3.3.5 Differential Scanning Calorimeter (DSC) . . . . . . . . . . . . . . 41 3.3.6 Thermal Gravimetric Analysis (TGA) / DSC . . . . . . . . . . . . 42 4 Results and Discussion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 44 4.1 The Proposed Binary System: NaBH4 / Mg2NiH4 . . . . . . . . . . . . . 44 4.2 Preparation and Characterization of Mg2NiH4 . . . . . . . . . . . . . . . . 46 4.3 Crystallographic Phases and Reaction Pathways . . . . . . . . . . . . . . . 54 4.4 Thermodynamic Studies (Pressure-Composition-Temperature) . . . . . . . 63 4.5 Hydrogen Back Pressure Experiments . . . . . . . . . . . . . . . . . . . . 74 4.6 Ball Milling Experiments of NaBH4 / Mg2NiH4 Mixtures . . . . . . . . . 80 4.7 Hydrogen Absorption and Desorption Cycling Studies . . . . . . . . . . . 92 4.8 Hydrogen Desorption Kinetics of 4NaBH4 / 5Mg2NiH4 . . . . . . . . . . 96 4.9 Comparison of Different Complex Metal Hydride Composites . . . . . . . 100 4.10 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102 5 Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 105 5.1 Future Work . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 106 Bibliography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 108 vii A Experimental Techniques . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 114 A.1 Ball Milling for Materials Preparation . . . . . . . . . . . . . . . . . . . . 114 A.2 Pressure-Composition-Temperature (PCT) . . . . . . . . . . . . . . . . . 117 A.2.1 Van’t Hoff Analysis . . . . . . . . . . . . . . . . . . . . . . . . . 119 A.2.2 Sample Volume Calibration . . . . . . . . . . . . . . . . . . . . . 121 A.2.3 Volume Selection . . . . . . . . . . . . . . . . . . . . . . . . . . 124 A.2.4 Desorption Pressure . . . . . . . . . . . . . . . . . . . . . . . . . 126 A.2.5 Error due to Temperature Fluctuations During the Manometric Mea- surements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 127 A.3 X-Ray Diffraction (XRD) . . . . . . . . . . . . . . . . . . . . . . . . . . 129 A.3.1 Crystallographic Phases . . . . . . . . . . . . . . . . . . . . . . . 130 A.3.2 Grain and Particle Size . . . . . . . . . . . . . . . . . . . . . . . 131 A.3.3 Effect of Hydrogen Back Pressure During Sample Cooling . . . . . 132 A.4 Differential Scanning Calorimetry (DSC) . . . . . . . . . . . . . . . . . . 133 B Estimation of Na Loss From the Sample Holder . . . . . . . . . . . . . . . . 135 C Acoustic Box for Ball Mill . . . . . . . . . . . . . . . . . . . . . . . . . . . . 137 D X-Ray Diffraction Sample Holder . . . . . . . . . . . . . . . . . . . . . . . . 141 E XRD Analysis Program . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 149 viii Appendices  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  .  114 List of Tables 2.1 DOE hydrogen storage targets for vehicular application (adapted from [11]) 5 2.2 Compressed Hydrogen Storage Parameters ([13]) . . . . . . . . . . . . . . 6 2.3 Storage Parameters ([14]) . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 2.4 Overview of different hydrogen storage systems ([11]) . . . . . . . . . . . 9 2.5 Thermodynamic Data For LiBH4 [22] . . . . . . . . . . . . . . . . . . . . 24 2.6 Thermodynamic Data For 2NaBH4 / MgH2 . . . . . . . . . . . . . . . . . 25 2.7 Thermodynamic Data For NaBH4 . . . . . . . . . . . . . . . . . . . . . . 27 2.8 Thermodynamic Data For 2NaBH4+MgH2 . . . . . . . . . . . . . . . . . 28 2.9 Thermodynamic Data for Mg2NiH4 . . . . . . . . . . . . . . . . . . . . . 29 2.10 Enthalpy of desorption shown for the dehydrogenation of three composite systems: LiBH4/MgH2, LiBH4/MgH2/Ni, and LiBH4/Mg2NiH4 . . . . . 30 2.11 Thermodynamic Data for 4LiBH4+5Mg2NiH4 [4] . . . . . . . . . . . . . 30 3.1 Ball mill study ball to powder ratio . . . . . . . . . . . . . . . . . . . . . . 34 3.2 PCT isotherm desorption parameters used with the PCT Pro 2000 . . . . . 35 3.3 PCT isotherm absorption parameters used with the PCT Pro 2000 . . . . . 35 3.4 Parameters for the isothermal cycling studies . . . . . . . . . . . . . . . . 37 4.1 Thermodynamic data for NaBH4 and Mg2NiH4 under 1 bar hydrogen . . . 45 4.2 Thermodynamic data used for the 4NaBH4 / 5Mg2NiH4 desorption reaction 46 4.3 Results for enthalpy of desorption calculation . . . . . . . . . . . . . . . . 46 ix 4.4 Thermodynamic properties of various hydrides and mixtures of hydrides for 1 bar hydrogen back pressure . . . . . . . . . . . . . . . . . . . . . . . 66 4.5 Desorption Equilibrium Pressures for the 4NaBH4 / 5Mg2NiH4 system . . . 68 4.6 Comparison of desorption enthalpies and entropies, and temperatures ob- tained for 4NaBH4 / 5Mg2NiH4 . . . . . . . . . . . . . . . . . . . . . . . 73 4.7 Isothermal plateau temperatures and pressures from TPD measurements. . 79 4.8 Parameters for the isothermal cycling studies . . . . . . . . . . . . . . . . 95 4.9 Comparison of several hydride storage systems (under 1 bar hydrogen pres- sure) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 102 x List of Figures 2.1 Hydrogen storage cycle with N-ethylcarbazole [18] [©Industrial & Engi- neering Chemistry Research, 2012, by permission] . . . . . . . . . . . . . 8 2.2 Schematic of the hydrogenation of a conventional hydride. Hydrogen is first adsorbed onto the metal surface where it dissociates. It diffuses into the metal and creates the solid solution α phase, upon saturation of the α phase the β phase hydride is formed [19] [©Nature, 2001, by permission] . 10 2.3 Schematic of the orthorhombic LiBH4 showing the atomic locations of Li (red - large), B (green - medium) and H (blue - small) . . . . . . . . . . . 11 2.4 Palcan metal hydride hydrogen storage available in 600-6000 SL sizes using LaNi5H6 based hydrides. It has about 1 wt% H2 system capacity [©Palcan, 2012, by permission] . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 2.5 Pressure-Composition-Temperature (PCT) curve . . . . . . . . . . . . . . 14 2.6 PCT (a) and van’t Hoff (b) plots for a hydride. (adapted from [19]) [©Nature, 2001, by permission] . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 2.7 The effect of thermodynamic destabilization on enthalpy of formation [6] [©Scripta Materialia, 2007, by permission] . . . . . . . . . . . . . . . . . 18 2.8 Activation energy (Ea) and enthalpy (∆H) diagram of the reaction AH2 ⇋ A+H2 . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 2.9 a) MgH2 heated at several different rates for Kissinger analysis. b) closeup showing increased peak temperature with heating rate . . . . . . . . . . . 21 xi 2.10 Kissinger analysis for as received MgH2 for the data shown in Figure 2.9 . 23 2.11 Enthalpy diagram for LiBH4 showing the enthalpy for each step of a des- orption as well as the structural and phase changes. . . . . . . . . . . . . . 24 3.1 Setaram PCT Pro 2000 Pressure Composition Temperature (PCT) machine 38 3.2 Autoclave diagram showing spacers and sample holder location. . . . . . . 39 3.3 Argon glovebox . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 40 3.4 TGA/DSC temperature program . . . . . . . . . . . . . . . . . . . . . . . 43 4.1 XRD scan of synthesized Mg2NiH4 showing a dominant presence of the monoclinic Mg2NiH4 phase and a small fraction of the hexagonal Mg2NiH0.3 phase (deg in 2θ ). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47 4.2 XRD scan of Mg2NiH4 after hydrogen desorption at 450 ºC (deg in 2θ ). . . 48 4.3 The dehydrogenation of Mg2NiH4 may be limited through mass transport of hydrogen. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 48 4.4 Mg2NiH4 absorption and desorption isotherms with the plateau pressures and standard deviations shown for the 350, 400 and 450 ºC tests. . . . . . . 50 4.5 Van’t Hoff analysis for Mg2NiH4 . . . . . . . . . . . . . . . . . . . . . . . 51 4.6 DSC scans at the heating rates 4, 6, 8 and 10 ºC min−1 for Kissinger analysis 52 4.7 Kissinger analysis of the Mg2NiH4 DSC data presented in Figure 4.6. . . . 53 4.8 XRD plot of 4NaBH4 / 5Mg2NiH4 showing the phases at each of the des- orption steps for as synthesized, 300, 400 and 450 ºC. Powder diffraction file (PDF) numbers are shown in brackets to the right of the respective com- pounds (deg in 2θ ). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 55 4.9 Photo of sample holder spacer with Na compounds after heating the 4NaBH4 / 5Mg2NiH4 under vacuum. . . . . . . . . . . . . . . . . . . . . . . . . . . 57 4.10 XRD scan of the Na compounds formed on the spacer (deg in 2θ ). . . . . 58 xii 4.11 Photo of Mg crystallized on underside of sample holder lid . . . . . . . . . 59 4.12 Diagram of the sample holder, lid and spacer with the location of the Na deposit on the spacer and the Mg deposit on the lid. . . . . . . . . . . . . . 60 4.13 XRD analysis of 4NaBH4 / 5Mg2NiH4 composite after the first and fourth absorptions. The hydride mixture was hydrogenated at 57 bar and 350 ºC and dehydrogenated at about 5 bar and 450 ºC (see TPD cycling in section 4.7) (deg in 2θ ). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 62 4.14 Desorption PCTs for 4NaBH4 / 5Mg2NiH4 at 350, 400, 450 and 475 ºC. An upper plateau (A) is shown for Mg2NiH4 decomposition and low pressure plateaus for NaBH4 (B and C). . . . . . . . . . . . . . . . . . . . . . . . . 67 4.15 Absorption PCT measurements for 4NaBH4 / 5Mg2NiH4 at 400, 450 and 475 ºC. An upper plateau related to Mg2NiH4 absorption, and a lower plateau related to NaBH4 absorption can be observed . . . . . . . . . . . . 69 4.16 The natural logarithm of the wt % H absorbed and the pressure for the absorption PCT measurements of 4NaBH4 / 5Mg2NiH4 at 400, 450 and 475 ºC. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 70 4.17 Van’t Hoff analysis for the desorption PCT curves, the upper pressure plateau corresponding to the decomposition of Mg2NiH4 has the lowest enthalpy, while the decomposition of NaH has the highest enthalpy. . . . . . . . . . 71 4.18 Van’t Hoff analysis for absorption plateau I. . . . . . . . . . . . . . . . . . 73 4.19 XRD scans of 4NaBH4 / 5Mg2NiH4 after heating to 450 ºC under hydrogen pressures of vacuum, 1 bar and 5 bar (deg in 2θ ). . . . . . . . . . . . . . . 75 4.20 Impact of hydrogen back pressure on the plateaus observed during TPD experiments. Hydrogen release is shown in A, and pressure is shown in B. . 76 xiii 4.21 TPD for 3 different pressures, vacuum, 1 bar and 5 bar. The horizontal lines represent the theoretical weight percent obtained from the decomposition of the Mg2NiH4, NaBH4 and NaH components of 4NaBH4 / 5Mg2NiH4. Enthalpies are results of the PCT measurements discussed in section 4.4. . 78 4.22 Combined TGA and DSC analysis under Argon flow for 4NaBH4 / 5Mg2NiH4. 79 4.23 Van’t Hoff plot with the data from plateau II of the TPDs and the data from the PCT isotherms for plateau B compared. The data are almost collinear indicating that plateau II is not related to a new reaction step. . . . . . . . 80 4.24 TPD study of ball milled 4NaBH4 / 5Mg2NiH4 samples showing negligible impact on the onset of the desorption temperature for milling times up to 2 hours. The dashed line shows the ’apparent’ hydrogen released due to gas heating. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 82 4.25 The variation in hydrogen release and temperature with respect to the milling time is about 0.15 wt% H2. . . . . . . . . . . . . . . . . . . . . . . . . . . 83 4.26 DSC Ball Mill study for 4NaBH4 / 5Mg2NiH4 showing the major endotherms (II and III) for the powder mixture at between 325 - 350 ºC and 430 - 460 ºC. Endotherm designation: I: phase transition from monoclinic to fcc cu- bic, II: decomposition of Mg2NiH4, III: decomposition of NaBH4. . . . . . 84 4.27 XRD patterns of the ball milled 4NaBH4 / 5Mg2NiH4 powder at 10, 30, 60, and 120 minutes. The patterns are very similar, indicating no phase change. A broadening of the diffraction peaks can be observed (deg in 2θ ). . . . . 86 4.28 XRD grain size analysis of ball milled samples at 10, 30, 60 and 120 min- utes for the mixture 4NaBH4 / 5Mg2NiH4, and the grain size for the as received NaBH4 and just synthesized Mg2NiH4 are shown at time zero. . . 87 xiv 4.29 SEM analysis of as synthesized Mg2NiH4 and as received NaBH4 shown at the same scale. The Mg2NiH4 particle size is much smaller than the NaBH4. Scale 300 µm. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 88 4.30 NaBH4 ball milled for 10 minutes and 2 hours shows little difference, both show a fractured and jagged surface. Scale 300 µm. . . . . . . . . . . . . 90 4.31 SEM of 4NaBH4 / 5Mg2NiH4 ball milled for 10 and 120 minutes showing little impact on particle size. Scale 100 µm. . . . . . . . . . . . . . . . . . 91 4.32 Analysis of cycled 4NaBH4 / 5Mg2NiH4. A loss of 0.7 wt% H is observed due to the formation of MgNi2. . . . . . . . . . . . . . . . . . . . . . . . 93 4.33 4NaBH4 / 5Mg2NiH4 H2 capacity as a function of cycling. Two identical tests with different batches were run: Test A and Test B. Both tests show a similar drop in capacity. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 94 4.34 Capacity as a function of the number of cycles for isothermal cycling. . . . 96 4.35 DSC scan of the 4NaBH4 / 5Mg2NiH2 composite for different heating rates (β ). The first endotherms’ (I) peak positions increases with heating rate. The second endotherms (II) do not show a similar behavior. . . . . . . . . 97 4.36 Fitted curve for Kissinger analysis of the first DSC endotherm (I). . . . . . 99 4.37 Comparison of LiBH4 and NaBH4 based composites and the destabilizing effect of MgH2 and Mg2NiH4. . . . . . . . . . . . . . . . . . . . . . . . . 101 A.1 Schematic of high energy ball milling process . . . . . . . . . . . . . . . . 114 A.2 Mechanical alloying steps [47] [©Progress in Materials Science, 2001, by permission] . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 116 A.3 Particle and grain size reduction as a function of milling time and ball to powder ratio ([47]) [©Progress in Materials Science, 2001, by permission] 117 A.4 Simplified schematic of a Sievert apparatus . . . . . . . . . . . . . . . . . 118 A.5 PCT measurement for desorption . . . . . . . . . . . . . . . . . . . . . . . 120 xv A.6 Schematic of temperature gradient in hydrogen above the loaded sample. The temperature gradient starts at 300 ºC in the heater and ends at 25 ºC in the reservoir. The sample dead volume and spacer can be seen directly above the sample. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 122 A.7 Simplified Sieverts schematic for volume calibration . . . . . . . . . . . . 123 A.8 Illustration of driving force for absorption where the equilibrium pressure is the solid black curve and two absorption pressures are applied: a very small driving force and a large driving force. . . . . . . . . . . . . . . . . . 126 A.9 A simplified schematic diagram of PCT equipment . . . . . . . . . . . . . 127 A.10 PCT Pro 2000 reservoir temperature over about 2.5 days . . . . . . . . . . 128 A.11 Gas expansion during a TPD measurement of an empty sample holder. The rise in pressure can be interpreted as hydrogen release. Data are shown for 3 different reservoir volumes. . . . . . . . . . . . . . . . . . . . . . . . . 129 A.12 Multiple crystallites (grains) and particles. . . . . . . . . . . . . . . . . . 131 A.13 Effect of hydrogen back pressure during cooling on hydride composition . . 132 A.14 DSC scan for the desorption of the MgH2 . . . . . . . . . . . . . . . . . . 133 B.1 Photo of Na on spacer in reactor . . . . . . . . . . . . . . . . . . . . . . . 135 B.2 Solidworks rendering of Na on sample holder spacer . . . . . . . . . . . . 136 C.1 Rendered drawing of acoustic box with ventilation system exposed . . . . . 138 C.2 Acoustic box after completion . . . . . . . . . . . . . . . . . . . . . . . . 139 C.3 Electronics control box on side of acoustic box . . . . . . . . . . . . . . . 140 D.1 XRD scan of aluminum with and with out Kapton (deg in 2θ ). . . . . . . . 142 D.2 Photograph of XRD airtight sample holder . . . . . . . . . . . . . . . . . . 143 D.3 Sample loading tray for airtight XRD sample holder . . . . . . . . . . . . 143 D.4 Simple XRD sample holder . . . . . . . . . . . . . . . . . . . . . . . . . 144 xvi D.5 Ring part for simple XRD sample holder . . . . . . . . . . . . . . . . . . . 145 D.6 Base part for simple XRD sample holder . . . . . . . . . . . . . . . . . . . 146 D.7 Sample pan part for simple XRD sample holder . . . . . . . . . . . . . . . 147 D.8 Screw part for simple XRD sample holder . . . . . . . . . . . . . . . . . . 148 E.1 Fit page for XRD program . . . . . . . . . . . . . . . . . . . . . . . . . . 150 E.2 Graphical output from XRD program showing the original XRD scan (solid blue) overlaid with the fitted peaks with baseline removal (multiple colours). The PDF file is also shown for peak matching along with the grain size is placed to the right side of each peak . . . . . . . . . . . . . . . . . . . . . 152 xvii Nomenclature DOE Department of Energy α The unsaturated hydrogenated phase a hydride β The saturated hydrogenated phase of a hydride βFWHM Full width half max of peak height βramp Heating rate for kissinger analysis ∆H Change in enthalpy ∆S Change in entropy η Weight factor for pseudo - Voight function γ - MgH2 High pressure gamma phase magnesium hydride λ Wavelength of x-ray θ angle of incidence d Interplanar distance between atoms dg Grain size DSC Differential scanning calorimetry Ea Activation energy xviii fcc Face centered cubic FWHM Full width half maximum G Gaussian function k0 Constant for Kissinger analysis L Lorentzian function LOHC Liquid organic hydrogen carriers n number of moles NSERC Natural Sciences and Engineering Council P Pressure P0 Atmospheric pressure PR Reactor pressure PS Reservoir pressure PT Total pressure PCT Pressure composition temperature PDF Powder diffraction file PEM Polymer electrolyte membrane R Universal gas constant SEM Scanning electron microscopy T Temperature xix Tplateau Temperature of PCT isotherm plateau TGA Thermogravimetric analysis TPD Temperature programmed desorption V Volume VR Reactor volume VS System volume XRD X-ray diffraction xx Acknowledgments I wish to thank my supervisors Dr. David Wilkinson and Dr. Arman Bonakdarpour for the opportunity to participate in trying to make this world a better place. I would also like to thank them for their endless patience and guidance; it is truly an honour to have them as my mentors. In addition, I wish to thank Rob Hui at the National Research Council Institute for Fuel Cell Innovation (now Energy, Mining and Environment) for putting up with my constant requests for building access and allowing me the use of their facilities, without which I would not have been able to complete my research. To Jasna Jankovich, who planted the seed that started me off on this journey. I would also like to acknowledge NSERC and H2Can for funding to make this research possible. To my mom Christine Zarazun who has shown me throughout my life that anything can be overcome. Her single handed determination to never settle for what life has given her and to strive for better despite the odds has been a continuing inspiration for me as I move through my life. To Kar-Yee whose support and shelter from the storm was greatly needed. And to my son Rowan, who makes every day worth living. I can’t wait until he can write about how inspirational I was to him in his own thesis acknowledgments! xxi Chapter 1 Introduction Energy carriers are currently under intense research and hydrogen is one of the top con- tenders because it can be produced with renewable energy technologies (e.g., water elec- trolysis), it is very light, non-toxic, and can be used to generate power without harmful emissions. The interest in new energy carriers stems from our growing use of energy and the increasingly limited available resources. Although hydrogen is not a primary energy source, as a carrier, it will enable energy sources such as wind and solar to cleanly and safely store and / or transport the energy they produce upon demand. Concerns over energy security are also playing an important role in the push to intro- duce hydrogen as an energy carrier. Traditional sources of energy such as fossil fuels are only available in specific locations which are often politically volatile, leading to instabil- ities in their price and availability. Renewable energy sources such as solar energy can lead to a decentralized energy system, effectively allowing energy production to occur any- place a renewable energy source can be situated. This is increasingly important as fossil fuel resources are finite and will eventually become scarce, leading to intense competition and political strain. Industrialization of countries such as China, India and other develop- ing countries mean rising energy demands for a population seeking the same lifestyle and standard of living seen in industrialized nations like the USA. Decentralized energy production and storage through the use of hydrogen is a goal that many researchers have been striving to make possible, however, there are a number of problems associated with the use of this new energy carrier. One particular problem is the 1 storage of hydrogen: while it has the highest gravimetric energy density of common fuels, it also has the lowest volumetric energy density. This means that while the weight associated with the storage of hydrogen may be low, it requires a large volume. Safe storage and handling of hydrogen is also very important in any potential appli- cation. Pressurized and liquified hydrogen are the current methods for storing hydrogen. However, for mass production and use of hydrogen these methods of storage are not con- sidered suitable due to the extreme pressures or low temperatures required [1, 2, 3]. As well, the energy required to compress or liquefy hydrogen, which can be up to 1/3 of the total energy stored, lowers the efficiency of hydrogen as an energy carrier making it an even more unattractive solution [2]. Currently there is significant research effort on developing materials that can store hy- drogen reversibly. One of these methods is solid state storage through metal hydrides. This approach differs from the more traditional methods of pressurization and liquefaction in that chemical reactions typically take place to liberate the hydrogen from the storage medium and are also required to store it back again. In general, heat is added to the material to release hydrogen, and pressure and heat removal are required for rehydrogenation. 1.1 Thesis Overview The overall approach is to develop a complex metal hydride system (mixture of complex metal hydrides) capable of absorbing and desorbing hydrogen at a temperature of about 100 ºC or lower. The features of such a system would be as follows: • Absorption and desorption of hydrogen at or below 100 ºC • Mixed complex hydride system • Use of non-scarce materials • Low pressure storage 2 In order to achieve these goals the thermodynamics of currently existing complex hydrides must be altered such that the enthalpy of reaction for the hydrogen release and storage is lowered. The goal of the work in this thesis is to develop a combination of complex hydrides (i.e. hydride system) to achieve this lower enthalpy and thus lower operating temperatures. 1.2 Thesis Layout This thesis is arranged into 5 main sections: i) the introduction, ii) the background, lit- erature review and thesis objectives, iii) experimental techniques and analysis, iv) results and discussion, v) conclusions and future work. The introduction will present a brief back- ground to the topic as well as the motivations involved. The literature review will survey the current state of the art in the field of hydrogen storage pertaining to metal hydrides. Exper- imental techniques will introduce and explain all of the equipment and the analysis used to characterize and understand the metal hydride systems under study. Results and discussion will go in depth into the analysis of the data collected and explain what was discovered, and finally any conclusions made from the results and discussion section will be highlighted in the conclusions section along with recommendations for future work. 3 Chapter 2 Background, Literature Review and Thesis Objectives Much research has been conducted on altering the thermodynamics of metal hydrides [4, 5, 6, 7, 8]. In particular destabilized LiBH4 mixtures (either LiBH4 / MgH2, LiBH4 / Mg2NiH4 etc.) have been studied in the past few years. LiBH4 is a readily available chemical commonly used as a reducing agent for esters [9]. Unfortunately, lithium is a relatively scarce element, and its use in the lithium ion battery is continuing to rise. It is projected that by 2025 its use alone in the lithium ion battery will exceed the amount available on the planet [10] creating the potential for limited supply. An alternative to LiBH4 is NaBH4 and although not as gravimetrically appealing (18.4 vs 10.66 wt%) it has the advantage of utilizing one of the most abundant elements on earth, sodium. The compound NaBH4 is also used commonly as a reducing agent and as such is already available on a commercial scale as well. 2.0.1 US Department of Energy (DOE) Targets for Hydrogen Storage For automotive applications, the US Department of Energy (DOE) has set requirements for storage capacity, reaction rates, and cycling reversibility as presented in table 2.1 [11]. In Canada there are various research groups working on this subject and there is also a nation-wide NSERC research network (H2Can) with a five year time frame which brings researchers across Canada to work on hydrogen storage materials, production, handling and 4 safety [12]. Table 2.1: DOE hydrogen storage targets for vehicular application (adapted from [11]) Targets (as of 2009) Storage parameter Units 2010 2015 Ultimate System gravimetric kWh kg−1 1.5 1.8 2.5 capacity wt% H2 4.5 5.5 7.5 System volumetric kWh L−1 0.9 1.3 2.3 capacity g H2 L−1 28 40 70 H2 delivery ºC -40 to 85 -40 to 85 -40 to 95-105temperature (to FC) Operating pressure MPa 0.5/1.2(min/max) Kinetics (g H2 s−1) kW−1 0.02 0.02 0.02 2.1 Hydrogen Storage Technologies 2.1.1 Compressed Hydrogen Gas Compressing hydrogen is the most common way of storing it. The present generation of fuel cell vehicles use 70 MPa compressed hydrogen tanks to store their hydrogen, which is about 5.6 kg of usable hydrogen [11]. The system volumetric capacity numbers from Table 2.2 show that for 5.6 kg of H2 the volume of tank and system is about 213 L. While the gravimetric capacity at 350 bar meets the 2015 DOE targets, the volumetric capacity is quite low, and is far from meeting the 2010 targets. At 700 bar the gravimetric capacity of the system decreases and does not meet the 2015 targets due to the extra weight of the tank required to hold the higher pressure. The volumetric system capacity at 700 bar is closer to the 2010 targets. 5 Table 2.2: Compressed Hydrogen Storage Parameters ([13]) Compressed H2 DOE Targets Storage Parameter Units 350 bar 700 bar 2010 2015 Ultimate System Gravimetric wt % 5.5 5.2 4.5 5.5 7.5Capacity System Volumetric g H2 L−1 17.6 26.3 28 40 70Capacity Unfortunately, there is no foreseeable way to increase the capacity of compressed hy- drogen systems. Heavier tanks are required as the pressure is increased because thicker tank walls are required, as well the energy losses incurred by going to higher pressures will also lower the overall efficiency [11]. 2.1.2 Liquid Hydrogen Liquid hydrogen is more effective for storage of large amounts of hydrogen. The surface of the tank acts as a heat sink for the tank and as the surface area of the tank increases, the transfer of heat to the hydrogen also increases. This implies that the boil off of the cryogenic liquid is proportional to the surface area to volume ratio of the container. A larger volume has a smaller surface area to volume ratio. This means the heat loss per volume of hydrogen will be less and a large sphere will minimize the surface area for the most hydrogen stored. For a small non spherical container, which may be used in an automobile, the losses increase dramatically [9]. Reported hydrogen losses are about 25% per month for a tank left outside unused in a vehicular application [14]. Table 2.3 shows the gravimetric and volumetric capacities of liquid hydrogen. 6 Table 2.3: Storage Parameters ([14]) Liquified H2 DOE Targets Storage Parameter Units 2010 2015 Ultimate System Gravimetric kWh kg−1 2.36 1.5 1.8 2.5Capacity System Volumetric kWh L−1 1.48 0.9 1.3 2.3Capacity The efficiency of storing hydrogen as a liquid is poor in relation to other methods be- cause it requires about 1/3 of the energy stored to liquify it. This is due in part to the extremely low boiling point temperature of 22 K which requires a significant amount of en- ergy to reach. Hydrogen also has two nuclear states, one called para-hydrogen and the other ortho-hydrogen. At room temperature the ratio of para to ortho-hydrogen is 1:3. The sig- nificance of these two different states becomes apparent during hydrogen liquification. The ground state is para-hydrogen at 0 K, meaning that when hydrogen is cooled to 22 K ortho- hydrogen will convert to para-hydrogen spontaneously. Unfortunately this is an exothermic conversion and the magnitude of the enthalpy increases at lower temperatures. The enthalpy of conversion is higher than the latent heat of vaporization at 22 K, which means that the hydrogen will vaporize no matter how good the insulation of the tank is because the heat generation is originating from within the tank due to this nuclear state change. This requires that cryogenically stored hydrogen be converted to para-hydrogen while it is being stored, and is generally done with the use of catalysts to speed up the change [9]. 2.1.3 Liquid Organic Hydrogen Carriers Liquid Organic Hydrogen Carriers (LOHC) are unsaturated organic compounds that store hydrogen through hydrogenation. A large number of candidates exist, and examples of those with the appropriate requirements such as capacity and vapour pressure are hetero- cyclic aromatic compounds such as N-ethylcarbazole [15, 16]. The hydrogen capacity of 7 these systems is typically about 5-8 wt% H, and they have physical and handling charac- teristics similar to diesel or gasoline [16]. These systems require the use of precious metal catalysts such as Ru or Pd for the release/storage of hydrogen [17]. They can also be some- what volatile leading to potential evaporation of the storage media over time and so require a low vapour pressure and boiling point [16]. Figure 2.1 shows the proposed hydrogenation and dehydrogenation of N-ethylcarbazole. Figure 2.1: Hydrogen storage cycle with N-ethylcarbazole [18] [©Industrial & Engineering Chemistry Research, 2012, by permission] 2.1.4 Solid-State Storage Hydrogen can also be stored in solid material in atomic or molecular form. The advan- tages are very high volumetric density and low pressure, however, low gravimetric density and high temperature for hydrogen release remain challenges to be solved. The candidate materials for solid-state hydrogen storage can be grouped into four categories: (1) conven- tional metal hydrides, (2) complex hydrides, (3) sorbents, and (4) chemical hydrides. Table 2.4 provides a summary of these four groups [11]. The criteria listed on the left are the metrics by which hydrogen storage materials can be compared. Each of the groups have their strengths and weaknesses, and choosing one particular group over another will largely depend on application. 8 Table 2.4: Overview of different hydrogen storage systems ([11]) Conventional Complex Sorbent Chemical Hydrides Hydrides Systems Hydrides Vol. Capacity Good Good Fair Good Grav. Capacity Poor Good Good Good Reversibility Good Good Good Poor Thermodynamics Good Good Poor Fair Kinetics Good Poor Good Fair Efficiency Good Good Good Poor Conventional Hydrides Conventional hydrides are conceptually the simplest systems. They generally consist of a metal host with hydrogen atoms interstitially situated in the crystal lattice structure. Figure 2.2 shows a simple schematic of hydrogen absorption into a metal. Generally, hydrogen is heated in a pressurized vessel containing a metal powder. First molecular hydrogen is adsorbed onto the surface of the metal where it dissociates and diffuses into the host metal. Initially, a so-called α phase is formed where the hydrogen atoms are randomly distributed through out the grain. No change in the crystal structure is associated with this phase. When the solubility limit of dissolved hydrogen in the α phase is reached, a second phase, (β phase) begins to form. The β phase has a different crystallographic structure than that of the α phase. When the metal is completely transformed to the β phase it can no longer hold any more hydrogen. 9 Figure 2.2: Schematic of the hydrogenation of a conventional hydride. Hydrogen is first adsorbed onto the metal surface where it dissociates. It diffuses into the metal and creates the solid solution α phase, upon saturation of the α phase the β phase hydride is formed [19] [©Nature, 2001, by permission] The general reaction for conventional hydrides is shown in equation 2.1, where M is the metal atom(s) which forms the hydride. M+(x/2)H2⇋MHx (2.1) An example system would be LaNi5H6 which has a capacity of about 1.4 wt% H. It has excellent kinetics and thermodynamics, but the low gravimetric capacity and high cost limit its use [2, 20]. Complex Metal Hydrides Complex metal hydrides are distinguished from conventional hydrides by having a “com- plex” anion and a metal cation. An example is LiBH4 which is comprised of the complex anion BH−4 and the cation Li+. Figure 2.3 shows a schematic of the LiBH4 structure where the complex anion is clearly seen as a distinct and separate repeating unit in the overall LiBH4 unit cell [21]. 10 Figure 2.3: Schematic of the orthorhombic LiBH4 showing the atomic locations of Li (red - large), B (green - medium) and H (blue - small) Complex hydrides are known for their high hydrogen capacities with LiBH4 having a total hydrogen capacity of 18.5 wt%. There are a number of distinct families among the complex hydrides: alanates (AlH−4 ), borohydrides (BH−4 ), and amides (NH−2 ) and imides (NH2−) [22]. It is difficult to achieve hydrogen absorption / desorption reversibility in complex metal hydride based compounds. Alanate based materials consist of the alanate anion (AlH−4 ) and a metal cation. NaAlH4 is one of the most studied alanate based compounds. Bogdanovic et al. performed extensive research on NaAlH4 and found that mixing it with TiCl3 drastically improved the kinetics of desorption and enabled reversible operation of the system under moderate temperatures and pressures [23]. The dehydrogenation of NaAlH4 to NaH is shown in equations 2.2 and 2.3, the dehydrogenation is a two step reaction involving an intermediate species alanate (Na3AlH6). NaAlH4⇋1/3Na3AlH6 + 2/3Al+H2 (2.2) 11 1/3Na3AlH6⇋ NaH+ 1/3Al+ 1/2H2 (2.3) Borohydrides are a group of complex hydrides consisting of a BH−4 complex anion and a metal cation. These materials have high gravimetric capacities because of the presence of light weight boron and a high ratio of H in the complex anion. LiBH4 and NaBH4 are among the most common borohydrides studied. LiBH4 has the highest capacity of the borohydrides. Equation 2.4 shows the partial dehydrogenation of LiBH4 to release 13.9 wt% hydrogen (complete decomposition is 18.5 wt%). LiBH4⇋ LiH+B+ 3/2H2 (2.4) Amides and imides are a group of nitrogen containing hydrides. LiNH2 is lithium amide and Li2NH is lithium imide. The reaction in equation 2.5 shows the dehydrogenation of the amide to form the imide for about 6.5 wt% hydrogen [2]. LiNH2 +LiH⇋ Li2NH+H2 (2.5) The low temperature and pressure of desorption for this system has brought it attention as a possible candidate for use in fuel cell systems; however, trace amounts of NH3 has to be mitigated before use as it is poisonous to the fuel cell’s catalyst even in extremely small amounts [2]. 2.1.5 Applications of Metal Hydride Storage Systems Commercial products based on metal hydrides for storage of hydrogen already exist. For example, Palcan Energy Co. has commercialized 600 and 6000 SL sized metal hydride canisters using LaNi5 intermetallic hydride with a gravimetric capacity of about 1 to 1.5 wt%. LaNi5 alloys (and the related AB5 alloy family) are excellent materials for reversible 12 absorption / desorption of hydrogen, and form the basis of anode electrodes in rechargeable NiMH (nickel metal hydride) batteries. However, their gravimetric capacity is low (~1.5 wt%) due to presence of heavy metals. Figure 2.4 shows one of Palcan’s LaNi5 based storage system. Figure 2.4: Palcan metal hydride hydrogen storage available in 600-6000 SL sizes using LaNi5H6 based hydrides. It has about 1 wt% H2 system capacity [©Palcan, 2012, by per- mission] 2.2 Thermodynamics of Metal Hydrides An important figure of merit for metal hydrides is the enthalpy of desorption or absorption. The enthalpy correlates the pressure of hydrogen gas to the temperature of the metal hydride at equilibrium and can be expressed by the van’t Hoff equation (2.6): ln ( P P0 ) =− ∆H RT + ∆S R (2.6) Where P0 is the atmospheric pressure, ∆H and ∆S are the enthalpy and entropy change respectively during hydrogenation or dehydrogenation, T is the temperature in Kelvin, and R is the gas constant. 13 Figure 2.5 shows 3 regions of a typical pressure-composition-temperature (PCT) plot. In region I hydrogen is introduced to the metal and the so called α phase is formed. This is a solid solution phase where the hydrogen is randomly distributed in the metal host’s matrix and does not change crystal structure. As the concentration of hydrogen in the metal host increases a new phase (referred to as the β phase) begins to nucleate. The β phase has a different crystal structure than the host metal. The β phase is associated with region II on the PCT diagram and has a fairly constant pressure as the α phase is turned into the β phase. This plateau or constant pressure region signifies a co-existence of both α and β phases. In region III the pressure starts to increase; this corresponds to the full conversion of the α phase to the β phase. The pressure then increases from this point on because no more absorption can take place. Figure 2.5: Pressure-Composition-Temperature (PCT) curve For most metal hydrides the enthalpy and entropy of hydrogenation are negative (exother- 14 mic) and positive for dehydrogenation (endothermic). This means that heat is required to desorb the hydrogen, and it is released during the hydrogen absorption. Figure 2.5 also shows information about how the reversible capacity of a metal hydride is determined. Reversible capacity of a hydride is the plateau width of the PCT diagram (labeled as region II), and is less than the maximum capacity. Another phenomena seen in Figure 2.5 is the hysteresis between the absorption and desorption curves. It shows that the hydrogen pressure required for absorption (hydride formation) is greater than that of desorption (decomposition) [2]. The enthalpy value dominates the thermodynamic behavior of a given material. The entropy term is nearly constant (approximately 130 Jmol−1K−1) for solid-H2 gas systems [24, 20]. If the pressure in equation 2.6 is equal to P0 (atmospheric pressure), the Van’t Hoff relationship can be rearranged to equation 2.7: ∆H = ∆STplateau (2.7) Figure 2.6a shows the PCT data for several different temperatures. The mid-point pres- sure of each plateau enables one to determine the heat of desorption through a Van’t Hoff plot shown in figure 2.6b [2]. The slope and intercept provides ∆H and ∆S, respectively. 15 Figure 2.6: PCT (a) and van’t Hoff (b) plots for a hydride. (adapted from [19]) [©Nature, 2001, by permission] The relationship between ∆H and Tplateau at atmospheric pressure has shown that as the plateau temperature increases, the magnitude of the hydrogenation enthalpy also increases [2]. It also shows that the enthalpy of decomposition can be used to determine the temper- ature which corresponds to an equilibrium pressure of 1 bar. This is close to the pressure that a fuel cell will be operating and is a good indication of candidate materials for use in such systems. This illustrates that the maximum useful enthalpy of formation is limited to about -30 kJ mol−1H2 so that the plateau temperature is kept below 100 ºC to enable the waste heat from a PEM fuel cell to power the dehydrogenation process. Equations 2.8, 2.9 and 2.10 show the dehydrogenation of MgH2 along with the calcula- tion of the H2 desorption enthalpy, respectively. 16 MgH2⇋Mg+H2 (2.8) ∆Hdes = ∆Hf,products−∆Hf,reactants (2.9) = ( ∆HMg +∆HH2 ) − ( ∆HMgH2 ) = 74 kJ mol−1H2 (2.10) The enthalpy of desorption specifically refers to the change in energy that occurs from the release of hydrogen from a hydride and is the enthalpy of formation of the products minus the enthalpy of formation of the reactants. It is usually normalized by dividing by the number of moles of hydrogen released (∆Hdes mol−1H2 ). 2.2.1 Thermodynamic Destabilization One of the major goals in solid state hydrogen storage research today is to tune the thermo- dynamics of a hydride system such that a suitable temperature of desorption for use with a polymer electrolyte membrane (PEM) fuel cell is obtained. This translates into a hydride that can use the waste heat from the fuel cell to evolve the hydrogen, which means that it has an equilibrium pressure above 1 bar at temperatures lower than 100 ºC. Many com- pounds of interest (such as MgH2, LiBH4, ...) have desorption temperatures well above 100 ºC making them difficult to use. These high temperatures stem from the high enthalpies of desorption and absorption. One method to alter the thermodynamics of absorption and desorption is referred to as thermodynamic destabilization. A metal hydride can be destabilized by introducing other compounds that will alter the end products of dehydrogenation. This idea is conceptually presented in Figure 2.7. Dehydrogenation of two systems, AH2 and AHx + xB are shown in Figure 2.7. The reactions and the corresponding enthalpies are: 17 AH2⇋ A+H2 ∆Hdes = (∆HA +∆HH2)−∆HAH2 AH2 +xB⇋ ABx +H2 ∆Hdes = (∆HABx +∆HH2)− (∆HAH2 +∆HxB) It is the addition of B to the system that lowers the enthalpy of the reaction through the formation of ABx, and this is shown schematically by the middle line in Figure 2.7. A decrease in desorption temperature proportional to the change in enthalpy occurs. Figure 2.7: The effect of thermodynamic destabilization on enthalpy of formation [6] [©Scripta Materialia, 2007, by permission] One example of thermodynamic destabilization is for the LiBH4 / MgH2 system. The temperature of desorption for LiBH4 is about 450 ºC and the temperature of desorption for MgH2 is about 280 ºC. When the two are combined the temperature of desorption for the system is 225 ºC, showing a lowering of the temperature required to obtain hydrogen absorption and desorption [4]. 2.3 Kinetics Enthalpy of reaction provides no information about the reaction kinetics (rate of hydrogen desorption / absorption). The rate of many chemical reactions is dominated by the pres- 18 ence of activation barriers in the reaction pathway. The H2 absorption and desorption of many compounds are examples of chemical reactions which show significantly high activa- tion barriers. Figure 2.8 shows a typical activation barrier for a generic system, where the activation energy, Ea, quantifies the value of the barrier. Figure 2.8: Activation energy (Ea) and enthalpy (∆H) diagram of the reaction AH2⇋A+H2 Figure 2.8 also shows the effect of a catalyst on the system. Addition of a catalyst has no effect on the thermodynamics of the system as the nature of the reactants and products (i.e., ∆H) is unchanged, but the activation energy is reduced. Although a system may have favourable thermodynamic properties, the reaction kinetics may be too slow for reasonable dehydrogenation (and rehydrogenation) to occur. 2.3.1 Kissinger Analysis The activation energy for the desorption of a hydride can be obtained through performing the Kissinger analysis. A series of samples are tested at different heating rates βramp, where 19 the maximum temperature T of desorption for each sample shifts with the change in heating rate [24, 2]. Figure 2.9a shows the effect of heating rate on the temperature of desorption in this research for as received MgH2. As the heating rate increases, the peak temperature also increases. 20 Figure 2.9: a) MgH2 heated at several different rates for Kissinger analysis. b) closeup showing increased peak temperature with heating rate 21 Equation 2.11 shows the relationship between the activation energy Ea and the peak’s temperature (Tmax) [24]. Here βramp is the heating rate, Tmax is the temperature of desorp- tion, Ea is the activation energy, R is the universal gas constant, and k0 is a constant. ln (βramp Tmax ) =− ( Ea RTmax ) + ln(k0) (2.11) A plot of ln(βramp/Tmax) against 1000/Tmax has a slope of Ea/R from which the activation energy can be determined. Alternatively we could differentiate equation 2.12 to obtain [2]: d ln ( β T2max ) d ( 1 Tmax ) = Ea R (2.12) Now a plot of ln(βramp/T2max) vs. 1000/Tmax will be linear with a slope of Ea/R. Figure 2.10 shows the the Kissinger plot from the DSC scans in Figure 2.9. 22 Figure 2.10: Kissinger analysis for as received MgH2 for the data shown in Figure 2.9 2.4 Previously Studied Complex Hydride Systems 2.4.1 LiBH4 LiBH4 has the highest known gravimetric hydrogen capacity (18.5 wt%) among hydrides (excluding beryllium containing compounds as they are toxic). However, it has a high des- orption temperature and is very difficult to regenerate, requiring temperatures and pressures around 690 ºC and 200 bar [20]. The temperature at which LiBH4 desorbs hydrogen at 1 bar is about 450 ºC [4]. The bonding in LiBH4 is described as ionic between the [BH4]− and [Li]+ ions and as covalent between the B and H atoms [25]. The high desorption temper- ature and the difficult regeneration make this hydride unsuitable for vehicular applications (see Table 2.1 in introduction). Table 2.5 is a summary of hydrogen storage parameters for 23 LiBH4. Table 2.5: Thermodynamic Data For LiBH4 [22] ∆Hdes Measured Wt % [kJ molH2−1] Temperature [ºC] LiBH4⇋ LiH+B+ 3/2H2 67-69 [5, 9] 400-500 [2] 13.9 LiH⇋ Li+H 90.625 [26] >727 [2] 4.6 Figure 2.11 shows the desorption path and associated enthalpy for each step of desorp- tion. LiBH4 melts before the release of hydrogen. Figure 2.11: Enthalpy diagram for LiBH4 showing the enthalpy for each step of a desorption as well as the structural and phase changes. The differential scanning calorimetry (DSC) analyses for LiBH4 has three endothermic peaks. The first peak corresponds to a structural change from orthorhombic to hexagonal, the second peak is the fusion (melting) of LiBH4, and the third peak is due to the release of hydrogen [27]. 24 2.4.2 LiBH4 / MgH2 LiBH4 has one of the highest gravimetric hydrogen contents known (18.5 wt%) and would make an excellent candidate for hydrogen storage if a) the temperature of desorption was lower and b) if rehydrogenation was more practical. MgH2 is another compound with a relatively high gravimetric hydrogen content of 7.66 wt% with desorption / absorption temperatures and pressures substantially lower than that of LiBH4. The composite mixture of LiBH4+MgH2, however, has a theoretical equilibrium des- orption temperature of 225 ºC at 1 bar. This desorption reaction is shown below: 2LiBH4 +MgH2⇋MgB2 +2LiH+4H2 (2.13) The exothermic formation of the compound MgB2 reduces the overall desorption en- thalpy through thermodynamic destabilization. The enthalpy of desorption for reaction 2.13 is 46 kJ mol−1H2 [4], which is much lower than that of LiBH4 at 69 kJ mol −1 H2 . Table 2.6 is a summary of hydrogen storage parameters for the 2LiBH4 / MgH2 system. Table 2.6: Thermodynamic Data For 2NaBH4 / MgH2 ∆Hdes Measured Wt % [kJ mol−1H2 ] Temperature [ºC] 2LiBH4 +MgH2⇋ n/a n/a 14.42Li+MgB2 +5H2 2LiBH4 +MgH2⇋ 46 270 [5] 11.52LiH+MgB2+4H2 Calorimetric studies of LiBH4 / MgH2 show three endothermic peaks similar to those observed in LiBH4. The peaks correspond to a phase change, melting of LiBH4, and de- hydrogenation. No distinct peaks related to the formation of MgB2 were observed [28]. When compared to LiBH4 the temperature of desorption is about 100 ºC lower indicating thermodynamic destabilization has occurred. 25 In order to form MgB2 in the LiBH4 / MgH2 system, a back pressure of about 2-3 bar of hydrogen is required during desorption [5]. Under 1 bar back pressure amorphous B is formed and a loss of capacity is observed because the amorphous boron will not reform into LiBH4 during absorption [5, 29]. When the back pressure was too low, the dehydrogenation of LiBH4 to LiH and amorphous B was kinetically preferred over the reaction with Mg. Pinkerton et. al. argue that the back pressure kinetically suppresses the decomposition of LiBH4 and allows Mg to react and form MgB2 before LiH and B are formed [29]. While the thermodynamics and reversibility are improved, the desorption and absorp- tion temperature are still too high to meet the DOE targets. However, this system is an excellent example that thermodynamic destabilization is an effective way of improving the unfavourable thermodynamic properties of hydrogen storage materials. 2.4.3 NaBH4 Sodium borohydride (NaBH4) is a complex hydride with a hydrogen content of about 10.66 wt%. The desorption temperature of NaBH4 is about 200 ºC higher than that of LiBH4 thus making it less attractive as a hydrogen storage material. What makes it more attractive however, is the use of Na instead of Li. There are concerns about the abundance of lithium because of its use in lithium ion battery applications [10], and Na is a cheaper and much more abundant replacement. Urgnani et al. reported that the standard enthalpy of formation for NaBH4 is 191.8 kJ molH2−1 [30]. According to Milanese et al. the decomposition enthalpy of NaBH4 to Na and B is 106.8 kJ mol−1H2 [31]. A summary of thermodynamic data for NaBH4 is shown in Table 2.7. 26 Table 2.7: Thermodynamic Data For NaBH4 ∆Hdes Measured Wt % [kJ molH2−1] Temperature [ºC] NaBH4⇋ NaH+B+ 3/2H2 88.2 [32] - 8 NaH⇋ Na+ 1/2H2 114 [33] > 400 [33] 4.2 NaBH4⇋ Na+B+2H2 106.8 [31] 577 [2, 31] 10.8 Calorimetric studies of NaBH4 decomposition show that desorption occurs at tempera- tures approaching 600 ºC. A small amount of hydrogen (0.25 wt%) evolution occurs around 500 ºC when NaBH4 begins to melt [31, 2]. 2.4.4 NaBH4 / MgH2 Similar to LiBH4+MgH2, thermodynamic destabilization also occurs in the NaBH4+MgH2 system. Reaction 2.14 shows the hydrogen desorption reaction of this composite [32]. The formation of MgB2 is similar to that of the LiBH4+MgB2 system. MgB2 acts as a destabilizing agent by lowering the temperature of desorption to 450 - 490 ºC [32, 34], which is about 100 ºC lower than for NaBH4 alone. Calorimetric studies of this system show one endotherm which corresponds to the decomposition of MgH2 and a second endotherm at ≃ 500 ºC corresponding to the desorption of hydrogen from NaBH4 [31]. 2NaBH4 +MgH2⇋MgB2 +2Na+5H2 (2.14) According to Mao et al., the enthalpy of desorption for 2NaBH4 + MgH2 is 100.58 kJ mol−1H2 with 5 mol% TiF3 added as a catalyst [34], which is lower than that of NaBH4 which has a decomposition enthalpy of 106.8 kJ mol−1H2 [31]. The enthalpy of decomposition determined by Milanese et al. is 90.9 kJ mol−1H2 [31] for 2NaBH4 + MgH2. According to Garroni et al. the decomposition enthalpy for 2NaBH4+MgH2 is 62 kJ mol−1H2 with a corresponding theoretical desorption temperature of 350 ºC at 1 bar [35]. Thermodynamic 27 data for the 2NaBH4+MgH2 system are summarized in Table 2.8. Table 2.8: Thermodynamic Data For 2NaBH4+MgH2 ∆Hdes Measured Wt % [kJ mol−1H2 ] Temperature [ºC] 2NaBH4 +MgH2⇋ 62-100.58 [35, 31, 34] 450 - 490 [35, 34, 31] 9.882Na+MgB2 +5H2 2NaBH4 +MgH2⇋ 64 [26] n/a 7.912NaH+MgB2 +4H2 Garroni et al. report that varying the stoichiometry of the system from 2:1 NaBH4:MgH2 to 1:2 significantly changed the sorption properties of the system [35] changing the reac- tion pathways and corresponding temperatures of desorption. It was reported that hydrogen desorption in the 1:2 mixture completes significantly faster than the 2:1 mixture, although the 2:1 mixture has a lower calculated enthalpy (62 vs. 69 kJ molH2−1) [35, 31]. Powder agglomeration and the presence of more MgH2, which has been shown to correlate with the kinetics of NaBH4 decomposition, are two reasons for the observed difference in the kinetics [35]. Absorption for the 2:1 and 1:2 systems were carried out isothermally at 450 ºC under 50 bar H2 pressure for 1 hour. The 2:1 mixture has significantly better reversibility. It is speculated that the formation NaMgH3 is one of the reasons that full absorption is not achieved as it is not can not hold as much hydrogen [35]. 2.4.5 Mg2NiH4 Mg2NiH4 has been the focus of intense research for hydrogen storage due to its relatively low desorption temperature when compared to MgH2 [36, 37, 38, 22, 39, 4, 3]. Recently, it has also been shown to be a very effective destabilizing agent for LiBH4. Thermodynamic data of Mg2NiH4 and MgH2 are is shown in Table 2.9. 28 Table 2.9: Thermodynamic Data for Mg2NiH4 ∆Hdes Measured Wt % [kJ molH2−1] Temperature [ºC] Mg2NiH4⇋Mg2Ni+2H2 64.5 - 65.8 [2, 40] 280 [20] 3.6 MgH2⇋Mg+H2 74 [2] 310 7.66 Calorimetric measurements of Mg2NiH4 show that at about 235 ºC Mg2NiH4 undergoes a phase transformation from monoclinic to fcc cubic [38], after which dehydrogenation occurs at about 280 ºC. Mg2NiH4 is not commercially available and can be synthesized through the ball milling of MgH2 and Ni and high temperature and pressure cycling of the milled mixture at a temperature of 495 ºC and a pressure of 95 bar. As well the ball milling of Mg and Ni in a 2:1 stoichiometric ratio respectively to form Mg2Ni and hydrogenating it under 30 bar hydrogen and 350 ºC will form Mg2NiH4 [38]. 2.4.6 LiBH4 / Mg2NiH4 Mg2NiH4 has been shown to destabilize LiBH4 more effectively than MgH2 [4]. The pro- posed desorption steps for the system are: Step 1 at 250ºC: 4LiBH4 +5Mg2NiH4⇋ 2MgNi2.5B2 +4LiH+8MgH2 +8H2 (2.15) Step 2 at 340ºC: MgH2⇋Mg+H2 (2.16) The formation of MgNi2.5B2 leads to a lower enthalpy of desorption than the LiBH4 + MgH2 system. This can be seen in table 2.10 where row C shows hydrogen release at an 29 enthalpy of 15 kJ mol−1H2 . According to Vajo et al. [4], the compound MgNi2.5B2 is twice as stable as MgB2 on a per atom basis, meaning that it will destabilize the system more. This follows Miedema’s rule less stable alloys form more stable hydrides [20]; or in this case more stable alloys form less stable hydrides. It can also be seen in table 2.10 (row B) that the addition of Ni not alloyed with MgH2 has no effect on the thermodynamics because in this case it is only acting as a catalyst for the system. Table 2.10: Enthalpy of desorption shown for the dehydrogenation of three composite sys- tems: LiBH4/MgH2, LiBH4/MgH2/Ni, and LiBH4/Mg2NiH4 Composite System ∆Hdes[kJ mol−1H2 ] A 2LiBH4 +MgH2⇋ 462LiH+MgB2 +4H2 B 2LiBH4 +MgH2 +2.5Ni⇋ 462LiH+MgB2 +2.5Ni+4H2 C 4LiBH4 +5Mg2NiH4⇋ 154LiH+8MgH2 +2Mg2.5NiB2 +8H2 Hydrogen desorption in 4LiBH4+5Mg2NiH4 initiates at a temperature of about 250 ºC and has a ∆Hdes of about 15 kJ mol−1H2 . This corresponds to a theoretical desorption temperature of -25 ºC [4]. However, room temperature desorption is not observed due to kinetic limitations. Table 2.11: Thermodynamic Data for 4LiBH4+5Mg2NiH4 [4] ∆Hdes Measured Wt % [kJ mol−1H2 ] Temperature [ºC] 4LiBH4 +5Mg2NiH4⇋ 15 250 2.54LiH+8MgH2 +2Mg2.5NiB2 +8H2 30 2.5 Thesis Objectives The main objective of this thesis was to use the concept of thermodynamic destabilization to lower the enthalpy of the NaBH4 / Mg2NiH4 system to a temperature of about 100 ºC for use with a PEM fuel cell. The objectives were as follows: 1. Develop and synthesize novel complex hydride system based on Vajo et al.’s work [4]with LiBH4 using Mg2NiH4 as a thermodynamic destabilizer. 2. Synthesize Mg2NiH4 based on literature review and characterize it through XRD, DSC, and Sievert’s apparatus. 3. Determine the effect of ball milling on the complex hydride system in terms of kinet- ics, hydrogen capacity, phase changes, grain size and particle size. 4. Determine the need for a hydrogen back pressure in the system for the formation of destabilizing compounds (i.e., MgB2). 5. Demonstrate thermodynamic destabilization of the complex hydride system. 6. Evaluate the cyclability of the complex hydride system in terms of reversible capacity and phase changes during cycling. 7. Evaluate the hydrogen desorption kinetics of the complex hydride system. 8. Examine the crystallographic phase changes that occur during decomposition of the complex hydride system to determine the reaction path during desorption and absorp- tion of hydrogen. 31 Chapter 3 Experimental Methods A number of experimental techniques and analytical tools are available to synthesize and characterize materials for hydrogen storage research. Here, an overview of the techniques used in this research are presented. For synthesis of metal hydrides and complex metal hydride systems, ball milling is an effective and relatively simple mechanism to mix and prepare new materials. It is also the most common method of sample preparation for metal hydride based hydrogen storage materials. Manometric methods, where one measures the pressure changes during absorption and desorption of hydrogen is commonly used to mea- sure the hydrogen capacity of a storage material. The data obtained from manometric mea- surements can be used to determine the thermodynamics and kinetics of hydrogen uptake / release, making it ideal for such studies. The crystallographic phase changes that occur dur- ing hydrogenation and dehydrogenation are determined through X-ray diffraction (XRD). Differential scanning calorimetry (DSC) analysis is also very helpful in determining the heat of absorption and desorption. This technique can also be used to obtain the activation energy. Thermogravimetric analyses () measures the change in mass of a sample during heating and can directly measure the hydrogen storage capacity of a metal hydride. 32 3.1 Materials Synthesis 3.1.1 Mg2NiH4 synthesis Mg2NiH4 is not commercially available and had to be synthesized and characterized in the lab. Mg2NiH4 was synthesized with ball milling appropriate amounts of MgH2 (Acros 90% rest Mg) and Ni (Acros 99.9%, -325 mesh) powders. The stoichiometric ratio for MgH2 to Ni was 2:1. For a total mass of 2.4 g, 1.135 g MgH2 and 1.265 g of Ni are required. The powders were weighed and placed into a stainless steel ball mill jar with three 8 g stainless steel balls and milled for one hour in a SPEX 8000TM shaker mill. After milling the mixture was loaded into a stainless steel sample holder and heated to 495 ºC under 95 bar of pressure for 4 hours in the PCTPro 2000 apparatus. The powder was then kept at 495 ºC for 4 hours to desorb all the hydrogen, and then was at 495 ºC and 95 bar for 4 hours for hydrogen absorption to take place. All powder handling was performed in an Ar glove box with less than 0.1% O2 and a constant Ar flow rate of about 100 ml min−1. XRD was used to verify that Mg2NiH4 was made. 3.1.2 4NaBH4 / 5Mg2NiH4 synthesis NaBH4 (Acros, 98+%) was used as-received and mixed in an argon filled glovebox with the synthesized Mg2NiH4 in a 4:5 stoichiometric ratio. The mixture was then ball milled for up to 2 hours in a SPEX 8000™ ball mill with a ball to powder ratio of 10:1. 3.2 Experimental Methodologies 3.2.1 Ball Milling Experiments Ball milling experiments with different time periods were conducted to investigate the im- pact of high energy ball milling on the hydrogen desorption / absorption properties. Data 33 were collected from manometric measurements in the PCTPro 2000 to determine variations in the hydrogen desorption temperature. XRD was used for compositional or phase analy- sis, as well as grain size analysis. DSC was used to correlate phase changes with PCT data as well as determining the heat of reaction. Particle size analysis was performed with an SEM. A 2.4 g mixture of 4NaBH4 / 5Mg2NiH4 was ball milled for 10, 30, 60 and 120 minutes. A small amount of powder (about 0.31 g) was removed after each interval for XRD, SEM and DSC analysis, while the rest remained in the ball mill jar for further milling. The powder to ball ratio was kept approximately the same by removing a ball when the ball to powder ratio was changed significantly. Table 3.1 shows the ball to powder ratio and milling times for these experiments. Table 3.1: Ball mill study ball to powder ratio mill time amount of powder number of balls ball to powder ratio [min] [g] (8 g each) 10 2.4 3 10:1 30 2.09 3 11.5:1 60 1.78 2 9:1 120 1.47 2 10.9:1 Ball mills can be quite loud. In order to lower the noise from the ball mill an acoustic isolation chamber was constructed as discussed in appendix C.2. 3.2.2 Thermodynamic Studies (Pressure Composition Temperature) Pressure composition isotherm studies were made in order to determine the enthalpy of hydrogen absorption / desorption reactions. The PCT Pro 2000 was used to collect the isotherm data. Desorption and absorption isotherms were made at 350, 400, 450 and 475 ºC and at 400, 450 and 475 ºC, respectively. Desorption and absorption parameters used with the PCTPro 2000 are summarized in tables 3.2 and 3.3, respectively. The heading 34 f(∆P) refers to the algorithm used to compute the aliquot pressure in the PCTPro 2000 during desorption and absorption isotherms and is dependent on the change in concentration versus change in pressure (dC/dP). Table 3.2: PCT isotherm desorption parameters used with the PCT Pro 2000 Sample Temperature Preservoir Psample f(∆P) Reservoir Volume [ºC] [bar] [bar] [bar] [ml] 350 20 20 1.5 11.7 400 50 50 1.5 11.7 450 75 75 1.5 11.7 475 95 95 1.5 11.7 Table 3.3: PCT isotherm absorption parameters used with the PCT Pro 2000 Sample Temperature Preservoir Psample f(∆P) Reservoir Volume [ºC] [bar] [bar] [bar] [ml] 400 vacuum vacuum 1.5 11.7 450 vacuum vacuum 1.5 11.7 475 vacuum vacuum 1.5 11.7 3.2.3 Hydrogen Back Pressure Experiments The impact of back pressure on the formation of boride (MgB2, MgNi2.5B2.5) phases dur- ing the hydrogen desorption experiments will be studied. Back pressures examined were vacuum, 1 bar and 5 bar. TPD experiments were used to determine the amount of hydro- gen released as a function of time and temperature. XRD will be used to determine the crystallographic phases formed upon dehydrogenation. 3.2.4 Reservoir Volume Effects The impact of the PCTPro 2000 reservoir volumes were examined using the system’s 11.7, 167.7 and 1056.3 ml reservoirs. The size of the reservoirs impacts the gas pressure and is 35 important in hydrogen desorption of 4NaBH4 / 5Mg2NiH4 composites. 3.2.5 Hydrogen Absorption and Desorption Cycling Studies The 4NaBH4 / 5Mg2NiH4 composite system was cycled several times to determine loss of capacity. XRD was used to determine the phases present in the system in the hydrogenated and dehydrogenated state after cycling. The PCTPro 2000 was used for both non-equilibrium TPD and isothermal cycling ex- periments. For the TPD cycling experiments the sample started with a hydrogen back pres- sure of 1 bar in the 11.7 ml reservoir volume and was heated at a rate of 6 ºC min−1 to 450 ºC for desorption to happen. After this an absorption step was conducted at 400 ºC and 55 bar. The sample was then cooled under pressure. This desorption / absorption cycle was repeated four times. The difference between the TPD cycling tests and the isothermal tests were the heating and cooling of the sample between the cycles. The isothermal cycling tests in the PCTPro 2000 were set to automatically charge and discharge the sample chamber with hydrogen at a constant temperature of 450 ºC. Initially, the sample was heated at 6 ºC min−1 to 450 ºC under a hydrogen back pressure of 1 bar to desorb hydrogen. Three isothermal cycling tests were performed: • Experiment a: A hydrogen back pressure of 1-2 bar in the 1056.3 ml reservoir volume for all desorptions, and a hydrogen absorption of about 84 bar with the 167.7 ml reservoir. • Experiment b: The first desorption was into 1 bar hydrogen back pressure while the remaining desorptions were under a hydrogen back pressure of about 4.5 bar. All desorptions were into the 167.7 ml reservoir, while the absorption pressure was about 84 bar in the 167.7 ml reservoir as well. • Experiment c: The first desorption was into 1 bar hydrogen back pressure, but the 36 remaining isothermal back pressures for desorption were at about 3.5 bar with the absorption set at about 57 bar. The desorption reservoir volume for test C was 167.7 ml and the absorption volume was 11.7 ml. Table 3.4: Parameters for the isothermal cycling studies Test 1st TPD Deso Isothermal Deso Desorption Absorption Absorption back pressure back pressure Volume Pressure Volume [bar] [bar] [ml] [bar] [ml] a 1 1-2 1056.3 84 167.7 b 1 4.5 167.7 84 167.7 c 1 3.5 167.7 57 11.7 3.2.6 Hydrogen Desorption Kinetics The kinetics of hydrogen desorption were examined by the Kissinger analysis. Heat of reactions were measured with a DSC equipment using heating rates of 4, 6, 8 and 10 ºC min−1. See section 3.3.5 for more details on the DSC experiments. 3.2.7 Crystallographic Phase and Reaction Path Analysis X-Ray diffraction studies were performed for various stages of hydrogen desorption in the 4NaBH4 / 5Mg2NiH4 composite. Samples were heated to a desired temperature (e.g., 350, 400, ...ºC), and equilibriated for 1 hour. Then the sample was evacuated and allowed to cool down. The sample was placed in an airtight holder and directly placed in the XRD machine. 3.3 Equipment Used 3.3.1 PCTPro 2000 Pressure Composition Temperature (PCT) machine A PCTPro 2000 Sieverts instrument by Setaram was used to measure gas sorption properties of materials studied in this thesis. Thermodynamic properties and capacity analysis of 37 the complex hydride system were derived from the data. For temperature programmed desorption (TPD) analysis a ramp rate of 6 ºC to a temperature of 450 ºC or 500 ºC under a back pressure of vacuum, 1 bar or 5 bar was used. For more information on the use of the PCT machine see Appendix A.2. Figure 3.1 shows the PCT Pro 2000 machine used. Figure 3.1: Setaram PCT Pro 2000 Pressure Composition Temperature (PCT) machine Figure 3.2 shows a cross section of the autoclave and sample holder in the PCTPro 2000. The sample holder is inside the autoclave with two spacers on top of it to minimize 38 the dead volume. The sample holder’s lid is a porous stainless steel filter which prevents the spread of powders to the rest of the machine, and allows for the passage of gas in/out of the sample holder. The autoclave is heated with a band heater which is wrapped in an insulation jacket (Figure 3.1). Figure 3.2: Autoclave diagram showing spacers and sample holder location. Accuracy and Resolution The PCTPro 2000 is equipped with 4 capacitance pressure transducers. Two of these trans- ducers are used by the PID controlled electronic pressure regulator in the PCT Pro 2000. The other two are used for data collection, and include a high pressure sensor with a max- imum range of 0 - 200 bar (MKS 870 microbaraton) and a low pressure sensor (MKS 890 39 microbaraton) for high sensitivity with a range of 0 - 5 bar. The accuracy of the transducers is 1.0% of the reading. The resolution of the lower pressure transducer (MKS 890 micro- baraton) is about 1.1 × 10−4 bar and the resolution of the high pressure transducer (MKS 870 microbaraton) is about 33 × 10−4 bar. 3.3.2 Argon Glove Box An argon glove box (Hydrogen Electrics LLC) was used to keep all of the samples in an air and moisture reduced environment. The oxygen content was measured using an Expedition Oxygen Analyzer sensor by OxyCheq with a resolution of 0.1% O2. The oxygen level was kept below 0.3% and was typically at 0.1% or lower. The argon glove box is designed to have a constant argon purge through it at a rate of about 100 ml min−1. The purity of argon used was 99.9998%. Figure 3.3: Argon glovebox 40 3.3.3 Ball Mill A SPEX 8000® shaker ball mill was used to prepare the samples. The ball to powder ratio was 10:1, and 10 minutes to 2 hour milling times were used. The jar used with the machine is loaded in the glovebox and sealed with an o-ring under the lid to keep it air tight. A stainless steel jar and balls were used. See Appendix A.1 for more details on ball milling. 3.3.4 X-Ray Diffractometer A Bruker D8 Advance with a 0.6 mm slit, a Vantec-1 X-ray detector and a Cu Kα (λ = 1.5418 Å) X-ray source were used. The data was analyzed by EVA (version 14) and powder diffraction file (PDF) database version 9. The diffraction scans were made with a 0.02º step increment and a speed of 0.1 sec step−1 from 10 to 90º. See Appendix A.3 for more information on the fundamentals of XRD. 3.3.5 Differential Scanning Calorimeter (DSC) DSC measurements were conducted with a TA instruments DSC Q10 (V9.9 Build 303). An argon flow of 50 ml min−1 was used during the measurements. The software used for determining peak positions was Universal Analysis 2000 version 4.5A. Typically about 8- 11 mg of sample was loaded into an aluminum sample holder in an argon glove box and an aluminum lid was placed on top. This was then placed into an air tight transfer container (see section D) and transported to the DSC machine where the airtight holder was opened and the sample inside was placed into the DSC’s sample chamber. The sample chamber was then purged with argon for at least 5 minutes. Ramp rates used varied from 4 - 10 ºC min−1. The samples were all heated to 500 ºC. 41 3.3.6 Thermal Gravimetric Analysis (TGA) / DSC A Setaram SETSYS Evolution TGA with simultaneous TGA / DSC capability was used for thermogravimetric and DSC analysis. The TGA sample chamber was evacuated and then purged with an argon flow of 50 ml min−1. Heating rates of 4, 6, 8 and 10 ºC min−1 were used. The samples were all heated to 500 ºC. Typically about 6-10 mg of sample was loaded into an alumina sample holder with and without an alumina lid. This was then placed into an air tight transfer container (see section D) and transported to the TGA / DSC machine where the airtight holder was opened and the sample inside was placed onto the TGA / DSC’s balance. The chamber was then closed and evacuated. The TGA / DSC was programmed in zones as follows: • Zone 1 – Evacuate sample for 20 minutes – Purge with argon for 30 minutes – Heat to 120 ºC at about 12 ºC min−1 – Hold for 30 minutes • Zone 2 – Heat at 4, 6, 8 or 10 ºC min−1 to 500 ºC – Hold for 10 minutes at 500 ºC – Cool to 50 ºC at 10 ºC min−1 • Zone 3 – Hold at 50 ºC for 40 minutes – Ramp at 12 ºC min−1 to 120 ºC – Hold for 30 minutes • Zone 4 – Heat at 4, 6, 8 or 10 ºC min−1 to 500 ºC 42 – Hold for 10 minutes at 500 ºC – Cool to room temperature at 10 ºC min−1 – End Figure 3.4 shows the program schematically. Vacuum 20 mins 500 C 120 C 50 C 12 C/min 4, 6, 8, 10 C/min Hold 10  mins 10 C/min Zone 1 Zone 2 Zone 3 Zone 4 Figure 3.4: TGA/DSC temperature program To correct for the buoyancy effects, which arises from heating of the chamber flow gas, as empty sample holder is scanned using the same heating profile of the sample. The empty holder data are then subtracted from the sample measurements to obtain the sample mass loss during the heating. 43 Chapter 4 Results and Discussion 4.1 The Proposed Binary System: NaBH4 / Mg2NiH4 Vajo et al.’s success with the thermodynamic destabilization of the LiBH4 / Mg2NiH4 sys- tem (see section 2.4.6) naturally leads to the question of what other compounds could show similar destabilization effects. Given lithium’s relative scarcity compared to sodium, it seems relevant to consider the performance of a NaBH4 / Mg2NiH4 composite system. The theoretical hydrogen storage capacity of this system is 4.5 and 5.1 wt%, when the end products consist of NaH and Na, respectively. A possible dehydrogenation reaction could be: Step 1: 4NaBH4 +5Mg2NiH4⇋ 2MgNi2.5B2 +4NaH+8MgH2 +8H2 (4.1) Step 2: MgH2⇋Mg+H2 (4.2) 44 Table 4.1: Thermodynamic data for NaBH4 and Mg2NiH4 under 1 bar hydrogen ∆Hdes Measured Hydrogen [kJ molH2−1] Temperature [ºC] Capacity [wt%] NaBH4 → NaH+B+ 3/2H -88.2 [32] 240 - 500 [32, 20] 7.9 Mg2NiH4⇋Mg2Ni+2H2 -65 [4, 22] 280 [20] 3.6 The reaction NaBH4 ⇀ NaH + B + 3/2H2 is irreversible due to the formation of amor- phous elemental boron. However, according to Garroni et al. [32], MgB2 can be formed when MgH2 is added to NaBH4 as seen in reaction 4.3. Vajo et al.’s research with LiBH4 and LiBH4 / MgH2 shows that these systems decompose to form amorphous boron and MgB2, respectively, as well [5]. This indicates some similarity in the reactions of LiBH4 and NaBH4 based compounds. Because the NaBH4 / MgH2 system is reversible, due to the formation of MgB2, Barkhordarian et al. achieved direct absorption in the NaBH4 / MgH2 system at 300 ºC and 200 bar of hydrogen in 48 hours [8]. 2NaBH4 +MgH2⇋ 2NaH+MgB2 +4H2 (4.3) Vajo et al.’s work with 4LiBH4 / 5Mg2NiH4 has shown the formation of the ternary boride compound MgNi2.5B2 after decomposition. The similarity between the previously mentioned NaBH4 / LiBH4 / MgH2 systems suggests that a 4NaBH4 / 5Mg2NiH4 sys- tem will emulate the 4LiBH4 / 5Mg2NiH4 system and produce MgNi2.5B2 as a product of decomposition. The enthalpy of desorption for the 4NaBH4 / 5Mg2NiH4 system can be readily calculated when the product of decomposition is MgNi2.5B2, and is shown below in equation 4.4. ∆Hdes = ( 2∆Hf,MgNi2.5B2 +4∆Hf,NaH +8∆Hf,MgH2 ) − ( 4∆Hf,NaBH4 +5∆Hf,Mg2NiH4 ) (4.4) Using the data summarized in table 4.2, ∆Hdes for reaction 4.4 is calculated and pre- 45 sented in table 4.3. The data for NaH, NaBH4 and MgH2 were obtained from the NIST- JANAF Thermochemical Tables [26], and the data for Mg2NiH4 and MgNi2.5B2 were ob- tained from Vajo et al. [4]. Table 4.2: Thermodynamic data used for the 4NaBH4 / 5Mg2NiH4 desorption reaction Compound ∆Hf Standard Deviation [kJ mol−1] [kJ mol−1] NaH -56.4 ± 0.08 NaBH4 -191.84 ± 0.39 Mg2NiH4 -196 n/a MgNi2.5B2 -323 n/a MgH2 -76.15 ± 9.2 Table 4.3: Results for enthalpy of desorption calculation ∆Hdes [kJ mol−1H2 ] ∆Hdes low 24 ∆Hdes high 43 It can be seen in Table 4.3 that depending on the choice of enthalpies from literature, there are two values for ∆Hdes. They are within the range of enthalpies considered ac- ceptable for hydrogen storage (see section 2.2) and therefore further investigation into this system is needed. The NaBH4 / Mg2NiH4 composite was characterized by Temperature Programmed Des- orption (TPD), differential scanning calorimetry (DSC), X-ray diffraction (XRD) and scan- ning electron microscopy (SEM). The results are discussed in this chapter. 4.2 Preparation and Characterization of Mg2NiH4 Figure 4.1 shows the XRD profile of the Mg2NiH4 synthesized in the lab. The sample is dominantly monoclinic Mg2NiH4 with a small presence of hexagonal Mg2NiH0.3, which is 46 the incomplete transformation of the MgH2 and Ni into Mg2NiH4. The Mg2NiH4 synthe- sized is orange in color as mentioned by Sato et al. [41]. Figure 4.1: XRD scan of synthesized Mg2NiH4 showing a dominant presence of the mono- clinic Mg2NiH4 phase and a small fraction of the hexagonal Mg2NiH0.3 phase (deg in 2θ ). Figure 4.2 shows the XRD scan of Mg2NiH4 after a TPD to 450 ºC. Mg2Ni is dom- inant, along with a small amount of Mg2NiH4 and Mg2NiH0.3. It is not clear why there is Mg2NiH4 and Mg2NiH0.3 remaining after desorption, however, it is possible that hydro- gen diffusion from the cores of Mg2NiH4 and Mg2NiH0.3 becomes the limiting step in the desorption of hydrogen. Figure 4.3 illustrates this scenario. 47 Figure 4.2: XRD scan of Mg2NiH4 after hydrogen desorption at 450 ºC (deg in 2θ ). Varin et al. have reported that after ball milling MgH2, a γ-MgH2 phase is formed, and when the γ-MgH2 decomposes a small fraction of it forms an ultra-stable β -MgH2 that does not break down during heating [2]. The Mg2NiH4 and Mg2NiH0.3 phases, however, are not the product of another phase introduced by ball milling and may have formed during the cooling period before sample removal with the small amounts of hydrogen left in the sample chamber after evacuation. Figure 4.3: The dehydrogenation of Mg2NiH4 may be limited through mass transport of hydrogen. Figure 4.4 shows the PCT isotherms that were measured for Mg2NiH4. The absorption 48 plateaus are higher than the desorption plateaus indicating a hysteresis between absorption and desorption enthalpy is present. This phenomenon is common in metal hydrides and is not fully understood. One explanation is plastic or elastic deformation of the crystal lattice during hydrogen absorption and desorption. As the hydrogen enters the crystal lattice it must expand to accommodate it, and conversely as it leaves during desorption the crystal lattice contracts behind it. The plastic and elastic energies required for the insertion and removal of hydrogen are different leading to different values of enthalpy for hydrogenation and dehydrogenation [20]. The theoretical capacity for Mg2NiH4 is 3.6 wt % H, and the observed capacity ranged from about 3 to 3.6 wt %. The desorption isotherm at 450 ºC shows a capacity of about 0.5 wt % H. The reason for this is unclear, but may be due to an incomplete absorption before the desorption test. However, the plateau pressure could still be read so it was not repeated due to time constraints. 49 Figure 4.4: Mg2NiH4 absorption and desorption isotherms with the plateau pressures and standard deviations shown for the 350, 400 and 450 ºC tests. 50 Figure 4.5 shows the van’t Hoff analysis for Mg2NiH4 obtained from the results shown in Figure 4.4. The desorption and absorption enthalpies obtained are 74 ± 5 kJ mol−1H2 and -60 ± 3 kJ mol−1H2 respectively. The absorption value is close to the value reported by Reilly and Wiswall of -64.5 kJ mol−1H2 for enthalpy of formation [39]. The error bars are the standard deviation of the data set chosen to represent the plateaus from the absorption and desorption isotherms. The change in entropy associated with the decomposition of Mg2NiH4 is 135 ± 7 J mol−1H2 K −1 , which is very close to 130 J mol−1H2 K −1 that most metal hydrides exhibit [2]. Figure 4.5: Van’t Hoff analysis for Mg2NiH4 The kinetics of hydrogen desorption from Mg2NiH4 were studied by the differential calorimetric method. The activation energy of hydrogen desorption from Mg2NiH4 was 51 obtained from a Kissinger analysis as described in section 2.3.1. Figure 4.6 shows the DSC scans for Mg2NiH4 at four different scan rates β = 4, 6, 8 and 10 ºC min−1. The peaks at I occur at about 235 ºC and according to Cermak and Kral, indicate a monoclinic to face centered cubic (fcc) phase transformation [41]. The peaks at II are the dehydrogenation of Mg2NiH4 and occur at about 280 ºC. The peaks at location III are unknown and are not reported in literature. These peaks are likely a result of the Mg2Ni reacting with the aluminum sample holder used for the DSC experiment. An alumina crucible was used in a different DSC machine with Mg2NiH4 and no peak at III was observed. Figure 4.6: DSC scans at the heating rates 4, 6, 8 and 10 ºC min−1 for Kissinger analysis Figure 4.7 shows the fitted data for the calculation of activation energy and it is 73 ± 9 kJ mol−1. For an estimated error of ± 2 ºC in the peak position, the resulting error in activation energy is ± 9 kJ mol−1. Errors are due to inhomogeneities in the powder resulting in inconsistencies in its hydrogen storage properties. The history of the sample 52 can influence the properties as well, for instance if it has been exposed to air or moisture either while in transport or through long periods of storage. The activation energy for the synthesized Mg2NiH4 is different from that reported by Selvam et al. They obtained 183 kJ mol−1 for their Mg2NiH4 [40]. The discrepancy might be a consequence of the methodology used to synthesize the powder. Selvam et al. did not use ball milling in their procedure but rather started from finely ground Mg2Ni which would have significantly different kinetic properties than ball milled powder. The difference in the size of the particles, the number of grain boundaries, and the presence of an oxide layer can play a large role in the activation energy of a metal hydride [2] and is discussed further in the section 4.6 on ball milling. Figure 4.7: Kissinger analysis of the Mg2NiH4 DSC data presented in Figure 4.6. 53 4.3 Crystallographic Phases and Reaction Pathways Figure 4.8 shows the XRD patterns of the as synthesized 4NaBH4 / 5Mg2NiH4 composite (ball milled for 1 hour), heated to 300, 400 and 450 ºC under 1 bar hydrogen. The XRD pattern of the as-synthesized 4NaBH4 / 5Mg2NiH4 composite shows peaks for Mg2NiH4 and NaBH4, as well as the partially hydrogenated Mg2Ni compound Mg2NiH0.3. The pres- ence of Mg2NiH0.3 was observed in the as-synthesized Mg2NiH4 and is due to incomplete hydrogenation. By 300 ºC the Mg2NiH4 has decomposed into Mg2Ni, while the NaBH4 remains unchanged at this temperature. By 400 ºC the ternary boride phase MgNi2.5B2 has formed and only trace amounts of NaBH4 and Mg2Ni remain. A small amount of NaH is observed, suggesting that up to 400 ºC the NaBH4 has not decomposed to Na. Excess Mg from the decomposition of Mg2Ni to MgNi2.5B2 is also observed. By 450 ºC NaBH4 and Mg2Ni have completely decomposed into MgNi2.5B2, Mg, and Na / NaH. The XRD scan shows NaOH, which is likely due to reduction with oxygen and moisture contamination during handling and transport for analysis. 54 Figure 4.8: XRD plot of 4NaBH4 / 5Mg2NiH4 showing the phases at each of the desorption steps for as synthesized, 300, 400 and 450 ºC. Powder diffraction file (PDF) numbers are shown in brackets to the right of the respective compounds (deg in 2θ ). 55 The following reactions are proposed for the decomposition of the composite under 1 bar hydrogen pressure based on the XRD data: 4NaBH4 +5Mg2NiH4⇋ 4NaBH4 +5Mg2Ni+10H2 T≤ 300◦C (4.5) 4NaBH4 +5Mg2Ni⇋ 4NaH+2MgNi2.5B2 +8Mg+6H2 T≤ 400◦C (4.6) The final step proposed for the decomposition of the composite was observed under vacuum: NaH⇋ Na+H T≤ 425◦C (4.7) We noticed the presence of Na compounds (e.g. NaOH, NaO2 etc) on top of the stainless steel spacer which lies above the sample holder. The sample holder’s lid is a porous stainless steel filter (5 micron) which allows for the gas to flow in and out of the sample tube. During decomposition (~450 ºC), some Na diffuses out and deposits on the spacer piece (Na has a melting temperature of about 98 ºC [35]). Figure 4.9 shows a photo of Na compounds deposited on the spacer. The amount of Na seen here is estimated to be about 7 mg (See Appendix B) or about 20% of the total Na available in the composite. 56 A: Desorption under vacuum B: Desorption at 1 bar Figure 4.9: Photo of sample holder spacer with Na compounds after heating the 4NaBH4 / 5Mg2NiH4 under vacuum. When desorption experiments were carried out under 1 bar of hydrogen significantly less Na diffusion occurs. Figure 4.9B shows the formation of Na - based compounds for a desorption made under 1 bar. We estimate the Na loss is about 5% in this case. Figure 4.10 shows the XRD scan of the deposit on the spacer that was obtained after a desorption experiment under vacuum. Several Na containing compounds can be identified. These are NaOH, NaO2 and Na2O2. No sign of other metals present in the composite (Mg, Ni, B) can be seen. 57 Figure 4.10: XRD scan of the Na compounds formed on the spacer (deg in 2θ ). Mg also appears to diffuse and desposit on the lid. Magnesium, however, does not dif- fuse through the porous lid onto the spacer. It crystallizes on the underside of the sample holder lid when the composite is heated to 450 ºC under vacuum (see Figure 4.11). This fur- ther corroborates reaction 4.6, as excess Mg is formed when the ternary boride MgNi2.5B2 is formed. As with the Na, the amounts of Mg that formed on the lid also varied. 58 Figure 4.11: Photo of Mg crystallized on underside of sample holder lid Figure 4.12 shows a diagram of the sample holder, lid and spacer with the location of the Na deposit on the spacer and the Mg deposit on the lid. The diagram shows how the sample holder, lid and spacer are situated in the autoclave. The diffusion of Na through the filter will cause a number of problems for use of the composite in applications of hydrogen storage. The loss of Na and Mg will reduce the capacity of the system. Safety hazards and contamination of components down stream of the hydrogen storage system are possible problems that will have to be overcome. 59 Figure 4.12: Diagram of the sample holder, lid and spacer with the location of the Na deposit on the spacer and the Mg deposit on the lid. Pinkerton et al. observed foaming of the LiBH4 / MgH2 system during decomposition [29] due to LiBH4 melting. This led to the composite overflowing the sides of their open sample holder. Garroni et al. have observed NaH contamination outside of their sample holders after absorption (for a NaBH4 / MgH2 composite), suggesting it may be due to the low melting temperature of Na (~98 ºC), which allows it to escape [35]. The melting point of NaBH4 is ~505 ºC and occurs before hydrogen release [30] indicating that the addition of a destabilizing compound may reduce its melting temperature. Varin et al. observed a slight decrease in the melting point for NaBH4 with increasing MgH2, however, DSC scans 60 showed that after about 63 vol.% MgH2 the peak associated with the melting of NaBH4 disappeared [42]. The 4NaBH4 / 5Mg2NiH4 DSC scans also show no peak associated with the melting of NaBH4. However, foaming of the sample has been observed similar to Pinkerton et al.’s experience. The foaming 4NaBH4 / 5Mg2NiH4 composite may completely fill the sample holder enabling Na and Mg to get to the lid. However, as Figure 4.10 shows, only Na seems to diffuse through the stainless steel filter in the lid. As the temperature rises and the reactions proceed, Na and Mg are liberated from the composite and rise to the top of the foam due to their lower density. Na’s melting point is ~98 ºC [35], which allows it as a liquid to flow through the stainless steel filter in the lid of the sample holder and onto the spacer. Mg, however, does not pass through the filter and deposits on the inside of the sample holder lid. More Na is seen on the spacer during decomposition under vacuum because NaH decomposes under vacuum in the composite. The melting point of NaH is 800 ºC so it won’t melt and pass through. Figure 4.13 shows X-ray diffraction patterns of the 4NaBH4 / 5Mg2NiH4 composite after absorption steps. Each absorption step was preceded by a desorption step. The for- mation of MgNi2, which does not hydrogenate [39] is observed, and partially accounts for the loss of capacity discussed in section 4.7. The formation of MgNi2 may be the result of the break down of MgNi2.5B2 because the Ni rich ternary boride is closer in stoichiometry to MgNi2 than Mg2Ni. When the ternary boride MgNi2.5B2 decomposes and releases B it must either acquire more Mg, or lose excess Ni before its stoichiometry is correct. 61 Figure 4.13: XRD analysis of 4NaBH4 / 5Mg2NiH4 composite after the first and fourth ab- sorptions. The hydride mixture was hydrogenated at 57 bar and 350 ºC and dehydrogenated at about 5 bar and 450 ºC (see TPD cycling in section 4.7) (deg in 2θ ). Analyzing the X-ray diffractograms of the composite after a rehydrogenation reaction is challenging because some materials are still present in their dehydrogenated form. This could signify that either a longer hydrogenation time is required, or full rehydrogenation is not possible. If full rehydrogenation is not possible, then reaction 4.8 shows a possible re- action pathway with decomposition chemicals still present during absorption. However, the stoichiometry of the chemicals is unknown and depends on the amount of rehydrogenation 62 possible and the amount of MgNi2.5B2 converted to MgNi2 and Mg2Ni. Na+MgNi2.5B2 +Mg+H2⇋ MgNi2 +Mg2NiH4 +NaBH4+MgNi2.5B2 +Mg (4.8) The products of absorption in reaction 4.8 show the unreacted species MgNi2.5B2 and MgNi2 which suggests some irreversibility. The loss of Na is likely the reason MgNi2.5B2 does not decompose. The excess Mg is likely because MgNi2.5B2 and MgNi2 are Mg poor compounds, leaving an excess of Mg in the reactor which then seems to collect on the un- derside of the lid. The excess Mg is not seen in the XRD analysis shown in Figure 4.13 because it collected on the lid and was not part of the sample tested. The Mg that collects on the underside of the lid does not contribute to the capacity of the system by forming MgH2 during absorption. This could be due to the large size of the Mg crystals which are kinetically limiting due to large hydrogen diffusion paths and low surface area for hydrogen dissociation to occur. 4.4 Thermodynamic Studies (Pressure-Composition-Temperature) The reaction enthalpy of the 4NaBH4 / 5Mg2NiH4 composite depends on the chemical reaction pathway. Table 4.4 summarizes the enthalpies of a number of possible reactions involving LiBH4, NaBH4, Mg2NiH4 and MgH2. Table 4.4 also presents the theoretical and the observed hydrogen desorption temperatures. The desorption temperature here refers to the temperature of the sample when hydrogen pressure is 1 bar. The theoretical desorption temperatures are calculated using the van’t Hoff relation as described in equation 2.7 of 63 Chapter 2.2 and using the ∆Hdes tabulated in table 4.4. The entropy of desorption, ∆S, is taken to be constant at 130 J mol−1 K−1. The observed temperatures listed are typically generated from a TPD measurement performed under 1 bar hydrogen. The theoretical Tdes correspond to the temperature of the sample when PH2 ≈ 1 bar. This temperature is obtained from the van’t Hoff equation and using the tabulated ∆Hdes and ∆Sdes of 130 J mol−1 K−1. Heating to Tdes, however, does not usually lead to a pressure of 1 bar, and higher temperatures are required. This is because most of these reactions are kinetically limited. The last column in table 4.4 shows the observed Tdes, and reflects this fact. The influence of pathways on the desorption enthalpy can be clearly seen for the NaBH4 / Mg2NiH4 system, where a range of enthalpies that span from 34 to 61 kJ mol−1H2 could exist. For reaction 12 in table 4.4, the equilibrium temperature corresponding to 1 bar of hydrogen is -12 ºC. This means that 2.3 wt% hydrogen should be spontaneously released at room temperature and atmospheric pressure. This, however, does not happen because the reaction is slow and heating is required for any desorption to take place. There are a number of possible desorption routes for the Mg2NiH4 / NaBH4 system and table 4.4 lists several of them (9 - 12). These are some of the expected routes based on proposed reaction pathways for similar systems such as LiBH4 / Mg2NiH4 and NaBH4 / MgH2. The complete decomposition of LiBH4 (e.g., 1 in table 4.4) has an enthalpy of 96 kJ mol−1H2 . The decomposition of LiBH4 to LiH (2), however, has an enthalpy of about 68 kJ mol−1H2 , which is much lower than that required for complete decomposition. Vajo et al. reported that the addition of MgH2 to LiBH4 (3) reduces the enthalpy of decomposition to 46 kJ mol−1H2 , and has a theoretical desorption temperature of about 225 ºC [4]. They also reported that the addition of Mg2NiH4 to LiBH4 (4) further reduced the enthalpy to 15 kJ mol−1H2 . The same trend was reported for NaBH4 by Mao et al. for the addition of MgH2 (7), which reduced the enthalpy by about 30 kJ mol−1H2 . 64 The theoretical temperature of desorption for many of the entries in table 4.4 were calculated using equation 2.7 with an entropy of 130 J mol−1 K−1 as it is constant across many metal hydride systems. However, Vajo et al. reported an entropy of 62.2 ± 3 J mol−1 K−1 H2 for the LiBH4 / Mg2NiH4 system and hypothesize that it may be different because of the two complex hydride anions in the hydrogenated state [4]. The NaBH4 / Mg2NiH4 system was found in the work completed for this thesis to have an entropy that is close to 130 J mol−1 K−1. Theoretical temperatures of desorption that have a reference listed beside them have been obtained from the literature. Enthalpies shown for reactions 9- 12 were calculated using data from NIST-JANAF Thermochemical Tables [26] and Vajo et al.’s data for Mg2NiH4 [4]. The experimental values for reactions are discussed further in this section. 65 Table 4.4: Thermodynamic properties of various hydrides and mixtures of hydrides for 1 bar hydrogen back pressure # Reaction Wt% ∆Hdes Theoretical ObservedH [kJ mol−1H2 ] Tdes [ºC] Tdes [ºC] 1 LiBH4→Li+B+2H2 18.5 96 [26] 465 - 2 LiBH4→LiH+B+3/2H2 13.9 67-69 410 400-500 [5, 9] [6] [2] 3 2LiBH4+MgH2→ 11.5 46 168 - 225 2702LiH+MgB2+4H2 [4] [5, 6] [5] 4 5Mg2NiH4+4LiBH4→ 5 15 [4] -25 [4] 250 [4]2MgNi2.5B2+4LiH +8MgH2+8H2 5 NaBH4→Na+B+2H2 10.8 96 - 114 465 - 604 577[43, 26] [31, 2] 6 NaBH4→NaH+B+3/2H2 8 88 - 90 [32] 667 [32] - 7 2NaBH4+MgH2→ 9.9 62-100.58 288 - 500 450-4902Na+MgB2+5H2 [35, 34, 31] [35, 31, 34] 8 2NaBH4+MgH2→ 7.9 64 [26] 219 -2NaH+MgB2+4H2 9 5Mg2NiH4+4NaBH4→ 5.1 61 196 4502MgNi2.5B2+4Na +8Mg+18H2 11 5Mg2NiH4+4NaBH4→ 4.5 55 150 3602MgNi2.5B2+4NaH +8Mg+16H2 10 5Mg2NiH4+4NaBH4→ 2.8 49 104 -2MgNi2.5B2+4Na +8MgH2+10H2 12 5Mg2NiH4+4NaBH4→ 2.3 34 -12 -2MgNi2.5B2+4NaH +8MgH2+8H2 13 MgH2→Mg+H2 7.6 75 [2] 304 310 14 Mg2NiH4→Mg2Ni+2H2 3.4 65.8 233 280 [20] Figures 4.15 and 4.14 show the absorption and desorption PCTs, respectively for the 4NaBH4 / 5Mg2NiH4 system. The desorption equilibrium PCT measurements exhibit a number of pressure plateaus: A, B, and C, and correspond to the decomposition of Mg2NiH4, NaBH4, and NaH, respectively (discussed further in section 4.3). The desorp- 66 tion isotherm at 475 ºC appears to be reabsorbing hydrogen as the pressure is decreased. The reason for this is unclear, but may be due to instrumental error, possibly a change in temperature of the reservoir volume. Figure 4.14: Desorption PCTs for 4NaBH4 / 5Mg2NiH4 at 350, 400, 450 and 475 ºC. An upper plateau (A) is shown for Mg2NiH4 decomposition and low pressure plateaus for NaBH4 (B and C). Table 4.5 shows the equilibrium pressures and the standard deviation for plateaus A-C in Figure 4.14. 67 Table 4.5: Desorption Equilibrium Pressures for the 4NaBH4 / 5Mg2NiH4 system Temperature Pressure Uncertainty [ºC] [bar] [bar] Plateau A 350 9.96 0.91 400 25.3 1.3 450 60.4 1.8 475 85.7 2.4 Plateau B 350 0.378 0.053 400 1.01 0.27 450 2.81 0.11 475 4.28 0.13 Plateau C 400 0.213 0.01 450 0.672 0.041 475 1.15 0.1 Figure 4.15 shows the absorption isotherms for the 4NaBH4 / 5Mg2NiH4 system at 400, 450 and 475 ºC. The plateaus for absorption are not as clear as for desorption, although an upper plateau for the absorption of Mg2Ni + H2⇋Mg2NiH4 (I) and a lower plateau related to NaBH4 absorption (II) are observed. The isotherm at 450 ºC is difficult to interpret because of the plateaus shown by the ♦ and△ symbols, either of which could represent the Mg2NiH4 absorption plateau. For this thesis the upper most plateau (♦) was taken as the Mg2NiH4 absorption plateau for two reasons: 1) the plateaus at I for 400 and 475 ºC show large amounts of hydrogen absorbed with few data points, whereas the △ plateau shows a much more sloped and slow absorption, and 2) The van’t Hoff analysis has a better fit statistic for the ♦ plateau of R2 = 0.9976, compared to 0.984 for the △ plateau. 68 Figure 4.15: Absorption PCT measurements for 4NaBH4 / 5Mg2NiH4 at 400, 450 and 475 ºC. An upper plateau related to Mg2NiH4 absorption, and a lower plateau related to NaBH4 absorption can be observed To clarify the lower plateau of the absorption isotherms a log-log plot of Figure 4.15 is shown in Figure 4.16. The lower plateau pressure (II) shows little distinction between the three temperatures tested indicating that further testing at the lower pressure plateau is required. Smaller aliquots (hydrogen doses) and longer equilibration times may help differentiate the lower plateaus. 69 Figure 4.16: The natural logarithm of the wt % H absorbed and the pressure for the absorp- tion PCT measurements of 4NaBH4 / 5Mg2NiH4 at 400, 450 and 475 ºC. Figure 4.17 shows the van’t Hoff analysis of the desorption PCT curves. For the de- composition of Mg2NiH4 (plateau A) we obtain ∆H = 67 ± 4 kJ mol−1H2 . Selvam et al. have reported the value 65.8 kJ mol−1H2 [40] for the decomposition enthalpy of Mg2NiH4 alone. For the plateau associated with the decomposition of NaBH4 to NaH + B (plateau B), ∆H = 76 ± 5 kJ mol−1H2 , which is lower than, but comparable to the value of 88.2 kJ mol −1 H2 reported by Garroni et al. [32] for the decomposition of NaBH4 to NaH. For the plateau associated with the decomposition of NaH to Na, ∆H = 95 ± 7 kJ mol−1H2 , and according to Banus et al. [43] it has a decomposition enthalpy of 114 kJ mol−1H2 . The larger error for the higher enthalpies is a result of the sensitivity of the van’t Hoff plot to variations in pressure at lower pressures. The plateaus associated with NaBH4 decomposition show a reduction 70 of the decomposition enthalpy indicating thermodynamic destabilization was successfully achieved in this composite system. Figure 4.17: Van’t Hoff analysis for the desorption PCT curves, the upper pressure plateau corresponding to the decomposition of Mg2NiH4 has the lowest enthalpy, while the decom- position of NaH has the highest enthalpy. Desorption of hydrogen from Mg2NiH4, under 1 bar hydrogen back pressure initiates at 280 ºC according to Hirscher et al. [20]. In our experiment with Mg2NiH4 prepared here, desorption is observed at 300 ºC. The higher desorption temperature for the as synthesized Mg2NiH4 may be due to kinetic limitations rather than a change in thermodynamics. In the 71 4NaBH4 / 5Mg2NiH4 composite however, Mg2NiH4 decomposition under 1 bar hydrogen, commenced at 275 ºC. This shows a change in onset of about 25 ºC with respect to the same Mg2NiH4 tested under similar conditions indicating a possible kinetic effect. The decomposition of NaBH4 in the 4NaBH4 / 5Mg2NiH4 composite begins at about 360 ºC. This is significantly lower than 577 ºC for NaBH4 alone as reported by Varin et. al. [2]. It is also significantly lower than the 2NaBH4 / MgH2 system which occurs at about 450 ºC [34]. Based on the experiments performed, it is apparent that the first step of the re- action, i.e., decomposition of Mg2NiH4 does not show any thermodynamic destabilization, whereas the next step, i.e., decomposition of NaBH4, shows significantly more favourable thermodynamics. This is what we expected. The third plateau (C) associated with the de- composition of NaH to Na and H occurred at an observed temperature of about 420 ºC, and only occurred at lower than 1 atm pressures. For equilibrium pressures above atmospheric pressure, temperatures above 450 ºC are required (see plateau C in table 4.5). Figure 4.18 shows the van’t Hoff analysis of the absorption PCT curve for the hydro- genation of Mg2NiH4 (plateau I). The plateaus associated with the formation of NaBH4 could not be resolved from the PCT data collected as mentioned earlier. However, the en- thalpy of formation for the plateau associated with the hydrogenation of Mg2NiH4 is 59 ± 4 kJ mol−1H2 , which is similar to the enthalpy obtained in section 4.2 for Mg2NiH4 alone (60 ± 3 kJ mol−1H2 ). This indicates that Mg2NiH4 is likely formed from Mg2Ni, and not from an intermediate compound containing boron or other stoichiometries of Mg and Ni. 72 Figure 4.18: Van’t Hoff analysis for absorption plateau I. The equilibrium temperatures corresponding to 1 bar H2 for the enthalpies obtained by PCT measurements for plateaus A, B and C are 242 ± 17, 311 ± 39, and 450 ± 50 ºC, respectively. Table 4.6 shows a summary of the enthalpies, entropies and temperatures associated with the 4NaBH4 / 5Mg2NiH4 system. The observed temperature of desorption for plateau C was performed under vacuum. Table 4.6: Comparison of desorption enthalpies and entropies, and temperatures obtained for 4NaBH4 / 5Mg2NiH4 Plateau ∆Hdes ∆Sdes Theoretical Tdes (1 bar) Observed Tdes[kJ mol−1H2 ] [J mol−1K−1] [ºC] [ºC] A 67 ± 4 127 ± 5 242 ± 17 275 B 76 ± 5 113 ± 7 311 ± 39 360 C 94 ± 7 128 ± 10 450 ± 50 420 (vacuum) 73 As mentioned earlier, Vajo et al. reported a value of 62.2 ± 3 J mol−1 K−1 for the change in entropy of the 4LiBH4 / 5Mg2NiH4 system after desorption. They attributed it to the additional entropy associated with having two complex anions [BH4]− and [NiH4]4−. However, the entropy of the 4NaBH4 / 5Mg2NiH4 system, with the same complex anions, was close to the 130 J mol−1 K−1 that is typically reported for metal hydrides [9]. This may indicate that the reactions taking place to liberate hydrogen are quite different. 4.5 Hydrogen Back Pressure Experiments In some composites such as LiBH4 / MgH2 a hydrogen back pressure of greater than 1 bar is required during the desorption in order for the intermediate phase (e.g., MgB2) to form [6, 29]. Three pressures, vacuum, 1 bar, and 5 bar were used to investigate the effects of back pressure on the formation of MgNi2.5B2. According to Pinkerton et al. [29] hydrogen back pressure is required in the LiBH4 / MgH2 system to achieve reversibility. They report that the dehydrogenation of LiBH4 to LiH and amorphous B is kinetically favored over the reaction of Mg to form MgB2 unless the hydrogen pressure is above about 3 bar. In other words, the hydrogen pressure kinetically suppresses LiBH4 decomposition enough to allow the formation of MgB2 to occur before that of amorphous B. The 4LiBH4 / 5Mg2NiH4 system Vajo et. al. studied also required a hydrogen pressure above 3 bar to form MgNi2.5B2 [4]. This indicates that the formation of the ternary boride is kinetically limited in relation to the decomposition of LiBH4 as well. Figure 4.19 shows the XRD patterns for the three different hydrogen back pressures after heating to 450 ºC. The pressures of vacuum, 1 bar, and 5 bar hydrogen all resulted in the formation of the ternary boride phase MgNi2.5B2. This is because pure NaBH4 does not decompose before 500 ºC, and therefore will not decompose to form amorphous boron before it can react with the Mg2Ni and form MgNi2.5B2. 74 Figure 4.19: XRD scans of 4NaBH4 / 5Mg2NiH4 after heating to 450 ºC under hydrogen pressures of vacuum, 1 bar and 5 bar (deg in 2θ ). During desorption the pressure in the sample holder will increase as the composite is heated and releases hydrogen. The increase in the pressure depends on the size of the reservoir volume that is chosen, a larger increase in pressure is observed for a smaller reservoir size. Figure 4.20B shows the change in the pressure of the sample holder for three different reservoir sizes. The smallest volume (11.7 ml) approaches almost 6 bar during hydrogen desorption, while the largest volume (1056.3 ml) only reaches about 1.2 bar. Figure 4.20A shows the corresponding hydrogen release for each of the reservoir volumes. The 167.7 and 1056.3 ml reservoirs show more hydrogen release than the 11.7 ml volume. The lower hydrogen release is due to the equilibrium pressure of desorption for NaBH4 in 75 the composite. The equilibrium pressure of desorption for NaBH4 is quite low and is about 1 bar at 400 ºC and 2.8 bar at 450 ºC. Thus it will not decompose if the hydrogen back pressure is higher than these equilibrium desorption pressures. Neither the 1056.3 or the 167.7 ml reservoirs at 450 ºC are over 2.8 bar, hence the decomposition of NaBH4 can proceed, as seen by the rise in pressure after the plateaus marked I in Figure 4.20A. However, the 11.7 ml reservoir at 350 ºC is at almost 4.5 bar from the decomposition of Mg2NiH4 (which has an equilibrium desorption pressure of about 60 bar). This is above the equilibrium desorption pressure of 2.8 bar, and therefore the decomposition of NaBH4 cannot proceed until the temperature in the reactor rises to a point where the equilibrium desorption pressure is 4.5 bar. The onset of the pressure rise after plateau I is shifted to a higher temperature so that the equilibrium pressure of desorption is above the pressure in the sample holder. Figure 4.20: Impact of hydrogen back pressure on the plateaus observed during TPD ex- periments. Hydrogen release is shown in A, and pressure is shown in B. Figure 4.21 shows the TPD profile of the 4NaBH4 / 5Mg2NiH4 composite for three different hydrogen back pressures: vacuum, 1 bar and 5 bar in a constant volume reservoir (1056.3 ml). The different steps in decomposition that were determined in section 4.3 are 76 labeled: section a shows the decomposition of Mg2NiH4, section b shows the decomposi- tion of NaBH4 to NaH, and section c shows the decomposition of NaH into Na. The effect on the amount of hydrogen released from the composite due to hydrogen back pressure is similar to that observed in Figure 4.20A. The desorption under 5 bar back pressure is limited to about 3.5 wt% H for desorption to 450 ºC, which is the same as that observed for desorption into the 11.7 ml reservoir (which runs to 500 ºC). The desorption under 1 bar and the larger volumes are also the same. However, there are two major differ- ences when the composite is decomposed under vacuum: 1) - a third decomposition step is observed at c, related to the decomposition of NaH, and 2) - during the step labeled b in Figure 4.21, there is no plateau. The third decomposition step, labeled c, is observed because the pressure in the sam- ple holder is below the equilibrium pressure of desorption for NaH decomposition in the composite, which is about 0.67 bar at 450 ºC. This is a problem for any hydrogen storage system that will run under atmospheric (1 bar) operating conditions if the maximum capac- ity of the composite is required, as the temperature will have to be increased to raise the equilibrium pressure. For the decomposition of NaH a temperature of 475 ºC is required for an equilibrium desorption pressure of about 1.15 bar. The extra plateau observed during the heating of NaBH4 in step b under 1 and 5 bar hydrogen back pressure is explained by the equilibrium pressure of desorption. During des- orption, when the pressure in the sample holder reaches that of the equilibrium pressure of desorption, the result is a plateau in the TPD measurements. The pressure and the tempera- ture at the end of the plateau corresponds very well with the PCT equilibrium measurements made (section 4.4). 77 Figure 4.21: TPD for 3 different pressures, vacuum, 1 bar and 5 bar. The horizontal lines represent the theoretical weight percent obtained from the decomposition of the Mg2NiH4, NaBH4 and NaH components of 4NaBH4 / 5Mg2NiH4. Enthalpies are results of the PCT measurements discussed in section 4.4. Figure 4.22 shows a combined TGA / DSC analysis of the 4NaBH4 / 5Mg2NiH4 com- posite. The measurements were performed under flowing Argon. Similar to TPDs under vacuum, two dominant processes, namely decomposition of Mg2NiH4 and that of NaBH4 occur. Because no hydrogen back pressure is present, steps, as observed in TPD profiles are not observed here. These experiments independently support our explanation that steps at ~400 ºC (step b of Figure 4.21), observed in TPDs under pressure, are a consequence of back pressure and do not signify reaction steps. 78 Figure 4.22: Combined TGA and DSC analysis under Argon flow for 4NaBH4 / 5Mg2NiH4. Table 4.7 shows the pressures and temperatures associated with hydrogen desorption at the end of plateau II in Figure 4.20. The pressures and temperatures observed were used to construct a van’t Hoff plot. Table 4.7: Isothermal plateau temperatures and pressures from TPD measurements. Volume Pressure [bar] Temperature [ºC] 1056.3 ml 1.1 410 167.7 ml 1.5 425 11.7 ml 5.5 500 The pressures and temperatures listed in table 4.7 are a consequence of the pressure build up in the sample holder reaching that of the equilibrium desorption pressure for the composite hydride. Therefore, each pressure and temperature point should lie on the PCT 79 isotherm curve observed in plateau B of Figure 4.17. Figure 4.23 shows the van’t Hoff plot with the data from plateau II of the TPDs and the data from the PCT isotherms for plateau B compared. The results are very similar indicating that plateau II observed in TPD measurements are a consequence of the system reaching the equilibrium pressure. These plateaus are not related to formation of stable compounds. Figure 4.23: Van’t Hoff plot with the data from plateau II of the TPDs and the data from the PCT isotherms for plateau B compared. The data are almost collinear indicating that plateau II is not related to a new reaction step. 4.6 Ball Milling Experiments of NaBH4 / Mg2NiH4 Mixtures Ball milling has been shown to have a strong impact on the hydrogen sorption and des- orption properties in many hydride systems such as MgH2 [2]. In general the kinetics of 80 hydrogen absorption / desorption is improved. Varin et al. have attributed the kinetic en- hancement of ball milled MgH2 to: i) grain size reduction, ii) particle size reduction, iii) formation of new phases, and iv) break up of oxide shell surrounding the particle [2]. Grain size reduction improves hydrogen diffusion into and out of the metal hydride by moving easier through an increased number of grain boundaries. Decreased particle size increases the surface area of the hydride and reduces the diffusion distances required to achieve hydro- gen saturation in the hydride phase, and the formation of new phases with lower activation energy can improve kinetics, e.g., the γ - MgH2 phase that appears after ball milling com- mercial β - MgH2. The breakup of the oxide shell allows hydrogen dissociation to occur easier as it is much faster on a metal surface, and it also mitigates the slow diffusion through the oxide shell before entering the metal part of the particle, improving the kinetics. Figure 4.24 shows the impact of ball milling on the hydrogen desorption temperature of the ball milled Mg2NiH4 / NaBH4 mixture. Ball milling for up to 120 minutes has apparently little impact on the onset temperature of desorption. The Mg2NiH4 was also prepared fresh and had little exposure to oxygen or moisture, thus oxide phases are negligi- ble. Therefore it may not be surprising that ball milling the mixture Mg2NiH4 / NaBH4 up to 120 minutes does not show any significant impact on the hydrogen desorption properties. NaBH4 has shown little change in structure and decomposition behavior from ball milling according to Varin et al. [44, 2]. Thus overall changes are insignificant when both powders are ball milled together for up to 120 minutes. 81 Figure 4.24: TPD study of ball milled 4NaBH4 / 5Mg2NiH4 samples showing negligible impact on the onset of the desorption temperature for milling times up to 2 hours. The dashed line shows the ’apparent’ hydrogen released due to gas heating. There is a small loss in capacity with increasing the ball milling time, which has also been reported for other systems such as MgH2 [2]. Ball milling for ten minutes shows the highest capacity at about 3.6 wt%, whereas 1 to 2 hours of ball milling shows 3.4 wt% H2 release as seen in Figure 4.25. The total amount of hydrogen released from this system at 450 ºC is about 3.35 - 3.5 wt%. This is about 1 wt% less than the expected value of 4.5 wt%, when the decomposition proceeds to NaH (chapter 4.1). For these measurements, a small (11.67 ml) reservoir volume was used. The pressure rise, in these measurements, soon reached the equilibrium solid / gas pressure hence preventing further release of hydrogen. If we were to repeat these experiments a larger reservoir volume could be used to reduce 82 the impact of hydrogen back pressure. The impact of hydrogen back pressure and choice of reservoir volumes is discussed in section 4.5 in more detail. Figure 4.25: The variation in hydrogen release and temperature with respect to the milling time is about 0.15 wt% H2. DSC analysis was also conducted for each of the ball milled samples. Some small differences between the milling times were observed, however, the major endotherms, cor- responding to hydrogen desorption, were centered in the range of 325 - 350 ºC and 430 - 460 ºC (Figure 4.26). The first major endotherm (II) corresponds to the decomposition of Mg2NiH4 and the second major endotherm (III) is that of NaBH4 decomposition. The determination of the phases is discussed further in section 4.3. 83 Figure 4.26: DSC Ball Mill study for 4NaBH4 / 5Mg2NiH4 showing the major endotherms (II and III) for the powder mixture at between 325 - 350 ºC and 430 - 460 ºC. En- dotherm designation: I: phase transition from monoclinic to fcc cubic, II: decomposition of Mg2NiH4, III: decomposition of NaBH4. There is a slight endothermic event (I) which decreases with the milling time of the composite system (Figure 4.26). This feature corresponds to a phase transition of Mg2NiH4. Before decomposition, Mg2NiH4 converts from a monoclinic phase to a fcc cubic structure [41]. This may indicate that as the mixture is ball milled for longer periods of time, the 84 phase change may no longer occur. This may be because the phase change to fcc occurs during ball milling, or because ball milling suppresses the phase change to fcc. XRD scans of the powder taken after ball milling do not show the fcc phase of Mg2NiH4, indicating that a phase change from monoclinic to fcc does not occur during ball milling. Whether the fcc transformation occurs during heating is unknown. In-situ XRD may help clarify this, however, we did not have access to this equipment. Figure 4.27 shows the XRD patterns of 4NaBH4 / 5Mg2NiH4 ball milled for different periods of time. Varin et al. have reported that upon ball milling γ-MgH2 forms which has a slightly lower dehydrogenation temperature than MgH2 [2]. The γ-MgH2 usually forms under intense static pressure (8 GPa). For the 4NaBH4 / 5Mg2NiH4 mixture no phase changes were observed after ball milling. 85 Figure 4.27: XRD patterns of the ball milled 4NaBH4 / 5Mg2NiH4 powder at 10, 30, 60, and 120 minutes. The patterns are very similar, indicating no phase change. A broadening of the diffraction peaks can be observed (deg in 2θ ). Grain sizes were obtained using the broadening of the Bragg peaks (see section A.3.2) determined through a program written in Python (see appendix E) and the results are shown in Figure 4.28. For Mg2NiH4 and NaBH4, scattering angles of 2θ = 23.4◦ (A) and 2θ = 29◦ (B) were used, respectively. Each of these scattering angles is unique to its respective phase and does not overlap with other peaks (Figure 4.27). Varin et al. have reported that 86 NaBH4 is remarkably tough and resistant to ball milling resulting in little change in grain size after several hours of milling [2]. It can be seen that after 30 minutes of ball milling the grain size plateaus and does not change significantly for NaBH4. The Mg2NiH4 grain size also remains unchanged after 30 minutes. This is most likely because the powder has already been milled for an hour during synthesis. Figure 4.28: XRD grain size analysis of ball milled samples at 10, 30, 60 and 120 minutes for the mixture 4NaBH4 / 5Mg2NiH4, and the grain size for the as received NaBH4 and just synthesized Mg2NiH4 are shown at time zero. Scanning electron microscopy (SEM) analysis of the ball milled samples after each stage of milling was performed to determine the effect of milling time on the particle size and shape. Figure 4.29 shows SEM images of both Mg2NiH4 and NaBH4 before ball milling. The Mg2NiH4 particle size is much smaller than that of NaBH4. 87 Mg2NiH4 NaBH4 Figure 4.29: SEM analysis of as synthesized Mg2NiH4 and as received NaBH4 shown at the same scale. The Mg2NiH4 particle size is much smaller than the NaBH4. Scale 300 µm. Figure 4.30 shows that ball milling NaBH4 for 10 minutes to 2 hours does not yield 88 significant changes. The as received NaBH4 shows some particle size reduction and the shape has changed from a smooth rounded surface to a jagged fractured surface with ball milling. 89 NaBH4 ball milled 10 minutes NaBH4 ball milled 2 hours Figure 4.30: NaBH4 ball milled for 10 minutes and 2 hours shows little difference, both show a fractured and jagged surface. Scale 300 µm. Figure 4.31 shows that ball milling the NaBH4 / Mg2NiH4 mixture from 10 to 120 minutes also does not show any noticeable change in particle size. The insignificant change 90 in grain size and particle size due to ball milling may help explain why ball milling has no significant impact on the desorption kinetics. 10 mins 120 mins Figure 4.31: SEM of 4NaBH4 / 5Mg2NiH4 ball milled for 10 and 120 minutes showing little impact on particle size. Scale 100 µm. 91 4.7 Hydrogen Absorption and Desorption Cycling Studies The hydrogenation reversibility of a metal hydride system is important. Many systems have shown a loss of capacity after repeated cycling. In simpler systems like metal alloys (e.g., LaNi5H6, Mg2NiH4) decrepitation is the main cause [45, 2, 20]. Decrepitation is the break- ing down of the hydride particles into smaller particles, often with the segregation of the atoms in the powder, as for example in LaNi5 the separation of Ni [24]. In complex metal hydrides like borohydrides reversibility is difficult due to the high enthalpy of absorption which requires high temperatures and pressures [22]. The first cycling experiments for 4NaBH4 / 5Mg2NiH4 composite consisted of a set of four experiments as shown in Figure 4.32. The details of the experimental parameters are discussed in section 3.2.5. The desorption experiments were performed with a small (11.7 ml) reservoir volume, which limited the desorption of the 4NaBH4 / 5Mg2NiH4 composite as discussed in section 4.5. The loss of 0.7 wt% H after the first desorption occurs because of the formation of MgNi2, which is a non hydriding alloy that forms during the hydrogen absorption step from the Ni rich MgNi2.5B2 (see section 4.3). 92 Figure 4.32: Analysis of cycled 4NaBH4 / 5Mg2NiH4. A loss of 0.7 wt% H is observed due to the formation of MgNi2. Figure 4.33 shows the hydrogen capacity of two different batches of 4NaBH4 / 5Mg2NiH4 composites as a function of cycling. The capacity loss is about 0.7 wt % H during the first 4 cycles. Vajo et al. observed a similar loss in capacity of 0.6 wt% H during cycling tests for the 4LiBH4 / 5Mg2NiH4 composite, however they did not provide any explanations. They also observe a stabilization in reversible capacity after the first desorption from about 1.8 to 1.2 wt% H. They also noticed that longer ball milling times, up to 48 hours, improved the reversible capacity by up to 1.7 wt% H [4]. 93 Figure 4.33: 4NaBH4 / 5Mg2NiH4 H2 capacity as a function of cycling. Two identical tests with different batches were run: Test A and Test B. Both tests show a similar drop in capacity. In addition to these TPD cycling experiments, isothermal cycling experiments were also performed at 450 ºC. The first step consisted of a TPD scan made at 6 ºC min−1 and subsequent desorption and absorption steps were all done at 450 ºC. Figure 4.34 shows the isothermal hydrogen desorption /absorption cycling in the 4NaBH4 / 5Mg2NiH4 composite. For these measurements larger reservoir volumes (167.7 and 1056.3 ml) were used to fully desorb the hydrogen content of the composite. The first two desorption measurements show capacity of 4-5 wt% H. However, after a few cycles, capacity drops to 3 wt% H (similar to data shown in Figure 4.33). Table 4.8 shows the volumes and pressures used for the isothermal cycling experiments. For test a the pressure was maintained at about 1-2 bar during desorption and 84 bar for absorption. This allowed for complete decomposition of NaBH4 to NaH as the reservoir 94 pressure was kept below the equilibrium desorption pressure of NaBH4, but above that of NaH. An absorption pressure of 84 bar allowed for complete absorption of the 4NaBH4 / 5Mg2NiH4 composite. For test b, a pressure of about 4.5 bar was maintained throughout the desorptions. The desorption of hydrogen from NaBH4 was limited under this hydrogen back pressure to plateau II as shown in Figure 4.20. Therefore, only partial decomposition of NaBH4 to NaH was attained. Absorption at 84 bar allowed for complete absorption. The hydrogen pressure for absorption for test c was lowered to 57 bar, which is below the equilibrium absorption pressure for Mg2NiH4 (about 67 bar). Accordingly, only the NaBH4 phase of the composite was hydrogenated. The hydrogen back pressure for desorp- tion was maintained at about 3.5 bar. The partial decomposition of NaBH4 similar to test b was observed. Table 4.8: Parameters for the isothermal cycling studies Test 1st TPD Desorp. Isothermal Desorp. Abs. Abs. pressure Desorp. pressure Volume Pressure Volume [bar] [bar] [ml] [bar] [ml] a 1 1-2 1056.3 84 167.7 b 1 4.5 167.7 84 167.7 c 1 3.5 167.7 57 11.7 Figure 4.34 shows the gravimetric capacity as a function of cycle number for three different sets of absorption and desorption settings. The first curve (a) shows a gradual loss of capacity from about 4.75 to about 2.5 wt% H, which is consistent with the TPD cycling experiments. Curve (b) is similar to curve (a), however the loss in capacity seems to occur much sooner. This is due to the limited desorption of NaBH4 in experiment b. This implies that there is a gradual loss in capacity due to the decomposition of NaBH4 after plateau II. Curve (c) shows an immediate loss of capacity to about 1 wt% H. The capacity is constant over four cycles and it is representative of the partial decomposition of NaBH4 95 as shown in Figure 4.20 from plateau I to Plateau II. This indicates that there is no loss in capacity from this step. Figure 4.34: Capacity as a function of the number of cycles for isothermal cycling. 4.8 Hydrogen Desorption Kinetics of 4NaBH4 / 5Mg2NiH4 The hydrogen desorption kinetics of the 4NaBH4 / 5Mg2NiH4 system were examined using differential scanning calorimetry and the Kissinger method as described in section 3.3.5 and section 2.3.1. For these measurements the sample is heated at a constant heating rate (β = 4, 6, 8, 10 ºC min−1) and the heat flow in and out of the sample is recorded. Figure 4.35 shows the DSC scans for different heating rates β . The first major endotherm (I), associated with the decomposition of Mg2NiH4, showed a systematic shift to higher temperatures when the heating rate β was increased. The second endotherm (II), associated with the decomposition of NaBH4, however, decreases with the first two increases in β and then 96 increases significantly when β is increased to 10 ºC min−1. This makes the application of the Kissinger method difficult. The scatter in endotherm II’s position may be a consequence of experimental conditions such as: reaction of the sample with the pan, the presence of residual air in the instrument, or sample loss during the heating. Sample loss during the heating may be a consequence of the sample foaming when it is heated beyond about 425 ºC. Figure 4.35: DSC scan of the 4NaBH4 / 5Mg2NiH2 composite for different heating rates (β ). The first endotherms’ (I) peak positions increases with heating rate. The second en- dotherms (II) do not show a similar behavior. Figure 4.36 shows the Kissinger analysis for the 4NaBH4 / 5Mg2NiH4 mixture for the first endotherm. From the slope of the data an activation energy Ea = 131 ± 24 kJ mol−1 is obtained with an R2 value of the linear fit of about 0.988, indicating a good fit. 97 However, when compared to the activation energy obtained from the ball milled Mg2NiH4 in section 4.2 of 73 ± 9 kJ mol−1 this is a substantial increase. The Mg2NiH4 activation energy as reported by Selvam et al. [40] is 183 kJ mol−1. They prepared their Mg2NiH4 through the mixing of stoichiometric amounts of Ni and Mg which was then melted and powdered several times under argon. The resulting Mg2Ni powder was then hydrogenated. The particle and grain size of their Mg2NiH4 particles would likely be larger than those created in our lab and could account for the discrepancy in activation energy. The discrepancy in activation energy of hydrogen desorption has been reported in other systems such as MgH2, NaAlH4 [40, 2], and is most likely related to the surface proper- ties of the system including surface area, grain size, oxidation, presence of catalysts and activation procedures [2, 20]. Varin et al. have reported activation energies for MgH2 that range from 118 - 160 kJ mol−1 depending on the manufacturer of the powder, whether it was activated or not, whether it was ball milled or not, etc. They also reported that samples exposed to air had activation energies of 230 - 300 kJ mol−1, which they reported were high because of moisture and air reacting with the Mg particles to form Mg(OH)2 and MgO shells around the particles [2]. 98 Figure 4.36: Fitted curve for Kissinger analysis of the first DSC endotherm (I). The activation energy determined for the Mg2NiH4 component of the 4NaBH4 / 5Mg2NiH4 mixture indicate that the hydrogen desorption kinetics of Mg2NiH4 are more sluggish than the kinetics of Mg2NiH4 alone. The reasons for the difference are unclear as the Mg2NiH4 in the composite seems to decompose alone in the 4NaBH4 / 5Mg2NiH4 mixture. It could be that any free Ni that may have existed in the Mg2NiH4 after synthesis and may have acted as a catalyst was now spread out over both compounds and thus not as concentrated and available to the Mg2NiH4. The theoretical thermodynamic temperature of desorption for the NaBH4 component in the 4NaBH4 / 5Mg2NiH4 composite is 311 ºC, however the observed temperature of desorption is 360 ºC. The difference is perhaps due to poor kinetics. Milanese et al. reported an improvement in kinetics from the addition of 3 wt% MgF2 to the NaBH4 / MgH2 system 99 by decreasing the onset temperature from 440 to 410 ºC for the NaBH4 component [31]. Bosenberg et al. reported that the addition of titanium isopropoxide to the LiBH4 / MgH2 system improved the kinetics by decreasing the onset of desorption of the MgH2 component of the composite from 375 to 355 ºC [28]. The addition of catalysts such as MgF2 to the 4NaBH4 / 5Mg2NiH4 composite may improve the kinetics and improve the temperature of desorption so that it is closer to that calculated thermodynamically. 4.9 Comparison of Different Complex Metal Hydride Composites Figure 4.37 shows the TPD analysis of the compounds Mg2NiH4, MgH2, NaBH4, LiBH4 and their combinations. Hydrogen desorption in NaBH4 occurs above 500 ºC. Addition of MgH2, however reduces the onset of desorption to about 460 ºC. Thermodynamic desta- bilization has occurred. However, the kinetics were too sluggish during testing and the full curve is not shown. Addition of Mg2NiH4 to NaBH4 has a much more pronounced effect on destabilizing the NaBH4, consequently the onset of desorption is now about 360 ºC. There is no concerted reaction between the Mg2NiH4 and NaBH4 during the first step of the dehydrogenation. In the 4LiBH4 / 5Mg2NiH4 system, however, Vajo et al. reported that hydrogen release commences at 250 ºC via a concerted reaction between Mg2NiH4 and LiBH4 [4]. 100 Figure 4.37: Comparison of LiBH4 and NaBH4 based composites and the destabilizing effect of MgH2 and Mg2NiH4. Table 4.9 shows a comparison of hydrogen storage properties for the systems shown in Figure 4.37. 101 Table 4.9: Comparison of several hydride storage systems (under 1 bar hydrogen pressure) System ∆Hdes Tdes (observed) H weight [kJ molH2−1] [ºC] [%] 4NaBH4/5Mg2NiH4 75 360 4.5 Mg2NiH4 65 280 3.6 4LiBH4/5Mg2NiH4 15 250 2.5 2NaBH4/MgH2 62-100.58 [35, 31, 34] 450 9.88 NaBH4 106.8 577 10.8 4.10 Summary Monoclininc Mg2NiH4 was synthesized and characterized. The hydrogen desorption and absorption enthalpy were 74 ± 5 and 60 ± 3 kJ mol−1H2 , respectively. The activation energy of the hydrogen desorption reaction was about 73 ± 9 kJ mol−1. Heating Mg2NiH4 to 300 ºC and 1 bar led to formation of Mg2Ni and hydrogen release. Mg2Ni can be reversibly hydrogenated to form Mg2NiH4. Crystallographic phase analysis shows that 4NaBH4 / 5Mg2NiH4 decomposes in at least three steps with the first step being the decomposition of Mg2NiH4 starting at around 275 ºC, the second step the decomposition of NaBH4 to form NaH (at about 360 ºC), and the third step the decomposition of NaH to Na at around 425 - 450 ºC. The MgNi2.5B2 com- pound forms by 400 ºC. The thermodynamics of 4NaBH4 / 5Mg2NiH4 were studied using a manometric method and the enthalpy of decomposition was determined to be 67 ± 4, 76 ± 5 and 95 ± 7 kJ mol−1H2 for the decomposition of Mg2NiH4, NaBH4 and NaH respectively. This compares to a decomposition enthalpy of 66, 90 and 96-114 kJ mol−1H2 for Mg2NiH4, NaBH4, and NaH, respectively. The absorption enthalpy of 59 ± 4 kJ mol−1H2 corresponding to the Mg2NiH4 phase of the composite was obtained. The onset temperature of hydrogen desorption for NaBH4, in the composite, was reduced to about 360 ºC from 500 ºC for pure NaBH4 and 102 450 ºC for 2NaBH4 / MgH2 indicating thermodynamic destabilization. 4NaBH4 / 5Mg2NiH4 was tested under a hydrogen back pressure of vacuum, 1 and 5 bar. Formation of the ternary boride MgNi2.5B2 did not show any dependance on hydrogen back pressures. The 4NaBH4 / 5Mg2NiH4 composite was ball milled for 10, 30, 60 and 120 minutes. No effect on the desorption kinetics or thermodynamics was found due to ball milling. A small capacity loss of about 0.15 wt% was noticed for ball milling up to 120 minutes. Grain size of the composite 4NaBH4 / 5Mg2NiH4 was affected slightly due to ball milling with a reduction from about 90 nm to about 40 nm for the NaBH4, and 55 nm to 20 nm for the Mg2NiH4. Both of these reductions in grain size occurred within 30 minutes of ball milling. Particle size analysis showed little change in particle size for ball milling up to 120 minutes. The desorption kinetics may not be affected due to the small change in particle size. Hydrogen absorption and desorption into the 4NaBH4 / 5Mg2NiH4 composite was stud- ied up to 5 cycles. A loss in capacity of about 0.7 wt% hydrogen was observed due to the formation of the non-hydriding phase MgNi2 and the loss of Na outside of the sample holder. Isothermal cycling tests were conducted under back pressures of 1 to 4.5 bar and absorption pressures of 57 and 84 bar. Hydrogen capacity during these cycles, reduced from about 4.5 to 2.5 wt% H for experiments rehydrogenated under 84 bar hydrogen, and 4.5 to 1 wt% H for experiments rehydrogenated under 57 bar hydrogen pressure. Kinetic analysis of 4NaBH4 / 5Mg2NiH4 was undertaken using the Kissinger method and an activation energy of 131 ± 24 kJ mol−1 was obtained for the decomposition of the Mg2NiH4 composite. Determination of the desorption activation energy for the NaBH4 component in the composite was not possible with DSC. The activation energy of the NaBH4 component in the composite may be determined by the Arrhenius method instead of the Kissinger method. The effect of hydrogen absorption / desorption cycling on the phases formed was stud- 103 ied. The loss in hydrogen capacity of the Mg2NiH4 phase during cycling resulted from: a) the formation of MgNi2 (non-hydriding binary alloy), b) the formation of non-hydrogenable Mg and, c) the loss of Na from the sample holder. The TPD analysis of compounds like Mg2NiH4, MgH2, NaBH4, LiBH4 and their com- binations were compared. The 4NaBH4 / 5Mg2NiH4 composite had an observed dehydro- genation of the NaBH4 component starting at 360 ºC, the NaBH4 component of the NaBH4 / MgH2 composite started desorbing hydrogen at 450 ºC. In the 4LiBH4 / 5Mg2NiH4 com- posite the concerted reaction between the two components of the system started desorbing hydrogen at 250 ºC. The individual complex hydrides NaBH4 and Mg2NiH4 showed hy- drogen desorption at 577 and 280 ºC respectively. 104 Chapter 5 Conclusions A novel complex hydride composite 4NaBH4 / 5Mg2NiH4 was successfully prepared, and examined as a potential hydrogen storage material. Single phase Mg2NiH4 was synthesized by ball milling and solid state reactions at 495 ºC and under 95 bar hydrogen pressure. Mix- tures of 4NaBH4 / 5Mg2NiH4 were then prepared by ball milling, and hydrogen absorption / desorption experiments were performed. The addition of Mg2NiH4 led to a successful thermodynamic destabilization of NaBH4 through formation of MgNi2.5B2. This destabi- lization reduced the hydrogen desorption temperature of NaBH4 from 577 to 360 ºC under 1 bar hydrogen. The initial desorption capacity of this composite is about 5 wt%. Hydrogen desorption from 4NaBH4 / 5Mg2NiH4 was determined to occur in three steps: Reaction 1 4NaBH4 +5Mg2NiH4⇋ ∆H = 67 ± 4 kJ mol−1H24NaBH4 +5Mg2Ni+10H2 Reaction 2 4NaBH4 +5Mg2Ni⇋ ∆H = 76 ± 5 kJ mol−1H24NaH+2MgNi2.5B2 +8Mg+6H2 Reaction 3 NaH⇋ Na+H ∆H = 95 ± 7 kJ mol−1H2 The enthalpies of hydrogen desorption for NaBH4 and NaH are reported to be 106.8 and 114 kJ mol−1H2 respectively. Thus the addition of Mg2NiH4 has led to the reduction of the enthalpies of NaBH4 and NaH. Formation of the stable ternary boride MgNi2.5B2 leads to this destabilization. The enthalpy of absorption corresponding to Mg2NiH4 in the composite was 59 ± 4 kJ mol−1H2 . The impact of ball milling time, desorption back pressure, and hydrogen absorption / 105 desorption cycling was also studied. The following conclusions could be drawn: • Ball milling the 4NaBH4 / 5Mg2NiH4 for up to 2 hours had little impact on the crystallinity and the hydrogen desorption / absorption properties of this mixture. • MgNi2.5B2, was formed under 0, 1, and 5 bar of hydrogen back pressure during the desorption step. This is in contrast to 4LiBH4 / 5Mg2NiH4 where a back pressure of at least 3 bar is required. • The 4NaBH4 / 5Mg2NiH4 composite was cycled a few times to assess the repeated hydrogen release / uptake. The initial capacity of about 4.8 wt% H was reduced to about 2.5 wt% H after 4 cycles. The loss in capacity is attributed to the formation of MgNi2 and loss of Na. • Measurements of the activation energy of hydrogen desorption from Mg2NiH4 alone and mixed with NaBH4 yielded 73 ± 9 and 131 ± 24 kJ mol−1 respectively. The use of NaBH4 (with more abundant Na) to replace LiBH4 in the complex metal hydride composite looks promising and generally provided better performance. 5.1 Future Work • Formation of other possible ternary boride phases can be investigated. For these ex- periments, binary magnesium hydrides such as Mg2FeH6 or Mg2CoH5 and complex metal hydrides like NaBH4 or LiBH4 can be used. There are no reports of mixtures of these composites. • Mixtures of alanates, such as NaAlH4, and binary hydrides like Mg2NiH4 can also be studied. Here Al would form a ternary phase with the Mg and Fe or Co. One of the more attractive reasons for using aluminum is its relative abundance and low cost 106 relative to boron. Mg2AlNi2 alloys are known to form and have been reported by Parente et al. [46]. • To improve the kinetics of hydrogen desorption and absorption the addition of cat- alysts can be investigated. Catalysts such as TiF3 and TiCl3 have been shown to improve hydrogen desorption and absorption rates in NaAlH4 and LiBH4, as well as the composite system NaBH4 / MgH2 [23, 35]. • Preliminary tests have shown that ball milling the hydride system for up to 24 hours significantly reduces the hydrogen desorption temperature, with a concerted desorp- tion starting at around 100ºC. The necessity for long term ball milling of NaBH4 may be due to the hardness of NaBH4. Studies of extended ball milling times up to 24 hours or more are thus recommended. Ball milling NaBH4 alone for an extended time and then mixing with Mg2NiH4 is also suggested because the composite may release hydrogen during milling. This may occur because of the large energy input from the ball milling, which may cause the composite to reach the desorption tem- perature over extended milling times. This will be a concern if the temperature of desorption for the composite is lowered enough to bring it into the temperature range of the ball milling jar during milling. • Because of the difficulty in determining the phase transitions and reaction pathways of the hydride system during decomposition and absorption, in-situ thermal XRD would be very useful. During in-situ XRD diffraction, patterns of the sample are collected while the temperature of the sample is changed. With ex-situ XRD, there is always a concern that when a sample is cooled, while hydrogen is not fully desorbed, some absorption may occur. See Appendix A.3.3. 107 Bibliography [1] A Ter-Gazarian. Energy storage for power systems. P. Peregrinus-Institution of elec- trical engineers, Herts (GB), 1994. [2] R.A. Varin, T. Czujko, and Z.S. Wronski. Nanomaterials for Solid State Hydrogen Storage. Fuel Cells and Hydrogen Energy. Springer, 2008. [3] Professional Engineering Publishers. 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Pulz, and Z. S. Wronski. Synthesis of nanocom- posite hydrides for solid-state hydrogen storage by controlled mechanical milling techniques. Journal of Alloys and Compounds, 483(1-2), 2009. 112 [43] M. D. Banus, J. J. McSharry, and E. A. Sullivan. The sodium-sodium hydride- hydrogen system at 500-600-degrees. Journal of the American Chemical Society, 77(7), 1955. [44] R. A. Varin and C. Chiu. Structural stability of sodium borohydride (NaBH4) during controlled mechanical milling. Journal of Alloys and Compounds, 397(1-2), 2005. [45] Karl J Gross, K. Russel Carrington, Steven Barcelo, Abhi Karkamkar, Justin Purewal, and Philip Parilla. Recommended Best Practices for the Characterization of Storage Properties of Hydrogen Storage Materials. Technical report, V2-64 U.S. D.O.E. Hy- drogen Program, 2009. [46] A. Parente, A. Nale, M. Catti, E. Kopnin, and P. Caracino. Hydrogenation properties of Mg2AlNi2 and mechanical alloying in the Mg-Al-Ni system. Journal of Alloys and Compounds, 477(1-2), 2009. [47] C. Suryanarayana. Mechanical alloying and milling. Progress in Materials Science, 46(1-2), 2001. [48] F. D. Maslan and T. M. Littman. Compressibility chart for hydrogen and inert gases. Industrial and Engineering Chemistry, 45(7), 1953. [49] B.D. Cullity and S.R. Stock. Elements of x-ray diffraction. Pearson education. Prentice Hall, 2001. [50] O. Palchik, A. Gedanken, and Y. Rubinstein. Simple air-tight powder x-ray diffraction cell. Review of Scientific Instruments, 74(6), 2003. 113 Appendix A Experimental Techniques A.1 Ball Milling for Materials Preparation Ball milling has traditionally been used for mixing, changing the particle shapes, and size reduction. Since the 1960s ball milling has been used as a method for alloying metals and creating nano-materials. A very common type of high energy ball mill is a shaker or vibrational mill popularized by its manufacturer SPEX®. This milling method combines a back and forth motion with a lateral movement of the ends of the vial so that it moves in a figure eight [47]. A schematic of the jar’s motion and that of the ball in a shaker / vibrational mill is shown in Figure A.1. Figure A.1: Schematic of high energy ball milling process The SPEX 8000® is classified as a high energy ball mill due to the amount of energy transferred from the balls to the powder being milled inside. The kinetic energy of the balls 114 in the system is given by equation A.1. E = 1/2mv2 (A.1) where m is the mass of the balls and v is their velocity. With ball speeds on the order of 5 m s−1 and a ball mass typically 8 g, a high energy ball mill will have about 20 mJ of energy per ball [47]. Most of this energy is dissipated as heat, with less than 1 % of the energy transmitted through impact [47]; however, some amount of shear is realized as the balls move along the walls of the jar. The jars used for milling come in a variety of materials. Stainless steel is among the most common material. The size of the balls used for milling as well as the ball to powder mass ratio are impor- tant. Typically a mass ratio of 10:1 ball:powder is used for the preparation of metal hydrides, but ratios from 5:1 up to 50:1 have been reported in the literature. A typical SPEX milling is performed for 1 to 5 hours, but times have been reported to be in the hundreds of hours [2]. Mechanical alloying is achieved through the milling of more than one type of elemental metal powder. The method is a non-equilibrium, low temperature solid-state processing method. Mechanical alloying is basically a repeated stressing, deformation, fracture, and cold welding of powder particles, which can be seen schematically in Figure A.2. 115 Figure A.2: Mechanical alloying steps [47] [©Progress in Materials Science, 2001, by permission] Particle size reduction and grain size reduction occur in ball milling, which improves the ability of hydrogen to diffuse into or out of the metal powder particles. Figure A.3 shows the effect of milling time on particle size and grain size. In general a decrease in particle and grain size occurs, however a plateau occurs where the particle and grain size no longer become any smaller, this is because an equilibrium is reached between the fracturing of particle into smaller sizes and the welding of particles into larger sizes [47]. 116 Figure A.3: Particle and grain size reduction as a function of milling time and ball to powder ratio ([47]) [©Progress in Materials Science, 2001, by permission] For metal hydrides, the oxide shell that surrounds a lot of metals is broken up by the milling process enabling fresh surfaces to be exposed to the hydrogen. With a decrease in particle size there is an increase in specific surface area of the powder. The extra surface area that is not oxidized leads to significant improvements in the kinetics of the hydrogen desorption. A.2 Pressure-Composition-Temperature (PCT) A PCT (or Sievert’s apparatus) is a manometric method for studies of gas absorption / desorption into solids and liquids. It consists of a temperature controlled autoclave which houses the sample reactor, gas manifold and valves, a vacuum pump, pressure transducers, gas (H2, He) supplies, and reservoir volumes. Figure A.4 shows the schematic of a PCT 117 system. Figure A.4: Simplified schematic of a Sievert apparatus The sample, which is usually in powder form, is placed into the temperature controlled reactor. The hydrogen pressure in the reservoirs is controlled by an electronic pressure regulator. The basis of PCT measurements is the ideal gas law: PV = nRT (A.2) where P = pressure, V = volume, n = moles of gas, R = gas constant (0.082 L atm K−1 mol−1) and T = temperature. The ideal gas law assumption is valid here because the com- pressibility factor for hydrogen for the pressures used here (<200 bar) is negligible [2]. Masla et. al. [48] show that the compressibility of hydrogen between 1 and 150 bar and temperatures between 20 and 400 ºC is between 1 and 1.10. During a hydrogen absorption process gas pressure drops. This change in pressure (∆P) 118 can be used to determine the quantity of hydrogen that is absorbed. The converse is true for desorption, where a rise in pressure signifies the release of hydrogen. Equation A.3 shows the relationship between a change in pressure and the change in the number of moles associated with it. ∆nabs/des = ∆Pdec/inc V RT (A.3) The weight percent (wt%) of hydrogen released can then be calculated from equation A.4. wt% = 2MH2△nabs/des msample (A.4) where MH2= molar mass of hydrogen (1.008 g/mol), △nabs/des is the moles of hydrogen ab- sorbed or desorbed and msample is the mass of the sample in the reactor. A common and quick experiment is to heat the sample temperature from room tem- perature to a desired level, at a constant heating rate. This is referred to as a temperature programmed desorption or TPD. The rise in the hydrogen pressure provides a measure of the amount of hydrogen stored in the materials. Different metal hydride systems can also be compared by the temperature at which hydrogen starts to release. A.2.1 Van’t Hoff Analysis The enthalpy of desorption / absorption is related to the equilibrium pressure at which the hydrogen is desorbed / absorbed and the relationship can be described by the van’t Hoff equation (equation 2.6). The equilibrium pressure for a fixed temperature can be obtained from a pressure-composition-temperature (PCT) plot, sometimes called isothermal pressure-concentration curve (seen in Figure A.5a). 119 Figure A.5: PCT measurement for desorption Figure A.5a shows how single step kinetic hydrogen desorption reactions are used to produce an absorption PCT isotherm. For these measurements the sample is first heated to the desired temperature and pressurized. A small dose of hydrogen is then subtracted from the sample (this can be observed in Figure A.5a by the several sharp peaks) and allowed to desorb until it reaches equilibrium; this is the point where the hydrogen gas and the hydrogen in the solid are in equilibrium. This process of subtracting a known amount of hydrogen to the sample and reaching equilibrium is repeated several times until the metal hydride no longer has any hydrogen. At this point the pressure in the sample starts to fall with each subtraction of hydrogen(Figure A.5a). Figure A.5b shows the PCT curve derived from the kinetic hydrogen absorption graph shown in Figure A.5a. The equilibrium pressure for each dose of hydrogen is plotted ver- sus the amount of hydrogen added for that dose. The result is a PCT curve showing the equilibrium absorption pressure of the metal hydride for the temperature examined. 120 A series of isothermal pressure-concentration (or PCT) curves at different temperatures allows the computation of the desorption / absorption enthalpy by using the van’t Hoff relationship. Figures 2.6a and 2.6b in Chapter 1 show this. A.2.2 Sample Volume Calibration In the PCTPro 2000 apparatus, there is a temperature gradient from the sample holder to the reservoir. This implies that the gas temperature is not uniform. The temperature gradient makes it difficult to calculate the change in moles of gas due to absorption and desorption using the ideal gas law [45]. Fortunately, a correction to the sample’s dead volume would be sufficient to correct for this non-uniformity. The dead volume is the extra space surrounding the sample in the heated reactor up to the valve. It is usually minimized through the use of spacer bars placed in the reactor, but is not completely eliminated. The dead volume is schematically shown in Figure A.6, where the space around the sample and spacer is occupied by a small amount of gas only. 121 Figure A.6: Schematic of temperature gradient in hydrogen above the loaded sample. The temperature gradient starts at 300 ºC in the heater and ends at 25 ºC in the reservoir. The sample dead volume and spacer can be seen directly above the sample. Higher temperature leads to gas expansion and makes the dead volume appear smaller. A simple calibration procedure is then used to determine the apparent volume at the tem- perature of interest. The calibration procedure involves determining an apparent sample dead volume to compensate for the temperature gradient. Figure A.7 shows a simplified Sieverts apparatus schematic. 122 Vacuum Inert Gas ReactorReservoir Volume Valve 1 Valve 3 Valve 2 Figure A.7: Simplified Sieverts schematic for volume calibration Volume Calibration Steps There are two steps for calibrating the reactor volume (using Figure A.7). 1. Evacuate the reactor volume and system volume and close the valve between them (Valve 1). 2. Introduce inert gas (argon or helium) into the reservoir volume to a prescribed pres- sure P1. 3. Close all valves to system volume (Valves 2 and 3). Open valve to the reactor and record the change in pressure. 1. Close the valve to the reactor with the inert gas still in it at pressure P2 and evacuate the system volume. 2. The reactor valve (Valve 1) is then opened and the pressure (P3) is noted. Here system volume (VS) refers to the volume of the piping connecting the various com- ponents in the Sieverts apparatus between the valves as well as the reservoir volume (V) shown in Figure A.7. Care should be taken to ensure that all the components are at thermal 123 equilibrium when the pressure readings are made. Equation A.5 shows the relationship be- tween pressure change, the volume being calibrated (VR) and the system volume (VS which is known). P1VR = P2 (VR +VS) (A.5) Rearranging equation A.5 for VR leads to equation A.6. VR = (P2VS) (P1−P2) (A.6) Because of the difference in temperature between the reactor and the rest of the system, calibration must be repeated for each temperature that an absorption or desorption is done. This calibration procedure can only be performed for absorption as hydrogen will evolve from the sample if the pressures used are below the equilibrium pressure of the hydride. A typical Sieverts system may have several calibrated volumes of different sizes in order to be able to handle various amounts of hydrogen being desorbed or absorbed. This relates to the equilibrium pressure of the hydride sample, as for a given temperature the hydride will absorb or desorb hydrogen until it reaches the equilibrium temperature, at which point it will no longer absorb or desorb hydrogen. This pressure increases with increases in temperature. A.2.3 Volume Selection Care should be taken when selecting the volume of the system during an experiment because the rate of absorption or desorption will vary with the proximity to the equilibrium pressure. For desorption, if the volume chosen is too small, it will only release hydrogen until it reaches its equilibrium pressure. At this point the rate of hydrogen release will decrease as the driving force is reduced. Eventually hydrogen release will stop at the equilibrium 124 pressure where an equal amount of hydrogen is being desorbed as is being absorbed. If the volume chosen is too large then the change in pressure from desorption will be too small for the pressure sensor to resolve accurately and the resolution of the data will suffer. A similar effect can be seen for absorption. If the volume chosen is too small for the experiment the pressure drops to the equilibrium pressure and absorption will stop, even if the hydride holds more hydrogen. If the volume is too high the pressure drop measured will be small and will suffer from poor data resolution in a similar manner to desorption. Proper volume selection is required such that the largest constant driving force is ob- tained by choosing as large a volume as possible, but the resolution in data is maximized by choosing as small a volume as possible. These are conflicting requirements and both need to be satisfied as much as possible. The concept of driving force can be clearly seen in Figure A.8 for desorption where the farther away from the equilibrium pressure (solid black curve) the larger the driving force for hydrogen release there is. 125 Figure A.8: Illustration of driving force for absorption where the equilibrium pressure is the solid black curve and two absorption pressures are applied: a very small driving force and a large driving force. A.2.4 Desorption Pressure The initial pressure at which desorption was performed varied from 1 bar to about 5 bar. Figure A.9 shows an apparatus with a sample reactor (VR) and reservoir volume (VS) to illustrate the procedure. 126 Figure A.9: A simplified schematic diagram of PCT equipment Usually the sample reactor (VR) is pressurized to stop the hydrogen from desorbing at temperatures above the equilibrium temperature for isothermal measurements. The calcula- tion to determine the pressure required in the sample reactor (PR) before opening the valve is shown in equation A.7. PR = PT (VS +VR) VS (A.7) where PS is the sample reactor pressure and PT is the pressure in the total system after opening the valve. This calculation is done after the system volume calibration has been completed. Equation A.7 is valid only if the reservoir volume (VS) is evacuated before opening the valve. A.2.5 Error due to Temperature Fluctuations During the Manometric Measurements The temperature of the reservoir in the PCT Pro 2000 is controlled through a heater and is set to a constant temperature about 5 ºC above ambient temperature 22 - 27 ºC. The temperature fluctuations are within ±0.5 ºC. Figure A.10 shows the reservoir temperature variations over 2.5 days. The temperature varies from 25.8 to 27 ºC over this period. The temperature fluctuations directly impact the pressure measurements. Using the ideal gas law, a±0.5 ºC fluctuation corresponds to about±0.03 wt % error in the hydrogen capacities determined. 127 Figure A.10: PCT Pro 2000 reservoir temperature over about 2.5 days During a TPD measurement, when the sample is heated, pressure increases because of both the gas expansion and by any hydrogen which is desorbed during the heating. For the different reservoir sizes a temperature ramp of 6 ºC min−1 under 1 bar of hydrogen back pressure was completed to measure the gas expansion part of the pressure increase. Figure A.11 shows this apparent hydrogen desorption effect. During the temperature ramp from 25 to 450 ºC, the apparent desorption is about 0.08 wt %, assuming a sample size of 0.2 g, which is typically used in the experiments reported in this thesis. When the temperature is maintained a 450 ºC for 5 hours, gas continues to expand, and the apparent desorption reaches 0.16 wt %. 128 Figure A.11: Gas expansion during a TPD measurement of an empty sample holder. The rise in pressure can be interpreted as hydrogen release. Data are shown for 3 different reservoir volumes. A.3 X-Ray Diffraction (XRD) Through XRD powder analysis, bulk phase properties such as crystallographic phases, grain size and lattice spacings can be determined. Crystallographic phases can be obtained by us- ing a database of known powder diffraction files. Grain size is calculated from the Scherrer equation and lattice spacings are determined from Bragg’s law [49]. For this thesis a Bruker D8 Advance with a 0.6 mm slit, a Vantec-1 X-ray detector and a Cu Kα source for X-rays were used. The analysis software was EVA version 14 with powder diffraction file (PDF) database version 9. 129 A.3.1 Crystallographic Phases The crystallographic phase can be determined through XRD powder analysis. X-rays di- rected at a sample will reflect back at the same angle that they were hit. Because X-rays are electromagnetic radiation waves, their intensity can be attenuated and accentuated by the addition of other X-rays waves with different phase angles. This is called constructive and destructive interference respectively. Constructive interference is governed by Bragg’s law of diffraction (equation A.8), where λ is the wavelength of the X-rays, d is the interplanar distance between atoms in a crystal and θ is the angle of incidence. λ = 2dsinθ (A.8) Powder diffraction X-ray analysis takes advantage of the multiple orientations of crys- tallites (grains) to provide information about the crystallographic data of the materials. Fig- ure A.12 shows several separate particles with several crystallites in each of the particles. Each crystallite has a single crystal orientation, with each particle having several crystallites and several orientations. The crystallites are separated by a grain boundary. 130 Figure A.12: Multiple crystallites (grains) and particles. The XRD pattern is unique for every material with a crystal structure so that the iden- tification of an unknown sample, through the use of a database of known patterns, can be obtained. A.3.2 Grain and Particle Size The grain size refers to the size of the individual crystallites that compose a larger particle. The size of these grains can be determined by the Scherrer equation shown in equation A.9. dg = 0.9λ βFWHMcosθ (A.9) where dg is the grain size, λ is the wavelength of radiation used, βFWHM is the full width at half the maximum peak height (FWHM) used to determine the grain size, and 2θ is the diffraction angle of the peak. The multiplier can vary from 0.9 to 1 and depends on the shape of the crystallites [49]. Typically only grain sizes less than 100 nanometers can be 131 determined accurately. A.3.3 Effect of Hydrogen Back Pressure During Sample Cooling The hydrogen pressure in the sample holder could affect the composition of the hydride composite during cool down. All the samples studied here in this thesis were evacuated before cooling, otherwise some hydrogen absorption would have occurred during the cool- ing. This is shown in Figure A.13. Two samples were cooled under different hydrogen back pressures. For the sample cooled down under 1 bar hydrogen back pressure NaBH4 and Mg2NiH4 phases are observed. As the sample cools down the equilibrium pressure for absorption reached and the absorption of hydrogen can take place. Figure A.13: Effect of hydrogen back pressure during cooling on hydride composition 132 A.4 Differential Scanning Calorimetry (DSC) During a chemical reaction, heat is either released (exothermic) or absorbed (endothermic). Thus formation of new phases or compounds can be detected by measurement of the heat of reaction, i.e., exotherms / endotherms. The phase changes that a hydride undergoes as it releases and absorbs hydrogen can be measured through the use of a differential scanning calorimeter (DSC). Here the sample and an empty reference sample are heated at a constant rate and the difference in heat required to keep the sample at the same temperature as the reference sample is recorded as a function of temperature or time [24]. Figure A.14: DSC scan for the desorption of the MgH2 Figure A.14 shows the calorimetric desorption profile for the MgH2. The desorption scans are shown with downward peaks indicating that the phase changes are endothermic, while for absorption peaks are pointing up thus indicating exothermic reaction. Not all of the peaks represent hydrogen release and absorption. Changes in phase from solid to liquid or from one crystal structure to another are also seen as these changes require energy 133 transfer into or out of the system. 134 Appendix B Estimation of Na Loss From the Sample Holder A simple image processing method was used to estimate the Na loss during the hydrogen desorption reaction of 4NaBH4 / 5Mg2NiH4. Figure B.1 shows the picture of the Na com- pounds deposited on the spacer. Using Irfanview software and the known diameter of the spacer, the diameter and height of the Na-based dome was estimated. This data was used in Solidworks to reconstruct the dome and calculate its volume. The mass of Na was then calculated from the density and the estimated value of volume. Figure B.1: Photo of Na on spacer in reactor 135 Figure B.2: Solidworks rendering of Na on sample holder spacer 136 Appendix C Acoustic Box for Ball Mill A Spex 8000™ ball mill generates considerable noise. In order to keep the noise level in the lab to a minimum an acoustic box was designed and built to house the ball mill. Because the acoustic box is sealed, a ventilation system was incorporated with fans pushing air into the box on one side and out of the box on the other side through a specially designed ducting for noise mitigation. A microcontroller was programmed and used to turn the fans on and off when the acoustic box reached a certain temperature (about 30 ◦C). Figure C.1 shows the rendered drawing of the enclosure made before construction began with the cover on the ventilation system made clear so that the ventilation design can be seen. 137 Figure C.1: Rendered drawing of acoustic box with ventilation system exposed The actual constructed box can be seen in Figure C.2. The acoustic attenuation is so good that the ball milling cannot be heard when it is running. 138 Figure C.2: Acoustic box after completion Figure C.3 shows the electronics installed on the side of the box to turn the ball mill on and off, as well, the microcontroller that controls the fans is here (under the top level printed circuit board (PCB)). 139 Figure C.3: Electronics control box on side of acoustic box 140 Appendix D X-Ray Diffraction Sample Holder Because the samples prepared and studied here were oxygen and moisture sensitive, a spe- cial sample holder was required to handle the powder during x-ray diffraction measure- ments. Commercial air-sensitive sample holders did not fit the XRD system (Bruker D8) used. Thus in-house holders were developed and used. We adopted a design proposed by Palchik et al.[50]. Their design was partially modified by making the assembly shorter thereby eliminating one of the o-rings. We also used Kapton film instead of Mylar film. Figure D.1 shows the XRD scan of an aluminum plate with and without a Kapton film. It can be seen that the Kapton is quite transparent to x-rays and little attenuation of the peaks occur. No additional background peaks are observed. 141 Figure D.1: XRD scan of aluminum with and with out Kapton (deg in 2θ ). Figure D.2 shows a photograph of the XRD samples holders made.There are three part: A - Base, B - Ring, C - Sample pan and D - Screw. The powder sample is placed into the Sample pan (C) and then placed onto the base. The lid is then secured over the base and the screw (D) tightened. 142 Figure D.2: Photograph of XRD airtight sample holder A sample loading tray was also designed to work with the XRD sample holder and is shown in Figure D.3. Figure D.3: Sample loading tray for airtight XRD sample holder The following are the engineering drawings for the XRD sample holders as made for 143 this thesis. Figure D.4: Simple XRD sample holder 144 Figure D.5: Ring part for simple XRD sample holder 145 Figure D.6: Base part for simple XRD sample holder 146 Figure D.7: Sample pan part for simple XRD sample holder 147 Figure D.8: Screw part for simple XRD sample holder 148 Appendix E XRD Analysis Program A program was written to analyze the data collected from the XRD machine. Grain size and peak matching were incorporated into the program. Figure E.1 shows the fitting page of the program. Choices such as which data to show and how much smoothing to apply to the raw file are given. At the bottom the choice to output the all the data to a file for later use is also given. 149 Figure E.1: Fit page for XRD program The XRD fitting program can read an entire scan at once and split it into individual peaks that it then sequentially fits. The fit algorithm in the program uses a pseudo - Voight function which is a function of the form 150 I(2θ) = Ihkl[ηL(2θ −2θ0)+ (1−η)G(2θ −2θ0)] (E.1) where L and G are the Lorentz and Gaussian functions respectively, and η represents how much of each function is used. Ihkl is the intensity and is a scaling factor that multiplies through everything to match the peak height of the fitted data. The peak search algorithm uses a Gaussian smoothing algorithm to smooth the data so that the first derivative can be used to find the start and end of each peak. It looks for a change in the sign of the derivative to determine the peak location and the beginning and end of the peak. The grain size is calculated by using the Scherrer equation shown in Equation E.2. d = 0.9λβFWHMcosθ (E.2) Where d is the grain size, l is the wavelength of radiation used, βFWHM is the FWHM of the peak used to determine the grain size, and θ is the diffraction angle of the peak. The value 0.9 can vary up to 1 and depends on the shape of the crystallites. By fitting the data and obtaining the FWHM which is the full width of the peak at half of the maximum height of the peak (full width half max) an estimate of the grain size can be obtained. However, other factors must be taken into account first, specifically instrumental resolution. Instrumental resolution can be mitigated by obtaining a reference scan from a mono-crystalline material such that the peaks will be as narrow as possible. This will define the instrumental resolution of the machine. Program output is shown in Figure E.2. There is a lot of information on the graph. In solid blue is the original XRD scan as received from the XRD machine, and superimposed just beneath it is a multitude of coloured peaks which are the peaks that the search algorithm has identified. Each peak that is fit is shown as a different coloured series of small crosses 151 and the fit is shown as a solid line just behind it. To the right of the fitted peak is the calculated grain size. A PDF (powder diffraction file) can be used to do peak identification and can be seen here for NaAlH4 by the long vertical blue lines with red diamonds on top of them. 25 30 35 40 45 50 55 60 65 2 Theta 0 10000 20000 30000 40000 50000 60000 70000 80000 90000 In te n si ty  [ a rb . u n it s] 89.44 68.43 84.83 57.71 32.2 39.44 34.86 19.5 27.6134.32 30.6 nan Figure E.2: Graphical output from XRD program showing the original XRD scan (solid blue) overlaid with the fitted peaks with baseline removal (multiple colours). The PDF file is also shown for peak matching along with the grain size is placed to the right side of each peak 152

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