（千葉大学審査学位論文）

Enhancement of the reactivity of magnesiumhydroxide as a chemical heat storage material andinvestigation of the reaction mechanism
水酸化マグネシウム系化学蓄熱材の高性能化とその反応機構に関する検討
2021 年8月

千葉大学大学院融合理工学府 地球環境科学専攻
都市環境システムコース

19WD1305
劉研究室
指導教員

劉

黒沢

諒

醇一

准教授

Table of Contents
Chapter 1. Introduction......................................................................................................................1
1.1. General Introduction ...................................................................................................................1
1.2. Utilizing waste heat ......................................................................................................................2
1.3. Previous studies of heat storage techniques ...............................................................................3
1.3.1. Sensible heat storage .............................................................................................................3
1.3.2. Latent heat storage ...............................................................................................................4
1.3.3. Chemical heat storage...........................................................................................................4
Redox system ...............................................................................................................................5
Carbonates system ......................................................................................................................6
Salt hydrates system ...................................................................................................................7
Metal hydrides system ................................................................................................................8
Hydroxides system ....................................................................................................................10
1.4 Properties of hydroxides for chemical heat storage .................................................................14
1.5. Applications for chemical heat storage materials....................................................................17
1.6. Purpose and structure of this study..........................................................................................18
References ..........................................................................................................................................19
Chapter 2. Effect of LiOH Addition on Dehydration Reaction of Magnesium Hydroxide ........25
ABSTRACT.......................................................................................................................................25
2.1 Introduction .................................................................................................................................25
2.2 Experimental ...............................................................................................................................26
2.2.1 Sample preparation and XRD characterization................................................................26
2.2.2 Dehydration reaction measurements..................................................................................27
2.2.3 Dehydration reaction rate analysis .....................................................................................27
2.3. Results and Discussion...............................................................................................................28
2.3.1 Characterization of Mg(OH)2-W and LiOH/Mg(OH)2 using XRD .................................28
2.3.2. Effect of LiOH addition on the dehydration reaction of Mg(OH)2.................................29
2.3.3. Effect of LiOH addition on the rate of the dehydration reaction ...................................30
2.4 Conclusions ..................................................................................................................................40
References ..........................................................................................................................................40
Chapter 3. Comparison of Effect of Co–addition of Li Compounds and Addition of Single Li
Compound on Reactivity and Structure of Magnesium Hydroxide .............................................42
ABSTRACT.......................................................................................................................................423.1. Introduction ................................................................................................................................423.2. Experimental ..............................................................................................................................43i3.2.1. Sample preparation.............................................................................................................43
3.2.2. Evaluation of reactivity using thermobalance ..................................................................44
3.2.2.1. Dehydration reaction ...................................................................................................44
3.2.2.2. Hydration reaction .......................................................................................................44
3.2.2.3. Evaluation of heat output density ...............................................................................47
3.2.2.4. Sample characterization by XRD ...............................................................................47
3.3. Results and discussion ...............................................................................................................48
3.3.1. Comparison of the effect of addition of single Li compound versus co–addition of the Licompounds on dehydration reactivity of Mg(OH)2....................................................................48
3.3.2. Effect of co-addition of the Li compounds on hydration reactivity of MgO..................49
3.3.3. Sample characterization by XRD analysis for evaluating the effect of addition of the Licompounds on structure of Mg(OH)2 ..........................................................................................55
3.3.4. Comparison of heat storage/output performances between our samples and reported
samples ...........................................................................................................................................60
3.4. Conclusions .................................................................................................................................62
References ..........................................................................................................................................63
Chapter 4. Fourier-Transform Infrared Analysis of the Dehydration Mechanism of Magnesiumhydroxide and Chemically Modified Magnesium Hydroxide .......................................................66
ABSTRACT.......................................................................................................................................66
4.1. Introduction ................................................................................................................................66
4.2. Experimental ..............................................................................................................................69
4.2.1. Sample preparation.............................................................................................................69
4.2.2. Dehydration reaction tests..................................................................................................70
4.2.3. XRD characterization .........................................................................................................70
4.2.4. FT-IR characterization .......................................................................................................70
4.3. Results .........................................................................................................................................71
4.3.1. XRD characterization of the prepared samples ...............................................................71
4.3.2. FT-IR characterization .......................................................................................................72
4.3.2.1. FT-IR characterization of the prepared samples ..........................................................72
4.3.2.2. FT-IR analysis of samples after dehydration .................................................................75
4.4. Discussion ...................................................................................................................................82
4.4.1. FT-IR peak assignment .......................................................................................................82
4.4.2. Effects of LiCl and/or LiOH addition on the dehydration reactivity of Mg(OH)2 ........84
4.5. Conclusions .................................................................................................................................87
References ..........................................................................................................................................87
Chapter 5. Fourier-Transform Infrared and X-ray Diffraction Analyses of the Hydrationii

Reaction of Pure Magnesium Oxide and Chemically Modified Magnesium Oxide ....................93
ABSTRACT.......................................................................................................................................93
5.1. Introduction ................................................................................................................................93
5.2. Experimental ..............................................................................................................................96
5.2.1. Sample preparation.............................................................................................................96
5.2.2. Hydration reaction test conducted using thermobalance ................................................96
5.2.3. Sample characterization using the XRD technique .........................................................97
5.2.4. Sample characterization using the FT-IR technique........................................................97
5.3. Results .........................................................................................................................................98
5.3.1. Sample (as-synthesized) characterization using the XRD and FT-IR techniques .........98
5.3.1.1. XRD patterns recorded for all the as-synthesized samples ......................................98
5.3.1.2. Characterization of the as-synthesized samples using FT-IR technique .................99
5.3.2. Characterization of hydrated samples using the XRD and FT-IR techniques ............ 101
5.3.2.1. XRD patterns recorded for the hydrated samples .................................................. 101
5.3.2.2. Analysis of the FT-IR spectral profiles of the hydrated samples ........................... 105
5.4. Discussion ................................................................................................................................. 115
5.4.1. Effect of the addition of LiCl and/or LiOH on the hydration of MgO ......................... 115
5.4.2. Hydration of pure and chemically-modified MgO: Probable hydration pathway ...... 117
5.5. Conclusions ............................................................................................................................... 120
References ........................................................................................................................................ 121
Chapter 6. Screening of the Effective Additives on the Dehydration and Hydration Reactivity ofMagnesium Hydroxide and Magnesium Oxide System Based on the Properties of Additives 125
ABSTRACT..................................................................................................................................... 125
6.1. Introduction .............................................................................................................................. 125
6.2. Experimental ............................................................................................................................ 126
6.2.1. Sample preparation........................................................................................................... 126
6.2.2. Sample characterization by XRD .................................................................................... 127
6.2.3. Sample characterization by FT-IR .................................................................................. 127
6.2.4. Dehydration reaction test ................................................................................................. 128
6.2.5. Hydration reaction test ..................................................................................................... 128
6.2.6. pH measurement of sample solution ............................................................................... 128
6.3. Results ....................................................................................................................................... 129
6.3.1. Characterization by XRD and FT-IR measurements for as-synthesized, samples ..... 129
6.3.2. Effect of the additives on the dehydration of Mg(OH)2 ................................................. 129
6.3.3. Effect of the additives on the hydration of MgO ............................................................ 133
6.4. Discussion ................................................................................................................................. 133iii

6.4.1. Relationship between the dehydration peak temperature and the properties of theadditives ....................................................................................................................................... 133
6.4.2. Properties related to hydration reactivity of MgO......................................................... 137
6.5. Conclusions ............................................................................................................................... 141
References ........................................................................................................................................ 142
Chapter 7. Conclusions and future issues ..................................................................................... 145
Supporting Information ................................................................................................................. 148
S1. Supporting information for Chapter 3 ................................................................................... 148
S2. Supporting information for Chapter 4 ................................................................................... 152
References ........................................................................................................................................ 162
S3. Supporting information for Chapter 5 ................................................................................... 162
S3.1. Experimental procedure.................................................................................................... 162
S3.1.1. Reactivity evaluation for hydration reaction by using thermobalance .................. 162
S3.2. Results ................................................................................................................................. 164
References ........................................................................................................................................ 173
Appendix .......................................................................................................................................... 175
A1. The Dehydration and Hydration Reactivity of the Samples ................................................ 175
A1-1. Dehydration reactivity of the samples............................................................................. 175
A1-2. Hydration behavior of the samples ................................................................................. 181
A2. SEM images .............................................................................................................................. 185
A2.1. SEM images for the prepared samples ............................................................................ 185
A2.2. SEM images for the reacted samples ............................................................................... 185
A3. DSC measurement.................................................................................................................... 190
A4. CO2 TPD measurement ........................................................................................................... 194
A5. Nitrogen gas adsorption measurement ................................................................................... 196
A6. Dehydration reaction kinetics ................................................................................................. 206
References ........................................................................................................................................ 209
Achievements................................................................................................................................... 211
Acknowledgement ........................................................................................................................... 214

iv

Chapter 1. Introduction
1.1. General Introduction
Nowadays, fossil fuels such as coal, oil, and natural gas have been used as the main energy
resources and the products, electricity, and heat derived from these resources are supplied in our life
(Figures 1-1 and 1-2[1]).

However, global environmental problems such as global warming, anddepletion of energy resources that resulted using too many fossil fuels have been critical issues for not
only present generation but also future generation. In particular, Japan has only few fossil fuels, wehave only 4% of energy self-efficiency rate[2]. Therefore, we need to pay attention to how to use
these materials.

Figure 1-1. Total energy supply in the world in 1990–2016[1]All over the world, renewable energy and sustainable energy technology have been focused
towards environmental problems mentioned above.

Power generation techniques by renewableenergies such as solar, wind, geothermal energy and so on have been researched. These techniques
do not depend on fossil fuels, meaning that every country is able to use those techniques. However,these techniques have some problems: restriction of installation, low efficiency and the gap between
energy supply and demand. On the other hand, the efficiency of nuclear power generation is higher
than the other power generation techniques, and the amount of CO2 emission is quite low. Therefore,
the nuclear power generation will be able to address the issues mentioned above. In Europe, the1power generation ratio of nuclear power is quite high and it also used to be high in Japan[3]. However,since 2011, the nuclear power generation have been shrunk because of Tohoku earthquake in Japan.
Not only in Japan but also in other countries, this movement is being actualized.

Therefore,

alternative techniques are required.
In Sweden, the estimation of economic and environmental costs of replacing nuclear fission withsolar and wind energy was carried out in previous study[4]. This work indicates that it is quite difficult
to replace nuclear energy by solar and wind energy because of the large installation cost for theachievement both environmental and economic goal.
Under such back ground, to increase yield of energy resources, to utilize renewable energy
resources to create electricity, and to enhance efficiency of energy creation are critical forestablishment of the independent energy supply system in Japan. However, as the author mentioned
above, we especially need to focus on utilizing unused energy and renewable energy because we havefew energy resources in Japan.

Figure 1-2. Total final consumption in the world in 1990–2016[1]1.2. Utilizing waste heat
There are a lot of waste heat from industrial department in Japan[5]. Figure 1-3 shows calculated
waste heat from industrial department in Japan. In particular, waste heat under 300 °C accounts forlarge portion, as shown in Figure 1-3. So then, these waste heat can be used as a heat resource. The
portion of heat consumption to total final consumption is 2.96%, that of heat and electricity is 21.7%
2

based on Figure 1-2. These values are not small, therefore, it is essential to use waste heat as aresource for heat and electricity.
In a past decade, utilizing waste heat by chemical heat storage (CHS) materials is a potential
technique to utilize waste heat. The CHS material is a material that can store/release thermal energy
by the chemical reaction. It can address the problems mentioned in the Section 1.1. and the gap
between energy supply and demand especially considered as one of the main issues for solar power
generation.

Figure 1-3. Calculated waste heat from industrial department in Japan[5]1.3. Previous studies of heat storage techniques
The heat storage technique consists of three techniques: sensible, latent, and chemical heat storage
technique.

1.3.1. Sensible heat storage
Sensible heat storage technique is the technique to store thermal energy by temperature rising, torelease it by temperature decreasing.[6–12] For sensible energy, water, thermal oil, liquid metals, earthmaterials, concretes, and molten salts are generally used for storage medium.[12–15]

Sensible heatstorage is quite easy way compared with other techniques. Therefore, the technique is applied forpractical use, for example, in concentrate solar plants.[13–15]

31.3.2. Latent heat storage
Phase change of material, for instance solidification of liquid to solid, is used for latent heat storage
technique, therefore, the material is generally called phase change materials (PCM). The concept of
PCM is that heat storage corresponds to melting of solid to liquid, heat release corresponds tosolidification of liquid to solid.
sensible heat storage.

[6–12]The energy density of latent heat storage is higher than that ofCommon PCMs are classified into organic, inorganic, eutectic, and sugar

alcohols.[14,16] Pure n–alkane, fatty acids, and esters are widely known as organic PCM.[17,18,19] Salt
hydrates is well-known as inorganic PCM,[17,19,20] eutectic PCMs consist of inorganic with organic,organic with inorganic, and organic with organic.[17,19]

The temperature range for these PCM

materials is ~100 °C. Microencapsulated PCM consisting of Al-25wt% Si microspheres as the core(PCM) and Al2O3 as the shell have been studied to utilize 500–600 °C.[21]These two techniques (Sensible and latent heat storage) are easy enough to apply practical use.
However, the energy density of present known organic PCMs is only ~350 kJ kg−1 and that of nonatural sugar alcohol PCMs is up to 450–500 kJ kg−1, which was predicted by computational
method.[16] Moreover, the period of energy storage of these two techniques is limited and easy tolose stored heat.

1.3.3. Chemical heat storage
On the other hand, chemical heat storage has much larger energy density (~0.5–1 kWh kg−1 (1.8–
3.6 GJ kg−1)) than that of the other two methods.[6–12]

Moreover, the period of energy storage

chemical heat is theoretically unlimited.[6–12] Thus, chemical heat storage realizes the long-distancetransport of stored heat and enable us to choose the timing to reuse the stored heat. Therefore, in thisstudy, chemical heat storage (CHS) technique is focused. Table 1-1 shows the properties of the threetechniques.
In the concept of CHS technique, decomposition process (dehydration, decarboxylation, anddeoxygenation, etc.) corresponds to heat storage operation, formation reaction (hydration, carbonation,and oxidation, etc.) corresponds to heat output operation (Eq. (1)). Many CHS materials have beenstudied for reasons mentioned above. The most suitable materials for a usage are different by therange of the application heat temperature. Here, various CHS systems which have been explored are
introduced.
(1)

𝐴𝐴 + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝐵𝐵 + 𝐶𝐶4

Table 1-1. Properties of each technique[6–12]
Heat storage

Sensible heat

Latent heat

Chemical heat

~0.02–0.03 kWh kg−1

~0.05–0.1 kWh kg−1

~0.5–1 kWh kg−1

-material-material-reactants

Maturity

Industrial scale

Pilot scale

Laboratory scale

Storage period

LimitedLimitedUnlimited(Thermal loss)(Thermal loss)(Theoretically)

Small distance

Small distance

Unlimited

technique
Energy density

Transport(Theoretically)
Technology

SimpleSimpleComplex

Disadvantage

Significant heat loss

Significant heat loss

High capital cost

Large volume

Corrosiveness of theTechnically complex

material
Low heat conductivity

Redox system
In the redox system, metal oxides are applied and the basic reaction is shown in below:𝑦𝑦(2)𝑀𝑀𝑥𝑥 𝑂𝑂𝑦𝑦 (𝑠𝑠) + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝑥𝑥𝑥𝑥(𝑠𝑠) + 𝑂𝑂2
2

In this system, cobalt oxide, manganese oxide, barium oxide, and copper oxide have been studied

as CHS materials due to their large energy density.[22–28] However, their decomposition temperature(temperature for heat storage) is quite high, therefore, it is required to lower the temperature. Table
1-2 shows the parameter of redox systems.
Mixed oxides such as perovskites have been studied to enhance the reactivity and overcome thedisadvantage of materials.[6,10,12,29] The second oxide of a mixed oxide is employed to these materials
as a impurity to increase the anion vacancies concentration.[6] The general formula of perovskites is
ABO3, where A is a generally a rare earth or an alkaline cation, while B is generally a transition metal
cation.[6,10,12]Babiniek et al. and Zhang et al. investigated LaxSr1−xCoyMn1−yO3 (LSCM) and

LaxSr1−xCoyFe1−yO3 (LSCF) perovskites with 0.1 ≤ x ≤ 0.9 and 0.1 ≤ y ≤ 0.9, and Ba and Sr series
perovskites (SrMnO3−δ, SrCoO3−δ, SrFeO3−δ, BaCoO3−δ, BaFeO3−δ, BaMnO3−δ, SrCo0.8Fe0.2O3−δ,SrCo0.2Fe0.8O3−δ,Ba0.5Sr0.5CoO3−δ,Ba0.5Sr0.5FeO3−δ,), respectively.[30,31]

These materials include expensive and heavy metals,

Ba0.5Sr0.5Co0.2Fe0.8O3−δ

Ba0.5Sr0.5Co0.8Fe0.2O3−δ,andleading increase of cost and decrease of energy density per unit weight. For these reason, CaMnO3
perovskites have been studied for more cost-effective material.[32]5Table 1-2. Turning temperature (Tturn), reaction enthalpy, and gravimetric energy density of metaloxides[6] Tturn can be determined as Tturn = ΔH/ΔS (ΔH is the enthalpy change, and ΔS is the entropychange).[8]Chemical formulaTemperature [°C]

Reaction enthalpy [kJGravimetric energymol−1]

density [kJ kg−1]

Cr5O12/Cr2O3

105

126.17279.1

Li2O2/Li2O

145

34.225745.9

MgO2/MgO

205

21.432

380.6

PbO2/PbO

295

55.234230.9

Sb2O5/SbO2

325

68.064

210.4

PtO2/PtO

420

62.802

276.5

UO3/U3O8670

35.232

123

BaO2/BaO

880

73.258432.6

Mn2O3/Mn3O4

915

90.038190.1

Co3O4/CoO

935

196.532816.1

Rh2O3/Rh2O

970

249.276

982.1

CuO/Cu2O

1115

64.446810.2

Fe2O3/Fe3O4

1360

232.613

485.6

Mn3O4/MnO

1700

194.63850.6

V2O5/V2O4

1710

176.594970.9

Carbonates system
The basic reaction in this system is as below;(3)

𝑀𝑀𝑀𝑀𝑀𝑀3 (𝑠𝑠) + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝑀𝑀𝑀𝑀(𝑠𝑠) + 𝐶𝐶𝑂𝑂2 (𝑔𝑔)This system can be applied for both carbon dioxide capture and chemical heat storage materials.

CO2 is used as working medium in this system, enabling materials to be used in open system. Metal
carbonates have high reaction enthalpy.
studied,[33–35] showing in Table 1-3.

In this system, CaCO3, MgCO3, and SrCO3 have been

However, some metal oxides; MgO, has only poor reversibility,

meaning that MgO hardly reacts with CO2. To address this disadvantage, Shkatulov et al. studied
eutectic mixture of LiNO3 and KNO3 as an additive to MgCO3.[36] MgCO3/MgO system with doped
the eutectic showed much higher reactivity with CO2 compared with authentic MgCO3/MgO system
under the condition of 300 °C with PCO2 = 10 bar.[36] Addition of binary salts to the carbonates system
can be a solution for reactivity enhancement.

6

Table 1-3. Turning temperature, reaction enthalpy, and gravimetric energy density of metal
carbonates[6]Chemical formulaTemperature [°C]

Reaction enthalpy [kJGravimetric energymol ]density [kJ kg−1]−1

ZnCO3/ZnO

120

68.326

544.9

CdCO3/CdO

290

96.573560.1

MgCO3/MgO

300

98.984

1174

PbCO3/PbO

310

83.906

314

CaCO3/CaO

885

165.832

1656.8

SrCO3/SrO

1220

202.349

1370.65

BaCO3/BaO

1555

165.096

836.6

Salt hydrates system
Salt hydrates have gained focus to utilize low temperature waste heat.

The basic reaction isshown in below;
𝑆𝑆𝑆𝑆𝑆𝑆𝑆𝑆 ∙ (𝑚𝑚 + 𝑛𝑛)𝐻𝐻2 𝑂𝑂(𝑠𝑠) + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝑆𝑆𝑆𝑆𝑆𝑆𝑆𝑆 ∙ 𝑚𝑚𝐻𝐻2 𝑂𝑂(𝑠𝑠) + 𝑛𝑛𝐻𝐻2 𝑂𝑂(𝑔𝑔)(4)m, n: Stoichometric number

The de/hydration reaction of salt hydrates occurs in deferent steps and intermediate hydrates canappear.[12,37–40] CaSO4∙0.5H2O, MgCl2∙6H2O, MgSO4∙7H2O, CaCl2∙6H2O, and LiOH∙H2O have beenreported. The reversibility of these materials is usually poor, shortly, the rate of the hydration issluggish compared with those of the rate of the dehydration.

Table 1-4 shows the properties ofseveral salt hydrates.
Rammelberg et al. studied the cycle stability of MgCl2∙6H2O, MgSO4∙7H2O, CaCl2∙6H2O, andmixtures of those salts.[37]A mixture of MgCl2 and CaCl2 showed good cycle stability, resulting

almost no decrease over 55 cycles.[37] Li et al. chose polyethylene glycol (PEG), LiCl, 13X-zeolite,
and NaY-zeolite as an additive to improve the hydration reactivity of LiOH. LiOH with added these
additives showed higher hydration reactivity than pure LiOH and the activation energy for dehydrationof LiOH∙H2O, LiOH∙H2O/PEG, LiOH∙H2O/LiCl, LiOH∙H2O/NaY-zeolite, and LiOH∙H2O/13Xzeolite is 52.1 kJ mol−1, 33.2 kJ mol−1, 30.0 kJ mol−1, 25.1 kJ mol−1, and 21.5 kJ mol−1, respectively.[38]Donkers et al. investigated the mechanism of dehydration of MgSO4∙7H2O by nuclear magnetic
resonance (NMR) which showed pore water (an aqueous solution of MgSO4) was produced during
dehydration.[39] K2CO3∙1.5H2O have been studied as a CO2 capture material and showed hydration
reactivity reversibly.[41–43]

Therefore, it is a potential chemical heat storage material. Ogura et al.

investigated the CaSO4/CaSO4∙0.5H2O system for chemical heat pump. They indicated that the system7

can store 373 K level heat and release 273 K level cold heat.[44] Esaki et al. evaluated the CaCl2∙H2O/CaCl2∙2H2O system combined with chemical heat pump system for upgrading waste heat temperaturelevel.[45] They employed an aluminum heat exchanger with corrugated fins and a porous Ni plate fora packed bed reactor. The reactor upgraded waste heat from the hydration reaction of CaCl2∙H2O to
155 °C from a heat storage temperature of 100 °C.[45] Recently, La2(SO4)3 has been focused as a newchemical heat storage material.

The dehydration temperature of La2(SO4)3∙H2O is ~200 °C and

La2(SO4)3∙H2O/ La2(SO4)3 reaction system shows de/hydration cycle with a small hysteresis.[46,47] It
means high de/hydration reaction rate and reduce heat loss during heat storage and output operation.
To elucidate the reaction mechanism of La2(SO4)3∙H2O/La2(SO4)3 reaction system, diffusion
mechanism of water in La2(SO4)3 system was analyzed by first principles calculation, which clarified
that the water maintains its shape of the water molecule in the process of migration and diffuse in the[010] direction.[48]Table 1-4. Dehydration temperature, reaction enthalpy, and gravimetric energy density of salt
hydrates.[40]Chemical formulaTemperature [°C]

Reaction enthalpy

Gravimetric energy[kJ mol−1]

density [kJ kg−1]

MgSO4∙7H2O/MgSO4∙H2O

28

56.0

227.2

CaCl2∙6H2O/CaCl2∙4H2O

36

66.1

301.7

LiOH∙H2O/LiOH

58

64.3

1532.4

CaCl2∙2H2O/CaCl2∙H2O

71

50.6

344.2

MgCl2∙6H2O/MgCl2∙4H2O

72

52.0

255.8

CaSO4∙0.5H2O/CaSO4

102

65.0

477.9

Metal hydrides system
In this system, the basic reaction is shown as below;(5)

𝑀𝑀𝑀𝑀2 (𝑠𝑠) + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝑀𝑀(𝑠𝑠) + 𝐻𝐻2 (𝑔𝑔)In this system, the forward reaction is the dehydrogenation of a metal hydrides corresponding toheat storage process, while the backward reaction is the hydrogenation of a metal corresponding toheat releasing process.

Metal hydrides are one of the promising system for thermal energy storage

due to its high energy density and wide range of operating temperature (Table 1-5).[49] As shown in
Table 1-5, the energy density of those metal hydrides is significantly higher than that of molten salt
and other CHS systems.[49] Those metal hydrides are required for a high hydrogen-storage capacity,
flat plateau meaning that amount of hydrogen that can be stored reversibly and independently with
8

pressure and time, low hysteresis, fast kinetics, easy activation, and high melting point.[11]Furthermore, Metal alloys with a high heat of formation and low disassociation pressure at higher
temperatures are preferable.[11] However, one of the disadvantages for metal hydrides system is theneed of a hydrogen storage subsystem (Figure 1-5).[10]For the sake of overcome for thisdisadvantage, Harries et al. reported the use of a secondary “low-temperature” hydrogen storage
system to capture hydrogen that is released from “high-temperature” hydrides.[50] As the primary
metal hydride is heated with solar heat, the hydrogen will be released. When the primary hydride cools
down, it will absorb hydrogen with an exothermic reaction, producing thermal energy that can be used
for a heat engine.[50] The secondary “low-temperature” hydride play a role for store the hydrogen
released from the “high-temperature” hydride with endothermic hydrogen releasing step.[50]Sheppard et al. evaluated the installed cost including the material cost, material handling, processing,and placement cost in the vessel for four metal hydrides (MgH2, Mg2FeH6, NaMgH3, NaMgH2F).[51]Table 1-5. Comparison of metal hydrides for their main features[49,50]
Chemical formulaTemperature [°C]

Gravimetric energy

Reaction enthalpy [kJ
−1]

mol

density [kJ kg−1]

300–550

15

156

MgH2

450–500

75

2814

TiH

650–700

150

3190

CaH2

1100–1400

186

4426

Molten salt
(NaNO3/KNO3, 60:40
mixture)Table 1-6. Comparison of temperature range and energy density for several metal hydrides system[52]SystemTemperature [°C]

Energy to volume(practical) [kJ dm−3]

Mg-Ni/Mg2NiH4

250–350

824

Mg/MgH2+2 wt% Ni

290–4201806

Mg/MgH2

350–450

1469

Mg-Fe/Mg2FeH6

450–550

2217

Mg-Co/Mg2CoH5

450–5501386

9

Figure 1-4. Layout of a CSP plant working with metal hydride TCS as proposed by PNNL [10] (source
http://energy.gov/eere/sunshot)
They concluded that the cost of NaMgH2F was the lowest, however, the heat storage capacity wasestimated as 1416 kJ/kg which is lower than that of MgH2 (2814 kJ/kg).[51] Reiser et al. investigatedseveral Mg-based metal hydrides (Mg/MgH2, Ni-doped Mg/MgH2, Mg-Ni/Mg2NH4, Mg-Fe/Mg2FeH6,
and Mg-Co-H system) and compared temperature range and energy density for each system (Table 16).[52]Hydroxides system
In this system, the basic reaction is shown as below;(6)

𝑀𝑀(𝑂𝑂𝑂𝑂)2 (𝑠𝑠) + 𝐻𝐻𝐻𝐻𝐻𝐻𝐻𝐻 ⇄ 𝑀𝑀𝑀𝑀(𝑠𝑠) + 𝐻𝐻2 O(𝑔𝑔)The dehydration temperature (for heat storage operation) of the materials in this system is usually

200–450 °C, which is comparably lower than that of other system mentioned above except for salt
hydrides system.[6–12] In this system, the forward reaction is dehydration reaction corresponding toheat storage step while the backward reaction is hydration reaction corresponding to heat output step.
This system requires water vapor for heat output operation, therefore, it is considered as a nontoxic
system.[11] Moreover, hydroxides system is usually cheap, and abundant.[11] Table 1-7 shows theproperties for hydroxides systems.
Magnesium hydroxide (Mg(OH)2) and calcium hydroxide (Ca(OH)2) have been particularly
investigated for chemical heat storage materials. Ogura et al. investigated Ca(OH)2/CaO chemical
10

heat pump system for heat storage and heating/cooling.[53]They mentioned that Ca(OH)2/CaO

chemical heat pump can upgrade the low temperature heat to the high one much more effectively thanmechanical heat pump.[53]The chemical heat pump system achieved higher coefficient of

performances (COP) in the heat enhancement mode and refrigeration mode compared with absorption
and adsorption heat pump.[53] Schaube et al. studied the kinetic of Ca(OH)2 dehydration and CaO
hydration in detail.[54] The dehydration of Ca(OH)2 progressed as a two-step (0.2 < X, X < 0.2, X:
conversion) and they showed an appropriate model for the dehydration under water vapor by taking
into account pressure term based on their study.[54]

Likewise, the hydration of CaO could be

described well by taking into account pressure term.[54] Yan et al. reported that the effect of Li dopingon the dehydration of Ca(OH)2 by preparing a mixture of Ca(OH)2 and LiOH by using a ball mill.[55]The results showed that the dehydration of Ca(OH)2 with Li shifted to a lower temperature comparedwith pure Ca(OH)2 while Li doping had little influence on the reaction enthalpy and the heat storage
capacity.[55] This report implied that the dehydration of Ca(OH)2 with Li progressed as a two-stepreaction,[55] however, the effect can be expected to differ by using other Li compounds (LiCl, LiBr,and so on). Shkatulov et al. investigated the effect of various additive on dehydration of Ca(OH)2.[56]Addition of KNO3 lowered the onset temperature of dehydration by about 40 °C.[56]Table 1-7. Turning temperature, reaction enthalpy, and gravimetric energy density of metalhydroxides[6]Chemical formulaTemperature [°C]

Reaction enthalpy

Gravimetric energy[kJ mol−1]

density [kJ kg−1]

Zn(OH)2/ZnO

55

49.609498.96

Ni(OH)2/NiO

70

47.846516

Be(OH)2/BeO

70

51.276

1191

Cd(OH)2/CdO

125

59.952

409.4

Mn(OH)2/MnO

190

67.072754

Mg(OH)2/MgO

265

77.7451333

Ca(OH)2/CaO

515

100.177

1352

Sr(OH)2/SrO

755

88.581728.3

Ba(OH)2/BaO

1005

93.462545.47Mg(OH)2 is one of the promising chemical heat storage materials for medium temperature (around
300 °C). Many reports have submitted so far about enhancement and investigation of reactivity ofMg(OH)2.
At an early stage, Kato et al. investigated Mg(OH)2/MgO system in a chemical heat pump.[57]They assumed that the hydration reaction of MgO had four reaction regimes for water: (1) containment11

of water as fixed structural water, (2) physical adsorption of water, (3) chemical reaction with waterproducing Mg(OH)2, and (4) inert portion to water.[57]This reaction model was reasonable toAfter that, Ryu et al. proposed a composite of

reproduce the hydration behavior of MgO.[57]Mg(OH)2 mixed with Co(OH)2 or Ni(OH)2 to improve the dehydration reactivity of Mg(OH)2 below
300 °C.[58]As a result, Mg0.5Co0.5(OH)2 and Mg0.5Ni0.5(OH)2 showed the higher dehydration

reactivity at a lower temperature compared with pure Mg(OH)2, it is probably because reaction
equilibrium shifted to a lower temperature, therefore, the hydration reactivity for the oxide of thehydroxide was poor.[58] Moreover, they proposed LiCl-modified Mg(OH)2 (LiCl/Mg(OH)2) by theimpregnation method for a newly promising chemical heat storage material.[59] The authors reported
that LiCl modified Mg(OH)2 showed marked dehydration reactivity compared with pure Mg(OH)2,
NaCl/Mg(OH)2, CaCl2/Mg(OH)2, Mg0.5Co0.5(OH)2 and Mg0.5Ni0.5(OH)2 resulting in the lower onset
dehydration temperature by 44 °C than that for pure Mg(OH)2.[59]In addition, LiCl/Mg(OH)2

de/hydrated reversibly indicating it is a promising chemical heat storage material.

Kim et al.

analyzed hydration reaction kinetic of MgO with mixed expanded graphite (EG) and calcium chloride(CaCl2) by using ash diffusion control model in the shrinking core model.[60,61] While CaCl2 did nothave much influence on the reactivity of Mg(OH)2/MgO, EG enhanced the reactivity (especiallydehydration) due to its high thermal conductivity and porosity.[60] Ash diffusion control model wasable to describe the hydration behavior of MgO with mixed EG and CaCl2 by assuming water
adsorption by CaCl2 at the initial reaction stage.[60] Ishitobi et al. researched LiCl/Mg(OH)2 in detailfor dehydration and hydration behavior.[62,63] At first, they analyzed the dehydration reaction kinetic
and hydration reaction by using thermogravimetric methods and near infrared spectroscopy. The
dehydration behavior of authentic Mg(OH)2 was described by first-order reaction of integration
method, on the other hand, that of LiCl/Mg(OH)2 was successfully reproduced by two step reaction
model based on zero- and first-order reactions.[62] The mixing ratio α = 10 mol% of LiCl was themost optimal mixing ratio and showed the lowest activation energy for dehydration and the mostefficient dehydration.[62] The effect of addition of LiCl on the state of water cluster of the sample
surface was additionally discussed.[62]They also presented the heat output density by hydration

reaction of LiCl/Mg(OH)2 (α = 10 mol%) at a condition of dehydration at 300 °C, subsequently
hydration at 110 °C under water vapor of 57.8 kPa.[63] That paper showed the heat output density by
hydration reaction of LiCl/MgO (α = 10 mol%) was 1180 kJ kg−1 which was about twice as high asthat of pure MgO of 534 kJ kg−1.[63]However, the effect of the addition of LiCl on Mg(OH)2

reactivity has not been clearly understood. Myagmarjav et al. examined the reactivity of Mg(OH)2
with added LiBr and EG (EML).[64–67]

EML showed excellent dehydration reactivity andreversibility compared with pure Mg(OH)2.[64,66,67]

The hydration of MgO with added LiBr and EG

progressed more efficiently at higher temperature under fixed water vapor pressure or at fixed
temperature under lower water vapor than authentic MgO.[67] Tablet type EML showed compatible
12de/hydration reactivity with powder type EML, indicating this sample is likely to be applied forpractical use.[67] In recent years, Shkathlov et al. have investigated the Mg(OH)2/MgO system withvarious additive for chemical heat storage materials. Mg(OH)2 with added sodium nitrate showed more
efficient dehydration at 280 °C than pure Mg(OH)2, and the activation entropy dropped from – (249 ±16) J (mol∙K)−1 for pure Mg(OH)2 to – (382 ± 66) J (mol∙K)−1 for NaNO3/Mg(OH)2.[68] The paper
proposed a schematic model for substitution of hydroxyl group by nitrate anions.

The bidentate

insertion of the anions probably made a positively charged defect of the Mg(OH)2 lattice and it maybe
made Mg(OH)2 more defective.[68] Addition to that, they investigated various salts for Mg(OH)2 andshowed that the addition of LiCl, acetates (Li, Na, and KOAc), LiNO3, or NaNO3 dramatically lowered
the onset temperature of the dehydration of Mg(OH)2.[69]In the paper, they suggested a newmechanism for the effect of the addition of salts to Mg(OH)2. An alkali cation possibly occupied alattice position and the impurity cation defects were formed resulting in the onset temperature
decrease.[69] However, the effect of the hygroscopic salts such as LiCl, LiBr, CaCl2, and Li2SO4 on
dehydration of Mg(OH)2 have not been cleared. Moreover, they analyzed the hydration reaction of
LiNO3-doped MgO.[70] The large amount of LiNO3 leaded to decrease the hydration conversion, it
is probably because the size of MgO particles grew which made the hydration rate slow.[70] They
mentioned the small reaction area of MgO by LiNO3 doping might be also the reason why thehydration rate decreased.[70] Hence, the requirement on the enhancement of hydration reactivity of
MgO especially at high temperatures remains as an issue to address. Furthermore, Mg(OH)2 withadded Ce(NO3)3 was proposed and the addition of 10 wt% Ce(NO3)3 decreased the dehydration peaktemperature of 31 °C compared with pure Mg(OH)2.[71]As mentioned in Section 1.2 100–300 °C waste heat accounts for large amount from industrial
department in Japan. In this study, Mg(OH)2/MgO system was chosen as a CHS material because
the system is cheap, non-toxic, and the temperature for heat storage is 300–400 °C which is close to
the temperature range mentioned above and the hydration reactivity of MgO is relatively highcompared with hydrate system.

In addition, there are few materials for the heat storage at 200–

300 °C, thus, the temperature for heat storage of Mg(OH)2 is required to be lowered. As shown in
the previous studies, some additives were used for the system to utilize waste heat in the range of 200–
300 °C owing to the advantages mentioned above. Tbale 1-8 shows the cost of hydroxides shown in
Figures 1-5–1-7 and the crust abandance of each metal element. This table shows that the cost ofMg(OH)2 is relatively inexpensive and Mg abundantly exists on the Earth. Therefore, Mg(OH)2 canbe easily used as a chemical heat storage material from the view point of cost abundance.
However, considering the reversibility of the system especially the hydration reactivity of MgO athigher temperatures (180–200 °C), the stored heat cannot be utilized efficiently at the heat output
operation meaning that it takes long time to release the stored heat.

Actually, the dehydration

reactivity of the reported materials is not high enough for a practical use. In addition, the mechanism
13

Table 1-8. Cost of hydroxides as precursors and crust abundance of each metal element (1985).[72,73]
The costs of Be(OH)2 and Mn(OH)2 are those of their oxides.
SampleCostofhydroxides

Crust abundance [ppm, ppb, or %][yen/g]Ba(OH)2101

250,000 ppb (250 ppm)

Be(OH)2 (BeO, 99.95%)

5,112

1,500 ppb

Ca(OH)2 (99.9%)82

52,900 ppm (5.29%)

Cd(OH)2 (99.8%, metal basis)7896

98 ppb

LiOH∙H2O

28

13 ppmMg(OH)2 (99.9%)

126

3.20%

Mn(OH)2 (MnO, 99.9%, metal

3046

1,400 ppm

Ni(OH)282.8

0.0105%

Sr(OH)2∙8H2O (metal basis)149

260 ppm

Zn(OH)2

4212

80 ppmbasis)of Mg(OH)2/MgO reaction system and the effect of the addition of additives especially LiCl to thesystem are not clearly and essentially understood. Therefore, the author needs to resolve these things
for applying the CHS system for a practical use and energy conservation system.

Moreover, if thosemechanisms can be elucidated, other compounds which are not restricted will be available for thereactivity enhancement of Mg(OH)2/MgO system (Eq. (7)).

It also enables us to propose a suitablesystem for utilization of a target temperature heat.
Mg(OH)2(s) ⇄ MgO(s) + H2O(g)∆𝐻𝐻r0 = 81.0 kJ mol−1(7)∆𝐻𝐻r0: the standard enthalpy change for the dehydration reaction

1.4 Properties of hydroxides for chemical heat storage
As mentioned in the previous section, the author has focused on magnesium hydroxide (Mg(OH)2)for chemical heat storage.

In this section, the properties of the hydroxides will be confirmed.

Figure 1-5 shows the relationship between turning temperature and cationic radius, and thetemperature seems to be relative to the ionic radius as shown in Figure 1-5. The lower turningtemperature is required to storage lower waste heat. Therefore, the hydroxides with smaller cationicradii can be desirable. On the other hand, the author should also focus on the reaction enthalpy andreversibility for stable and valuable usage.

Figure 1-6 shows the relationship between reaction
14

Figure 1-5. Turning temperature[6,40] vs Shanonn′s cationic radius[74] plots of the hydroxides. The
turning temperature of LiOH is the dehydration temperature of LiOH·H2O.[40] The cationic radii of
Mn2+ are dependent on the state of spin.

Figure 1-6. Reaction enthalpy [kJ mol−1][6,40] vs Turning temperature[6,40] plots of the hydroxides.
The turning temperature of LiOH is the dehydration temperature of LiOH·H2O.[40]15Figure 1-7. Reaction enthalpy [kJ kg−1][6,40] vs Turning temperature[6,40] plots of the hydroxides. The
turning temperature of LiOH is the dehydration temperature of LiOH·H2O.[40]Figure 1-8. Comparison of DSC peak temperature for the reported samples. The DSC curve of pure
Mg(OH)2 and Mg(OH)2 with added LiCl are shown in Appendix A3. The amount of LiNO3, NaNO3,
LiOAc, NOAc, and KOAc were 7 mol%,[69] and that of Ce(NO3)3 was 10 wt%.[71]16enthalpy and turning temperature.

These two parameters can also be relative to each other.

However, the reaction enthalpy should be evaluated in the unit of kJ kg−1 and kJ m−3. Therefore, theenthalpy was described as a unit of kJ kg−1 and compared among the hydroxides as shown in Figure
1-7.
Figure 1-8 shows the plots of peak temperature for pure Mg(OH)2 and previously reported
Mg(OH)2 samples determined by differential scanning calorimetry measurement (the DSC curve forpure and LiCl-added Mg(OH)2 are described in Appendix A3). The peak temperatures were citedfrom previous studies.[69,71] The dehydration peak temperature is typically a temperature where thedehydration is the fastest.[69]Among the reported samples, Mg(OH)2 with added LiOAc showed thelowest peak temperature of apploximately 280 °C. Therefore, the Mg(OH)2 sample whose peaktemperature of below 280 °C can be a break-through material for heat storage at a temperature rangeof 200–250 °C. Hence the peak temperature of 280 °C can be a criterion for heat storage performance.

1.5. Applications for chemical heat storage materials
The chemical heat storage materials can be applied for consentrating solar power plants.[23,26,31,49,51]
In addition, the materials are assumed to be used in a chemical heat pump[36,57] for the usage ofheating/cooling syatem,[53] and energy saving during idling stops on a vehicle.[75] These application
can utilize thermal energy and can be one of the solutions for environmental plobrems.
those application, the reactor design is important.

Considering

Generally, the reactors are classified into twogroups; a packed-bed reactor and a continuous reactor. The packed-bed reactor can be included into
the existing systems by inducing heat transfer fluids.[76] The reaction chamber is separate from thematerial storage tank for the continuous reactor, meaning that the reactor can continuously operate
during heat storage and release steps.[76]
TOYOTA motor Corp., Chubu Electric Power Co., Inc, and TOHO GAS Co., Ltd corporated andinvestigaed the energy system for utilization of waste heat by themal energy transportation.[77] The
system finally estimated that 8,599 GJ year−1 could be saved compared with the original system. The
materials can be also used with renewable energy such as geothermal and solar energies.[78] The
usage with renewable energies can compensate advantages as the author mentioned in Section 1.1.
because the stored heat can be transported to demand area without any heat loss and utilized
independent on weather and time.

In this study, the author aims to enhance the material reactivity

for these applications.

171.6. Purpose and structure of this study
In this study, the author focuses on the three points as shown in below: (1) Improvement for thedehydration and hydration reactivities of Mg(OH)2/MgO system, (2) Analysis for the effect ofadditives on the reactivity of Mg(OH)2/MgO system, and (3) Investigation of the reaction mechanismof Mg(OH)2/MgO system. As mentioned in Section 1.2 the waste heat at 100–350 °C accounts forlarge amount in Japan, however, the dehydration rate of the reported materials at 200–250 °C is
actually sluggish or the reversibility is poor, therefore, the dehydration and hydration reactivity ofMg(OH)2 and MgO at the temperature range should be enhanced for the efficient heat storage andoutput.

For this aim, the peak temperature of 280 °C is one of the criterion for target temperature.

The stored heat must be also used on the heat output operation at higher temperatures (this study aims
for efficient heat output operation at 200 °C) for various usage.

The reaction mechanism of

Mg(OH)2/MgO system and the effect of the additive such as LiCl and LiOH on the system have notbeen cleared (the effect of the addition of LiOH on the dehydration reactivity of Mg(OH)2 is shown in
Chapter 2).

Those phenomena are required to be understood because the elucidation for themenables us to enhance the reactivity of the system more and propose other efficient methods using
unlimited resource because the amount of Li is relatively small.
Chapter 2 mentions the effect of the addition of LiOH on the dehydration of Mg(OH)2. The
dehydration reaction kinetic analysis is carried out to understand the effect of LiOH. In addition, theactivation energy for the dehydration of pure and LiOH-added Mg(OH)2 is investigated.

For aproposal of efficient way to enhance the reactivity of Mg(OH)2/MgO system, the effect of the coaddition of LiCl and LiOH on the reactivity is evaluated in Chapter 3. Furthermore, the chapter
mentions the effect of the addition of LiCl or LiOH on the dehydration based on X-ray diffraction(XRD) measurement. Chapter 4 focuses on the elucidation of the effect of the addition of LiCl or
LiOH on the dehydration.

The XRD and Fourier-transform infrared (FT-IR) spectroscopy

investigate the role of both Li compounds on the dehydration. The chapter proposes possible scheme
for the effect of LiCl or LiOH on the dehydration of Mg(OH)2. Chapter 5 aims to undestsnd theeffect of the addition of LiCl or LiOH on the hydration of MgO. In the chapter, XRD and FT-IR
measurements are used to discuss the effect in detail.

Furthermore, the hydration mechanism of pureand chemically-modified MgO is discussed based on the previous studies and results obtained by XRD
and FT-IR measurements. Chapters 4 and 5 propose possible schemes for the effect of the additionof LiCl and LiOH on the dehydration/hydration of Mg(OH)2/MgO system.

Meanwhile, it isimportant to understand what kinds of factor determin the effective and non-effective additives on thereactivity of Mg(OH)2/MgO system to design a chemical heat storage material. Recently, machinelearning has gained focus on field of desining materials.[79]Therefore, understanding which

properties affect on the reactivity can help to combine the machine learning and material science.
18

Chapter 6 focuses the relationship between sample reactivity and properties of additives.

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24

Chapter 2. Effect of LiOH Addition on Dehydration Reactionof Magnesium Hydroxide
ABSTRACTThe dehydration behavior of LiOH-modified Mg(OH)2 (LiOH/Mg(OH)2), as a new chemical heatstorage material, was investigated. Structural analysis of LiOH/Mg(OH)2 was carried out using an
X-ray diffraction (XRD) meter.

The profile of the dehydration reaction of LiOH/Mg(OH)2 wasobserved to analyze the dehydration kinetics of the LiOH/Mg(OH)2 samples. The XRD patterns of
the LiOH/Mg(OH)2 samples shifted toward higher angles than that of Mg(OH)2. This indicated thatthe lattice parameter might have changed and lattice defects such as oxygen vacancies might havebeen generated on the surface of Mg(OH)2. The rate of the dehydration reaction increased as theamount of added LiOH increased; therefore, LiOH/Mg(OH)2 could be a potential chemical heatstorage material. The author concluded that LO20 (Mg(OH)2 containing 20 mol% LiOH) was the
combination with the most optical mixing mole ratio. This result was determined by analyzing thedehydration activation energy values.

When the dehydration reactions of the LiOH/Mg(OH)2

samples were considered to be first-order reactions, the calculated mole fraction of Mg(OH)2 could
not be reproduced. However, the calculated mole fractions of Mg(OH)2 were in agreement with thosemeasured considering the reactions to be zero- and first-order reactions. Therefore, the kinetics ofthe dehydration reaction of the LiOH/Mg(OH)2 samples included both zero- and first-order reactions.

2.1 Introduction
Recently, global warming and the depletion of fossil fuel energy sources have become importantconcerns as mentioned in Chapter 1.

In Japan, the increased CO2 emission due to the increased useof thermal power plants is also an important factor.

These issues affect not only the present

generation but also future generations. Chapter 1 showed that utilizing waste heat by using chemicalheat storage materials is a potential solution for these concerns owing to its high energy density andunlimited heat storage period.[1–3]

Therefore, many chemical heat storage materials have been

researched for various temperature ranges. In particular, Ca(OH)2 has been researched as a chemicalheat storage material.[4–6]Mg(OH)2 is a chemical heat storage material that can dehydrate atmedium-low temperatures (300–400 °C) as shown in Chapetr 1, and many studies have been done toinvestigate it.[7–12]
The heat storage temperature range for Mg(OH)2 is 300–400 °C. The dehydration reaction of25

Mg(OH)2 corresponds to the heat storage process, while the hydration reaction corresponds to the heat
output process (Eq. 1).
Mg(OH)2(s) ⇄ MgO(s) + H2O(g) ∆𝐻𝐻r0 = 81.0 kJ mol−1(1)0H2O(l) ⇄ H2O(g) ∆𝐻𝐻cnd= 40.7 kJ mol−1(2)0
is the enthalpy change
where ∆𝐻𝐻r0 is the enthalpy change for the dehydration reaction and ∆𝐻𝐻cnd

for the condensation of water.

Since Mg(OH)2 cannot efficiently dehydrate below 300°C, to expand its uses and meet therequirements for using lower-temperature waste heat, the increase of the rate of its dehydration
reaction is needed. Therefore, mixed magnesium hydroxides MgxM1−x(OH)2; M = Ni or Co; x = 0–
1) and LiCl-modified Mg(OH)2 mixtures (LiCl/Mg(OH)2) were developed, which can efficiently store
thermal energy below 300°C.[7–9] However, these materials contain transition metals and a halogen,which is corrosive, and the use of these elements should be minimized.
In this study, LiOH was used as a new additive for Mg(OH)2 and the dehydration behavior of
LiOH-modified Mg(OH)2 (LiOH/Mg(OH)2) was observed.

Moreover, the dehydration kinetics of

LiOH/Mg(OH)2 and Mg(OH)2 were investigated. The author expected LiOH to have the same effectas LiCl on the dehydration of Mg(OH)2, and to reduce the temperature of the dehydration reaction.
The author chose LiOH as additive in our study since it does not contain transition metals or halogens,and therefore, it might be less corrosive than LiCl.

While the kinetics of the dehydration of

LiCl/Mg(OH)2 samples has already been investigated,[9] that of LiOH/Mg(OH)2 for designing a heatstorage system has not been investigated yet.

Moreover, elucidating the mechanism of thedehydration reaction of LiOH/Mg(OH)2 would be important.

2.2 Experimental
2.2.1 Sample preparation and XRD characterization
To prepare LiOH/Mg(OH)2, the author used LiOH·H2O (Wako Pure Chemical Industries, Ltd.)and Mg(OH)2 (99.9%, 0.07 µm, Wako Pure Chemical Industries, Ltd) as precursors. LiOH/Mg(OH)2
was prepared using the impregnation method.
LiOH·H2O and pure water.

Aqueous LiOH solutions were prepared using

Mg(OH)2 was impregnated with the aqueous LiOH solutions andevaporated at 40 °C using a rotary evaporator. Subsequently, the samples were dried overnight at
120 °C. A series of LiOH/Mg(OH)2 samples with Mg(OH)2:LiOH mole ratios of 100:5, 100:10,
100:20, 100:25, 100:33, and 100:50 were prepared. These samples will hereafter be referred to as26

LO5, LO10, LO20, LO25, LO33, and LO50, respectively.

For comparison, Mg(OH)2 without added

LiOH, hereafter referred to as Mg(OH)2-W, was also subjected to the same tests.
X-ray diffraction (XRD) analysis was carried out to analyze the crystal structure of Mg(OH)2-W
and the LiOH/Mg(OH)2 samples using an Ultima IV (Rigaku Corp.) XRD instrument. The 2θ range
was between 10° and 80°, the scanning rate was 10.0° min−1, and the scan width was 0.01°. The
XRD instrument used Cu Kα radiation, a 40 kV generator voltage and 40 mA current.

2.2.2 Dehydration reaction measurements
The weight changes of the samples during the dehydration reaction were measured using a TGD9600, (ULVAC RIKO, Inc) thermobalance. The samples were loaded into a Pt cell and dehydrated
at various dehydration reaction temperatures (Td) as follows. All samples, including Mg(OH)2-W,
were dehydrated by increasing the temperature from room temperature (20–25 oC) to 600 °C at aheating rate of 10 °C min−1 under a 100 mL min−1 flow of Ar.
To analyze the kinetics of the dehydration reaction, 4 samples (LO20, LO25, LO33, and LO50)
were dehydrated at 230, 250, 270, and 300 °C, whereas 3 samples (Mg(OH)2-W, LO5, and LO10)
were dehydrated at 270, 300, 320 and 350 °C. The dehydration reactions were carried out under a
100 mL min−1 flow of Ar. The heating rate was 20 °C/min when Td was 270, 300, 320 and 350 °C,
and 10°C min−1 when Td was 230 and 250 °C. To analyze the effect of removing the physicallyabsorbed and hydration water from LiOH, the weight of the sample at 200 °C was defined as the initialsample weight.
The mole fraction of Mg(OH)2 was calculated using the weight change data obtained using thethermobalance, as follows:𝑥𝑥 =(𝑤𝑤𝑖𝑖𝑖𝑖𝑖𝑖 − 𝑤𝑤𝑓𝑓𝑓𝑓𝑓𝑓 ) − (𝑤𝑤𝑖𝑖𝑖𝑖𝑖𝑖 − 𝑤𝑤)(3)(𝑤𝑤𝑖𝑖𝑖𝑖𝑖𝑖 − 𝑤𝑤𝑓𝑓𝑓𝑓𝑓𝑓 )𝑤𝑤𝑓𝑓𝑓𝑓𝑓𝑓 = 𝑤𝑤𝑖𝑖𝑖𝑖𝑖𝑖 × 𝑀𝑀

𝑀𝑀𝑀𝑀𝑀𝑀𝑀𝑀(4)

𝑀𝑀𝑀𝑀(𝑂𝑂𝑂𝑂)2

where wini is the initial weight of Mg(OH)2 (the weight of Mg(OH)2 at 200 °C) [mg], wfin is the weightthat reacted theoretically (the weight of MgO) [mg], w is the weight of the sample during thedehydration reaction [mg], MMg(OH)2 and MMgO are the molecular weights of Mg(OH)2 and MgO[g/mol], respectively, and x is the mole fraction of Mg(OH)2 [-].

2.2.3 Dehydration reaction rate analysis
Analyzing the kinetics of the dehydration reactions was carried out using the integration method.[9]27The author assumed that the dehydration reaction progressed as a first-order reaction and occurred
homogeneously within the particles. Therefore, the rate of the dehydration reaction can be written
as follows:

−

𝑑𝑑𝑑𝑑
𝑑𝑑𝑑𝑑

(5)

= 𝑘𝑘1 𝑥𝑥

where k1 [1/s] is the rate constant of a first-order reaction, x [-] is the mole fraction of Mg(OH)2, and t[s] is the reaction time.
Equation (5) can be integrated using the initial conditions (t = t0, x = x0) to yield Eq. (6) as follows:(6)

𝑙𝑙𝑙𝑙𝑙𝑙 = −𝑘𝑘1 𝑡𝑡 + 𝑙𝑙𝑙𝑙𝑥𝑥0

According to Eq. (6), the rate constant can be obtained from the slope of the plot of t vs. lnx refinedby a least square method. To remove the effect of the heating process on the dehydration reaction,
the initial x was set as 0.9. Subsequently, Eq. (6) was solved for x, where the calculated value for themole fraction of Mg(OH)2 was obtained from Eq. (7):𝑥𝑥 = 𝑒𝑒 −𝑘𝑘1𝑡𝑡 ∙ 𝑥𝑥0(7)

To evaluate the activation energy of each sample, the Arrhenius Equation (Eq. (8)) was used:𝐸𝐸(8)𝑙𝑙𝑙𝑙𝑙𝑙 = − 𝑅𝑅𝑅𝑅𝑎𝑎 + 𝑙𝑙𝑙𝑙𝑙𝑙

where Ea is the activation energy [kJ mol−1], R is the gas constant [J (K·mol) −1], T is temperature [K],and A [s−1] is the frequency factor.

2.3. Results and Discussion
2.3.1 Characterization of Mg(OH)2-W and LiOH/Mg(OH)2 using XRD
Figure 2-1 illustrates the XRD patterns of Mg(OH)2-W and the LiOH/Mg(OH)2 samples.

All

XRD patterns presented similar patterns in the (002), (102), and (110) planes. These characteristic
planes indicated that Mg(OH)2 presented a brucite structure.

No LiOH peaks were detected for the

LO5 and LO10 samples, which indicated that LiOH was well dispersed throughout the Mg(OH)2
particles, or LiOH dissolved into the Mg(OH)2 phase. The XRD patterns of LO20–LO50 included
28

the peaks of LiOH, which indicated the presence of LiOH on the surface of Mg(OH)2. Subsequently,
Figure 2-2 illustrates the expanded XRD patterns of the LiOH/Mg(OH)2 samples in the (002) plane.
The yellow dotted lines represent the XRD peak position of the (002) plane of Mg(OH)2-W. The
peaks of the (002), (102), and (110) planes of the LiOH/Mg(OH)2 samples shifted toward higher angles
compared to that of Mg(OH)2-W. These results indicated that the lattice parameters of the samples
decreased, as indicated by Bragg’s law (nλ = 2dsinθ), and that Li+ replaced Mg2+ in the crystal
lattices.[13] This could be attributed to the difference between the ionic radii of Mg2+ (72 pm) and Li+(76 pm).[13] Moreover, Li-doped MgO exhibited similar results. Li-doped MgO presented semiconductive properties.

If a system of a sample had good electrical affinity, the bond length would be

reduced, therefore, its lattice parameters would decrease.[13] Hence, the bond length of Mg(OH)2might also decrease as LiOH was added to it, which decreased the lattice parameter.

Lattice defects

such as oxygen vacancies might be generated on the surface of Mg(OH)2 due to the substitution of
Mg2+ by Li+, because the charges are no longer balanced when monovalent cations (Li+) replace
divalent cations (Mg2+). At these sites, water molecules might be easily released from the surface ofMg(OH)2; the rate of the dehydration reaction of the LiOH/Mg(OH)2 samples might be higher thanthat of Mg(OH)2-W.

Figure 2-1. XRD patterns of : (a) Mg(OH)2-W, (b) LO5, (c) LO10, (d) LO20, (e) LO25, (f) LO33, and(g) LO50

2.3.2. Effect of LiOH addition on the dehydration reaction of Mg(OH)2
Figure 2-3 presents the behavior of the LiOH/Mg(OH)2 samples subjected to dehydration
reactions as the temperature increased to 600°C.

Furthermore, Figure 2-3 demonstrates that adding
29

LiOH to Mg(OH)2 led to a decrease in the onset temperature of the dehydration reactions for all the
LiOH/Mg(OH)2 samples compared to that for Mg(OH)2-W. However, the behaviors of LO25–LO50
during the dehydration reactions were almost the same.
Since Mg(OH)2 can be dehydrated at 300–350 °C, this indicated that the LiOH/Mg(OH)2 samples
could efficiently be dehydrated at a lower temperature of 300 °C.

2.3.3. Effect of LiOH addition on the rate of the dehydration reaction
Figures 2-4 (a)–(g) illustrate the profile of the dehydration reaction for each sample at 270 °C.
The rate of the dehydration reaction increased as the amount of added LiOH increased. However,
the rates of the dehydration reactions of LO25–LO50 were almost the same.

Thus, the fastest

dehydration reaction was probably associated with a constant LiOH amount and the reaction rate did
not always increase as the amount of LiOH increased.
To understand the mechanism of the dehydration reactions of the LiOH/Mg(OH)2 samples, kinetic
analysis of the LiOH/Mg(OH)2 samples was carried out. To analyze the kinetics of the dehydration
reactions, the analysis plot of the reaction rate of the samples, which was based on Eq. (6) are presented
in Figures 2-5 (a)–(g). The reaction rate constant was obtained using the best fit line of the plot.
For example, the reaction rate constant of LO10 at 270 °C was calculated as 7.4 × 10−5 (s−1).
Figures 2-6 (a)–(g) show a comparison of the profiles of the calculated and measured dehydration
reactions of the samples at 270 °C. The calculated mole fraction of Mg(OH)2 was obtained using thereaction rate constant from Figures 2-5 (a)–(g).

The results indicated that the calculated value

mostly agreed with the measured value for Mg(OH)2-W, LO5, and LO10. Therefore, the dehydration
reaction of those samples proceeded as a first-order reaction.

Figures 2-6 (d)–(g) show comparisonof the profiles of the calculated and measured dehydration reactions of LO20, LO25, LO33, and LO50
at 270 °C. Here, the calculated value did not agree with the measured value. Thus, a first-order
reaction failed to reproduce the dehydration reaction profile of LO20, LO25, LO33, and LO50.
These results indicated that a zero-order reaction might influence the kinetics of the dehydration
reaction, just as it influenced the kinetics of the dehydration reaction of LiCl/Mg(OH)2.[9] Ishitobi etal. demonstrated that both zero-order and first-order reactions affected the kinetics of the dehydration
reactions.[9]The LiOH/Mg(OH)2 samples were also analyzed using the zero- and first-order

reactions to investigate the kinetics of their dehydration reactions.

30

Figure 2-2. XRD peaks of (002) planes of: (a) LO5, (b) LO10, (c) LO20, (d) LO25, (e) LO33, and (f)LO50. The yellow dotted lines indicate the positions of the peak of the (002) plane for Mg(OH)2-W

31

Figure 2-3. Behavior of dehydration of the LiOH/Mg(OH)2 samples upon heating 600 oC.

Figure 2-4. Dehydration reaction profile of all samples at 270 oC.

32

Figure 2-5. Reaction rate analysis plot of (a) Mg(OH)2-W, (b) LO5 (c) LO10, (d) LO20, (e) LO25,(f) LO33, and (g) LO50 at 270 oC.

33

Figure 2-6. Calculated and measured mole fractions of Mg(OH)2 for (a) Mg(OH)2-W, (b) LO5 (c)LO10, (d) LO20, (e) LO25, (f) LO33, and (g) LO50 at 270 oC.

34

The author analyzed the kinetics of the dehydration reactions of the LiOH/Mg(OH)2 samples aszero-order reactions when 0.6 ≤ x ≤ 0.9, 0.5 ≤ x ≤ 0.9, and 0.4 ≤ x ≤ 0.9.[9] Table 2-1 lists thecoefficient of determination (R2) values when each range of x values was approximated as a straightline. These results indicated that when the author analyzed the kinetics of the dehydration reactionsof the LiOH/Mg(OH)2 samples as zero-order reactions for 0.6 ≤ x ≤ 0.9, the calculated mole fractions
reproduced the measured values the best. Thus, the author assumed that the dehydration reaction of
LiOH/Mg(OH)2 proceeded as a zero-order reaction when 0.6 ≤ x ≤ 0.9. A zero-order reaction can be
described as follows:

−

𝑑𝑑𝑑𝑑
𝑑𝑑𝑑𝑑

(9)

= 𝑘𝑘0

where k0 [s−1] is the rate constant of the zero-order reaction.

Equation (9) can be integrated asfollows:(10)

𝑥𝑥 = −𝑘𝑘0 𝑡𝑡 + 𝑥𝑥0

It was assumed that the dehydration reaction proceeded as a zero-order reaction when the mole

fraction of Mg(OH)2 varied between 0.9 and 0.6, and as a first-order reaction when the mole fraction
ranged between 0.5 and 0.2. The calculated mole fractions of Mg(OH)2 when 0.5 ≤ x ≤ 0.6 were
weighted average values obtained from the values based on zero- and first-order reactions.
Figures 2-7 (a)–(d) show comparison of the measured and calculated mole fractions as one-step(first-order reactions) and two-step processes (zero- and first-order reactions) for LO20, LO25, LO33,
and LO50. These figures indicate that the two-step process was able to reproduce the measured mole
fractions.

Other LiOH/Mg(OH)2 samples exhibited similar behaviors. Therefore, the kinetics ofthe dehydration reaction of the LiOH/Mg(OH)2 samples (LiOH amount ≥ 20 mol%) included zeroand first-order reactions.

Furthermore, it is possible that the rate controlling step was changed during

the transition from the zero- to the first-order reaction.

These results were in agreement with thosereported in a previous study.[9] However, further studies are required to reveal why a zero-order
reaction changed into a first-order reaction and if the transition of reaction order was simple enoughto describe as weighted average values nor not during the dehydration reaction. In addition, the range
of zero- and first-order reactions is unclear to define.

35

Figure 2-7. Calculated and measured mole fractions of Mg(OH)2 for (a) LO20, (b) LO25, (c) LO33,
and (d) LO50 at 270 oC: one-step (green line) and two-step process (red line).
The author confirmed that the onset temperature of the dehydration reaction of Mg(OH)2 waslowered by adding LiOH. To clarify the effect of adding LiOH, Arrhenius plots for the samples areillustrated in Figure 2-8, where the blue, orange, and green lines represent the Arrhenius plots of thesamples for the one-step process, zero-, and first-order reactions of the two-step process, respectively.
The activation energy of each sample can be calculated from the slope of the Arrhenius plot in
Figure 2-8. Tables 2-2 (a), (b) and 3 (a), (b) list the activation energies and frequency factors of the
one-step (only first-order reaction) and two-step models (zero- and first-order reactions) for all thesamples, respectively.
Since the LO20 sample presented the lowest activation energy (Table 2-2 (b)), the rate of thedehydration reaction of the LO20 sample should have been the highest.

However, Figure 2-4

illustrates that the rate of the dehydration reaction of LO20 was lower than those of the LO25–LO50
samples at 270 °C. These results indicated that the rate of the dehydration reaction did not depend
only on the activation energy.

36Figure 2-8. Arrhenius plot of dehydration constant for (a) Mg(OH)2-W, (b) LO5. (c) LO10, (d) LO20,(e) LO25, (f) LO33, (g) LO50.
37

Table 2-1. Evaluating linearity for each range of y values using the coefficient of determination (R2)for LO20 at 270 °C
Range of zero-order reaction

0.6 ≤ x ≤ 0.9

0.5 ≤ x ≤ 0.9

0.4 ≤ x ≤ 0.9

R2

0.99600.98770.9691

Table 2-2 (a). Activation energies of Mg(OH)2-W, LO5, and LO10
SampleEa [kJ mol−1]
One-step

Mg(OH)2-W

156

LO5

152

LO10

148
Table 2-2 (b). Activation energies of LO20, LO25, LO33, and LO50
Ea [kJ mol−1]

SampleOne-stepTwo-step, Zero-orderTwo-step, First-order

LO20

135

134

135

LO25

156

158

153

LO33

145

144

145

LO50

157

154

158

Table 2-3 (a). Frequency factors of Mg(OH)2-W, LO5, and LO10
SampleA [s−1]One-step

Mg(OH)2-W

4.91 × 1010

LO5

2.84 × 1010

LO10

1.40 × 1010
Table 2-3 (b). Frequency factors of LO20, LO25, LO33, and LO50
A [s−1]

SampleOne-stepTwo-step, zero-orderTwo-step, first orderLO20

2.18 × 109

1.37 × 109

1.96 × 109

LO25

4.13 × 1011

6.41 ×1011

1.82 × 1011

LO33

3.37 × 1010

2.37 × 1010

3.02 × 1010

LO50

5.86 × 1011

2.35 ×1011

6.05 × 1011

38

When the amount of added LiOH was lower than 20 mol%, the added LiOH was not adequate tosignificantly increase the rate of the dehydration reactions; the number of defect sites on the surfaceof Mg(OH)2 would be small (water molecules might be easily released from the surface of Mg(OH)2at these sites, thus adding LiOH to Mg(OH)2 promoted the dehydration reaction of Mg(OH)2). Whenthe amount of added LiOH was greater than 20 mol%, the added LiOH significantly increased the rateof the dehydration reaction; furthermore, the number of defect sites on the surface of Mg(OH)2 would
be large.
When adding LiOH to Mg(OH)2, LiOH particles might partially cover the surface of Mg(OH)2,
therefore this might prevent the dehydration of Mg(OH)2. When the amount of added LiOH was
smaller than 20 mol%, the obstruction effect caused by the LiOH particles might be greater than the
promoting effect of LiOH on the rate of the dehydration reaction of Mg(OH)2. Thus, the activationenergies of LO5 and LO10 were higher than that of Mg(OH)2-W. However, no peaks for LiOH weredetected when the samples were subjected to XRD analysis (Figure 2-1). These results were similar
to those reported in previous studies conducted using LiCl/Mg(OH)2.[8] Therefore, analyzing thesurface of Mg(OH)2 would be required to clarify whether LiOH species were present on the surfaceof Mg(OH)2 or not.
When the amount of added LiOH was higher than 20 mol%, the number of lattice defects would
be large, however large LiOH particles might be generated on the surface of Mg(OH)2 as detected by
XRD (Figure 2-1). Therefore, more energy would be required to pass water molecules between
these large LiOH particles, and the activation energy of the dehydration reactions increased for the
samples containing more than 20 mol% LiOH. When the amount of added LiOH was 20 mol%, thenumber of defects could be large (water molecules could be easily released easily from many sites),
and the diffusion resistance of the LiOH particles would be insignificant; the balance between thenumber of defects and diffusion resistance of the LiOH particles was reached when the amount ofadded LiOH was 20 mol%. Thus, the activation energy of LO20 was the lowest of all the analyzed
samples.
Therefore, the rate constant was a function of the activation energy and temperature (Eq. 2-8),
while the activation energy could be considered to be a function of the number of defects and diffusion
resistance. These results indicated that LO20 presented the best mixing ratio with respect to theactivation energy.
In this study, the author did not investigate the variability of kinetic factors between different
sample lots. Thus, further studies of the reaction kinetics are required for practical use.

39

2.4 Conclusions
The effect of adding LiOH on the structure and dehydration reaction of Mg(OH)2 was investigatedin this study. The (002) plane of the LiOH/Mg(OH)2 samples shifted toward higher angles comparedto that of Mg(OH)2-W. This indicated that Li+ entered the brucite lattice sites and replaced Mg2+.
This might have led to lattice defects such as oxygen vacancies. At these sites, water molecules could
be easily released from the surface of Mg(OH)2. However, the dehydration rate did not necessarily
increase as the amount of added LiOH increased.
Adding LiOH lowered the onset temperature and increased the rate of the dehydration reaction ofMg(OH)2. Therefore, adding LiOH promoted the dehydration reaction of Mg(OH)2 and, therefore,
LiOH/Mg(OH)2 could be considered as a new potential chemical heat storage material that does notcontain any transition metals or halogens. A first-order reaction could not necessarily reproduce thedehydration reaction profile of Mg(OH)2, while a two-step process including zero- and first-order
reactions could reproduce the dehydration reaction of Mg(OH)2. Thus, the dehydration reaction ofMg(OH)2 might involve both zero- and first-order reactions. Meanwhile, further studies are required
to clarify why a zero-order reaction changed into a first-order reaction and if the transition of reaction
order was simple enough to describe as weighted average values nor not during the dehydration
reaction.

In addition, the range of zero- and first-order reactions is unclear to define. The activationenergy for the dehydration reaction of LO20 was the smallest of all the analyzed reactions, therefore,
LO20 had the best mixing ratio.

However, further study is required to clarify why a zero-order

reaction changed into a first-order reaction and if the transition of reaction order was simple enoughto describe as weighted average values nor not during the dehydration reaction, and the range of zeroand first-order reaction.

References[1] P. Pardo, A. Deydiera, Z. Anxionnaz-Minvielle, S. Rougéa, M. Cabassud, P. Cognetb; “A Review
on High Temperature Thermochemical Heat Energy Storage”, Renewable and Sustainable Energy
Reviews, 32, (2014), 591–610[2] D. Aydin, S. P. Casey, S. Riffat; “The latest advancements on thermochemical heat storage systems”,
Renewable and Sustainable Energy Reviews, 41, (2015), 356–367[3] L. André, S. Abanades, G. Flamant; “Screening of thermochemical systems based on solid-gas
reversible reactions for high temperature solar thermal energy storage”, Renewable and Sustainable
Energy Reviews, 64, (2016), 703–715[4] H. Ogura, T. Yamamoto, H. Kagea, Y. Matsuno, A. S. Mujumdar; “Effects of heat exchange
condition on hot air production by a chemical heat pump dryer using CaO/H2O/Ca(OH)2 reaction”,
40Chemical Engineering Journal, 86, (2002), 3–10
[5] F. Schaube, A. Kohzer, J. Schutz, A. Worner, H. Muller-Steinhagen; “De- and rehydration of
Ca(OH)2 in a reactor with direct heat transfer for thermo-chemical heat storage. Part A: experimental
results”, Chemical Engineering Research and Design, 91, (2013), 856–864[6] A. Shkatulov, Y. Aristov; “Modification of magnesium and calcium hydroxides with salts: An
efficient way to advance for materials for storage of middle-temperature heat”, Energy, 85, (2015),
667-676[7] J. Ryu, R. Takahashi, N. Hirao, Y. Kato; “Effect of transition metal mixing on reactivities ofmagnesium oxide for chemical heat pump”, Journal of Chemical Engineering of Japan, 40, No. 13,(2007), 1281–1286[8] H. Ishitobi, N. Hirao, J. Ryu, Y. Kato; “Evaluation of heat output densities of lithium chloridemodiﬁed magnesium hydroxide for thermochemical energy storage”, Industrial & Engineering
Chemistry Research, 52, (2013), 5321–5325[9] H. Ishitobi, K. Uruma, M. Takeuchi, J. Ryu, Y. Kato; “Dehydration and hydration behavior of
metal-salt-modified materials for chemical heat pumps”, Applied Thermal Engineering, 50, (2013),
1639–1644[10] A. Shkatulov, T. Krieger, V. Zaikovskii, Y. Chesalov, Y. Aristov; “Doping magnesium hydroxidewith sodium nitrate: A new approach to tune the dehydration reactivity of heat-storage materials”, ACS
Applied Materials & Interfaces, 6, (2014), 19966–19977[11] O. Myagmarjav, J. Ryu, Y. Kato; “Waste heat recovery from iron production by using magnesiumoxide/water chemical heat pump as thermal energy storage”, ISIJ International, 55, No. 2, (2015),
464–472[12] Z. Pan, C. Y. Zhao; “Dehydration/hydration of MgO/H2O chemical thermal storage system”,
Energy, 82, (2015), 611–618[13] J. Sivasankari, S. Selvakumar, K. Sivaji, S. Sankar; “Structural and optical characterization of
MgO: X (X = Li, Na, and K) by solution combustion technique”, Journal of Alloys and Compounds,
616, (2014), 51–57

41

Chapter 3. Comparison of Effect of Co–addition of Li
Compounds and Addition of Single Li Compound on
Reactivity and Structure of Magnesium Hydroxide
ABSTRACT
Mg(OH)2 is a chemical heat storage material that is studied for the utilization of 300–350 °C wasteheat.

In this study, LiCl and LiOH were co-added to Mg(OH)2 and the reactivity and structural

evolution were investigated. In the hydration of the samples at 200 °C subsequent to dehydration at
270 °C, Mg(OH)2 with co-added LiCl and LiOH showed excellent hydration reactivity, with a heat
output density of 1053 kJ kg−1.

LiCl and LiOH co-addition enhanced both the dehydration andhydration reactivity of Mg(OH)2.

X-ray diffraction analysis indicated that addition of LiOH to

Mg(OH)2 promoted the decomposition of Mg(OH)2 and the diffusion of water on the surface ofMg(OH)2, while addition of LiCl to Mg(OH)2 promoted these processes in the bulk phase of Mg(OH)2.

3.1. Introduction
As mentioned in previous sections, chemical heat storage materials can utilize waste heat via
chemical reactions, especially gas-solid reactions.[1–4]

Several reaction systems such as

Ca(OH)2/CaO,[5–7] MgCl2∙6H2O/MgCl2, and MgSO4∙7H2O/MgSO4[8, 9] have been investigated.

For

storage of heat in the 450–550 °C range, the dehydration of Ca(OH)2 is suitable, but is unsuitable forstorage of low-medium temperature heat (200–300 °C). Despite the low temperature of ~100 °C
required for dehydration of salt hydrates such as MgSO4∙7H2O and MgCl2∙6H2O, the reversibility ofthis process for these materials is quite low, where hydration (heat output) requires a long time. The
Mg(OH)2/MgO system has been studied for storage of low-medium (300–350 °C) temperature
heat.[10–17] The dehydration temperature for Mg(OH)2 is quite low compared with that for Ca(OH)2,and the hydration reactivity of MgO is much higher than that of MgSO4.
In previous studies, several additives such as cetyl trimethyl ammonium bromide, lithium chloride,and lithium hydroxide-modified Mg(OH)2 have been shown to enhance the reactivity of Mg(OH)2.[11–15]In particular, LiCl-added Mg(OH)2 and LiOH-added Mg(OH)2 (shown in Chapter 2) were much

more efficiently dehydrated at 270–300 °C than pure Mg(OH)2.[12–14] Therefore, the author believethat the Mg(OH)2/MgO system has significant potential for thermal energy storage at 200–300 °C,
which constitutes a large part of the industrial waste heat in Japan. However, the hydration rate of42

MgO with these added salts and that of pure MgO at 180–200 °C are sluggish. Thus, the applicable
reaction conditions for the Mg(OH)2/MgO system are still limited.
C. Knoll et al. studied the relationship between the specific surface area and the reactivity of the
Mg(OH)2/MgO system.[16] However, there are few studies on the effect of additives on the reactivityof materials; one such study was carried out by Shkatulov et al..[15] The effect of LiCl or LiOH
addition on the reactivity and structure of the Mg(OH)2/MgO system has not yet been reported. If
the mechanism can be elucidated, it can be extended to other chemical heat storage materials withadded salts.
In this study, the author propose the use of LiCl and LiOH co-added Mg(OH)2 to enhance thereactivity of Mg(OH)2 and investigate the effect of adding the Li compounds on the reactivity andstructure of Mg(OH)2. In several similar studies, the effect of the two different Li compounds (LiCland LiOH) on the reactivity of the Mg(OH)2/MgO system has not been reported. Moreover, the roleof LiCl or LiOH on the dehydration reactivity of Mg(OH)2, which has not been revealed earlier, wasclarified in this study.

3.2. Experimental
3.2.1. Sample preparation
To prepare LiCl and LiOH co-added Mg(OH)2 (referred to as LiCl/LiOH/Mg(OH)2) and LiCl- or
LiOH-added Mg(OH)2 (referred to as LiCl/Mg(OH)2 and LiOH/Mg(OH)2, respectively), LiCl·H2O(99.9%, Wako Pure Chemical Industries, Ltd.), LiOH·H2O (Wako Pure Chemical Industries, Ltd.),
and Mg(OH)2 (99.9%, 0.07 µm, Wako Pure Chemical Industries, Ltd) were used as precursors.
LiCl/Mg(OH)2, LiOH/Mg(OH)2, and LiCl/LiOH/Mg(OH)2 were prepared by the impregnation
method. Detail explanations for the method was described in Chapter 2. LiCl-added Mg(OH)2
with a Mg(OH)2:LiCl mole ratio of 100:10 (referred to as L10), LiOH-added Mg(OH)2 with
Mg(OH)2:LiOH = 100:20 (referred to as LO20), LiCl and LiOH co-added Mg(OH)2 with
Mg(OH)2:LiCl:LiOH = 100:5:5 and 100:10:10 (referred to as L5/LO5 and L10/LO10, respectively)
were prepared by this method.

For comparison, Mg(OH)2 without the Li compounds, referred to as

Mg(OH)2-W, was prepared by the same method. All the prepared samples are shown in Table 3-1.
The previous studies showed that L10 and LO20 were the best mixing ratios with respect to theactivation energy for dehydration and for enhancing the dehydration and hydration reactivity of
LiCl/Mg(OH)2 and LiOH/Mg(OH)2, respectively.[12–14]
comparison of the reactivity in this study.

43

Hence, these samples were used forTable 3-1. Mixing ratio of the prepared samples
SampleMixing ratio [mole ratio]

Mg(OH)2-W

Mg(OH)2 without the Li compounds

L10

Mg(OH)2:LiCl = 100:10

LO20

Mg(OH)2:LiOH = 100:20

L5/LO5Mg(OH)2:LiCl:LiOH = 100:5:5

L10/LO10

Mg(OH)2:LiCl:LiOH = 100:10:10

3.2.2. Evaluation of reactivity using thermobalance
The reactivity of all samples was measured by using a thermobalance (TGD–9600 series,
ADVANCE RIKO, Inc.). The samples (~20 mg) were charged into a Pt cell and then heated underconstant Ar flow. The mole fraction of Mg(OH)2 was calculated as shown in Chapter 2 based on
the thermogravimetric (TG) data measured by using the thermobalance.[12–14]
Figure 3-1 shows the TG curve for the sample dehydrated at 270 °C and then hydrated at 110 °C.
The first weight loss is attributed to the removal of physically adsorbed water. The sample weight
did not change at 200 °C; therefore, the sample weight at 200 °C is used as the initial sample weightin this study.
Figure 3-2 is the schematic diagram for reaction test. For the dehydration reaction test, an Ar
folw was carried ([3]→[4]) into the thermobalance, and the sample in the Pt cell was heated. For thehydration reaction test, the reaction gas mixture of water vapor and Ar gas was induced into thethermobalance through 4-way valve ([1]→[4]).

After the hydration, the Ar gas was induced([3]→[4]) for drying process of the sample.

3.2.2.1. Dehydration reaction
The conditions for the dehydration reaction are shown below; the sample was heated from room
temperature (20–25 °C) to various temperatures (220–800 °C) at a rate of 10 °C min−1 under Ar
flowing at 100 mL min−1.

3.2.2.2. Hydration reaction
For the hydration reaction, to remove physically adsorbed water, the sample was heated at 120 °C
for 30 min at a rate of 20 °C min−1 under Ar flowing at100 mL min−1.

The sample was thendehydrated at 270 or 350 °C for 30 min at a heating rate of 20 °C min−1 under Ar flowing at 100 mL
min−1. The sample was then hydrated at constant hydration temperatures (110, 170, and 200 °C) for
80 min at a heating rate of −20 °C min−1 by introducing a gas mixture of water vapor at 57.8 kPa andAr gas. Subsequently, the introduction of water vapor was stopped and the sample was dried at thehydration reaction temperature for 30 min.
44

Figure 3-1. TG curve for sample dehydrated at 270 °C under 100 mL min−1 Ar flow and subsequently
hydrated at 110 °C with Ar gas and water vapor mixture at 57.8 kPa.

Figure 3-2. Schematic diagram for reaction test.

45

Figure 3-3. Data for hydration reaction test; dehydration temperature (Td) = 270 °C under 100 mL
min−1 Ar flow, hydration temperature (Th) = 110 °C with Ar gas and water vapor mixture at 57.8 kPa.
Figure 3-3 shows the dehydration and hydration profiles of L10/LO10. As shown in Figure 33, x0 is the mole fraction of Mg(OH)2 before the hydration reaction [–], xh is that of Mg(OH)2 at theend of hydration [–], xc is that of Mg(OH)2 at 10 min after the end of hydration [–], ∆xd is thedehydration reaction conversion [%], ∆x1 is the hydration reaction conversion [%], and ∆x2 is thereaction conversion by physical water adsorption [%]; ∆xd, ∆x1, and ∆x2 were calculated as percentages,
as shown below:(1)

𝛥𝛥𝛥𝛥𝑑𝑑 = (1 − 𝑥𝑥0 ) × 100(2)

𝛥𝛥𝑥𝑥1 = (𝑥𝑥𝑐𝑐 − 𝑥𝑥0 ) × 100(3)

𝛥𝛥𝛥𝛥2 = (𝑥𝑥ℎ − 𝑥𝑥𝑐𝑐 ) × 100

To evaluate the effect of the hydration reaction temperature (Th) on the hydration of MgO, Th wasvaried as 110, 170, and 200 °C and the water vapor pressure was fixed at 57.8 kPa. To evaluate theeffect of the water vapor pressure on the hydration of MgO, the water vapor pressure was varied as
7.4–57.8 kPa by controlling the water flow using a micro feeder (NP-KX-101, NIHON SEIMITSUKAGAKU CO.LTD) and the Ar balance, and Th was fixed at 110 °C.

Table 3-2 shows therelationship between the water vapor pressure (PH2O) and saturated temperature (Ts).

46

Table 3-2. Relationship between PH2O and the saturated vapor temperature (Ts)Ts [°C]PH2O [kPa]

40

7.4

50

12.3

60

19.9

70

31.2

80

47.4

85

57.8

3.2.2.3. Evaluation of heat output density
The heat output densities of all samples were calculated from the enthalpy change of the hydration
reaction and that during the sorption of water vapor (eqs (4) and (5)). Herein, the heat output density
is expressed as the heat output density per unit weight (kg) of the sample.

It was assumed that theenthalpy change for the sorption of water vapor corresponds to that of the condensation of water. The
heat output densities were calculated as follows:

𝑄𝑄𝑟𝑟 = −Δ𝐻𝐻ℎ𝑦𝑦𝑦𝑦 ×
𝑄𝑄𝑠𝑠 = −Δ𝐻𝐻𝑎𝑎𝑎𝑎 ×

∆𝑥𝑥1

100

∆𝑥𝑥2

100

(4)

× 𝑀𝑀ℎ𝑦𝑦𝑦𝑦(5)× 𝑀𝑀ℎ𝑦𝑦𝑦𝑦

where Qr is the heat output density of the hydration reaction [kJ kg−1], Qs is the heat output density forthe condensation of water [kJ kg−1], Mhyd is the weight of magnesium hydroxide per kg of sample [kg].
ΔHhyd and ΔHad are the enthalpy change of the reaction and condensation per kg magnesium hydroxide[kJ kg−1], respectively.[13,14]

3.2.2.4. Sample characterization by XRD
To investigate the crystal structure of the samples, X-ray diffraction (XRD) analysis was carriedout in triplicate to reduce the error by using an Ultima IV (Rigaku Corp.) X-ray diffractometer. The
2θ range was 10−150°, the scan rate was 10.0° min−1, and the scan width was 0.01°. Cu-Kα radiation
was used with a generator voltage of 40 kV and 40 mA current.

Before measurement, the sample

was dried at 120 °C overnight to remove physically adsorbed water.
To analyze the structure of pure Mg(OH)2 and the Li compounds-added Mg(OH)2, the lattice
parameters and volumes of all the samples were calculated by the least-squares method.

These

values were adjusted to match the measured peak positions by this method.[18] The 90% confidence
interval for the calculations was estimated by Student’s t-distribution based on eq (6).
47𝑥𝑥(𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚𝑚) = 𝑥𝑥̅ ± 2.920

𝑢𝑢

(6)

√𝑁𝑁

where N is the number of measurements (in this study, N = 3), and u is the standard deviation.

3.3. Results and discussion
3.3.1. Comparison of the effect of addition of single Li compound versus co–additionof the Li compounds on dehydration reactivity of Mg(OH)2
Figure 3-4 shows the dehydration behavior of all samples when heated to 600 °C. The y-axis
represents the mole fraction of Mg(OH)2 based on eq (3), the x-axis represents the temperature. The
purple line shows the dehydration behavior of Mg(OH)2-W.
added LiCl and/or LiOH progressed at lower temperatures.

The dehydration of Mg(OH)2 with

Interestingly, the co-addition of LiCl and

LiOH to Mg(OH)2 (especially L10/LO10) resulted in dehydration at lower temperature than in thecase of single addition (L10, LO20). For example, at 270 °C, the slop of the TG curve of L10/LO10
was significantly larger than that observed for any other sample. These results show that the coaddition of LiCl and LiOH to Mg(OH)2 enhanced the dehydration reactivity of Mg(OH)2 to a greater
extent than the addition of LiCl or LiOH alone.

Figure 3-4. Dehydration behavior of all samples with heating to 600 °C.

48

Although Mg(OH)2-W did not undergo complete dehydration, the XRD patterns of all the samples
heated to 600°C showed no peaks derived from Mg(OH)2 (Figure S1-1). A previous study showedthat pretreatment at temperatures lower than 600 °C did not lead to the complete removal ofstructurally bound water.[19]An IR analysis revealed that the peak at around 3700 cm−1

corresponding to isolated hydroxyl groups on the Mg(OH)2 surface did not disappear (Figure S1-2).[20,21]This peak was shifted to 3730 cm−1 after Mg(OH)2-W was heated at higher than 400 °C. Thisbehavior clearly indicates that the hydroxyl groups of Mg(OH)2 surface were converted to the ones of
MgO surface at higher than 400 °C. However, when we carefully looked at the absorption band ataround 3730 cm−1, a small component due to the hydroxyl groups of Mg(OH)2 still existed at 3700
cm−1.[20, 21] Therefore, this result might cause the incomplete dehydration of Mg(OH)2. Thus, themole fraction did not reach 0%, as seen in Figure 3-4.
samples, however, were more than 90%.
complete.

The dehydration conversions of all theHence, the author considered that the dehydration was

Further research is required to elucidate this effect in detail and to investigate othersamples.
Table 3-3 shows the peak temperature for dehydration of all the samples, derived from thedifferential thermogravimetry (DTG) curve (Figure 3-5 (a)–(e)) based on Figure 3-4.

The peaktemperature is the temperature at which the dehydration rate is the fastest. Here, the dehydration rate
was calculated from the measured weight change and temperature, and is the sample weight
differentiated by the temperature.

A lower peak temperature means that the sample can be

dehydrated at lower temperature, which indicates that the sample is suitable for lower-temperature
heat storage.

L10/LO10 showed the lowest peak temperature of all the samples.

Therefore,

L10/LO10 is expected to store low grade waste heat more efficiently than L10 and LO20, which havebeen studied previously.[12–14] These results indicated that the temperature range at which Mg(OH)2
can be efficiently dehydrated was increased by the co-addition of LiCl and LiOH. Moreover, detailed
investigation of the dehydration kinetics analysis at 200–300 °C is required in future work.

3.3.2. Effect of co-addition of the Li compounds on hydration reactivity of MgOFigure 3-6 shows the dehydration and hydration behavior of all the samples. The sample was
first heated at 270 °C for 30 min and subsequently hydrated at 110 °C for 80 min under a mixture of
Ar gas and saturated water vapor at 57.8 kPa. Under these conditions, the ∆xd value for Mg(OH)2W and that for LO20 was only 9.50 and 55.2%, respectively. The ∆xd value for L10, L5/LO5, andL10/LO10 was 89.7, 89.0, and 95.0% respectively, indicating that these three samples are suitable foruse under the stated conditions. However, L10/LO10 was dehydrated the fastest of all the samples,
similar to the case in Figure 3-4. All samples could be reversibly (de)hydrated. Thus, Mg(OH)2
with co-added LiCl and LiOH should be useful as a chemical heat storage material.

49

Figure 3-5. Differential thermogravimetry (DTG) curve of (a) Mg(OH)2-W, (b) L10, (c) LO20,(d) L5/LO5, and (e) L10/LO10 upon heating to 600 °C.
Table 3-3. Peak temperature of all samples
SamplePeak temperature [°C]

Mg(OH)2-W

371

L10

317

LO20

342

L5/LO5314

L10/LO10305

50

Figure 3-6. Dehydration and hydration behavior of all the samples; Td = 270 °C, Th = 110 °C, PH2O =57.8 kPa.
Table 3-4 shows the reaction conversion for all samples. The ∆x2 value for L10 was much higher
than that of Mg(OH)2-W and LO20. This is probably because of the high hygroscopicity of LiCl.
The mole fraction change ∆x2 indicates the formation of the n hydrate of LiCl or LiOH, as expressedin eqs. (7) and (8). This result demonstrates that LiCl is easily converted into LiCl∙nH2O than is
LiOH, so that the hydration of MgO is promoted.[13] nH2O in LiCl∙nH2O probably interacted with
the interface between MgO and LiCl∙nH2O.[13](7)

𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 + 𝑛𝑛𝑛𝑛2 𝑂𝑂 ↔ 𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 ∙ 𝑛𝑛𝑛𝑛2(8)

𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 + 𝑛𝑛𝑛𝑛2 𝑂𝑂 ↔ 𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 ∙ 𝑛𝑛𝑛𝑛2 𝑂𝑂

Table 3-4. Reaction conversion for all samples; Td = 270 °C, Th = 110 °C, PH2O = 57.8 kPa.

Sample∆xd [%]∆x1 [%]∆x2 [%]

Mg(OH)2-W

9.50

7.80

2.30

L10

89.7

98.0

52.5

LO20

55.2

54.5

4.40

L5/LO5

89.0

81.1

26.4

L10/LO10

95.0

90.1

51.1

Figure 3-7 shows the hydration behavior of L10, L5/LO5, and L10/LO10 at 110 and 200 °C for
80 min after dehydration at 270 °C under Ar flow for 30 min. Although L10 and L5/LO5 were well51

hydrated at 110 °C, only 30–40% hydration was achieved at 200 °C. Generally, hydration does notprogress effectively at higher temperature because it is an exothermic reaction. Thus, it is naturalthat the hydration reactivity of the samples declined with increasing temperature.[13]However,

L10/LO10 was well hydrated at 110 and 200 °C, although the apparent hydration reaction rate
decreased at 200 °C. Therefore, the hydration reactivity of L10/LO10 was higher at high hydration
temperatures as compared with that of the other samples.
Figure 3-8 shows the heat output density for the hydration reaction of all samples at three different
hydration temperatures (110, 170, and 200 °C); the orange dotted line shows the heat output densityof 1000 kJ kg−1, which is the recommended target value considering the economy and repayment of
the initial cost in Japan.[22] The heat output density at 110 °C exceeded 1000 kJ kg−1, except in thecase of LO20 (only 55.2% dehydration of LO20 was achieved at 270 °C; Figure 3-6).

However,

except for LO20 and L10/LO10, the heat output density decreased dramatically with increasing
temperature for the reason mentioned above. The only sample for which the heat output density
exceeded 1000 kJ kg−1 at all temperatures was L10/LO10 (1185, 1069, and 1053 kJ kg−1 at 110, 170,
and 200 °C, respectively). The amount of LiCl and LiOH might be insufficient for improving thehydration reactivity of MgO for L5/LO5 under the stated conditions.

Thus, L10/LO10 had highhydration reactivity even at high temperature (200 °C), which can reduce heat loss during thedehydration and hydration cycles. Thus, this sample can release high-temperature and high-qualified

Figure 3-7. Hydration behavior of L5/LO5, L10/LO10, and L10 at 110 and 200 °C; Td = 270 °C, Th =110 and 200 °C, PH2O = 57.8 kPa. The Blue line shows the hydration behavior at 110 °C, the red line
shows that at 200 °C. “◆”, “●”, and “■” indicate the hydration behavior of L5/LO5, L10/LO10,
and L10 respectively.
52

Figure 3-8. Heat output density at various hydration temperatures for all the samples; Td = 270 °C, Th= 110, 170, and 200 °C; PH2O = 57.8 kPa.

Figure 3-9. Hydration behavior of all samples at 200 °C; Td = 350 °C, Th = 200 °C, PH2O = 57.8 kPa.
MgO-W was obtained by a dehydration at 350 °C.
heat in the heat output operation. This feature is very valuable for the heat output operation, and ismuch better than that of L10, which was previously studied.[12, 13]These results show that thedehydration and hydration reactivity of Mg(OH)2/MgO can be more effectively enhanced by coaddition of LiCl and LiOH than by the individual addition of LiCl or LiOH. In the former case, the53

dehydration temperature shifts more dramatically toward a lower level and the hydration reactivity of
MgO is enhanced to a greater extent, as seen in Figures 3-4 and 3-9. MgCO3/MgO with an added
LiNO3–KNO3 eutectic showed much better reactivity than the pure MgCO3/MgO system in a previous
study.[23] Hence, the addition of binary salts can improve the reactivity of other chemical heat storage
materials.
The heat output density of LO20 was almost constant, independent of the temperature. Therefore,
LiOH possibly has good effects on the hydration reactivity of MgO, such as maintaining highhydration reactivity at high hydration temperatures. Thus, LiCl and LiOH might exert independent
effects on the hydration of MgO. While MgO can plausibly absorb water due to the hygroscopicityof the added LiCl,[13] the hydration of MgO at high temperature may be promoted by LiOH addition(Figure 3-9).
After dehydration at 350 °C, hydration of the samples was evaluated to clarify the effect of theaddition of LiCl and/or LiOH on the hydration reactivity of MgO. Figure 3-9 shows a comparisonof the reactivity of all samples when hydrated at 200 °C for 80 min under a mixture of water vapor (at57.8 kPa) and Ar gas after dehydration at 350 °C for 30 min. The ∆x1 value for Mg(OH)2-W and
L10 was only 2.89 and 16.8% respectively. On the other hand, that of ∆x1 for LO20 was 49.7%. The∆x1 value for L5/LO5 and L10/LO10 was 89.4% and 94.3%, respectively, and the hydration of L5/LO5and L10/LO10 was almost complete at 200 °C, even after dehydration at 350 °C. The hydration
reactivity of MgO was especially improved by LiCl and LiOH co-addition. From this result, theauthor conclude that the co-addition of LiCl and LiOH enhanced not only the dehydration reactivity,
but also the hydration reactivity at 200 °C, of Mg(OH)2/MgO. In particular, LiOH addition promoted
this enhancement at 200 °C to a greater extent than LiCl addition.

Surprisingly, ∆x1 for L5/LO5 after

dehydration at 350 °C was much higher than that after dehydration at 270 °C. After dehydration at
350 °C, the specific surface area might have increased, because of which the reactive area for thehydration of MgO was larger than that in the case of dehydration at 270 °C.[24] As the hydration of
L10/LO10 after dehydration at 270 °C was almost complete, the reactivity of L10/LO10 was higher
than that of L5/LO5.
This result shows that LiOH addition promoted the hydration of MgO at 200 °C.

Interestingly,

the co-addition of LiCl and LiOH promoted the dehydration and hydration of Mg(OH)2/MgO to amuch greater extent than LiCl or LiOH singly addition.

However, further studies are required tounderstand why LiCl and LiOH co-addition greatly promoted the hydration of MgO at 200 °C.
Figure 3-10 shows the hydration behavior of L10/LO10 at 110 °C with variation of the water vapor
pressure (7.4–57.8 kPa); the relationship between the water vapor pressure and saturated temperature
is shown in Table 3-2. As shown in Figure 3-10, the extent of hydration of MgO decreased when
the water vapor pressure was lowered. This behavior agrees with the fact that the number of watermolecules decreased at lower water vapor pressure and with a previous study of L10.[13] For PH2O =54

19.9, 12.3, and 7.4 kPa, the ∆x1 value for L10/LO10 was only 59.1, 27.1, and 11.5%, respectively.
Below PH2O = 19.9 kPa, the hydration reactivity of L10/LO10 declined dramatically. Thus, the coaddition of LiCl and LiOH could not enhance the hydration reactivity at low water vapor pressure.
Heat output at higher temperature can be achieved by increasing the water vapor pressure.[25]However, the hydration reactivity at low water vapor pressure must be enhanced for utilization of lowgrade waste heat. For instance, the temperature of waste heat is supposed to be as low as 30 °C.
Therefore, if the low water vapor pressure can be used efficiently in the heat output operation, theapplication scope of the Mg(OH)2/MgO system can be widened, e.g., a cooling system.[26]3.3.3. Sample characterization by XRD analysis for evaluating the effect of additionof the Li compounds on structure of Mg(OH)2
Figure 3-11 shows the XRD patterns of all the samples before the reaction. This figure shows
the peaks of the brucite structure of Mg(OH)2 for all samples. The data indicate the dominance ofthe brucite structure of Mg(OH)2 in all the samples, although LiCl and/or LiOH were added to
Mg(OH)2. No peaks of LiCl were detected in the LiCl-added samples (L10, L10/LO10, and L5/LO5)
because the LiCl species were probably dissolved in the Mg(OH)2 phase, well dispersed throughout
the Mg(OH)2 particles, or converted into solution, which is not detectable during the experiment.

Figure 3-10. Hydration behavior of L10/LO10 at various water vapor pressures; Td = 270 °C, Th =110 °C, PH2O = 7.4–57.8 kPa.

55

In

Figure 3-11. XRD patterns of all samples.
the future work, the author must investigate the presence of LiCl by techniques such as XAFS. LiOH
might be partially converted to Li2CO3 during preparation, which agrees with the fact that LiOH isgenerally contaminated with Li2CO3.[27] This accounts for the detection of the Li2CO3 peak.
Because of the difficulty in detecting any changes in the XRD pattern in Figure 3-11, two of thestrongest peaks derived from the Mg(OH)2 (001) and (002) planes of each sample are expanded in
Figures 3-12 (a) and (b). The black dotted line indicates the position of the Mg(OH)2 (001) or (002)
peaks of Mg(OH)2-W. For Mg(OH)2 with added LiCl and/or LiOH, the peaks of the (001) and (002)planes shifted towards lower and higher angle, respectively.

Figure 3-12. Expanded XRD patterns of all samples; black dotted line shows peak position ofMg(OH)2 (001) or (002) plane of Mg(OH)2-W: (a) (001) plane, (b) (002) plane.

56

The author evaluated this change from only two planes, but evaluation of all measured planes isrequired.

The lattice parameters and volumes of all samples were calculated based on all themeasured XRD peak positions to clarify the factors causing the peak shift. Table 3-5 shows thecalculated lattice parameters and volumes for all samples. These values were calculated by the leastsquares method using all measured XRD peak positions.

The measurement was carried out in

triplicate to reduce the error.
Table 3-5. Calculated lattice parameters and lattice volume of all samples based on measured peak
position
SampleLattice parameter: a [Å]Lattice parameter: c [Å]Lattice volume: V [Å3]Mg(OH)2-W

3.145 ± 0.003

4.763 ± 0.007

47.12 ± 0.15

L10

3.147 ± 0.001

4.787 ± 0.031

47.42 ± 0.30

LO20

3.146 ± 0.004

4.783 ± 0.007

47.33 ± 0.08

L5/LO5

3.146 ± 0.002

4.775 ± 0.005

47.27 ± 0.12

L10/LO10

3.145 ± 0.001

4.766 ± 0.013

47.15 ± 0.16

The lattice parameters and volumes of Mg(OH)2 with added LiCl and/or LiOH were larger thanthose of Mg(OH)2-W. Thus, the values seemed to increase by addition of LiCl and/or LiOH, where
Li+ ions may have been substituted into the Mg2+ sites.[28] The effective ionic radius of Li+ (0.76 Å)
is larger than that of Mg2+ (0.72 Å),[29] indicating that the lattice parameters increased due to Li+ ion
substitution.

Li+ substitution could disrupt the charge balance at the substituted sites due to thereplacement of a divalent cation (Mg2+) by a monovalent cation (Li+). Therefore, lattice defects suchas oxygen vacancies might be generated to compensate the disruption in the charge balance caused by
Li+ ion substitution. At such sites, the decomposition of Mg(OH)2 may occur readily.[14]The increase in the lattice parameters was small.

However, even this small difference might

influence the reactivity of the Mg(OH)2/MgO system. Further, the effective ionic radii of Li+ andMg2+ are very close to each other, and even a slight change in the lattice parameters is reasonable.
The increase in the lattice parameter c, as shown in Table 3-3, is indicative of expansion of theinterlayer distance in the brucite structure based on first-principles calculation, as cited from a previous
article.[30] This plausibly promoted release/insertion of water molecules from/into the structure ofMg(OH)2/MgO, which could lead to enhancement of the rate of dehydration and hydration ofMg(OH)2/MgO.
In summary, the (1) lattice defects and (2) the expansion of the inter-layer distance in the brucite
structure by Li+ ion substitution plausibly influence the dehydration of Mg(OH)2. On the other hand,point (2) has a more profound effect on the hydration of MgO.
The results of this study qualitatively indicate that the lattice parameters increase upon LiCl and/or
57

LiOH addition as compared to those of pure Mg(OH)2. Nonetheless, more analysis is required toclearly reveal the effects mentioned in this study, by drastically increasing the amount of the Licompounds added to Mg(OH)2 or by simulation using first-principles calculations.
Table 3-3 shows that L10 had a lower peak temperature for the dehydration reaction and thatdehydration progressed at lower temperature compared with LO20, meaning that LiCl and LiOH havedifferent effects on the dehydration of Mg(OH)2. To elucidate the effect of the addition of LiCl and
LiOH on the dehydration of Mg(OH)2, each sample was heated to various temperatures and XRD
analysis was carried out for the samples after dehydration.
Figure 3-13 (a) shows the XRD patterns of Mg(OH)2-W after heating to various temperatures.
The black dotted line shows the peak position of Mg(OH)2-W after heating to 310 °C and the red line
shows that of MgO after heating to 380 °C. This figure shows the appearance of the highest intensity
peak at 42° corresponding to the MgO (200) plane in the XRD pattern of the sample treated at 310–
320 °C, whereas the highest intensity peak corresponding to the Mg(OH)2 (002) plane at 38°
disappeared for the sample treated at 370–380 °C.

These observations are indicative of thedecomposition of Mg(OH)2 and that the diffusion of water molecules on the surface of Mg(OH)2
started at 310–320 °C.

Likewise, the decomposition of Mg(OH)2 and the diffusion of watermolecules in the bulk phase were completed at 370–380 °C.
The same measurement was carried out for other samples, and the XRD patterns are presented in
Figures 3-13 (b)–13 (e). Table 3-6 shows the temperatures at which the peak corresponding to the
MgO (200) plane at around 42° appeared and that of the Mg(OH)2 (002) plane at around 38°
disappeared for all samples.

For LO20, the MgO (200) peak at ~42° appeared at much lower

temperature compared with that of Mg(OH)2-W, and the Mg(OH)2 (002) peak at ~38° disappeared atmuch lower temperature for L10 compared with that of Mg(OH)2-W. These results indicate that
LiOH promoted the decomposition of Mg(OH)2 and the diffusion of water molecules on the surfaceof Mg(OH)2, whereas LiCl promoted the decomposition of Mg(OH)2 and the diffusion of watermolecules in the bulk phase. Therefore, it is possible that LiCl and LiOH act cooperatively in thedehydration of Mg(OH)2.
From the discussion presented above, it is deduced that the decomposition of Mg(OH)2 anddiffusion of water molecules both on the surface and in the bulk phase of Mg(OH)2 were promoted in
the case of L10/LO10. Thus, the dehydration of L10/LO10 progressed at the lowest temperature(Figure 3-3).
Figure 3-14 shows a plot of the heat-treatment temperature vs. the logarithm of the intensity ratioof the MgO (200) and Mg(OH)2 (002) peaks from the XRD patterns (Figure 3-13). The orange
dotted line indicates the point where the y axis value is equal to 0, meaning that the intensity of the
MgO (200) and Mg(OH)2 (002) peaks are equal.
In case of Mg(OH)2-W, the plot was in the highest temperature side, indicating that the dehydration
58

Figure 3-13 (a). XRD patterns of Mg(OH)2-W after heating; the black dotted line shows the peak position of Mg(OH)2after heating to 310 °C, the red line shows that of MgO after heating to 380 °C.

(b) XRD patterns of L10 after heating

to various temperatures (250–320 °C); black dotted line shows peak position of Mg(OH)2 after heating to 250 °C, thered line shows that of MgO after heating to 310 °C.

(c) XRD patterns of LO20 after heating to various temperatures(230–300 °C); black dotted line shows the peak position of Mg(OH)2 after heating to 300 °C, the red line shows thatof MgO after heating to 300 °C. (c)' XRD patterns of LO20 after heating to various temperatures (300–370 °C); black
dotted line shows the peak position of Mg(OH)2 after heating to 300 °C, the red line shows that of MgO after heating
to 370 °C.

(d) XRD patterns of L5/LO5 after heating to various temperatures (240–320 °C); black dotted line shows

the peak position of Mg(OH)2 after heating to 240 °C, the red line shows that of MgO after heating to 310 °C.

(e)

XRD patterns of L10/LO10 after heating to various temperatures (220–310 °C); black dotted line shows the peak
position of Mg(OH)2 after heating to 220 °C, the red line shows that of MgO after heating to 300 °C.

59

Figure 3-14. Heating temperature vs. log plot; y axis shows logarithm of intensity ratio for Mg(OH)2and MgO, x axis shows the heating temperature [°C].
of Mg(OH)2-W progressed at higher temperature as compared with the case of the other samples.
The figure clearly shows that the dehydration of L10/LO10 progressed at the lowest temperature
because the plot fell on the lowest temperature side. This trend is consistent with the TG analysis
presented in Figure 3-4 and Table 3-3.
Table 3-6. Temperature of disappearance of Mg(OH)2 (002) peak at 38° and appearance of MgO (200)peak at 42° for all samples
SampleTemp. of disappearance

Temp. of appearance

[°C]

[°C]

Mg(OH)2–W

370–380

310–320

L10

310–320

250–260LO20

360–370

230–240

L5/LO5

310–320

240–250

L10/LO10

300–310220–230

3.3.4. Comparison of heat storage/output performances between our samples andreported samples
Here, the heat storage and output performances of our and reported samples are compared as shownin Figures 3-15 and 3-16 to highlighten the validity of the co-addition of LiCl and LiOH. Figure 315 shows the comparison of DSC peak temperature between the samples previously reported and60

proposed in our sudy (The DSC curves of Mg(OH)2-W, L10, LO20, and L10/LO10 are described inAppendix A3).

The peak temperature means the temperature at which the dehydration rate ofMg(OH)2 is the fastest,[6] which is the same definition as that of the dehydration peak temperature(described in p. 47). The DSC peak temperature of L10/LO10 was 305 °C, which was the third
lowest of the samples.

The reaction condition was different with each study while these results

indicate that the co-addition of LiCl and LiOH was efficient way to improve Mg(OH)2 dehydration
reactivity. To the contrary, the peak temperature should be lowered to apploximatetly 250 °C for theefficient heat storage at 200–250 °C as mentioned in Section 1.1.
Figure 3-16 shows the comparision of heat output density by hydration between the samples
previously reported and proposed in our sudy. The heat output densities of MgO-W, L10, LO20,
L10/LO10, and EML were obtained by the hydration at 200 °C under water vapor of 57. 8 kPa for 140
min[33] while that of LiNO3/MgO was obtained by the hydration at 200 °C under water vapor of 33.5
kPa for 90 min.[34] Here, the hydration condition was not the same with each other, the heat output
density of L10/LO10 was, however, drastically high among the samles.

In addition, the highhydration reactivity of L10/LO10 at 200 °C was advantageous with regard to the heat output operation
for wide-spred usage.[5] Therefore, these results clealy show that the co-addition of LiCl and LiOH
is effective way to improve both Mg(OH)2 dehydration and MgO hydration reactivities.

Figure 3-15. Comparison of DSC peak temperature between the samples previously reported andproposed in our sudy. The amount of additive was 7 mol% for [6], that of Ce(NO3)3 was 10 wt%,[31]and those of Ce(NO3)3 and LiOH were 8 and 6 wt%, respectively.[32]61Figure 3-16. Comparison of heat output density by hydration between the samples previously reported
and proposed in our sudy. The densites of MgO-W, L10, LO20, and L10/LO10 were obtained fromresults in Figure 3-8.

EML is the Mg(OH)2 with added LiBr of 10 mol% and expanded graphite of

33 wt%.[33] This sample was heated at 300 °C and then, was hydrated at 110 °C under water vapor
MgO with added LiNO3 of 10 wt% was heated at 400 °C, andof 57.8 kP for 140 min.[33]subsequently hydrated at 150 °C under water vapor of 33.5 kPa for 90 min.[34]3.4. Conclusions
Mg(OH)2 was singly doped with LiCl or LiOH and co-doped with LiCl & LiOH to enhance its
reactivity for application as a chemical heat storage material. The reactivity of LiCl/LiOH/Mg(OH)2
was also investigated and the effect of the addition of LiCl and/or LiOH on the structure of Mg(OH)2
was analyzed.
LiCl/LiOH/Mg(OH)2 (L10/LO10) showed much better dehydration and hydration reactivity than
LiCl/Mg(OH)2 (L10) and LiOH/Mg(OH)2 (LO20).

In particular, almost complete hydration of

L10/LO10 was achieved at 200 °C after dehydration at 350 °C, while pure Mg(OH)2, L10, and LO20
were hardly hydrated. Thus, L10/LO10 can be used to supply high-temperature and qualified-heatin heat output operations.

Therefore, the co-addition of LiCl and/or LiOH enhances both thedehydration and hydration reactivity to a greater extent than the addition of LiCl or LiOH individually.
On the other hand, the co-addition of LiCl and LiOH could not enhance the hydration reactivity of
MgO at low water vapor pressure. This is one of the main issues to be resolved for heat utilization
over a wide temperature range.

In addition, the dehydration temperature of Mg(OH)2 is required to62

be lowered more for the heat storage at 200–250 °C.

XRD analysis indicates that the Li+ ionssubstituted into the Mg2+ site cause inter-layer expansion in the brucite structure, where the lattice
parameters increased upon addition of LiCl and/or LiOH.

Li substitution plausibly induced theformation of lattice defects such as oxygen vacancies due to disruption of the charge balance at thesubstituted sites.

These effects promoted the dehydration and hydration of Mg(OH)2 and MgO.

However, further studies are required for the clear elucidation of these effects by dramatically additionof the Li compounds (LiCl and/or LiOH) or by simulation such as first-principles calculations.

It isplausible that LiOH promoted the decomposition of Mg(OH)2 and the diffusion of water molecules on
the surface of Mg(OH)2, while LiCl promoted the decomposition of Mg(OH)2 and the diffusion ofwater molecules in the bulk phase. Further studies are required to elucidate the effect of co-additionof LiCl and LiOH on the hydration of MgO. The author must also confirm the cycle repeatability of
dehydration and hydration in the future work.

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Applied Energy, 253, (2019), 113562[25] Y. Kato, Y. Sasaki, Y. Yoshizawa, “Magnesium oxide/water chemical heat pump to enhanceenergy utilization of a cogeneration system”, Energy, 30, (2005), 2144–2155[26] H. Ogura, T. Yamamoto, H. Kage, “Efficiencies of CaO/H2O/Ca(OH)2 chemical heat pump forheat storing and heating/cooling”, Energy, 28, (2003), 1479–1493[27] Y. Hase, “Raman spectroscopic study of four isotopically substituted lithium hydroxidemonohydrates”, Monatshefte für Chemie, 112, (1981), 73–82[28] J. Sivasankari, S. Selvakumar, K. Sivaji, S. Sankar, “Structural and optical characterization of
MgO: X (X = Li, Na, and K) by solution combustion technique” Journal of Alloys and Compounds,
616, (2014), 51−57[29] R. D. Shannon, “Revised effective ionic radii and systematic studies of interatomic distance in
halides and chalcogenides” Acta Crystallographica Section A, 32, (1976), 751–767[30] Pascale, F.; Tosoni, S.; Zicovich-Wilson, C.; Ugliengo, P.; Orlando, R.; Dovesi, R. Vibrational
spectrum of brucite, Mg(OH)2: A periodic ab initio quantum mechanical calculation including OH
anharmonicity. Chem. Phys. Lett. 2004, 396, 308–315.
[31] L. Sun, Q. Wu, L. Zhang, Y. Li, M. Li, T. Gao, S. Guo, D. Wang, ”Doping magnesium hydroxidewith Ce(NO3)3 : A promising candidate thermochemical energy storage materials”, IOP Conf. Ser.:
Earth Environ. Sci., 295, (2019), 032068.

[32] “Tremendous enhancement of heat storage efficiency for Mg (OH)2-MgO-H2O thermochemicalsystem with addition of Ce(NO3)3 and LiOH”, Nano Energy, 81, (2021), 105603.
[33] O. Myagmarjav, J. Ryu, Y. Kato, “Lithium bromide-mediated reaction performance enhancementof a chemical heat-storage material for magnesium oxide/water chemical heat pumps”, Appl. Therm.
Eng., 63, (2014), 170–176.
[34] A. Shkatulov, H. Takasu, Y. Kato, Y. Aristov, “Thermochemical energy storage by LiNO3-doped
Mg(OH)2: rehydration study”, Journal of Energy Storage, 22, (2019), 302–310.

65

Chapter 4. Fourier-Transform Infrared Analysis of the
Dehydration Mechanism of Magnesium hydroxide andChemically Modified Magnesium Hydroxide

ABSTRACT
Mg(OH)2 is a medium-temperature chemical heat storage material.

The addition of LiCl and/or

LiOH to Mg(OH)2 has been reported to promote the dehydration of Mg(OH)2.

However, themechanism underlying the effect of the addition of Li compounds on the dehydration of Mg (OH)2 has
not yet been elucidated. This study focused on the dehydration of pure Mg(OH)2 and Mg(OH)2 withadded LiCl and/or LiOH. The author analyzed the role of the Li compounds in the dehydration ofMg(OH)2 using X-ray diffraction and Fourier-transform infrared (FT-IR) spectroscopy.

Our FT-IR

measurements assigned the peaks of the bulk OH groups of Mg(OH)2, which are not referred to inprevious studies.

FT-IR measurements revealed that the surface state of Mg(OH)2 with added LiOH

with a Mg(OH)2:LiOH mole rate of 100:20 was similar to that of MgO.

This indicated that theaddition of LiOH to Mg(OH)2 affected the dehydration of the Mg(OH)2 surface.

Moreover, theauthor confirmed that LiCl affected the dehydration of the bulk Mg(OH)2 phase. Using these results,
the author proposed possible mechanisms for the effects of the addition of LiCl and/or LiOH on thedehydration of Mg(OH)2. Lastly, the author discuss the reason for the Li compounds promoting thedehydration of Mg(OH)2.

4.1. Introduction
The utilization of unused heat, such as waste heat (industrial waste and surplus solar heat) is an
attractive option. The calculated annual amount of waste heat in Japan is approximately 740 000
TJ,[1] and waste heat is a type of thermal energy resource. Consequently, heat storage techniques
have gained increasing attention, and the main heat storage techniques nowadays are sensible, latent,
and chemical heat storage (CHS) techniques. Water, molten salts, basalt, and concrete have beenreported as sensible heat storage materials, whereas mannitol, paraffin, fatty acids, polyethylene glycol,
and microcapsule phase-change materials have been used as latent heat storage materials.[2–8]In

addition, the heat storage density and long-term heat storage parameters for the CHS technique are
typically better than those for the other two techniques.[9–14]
66

To date, several reports on CHS materials have been published.
MgCl2·6H2O,[15,16]

LiOH·H2O,[17]

and metal sulfates,[16,18–20]

Salt hydrates, such as

have been investigated as CHS materials

in the relatively low temperature range of 100–250 °C. In contrast, Mg(OH)2,[21–33] Ca(OH)2,[27,34–
36]

and MgCO3[37] are suitable for thermal energy storage at 200–500 °C and have been explored owing

to their high energy density and reversibility.

Among the CHS materials, metal oxides andcarbonates have been examined for thermal energy conservation in the temperature range of 800–
1000 °C, owing to their relatively high energy density.[38–40] Therefore, we focused on Mg(OH)2 as
a medium-temperature CHS material.

The dehydration of Mg(OH)2 and hydration of MgO

correspond to the heat storage and output operations, respectively.

The system is inexpensive,

abundant, and nontoxic; moreover, unlike the MgCO3 system, it does not require a sorbent to capture
the working fluid when combined with a chemical heat pump.[37]The previous chapters mentioned that the use of dopants to increase the dehydration reactivity ofMg(OH)2 has already been studied.[23,24,26–33] However, the reason some dopants can increase thedehydration reactivity of Mg(OH)2 and the mechanism behind their ability to do so have not yet been
elucidated.
Several reports have been published on the dehydration mechanism of Mg(OH)2. McKelvy et al.
suggested the formation of an intermediate during the dehydration of Mg(OH)2 using in situ
transmission electron microscopy (TEM) observations, and reported that the dehydration of Mg(OH)2
formed oxides, hydroxides, and intermediate oxyhydroxides (Mgx+yOx(OH)2y, x oxide layers for every
2y hydroxide layers).[41] The oxide lamellae typically nucleate and grow from the parent hydroxide
lamella until complete dehydration.

During dehydration, the oxyhydroxides nucleate and grow, andsubsequently yield oxides. In addition, they revealed that two hydroxide layers provided two OH−
ions and formed one H2O molecule and one O2− ion.[41] Pimminger et al. evaluated the internal
nanostructure of Mg(OH)2 during dehydration and indicated that deformation occurred parallel to the(0001) planes of Mg(OH)2, which corresponded to the (111) planes of the newly formed MgO.[42]Mironyuk et al. indicated that the lattice constants a and c of Mg(OH)2 increased by 1.2% and 3.8%
respectively, and the Mg–O bond length of Mg(OH)2 increased from 0.2152 to 0.2194 nm, during
dehydration from 20 to 440 °C.[43]Moreover, according to thermogravimetry and XRD data,

dehydration in the temperature range of 440–900 °C caused the Mg–O bond length to decrease from
0.21157 to 0.21112 nm.[43] Their TEM results demonstrated that the dehydration of Mg(OH)2 caused
O atom defects, which led to the formation of MgOx (x < 1) species. In addition, a computational
method revealed that this structure corresponded to Mg4O3 with oxygen vacancies in the (1/2, 1/2, 0)
and (1/2, 1/2, 1) positions of the cell.[43] However, these analyses focused only on pure Mg(OH)2.
The author proposed the addition of LiCl and/or LiOH to Mg(OH)2 to improve dehydration
reactivity.[23,24,31,32] Chapter 3 indicated that LiCl and LiOH play different roles in the dehydration
reactivity of Mg(OH)2, and suggested that LiOH and LiCl promoted the decomposition of Mg(OH)267

at the surface and in the bulk phase, respectively.[32]Fourier-transform infrared (FT-IR) measurements can be used for qualitative analysis. Moulin etal. determined the molar adsorption coefficient of three types of methoxy species by in situ IR spectra
combined with volumetric and thermogravimetric methods.[44] Several researchers have investigated
the role of the OH groups on MgO and proposed the following stoichiometric equation for theadsorption and dissociation of a H2O molecule:[45–50]

2+− 𝑂𝑂𝐿𝐿2−
𝑀𝑀𝑀𝑀𝐿𝐿𝐿𝐿′ 𝐶𝐶 + 𝐻𝐻2 𝑂𝑂 ⇄𝑂𝑂𝑂𝑂− 𝐻𝐻 +
|
|
2+
2−
𝑀𝑀𝑀𝑀𝐿𝐿𝐿𝐿 −𝑂𝑂𝐿𝐿𝐿𝐿(1)

where LC and LʹC denote low-coordinate.
Knözinger et al. identified four types of OH groups on the MgO surface based on their coordination
numbers and the presence or absence of H bonds (types A–D, Figure 4-1).[47] Chizallet et al.[48] andChizallet et al.[50] proposed a new classification of the OH groups using a combination of measured
infrared (IR) spectra and density functional theory calculations (types Aʹ and Cʹ, Figure 4-1).
Moreover, Chizallet et al. investigated the effects of H bonds, surface morphology, and coordination
number on the frequency of the OH groups and associated the frequency with the type of OH group.[50]Figure 4-1. Four types of OH groups (A–D) defined by Knözinger et al.,[47] and two additional typesof OH groups (Aʹ and Cʹ) defined by Chizallet et al..[50] The description of these OH groups wascited from their papers[47,50] and a literature.[51] Here, LC and LʹC denote low-coordinated.

68Because only a few studies have been conducted on the decomposition of pure and doped Mg(OH)2
using IR measurements, our goal was to analyze the decomposition of Mg(OH)2 using previously
reported data on the characteristics of the OH groups on the MgO surface. In addition, this study
aimed to elucidate the effect of the addition of LiCl and/or LiOH on the dehydration reaction (heatstorage operation) of Mg(OH)2 using XRD and FT-IR analyses.

If the effect of Li compounds on thedehydration of Mg(OH)2 was elucidated, researchers should be able to identify other dopants that
could effectively enhance the reactivity of the Mg(OH)2/MgO system and apply it to other CHS
materials.

4.2. Experimental
4.2.1. Sample preparation
To prepare LiCl and LiOH–co-added Mg(OH)2 (LiCl/LiOH/Mg(OH)2) and LiCl- or LiOH-added
Mg(OH)2 (LiCl/Mg(OH)2 and LiOH/Mg(OH)2, respectively), LiCl·H2O (99.9% purity), LiOH·H2O,
and Mg(OH)2 (99.9% purity, 0.07 µm) were purchased from Wako Pure Chemical Industries, Ltd.
(Japan), and were used as precursors.
LiCl/Mg(OH)2, LiOH/Mg(OH)2, and LiCl/LiOH/Mg(OH)2 were prepared using an impregnation
method as described in previous chapters. LiCl-added Mg(OH)2 with a Mg(OH)2:LiCl molar ratioof 100:10, LiOH-added Mg(OH)2 with a Mg(OH)2:LiOH molar ratio of 100:20, and LiCl and LiOH–
co-added Mg(OH)2 with a Mg(OH)2:LiCl:LiOH molar ratio of 100:10:10 (hereafter L10, LO20, andL10/LO10, respectively) were prepared using this method.

For comparison, Li-free Mg(OH)2 andMgO (99.9%, 0.05 mm, Wako Pure Chemical Industries, Ltd., Japan) (hereafter, Mg(OH)2-W and
MgO-W, respectively) were prepared using the same method. The mixing ratios of all preparedsamples are listed in Table 4-1.

Because previous studies and Chapter 2 indicated that L10, LO20,

and L10/LO10 presented the most optimal mixing ratios with respect to dehydration and hydration
reactivity and the activation energy for dehydration,[24,31,32] the author compared their dehydration
reactivities in this study. In addition, the specific surface area and mean particle size of the prepared
Table 4-1. Mixing ratios of the prepared samples.
SampleMixing ratio [mole ratio]

Mg(OH)2-W

Mg(OH)2 without Li compounds

L10

Mg(OH)2:LiCl = 100:10

LO20

Mg(OH)2:LiOH = 100:20

L10/LO10

Mg(OH)2:LiCl:LiOH = 100:10:10

MgO-W

MgO without Li compounds
69

samples derived from BET measurement are shown in Table A5-1 in Appendix A5.

4.2.2. Dehydration reaction tests
All CaF2-added samples (hereafter sample+CaF2) were heated using a TGD–9600 series(ADVANCE RIKO, Inc., Japan) thermobalance. The samples (~20 mg) were added to Pt cells andwere subsequently heated from 20 to 25 °C to different temperatures in the range of 220–400 °C at a
rate of 10 °C min−1 under an Ar flow of 100 mL min−1. To perform FT-IR measurements quickly
after the dehydration test, the samples were mixed with CaF2 (sample-to-CaF2 mixing ratio of 0.3:0.6(w/w)) before the reaction test.

4.2.3. XRD characterization
XRD measurements were performed in air in the temperature range of 20–25 °C, in the 2θ rangeof 10−150°, at a scan rate of 1.0° min−1, and at a scan width of 0.01° as shown the previous chapters.
The XRD profiles of the prepared samples were used to confirm their successful preparation. To
avoid the subsequent hydration of the samples during measurements, the CaF2-added samples were
scanned after reaction at a scan rate of 10° min−1 using Cu Kα radiation at a voltage of 40 kV and acurrent of 40 mA. As the 2θ range of 10–80° was sufficient to characterize Mg(OH)2, the data were
described in this range.

Before the measurements, the prepared samples were dried at 120 °C

overnight to remove the physically adsorbed water.

4.2.4. FT-IR characterization
To gain additional information on the dehydration of pure and doped Mg(OH)2, FT-IR
measurements were performed using an FT/IR-4200 (JASCO Corp., Japan) spectrometer in diffuse
reflectance mode in air in the temperature range of 20–25 °C.

The background for the FT-IR

measurements was calibrated using CaF2 powder. The scans were obtained in the wavenumber rangeof 7800–350 cm−1, at a resolution of 4.0 cm−1, and all results were reported as averages of 64 scans.
Similar to the XRD measurements, the prepared samples were dried at 120 °C overnight prior to the
FT-IR measurements to remove any physically adsorbed water.

Because the author was mainlyinterested in the peaks of Mg(OH)2, in this study, the author focused on the FT-IR spectra in thewavenumber range of 5000–2750 cm−1. The effect of physical adsorption of water and CO2 could
be ignored because the spectra did not change upon exposure to air[52] (Figure S2-1 and S2-2 in the
Supporting Information).

In addition, we considered that the carbonation in the air did not affect thereactivity of the samples and experimental results because we did not observe the desorption of CO2
from our Mg(OH)2 samples in TG measurement[53] (Figure S2-3 and S2-4 in the Supporting
Information).
The samples premixed with CaF2 were used for the FT-IR measurements. After dehydration tests
70

Figure 4-2. Experimental procedure for the analysis of the dehydrated samples. Here FT-IR, XRD,
and TG denote Fourier-transform infrared, X-ray diffraction, and thermogravimetry, respectively.
at different temperatures, the samples premixed with CaF2 were quickly placed on a diffuse reflectanceunit and their FT-IR spectra were obtained in air in the temperature range of 20–25 °C.

A flow chartof the experimental procedure is illustrated in Figure 4-2.

4.3. Results
4.3.1. XRD characterization of the prepared samples
As previously reported, the peaks of brucite Mg(OH)2 were observed in the XRD patterns of all
prepared samples (before the addition of CaF2) (Figure 4-3 (a)), and indicated the predominance of
brucite Mg(OH)2 in all samples, including those with added LiCl and/or LiOH.[24,31,32]

Thecharacteristic peaks of LiCl were not detected in the LiCl-added samples (L10, and L10/LO10),
probably because the LiCl species were dissolved in the Mg(OH)2 phase, were well dispersed
throughout the Mg(OH)2 particles, and were converted into aqueous solution, which rendered them
undetectable, or because a LiCl and Mg(OH)2 solid solution formed.[24,32] In addition, Figure 4-3(a) does not detect the peaks of LiOH and LiOH·H2O; however, the author confirmed those peaks
previously reported,[31,32] the author concluded that LiOH or LiOH·H2O could exist on the surfaces of
the LiOH-added samples (LO20, and L10/LO10).

LiOH might be partially converted to Li2CO3

during preparation, which confirmed that LiOH was typically contaminated with Li2CO3,[32,54] andaccounted for the presence of the Li2CO3 peak in the XRD profiles of the LiOH-added samples.
71Figure 4-3. X-ray diffraction patterns of (a) the prepared samples, and (b) the CaF2-added samples
before dehydration. Here, Mg(OH)2-W denotes Mg(OH)2 pretreated using the impregnation method
without Li compounds, L10 denotes LiCl-added Mg(OH)2 with a Mg(OH)2:LiCl molar ratio of 100:10,
LO20 denotes LiOH-added Mg(OH)2 with Mg(OH)2:LiOH molar ratio of 100:20, and L10/LO10
denotes LiCl and LiOH–co-added Mg(OH)2 with a Mg(OH)2:LiCl:LiOH molar ratio of 100:10:10.
The peaks of Mg(OH)2 observed in the XRD profiles of the prepared samples (Figure 4-3 (a))
were still present in the XRD patterns of the CaF2-added samples (Figure 4-3 (b)). Conversely, thepeaks of LiOH, LiOH·H2O, and Li2CO3 were not detected in the XRD profiles of the samples,
probably because the intensities of those peaks were too weak compared to that of CaF2. These
results confirmed that the addition of CaF2 to the samples did not affect the sample structure.

4.3.2. FT-IR characterization
4.3.2.1. FT-IR characterization of the prepared samples
Figure 4-4 (a) presents the FT-IR spectra of the prepared samples (before the addition of CaF2).
The peak at 3695 cm−1 was assigned to the OH groups on the surface, at the edge, or in the corner sites
with low coordination numbers.[45–48,50]

Knözinger et al. assigned this peak to the isolated

monocoordinated OH groups.[47] In contrast, Chizallet et al.[48] and Chizallet et al.[50] assigned thispeak to the H-bonded monocoordinated OH groups that served as acceptors. Typically, the broad
peak in the range of 3600–3200 cm−1 was assigned to the H-bonded OH groups that served as
donors.[47,48,50]

This broad peak offered information on the state of the H2O molecules.[46]The

intensities of the broad peak in the spectra of L10 and L10/LO10 were significantly higher than thoseof the corresponding peak in the spectra of the other two samples, which indicates that H2O molecules
were adsorbed on the sample surface.[55–57]

This broad peak can be attributed to the highhygroscopicity of LiCl, indicating that LiCl easily adsorbed water to form LiCl hydrate species.
Therefore, LiCl hydrate species could form on the surface of Mg(OH)2 (no LiCl of LiCl hydrate
species was observed in the XRD profiles of the samples).
72

However because the author could notperform in situ measurements, the author did not analyze the behavior of this peak. The peaks at
4308, 4033, 3950, and 3794 cm−1 can be assigned to the OH groups of the bulk phase of Mg(OH)2, as
explained in detail in Section 4.4.1. of FT-IR peak assignment.
To investigate the states of LiCl and LiOH in the samples, the author performed several supporting
characterizations.

The SEM images of the samples could not distinguish Mg(OH) 2 and Li

compounds (Figure S2-5). To determine the state of LiOH in LO20, the author prepared a LO20
sample by physically mixing LiOH and Mg(OH)2 in an agate mortar. The FT-IR spectra of the LO20
sample prepared using the impregnation method, LO20 sample prepared by physically mixing LiOH
and Mg(OH)2, and LiOH·H2O are presented in Figure S2-6.

Both LO20 samples were dried at

120 °C overnight before measurements, as mentioned in the experimental section. The peaks at3676 and 3564 cm−1 in Figure S2-6 were assigned to the OH groups of LiOH and LiOH·H2O,
respectively. The peak at 3564 cm−1 was also detected in the spectrum of the physically mixed LO20
sample, which indicated that LiOH·H2O was present on the Mg(OH)2 surface even after drying at
120 °C. When the LiOH·H2O phase was deposited on the Mg(OH)2 surface, the absorption peaksof LiOH·H2O and Mg(OH)2 should be detected separately.

However, the peak of LiOH·H2O was

not detected in the spectrum of the LO20 sample prepared using the impregnation method. These
results indicate that a solid solution of Mg(OH)2 and LiOH or LiOH·H2O formed on the surface ofMg(OH)2 in the LO20 samples. This solid solution hindered the detection of the characteristic peaksof LiOH or LiOH·H2O in the XRD profile of LO20, and probably accounted for the absence of thecharacteristic peak of LiCl. The possible formation of solid solution between LiCl and Mg(OH)2 is
supported by the XRD peak shift in the Supporting Information (Figure S2-7 (a) and (b)) because
there might be no peak shifts if no solid solution was produced by the substitution of a Li+ ion to a
Mg2+ site.[31,32,58] Even a small difference in the peak position might have influenced the reactivityof the Mg(OH)2/MgO system because the effective ionic radii of Li+ (0.76 Å) and Mg2+ (0.72 Å)[59]Figure 4-4. Fourier-transform infrared spectra of (a) all prepared samples, and (b) all CaF2-added
samples before dehydration. The mixing ratio of the samples is shown in the experimental section.
73

Figure 4-5. Fourier-transform infrared spectra of (a) Mg(OH)2-W+CaF2, (b) L10+CaF2, (c)LO20+CaF2, (d) MgO-W+CaF2, and (e) L10/LO10+CaF2 after dehydration under an Ar flow of
100 mL min−1. (c)′ and (d)′ are the expanded spectra of (c) and (d), respectively.
74are similar and the number of substituted ions cannot be significant. Hence, the author hypothesizedthat a LiCl and Mg(OH)2 solid solution could form in L10. However, the author need to analyze thestate of LiCl in detail using other techniques such as computational methods.
The same peaks observed in the FT-IR spectra of the CaF2-free samples (Figure 4-4 (a)) weredetected in the spectra of the CaF2-added samples (Figure 4-4 (b)), which indicated that the additionof CaF2 did not affect the sample structure.

4.3.2.2. FT-IR analysis of samples after dehydration
Figure 4-5 (a) presents the FT-IR spectra of Mg(OH)2-W+CaF2 after heating to different
temperatures in the range of 310–400 °C under an Ar flow. The peaks observed before the reaction(Figures 4-4 (a) and (b)) were also detected in the samples heated at temperatures below 310 °C.
Only the peaks of Mg(OH)2 and CaF2 were observed in the XRD profile of the sample heated at 310 °C,
which indicated that the decomposition of Mg(OH)2 hardly progressed (Figure S2-8).
The diffraction peaks of MgO were detected in the XRD profile of the Mg(OH)2-W+CaF2 sample
heated at 320 °C (Figure S2-8, Table 4-2). The absorption intensity of the FT-IR peaks decreased
with increasing temperature, and the absorption bands at 3788, 4315, 3953, and 4036 cm−1 disappeared
at 360, 370, 370, and 380 °C, respectively (Table 4-2).

Furthermore, the peaks of Mg(OH)2

disappeared from the XRD pattern of the Mg(OH)2-W+CaF2 sample heated at 370 °C. This behavior
was consistent with that reported in Chapter 3,[32] and indicated that the transition from Mg(OH)2 to
MgO was almost complete at 370–380 °C.

Moreover, the author concluded that the peaks at

approximately 4300–3790 cm−1 (Figure 4-5 (a)) reflected the state of the bulk phase rather than thatof the surface phase. Table 4-2 summarizes the observed FT-IR peak positions and the corresponding
heating temperatures for Mg(OH)2-W+CaF2.
The intensity of the peak at 3695 cm−1, which was assigned to the isolated OH groups on the
Mg(OH)2 surface, significantly decreased when the Mg(OH)2-W+CaF2 samples were heated in thetemperature range of 360–370 °C. Furthermore, the intensity decreased and the peak broadened forthe samples heated in the temperature range of 370–380 °C (Figure 4-5 (a)).

A similar behavior

was observed in another study.

[47]

Although the position of the peak at 3695 cm−1 was unchanged until 360 °C, the peak shifted tomuch higher wavenumber regions when the temperature exceeded 370 °C (Table 4-2). This shift
was attributed to the decrease in the coordination number of the OH groups, Mg–O distance, andnumber of hydrogen bonds.[47,50]
Figure 4-5 (b) illustrates the FT-IR spectra of L10+CaF2 after heating at different temperatures in
the range of 260–400 °C. The peak of MgO was detected in the spectrum of the sample heated at
270 °C, and the peak of Mg(OH)2 disappeared from the XRD profile of the sample heated at 320 °C(Figure S2-9). These results agreed well with those observed in Chapter3.[32] The intensities of75

the FT-IR peaks decreased with increasing temperature, and the absorption bands at 4307, 3787, 4034,
and 3950 cm−1 disappeared at 310, 310, 330, and 330 °C, respectively (Table 4-3). The intensity ofthe peak at 3696 cm−1, which was assigned to the isolated OH groups on the Mg(OH)2 surface,significantly decreased for the samples heated in the temperature range of 320–340 °C; moreover, thepeak broadened (Figure 4-5 (b)). These changes were observed at lower temperatures than thosereported for Mg(OH)2-W+CaF2, which indicated that the addition of LiCl to Mg(OH)2 promoted its
dehydration.[23,24,32]
The sharp peak at approximately 3700 cm−1 was shifted to lower wavenumbers after the sample
was heated at 350 °C and 400 °C (Table 4-3). This behavior was different from that of Mg(OH)2W+CaF2, and indicated the formation of H bonds and an increase in the coordination number of the
OH groups.[47,50] This can be attributed to the aggregation of the sample and the formation of bridges
between surface OH groups and Mg2+ or O2− ions, which led to an increase in the coordination numberof the surface OH groups to more than one.[48] Therefore, the surface state of Mg(OH)2 was changedby the addition of LiCl.
Figure 4-5 (c) presents the FT-IR spectra of LO20+CaF2 after heating at 300–400 °C. The peaksof MgO were detected at 310 °C, and the peaks of Mg(OH)2 disappeared from the XRD profile of thesample at 350 °C (Figure S2-10). The temperature at which the diffraction peaks of the Mg(OH)2phase disappeared (350 °C), was in agreement with that observed in Chapter3.[32] However, thetemperature at which the diffraction peaks of the MgO phase appeared (310 °C) was 70 °C higher thanreported in Chapter 4(Table 4-4 and Figure S2-10).[32] This could be attributed to the peak intensityof CaF2 being much stronger than that of MgO. Therefore, the peaks of MgO were undetectable (inthis study, a MgO phase was only present in the XRD profile of LO20 after heating at 240 °C, and thistemperature was in good agreement with the results in Chapter3[32]).

Because this study focused on

the behavior of the peaks of the OH groups, this effect will not be analyzed in detail.
The absorption bands at 3789, 4312, 4034, and 3956 cm−1 disappeared at 340, 350, 360, and 360 °C,
respectively (Table 4-4). Moreover, the intensity of the peak at 3696 cm−1 significantly decreased
and peak broadened in the heating temperature range of 350–370 °C (Figure 4-5 (c)).

Thesebehaviors occurred in slightly lower temperature ranges than those observed for Mg(OH)2-W+CaF2.
A shoulder peak at approximately 3760 cm−1 was observed at 340 °C, and its intensity increased at
higher temperatures (Figure 4-5 (c)ʹ).
Moreover, Figure 4-5 (c)ʹ illustrates that the peak at 3696 cm−1 was split into two components at
temperatures higher than 370 °C, which was in agreement with the data reported in a previous paper.[47]Knözinger et al. assigned the component at higher and lower wavenumbers to isolated
monocoordinated OH groups (type A, Figure 4-1), and isolated multicoordinated (type B, Figure 41) or H-bonded monocoordinated acceptor OH groups (type C, Figure 4-1), respectively.[47]To assign these split peaks and the shoulder peak at 3756 cm−1, a MgO-W sample was heated to76different temperatures.
The FT-IR spectra of the MgO-W+CaF2 samples heated at different temperatures in the range of
310–400 °C are presented in Figure 4-5 (d).

Because MgO-W was prepared using a typicalimpregnation method, Mg(OH)2 was partially formed (Figure S2-13). A shoulder peak at 3766 cm−1
was observed in the FT-IR spectrum of the sample heated at 310 °C (Figure 4-5 (d)), which was
similar to that observed in the spectrum of LO20+CaF2 (Figure 4-5 (c)). The author confirmed thatthe intensity of the peak at 3702 cm−1 significantly decreased at 370 °C. Moreover, after the sample
was heated at 380 °C, the absorption peak was split, as illustrated in Figures 4-5 (d) and (d)ʹ.
Figure 4-5 (d)ʹ shows the expanded FT-IR spectra of MgO-W+CaF2 after heating in thetemperature range of 360–400 °C. As illustrated in Figures 4-5 (c) and (c)ʹ, the shoulder peak at
3755 cm−1 and split peaks at 3700 cm−1 can be assigned to MgO species. The behavior of MgOW+CaF2 and the reasonable assumption that the decomposition of Mg(OH)2 proceeds from thesurface[60,61] indicated that the surface of LO20+CaF2 likely consisted of MgO compared with
Mg(OH)2-W+CaF2 and L10+CaF2 after dehydration. Therefore, these results can explain why, in
Chapter 3, a MgO phase was detected in the XRD profile of LO20 at a low temperature of 240 °C.[32]Figure 4-5 (e) illustrates the FT-IR spectra of L10/LO10+CaF2 after heating at various
temperatures in the range of 260–400 °C. The peak of MgO was observed at 270 °C, and the peakof Mg(OH)2 disappeared from the XRD profile of this sample at 310 °C (Figure S2-12 and Table 47). The temperature at which the diffraction peaks of the Mg(OH)2 phase disappeared (310 °C), wasin agreement with that observed in Chapter 3. However, the temperature at which the diffraction
peaks of the MgO phase were observed (270 °C), was 40 °C higher than that reported in Chapter 3(Figure S2-12 and Table 4-7).[32]This can be caused by the aforementioned significant peakintensity of CaF2 in the XRD profile of the sample.
The absorption bands at 4304, 4033, 3956, and 3797 cm−1 disappeared at 310, 320, 320, and 320 °C,
respectively (Table 4-6). Moreover, the intensity of the peak at 3695 cm−1 significantly decreased
and the peak broadened in the temperature range of 320–340 °C (Figure 4-5 (e)). These behaviors
occurred at lower temperatures than those observed for Mg(OH)2-W+CaF2. The shoulder and split
peaks observed in the FT-IR spectra of LO20+CaF2 and MgO-W+CaF2 (Figures 4-5 (c)ʹ and (d)ʹ)
were not detected in the spectra of L10/LO10+CaF2.

This might be ascribed to the insufficientamount of LiOH. The changes in the spectra of L10/LO10+CaF2 were similar to those in the spectraof L10+CaF2 (Figure 4-5 (b)).
The sharp peak at 3695 cm−1 shifted to higher wavenumber regions after heating from 290 to
340 °C (Table 4-6). Conversely, when the sample was heated at 350–400 °C, the absorption peakshifted to a lower wavenumber (3683 cm−1; Table 4-6).

This behavior was similar to that of

L10+CaF2. Therefore, the surface state of Mg(OH)2 in the L10/LO10+CaF2 sample was very close
to that in the L10+CaF2 sample.
77

Figure 4-6. Second derivative of the intensities of the peaks at 3700 cm−1 of all samples. The mixing
ratio of the samples is shown in the experimental section.
Figure 4-6 presents the changes in the intensity of the peak at approximately 3700 cm−1 in the
spectra of all samples using the second derivatives for all samples (the second derivative of theintensity of the peaks of Mg(OH)2-W+CaF2 after heating to 310 °C is illustrated in Figure S2-14.).
The y- and x-axes represent the second derivative of the intensity of the peak at 3700 cm−1 and heating
temperature, respectively. The results revealed that the intensities of the peaks of Mg(OH)2-W+CaF2,
L10+CaF2, LO20+CaF2, and L10/LO10+CaF2 were leveled off at ~380, 340, 370, and 340 °C,
respectively.

Considering the changes in peak shape, these results indicate that the transition from

Mg(OH)2 to MgO was almost complete when the peak intensity was leveled off (Figures 4-5 (a), (b),(c), and (e), respectively).
Table 4-2–4-7 summarizes the temperatures at which the Mg(OH)2 phases disappeared and the
MgO phases were first observed in the XRD profiles of the samples, and the temperatures at which
the second derivatives of the intensity of the peak at 3700 cm−1 in the FT-IR spectra of the samples
leveled off.

78

Table 4-2. Fourier-transform infrared (FT-IR) peaks detected after dehydration at different heating
temperatures for Mg(OH)2-W+CaF2. The circles indicate that a phase was detected in the X-ray
diffraction (XRD) pattern.
Temperature

Peak position [cm−1]

[°C]

XRD:

XRD:

Mg(OH)2

MgO

400

3712

○

390

3714

○

380

3711

○

3698

○

370

4027

3696

○

○

3788

3696

○

○

3954

3792

3696

○

○

4034

3954

3798

3696

○

○

4308

4033

3953

3791

3695

○

○

4315

4036

3953

3788

3695

○

360

4303

4033

3952

350

4308

4033

3954

340

4308

4035

330

4308

320
310

Table 4-3. FT-IR peaks detected after dehydration at different heating temperatures for L10+CaF2.
The circles indicate that a phase was detected in the XRD pattern.
Temperature

Peak position [cm−1]

[°C]

XRD:

XRD:

Mg(OH)2

MgO

400

3681

○

350

3677

○

340

3698

○

330

3697

○
○

320

4032

3948

3696

310

4035

3950

3696

○

○

300

4304

4035

3952

3793

3696

○

○

290

4303

4034

3950

3792

3696

○

○

280

4304

4032

3950

3791

3696

○

○

270

4305

4034

3956

3793

3696

○

○

260

4307

4034

3950

3787

3696

○

79

Table 4-4. FT-IR peaks detected after dehydration of LO20+CaF2 at different heating temperatures.
The circles indicate that a phase was detected in the XRD pattern.
Temperature

Peak position [cm−1]

[°C]

XRD:

XRD:

Mg(OH)2

MgO

3712

3688

○

3720

3688

○

3717

3688

○

370

3714

3697

○

360

3697

○
○

400

3756

390
380

3757

350

4029

3952

3696
3697

○

○

340

4301

4034

3856

330

4306

4032

3954

3797

3696

○

○

320

4306

4037

3954

3788

3696

○

○

310

4313

4034

3955

3789

3696

○

○

300

4304

4036

3955

3790

3696

○

Table 4-5. FT-IR peaks detected after dehydration of MgO-W+CaF2 at different heating temperatures.
The circles indicate that a phase was detected in the XRD pattern.
Temperature

XRD:

Peak position [cm−1]

[°C]
400

3755

390
380

3749

370

MgO
3709

3686

○

3711

3689

○

3711

3689

○

3701

○

360

3753

3702

○

350

3751

3702

○

3702

○

3765

3702

○

340

4031

330

4306

4034

320

4306

4034

3950

3766

3702

○

310

4307

4033

3951

3766

3702

○

80

Table 4-6. FT-IR peaks detected after dehydration of L10/LO10 at different heating temperatures.
The circles indicate that a phase was detected in the XRD pattern.
Temperature

Peak position [cm−1]

[°C]

XRD:

XRD:

Mg(OH)2

MgO

400

3683

○

390

3677

○

380

3687

○

370

3678

○

360

3677

○

350

3679

○

340

3699

○

330

3697

○

320

3696

○
○

310

4035

3950

3799

3696

300

4313

4036

3951

3787

3696

○

○

290

4310

4034

3952

3800

3696

○

○

280

4312

4036

3954

3799

3695

○

○

270

4304

4033

3950

3794

3695

○

○

260

4304

4033

3956

3797

3695

○

Table 4-7. The temperatures at which the XRD peaks of Mg(OH)2 were no longer detected, the peaksof MgO appeared, and the second derivative of the intensity of the FT-IR peak leveled off (based on
the data in Figures S4-8–S12 and 4-6).
SampleTemperature at which theTemperature at which theTemperature at which

Mg(OH)2 peak was no

MgO

first

the intensity of the FT-

longer detected in theobserved in the XRD

IR peak at 3700 cm−1

XRD pattern [°C]

pattern [°C]

became saturated [°C]

peak

was

Mg(OH)2-W

370–380

310–320380

L10

310–320

260–270340

LO20

340–350

300–310370

L10/LO10

300–310

260–270340

81

4.4. Discussion
4.4.1. FT-IR peak assignment
The sharp peak at 3700 cm−1 was assigned to the low-coordinated OH groups on the edge and/or
at corner sites.[45–48,50] Because this peak was observed before dehydration, the author inferred that
it was derived from the Mg(OH)2 surface. The peaks at 4300, 4030, 3950, and 3790 cm−1 (Figures
4-5 (a)–(e)) disappeared at around the temperature at which the diffraction peaks of the Mg(OH)2phase disappeared. Therefore, it is reasonable to consider that these peaks were related to the OH
groups in the bulk Mg(OH)2 phase (Tables 4-2–6).

Moreover, the peaks at 4030 and 3950 cm−1

(Figures 4-5 (a)–(e)) disappeared at higher temperatures than the other two peaks, which indicatedthat they were related to the OH groups of the more stable bulk Mg(OH)2 phase (Tables 4-2–6).
Conversely, the sharp peak at 3700 cm−1 did not disappear, even at 400 °C. However, at higher
temperatures, the peak was broader (Figure 4-5 (a)), which indicated that the peak should be assigned
to MgO. The components of this peak at 3700 cm−1 were not only the OH groups on the edge and/or
at corner sites but also at kink and divacancy sites.[48–50]

Because computational methods revealed

that the thermal stability of the OH groups at the kink and divacancy sites were high,[49,50] the peak at
3700 cm−1 could be observed at higher temperatures.
The split peaks at higher and lower wavenumbers in the FT-IR spectra of LO20+CaF2 and MgOW+CaF2 after heating (Figures 4-5 (c)ʹ and (d)ʹ) can be assigned to the isolated monocoordinated OH
groups (type A in Figure 4-1) and isolated multicoordinated (type B in Figure 4-1) or H-bonded
monocoordinated acceptor OH groups (type C in Figure 4-1), respectively.[47,50] Therefore, these
split peaks were ascribed to MgO species.

In addition, the shoulder peak at 3766–3755 cm−1

(Figures 4-5 (c)ʹ and (d)ʹ) can also be assigned to MgO species.
It was reported that an intermediate was formed during the transition of Mg(OH)2 to MgO.[41–43]
In the temperature range at which the intensity of the peak at 3700 cm−1 decreased significantly, thepeak shape also changed noticeably (the peak broadened). This behavior was observed at 360–380,
320–340, 360–380, and 320–340 °C for Mg(OH)2-W+CaF2 (Figure 4-5 (a)), L10+CaF2 (Figure 4-5(b)), LO20+CaF2 (Figure 4-5 (c)), and L10/LO10+CaF2 (Figure 4-5 (e)), respectively, and the results
suggested that an intermediate formed during the transition of Mg(OH)2 to MgO.
Figure 4-7 illustrates the peak assignments based on the aforementioned results, and Figures 4-8(a) and (b) present the intensity ratios determined using the second derivatives of the intensities of thepeaks at 3700 and 4030 cm−1 for each sample. The intensity ratio was calculated as follows:

𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼𝐼 𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟 =𝐼𝐼3700,𝑇𝑇

𝐼𝐼3700,𝑃𝑃𝑃𝑃𝑃𝑃

𝐼𝐼

(2)

𝑜𝑜𝑜𝑜 𝐼𝐼 4030,𝑇𝑇

4030,𝑃𝑃𝑃𝑃𝑃𝑃

82

where I3700 and I4030 are the second derivatives of the intensities of the peaks at 3700 and 4030 cm−1,
respectively; T is the heating temperature; and I3700, Pre and I4030, Pre are the intensities of the peaks at
3700 and 4030 cm−1 of the prepared samples, respectively.

As the peak at 4030 cm−1, which could

be ascribed to the OH groups of the stable bulk Mg(OH)2 phase, was observed at higher temperatures,
as mentioned above, the intensity ratio of the peak at 4030 cm−1 was assumed to be representative of
the bulk state.

Figure 4-7. Fourier-transform infrared spectra of Mg(OH)2-W+CaF2.

Peak assignments were basedon this and previous studies.

45–48,50

Figure 4-8. Intensity ratio based on the second derivative of the intensities of the peaks at (a) 3700
cm−1 and (b) 4030 cm−1 of the Fourier-transform infrared spectra (Figures 4-5 (a)–(c), and (e)). The
mixing ratio of the samples is shown in the experimental section.

83

Figures 4-8 (a) and (b) provide the following insights.

The L10+CaF2, LO20+CaF2, andL10/LO10+CaF2 samples were dehydrated at temperatures lower than the Mg(OH)2-W+CaF2 sample.
This trend was in agreement with that observed in Chapter 3.[32] The ratio of the intensities of thesecond derivative of the peak at 4030 cm−1 decreased rapidly at low temperatures compared with thatof the peak at 3700 cm−1.

Therefore, as the heating temperature increased, the decrease in theintensities of the peaks ascribed to the bulk phase was larger than those of the peaks attributed to thesurface phase, which indicated that the bulk decomposition of Mg(OH)2 was prioritized over that at
the surface over a certain temperature range. During the initial step of the dehydration process, thedehydration of Mg(OH)2 occurred at the surface.[60] Subsequently, the bulk phase was decomposed,
followed by the dissociation of the OH groups at the kink and divacancy sites, which are thermallystable.[49]4.4.2. Effects of LiCl and/or LiOH addition on the dehydration reactivity ofMg(OH)2
The previous studies reported[23,24,31,32] that the addition of LiCl and/or LiOH to Mg(OH)2
promoted its dehydration. However, the effects of the addition of LiCl and/or LiOH to Mg(OH)2 on
its dehydration have not been elucidated yet. Therefore, the author analyzed the effect of the additionof Li compounds (LiCl and LiOH) to Mg(OH)2 on its dehydration.
The author confirmed that the Li-added Mg(OH)2 samples were dehydrated at temperatures lower
than Mg(OH)2-W using both XRD and FT-IR measurements.

In particular, L10+CaF2 andL10/LO10+CaF2 were dehydrated at low temperatures compared to the other samples (Figures 4-5(b) and (e), S4-9, and S4-12).

Sivasankari et al. analyzed the crystal structure of MgO doped withalkali metals (Li, Na, and K) and calculated the lattice parameters of MgO using the shift in the XRD
peak position caused by the alkali metal dopants.[58]substituted the Mg2+ sites.

Their results indicated that the alkali ionsShkatulov et al.[26] and Shkatulov and Aristov[27] investigated thedehydration of Mg(OH)2 with different added salts, and proposed that NO3− ions substituted OH− ions,and the cation impurity defects introduced by different valence cations (Na+) were caused by theaddition of NaNO3 to Mg(OH)2, which affected the dehydration of Mg(OH)2.

In Chapter 3, theauthor demonstrated that the lattice parameter changed (parameter c increased) after the addition of
LiCl and/or LiOH to Mg(OH)2, and this change might have promoted the dehydration of Mg(OH)2.[32]Although the difference in the lattice parameters and peak shift were slight in those studies, Li
substitution can be inferred.

Moreover, the author confirmed that a solid solution of LiOH or

LiOH∙H2O and Mg(OH)2 could form (Figure S2-6). A solid solution between LiCl and Mg(OH)2
could also be synthesized (see Section 4.3.2.1. in detail).

Therefore, assuming that the Li+ ionssubstituted the Mg2+ ion sites (because the effective ionic radii of Li+ (0.76 Å) and Mg2+ (0.72 Å)[59]
are similar),[31,32,58] the difference in valence between the Li+ and Mg2+ ions could lead to lattice defects
84such as oxygen vacancies or OH groups defects.[26,27]

In addition the hydration enthalpy of Li+(approximately –500 kJ mol−1) was much lower than that of Mg2+ (approximately −1900 kJ
mol−1).[62,63] Therefore, Li+ substitution could favor the release of water, and dehydration would bepromoted at the Li+-substituted sites.
The addition of LiCl to Mg(OH)2 did not have the same effect as the addition of LiOH.[32]Chapter 3 indicated that LiOH promoted the decomposition of Mg(OH)2 at the surface and LiCl
mainly affected dehydration in the bulk phase.[32]Therefore, LiCl and LiOH probably affected

dehydration at different sites.
A shoulder peak at 3760 cm−1 and split peaks at 3700 cm−1 were observed in the FT-IR spectra of
the heated LO20+CaF2 and MgO-W+CaF2 samples (Figures 4-5 (c)ʹ and (d)ʹ, respectively). Split
peaks have been previously observed by Knözinger et al. in their study on the role of the OH groups
on the MgO surface during dehydration.[47] Thus, the surface state of LO20+CaF2 was closer to thatof MgO-W+CaF2 than to those of the other samples. The peaks of the Mg(OH)2 phase in the XRD
profile of L10+CaF2 disappeared at lower temperatures than those of Mg(OH)2-W+CaF2 andLO20+CaF2 (Table 4-7).
LO20.

[32]

Moreover, L10 was dehydrated at a markedly lower temperature than

Therefore, the author concluded that the addition of LiOH mainly affected the dehydrationof the Mg(OH)2 surface.

Furthermore, the author inferred that the addition of LiCl to Mg(OH)2

possibly affected the dehydration of the bulk Mg(OH)2 phase, and the effect of LiCl was stronger thanthat of LiOH.[32] Therefore, the dehydration of both the bulk and surface phases of L10/LO10 could
be promoted by the addition of LiCl and LiOH, which led to a better dehydration reactivity than othersamples.[32]By combining the aforementioned results, the author proposed models for the dehydration ofMg(OH)2 with added LiCl and/or LiOH. Figures 4-9 (a) and (b) show the mechanism of the effectof the addition of LiCl or LiOH on the Mg(OH)2 structure (the Mg(OH)2 structure was cited fromother articles).[41,64,65] For LiOH/Mg(OH)2, only Li+ ions affected dehydration because its pair ion
was OH−. The FT-IR results (Figures 4-5 (c)ʹ and (d)ʹ) revealed that substituted Li+ ions existed on
the surface of Mg(OH)2, and unsaturated sites could be generated via Li+ ion substitution because thecationic valences of the Mg2+ and Li+ ions were different (Figure 4-9 (a)).[31,32] This could lead tothe formation of OH− defects to compensate for the valence difference. McKelvy et al. reported that
two hydroxide layers provided two OH− ions, forming one H2O molecule and one O2− ion.[41] Thus,
the OH groups in the upper surface layer can be combined with adjacent OH groups to form H2O
molecules, whereas the lower OH groups in the surface layer formed H2O molecules with the upper
OH groups in the second layer (Figure 4-9 (a)).

85

Figure 4-9. Possible scheme for the effect of the addition of (a) LiOH and (b) LiCl on the dehydrationof Mg(OH)2.
At the surface of LiCl/Mg(OH)2, the number of Li+ ions that could substitute the Mg2+ sites was
smaller than the number of Li+ ions at the surface of LiOH/Mg(OH)2.

In addition, the author

anticipated that Cl− ions affected the dehydration of Mg(OH)2. The ionic radius of the Cl− ion (1.72
Å) is much larger than that of the OH− ion,[66] and generates distorted sites, leading to the penetrationof Li+ ions into deeper sites (bulk phase), as illustrated in Figure 4-9 (b). Further, the addition of
LiBr promoted the dehydration of Mg(OH)2, as shown in Figure S2-15 and the previous studies (Br−:1.88 Å[66]).[28–30]

Hence, the counter halogen anion can affect the dehydration.
− anions can be expected.

distortion sites due to Cl

Two different

Case (1) shows that Cl− ions can exist as Cl− on

the Mg(OH)2 surface. Case (2) is that Cl− ions can substitute the OH− sites. These distorted sites
can be generated singly or simultaneously (Figure 4-9 (b)).

Thus, Li+ substitution presumably

occurred in the bulk phase and promoted dehydration in the bulk phase. Dehydration can also be
promoted at sites close to Li+ ions and OH− defects because the hydration enthalpy of Li+ ions is lower
than that of Mg2+ ions,[62,63] and the H bonds can weaken (Figure 4-9 (b)).

Moreover, Cl− ions appear

to interact less with water owing to the calculated hydration enthalpies of Cl− and OH− being −390
and −410 kJ mol−1, respectively.[62] For LiCl/LiOH/Mg(OH)2, these effects can occur simultaneously
at the surface and in the bulk. Hence, the dehydration of LiCl/LiOH/Mg(OH)2 occurred at the lowest
temperatures of all samples (Table 4-7).[32]Shkatulov and Aristov analyzed the effects of additives on the dehydration of Mg(OH)2 andCa(OH)2.[27] effects of chlorides, acetates, and nitrates on the dehydration of Mg(OH)2 were similar,and Li compounds significantly lowered the dehydration onset temperature.

The hydration

enthalpies of Li+, Na+, and K+ were approximately −500, −400, and −300 kJ/mol,[62] respectively.

In+, Na+, and K+ through Nafion-117

addition, Stenina et al. determined the diffusion coefficients of Li

membranes using conductivity measurements and revealed that they decreased as follows: Na+ > Li+> K+.[67] Lee et al. also investigated the mobility of alkali metal ions using molecular dynamics
simulations and determined that the mobilities at infinite dilution in water at 25 °C increased asfollows: K+ > Na+ > Li+.[68] The author also investigated the relationship between the pH of the86

sample solution and dehydration reactivity (Figure S2-15).
clear relationship between them.

However, the author could not find any

Therefore, these factors did not appear to directly affectdehydration reactivity; however, the similar ionic radii of Li+ (0.76 Å) and Mg2+ (0.72 Å)[59] could be
the main factor for the increase in dehydration reactivity. Because the ionic radii of Na+ (1.02 Å) andK+ (1.38 Å) ions[59] were much larger than those of Mg2+, Na+, and K+ ions cannot substitute Mg2+
sites efficiently.[27]This study demonstrated that the addition of LiCl and/or LiOH promoted the dehydration ofMg(OH)2, and the effects of LiCl and LiOH on the dehydration of Mg(OH)2 were different. However,
the author only focused on the dehydration of Mg(OH)2, and a future study should be conducted toanalyze the effects of LiCl and LiOH on the hydration of MgO.

4.5. Conclusions
In this study, we monitored the effect of the addition of LiCl and/or LiOH on the dehydration ofMg(OH)2 using FT-IR and XRD measurements. The author confirmed that Mg(OH)2 with added
LiCl and/or LiOH was dehydrated at lower temperatures than pure Mg(OH)2, and these results are inaccordance with those reported in the previous studies.[23,24,31,32] The FT-IR spectra of L10+CaF2 andLO20+CaF2 after heating revealed that the effects of LiCl and LiOH on the dehydration of Mg(OH)2
were different (Figures 4-5 (b) and (c), respectively).

In particular, the results of this study indicatedthat LiOH and LiCl affected the dehydration of the surface and bulk phases, respectively.

However,

the author need further studies to determine the state of LiCl on the Mg(OH)2, especially the role ofCl− anions more clearly.

Moreover, the author also monitored the dehydration reaction of pure andchemically modified Mg(OH)2 using FT-IR and found some significant differences in the FT-IR
spectra. To our knowledge, few FT-IR spectra during the decomposition of Mg(OH)2 have beenreported to date. In addition, the author obtained the FT-IR spectrum of bulk Mg(OH)2.

For further

development of our study, it is necessary to clarify which bulk states the FT-IR peaks at 4300, 4030,
3950, and 3790 cm−1 represent.
Further studies are required to investigate the effect of LiCl and LiOH on the hydration reactivityof MgO.

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92

Chapter 5. Fourier-Transform Infrared and X-ray Diffraction
Analyses of the Hydration Reaction of Pure Magnesium Oxide
and Chemically Modified Magnesium Oxide

ABSTRACTThe magnesium hydroxide/magnesium oxide (Mg(OH)2/MgO) system is a promising chemical heatstorage system that utilizes unused heat at the temperature range of 200–500 ℃. The author havepreviously reported that the addition of lithium chloride (LiCl) and/or lithium hydroxide (LiOH)
promotes the dehydration of Mg(OH)2.

The results revealed that LiOH primarily catalyzed thedehydration of the surface of Mg(OH)2, while LiCl promoted the dehydration of bulk Mg(OH)2.
However, the roles of Li compounds in the hydration of MgO have not been discussed in detail. Xray diffraction (XRD) and Fourier-transform infrared (FT-IR) techniques were used to analyze theeffects of adding the Li compounds. The results revealed that the addition of LiOH promoted the
diffusion of water into the MgO bulk phase and the addition of LiCl promoted the hydration of the
MgO bulk phase. It was also observed that the concentration (number) of OH− affected hydration.
The mechanism of hydration of pure and LiCl- (or LiOH)-added MgO has also been discussed.

5.1. Introduction
The author has focused on chemical heat storage materials to utilize waste heat as shown inprevious chapters. The author has focused on Mg(OH)2 (CHS material) to effectively store thermalenergy in the temperature range of 200–300 ℃.
conditions and it exhibits high reversibility.

This material can be dehydrated under mild

The dehydration reaction of Mg(OH)2 and the hydration

reaction of MgO correspond to the heat storage and output operations, respectively.

The rate ofdehydration of Mg(OH)2 is significantly slow when the temperature is ˂300 ℃. A significantly low
hydration reactivity is observed for MgO in the temperature range of 150–200 ℃.[1–3] Therefore, thedifference in the dehydration and hydration temperatures is one of the problems for the widespread
application of CHS materials because the difference leads to the waste of sensible heat.[4]The author proposed that the addition of LiCl and/or LiOH to Mg(OH)2 samples can help improvethe dehydration reactivity (Chapter 3).[1,3,5–7]

The rate of dehydration of Mg(OH)2 in thetemperature range of 270–300 °C increased significantly when Li compounds were added.[ 1,3,5–7]
93

The hydration reactivity of the MgO system where both LiCl and LiOH were used as additives (at200 °C) was significantly higher than the hydration reactivity of the MgO system containing either
LiCl or LiOH.[3] The author has identified the roles played by LiCl and LiOH during the dehydrationof Mg(OH)2 (Chapter 4). LiOH primarily catalyzes the dehydration of the Mg(OH)2 surface, while
LiCl helps increase the rate of dehydration of the bulk Mg(OH)2 system.[8] The effect of the additionof these Li compounds on the hydration of MgO is yet to be fully understood.
Refason et al. suggested that the surface structures of the samples under study were similar to thesurface structures of Mg(OH)2 (0001) when MgO (111) was protonated during the hydration process.[9]This suggestion was in agreement with the suggestions presented in other studies where it was revealed
that the Mg(OH)2 (0001) plane was parallel to the MgO (111) plane during the process ofdehydration.[10,11] Refason et al. calculated the energy required for the chemisorption of water on
MgO. It was reported that the hydroxylation of the (001) surface of MgO was energetically unstable,
whereas the (111) surface of MgO was stabilized by hydroxylation based on density functional theory(DFT) calculations.[9] It can be interpreted that compared to the (100) and (001) planes, it was easier
to hydrate the (111) plane.[12] Some researchers have reported similar results.[13,14] The surface wasreconstructed to obtain the lowest energy structure.[14] Finocchi and Goniakowski[15] performed DFT
calculations to investigate the hydration properties of MgO. They reported that a strong bond wasformed between the O atom of H2O and the surface Mg atom present on the defective surface.[15] The
O atom filled the oxygen vacancy.[15] Hu et al.[16] reported that the water dimer was adsorbed on the
MgO (001) surface.

They also reported that proton transfer between H2O and the MgO surface

resulted in the dissociation of one H2O molecule.[16] However, the adsorption of a water monolayer
is more natural to consider.

Asaduzzaman[17] performed the first-principle DFT calculations toinvestigate the hydration of MgO. The interaction between H2O and the (001) and (111) planes of
MgO was studied. The results revealed that the interaction between the monolayer of water and the(111) surface of MgO resulted in the formation of Mg(OH)2.[17] The author demonstrated that surface
defects helped reduce the hydration energy barrier.[17] Kato et al. reported the results of a kinetic
study of the hydration of MgO. They assumed that the hydration reaction had four regimes: (1)containment of water as fixed structural water, (2) physical adsorption of water, (3) chemical reaction
with water producing Mg(OH)2, and (4) the participation of the inert portion of water.[18]Asaduzzaman simulated the first two mechanisms.[17] Pimminger et al.[11] and Iwasaki et al.[19] used
the thermogravimetric analysis (TGA) technique to conduct their studies. The results revealed thepresence of structural water.

Sharma et al. studied the dehydroxylation/rehydroxylation ofMg(OH)2/MgO using the transmission electron microscopy (TEM) technique.[12]During therehydroxylation of the dehydroxylated Mg(OH)2, the associated hydroxide species exhibited highmobility.

The number of dislocations and other defects increased during the initial stages ofcrystallization of the rehydroxylated material.[12] They reported the mechanism of hydroxylation of94

dehydroxylated Mg(OH)2 system. The first step involves the formation of Mg(OH)2 nanocrystals.
Following this, the rapid intergrowth of Mg(OH)2 nanocrystals acheved. The next step involves the
rapid formation and annealing out of the resulting crystal defects (e.g., edge dislocations) associatedwith nanocrystal intergrowth. High mobility of the hydroxides can also be achieved at this stage.
The high surface mobility of the participating Mg-containing species facilitates the growth of the
Mg(OH)2 crystals.[12]Chizallet et al.[20] identified the sites where the units were present andpresented molecular descriptions of the different types of OH groups.

They also identified thedifferent catalytically active sites participating in the 2-methylbut-3-yn-2-ol reaction.[20] Rimsza et
al. investigated the reaction between commercial MgO and deionized water. They used the proton
magic-angle spinning nuclear magnetic resonance (1H MAS NMR) technique to characterize thecompounds.[21] They observed several MgOH-type phases that were not strongly hydrogen-bonded
to nearby atoms. These phases were the dominant intermediate species produced during hydration.
The MgOH-type species were identified to be Mg5O4(OH)2, Mg4O3(OH)2, Mg3O2(OH)2, Mg2O(OH)2,
Mg3O(OH)4, and Mg5O(OH)8.[21] The studies focused only on the hydration of pure MgO. The
effects of adding LiCl and/or LiOH during the hydration process should be studied to further improve
the reactivity of the Mg(OH)2/MgO system that can be potentially used for the effective utilization ofwaste heat.
Fourier-transform infrared (FT-IR) spectroscopy is a qualitative analytical technique.
Researchers have studied the role of hydroxyl groups on MgO.

They proposed a stoichiometricequation for the adsorption and dissociation reactions involving molecular H2O. The equation canbe represented as follows (Eq. (1),[22–27]):

2+
𝑀𝑀𝑀𝑀𝐿𝐿𝐿𝐿− 𝑂𝑂𝐿𝐿2−′ 𝐶𝐶 + 𝐻𝐻2 𝑂𝑂 ⇄𝑂𝑂𝑂𝑂− 𝐻𝐻 +
|
|
2+
2−
𝑀𝑀𝑀𝑀𝐿𝐿𝐿𝐿−𝑂𝑂𝐿𝐿𝐿𝐿(1)

where LC represents the low coordinate.
Knözinger et al. classified OH groups into four types based on the coordination number and thepresence and absence of hydrogen bonds.[24] Based on the results obtained from the experiments
conducted using the FT-IR and DFT techniques, Chizallet et al.[25,27] proposed a new OH groupclassification.
The hydration properties of pure MgO and doped MgO have rarely been investigated using the
FT-IR technique. This study provides information on hydration based on previously reported results[22–27] obtained by studying the characteristics of OH groups on the MgO surface.

The author aims

to determine the effects of the addition of LiCl and/or LiOH on the hydration properties (heat output
operation) of MgO using two characterization techniques: X-ray diffraction (XRD) and FT-IR.
95

Insight into the role of Li compounds can help identify dopants that can effectively catalyze theMg(OH)2 dehydration and MgO hydration processes. It can also help in the development of other
CHS systems.

5.2. Experimental
5.2.1. Sample preparation
LiCl·H2O (99.9%), LiOH·H2O, and Mg(OH)2 (99.9%, 0.07 µm) were purchased from Wako Pure
Chemical Industries, Ltd. (Japan) and used as precursors to prepare the LiCl and LiOH co-added
Mg(OH)2 (referred to as LiCl/LiOH/Mg(OH)2) and LiCl- or LiOH-added Mg(OH)2 (referred to as
LiCl/Mg(OH)2 and LiOH/Mg(OH)2, respectively) samples.
LiCl/Mg(OH)2, LiOH/Mg(OH)2, and LiCl/LiOH/Mg(OH)2 were prepared following an
impregnation method as described in previous chapters.

LiCl-added Mg(OH)2 with an

Mg(OH)2:LiCl mole ratio of 100:10 (referred to as L10), LiOH-added Mg(OH)2 with an
Mg(OH)2:LiOH mole ratio of 100:20 (referred to as LO20), and LiCl and LiOH co-added Mg(OH)2
with an Mg(OH)2:LiCl:LiOH mole ratio of 100:10:10 (referred to as L10/LO10) were preparedfollowing this method.

Mg(OH)2 and MgO (99.9%, 0.05 µm, Wako Pure Chemical Industries, Ltd.),

devoid of Li compounds, referred to as Mg(OH)2-W and MgO-W, respectively, were preparedfollowing the same method. Details of all the prepared samples have been presented in Table 5-1.
Results from Chapters 2 and 3 revealed that the best mixing ratios could be expressed as L10, LO20,
and L10/LO10 (with respect to the activation energy for dehydration).
hydration reactivities could also be achieved.[1,3,6,7]

Increased dehydration and

These samples were used to compare thereactivities of the samples under study. The samples used in this study were the same as those usedin Chapter 4.[8]Table 5-1. Mixing ratio of the prepared samples.
SampleMixing ratio [mole ratio]

Mg(OH)2-W

Mg(OH)2 sample devoid of Li compounds

L10

Mg(OH)2:LiCl = 100:10

LO20

Mg(OH)2:LiOH = 100:20

L10/LO10

Mg(OH)2:LiCl:LiOH = 100:10:10

MgO-W

MgO samples devoid of the Li compounds

5.2.2. Hydration reaction test conducted using thermobalance
The reactivities of all the samples were determined using a thermobalance (TGD–9600 series,
96

ADVANCE RIKO, Inc.). The samples (20 mg) were charged into a Pt cell.
The conditions for the hydration reaction have been presented: the sample was heated from room
temperature (r. t.) to 120 ℃ at a heating rate of 20 °C min−1. The temperature was maintained at
120 ℃ for 30 min.

The flow of Ar gas was maintained at 100 mL min−1 to remove traces ofphysically adsorbed water. Following this, the sample was dehydrated at a temperature of 350 °C at
a rate of 20 °C min−1. The dehydration process was allowed to proceed for 30 min. Dehydration at
350 °C can achieve dehydration of 90% or more, which enables the sample to be hydrated from almost
complete MgO.

Subsequently, the sample was hydrated at different hydration temperatures (Th;

110 °C, 170 °C, or 200 °C) at a rate of −20 °C min−1. The hydration process was allowed to proceedfor every 10 min up to 80 min under a mixture of Ar gas and water vapor. The saturated vapor
pressures were (PH2O) 31.2 and 57.8 kPa, respectively when the saturated vapor temperatures were
70 °C and 85 °C, respectively. The pressure was controlled by controlling the water flow using a
micro feeder (NP-KX-101, NIHON SEIMITSU KAGAKU Co. Ltd. Japan). An Ar balance was also
used during the process.

Subsequently, the sample was dried at the hydration temperature (110 °C,

170 °C, or 200 °C) for 30 min under an atmosphere of Ar (Figure S3-1). Following the completionof the hydration reaction, XRD and FT-IR techniques were used to characterize the obtained samples.
The experiments were conducted in air at r. t. The samples dehydrated at a temperature of 350 ℃(before carrying out the hydration process) have been represented as de-samples (dehydrated
Mg(OH)2-W has been represented as de-Mg(OH)2-W) to distinguish them from the synthesized
samples.

5.2.3. Sample characterization using the XRD technique
XRD technique was used to investigate the crystal structures of the samples. The experiments
were conducted using an X-ray diffractometer (Ultima IV, Rigaku Corp., Japan) in the air at r. t. The2θ values were in the range of 10°–150° and the scan rate was 1.0° min−1 to confirm the successful
preparation of the desired samples. The scan width was set at 0.01°. The scan rate was set at 10°
min−1 (for reacted samples) to avoid further hydration or adsorption of water vapor on the reacted
samples. The Cu-Kα radiation was used for the studies. The generator voltage was 40 kV and thecurrent was 40 mA. All XRD patterns were recorded in the range of 10°–80°. Prior to conducting
the experiments, the prepared samples were dried overnight at 120 ℃ to remove all traces of physicallyadsorbed water.

5.2.4. Sample characterization using the FT-IR technique
FT-IR technique was used to obtain additional information on the hydration reaction of pure anddoped MgO. The spectra were recorded using an FT-IR spectrophotometer (FT/IR-4200, JASCO
Corporation, Japan) in the diffuse reflectance mode in air at r. t. The spectra were recorded before
97and after the hydration reaction was carried out in the thermobalance.

Calcium fluoride (CaF2)

powder was used to calibrate the baseline values. The scan range was 5000–350 cm−1, resolution
was 4.0 cm−1, and 64 scans were accumulated. The prepared samples were dried overnight at 120 °C
to remove traces of physically adsorbed water prior to recording the FT-IR spectra.

The author

reported the FT-IR spectra in the range of 5000–2750 cm−1 because the author focused on the peaks
derived from Mg(OH)2. The experimental procedure followed has been presented in Figure 5-1.
The influence of physically adsorbed water and CO2 on the results was investigated.

The

influence exerted by water vapor on the experimental resultscould be ignored because the intensitiesof the OH groups did not change significantly (Figure S3-3).[8] The author also observed that theadsorption of CO2 did not affect the reactivity of the Mg(OH)2/MgO system (Figure S3-4).[8,28,29]

Figure 5-1. Experimental procedure for this study.

5.3. Results
5.3.1. Sample (as-synthesized) characterization using the XRD and FT-IR
techniques
5.3.1.1. XRD patterns recorded for all the as-synthesized samples
Figure 5-2 shows the XRD patterns recorded for all the samples after they had been dried at 120 ℃.
The same figure has been presented in Chapter 4.[8] The figure presents the peaks corresponding to
the brucite structure of Mg(OH)2 (for all the samples). The peaks at 18.4°, 33.3°, 38.0°, 50.7°, 58.8°,
68.2°, and 72.0° were assigned to the Mg(OH)2 brucite structure. The author could not assign thepeak at 62° to either Mg(OH)2 or MgO. Hence, the author has not discussed its origin. The peak
patterns indicate the dominance of the brucite structure of Mg(OH)2. Although LiCl was added to
Mg(OH)2, peaks corresponding to LiCl could not be detected in the spectra of the LiCl-added samples(L10 and L10/LO10). The incorporation of the LiCl species into the Mg(OH)2 lattice can potentially98

explain this observation. It can also be explained if the particles are well dispersed throughout the
Mg(OH)2 particles. When an aqueous solution was prepared, it could not be detected during theexperiment. The absence of the peaks could also be attributed to the partial formation of the solid
solution between LiCl and Mg(OH)2.[1,3,8] It was also difficult to detect an LiOH phase. The author
has previously reported similar diffraction patterns for LiOH (and LiOH·H2O).[3,7,8] Therefore, small
amounts of LiOH or LiOH·H2O could be present on the sample surface. The partial conversion of
LiOH to Li2CO3 could be attributed to the adsorption of CO2 during the preparation process. Thisobservation reflects the fact that LiOH is generally contaminated with Li2CO3.[3,30] Thus, the peakcorresponding to Li2CO3 could be accounted for.

Figure 5-2. XRD patterns recorded for prepared samples.

This figure is the same as the figure

previously reported in Chapter 4.[8]5.3.1.2. Characterization of the as-synthesized samples using FT-IR techniqueFigure 5-3 shows the FT-IR spectral profiles recorded for the samples dried at a temperature of
120 ℃. The same figure has been presented in Chapter 4.[8] The peak at 3696 cm−1 was assigned
to the OH groups (that exhibited low coordination numbers) present on the surface edges and/or cornersites.[22–25,27] Knözinger et al.[24] assigned it to the isolated mono-coordinated OH groups. Chizallet
et al.[25,27] assigned it to the hydrogen-bonded mono-coordinated OH groups. Generally, the broad
absorption band present in the range of 3600–3200 cm−1 is assigned to the hydrogen-bonded watermolecules adsorbed on the sample surface.[24,25,27] The intensities of the broad bands corresponding
to the LiCl-added samples (L10 and L10/LO10) were significantly higher than those of the other two
samples.

This indicated that a significantly large volume of water was adsorbed on the sample
99surface.[31–33] This notable broad band can be attributed to the high hygroscopicity of LiCl. This
indicated that LiCl could readily adsorb water to form the hydrated species of LiCl. Therefore, LiCl
hydrate species could be formed on the surface of Mg(OH)2 (LiCl and LiCl hydrate species were notdetected by XRD).[8] However, it is difficult to observe the behavior of the broad band unless using
by in situ.

Hence, the author has not reported these properties herein. The author has previously

reported that the peaks at 4308, 4033, 3950, and 3794 cm−1 can be ascribed to the OH groups of thebulk Mg(OH)2 phase.[8]Figure 5-3. FT-IR spectral profiles recorded for prepared samples. This figure is the same as thefigure previously reported in Chapter 4.[8]Chapter 4 has characterized the samples to investigate the states of LiCl and LiOH in the samples
and reported the results.[8]The chapter showed that the FT-IR spectra of LO20 prepared byphysically mixing presented the peak by LiOH∙H2O, while LO20 prepared by the impregnation
method has no peaks due to LiOH∙H2O.[8] Therefore, a solid solution consisting of Mg(OH)2 andLiOH∙H2O (or LiOH) could be produced by the impregnation method for LO20. A solid solution ofMg(OH)2 and LiCl could also be produced by the impregnation method because a peak shift in the
XRD pattern was detected when the spectra were recorded for the LiCl-added samples (L10 and
L10/LO10).[8] The absence of the diffraction patterns due to LiCl can potentially account for theformation of the solid solution in LiCl-added samples (L10 and L10/LO10).

It has been reported[28,34,35] that a metal ion can potentially occupy the site of a magnesium ion or formation of defect sites

by the metal compound doping. A literature by Housecroft et al.[36] reported that the properties ofmagnesium and lithium are typically similar (diagonal relationships). Therefore, the formation of100

the solid solution by a Li+ ion substitution into a Mg2+ ion site can be reasonable.

5.3.2. Characterization of hydrated samples using the XRD and FT-IR techniques
5.3.2.1. XRD patterns recorded for the hydrated samples
Figure 5-4 shows the XRD patterns of the de-Mg(OH)2-W, de-L10, de-LO20, and de-L10/LO10
samples post hydration at 110 ℃.

The black line represents the patterns immediately before

hydration at 110 ℃. The hydration behavior at 110 ℃ is shown in Figure S3-5. Analysis of Figure
5-4 revealed that the intensity of the peak corresponding to Mg(OH)2 increased, and the intensity ofthe peak corresponding to MgO decreased as the hydration time increased. The peaks corresponding
to MgO (in the de-Mg(OH)2-W, de-L10, and de-LO20 samples) were retained post hydration at 110 ℃.
The peaks could be observed even when the experiment was conducted for 80 min.

The results

agreed well with the results obtained from the TG experiments. Thus, the unreacted core and/or inert
portion of the water was retained by the sample.[18] The hydration conversion (∆x1) values wererecorded for the de-Mg(OH)2-W samples (hydration temperature: 110 ℃; time: 60–80 min). The
values leveled off at approximately 70% (Figure S3-8 (a)).
previously reported result.

[18,37]This result agrees well with the

The observations can be attributed to the transition process (frominterface reaction to diffusion control).[37] This could be ascribed to the successive clogging of theporous and shielding of the MgO surface by the newly formed Mg(OH)2.[37,38] Figure 5-4 (b) showsthat the intensity of the peak corresponding to Mg(OH)2 present in de-L10 (recorded after the first 10–
30 min of hydration) was significantly weaker than the intensity of the peak corresponding to Mg(OH)2in de-Mg(OH)2-W. Figure S3-8 (a) shows that the values of ∆x1 recorded for de-L10 during the first
20 min were lower than those recorded for de-Mg(OH)2-W. The ∆x2 values recorded for de-L10
were significantly higher than those recorded for de-Mg(OH)2-W. The results suggested that deMg(OH)2-W could be more easily hydrated than de-L10 when the hydration time was in the range of
10–30 min.

Results obtained from the experiments conducted using XRD and TG techniques(Figure S3-5 (b)) revealed that the hydration of LiCl was favored at the early stages of hydration.
This indicated that the reaction pathways for the hydration of de-Mg(OH)2-W and de-L10 weredifferent from each other. Hydrated LiCl species could be formed and the H2O molecules present in
the hydrated LiCl species (Eq. (1)) could potentially interact with the MgO surface, resulting in the
promotion of MgO hydration (Figure S3-8 (a)).[1,3] The hydration process can be represented asfollows:(1)

𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 + 𝑛𝑛𝑛𝑛2 𝑂𝑂 ↔ 𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 ∙ 𝑛𝑛𝑛𝑛2 𝑂𝑂

Figure 5-4 (c) shows that the peak intensity at 38°, corresponding to the Mg(OH)2 (002) plane inde-LO20, was stronger (post hydration; 10 min) than the peak intensity at 42° (corresponding to the101

MgO (200) plane). This result indicated that the addition of LiOH promoted the hydration of MgO
at an early stage. Interestingly, a weak peak at 36°, corresponding to the MgO (111) plane, wasdetected when the de-LO20 sample was hydrated and the spectra were recorded (time: 80 min).
Figure 5-4 (d) revealed that the peak corresponding to MgO could not be detected post hydration(time range: 50–80 min; Table 5-5). This indicated that the unreacted core and the inert portion ofwater were completely hydrated. Therefore, the addition of LiCl and LiOH resulted in the significant
enhancement of the hydration reactivity of MgO.[3]The patterns recorded immediately before de-L10 samples were hydrated (black line in Figure 54 (b)) present the full width at half maximum (FWHM) of the peak at 42°. This peak was sharper

Figure 5-4. XRD patterns recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and (d)de-L10/LO10 (samples were hydrated at a temperature of 110 ℃). Mg(OH)2-W, L10, LO20, andL10/LO10 were dehydrated at 350 ℃, following which the samples were hydrated at 110 °C (PH2O:
57.8 kPa; samples hydrated using a thermobalance). Following the completion of the reaction, thediffraction patterns were recorded in air at r. t.

Red dotted lines: positions of the peaks corresponding

to MgO before hydration. Black dotted lines: positions of the peaks corresponding to Mg(OH)2 post
hydration.

102

than the peak representing de-Mg(OH)2-W (black line in Figure 5-4 (a)). The FWHM recorded forde-LO20 (black line in Figure 5-4 (c)) was sharper than that recorded for de-Mg(OH)2-W. These
results indicated that the surface states of de-L10 and de-LO20 (before hydration) were different fromthe surface states of de-Mg(OH)2-W.

The crystallinities exhibited by de-L10 and de-LO20 weresignificantly higher than the crystallinity exhibited by de-Mg (OH)2-W. The order of crystallinity ofthe samples (after dehydration) was determined: de-Mg(OH)2-W < de-LO20 < de-L10.
Figure 5-5 shows the XRD patterns of the de-Mg(OH)2-W, de-L10, de-LO20, and de-L10/LO10
samples after hydration at 170 ℃.

The black line represents the patterns recorded immediately

before the samples were hydrated at 170 ℃.

The hydration behavior is shown in Figure S3-6.

Figure 5-5 shows that the intensity of the peak corresponding to Mg(OH)2 increased with increasing
hydration reaction time. Figure 5-5 (a) shows that the peak at 38° (in de-Mg(OH)2-W) appeared
when the sample was hydrated for 30 min and the peak was assigned to Mg(OH)2 when the hydration
time was 40 min (Table 5-2). The results revealed that the surface of MgO could be potentially
hydrated during the first 10–30 min of hydration. A peak at 38° was observed in the spectra recordedwith de-L10 (Figure 5-5 (b)). The peak was assigned to Mg(OH)2 (post hydration; 10 min; 170 °C)
for the de-L10. The time taken for the peak to assign to Mg(OH)2 was shorter when de-L10 was
studied than the time taken for the peak to assign to Mg(OH)2 when de-Mg(OH)2-W was studied.
Figure 5-5 (c) reveals that for de-LO20, the peak corresponding to Mg(OH)2 appeared when thehydration time was 10 min and the temperature was 170 ℃. The results obtained by analyzing the
XRD data revealed that the addition of LiCl (or LiOH) promoted the hydration of MgO. The results
agreed well with the results obtained by analyzing the data obtained using the TG technique (Figure
S3-8 (b)). Figure 5-5 (d) shows that the intensity of the peak corresponding to MgO (de-L10/LO10)
was significantly weaker than the intensity of the peak corresponding to MgO, recorded for othersamples. This indicated that the hydration process proceeded well at 170 ℃.

This result is inagreement with the results obtained using the TG technique (Figure S3-8 (b)).
Figure 5-6 shows the XRD patterns of de-Mg(OH)2-W, de-L10, de-LO20, and de-L10/LO10 post
hydration at 200 ℃. The black line represents the patterns recorded with the samples immediately
before hydration at 200 ℃.

The hydration behavior is shown in Figure S3-7.

Figure 5-6 (a)reveals that the peaks corresponding to Mg(OH)2 (in de-Mg(OH)2-W) could not be detected post
hydration. This result is in agreement with the results obtained using the TG technique (Figure S37).

Most of the samples remained as MgO.

hydration temperature was increased.

Lower hydration reactivity was achieved as theThis agrees well with the previously reported results.[1,2]These results can be attributed to the exothermic hydration reaction and/or the small driving force ofthe hydration reaction recorded at a high temperature (the temperature can approach theequilibrium).[2,38]

Figure 5-6 (b) shows the peaks corresponding to Mg(OH)2 in de-L10 wereobserved after hydration at 200 °C only for 80 min. Figure 5-6 (c) shows that a 10-min hydration103

process at a temperature of 200 ℃ can result in the formation of the Mg(OH)2 phase in de-LO20.
This result indicated that the addition of LiOH resulted in an increase in the hydration reactivity of
MgO. The extent of the increase in reactivity was larger when LiOH was added than when LiCl wasadded.

The results agreed well with the results obtained using the TG technique and the results

reported previously [3] (Figure S3-7). Figure 5-6 (d) shows that the Mg(OH)2 peak at 38° appears
post hydration (time: 10 min). The intensity of the Mg(OH)2 peak was higher than the intensity ofthe peaks of the other samples. The author used the XRD technique to confirm that de-L10/LO10
exhibited high hydration reactivity. These results agreed well with the results obtained using the TG
technique and results reported previously by us (Figures S3-5–S7).[3]Figure 5-5. XRD patters recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and(d) de-L10/LO10 (samples were hydrated at a temperature of 170 ℃). Mg(OH)2-W, L10, LO20,
and L10/LO10 were dehydrated at 350 ℃ to prepare de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10, respectively (Red dotted lines: MgO peak position before hydration, black dotted lines:
Mg(OH)2 peak position post hydration).

104

Figure 5-6. XRD patterns recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and (d)de-L10/LO10 (samples were hydrated at a temperature of 200 ℃). Mg(OH)2-W, L10, LO20, andL10/LO10 were dehydrated at 350 ℃ to prepare de-Mg(OH)2-W, de-L10, de-LO20, and de-L10/LO10,
respectively (Red dotted lines: MgO peak position before hydration, black dotted lines: Mg(OH)2 peak
position post hydration).

5.3.2.2. Analysis of the FT-IR spectral profiles of the hydrated samples
Figure 5-7 shows the FT-IR spectra recorded for the de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10 samples post hydration at 110 ℃.

The black lines represent the spectra recorded

immediately before the samples were hydrated at 110 ℃. The hydration behavior is shown in FigureS3-5. The absorption bands at 4300, 4030, 3950, 3790, and 3700 cm−1 were observed post hydration
at a temperature of 110 ℃ (time: 10 min; Figure 5-7).

Analysis of the figures revealed that theabsorption band at 3700 cm−1 (black lines in Figure 5-7; recorded before hydration) shifted to a lower
wavenumber region post hydration (Tables 5-2–5). This indicated the formation of the hydrogenbonded hydroxyl groups and an increase in the coordination number of surface hydroxyl groups.[24,27]
The black lne in Figure 5-7 (c) shows a small shoulder band at 3761 cm−1 and split peaks at
approximately 3700 cm−1 (3688 and 3677 cm−1). The shoulder and split peaks were assigned to the
MgO species based on the results of Chapter 4.[8] Therefore, the surface state of the sample was105similar to that of MgO (before hydration; de-LO20).

The FT-IR spectral profiles of de-MgO-W

before hydration are shown in Figure S3-9. The two split peaks coalesced to form one sharp peak,which shifted to the lower wavenumber region (Figure 5-7 (c)).
Figure 5-8 shows the FT-IR spectral profiles of de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10 post hydration at 170 ℃.

The black lines represent the spectral profiles recorded

immediately before hydration at 170 °C. The hydration behavior is shown in Figure S3-6. Figure
5-8 (a) revealed that the peak at 4030 cm−1 was detected post hydration (time: 30 min) in de-Mg(OH)2W. A diffraction peak corresponding to Mg(OH)2 was observed at 38° (Figure 5-4 (a)) at this time.
Surface hydration was favored when the hydration time was in the range of 10–20 min.
hydration proceeded at 170 °C when the hydration time was in the range of 30–40 min.

Bulk

Bands atFigure 5-7. FT-IR spectral profiles recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20,and (d) de-L10/LO10 (samples were hydrated at a temperature of 110 ℃). Mg(OH)2-W, L10, LO20,
and L10/LO10 were dehydrated at 350 ℃.

Following this, they were hydrated at a temperature of

110 ℃ (PH2O: 57.8 kPa; thermobalance was used).
IR spectra were recorded in air at r. t.

Following the completion of the reaction, the FT-

Red dotted lines: peak position before hydration. Blackdotted lines: peak position post hydration.

106

4298 and 3955 cm−1 were observed, while the peak at 3790 cm−1 could not be detected (Table 5-2)
when the samples were allowed to hydrate for 70 min. The peak at 3700 cm−1 shifted to a lower
wavenumber region as the hydration time increased.

The band at 3700 cm−1 (spectra recorded

immediately before hydration at 170 ℃; de-L10; black line in Figure 5-8 (b)) was absent.

Theabsence of the peak could be attributed to the low intensity of the band and the presence of a prominent
broad absorption band in the region of 3600–3200 cm−1. Peaks at 3699, 4028, 3953, 4305, and 3796
cm−1 were observed post hydration at a temperature of 170 ℃ when the hydration times were 10, 10,
10, 20, and 30 min, respectively (Table 5-3). The black line in Figure 5-8 (c) represents the shoulder(3761 cm−1) and split peaks (3687 and 3676 cm−1). These were detected in the spectra recorded with
de-LO20. The spectra were recorded immediately before hydration at 170 ℃. The two split peaks
coalesced to form one sharp peak, which shifted to a lower wavenumber region. The peaks at 4305,

Figure 5-8. FT-IR spectral profiles recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20,and (d) de-L10/LO10 (samples were hydrated at a temperature of 170 ℃). Mg(OH)2-W, L10, LO20,
and L10/LO10 were dehydrated at 350 ℃ to prepare de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10, respectively. (Red dotted lines: peak position before hydration, black dotted lines: peak
position post hydration).

107

4034, 3955, and 3797 cm−1 were observed post hydration when the hydration times were 10, 10, 10,
and 20 min, respectively (Table 5-4). The peak at 3761 cm−1 could be detected even after hydration(time: 50 min, Table 5-4). An analysis of the black line in Figure 5-8 (d) revealed that the intensityof the peak at 3700 cm−1 was significantly weak. The intensity was similar to the intensity of thepeak recorded for L10 (black line in Figure 5-8 (b)). The peaks at 4309, 4034, 3950, and 3796 cm−1
were observed post hydration at 170 °C (time:10 min; de-L10/LO10; Table 5-5). The time taken forthese peaks to appear in the spectral profile was shorter in de-L10/LO10 than the time taken for thesepeaks to appear in the profiles recorded for other samples.
Figure 5-9 shows the FT-IR spectral profiles of de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10. The spectra were recorded post hydration at a temperature of 200 ℃. The black lines
represent the spectra recorded immediately before the samples were hydrated. The hydration behavior

Figure 5-9. FT-IR spectral profiles recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20,and (d) de-L10/LO10 (samples were hydrated at a temperature of 200 ℃). Mg(OH)2-W, L10, LO20,
and L10/LO10 were dehydrated at a temperature of 350 ℃ to prepare de-Mg(OH)2-W, de-L10, deLO20, and de-L10/LO10, respectively.

(Red dotted lines: peak position before hydration, black

dotted lines: peak position post hydration (secondary differential method was used to detect the peakat 3761 cm−1 (de-Mg(OH)2-W)).
108

is shown in Figure S3-7. Absorption bands were observed at 3761 and 3715 cm−1 for de-Mg(OH)2W (Figure 5-9 (a)). The intensity and the position of the band originally at 3715 cm−1 (in the spectral
profile of de-Mg(OH)2-W) hardly changed when the temperature was 200 ℃. The shoulder band at
3761 cm−1 (assigned to the hydroxyl groups present on the MgO surface) appeared post hydration.
This indicated that the decomposition of Mg(OH)2 was favored over the hydration of MgO at 200 ℃.
Figure 5-9 (b) shows that the peak at 3700 cm−1 (in the spectral profile of de-L10) did not appear
during the first 20 min of hydration when the temperature was 200 ℃.
attributed to the weak peak intensity.

The absence could be

The intensity of the peak increased when the sample washydrated over a period of 30 min. This indicated that the sample surface was continuously getting
hydrated. However, bands in the range of 4300–3790 cm−1, attributable to the bulk OH groups, were
not observed. The black line in Figure 5-9 (c) shows a shoulder peak at 3761 cm−1 and split peaks
at approximately 3700 cm−1 (3686 and 3676 cm−1). These peaks appeared in the spectral profiles
when the spectra were recorded with de-LO20 immediately before hydration at 200 ℃. The two
split peaks coalesced to form one sharp peak post hydration (time: 10 min). The single peak shiftedto a lower wavenumber region. The peaks at 4032, 3956, 4302, and 3794 cm−1 were observed post
hydration when the hydration times were 10, 10, 20, and 60 min, respectively (Table 5-4).

Theshoulder peak at 3761 cm−1 was observed in the spectral profile (Figure 5-9 (c)). This indicated thatthe hydration reaction path was influenced by the addition of LiOH and/or the hydration temperature.
The author used the FT-IR technique to confirm that the addition of LiCl (or LiOH) promoted thehydration of MgO. Results from the experiments conducted using the FT-IR technique also revealed
that the extent of the positive effect exerted by LiOH on the hydration process was larger than theextent of the positive effect exerted by LiCl. These results agreed well with the results obtained using
the TG and XRD techniques.

Figure 5-9 (d) shows that the peaks at 4033, 3953, 4311, and 3787

−1 appeared post hydration when the hydration times were 10, 10, 20, and 20 min, respectivelycm(Table 5-5). Figures 5-7–9 (d) reveal that the addition of LiCl and LiOH result in a significant
increase in the hydration reactivity of MgO. This result agreed well with the results obtained using
the TG and XRD techniques and the results in Chapter 3.[3]The data presented in Tables 5-2–5 reveal the relationship between the peak positions (in the FTIR spectral profiles) and the hydration reaction time (at three different hydration reaction temperatures).
The circles in the tables indicate the presence of the Mg(OH)2 (or MgO) phase.

Peaks corresponding

to these phases were observed in the recorded XRD patterns.
Analysis of the Figures S3-8 (b) and (c) revealed that the measured ∆x1 values (expressed aspercentages) for de-Mg(OH)2-W hydrated at 170 ℃ (time: 30 min) and de-L10 hydrated at 200 °C for
70 min were 11% and 5%, respectively.

Peaks indicating the presence of the Mg(OH)2 phases were

not observed in the diffraction patterns recorded for de-Mg(OH)2-W and de-L10 (Figures 5-5 (a) and5-6 (b), respectively). The FT-IR spectra of the de-Mg(OH)2-W and de-L10 samples post hydration
109

Table 2. Relationship between the obtained peaks (de-Mg(OH)2-W, after hydration) and the hydration
time (circle in the table indicates the detection of a phase in the XRD pattern).
Hydration

Th=110 ℃ Peak position [cm−1]

time [min]
80

4298

4036

3953

3798

3698

70

4305

4033

3951

3785

3698

60

4312

4036

3957

3796

3698

50

4302

4034

3959

3799

3698

40

4309

4032

3955

3797

3698

30

4311

4032

3959

3796

3698

20

4308

4030

3961

3794

3698

10

4305

4034

3955

3795

3698

0

3712

Hydration

Th=170 ℃ Peak position [cm−1]

time [min]
80

4298

4031

3955

3699

70

4314

4034

3960

3700

60

4026

3700

50

4029

3701

40

4026

3701

30

4024

3701

20

3703

10

3707

0

3712

Hydration
time [min]Th=200 ℃ Peak position [cm−1]

80

3715

70

3761

3714

60

3711

50

3715

40

3715

30

3714

20

3716

10

3713

0

3712
110

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯

◯

◯
◯
◯

◯

◯

◯

XRD:

XRD:

Mg(OH)2

◯

MgO

◯

◯

◯

◯

◯

◯

◯
◯

◯

◯

◯
◯

◯

◯
◯

◯

XRD:

XRD:

Mg(OH)2

MgO
◯
◯

◯

◯
◯
◯
◯
◯

◯

Table 3. Relationship between the obtained peaks (de-L10, after hydration) and the hydration time(circle in the table indicates the detection of a phase in the XRD pattern).
Hydration

Th=110 ℃ peak position [cm−1]

time [min]
80

4305

4035

3950

3788

3699

70

4310

4034

3953

3788

3699

60

4306

4031

3954

3793

3699

50

4310

4035

3950

3790

3699

40

4308

4034

3954

3787

3699

30

4310

4034

3956

3792

3699

20

4297

4035

3950

3796

3700

10

4302

4032

3959

3792

3700

0

3700

Hydration

Th=170 ℃ peak position [cm−1]

time [min]
80

4305

4028

3953

3796

3699

70

4309

4036

3950

3789

3700

60

4309

4032

3956

3796

3700

50

4316

4035

3954

3792

3699

40

4306

4032

3956

3788

3700

30

4293

4030

3953

3789

3700

20

4291

4028

3953

3700

4031

3958

3701

10
0
Hydration
time [min]Th=200 ℃ peak position [cm−1]

80

3699

70

3699

60

3699

50

3699

40

3699

30

3699

20

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯
◯
◯

◯
◯
◯
◯
◯
◯

◯

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯

◯
◯
◯
◯

◯

◯

XRD:

XRD:

Mg(OH)2

◯

MgO

◯

◯

◯

◯

◯
◯
◯
◯
◯
◯

10

◯

0
111

◯

Table 4. Relationship between the obtained peaks (de-LO20; after hydration) and the hydration time(circle in the table indicates the detection of a phase in the XRD pattern).
Hydration

Th=110 ℃ peak position [cm−1]

time [min]
80

4315

4035

3951

3800

3697

70

4305

4034

3955

3792

3696

60

4308

4035

3953

3796

3697

50

4305

4036

3955

3791

3697

40

4307

4033

3956

3793

3696

30

4305

4035

3953

3786

3697

20

4300

4035

3955

3787

3697

10

4299

4034

3954

3798

3698

0

3761

Hydration

3688

80

4305

4034

3955

3797

3697

70

4312

4032

3953

3792

3697

60

4310

4033

3951

3794

3697

50

4314

4035

3953

3792

3757

3697

40

4304

4032

3955

3787

3757

3698

30

4298

4031

3951

3794

3756

3697

20

4301

4035

3958

3794

3757

3697

10

4302

4035

3953

3761

3697

3762

3687

0
Hydration
80

4302

4032

3956

3794

3758

3697

70

4304

4034

3951

3788

3761

3697

60

4300

4032

3951

3788

3763

3696

50

4301

4033

3957

3762

3696

40

4305

4033

3953

3761

3696

30

4297

4034

3952

3761

3696

20

4296

4034

3952

3761

3696

4031

3950

3761

3696

3761

3686

10
0

112

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯
3677

◯

◯
◯
◯
◯

◯

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯

◯
◯
◯
◯
◯
◯

◯

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯
3676

◯

XRD:

◯

3676

◯

XRD:

◯

Th=200 ℃ peak position [cm−1]

time [min]

XRD:

◯

Th=170 ℃ peak position [cm−1]

time [min]

XRD:

◯
◯

◯
◯
◯
◯
◯
◯

◯

Table 5. Relationship between the obtained peaks (de-L10/LO10; after hydration) and the hydration
time (circle in the table indicates the detection of a phase in the XRD pattern).
Hydration

Th=110 ℃ peak position [cm−1]

time [min]
80

4312

4033

3952

3788

3697

70

4308

4034

3954

3793

3697

60

4309

4033

3952

3793

3697

50

4307

4032

3951

3787

3698

40

4312

4031

3955

3788

3697

30

4300

4032

3950

3788

3698

20

4302

4034

3951

3799

3698

10

4316

4034

3956

3794

3698

0

3700

Hydration

Th=170 ℃ peak position [cm−1]

time [min]
80

4309

4034

3950

3796

3698

70

4306

4035

3953

3788

3698

60

4303

4034

3953

3793

3698

50

4307

4032

3955

3793

3698

40

4307

4033

3953

3787

3698

30

4312

4032

3957

3790

3698

20

4310

4036

3956

3793

3698

10

4308

4031

3954

3787

3699

0

3699

Hydration

Th=200 ℃ peak position [cm−1]

time [min]
80

4311

4033

3953

3787

3698

70

4303

4033

3952

3788

3698

60

4310

4033

3951

3794

3698

50

4308

4032

3951

3788

3698

40

4313

4033

3952

3788

3698

30

4303

4035

3954

3800

3698

20

4306

4031

3951

3788

3698

4031

3948

10

3699

0

3699
113

XRD:

XRD:

Mg(OH)2

MgO

◯
◯
◯
◯
◯

◯

◯

◯

◯
◯

◯
◯

◯

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯
◯
◯
◯
◯
◯

◯
◯
◯
◯
◯
◯

◯

XRD:

XRD:

Mg(OH)2

MgO

◯

◯

◯

◯

◯

◯

◯
◯
◯
◯

◯

◯
◯
◯
◯

◯

◯

at 170 ℃ (time: 30 min) and 200 ℃ (time: 70 min), respectively, revealed the presence of theabsorption bands at 3700 cm−1. These bands were assigned to the surface OH groups (Figures 58(a) and 5-9 (b)). [58–61,63] The results that bulk information could be obtained by analyzing the XRD
patterns.

Peaks corresponding to structural water (described as ∆xs; experimental section in the

Supporting Information for chapter 5) were not observed in the XRD patterns recorded with thesamples immediately before they were hydrated (black lines in Figures 5-4–6). A peak at 3700 cm−1
was observed in the FT-IR spectral profiles recorded immediately before the samples were hydrated(black lines in Figures 5-7–9). The absorption band corresponding to the surface OH groups present
on Mg(OH)2 or MgO did not completely disappear post dehydration at a temperature of 350 ℃.
Therefore, the mole fraction of Mg(OH)2 remaining as structural water post dehydration at 350 ℃ can
be attributed to such surface OH groups, indicating that magnesium oxide could potentially form thebulk state of the samples.
Mutch et al. investigated the ability of MgO to capture carbon. They reported that MgO calcinedat 800 ℃ exhibited a higher carbon capacity (for CO2) compared to that calcined at 400 ℃.[39] Intheir study, the peak at 3760 cm−1, in the spectral profile recorded for MgO calcined at 800 °C, wasassigned to a 100% H-covered O-terminated (111) plane. They used DFT calculations to arrive at
the conclusions.[39]The position of this peak and the position of the shoulder peak (3761 cm−1)

observed by us agreed well with each other (Figures 5-7–9 (c)). Figure S3-10 (reported in Chapter4[8]) was revised as shown in Figure 5-10.

Figure 5-10. FT-IR spectral profile recorded for Mg(OH)2-W.
peak assignment.

[8,22–25,27,39]

114

Literature reports were referred to for5.4. Discussion
5.4.1. Effect of the addition of LiCl and/or LiOH on the hydration of MgO
The XRD data presented in Section 5.3.2.1. suggested the presence of an unreacted core and/or
inert portion of water in de-Mg(OH)2-W, de-L10, and de-LO20. The ∆x1 values recorded for de-L10,
de-LO20, and de-L10/LO10 were always higher (at all hydration temperatures) than the ∆x1 values
recorded for the de-Mg(OH)2-W samples (Figure S3-8). These results indicated that the addition of
LiCl and/or LiOH promoted the hydration of MgO in the unreacted core and/or inert portion of water.
The addition can also help lower the activation energy of the process.
For the hydration of the MgO with added-LiCl (de-L10, and de-L10/LO10) at 110 °C, thehydration of LiCl was favored at this temperature. The water molecules present in hydrated LiCl
species interacted with MgO.[1,3]This can be one of the effects of the addition of LiCl on thehydration. The hydration of LiCl, however, unnecessarily positively affected the hydration of MgO,
because the value of ∆x1 for de-L10 at 110 °C and 31.2 kPa for 80 min was lower than that for deMg(OH)2-W (Figure S3-11).

At higher temperatures of 170 ℃ and 200 ℃, the hydration of LiCl

was not favored (the values of ∆x2 recorded for de-L10 and de-L10/LO10 (Figures S3-8 (b) and (c),respectively) were low).

Hence, the hydration of LiCl did not exert a significant effect on thehydration of MgO. However, the fact that the addition of LiCl enhanced the hydration reactivity of
MgO suggested that other factors affected the hydration of MgO. Finocchi et al.[15] performed DFT
calculations to investigate the process of water adsorption on a MgO (100) surface containing an
oxygen vacancy.

They reported that the lattice defect was occupied by the O atom of H2O.[15]Asaduzzaman performed DFT-based calculations. The results revealed that the adsorption energy
for water molecules adsorbed on the MgO (111) surface containing the lattice defects was significantly
low.[17] The author has previously reported that lattice defects (such as OH−) could be produced by
Li+ ion substitution into Mg2+ sites.[8] This indicated that dehydration could result in the creation ofoxygen vacancies. These results indicated that the substitution with Li+ could promote the processof hydration (in defect sites).

Chapter 4 reported[8] that OH− defects in bulk Mg(OH)2 can be createdin the presence of LiCl. The addition of LiCl probably results in the formation of O2− vacancies in
MgO bulk post dehydration. This process promotes the hydration process in the bulk phase. On
the other hand, effective hydration could not be achieved for de-L10 at a temperature of 200 ℃.
Rapid hydration can be observed in the bulk phase if the surface hydration and/or water adsorption
process proceeds smoothly.
The ∆x1 values recorded at 200 ℃ for de-LO20 and de-L10/LO10 were significantly higher thanthose recorded for de-Mg(OH)2-W and de-L10 (Figure S3-8 (c)).

It is worth noting that the ∆x1

value recorded for de-L10/LO10 reached 90% when the sample was hydrated at a temperature of
200 ℃ for 80 min. The addition of LiOH exerted a positive effect on the MgO hydration process.
115

A cumulative effect was observed when LiCl was added.

The author has discussed the reasons

behind the significant increase in the hydration reactivity of the samples (de-LO20 and de-L10/LO10)
containing LiOH as the additive.

Cahpter 4 reported[8] that the addition of LiOH results in theformation of surface defects on the surface of Mg(OH)2. The number of surface defects produced
was larger when LiOH was added than the number of surface defects produced when LiCl was added.
Thus, the surface defects possibly promoted the process of surface hydration (for MgO) in the presenceof LiOH. The shoulder peak at 3760 cm−1 (assigned to the isolated hydroxyl groups on the MgO(111) surface[39]; de-LO20) should have disappeared in the spectral profiles recorded during the earlystages of hydration if the surface defects produced following the process of Li+ ion substitution
promotes surface hydration (Figures 5-8 (c) and 5-9 (c)). Therefore, it is believed that the additionof LiOH promotes the diffusion of water into the bulk phase through the surface defects. The (111)plane can be observed in the XRD patterns recorded for de-LO20 (Figures 5-3–5 (c)) post hydration(time: 80 min).

Some researchers have reported that compared to the MgO (100) plane, it is easier

to hydrate the MgO (111) plane.[9,13,14,17]

These results indicate that the sites that can be easily

hydrated can be potentially generated by adding LiOH.
Amaral et al. reported a probable mechanism for the hydration of MgO. The hydration processproceeded in the presence of liquid water and the mechanism can be expressed as follows[40]:1. MgO-alkaline oxide plays an electron donator role in water:+−+ 𝑂𝑂𝑂𝑂(𝑎𝑎𝑎𝑎)
𝑀𝑀𝑀𝑀𝑀𝑀(𝑠𝑠) + 𝐻𝐻2 𝑂𝑂(𝑙𝑙) → 𝑀𝑀𝑀𝑀𝑀𝑀𝑀𝑀(𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠)(2)

2. OH– anions are adsorbed in the positively charged surface:+−−
𝑀𝑀𝑀𝑀𝑀𝑀𝑀𝑀(𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠)+ 𝑂𝑂𝑂𝑂(𝑎𝑎𝑎𝑎)→ 𝑀𝑀𝑀𝑀𝑀𝑀𝑀𝑀+ ∙ 𝑂𝑂𝑂𝑂(𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠)(3)

3. OH– anions are desorbed from the surface, releasing magnesium ions into the solution:

2+−−→ 𝑀𝑀𝑀𝑀(𝑎𝑎𝑎𝑎)+ 2𝑂𝑂𝑂𝑂(𝑎𝑎𝑎𝑎)
𝑀𝑀𝑀𝑀𝑀𝑀𝑀𝑀+ ∙ 𝑂𝑂𝑂𝑂(𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠𝑠)

(4)

2+
−
+ 2𝑂𝑂𝑂𝑂(𝑎𝑎𝑎𝑎)→ 𝑀𝑀𝑀𝑀(𝑂𝑂𝑂𝑂)2(𝑠𝑠)𝑀𝑀𝑀𝑀(𝑎𝑎𝑎𝑎)(5)

4. Ions concentration reaches the solution super-saturation, at which point the hydroxide starts toprecipitate on the oxide surface:

Amaral et al.[40] reported that the extent of hydration of MgO in the presence of a solution of KOH

increased with an increase in the concentration of the solution.

A high concentration of OH− (asolution of KOH is being used) can potentially cause a shift in the equilibrium, resulting in the
promotion of hydration (Eq. 5).[40] The data presented in Figure S3-12 and Table S3-1 show thatthe hydration reactivity of LiOH/MgO increase with an increase in the mole ratio of LiOH (used as anadditive).

Figures S3-13 (a) and (b) show the FT-IR spectral profiles of LO20, LO25, LO30, andLO50 prepared following the impregnation and physical mixing methods, respectively. The absenceof the peaks corresponding to LiOH∙H2O revealed that a solid solution of LiOH and Mg(OH)2 (for116

LO20) could be produced following the impregnation method. When the spectra were recorded forLO20 (prepared following the physical mixing method), peaks corresponding to LiOH∙H2O wereobserved.[8] A peak at 3654 cm−1 was observed in the spectra recorded for impregnated LO25, LO30,
and LO50 (Figure S3-13 (a)).

This indicated that limited numbers of Mg2+ ion sites could besubstituted by Li+ ions. It was observed that the number (concentration) of OH− in LiOH potentiallyaffected the hydration of MgO (Figure S3-12 and Table S3-1).

However, the effects of otherhydroxides (e.g., NaOH and KOH) on the extent of hydration should be studied to determine if theconcentration of OH− affects the hydration reactivity.
It was observed that the addition of LiOH could potentially promote the diffusion of H2O into the
MgO bulk phase (particularly at higher temperatures) because the number of surface defects produced
was large under the reported conditions and a shift in the equilibrium toward the direction of theformation of Mg(OH)2 could be achieved.
hydration of the MgO bulk phase.
hydration reaction.

The addition of LiCl could potentially promote the

The extent of hydration depended on the rate of the surface

For the hydration of LiCl and LiOH co-added MgO, these two effects

simultaneously contributed to the promotion of the diffusion of H2O into the bulk phase.

The

hydration process (MgO bulk phase) was also promoted.

5.4.2. Hydration of pure and chemically-modified MgO: Probable hydration
pathway
The data presented in Tables 5-2–5 reveal that the major peaks at 4300, 4030, 3950, 3790, and3700 cm−1 could be observed in the spectral profiles of all the samples hydrated at 110 ℃ (time: 10min). The peaks at 4030 and 3950 cm−1 appeared during the early stages of the hydration process at
170 and 200 ℃ (after the peak at 3700 cm−1 appeared). The peaks at 4030 and 3950 cm−1 wereassigned to the OH groups present in the stable bulk phase.[8] The peaks at 4300 and/or 3790 cm−1
should appear during the early stages of hydration if the hydration process progresses starting fromthe sample surface (the peaks at 4300 and 3790 cm−1 could not be detected during the early stages ofdehydration [8]).

FT-IR spectra were recorded for samples hydrated at 110 ℃ (PH2O: 31.2 kPa) todetect the order of peak appearance when the effect of hydration temperature could be neglected(Figure 5-11). Analysis of the figures revealed that the peaks at 4030 and 3950 cm−1 appeared after
the peak at 3700 cm−1 appeared.

Following this, the peak at 4300 cm−1 appeared and finally, the peakat 3790 cm−1 appeared. The trend of peak detection was found to be similar for all the samples andwas independent of the hydration temperatures.
Therefore, the hydration of MgO (stable bulk phase) could proceed when the surface hydration
process had proceeded to some extent. These results indicated that the hydration reaction was not
merely a backward reaction of the dehydration reaction involving Mg(OH)2. This can be potentiallyattributed to the fact that the dehydration process resulted in the formation of the defect sites[12,41,42]
117

and cracks. A reduction in the particle volume[10–12] and/or an increase in the specific surface area
could also be achieved.[11,41,43,44] These can potentially promote the process of diffusion of water into
the porous bulk structure.[37] Water diffusion could be ascribed to the drastic hydration conversion
change at the initial hydration stage[37] (Figure S3-5 (a)).
Some researchers have reported the mechanism of MgO hydration. Kato et al. studied the kineticsof the MgO hydration process and assumed that the process could be divided into four stages: (1)containment of water as fixed structural water, (2) physical adsorption of water, (3) chemical reaction
with water producing Mg(OH)2, and (4) the participation of the inert portion of water.[18] Sharma et
al. concluded that the hydroxylation of the dehydroxylated Mg(OH)2 occurs over several steps: In thefirst step, Mg(OH)2 nanocrystal form.

In the second step, rapid intergrowth of the Mg(OH)2

nanocrystals can be observed. The third step involves the rapid formation, mobility, and annealing

Figure 5-11. FT-IR spectral profile recorded for hydrated (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20,and (d) de-L10/LO10 (samples were hydrated at a temperature of 110 ℃; PH2O: 31.2 kPa).

Mg(OH)2-

W, L10, LO20, and L10/LO10 were dehydrated at a temperature of 350 ℃ to prepare de-Mg(OH)2W, de-L10, de-LO20, and de-L10/LO10, respectively.

(Red dotted lines: peak position before

hydration, black dotted lines: peak position post hydration).

118

Figure 5-12. Probable mechanism of hydration for pure MgO. Corresponding FT-IR spectral profiles
have been presented in Figure 5-10.
out of the resulting crystal defects (e.g., edge dislocations) formed during the intergrowth of thenanocrystals.

Fast mobility of the hydroxides can also be observed at this stage.

High surface

mobility of the participating Mg-containing species is observed in the fourth stage. This further
promotes the growth of the Mg(OH)2 crystals.[12]Rocha et al. investigated the MgO hydration

process in the presence of liquid water: (1) water is first adsorbed on the surface. Simultaneous
diffusion of water into the porous MgO particles can also be observed, (2) dissolution of oxide occurs
within the particles. The porosity of the samples changes with time, and (3) nucleation and growthof Mg(OH)2 on the surface of MgO occurs. The surface becomes supersaturated.[37] The author
referred to the previously reported results and the results presented in Figure 5-11 and modeled thehydration reaction pathway of pure MgO (Figure 5-12). The relevant FT-IR spectral profiles havebeen shown in Figure 5-10. Initially, the sample surface adsorbs the water molecules.

Following

this, the surface is hydrated. Subsequently, water vapor diffuses into the porous MgO particles andthe nucleation of Mg(OH)2 occurs.

Finally, the intergrowth of the Mg(OH)2 layer is observed.

A schematic representation of the probable mechanism of the hydration reaction involving
LiCl/MgO has been presented in Figure 5-13 (a). The author assumed that the first step of hydration
involved the adsorption of water on the (111) surface of MgO.[17] The hydration of LiCl was favored
under conditions suitable for the formation of the hydrated LiCl species (e.g., Th = 110 ℃, PH2O = 57.8kPa). The formation of the defect sites (Li+ substitution) promoted the process of surface hydration.
The hydration steps 4 and 5 in LiCl/MgO are the same as the steps 3 and 4 shown in Figure 5-12.
The pathway of hydration was the same as the pathway followed for the hydration of pure MgO under119

unfavorable conditions (Th = 170 or 200 ℃) of forming hydrated LiCl species (Figure 5-12).
A schematic representation of the probable mechanism of the hydration reaction involving
LiOH/MgO has been presented in Figure 5-13 (b).

The overall reaction steps involved in thehydration of LiOH/MgO are similar to the reaction steps involved in the hydration process of pure
MgO.

The first step involves the hydration of the sample.

samples (pure MgO and LiCl/MgO).

This step is different for different

It took 10–20 min for the peak at 3760 cm−1 to disappear when

the sample was de-LO20 (Figure 5-11 (c)) and ~5 min when the sample was de-MgO-W (Figure S314).

Despite the high hydration reactivity exhibited by de-LO20, the time taken for the peak todisappear was longer. As the peak at 3760 cm−1 (ascribed to the hydrogen-covered O-terminated(111) plane) did not disappear during the early stages of hydration, water vapor could diffuse from the(111) plane into the bulk MgO phase. This resulted in an increase in the rate of hydration (steps 3
and 4; Figure 5-12). The increase in the rate can be attributed to the addition of LiOH.
The hydration process of LiCl and/or LiOH-added MgO should be studied in detail using
computational simulation methods.

Figure 5-13. (a) Schematic representation of the predicted mechanism of hydration (for LiCl/MgO).
Steps involved in hydration: (1) Hydration of LiCl. (2) Physical adsorption on the MgO surface.
(3) Surface hydration. (4) Nucleation of Mg(OH)2.

(5) Intergrowth of the Mg(OH)2 layer.

Schematic representation of the predicted mechanism of hydration (for pure LiOH/MgO).
involved in hydration: (1) Physical adsorption on the MgO surface.
Nucleation of Mg(OH)2.

(4) Intergrowth of the Mg(OH)2 layer.

(2) Surface hydration.

(b)

Steps
(3)

Steps (3) and (4) can be promotedby the addition of LiOH.

5.5. Conclusions
The effect of the addition of LiCl and/or LiOH on the hydration reaction and the MgO hydration
mechanism was studied using the XRD and FT-IR techniques.

It was confirmed that the addition of

LiCl and/or LiOH resulted in an increase in the hydration reactivity of MgO. The addition of LiOH
120

promoted hydration at 200 ℃. The author indicated that the addition of Li compounds promoted thereaction occurring on the defect sites.

Substitution using Li+ helped the promotion of the reactions.

Similar observations have been reported in the literature.

The addition of LiCl resulted in theformation of defect sites in the bulk phase. This promoted the hydration of the samples in the bulk
phase of MgO. The addition of LiOH promoted the diffusion of H2O into bulk MgO through thesurface O2− defects. Results obtained from the experiments conducted using the TG technique revealed
that the concentration of OH− can potentially influence the hydration reactivity of the samples. The
hydration reactivity increased with an increase in the LiOH content. Therefore, the addition of LiCland LiOH affected the extent of hydration of LiCl and LiOH-added MgO.
It has been proposed that the hydration reaction (hydration of pure MgO) proceeds over four steps:(1) physical adsorption, (2) surface hydration, (3) nucleation of Mg(OH)2, and (4) inter-growth of the
Mg(OH)2 layer.
The hydration of LiCl (for LiCl/MgO) is favored under conditions of Th = 110 ℃ and PH2O = 57.8
kPa. After that, the MgO hydration reaction proceeded. The initial step involved in the hydrationof LiOH/MgO was different from the initial step involved in the hydration of pure MgO and LiCl/MgO
samples. The results and observations indicated that the diffusion of water from the (111) surface of
MgO into the bulk phase could be promoted.
Further studies on the reaction mechanism involving LiCl and/or LiOH-added MgO should be
conducted. The properties that affect the reactivity of the Mg(OH)2/MgO system should be studied
in detail for the further development of the system. This can also help in designing the materials thatcan be used for storing chemical heat.

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5321–5325.
[2] Shkatulov, A.; Takasu, H.; Kato, Y.; Aristov, Y. Thermochemical energy storage by LiNO3-doped
Mg(OH)2: Rehydration study. J. Energy Storage 2019, 22, 302–310.
[3] Kurosawa, R.; Takeuchi, M.; Ryu, J. Comparison of the effect of coaddition of Li compounds andaddition of a single Li compound on reactivity and structure of magnesium hydroxide. ACS Omega
2019, 4, 17752−17761.
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[5] Ryu, J.; Hirao, N.; Takahashi, R.: Kato, Y. Dehydration behavior of metal-salt-added magnesiumhydroxide as chemical heat storage media. Chem. Lett. 2008, 37, 1140–1141.
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124

Chapter 6. Screening of the Effective Additives on the
Dehydration and Hydration

Reactivity of Magnesium

Hydroxide and Magnesium Oxide System Based on theProperties of Additives

ABSTRACT
Magnesium hydroxide and magnesium oxide (Mg(OH)2/MgO) system is one of the promisingchemical heat storage material for utilizing unused heat at medium temperature of 200–500 °C.
Some previous studies have presented that the addition of dopants or mixing different metal ions to
Mg(OH)2 enhanced the dejhydration and hydration reactivity of Mg(OH)2/MgO system.

Ourprevious studies also revealed that the addition of LiOH or LiCl has different role in the dehydrationand hydration.

Meanwhile the properties of additives related to these reactivities have not beencleared, thus, the rules to enhance the reactivity of Mg(OH)2/MgO system have not been established.
This chapter focused on screening effective additives on the reactivities according to properties ofadditives, such as pH of sample solution, cation and anion radius, solubility.

The results in thischapter observed that the liner relationship between the dehydration peak temperature of Mg(OH)2and the solubility of the additives.

For the enhancement of the hydration reactivity of MgO, theaddition of hydroxides was an effective way because a site easily hydrated could be generated.

6.1. Introduction
The chemical heat storage technique has gained focus for our study as mentioned previous chapters.
The author has focused on the magnesium hydroxide/magnesium oxide (Mg(OH)2/MgO) system tostorage low-medium temperature (200–500 °C) heat as a chemical heat storage material.

The

dehydration of Mg(OH)2 corresponds to heat storage operation accompanied with an endothermic
reaction while the hydration of MgO corresponds to heat output operation accompanied with an
exothermic reaction. The previous studies have reported that mixed hydroxide with Mg(OH)2 andCo(OH)2 or Ni(OH)2,[1] and the addition of LiCl and/or LiOH to Mg(OH)2 enhanced the dehydration
reaction rate at 270–300 °C[2–5] because the rate of pure Mg(OH)2 below 300 °C was poor.

In

particular, the addition of LiCl and LiOH dramatically enhanced the hydration reactivity of MgO at
125

200 °C.[5]Therefore, the addition of LiCl and LiOH is obviously effective way to increase thereactivity of the Mg(OH)2/MgO system.

In addition, the role of the addition of LiCl or LiOH on thedehydration and hydration of Mg(OH)2/MgO system has been reported.[6,7]
However, Li is relatively rare metal compared with Na, Mg, K, and so on, probably leading tosoring price.

In addition, Li-ion batteries have been developed for energy conservation.

Other

alternative salts, therefore, should be investigated for the Mg(OH)2/MgO system. Shkatulov andAristov[8] investigated various salts (chlorides, acetates, sulfates, and nitrates) for the Mg(OH)2/MgO
and Ca(OH)2/CaO systems, showing that Li compounds of the salts decreased the onset dehydration
temperature of Mg(OH)2. They investigated the correlation between the three properties of additives
and the onset temperature. The first investigation was whether the salt contains oxygen resulting incation impurity defect. The second was whether the Tammann temperature at which the mobility
and reactivity of the molecules/atoms in a solid become appreciable was lower than the dehydration
temperature of pure Mg(OH)2.

The third was whether the salt is highly hygroscopic.[8]properties, however, are not universal for effective salts.

These

Chapter 4 reported that the role of LiCland LiOH on the dehydration of Mg(OH)2.[6] The chapter reported that the hydration enthalpy andmobility of Li+, K+, and Na+ did not the main factor and pH of sample solution did not show clear
relationship on the dehydration peak temperature.[6] Meanwhile, the cation radius can be main factor
for effective salt because the radius of Li+ and Mg2+ are close with each other.[6,9,10]
For the hydration reactivity of MgO, Chapter 5 showed a possibility that the number of OH−
affects the reactivity.[7] Actually, the hydration of MgO with water solution or water vapor promotedwith increasing the amount of OH−.[7,10]
To the contrary, the clear parameter for designing materials have not been revealed for
Mg(OH)2/MgO system. Therefore, this chapter aims to find reasonable guideline for enhancing thereactivity of Mg(OH)2/MgO system.

In this chapter, the author focused on pH of sample solution,

cation and anion radius, and solubility for effective salts on the dehydration, and pH and concentrationof OH− for the hydration. In addition, whether the scheme previously proposed by us can be adapted
any materials or not was investigated.

6.2. Experimental
6.2.1. Sample preparation
LiCl·H2O (99.9% purity), LiOH·H2O, LiBr·H2O (99.5% purity), Li2CO3, NaCl, NaOH, Na2CO3,
NaBr, KCl, KOH, K2CO3, KBr, and Mg(OH)2 (99.9% purity, 0.07 µm) were purchased from Wako
Pure Chemical Industries, Ltd. (Japan), and were used as precursors to prepare salt-added Mg(OH)2.
Salt-added Mg(OH)2 (salt/Mg(OH)2) were prepared using the impregnation method as shown in126

previous chapters. Salt-added Mg(OH)2 with a Mg(OH)2:salt molar ratio of 100:10 was preparedusing this method.

For a hydroxide- (LiOH, NaOH, or KOH) or LiBr-added Mg(OH)2, the mixing

ratios were 100:10, 20, and 30 to investigate the effect of added amount on the hydration of MgO.
For comparison, Li-free Mg(OH)2 (hereafter, Mg(OH)2-W) was prepared using the same method.
The mixing ratios and sample names of all prepared samples are listed in Table 6-1.
Table 6-1. Mixing ratios and sample names of the prepared samples.
SampleMixing ratio [mole ratio]

Mg(OH)2-W

Mg(OH)2 without Li compounds

L10

Mg(OH)2:LiCl = 100:10

LO10, 20, 30

Mg(OH)2:LiOH = 100:10, 20, 30

LC10

Mg(OH)2:Li2CO3 = 100:10

LB10, 20, 30

Mg(OH)2:LiBr = 100:10, 20, 30

N10

Mg(OH)2:NaCl = 100:10

NO10, 20, 30

Mg(OH)2:NaOH = 100:10, 20, 30

NC10

Mg(OH)2:Na2CO3 = 100:10

K10

Mg(OH)2:KCl = 100:10

KO10

Mg(OH)2:KOH = 100:10

KC10

Mg(OH)2:K2CO3 = 100:10

KB10

Mg(OH)2:KBr = 100:10

6.2.2. Sample characterization by XRD
XRD measurements were performed to investigate the crystal structure of the samples using an
Ultima IV (Rigaku Corp., Japan) XRD instrument in air at r. t., in the 2θ range of 10−150°, at a scanrate of 1.0° min−1, and at a scan width of 0.01°. The XRD profiles of the prepared samples were used
to confirm their successful preparation. To avoid the subsequent hydration of the samples during
measurements, the reacted samples were scanned after reaction at a scan rate of 10° min−1 using Cu
Kα radiation at a voltage of 40 kV and a current of 40 mA. As the 2θ range of 10–80° was sufficient
to characterize Mg(OH)2, the data were described in this range.

Before the measurements, theprepared samples were dried at 120 °C overnight to remove the physically adsorbed water.

6.2.3. Sample characterization by FT-IR
To gain additional information on the dehydration of pure and doped Mg(OH)2, and the hydrationof pure and doped MgO, FT-IR measurements were performed using an FT/IR-4200 (JASCO Corp.,
Japan) spectrometer in diffuse reflectance mode in air at r. t.

The background for the FT-IR

measurements was calibrated using CaF2 powder. The scans were obtained in the wavenumber range127

of 7800–350 cm−1, at a resolution of 4.0 cm−1, and all results were reported as averages of 64 scans.
Similar to the XRD measurements, the prepared samples were dried at 120 °C overnight prior to the
FT-IR measurements to remove any physically adsorbed water.

Because the author was mainlyinterested in the peaks of Mg(OH)2, in this study, the author focused on the FT-IR spectra in thewavenumber range of 5000–2750 cm−1.

6.2.4. Dehydration reaction test
All samples were heated using a TGD–9600 series (ADVANCE RIKO, Inc., Japan) thermobalance.
The samples (~20 mg) were added to Pt cells and were subsequently heated from room temperature(r. t.) to 600 °C at a rate of 10 °C min−1 under an Ar flow of 100 mL min−1 for dehydration reaction
test. In this study, the dehydration peak temperature was used to evaluate the dehydration reactivityof the samples as shown in the Experimental section in Chapter 3.

6.2.5. Hydration reaction test
For the hydration reaction, to remove physically adsorbed water, the sample was heated at 120 °C
for 30 min at a rate of 20 °C min−1 under Ar flowing at100 mL min−1.

The sample was thendehydrated at 350 °C for 30 min at a heating rate of 20 °C min−1 under Ar flowing at 100 mL min−1.
The sample was then hydrated at constant hydration temperatures (200 °C) for 80 min at a heating rateof −20 °C min−1 by introducing a gas mixture of water vapor at 57.8 kPa and Ar gas. Subsequently,
the introduction of water vapor was stopped and the sample was dried at the hydration reaction
temperature for 30 min.
The aauthor evaluated the reactivity of the samples as shown in the Experimental section in
Chapter 3. The sample dehydrated at 350 °C before the hydration reaction is described as a de-sample(for example, Mg(OH)2-W dehydrated at 350 °C before the hydration is referred to as de-Mg(OH)2).

6.2.6. pH measurement of sample solution
The author determined the pH of LiCl, LiOH, Li2CO3, LiBr, NaCl, NaOH, Na2CO3, NaBr, KCl,
KOH, K2CO3, and KBr solutions using a conventional pH meter D-50 series (HORIBA Ltd., Japan).
The concentration of the solutions was fixed at 0.04 mol/L because the concentration was close to thatof the sample prepared by the impregnation method.

128

6.3. Results
6.3.1. Characterization by XRD and FT-IR measurements for as-synthesized,
samplesFigure 6-1 shows the XRD patterns of the as-synthesized samples. The diffraction patterns foundthe Mg(OH)2 Li2CO3, NaCl, Na2CO3, NaBr, KCl, K2CO3, and KBr phases. The LiCl, LiOH, LiBr,
NaOH, and KOH phases were undetectable, meaning that those species were dissolved in the
Mg(OH)2 phase, were well dispersed throughout the Mg(OH)2 particles, and were converted intoaqueous solution, which rendered them undetectable, or because a compound and Mg(OH)2 solid
solution formed.[3–7] In addition, Figure 6-1 (a) does not detect the peaks of LiOH and LiOH·H2O;
however, the author confirmed those peaks previously reported,[4] the author concluded that LiOH or
LiOH·H2O could exist on the surfaces of the LiOH-added samples (LO10).[4–6]

LiOH might bepartially converted to Li2CO3 during preparation, which confirmed that LiOH was typicallycontaminated with Li2CO3,[12] and accounted for the presence of the Li2CO3 peak in the XRD profilesof the LO10. Na2CO3 and K2CO3 phases were also detected in those diffraction patterns of NO10
and KO10. It is probably because of the same reason for detection of Li2CO3 in the patterns of LO10.
Figure 6-2 shows the FT-IR spectral of the as-synthesized samples. The peak at 3700 cm−1 canbe assigned to the surface hydroxyl (OH) groups in the corner sites and/or edges with low coordination
numbers.[13–17] In addition, the peaks at 4311, 4034, 3953, and 3793 cm−1 can be derived from OH
groups present in Mg(OH)2 bulk phase.[6] Typical broad band at 3600–3200 cm−1 is assigned to thehydrogen-bonded OH groups due to water molecules adsorbed on the sample surface.[18–20]

Theintensity of the broad band for as-synthesized L10, LB10, NB10, KO10, and KC10 were comparablyhigh.

It can be because the additives of those sample (LiCl, LiBr, NaBr, KOH, and K2CO3) usually

have high hygroscopicity. Hence, these results indicate that a large volume of water adsorbed and/or
hydrated species of those additives were formed on those samples surface.

6.3.2. Effect of the additives on the dehydration of Mg(OH)2
Figure 6-3 (a) shows the dehydration behaviors of pure and a Li compound-added Mg(OH)2
heated from r. t. to 600 °C under Ar flowing of 100 mL min−1. This result shows that L10 and LB10
were dehydrated at dramatically low temperatures compared with Mg(OH)2-W, which agrees with theprevious studies. The addition of Li2CO3 did not affect the dehydration compared with the additionof LiCl or LiBr. LO10 was dehydrated at slightly low temperatures compared with Mg(OH)2-W,
meanwhile, the addition of LiOH lowered the dehydration temperature range of Mg(OH)2 withincreasing LiOH amount as described in Chapter 2.[4]Figure 6-3 (b) shows the dehydration

behaviors of pure and a Na compound-added Mg(OH)2 heated from r. t. to 600 °C under Ar flowing
129

Figure 6-1. XRD patterns of as-synthesized (a) Li compounds-added, (b) Na compounds-added,and (c) K compounds-added Mg(OH)2.

Figure 6-2. FT-IR spectra of as-synthesized (a) Li compounds-added, (b) Na compounds-added,and (c) K compounds-added Mg(OH)2.
130

Figure 6-3. Dehydration behavior of (a) Li compounds-added, (b) Na compounds-added, and (c) K
compounds-added Mg(OH)2 heated from r. t. to 600 °C under Ar flowing of 100 mL min−1.
of 100 mL min−1. The NO10 was dehydrated in a low temperature compared with other samples,while the N10, NC10, and NB10 were dehydrated in slightly low temperatures compared with
Mg(OH)2-W. Figure 6-3 (c) shows the dehydration behaviors of pure and a K compound-added
Mg(OH)2 heated from r. t. to 600 °C under Ar flowing of 100 mL min−1.

In this case, the KO10 was

dehydrated in dramatically low temperatures compared with other samples, meanwhile the K10, KC10,
and KB10 were dehydrated in slightly low temperatures compared with the Mg(OH)2-W.
Table 6-2 shows the dehydration peak temperature of the samples, which is derived from thedifferential thermogravimetry (DTG) curve (Figure 6-4).
temperature at which the dehydration rate is the fastest.

[5]

The peak temperature means the

The dehydration rate was calculated fromthe measured weight change and temperature, and is the sample weight differentiated by thetemperature.

A lower peak temperature means that the sample can be dehydrated at lower

temperature, which indicates that the sample is suitable for lower-temperature heat storage. These
results present the addition of LiCl, LiBr, NaOH, and KOH dramatically lowed the peak temperature.
In the later section, the properties of the compounds related to the peak temperature is discussed.
Figures 6-4 (b) and (c) show the two obvious peaks for NO10 and KO10, which means that the131dehydration of Mg(OH)2 proceeded as two steps reaction. The dehydration of LiOH-added Ca(OH)2
previously presented the two steps reaction.[21] The characteristic peaks were not detected for DTG
curve of LO10 (Figure 6-4 (a)).

On the other hand, the DTG curve of LO20 (Figure 6-4 (d))

demonstrated a shoulder peak. Therefore, the dehydration process of Mg(OH)2 can be changed by
the addition of hydroxides. Chapter 3 showed that the diffraction peaks of MgO for LO20 weredetected at marked lower temperatures than that for Mg(OH)2-W. Chapter 4 also presented that theaddition of LiOH promoted the surface dehydration of Mg(OH)2.

These results indicate that theonset temperature of the dehydration for hydroxides-added Mg(OH)2 was lowered. In addition, theseresults can be interpreted that the first peak corresponds to the surface dehydration while the second
peak corresponds to the bulk dehydration. The further studies, however, are required to clarify thereason that the dehydration process can be changed by the addition.

Figure 6-4. Differential thermogravimetry (DTG) curve of (a) Li compounds-added, (b) Na
compounds-added, (c) K compounds-added Mg(OH)2, and (d) LO20 heated from r. t. to 600 °C under
Ar flowing of 100 mL min−1.

132Table 6-2. Dehydration peak temperature of the Li, Na, or K compounds-added samples. The amountof added Li compounds was 10 mol%.
SamplePeak temperature [°C]

Mg(OH)2-W

376

L10

319

LO10

357

LC10

373

LB10

305

N10

358

NO10

337

NC10

364

NB10

350

K10

361

KO10

334

KC10

346

KB10

357

6.3.3. Effect of the additives on the hydration of MgOFigure 6-5 (a) shows the hydration behaviors of pure and a Li compound-added MgO. This
presents that the addition of LiOH or LiBr promoted the hydration of MgO at 200 °C compared withthe addition of LiCl or Li2CO3. Among the Na, K compounds, the addition of NaOH or KOH
dramatically enhanced the hydration reactivity (Figure 6-5 (b), (c)). The extent of reduction of themole fraction after the completion of hydration for de-KO10, however, was significant (Figure 6-5(c)). Therefore, the hydration of MgO itself seemed not to proceed efficiently for de-KO10. These
results indicate that the addition of hydroxides to MgO has positively large effect on the hydration of
MgO.

6.4. Discussion
6.4.1. Relationship between the dehydration peak temperature and the properties of
the additives
The relationship between the dehydration reactivity and properties of samples is required to beinvestigated to understand how the properties affect the reactivity and efficiently design a material fora chemical heat storage. In this section, the relationship between the dehydration reactivity and cationor anion radii, solubility of the additives, and pH of additive solution were discussed.
133

Figure 6-5. Hydration behavior of (a) Li compounds-added, (b) Na compounds-added, and (c) K
compounds-added MgO. All the samples (hydroxides) were dehydrated at 350 °C for 30 min, and then,hydrated at 200 °C under water vapor (PH2O) of 57.8 kPa.
Table 6-3 shows the pH of the sample solution used for the sample preparation. Figure 6-6 (a)shows the relationship between the dehydration peak temperature and pH of the sample solution.
This figure does not show the clear relationship, hence, the pH of the sample solution is not a main
factor related to the dehydration reactivity.

This is in accordance with results in Chapter 4.[6]Figure 6-6 (b) shows the relationship between the dehydration peak temperature and the cation radius.
Chapter 4 [6] indicated that the cation radius was one of the main factors related to the sample
dehydration reactivity because the ionic radius of Mg2+ (0.72 Å[22]) and that of Li+ (0.76 Å[22]) is close.
It could lead to the efficient substitution by a Li+ into a Mg2+ site. Therefore, the additives withsimilar cation radius compared with Mg2+ can possibly affect the dehydration reactivity.

However,

Figure 6-6 (b) does not present the clear relationship between the dehydration reactivity and the cation
radii. For example, the addition of a compound with larger radius compared ith Mg2+ such as NaOH
or KOH (the ionic radii of Na+ and K+ are 1.02 Å and 1.38 Å,[22] respectively) lowered the peaktemperature compared with Mg(OH)2-W.
Figure 6-6 (c) shows the relationship between the dehydration peak temperature and the anion
134

radius. Chapter 4 has reported the possible scheme of the affect of the addition of LiCl on thedehydration of Mg(OH)2.[6] The scheme indicated that a Cl− whose ionic radius (1.72 Å[23]) was
larger than that of OH− (1.33 Å[23]) formed distorted sites on the Mg(OH)2 surface, leading the Li+
penetration into bulk Mg(OH)2.

This presumably promoted the dehydration in the bulk phase.

Thus, the additives with larger anion radius compared with OH− can possibly promote the dehydrationof Mg(OH)2.

The dehydration peak temperature of LB10 (Mg(OH)2 with added LiBr) was

dramatically lowered compared with that of Mg(OH)2-W (Figure 6-3 (a)). This can be explained bylarge ionic radius of Br− (1.88 Å[23]).

On the other hand, Figure 6-3 (a) and Table 6-2 show that thepeak temperature of LC10 hardly changed compared with that of Mg(OH)2-W (the radius of CO32−
is 1.78 Å[23]). This result can be inconsistent with the assumption that the large anion radius affects
the dehydration. Therefore, the possible scheme reported in the Chapter 4 can be limited only tocertain samples or conditions.
Table 6-3. pH of the sample solutions
SamplepH of sample solution [-]

LiCl

5.8

LiOH

12.34

Li2CO3

11.21

LiBr

6.49

NaCl

5.75

NaOH

12.54

Na2CO3

11.21

NaBr

5.57

KCl

5.7

KOH

12.42

K2CO3

11.22

KBr

5.53

135

Table 6-4. Solubility of the additives.

Superscripts 0, and 20 mean the temperature.[24–26]

Solubility [g 100 mL or g-

Solubility [mol 100 mL or g-water−1]water−1]

LiCl

670[24]

1.58

LiOH

12.820[25]0.53

Li2CO3

1.540[24]

0.02

LiBr

142.7

0[24]

1.64

NaCl

35.70[24]

0.61

NaOH

420[24]

1.05

Na2CO37.10[24]

0.07

NaBr

79.50[24]

0.77

KCl

27.6

0[24]

0.37

KOH

970[24]

1.73

K2CO3

1050[24]

0.76

KBr

53.480[24]

0.45

LiNO3

47.5825[24]

0.69

LiOAc

40.820[26]

0.62

NaOAc

5020[26]

0.61

KOAc

269.420[26]2.74

Additive

Figure 6-6 (d) shows the relationship between the dehydration peak temperature and the solubilityof the additives. The solubility is described as the value at 0 °C except for LiOH. This figure shows
the liner relationship between them. The addition of the additives with higher solubility lowered thedehydration peak temperature.
These results suggest that the solubility of the additive can be a factor that effectively decide thedehydration reactivity of the sample. Here, the reason that the solubility can affect the dehydration
reactivity will be discussed. For the sample preparation by the impregnation method, the additives,
first of all, were dissolved into water to prepare an aqueous solution, secondly, the Mg(OH)2 powder
was impregnated with the solution.

At that time, Mg(OH)2 could be slightly dissolved into water.

Therefore, the higher solubility an additive has, the easier a metal cation can be substituted into a Mg2+
ion site as reported by the literatures.[4–10] This assumption indicates that when a compound withhigh solubility is used as an additive, the possible scheme reported in Chapter 4[6] can be adapted.
The previous study reported by Shkatulov and Aristov[8] has shown that the dehydration ofMg(OH)2 with added LiNO3, or acetates of Li, Na, or K (LiOAc, NaOAc, and KOAc) proceeded insignificantly low temperatures compared with pure Mg(OH)2. The solubility of LiNO3, LiOAc, and136

Figure 6-6. Relationship between the dehydration peak temperature and (a) pH of sample solution, (b)Shannon′s cationic radius, (c) Anion radius, and (d) solubility of the additives. The solubility wasdescribed as the value divided the solubility by mole weight. The orange dotted line means the peaktemperature of Mg(OH)2-W. The cation and anion radii were cited from previous literatures.[22,23]
NaOAc, however, are not drastically high (Table 6-4). Therefore, solubility is not the only factor
describing the sample dehydration reactivity.

Meanwhile, as described above, solubility can be aguide for designing chemical heat storage materials.

6.4.2. Properties related to hydration reactivity of MgOFigure 6-5 shows that the addition of LiOH, LiBr, NaOH, and KOH enhanced the hydration
reactivity of MgO compared with the addition of other additives. The pH of sample solution did notaffect on the hydration because the hydration reactivity of MgO with added salt carbonates was low(the pH of salt carbonates solution was high (approximately 11)).
Here, the FT-IR measurement was carried out before and after the hydration reaction to analyze
the effect of the addition of hydroxides (LiOH, NaOH, and KOH) and LiBr on the hydration of MgO.
Figure 6-7 shows the FT-IR spectra of the dehydrated samples (the samples were prepared by adehydration at 350 °C for 30 min).

The spectra of de-LO10, de-NO10, and de-KO10 presented
137

Figure 6-7. FT-IR spectral of (a) Li compounds-added, (b) Na compounds-added, and (c) K
compounds-added MgO. The Li, Na, or K-added Mg(OH)2 were dehydrated at 350 °C, furthermore,heated at 200 °C for 10 min. Consequently, the FT-IR measurements were carried out at r. t. in air
for the dehydrated samples.
characteristic shoulder peak at 3760 cm−1 (shown as red dotted line in the figure). This shoulder peak
could be derived from OH groups on MgO (111) surface.[27] Some previous studies reported that the(111) plane is easily hydrated than (100) MgO surface.[28–31] This is in agreement with the TG results
showing that a hydroxide- added MgO was easily hydrated at 200 °C (Figure 6-5). On the other
hand, the FT-IR spectra of de-LB10, which shows high hydration reactivity compared with deMg(OH)2-W, did not form the shoulder peak.

Meanwhile, the fact that the de-LB10 showed the highhydration reactivity suggest that other factors affect the hydration. Therefore, the MgO (111) surface(described as the FT-IR peak at 3760 cm−1) seems to be an important factor for a sample easily hydrated,in addition, other factors affected the hydration of MgOChapter 5 concluded that the concentration of OH− can affected the hydration reactivity of MgO.
Here, the dependence of the concentration of OH− on the hydration reactivity is firstly investigated(the effect of the amount of added KOH will not mentioned because the hydration behavior of deKO10 was too unique to investigate).

Figure 6-8 (a) and (b) show that the hydration reactivity of138

MgO was enhanced with increase of the amount of added LiOH or NaOH.

In Chapter 5, limited

Mg sites on Mg(OH)2 surface could be substituted by Li when LiOH was doped. Therefore, these2++

results indicate that the OH− in hydroxides shifted the equilibrium to the formation of Mg(OH)2.[11]Hence, the results in Figures 6-5 and 6-7 suggests that the exposure of MgO (111) surface and highconcentration of OH− are desirable for the high extent of hydration reactivity of MgO.
Actually, the fact that LiBr/MgO presented the high hydration reactivity compared with deMg(OH)2-W suggests that another factor affected on the hydration.

Figure 6-9 shows thedehydration and hydration behavior of LiCl∙H2O, LiOH∙H2O, and LiBr∙H2O. This figure shows that
LiCl and LiOH were hardly hydrated at 200 °C while LiBr was much adsorbed compared with othertwo compounds.

The hydration of LiBr (the formation of hydrated species of LiBr) presumably

promoted the hydration. This result can be interpreted that the additives with high hygroscopicity
such as LiBr, CaCl2, and MgCl2 are suitable for enhancement of the hydration reactivity. On theother hand, the diffraction patterns of LiBr/MgO (Figure 6-10 (c)) did not show any Mg(OH)2 phases,
meaning that the bulk hydration of MgO has not effectively proceeded. Only the hydration of LiBrand adsorption of water molecule on sample surface were favoured during the hydration period.

Figure 6-8. Hydration behavior of (a) LiOH/MgO, (b) NaOH/MgO, and (c) KOH/MgO. All thesamples (hydroxides) were dehydrated at 350 °C for 30 min, and then, hydrated at 200 °C underwater vapor (PH2O) of 57.8 kPa.
139

Figure 6-9. Hydration behavior of the Li compound monhydrates. All the samples were dehydrated
at 350 °C for 30 min, and then, hydrated at 200 °C for 80 min under water vapor (PH2O) of 57.8 kPa.

Figure 6-10. XRD patterns of (a) LiOH/MgO, (b) NaOH/MgO, and (c) LiBr/MgO after hydration at
200 °C. All the samples (LiOH/Mg(OH)2, NaOH/Mg(OH)2, and LiBr/Mg(OH)2) were dehydrated
at 350 °C for 30 min, and then, hydrated at 200 °C for 80 min under water vapor (PH2O) of 57.8 kPa.
140

For the hydration reactivity of MgO, the author observed that the addition of a hydroxide as anadditive to MgO is one of the suitable methods to enhance the reactivity. The addition of hydroxides
could effectively generate surface defects as shown in Chapters 4 and 5. The defects generated byhydroxides might lead the water diffusion into bulk MgO, which promoted the hydration of MgO.

6.5. Conclusions
This chapter has focused on the relationship between the dehydration/hydration reactivity ofMg(OH)2/MgO with added salts and salts properties such as, solubility, cation radius, anion radius,
pH of sample solution, and concentration of OH−. It can be necessary to understand the factor thatis related to the dehydration and hydration reactivity to built the design rule for chemical heat storage
materials. This chapter revealed that the liner relationship was observed in the relationship between
the dehydration peak temperature and the solubilty of additives.

It can be because the higher

solubility an additive has, the easier a metal cation can be substituted into a Mg2+ ion site when thesample is prepared by the impregnation method. Therefore, the compounds with high solubility aresuitable for enhancement of the dehydration reactivity of Mg(OH)2. The results presented in thischapter indicated that the previously proposed scheme for LiOH- or LiCl-added Mg(OH)2 can belimittd only to certain samples and/or conditions.
The hydration reactivity of MgO was enhanced by the addition of hydroxides (LiOH, NaOH, andKOH) or LiBr.

The FT-IR spectra of dehydrated LiOH-, or NaOH-added Mg(OH)2 showed thecharacteristic peak at 3760 cm−1 assigned as hydroxyl groups on MgO (111) plane. Some previous
studies have reported that the MgO (111) plane can be easily hydrated than MgO (100) plane.
Therefore, the results shown in this chapter is consistent with these previous studies.
number (concentration) of OH− could affect on the hydration reactivity.

In addition, theMeanwhile, a highhygroscopic compound such as LiBr prmomoted the hydration of MgO because the conpound adorbed
water, which interacts with the MgO surface.

The hydration of MgO itself, however, did noteffectively proceed as presented in the diffraction patterns. Hence, the addition of a hydroxide canbe an effective way to enhance the hydration reactivity of MgO.
Recently, the machine learning has gained focus on the material science for the material design by
prediction.[32] Kiyabu et al.[33] screened hydration reactions for thermal energy storage by density
functional theory calculations.

These approarches based on computational methods are getting

powerful and important tools. Therefore, the results in this chapter can be adapted to the methods
and help to design materials by prediction.

141

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144

Chapter 7. Conclusions and future issues
Chapter 1 mentioned the general introduction about the global environmental problems andthermal energy storage techniques.

Currently, there have been much unused thermal energy.

Therefore, the thermal energy storage techniques are promising for energy conservation and utilization.
Chapter 1 showed that the chemical heat storage technique was one of the potential techniques forenergy conservation owing to its high energy density and unlimited heat storage period compared withsensible and latent heat storage techniques.

In addition, the amount of unused heat at 100–300 °C isrelatively large in Japan, hence, the Mg(OH)2/MgO system has been focused for this study. The
dehydration temperature range of Mg(OH)2 is typically at the temperature of 300–400 °C, and thehydration reaction temperature range is at the temperature of 100–200 °C.

This study aimed toenhance the reactivity of the Mg(OH)2/MgO system by chemically modification and investigate theeffect of the chemical modification and the reaction mechanism.
The effect of the addition of LiOH on the dehydration of Mg(OH)2 was investigated in Chapter
2.

LiOH was selected as a new additive to enhance the dehydration reactivity. The LiOH/Mg(OH)2

samples were effectively dehydrated compared with pure Mg(OH)2. The isothermal dehydration testat 270 °C showed the rate of dehydration of LiOH/Mg(OH)2 samples was higher than that of pure
Mg(OH)2, meaning that the addition of LiOH is an effective way to enhance the dehydration reactivity.
The dehydration reaction kinetics by integration method revealed that the dehydration behavior of
LiOH/Mg(OH)2 could be described by zero- and first-order reactions.

In addition, the calculated

activation energy for the dehydration based on Arrhenius plot presented that the energy of LO20
(Mg(OH)2 with added LiOH of Mg(OH)2:LiOH=100:20 mol%) was the lowest of all the
LiOH/Mg(OH)2 samples. Therefore, this chapter concluded that the LO20 was the best mixing ratiowith respect to the activation energy.
The Mg(OH)2 with co-added LiCl and LiOH was proposed and the reactivity of the sample wasinvestigated in Chapter 3. Mg(OH)2 with co-added LiCl and LiOH (L10/LO10; (Mg(OH)2 withadded LiCl and LiOH of Mg(OH)2:LiCl:LiOH=100:10:10 mol%) showed the dehydration in lower
temperature ranges than Mg(OH)2 with singly added LiCl or LiOH. Furthermore, the co-addition of
LiCl and LiOH dramatically enhanced the hydration reactivity of MgO. To the contrary, another way
to enhance the dehydration reactivity of Mg(OH)2 is required for the efficient heat storage at 200–
250 °C. The XRD patterns of dehydrated pure Mg(OH)2 showed the appearance of MgO phase at
320 °C and the disappearance of Mg(OH)2 phase at 380 °C. The patterns of dehydrated Mg(OH)2
with added LiOH showed the appearance of MgO phase at 240 °C and those of dehydrated Mg(OH)2
with added LiCl showed the disappearance of Mg(OH)2 phase at 320 °C. These results indicated that145

the addition of LiOH or LiCl has different roles in the dehydration of Mg(OH)2. In particular, theaddition of LiCl possibly promoted Mg(OH)2 bulk decomposition based on the XRD results.
The investigation in Chapter 4 discussed the effect of the addition of LiCl and LiOH on thedehydration of Mg(OH)2 by using FT-IR and XRD measurements.

The FT-IR measurement

demonstrated that the spectra of dehydrated LO20 and MgO were significantly similar, indicating thatthe surface state of both samples were close with each other. The results in this chapter indicatedthat the addition of LiOH promoted the surface dehydration of Mg(OH)2 while that of LiCl promoted
the dehydration in bulk phase. The surface OH− defectes could efficiently be generated by LiOH
addition. When LiCl was added to Mg(OH)2, the OH− defectes could be generated in bulk phase.
The Cl− of LiCl probably made distorted sites on Mg(OH)2 surface, leading to the penetration of Li+
into the bulk phase.

In addition, the author assigned the absorption bands at 4300–3790 cm−1 as bulk

OH groups of Mg(OH)2, which have not been discussed ever.
Chapter 5 investigated the role of LiCl and LiOH on the hydration of MgO based on FT-IR and
XRD measurements.

These measurements confirmed the high hydration reactivity of MgO withadded LiCl and/or LiOH compared with pure MgO.

In particular, the measurements clearly showedthat MgO with added LiOH had drastically high hydration reactivity at 200 °C. When LiOH wasadded to Mg(OH)2, surface OH− defects could be efficiently generated based on the Chapter 4, which
led to the formation of surface oxygen vacancies after the dehydration. This can promote the water
diffusion into bulk MgO instead of surface hydration, promoting the formation of Mg(OH)2. When
LiCl was added to Mg(OH)2, OH− defects in bulk phase could be efficiently generated based on the
Chapter 4, which led to the formation of oxygen vacancies in bulk phase after the dehydration. This
could promote the hydrion of bulk MgO. The rate of bulk hydration, however, could be dependent
on the rate of surface hydration and/or water adsorption on surface because the hydration reactivity of
MgO with added LiCl at 200 °C was poor. Therefore, the water diffusion into bulk phase and thehydration of bulk MgO could be simultaneously promoted for the MgO with co-added LiCl and LiOH.
Chapters 4 and 5 investigated the role of the addition of LiCl and LiOH on thedehydration/hydration of Mg(OH)2/MgO system.

On the other hand, for the design ofMg(OH)2/MgO chemical heat storage material, the effective way to enhance the reactivity ofMg(OH)2/MgO system has been required.

Hence, Chapter 6 screened the suitable additives toenhance the rectivity of the system. The addition of LiCl, LiBr, NaOH, or KOH was effective way
to lower the dehydration peak temperature. The diffraction patterns of Mg(OH)2 with added LiCl,
LiBr, NaOH, and KOH did not detect the patterns of those additives, which can mean a solid solution
could be produced between Mg(OH)2 and those additives. The addition of LiOH, LiBr, NaOH, and
KOH efficiently enhanced the hydration reactivity of MgO. The investigation of the relationshipbetween the dehydration peak temperature and pH of sample splution, cationic radius, anionic radius,and solubility of additives showed the liner relationship between the peak temperature and solubility146

of additives (the higher solubility an additive has, the lower dehydration peak temperature is achieved
for a sample with the additive). Therefore, the solubility can be a guide to determine an effective toenhance the dehydration reactivity of Mg(OH)2. The results in this chapter presented the addition ofhydroxides (LiOH, NaOH, or KOH) is one of the efficient ways to increase the hydration reactivity.
In this study, Mg(OH)2/MgO system was investigated as a chemical heat storage material. The
co-addition of LiCl and LiOH dramatically enhanced the reactivity of Mg(OH)2/MgO system. The
discussion on the role of LiCl and LiOH in the dehydration and hydration is main part for this study
because the investigation related to the effect has not been carried out in detail.
Furthernore, considering the sample cost, Li can cause the sample to be expensive.

In addition,

Li ion batteries have gained focus for next energy conservation technique. Therefore, the amount ofadded Li to the system should be reduced. The results in Chapters 4, 5, and 6 can propose a newmaterial for chemical heat storage based on a prediction for future research. These chapters can help
the author to choose inexpensive and abundant material for chemical storage. This approarch canalso save time and cost to design the material based on a prediction.

If this approarch can be adapted

to other systems (Ca(OH)2/CaO, MgCO3/MgO, CaCO3/CaO, and so on), the usage of the chemicalheat storage technique will be promoted as an energy conservation technique. Furthermore, theMg(OH)2 with binary salts system (likewise, L10/LO10) for the heat storage at 200–250 °C can beproposed using the results obtained in this study.
This study has focused on the Mg(OH)2/MgO system from the view points of physical chemistry.
The points of engineering are required for the practical use.

For example, mechanical properties suchas cycle stability, morphology, and pressure strength of the sample should be investigated.
Furthermore, the investigation on combination with a reactor, for instance, the packed bed reactor or
the continuous reactor, is needed, meaning that consideration for which reactor is suitable and thereaction condition are necessary.

147Supporting Information
S1. Supporting information for Chapter 3
In Figure 3-3, the mole fraction of Mg(OH)2 of the samples were not achieve 0% at 600 °C.
Because of that, the author carried out the XRD and IR measurement to confirm if the existence ofresidual Mg(OH)2 after heating at 600 °C as shown in Figures S1-1 and S1-2. Figure S1-1 showsno peak of Mg(OH)2.

However, the peak around 3700cm−1 (black dotted line) corresponding toisolated OH peak on Mg(OH)2 surface was observed such as in Figure S1-2. This peak probably
attributed to the incomplete dehydration of Mg(OH)2-W in Figure 3-3. However, the dehydration
conversions of all samples were more than 90%. Hence, the author considered as the dehydration
reaction was completed
Reaction behavior of L20/LO20 (Mg(OH)2:LiCl:LiOH=100:20:20) is shown in Figures S1-3–5.
There were no significant differences between the dehydration and hydration reactivity of L20/LO20
and L10/LO10 according to Figures S1-2 and 1-3. Moreover, the heat output density of L20/LO20
was not reached 1000 kJ/kg due to large amount of LiCl and LiOH included in the sample.

In

addition, the added Li amount was fixed as 20 mol%, and the effect of LiCl and LiOH ratio on thereactivity of Mg(OH)2 and MgO was investigated. The mixing ratios of LiCl and LiOH were of
LiCl:LiOH=20:0, 18:2, 15:5, 10:10, 5:15, 2:18, or 0:20.
Figure S1-6 shows the dehydration and hydration behaviors of the samples. Figure S1-7 showsthe effect of LiCl and LiOH amount on (a) heat output density by hydration of MgO at 200 °C, (b)hydration reaction conversion at 200 °C, and (c) dehydration peak temperature. The trends for theheat output density and hydration reaction conversion were same because the heat output density wasdetermined by the reaction conversion.

The heat output densites (hydration conversions) for

L10/LO10 and L15/LO5 (LiOH amount of 5 mol%) were especially high meanwhile the peaktemperatures of L10/LO10 and L5/LO15 (LiOH amount of 15 mol%) were relatively low among thesamples. The most optimal amount of LiCl and LiOH are 10–15 mol% and 5–10 mol% for the heat
output density, respectively while that of LiCl and LiOH are 5–10 mol% and 10–15 mol% for the peaktemperature, respectively. With respect to the both heat storage and output operation, L10/LO10 was
the best mixing ratio. Therefore, in Chapter 3, L10/LO10 was studied in detail.

148

Figure S1-1. XRD patterns of all the prepared samples heated to 600 °C

Figure S1-2. IR spectrum for Mg(OH)2-W with heated to 650°C

149

Figure S1-3. Dehydration behavior of all samples with heating to 600 °C

Figure S1-4. Dehydration and hydration behavior of all the samples; Td = 270 °C, Th = 110 °C, PH2O
= 57.8 kPa.

150

Figure S1-5. Heat output density at various hydration temperatures for all the samples; Td = 270 °C,
Th = 110, 170, and 200 °C; PH2O = 57.8 kPa

Figure S1-6. Effect of LiCl and LiOH amount on (a) hydration of MgO at 200 °C, and (b) dehydrationof Mg(OH)2 with heated from r.t. to 600 °C. For the hydration reaction, the samples were firstly
dehydrated at 350 °C, and then they were hydrated at 200 °C under mixture of water vapor and Ar gasof 57.8 kPa. The mixing ratios of LiCl and LiOH were of LiCl:LiOH=20:0, 18:2, 15:5, 10:10, 5:15,
2:18, or 0:20.

151

Figure S1-7. Effect of LiCl and LiOH amount on (a) heat output density by hydration of MgO at
200 °C, (b) hydration reaction conversion at 200 °C, and (c) dehydration peak temperature. Blackdotted lines mean those of Mg(OH)2-W. x axis is the amount of added LiOH in LiCl/LiOH/Mg(OH)2and the amount of added Li was fixed as 20 mol%.

S2. Supporting information for Chapter 4
Figures S2-1 and S2-2 show the FT-IR spectra of L10+CaF2 and Mg(OH)2-W+CaF2, respectively,
after heating from room temperature to 350 °C with heating rate of 10 °C min−1 under an Ar flow of
100 mL min−1.
(r.t.).

First, the samples were heated in a thermobalance and cooled to room temperature

The FT-IR spectra were then recorded at r.t.

in air, as described in the main manuscript(Figures S2-1 (a) and S2-2 (a)). Then, the samples were stored in the IR cell at r.t. in air for (b)10 min, (c) 30 min, and (d) 1 h. L10 was chosen for the investigation of the effect of water becauseof a deliquescent property of LiCl in the air. As shown in Figure S2-1, the intensity of the peak at
3678 cm−1, which is assigned to surface OH groups, was not significantly changed even after exposure
to air at room temperature for 10–60 min. Thus the water vapor in the air did not largely affect theresults of FT-IR measurements.

The typical absorption bands at 2348 cm−1 in Figure S2-2 isassigned to CO2 (νgas) in gas phase.[1] The sharp absorption bands at 1498 and 1408 cm−1 could be152

assigned to CO32− and/or HCO3− species resulted from CO2 adsorption on the MgO surface.[1,2]
However, the intensities of these bands did not change significantly.

Furthermore, no carbonate

species were observed in the XRD patterns of the all samples. Therefore, small amount of carbonate
species adsorbed on the sample surface. Figure S2-3 shows the TG and DTA profiles of Mg(OH)2W [(a) and (d)], hydromagnesite (light form) [(b) and (e)], and hydromagnesite (heavy form) [(c) and(f)].

The endothermic peak at approximately 370 °C for Mg(OH)2-W was attributed to thedehydration of Mg(OH)2. The endothermic peaks at approximately 260 °C and 460 °C arose fromthe dehydration and decarbonation of hydromagnesite, respectively,[3] while Mg(OH)2-W showed noendothermic peak at 460 °C. These results clearly demonstrate that the decomposition behavior of
Mg(OH)2 differed from that of hydromagnesite. Furthermore, to investigate the effect of CO2 captureon the sample surface, fresh Mg(OH)2 was exposed to a flow of 100 mL·min−1 CO2 for 24 h at r.t.
(referred to as CO2-exposed Mg(OH)2) in a TAC1000SE thermobalance (NETZSCH JAPAN, Inc.,
Japan).

Figure S2-4 shows the TG and DTA profiles of the fresh Mg(OH)2 and CO2-exposed

Mg(OH)2. The TG curves and positions of the DTA peaks did not show significant changes. The
behavior and endothermic peak due to decarbonation (see Figure S2-3) at 460 °C were not observed.
These results show that the bulk component of the sample was Mg(OH)2, and that surface carbonation
did not influence dehydration at the target temperature (260–400 °C) in our study. Thus, the author

Figure S2-1. FT-IR spectra of L10+CaF2 (a) after the sample was heated from room temperature (r.t.)
to 350 °C with heating rate of 10 °C min−1 under Ar flow of 100 mL min−1 and then cooled to r.t. The
sample was then stored in an IR cell at r.t. for (b) 10 min, (c) 30 min, and (d) 1 h. All the spectra
were recorded under ambient conditions (in air at room temperature) after heating, as described in themain manuscript.
153

has concluded that the carbonate species, which existed only on the surface, did not affect thedehydration of Mg(OH)2.
Figure S2-5 shows SEM images of Mg(OH)2-W, L10, LO20, and L10/LO10 using a scanningelectron microscope (JEOL, JSM-6510A). These figures show that the morphology of the samples
was almost the same, and Mg(OH)2 and the Li compounds could not be distinguished by this technique.
Therefore, the author needs to conduct further studies on surface and bulk information by othertechniques in future work.
Figure S2-6 shows the FT-IR spectra of LiOH-added Mg(OH)2 with a Mg(OH)2:LiOH ratio of
100:20 (LO20) prepared using the impregnation method and physical mixing, and that of LiOH·H2O.
When the LiOH·H2O phase was deposited on the Mg(OH)2 surface, the absorption peaks of
LiOH·H2O and Mg(OH)2 should be detected separately. The characteristic peak of LiOH∙H2O at3564 cm−1 was observed in the spectrum of the LO20 sample prepared by physical mixing, on theother hand, was undetectable in the spectrum of the LO20 sample prepared using the impregnation
method. These results indicate that a solid solution of Mg(OH)2 and LiOH or LiOH∙H2O could besynthesized using the impregnation method. In addition, these results explain the absence of LiOH
or LiOH∙H2O, and the LiCl or LiCl∙H2O peaks in the X-ray diffraction (XRD) patterns of L10, LO20,
and L10/LO10 (Figure 4-3 (a)), which supported our hypothesis that Li+ ions substituted Mg2+ ions.

Figure S2-2. FT-IR spectra of Mg(OH)2-W+CaF2 (a) after the sample was heated from room
temperature (r.t.) to 350 °C with heating rate of 10 °C min−1 under Ar flow of100 mL min−1 and thencooled to r.t. The sample was then stored in an IR cell at r.t. for (b) 10 min, (c) 30 min, and (d) 1 h.
All the spectra were recorded under ambient conditions (in air at room temperature) after heating, as
described in the main manuscript.
154

The expanded XRD patterns of the prepared samples are shown in Figures S2-7 (a) and (b). The
peak shift observed in these figures indicates the Li+ ion substitution to the Mg2+ ion site because there
might be no peak shifts if the substitution has not been performed. Moreover, previous studies[1,4]have already reported the Li+ ion substitution based on the slight lattice parameter changes and XRD
peak shifts. However, even such small differences in the peak positions might have influenced thereactivity of the Mg(OH)2/MgO system because the effective ionic radii of Li+ (0.76 Å) and Mg2+
(0.72 Å) are similar and the number of substituted ions cannot be significant. Therefore, the author
assumed that a solid solution between LiCl and Mg(OH)2 was produced.
Figures S2-8, S2-9, S2-10, S2-11, and S2-12 present the XRD patterns of Mg(OH)2-W+CaF2,
L10+CaF2, LO20+CaF2, MgO-W+CaF2, and L10/LO10+CaF2, respectively, after heating at different
temperatures under an Ar flow of 100 mL/min (Mg(OH)2-W and MgO-W denote Mg(OH)2 and MgO
pretreated using the impregnation method without Li compounds, respectively, and L10 denotes LiCladded Mg(OH)2 with Mg(OH)2:LiCl molar ratio of 100:10, LO20 denotes LiOH-added Mg(OH)2 with
Mg(OH)2:LiOH molar ratio of 100: 20, and L10/LO10 denotes LiCl and LiOH–co-added Mg(OH)2
with a Mg(OH)2:LiCl:LiOH molar ratio of 100:10:10). The temperatures at which the Mg(OH)2phase completely disappeared for all samples except for MgO-W were in good agreement with thosereported in Chapter 4 (Table 4-7),[4] which indicated that CaF2 did not inhibit the dehydration ofMg(OH)2. These results demonstrate that the FT-IR measurements and methodology used in thisstudy can be used to monitor the dehydration of Mg(OH)2.

Figure S2-3. TG profiles of (a) Mg(OH)2-W (treated by an impregnation method without Li
compounds), (b) hydromagnesite (light form), and (c) hydromagnesite (heavy form); DTA profiles of(d) Mg(OH)2-W, (e) hydromagnesite (light form), and (f) hydromagnesite (heavy form).
155

Figure S2-13 shows the FT-IR spectra of MgO-W+CaF2 before dehydration.

As MgO was

prepared using the impregnation method, a fraction of Mg(OH)2 was produced. MgO-W+CaF2 washeated at different temperatures using a thermobalance to identify the shoulder and split peaks at 3767
and ~3700 cm−1, respectively (Figures 4-5 (c)′ and (d)′).
Figure S2-14 depicts the second-derivative FT-IR spectra of Mg(OH)2-W+CaF2 after heating at
310 °C, obtained using the data in Figure 4-5 (a).
Figure S2-15 shows the relationship between the pH of the solution and the peak temperature forthe samples. The author determined the pH of LiCl, LiOH, LiBr, Li2CO3, LiCl, and LiOH solutions
using a conventional pH meter D-50 series (HORIBA Ltd., Japan).

The concentration of thesolutions was fixed at 0.04 mol/L because the concentration was close to that of the sample preparedby the impregnation method. The peak temperature was defined to be a temperature at which thedehydration rate reached to local maximum, as shown in shown in Chapter 3.[4] The lower peaktemperature indicates that dehydration can progress at lower temperatures. This figure does not show
a clear relationship between them. Therefore, the pH of the solution did not affect the dehydration
reactivity.

Figure S2-4. TG profiles of (a) fresh Mg(OH)2, and (b) CO2-exposed Mg(OH)2; DTA profiles of (c)fresh Mg(OH)2, and (d) CO2-exposed Mg(OH)2.

156

Figure S2-5. SEM images of (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d) L10/LO10.

Theexperiments were carried out by a scanning electron microscope (JEOL, JSM-6510A) belonging to
the Center for Analytical Instrumentation, Chiba University, Japan. The acceleration voltage was 15
kV.

Figure S2-6. FT-IR spectra of LO20 (LiOH-added Mg(OH)2 with a Mg(OH)2:LiOH molar ratio of
100:20) samples prepared using the impregnation method and physical mixing, and that of LiOH·H2O.
Both LO20 samples were dried at 120 °C overnight before measurements.
157

Figure S2-7. Expanded XRD patterns of the prepared samples. (a) and (b) show the peak positionsof the (001) and (002) planes, respectively. The grey dotted line is the position of Mg(OH)2-W.

Figure S2-8. XRD patterns of Mg(OH)2-W+CaF2 after dehydration under an Ar flow of 100 mL min−1.
The gray dotted lines with circles at the top and red dotted lines with triangles at the top indicate thepeak positions of Mg(OH)2 after heating to 310 °C and MgO after heating to 400 °C, respectively.

158

Figure S2-9. XRD patterns of L10+CaF2 after dehydration under Ar flow of 100 mL min−1. The graydotted lines with circles at the top and red dotted lines with triangles at the top indicate the peak
positions of Mg(OH)2 after heating to 260 °C and MgO after heating to 400 °C, respectively.

Figure S2-10. XRD patterns of LO20+CaF2 after dehydration under an Ar flow of 100 mL min−1.
The gray dotted lines with circles at the top and red dotted lines with triangles at the top indicate thepeak positions of Mg(OH)2 after heating to 300 °C and MgO after heating to 400 °C, respectively.

159

Figure S2-11. XRD patterns of MgO-W+CaF2 after dehydration under an Ar flow of 100 mL min−1.
The red dotted lines with triangles at the top indicate the peak position of MgO after heating to 400 °C.

Figure S2-12. XRD patterns of L10/LO10+CaF2 after dehydration under an Ar flow of 100 mL min−1.
The gray dotted lines with circles at the top and red dotted lines with triangles at the top indicate thepeak positions of Mg(OH)2 after heating to 260 °C and MgO after heating to 400 °C, respectively.

160

Figure S2-13. FT-IR spectrum of MgO-W+CaF2.

Figure S2-14. Second derivative FT-IR spectrum of Mg(OH)2-W+CaF2 after heating at 310 °C (basedon the data in Figure 5 (a)).

161

Figure S2-15. Plots of pH of Li-salt solutions vs peak temperature (LB10 denotes LiBr-added
Mg(OH)2 with a Mg(OH)2:LiBr molar ratio of 100:10 and LC10 denotes Li2CO3-added Mg(OH)2 with
Mg(OH)2:Li2CO3 molar ratio of 100:10).

References[1] Sivasankari, J.; Selvakumar, S.; Sivaji, K.; Sankar; S. Structural and optical characterization of
MgO:X (X = Li, Na, and K) by solution combustion technique. J. Alloys Compd. 2014, 616, 51−57.
[2] Sutradhar, N.; Sinhamahapatra, A.; Roy, B.: Bajaj, H. C.; Mukhopadhyay, I.; Panda, A. B.
Preparation of MgO nano-rods with strong catalytic activity via hydrated basic magnesium carbonates.
Mater. Res. Bull. 2011, 46, 2163–2167.
[3] Vagvolgyi, V.; Frost, R. L.; Hales, M.; Locke, A.; Kristof, J.; Horvath, E. Controlled rate thermal
analysis of hydromagnesite. J. Therm. Anal. Calorim. 2008, 92, 893–897.
[4] Kurosawa, R.; Takeuchi, M.; Ryu, J. Comparison of the effect of co-addition of Li compounds andaddition of a single Li compound on reactivity and structure of magnesium hydroxide. ACS Omega
2019, 4, 17752−17761.

S3. Supporting information for Chapter 5
S3.1. Experimental procedure
S3.1.1. Reactivity evaluation for hydration reaction by using thermobalance
The reactivities of all the samples were determined using a thermobalance (TGD–9600 series,
ADVANCE RIKO, Inc.). The samples (20 mg) were charged into a Pt cell. The mole fraction of162

Mg(OH)2 was calculated based on the data obtained using the thermogravimetric (TG) analysis
technique. A thermobalance was used during the process.[1–6] The mole fraction was determined as
explaned in the main chapters.
Figure S3-1 shows the dehydration and hydration reaction profiles of Mg(OH)2-W. x0 [-] denotes
the mole fraction of Mg(OH)2 and was determined immediately before the sample was hydrated, xh [-]denotes the mole fraction of the sample and the value was calculated immediately after hydration, andxc [-] denotes the mole fraction of the sample determined 10 min after the completion of the hydration
process. ∆xd, ∆x1, and ∆x2 [%] represent the extent of dehydration reaction conversion, hydration
reaction conversion, and conversion change (physically adsorbed water), respectively (expressed aspercentages) (Figure S3-1). These values are calculated as follows:

𝛥𝛥𝛥𝛥𝑑𝑑 = (1 − 𝑥𝑥0 ) × 100(1)

𝛥𝛥𝑥𝑥1 = (𝑥𝑥𝑐𝑐 − 𝑥𝑥0 ) × 100(2)

𝛥𝛥𝛥𝛥2 = (𝑥𝑥ℎ − 𝑥𝑥𝑐𝑐 ) × 100(3)

Figure S3-1. Dehydration and hydration reaction profiles of Mg(OH)2-W (Td = 350 ℃, Th = 110 ℃,
PH2O = 57.8 kPa).

163

The temperature dependence of the hydration process (MgO hydration) was determined. Th wasvaried (110, 170, and 200 ℃).

Dehydration temperature (Td) was set at 350 ℃ and the water vapor

pressure was 57.8 or 31.2 kPa.
The non-zero mole fraction post dehydration can be potentially attributed to the presence ofstructural water.[7–9] The mole fraction (hydrated MgO) defined in this study denotes the net hydration
conversion. The hydration behavior was studied (Figure S3-2 and Eq. (4)). The mole fraction wasdetermined as follows:(4)𝑥𝑥ℎ𝑦𝑦𝑦𝑦 = 𝑥𝑥 − 𝛥𝛥𝛥𝛥𝑠𝑠

where xhyd [-] denotes the net change in the mole fraction due to hydration and ∆xs denotes the mole
fraction of structural water.

Figure S3-2. Hydration reaction profile of de-Mg(OH)2-W studied at a temperature of 110 ℃(Mg(OH)2-W was dehydrated at 350 ℃, following which it was hydrated at 110 ℃ (PH2O: 57.8 kPa)).

S3.2. Results
Figure S3-3 shows the Fourier transform-infrared (FT-IR) spectral profiles of L10. The spectrumof the as-prepared sample was recorded at room temperature (r. t.) in air. Following this, the sample
was stored in the IR cell at r. t. in air.

The storage time was varied as 10 min, 30 min, and 1 h,following which the spectra were recorded. L10 was the sample of choice to study the effect of water
as LiCl exhibited the property of deliquescence (in air). The intensities of the peaks at 3700, 3395,
and 1634 cm−1, which were assigned to the hydroxyl groups present on the surface of Mg(OH)2, νOH,
164

Figure S3-3. FT-IR spectral profile recorded for L10 (a) as-synthesized. The sample was stored in
an IR cell at r. t. for (b) 10 min, (c) 30 min, and (d) 1 h. All the spectra were recorded under ambientconditions (in the air at r. t.), as described in the manuscript.

Figure S3-4. TG curves recorded for fresh de-Mg(OH)2 and CO2-exposed de-Mg(OH)2.

Fresh

Mg(OH)2 and CO2-exposed Mg(OH)2 samples were dehydrated at a temperature of 350 ℃, following
which they were hydrated at 110 ℃ (PH2O: 57.8 kPa).

165

Figure S3-5. Hydration behavior of (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and (d) deL10/LO10 studied at 110 ℃.

Mg(OH)2-W, L10, LO20, and L10/LO10 were dehydrated at 350 ℃,

following which they were hydrated at 110 ℃ (PH2O:57.8 kPa).
and δ modes of H2O molecules adsorbed on the Mg(OH)2 surface,[10] respectively, did not change
significantly when they were exposed to air at r. t. for a period of 10–60 min (Figure S3-3). Thus,
the water vapor present in the air barely influenced the FT-IR spectral profiles. The absorption bandat 2359 cm−1 was assigned to CO2 in the gas phase (Figure S3-3).[10] The absorption bands at 1523
and 1446 cm−1 could be ascribed to the CO32− and/or HCO3− species produced when CO2 was adsorbedon the surface of MgO.[10,11] It was observed that the intensities of these bands did not change
significantly.

Diffraction patterns corresponding to carbonate species were not observed in the X-

ray diffraction (XRD) patterns of the samples. It was observed that a small amount of carbonate wasadsorbed on the sample surface.

Fresh Mg(OH)2 was exposed to a flow of CO2 (flow rate: 100

mL·min−1; 24 h) at r. t. (referred to as CO2-exposed Mg(OH)2) to investigate the effect of CO2 captureon the sample properties. A TAC1000SE thermobalance (Netzsch Japan, Inc., Japan) was used to
conduct the experiments. Figure S3-4 shows the hydration properties of fresh de-Mg(OH)2-W and
CO2-exposed de-Mg(OH)2.

Both samples were dehydrated at 350 ℃ (time: 30 min) and thenhydrated at 110 ℃ under an atmosphere of water vapor and Ar gas (mixture; 57.8 kPa; 80 min). The
166

hydration properties of the samples were not significantly different from each other. These results
revealed that the major constituent of the sample was MgO.

It was also observed that surface

carbonation did not influence the hydration behavior of MgO.
Figures S3-5–S3-7 show the hydration behaviors of de-Mg(OH)2, de-L10, de-LO20, and deL10/LO10 at 110, 170, and 200 ℃ under (PH2O: 57.8 kPa). Analyses of the data presented in the
figures reveal that the addition of Li compounds promotes the hydration of MgO. The co-additionof LiCl and LiOH enhanced the reactivity at 200 ℃. The data presented in Figures S3-5–S3-7 wereused to prepare Figure S3-8. The conversions of ∆x1 and ∆x2 were calculated using Eqs. (4) and (5),respectively.
Figure S3-9 shows the FT-IR spectral profile of de-MgO-W (before the hydration reaction at 110,
170, and 200 ℃). Two absorption bands at 3710 and 3691 cm−1 and a shoulder band at 3756 cm−1
were observed in this figure. These characteristic peaks were observed in the spectral profile of deLO20 (before hydration; black lines in Figures 5-7–5-9 (c)). The results indicated that the surface
state of de-LO20 was similar to that of MgO.

Figure S3-6. Hydration behavior of (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and (d) deL10/LO10 studied at 170 ℃.

Mg(OH)2-W, L10, LO20, and L10/LO10 were dehydrated at 350 ℃,

following which they were hydrated at 170 ℃ (PH2O: 57.8 kPa).
167

Figure S3-10 shows the FT-IR spectra recorded for Mg(OH)2-W.[12–17] This peak assignment was
the same as that in Chapter 4.[12] This figure was revised (Figure 5-10; main manuscript).
Figure S3-11 shows the hydration behavior of the de-Mg(OH)2-W, de-L10, de-LO20, and deL10/LO10 at 110 °C (PH2O: 31.2 kPa).

Initially, the mole fraction of Mg(OH)2 (present in de-L10

and de-L10/LO10) increased rapidly. The increase could be attributed to the adsorption of water by
LiCl. Under these conditions, the hydration of LiCl did not positively affect the hydration of MgO
because the mole fraction of de-L10 (determined after an 80 min-long hydration process) was lower
than the mole fraction of de-Mg(OH)2-W.
Figure S3-12 shows the hydration behaviors of LiOH/MgO and de-Mg(OH)2-W at 200 ℃. This
figure shows that the hydration reactivity of LiOH/MgO increases with an increase in the amount of
LiOH. The data on hydration conversion (determined by analyzing the data presented here) has beenpresented in Table S3-1. A limited number of Mg2+ ion sites could be substituted by Li+ ions (FigureS3-13) because the peak at 3564 cm−1 (corresponding to LiOH·H2O) could not be detected in the
spectra recorded with the impregnated LO20 sample.

The results revealed that the hydration

Figure S3-7. Hydration behavior of (a) de-Mg(OH)2-W, (b) de-L10, (c) de-LO20, and (d) deL10/LO10 studied at a temperature of 200 ℃.

Mg(OH)2-W, L10, LO20, and L10/LO10 were

dehydrated at 350 ℃, following which they were hydrated at 200 ℃ (PH2O: 57.8 kPa).
168

reactivity of MgO was influenced by the concentration of OH− ions.
Figure S3-14 shows the FT-IR spectra recorded for hydrated de-MgO-W.

The sample washydrated at a temperature of 110 ℃ (PH2O: 31.2 kPa).

Figure S3-8. Reaction conversion vs. hydration time plots to determine the hydration reaction (∆x1)
and physical adsorption (∆x2) values for dehydrated samples. Data were recorded at (a) 110 ℃, (b)170 ℃, and (c) 200 ℃.

Filled symbols and open symbols represent ∆x1 and ∆x2, respectively.

Mg(OH)2-W, L10, LO20, and L10/LO10 were dehydrated at a temperature of 350 ℃, following which
they were hydrated at 110, 170, or 200 ℃ (PH2O: 57.8 kPa).

169

Figure S3-9. FT-IR spectral profiles of de-MgO-W. MgO-W was dehydrated at a temperature of
350 °C, following which it was dehydrated again at temperatures of 110, 170, and 200 ℃ for 10 min.
(FT-IR spectral profiles were recorded at r. t. before the samples were hydrated).

Figure S3-10. FT-IR spectral profiles of samples.

Literature reports were referred to for peak

[12–17] The peaks were reassigned, as shown in Figure 4-10.

assignment.

in the main manuscript.

170

The data has been presented

Figure S3-11. Hydration behavior of (a) de-L10/LO10, (b) de-L10, (c) de-LO20, and (d) de-Mg(OH)2W (110 ℃; 31.2 kPa).

Mg(OH)2-W, L10, LO20, and L10/LO10 were dehydrated at a temperatureof 350 ℃, following which they were hydrated at 110 ℃ (PH2O: 31.2 kPa).

Figure S3-12. Hydration behavior of pure and LiOH-added MgO studied at 200 ℃ for (a) de-LO10,(b) de-LO20, (c) de-LO25, (d) de-LO30, (e) de-LO50, and (f) de-Mg(OH)2-W. LO10, LO20, LO25,
LO30, LO50, and Mg(OH)2-W were dehydrated at a temperature of 350 ℃, following which they
were hydrated at a temperature of 200 ℃ (PH2O: 57.8 kPa). LO10, LO20, LO25, LO30, and LO50
represent LiOH-added Mg(OH)2 samples where the Mg(OH)2:LiOH molar ratios were 100:10, 100:20,
100:25, 100:30, and 100:50, respectively.
171

Figure S3-13. FT-IR spectral profiles recorded for samples (as-synthesized LO20, LO25, LO30, andLO50) prepared following the (a) impregnation method and (b) physical mixing technique. LO20,
LO25, LO30, and LO50 represent LiOH-added Mg(OH)2 samples where the Mg(OH)2:LiOH molar
ratios were 100:20, 100:25, 100:30, and 100:50, respectively.

Figure S3-14. FT-IR spectral profiles recorded for de-MgO-W hydrated at 110 ℃ (PH2O: 31.2 kPa).
MgO-W was dehydrated at a temperature of 350 ℃, following which it was hydrated at 110 ℃ (PH2O:
31.2 kPa).

172

Table S3-1. Hydration conversion values recorded at a temperature of 200 ℃ for de-Mg(OH)2-W andLiOH/MgO. LO10, LO20, LO25, LO30, and LO50 represent LiOH-added Mg(OH)2 samples where
the Mg(OH)2:LiOH molar ratio were 100:10, 100:20, 100:25, 100:30, and 100:50, respectively.
SampleHydration conversion (∆x1) [%]de-Mg(OH)2-W

2.9

de-LO1010.8

de-LO20

31.7

de-LO25

38.1

de-LO30

45.3

de-LO50

58.3

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[9] Iwasaki, S.; Kodani, S.; Koga, N. Physico-geometrical kinetic modeling of the thermal
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[15] Knözinger, E.; Jacob, K. H.; Singh, S.; Hofmann, P. Hydroxyl groups as IR active surface probes
on MgO crystallites. Surf. Sci. 1993, 290, 388–402.
[16] Chizallet, C.; Costentin, G.; Lauron-Pernot, H.; Krafft, J. M.; Bazin, P.; Saussey, J.; Delbecq, F.;
Sautet, P.; Che, M. Role of hydroxyl groups in the basic reactivity of MgO: a theoretical andexperimental study. Oil Gas Sci. Technol.-Rev. IFP 2006, 61, 479–488.
[17] Chizallet, C.; Costentin, G.; Che, M.; Delbecq, F.; Sautet, P. Infrared characterization of hydroxyl
groups on MgO: A periodic and cluster density functional theory study. J. Am. Chem. Soc. 2007, 129,
6442–6452.

174

Appendix
These data were not discussed in the main sections, however, the author hopes these can be helps
and developments for the research in this laboratory.

A1. The Dehydration and Hydration Reactivity of the Samples
A1-1. Dehydration reactivity of the samples
This section refers to the dehydration reactivity for the samples listed in Table A1-1.

The

samples listed in the table were prepared by the impregnation method as shown in the main sections.
The dehydration reaction test was carried out by using a thermobalance (TGD–9600 series,
ADVANCE RIKO, Inc.). The condition for dehydration reaction test is heating temperature fromroom temperature to 600 °C under Ar flow of 100 mL min−1.
Table A1-1. Mixing ratio of the prepared samples
Samples

Mixing ratio（mole ratio）

L5, 10, 20, 30

Mg(OH)2 : LiCl = 100 : 5, 10, 20, 30

LO5, 10, 20, 25, 30, 33, 50

Mg(OH)2 : LiOH = 100 : 5, 10, 20, 25, 30, 33,
50

LB5, 10, 20, 30

Mg(OH)2 : LiBr = 100:5, 10, 20, 30

N5, 10, 20, 30

Mg(OH)2 :NaCl = 100 : 5, 10, 20, 30

NO10, 20, 30

Mg(OH)2 : NaOH = 100 : 10, 20, 30

K5, 10, 20, 30

Mg(OH)2 :KCl = 100 : 5, 10, 20, 30

KC10, 20, 30

Mg(OH)2 :K2CO3 = 100 : 10, 20, 30

L5/LO5, L10/LO10, L20/LO20

Mg(OH)2 : LiCl : LiOH = 100 : 5: 5, 10: 10,
20: 20Mg(OH)2–W

Treated via impregnation Mg(OH)2 without any
compounds

Figures A1-1–1-8 show the dehydration behavior of Mg(OH)2 with singly added LiCl, LiOH,
LiBr, NaCl, NaOH, KCl, K2CO3, or co-added LiCl and LiOH. LiCl and LiBr obviously have large
effect on the dehydration of Mg(OH)2. The dehydration peak temperatures (the explanation wasshown in Chapter 3) for LiCl/Mg(OH)2 and LiBr/Mg(OH)2 were dramatically lowered comparedwith that for Mg(OH) 2 -W (Table A1-2).

The peak temperatures for LiCl/Mg(OH) 2 andLiBr/Mg(OH)2 were almost the same between the samples, thus, the hydration reaction test wasneeded to determine the most optimal mixing ratio. The peak temperatures for NaCl/Mg(OH)2 wereslightly lowered and they were dehydrated at a little lower temperatures compared with that for
175

Figure A1-1. Dehydration behavior of LiCl/Mg(OH)2 samples.

Figure A1-2. Dehydration behavior of LiOH/Mg(OH)2 samples.
Mg(OH)2-W (Table A1-2). NaOH/Mg(OH)2 was also dehydrated in low temperature side comparedwith Mg(OH)2-W.

The dehydration behavior of NO10 was unique (two steps dehydration).

Ca(OH)2 with added LiOH also showed similar two step dehydration reported by Yan et al.[A1] On
the other hand, the behavior of NO20 and 30 were one step dehydration. This reason is required to
be investigated in the future work.

The addition of K2CO3 slightly lowered the dehydration

temperature range. The peak temperature and the dehydration behavior of KCl/Mg(OH)2 were176

Figure A1-3. Dehydration behavior of LiBr/Mg(OH)2 samples.

Figure A1-4. Dehydration behavior of NaCl/Mg(OH)2 samples.
hardly changed to those of Mg(OH)2-W (Table A1-2). Therefore, it can be concluded that the effectof the addition of KCl on dehydration of Mg(OH)2 is small. The co-addition of LiCl and LiOH also
dramatically lowered the dehydration reaction temperature range of Mg(OH)2. The peak temperatureof Mg(OH)2 with added LiCl and LiOH was lowered than that of Mg(OH)2 with singly added LiCl or
LiOH (Chapter 3). Therefore, the investigation of the combination of added compounds can berequired to enhance the dehydration reactivity of Mg(OH)2.
177

Figure A1-5. Dehydration behavior of NaOH/Mg(OH)2 samples.

Figure A1-6. Dehydration behavior of KCl/Mg(OH)2 samples.

178

Figure A1-7. Dehydration behavior of K2CO3/Mg(OH)2 samples.

Figure A1-8. Dehydration behavior of LiCl/LiOH/Mg(OH)2 samples.

179

Table A1-2. Dehydration peak temperature for the samples.
Samples

Dehydration peak temperature [°C]

L5

325

L10

319

L20

316

L30

320

LO5

371

LO10

357

LO20

339

LO25

337

LO30

328

LO33

334

LO50

339

LB5

317

LB10

305

LB20

299

LB30

300

N5

358

N10

358

N20

361

N30

359

K5

365

K10

361

K20

364

K30

363

KC10

346

KC20

357

KC30

361

NO10

337

L5/LO5314

L10/LO10305

L20/LO20

297

Mg(OH)2–W

376

180

A1-2. Hydration behavior of the samples
This section refers to the hydration reactivity for the samples listed in Table A1-1. The condition
for dehydration/hydration reaction was as follows: First, the sample was heated at 350 °C (with heating
rate of 20 °C min−1) for 30 min under Ar flow of 100 mL min−1, and then the sample was hydrated at
the different hydration temperature (Th = 110, 170, or 200 °C with heating rate of −20 °C min−1) for
80 min under mixture of water vapor of 57.8 kPa and Ar flow. Finally, the sample was dried at thehydration temperature for 30 min. The explanation for the procedure, reaction conversion, and heat
output density were shown in Chapter 3.
Figures A1-9 (a)–(c) show the hydration behaviors of LiCl/MgO at 110, 170, and 200 °C,
respectively.

Figure A1-9 (a) shows the hydration of dehydrated L10, L20, and L30 were progressedmore quickly than that for dehydrated Mg(OH)2-W. However, that of dehydrated L5 was slower thanthat for dehydrated Mg(OH)2-W until 3000 s. This result indicates the mixing ratio of 5 mol% toMg(OH)2 is not enough for enhancing the hydration reactivity of MgO.

The values of ∆x2 for

LiCl/MgO samples were quite higher than that for dehydrated Mg(OH)2-W at 110 °C. This can bebecause LiCl easily forms LiCl hydrate species due to its high hygroscopicity.[A2] Figures A1-9 (b)and (c) show the behaviors at 170, and 200 °C. The hydration reactivity was improved for all thecases at 170 °C. On the other hand, the reactivity was low for dehydrated L20 and L30 comparedwith Mg(OH)2-W at 200 °C (Figure A1-9 (c)). This result indicates that too much LiCl to Mg(OH)2
declines the reactivity because large amount of additive might cover the sample surface, resulting inan obstacle for water diffusion. With respect to the heat output density by hydration reaction, thoseat 110 °C for dehydrated L10, L20, and L30 were around 1200 kJ kg−1 (Table S1-3). The values of
the heat output density and ∆x1 at 170 and 200 °C were dramatically declined. Considering theseresults and the amount of Li (small amount of Li is desirable due to its rarity), the best mixing ratio isconcluded 10 mol%. This result agrees with the previous studies.[A2,A3]
Figures A1-10 (a)–(c) show the hydration behavior of LiOH/MgO at 110, 170, and 200 °C,
respectively.

Figure A1-10 (a) shows that the addition of LiOH enhanced the hydration reactivity of

MgO. The value of ∆x1 for LiOH/MgO at 110 °C was higher than that of dehydrated Mg(OH)2-W.
It was worth noticing that the hydration of MgO proceeded well by the addition of LiOH at 170 and
200 °C. Chapter 5 indicated that the addition of a hydroxide (LiOH, NaOH, or KOH) is one of theeffective way to enhance the reactivity.
Figures A1-11 (a)–(c) show the hydration behavior of NaCl/MgO at 110, 170, and 200 °C,
respectively.

The effect of amount of added NaCl can be ignored at 110 °C.

In particular, thehydration reactivity was higher for dehydrated Mg(OH)2-W than for NaCl/MgO at 170 °C. On theother hand, dehydrated N30 showed slightly high hydration reactivity at 200 °C compared withdehydrated Mg(OH)2-W.

181

Table A1-3. Reaction conversion of ∆x1, ∆x2 [%], and heat output density Qr [kJ kg−1] of the samples
Th = 110 °C

Th = 170 °C

Th = 200 °C

Samples

∆x1

∆x2

Qr

∆x1

∆x2

Qr

∆x1

∆x2

Qr

L5

68.6

15.1

919

37.3

5.7

500

5.6

4.9

75

L10

90.3

32.7

1170

59.3

7.4

768

5.5

5.3

71

L20

99.3

79.6

1204

65.3

8.5

791

2.0

5.0

25

L30

107.0

114.7

1220

53.2

7.7

607

3.4

4.9

38

LO10

84.6

4.2

1128

57.5

2.2

767

10.8

3.3

145

LO20

83.8

4.2

1075

73.7

2.3

954

31.7

3.3

407

LO25

83.9

4.3

1057

81.0

2.1

1020

38.1

3.2

480

LO30

84.4

4.6

1043

81.6

1.9

1009

45.3

2.2

560

LO50

89.3

4.0

1029

89.2

1.9

1028

58.3

2.0

672

N5

77.3

4.8

1022

4.4

3.7

58

3.3

2.7

44

N10

70.3

6.4

887

4.2

2.1

53

1.7

3.1

22

N20

69.7

5.3

807

3.5

3.8

40

2.1

3.2

24

N30

66.5

4.9

710

10.4

4.0

120

6.9

2.8

79

K5

71.5

5.3

933

8.9

3.1

116

1.5

2.4

20

K10

72.2

4.8

888

5.5

4.0

68

1.2

2.7

15

K20

70.9

5.5

784

3.9

3.7

44

2.4

3.8

28

K30

70.3

4.3

706

3.9

4.4

39

2.5

2.7

26

NO10

94.5

10.5

1149

89.5

2.0

1089

59.7

1.6

763

NO20

103.0

33.0

1258

100.0

2.6

1227

99.8

1.6

1219

L10/LO10

95.4

37.7

1190

93.5

4.2

1166

92.0

2.7

1146

Mg(OH)2–W

68.2

4.0

947

28.4

2.5

394

2.9

3.3

40

Figures A1-12 (a)–(c) show the hydration behavior of KCl/MgO at 110, 170, and 200 °C,
respectively. The behaviors were almost the same at 110 and 200 °C, hence, the addition of KCl did
not have effect on the hydration.

In contrast, the hydration did not progress well for KCl-added

samples at 170 °C, it indicates that KCl covered the sample surface, which hindered the hydration of
MgO. The values of heat output density at 110, and 170 °C decreased with increasing the amount ofadded KCl.

Therefore, the author concluded that the addition of KCl was unnecessary for thehydration.

182

Figure A1-9. Hydration behavior of LiCl/MgOat (a) 110, (b) 170, and (c) 200 °C (all samples were
dehydrated at 350 °C, and then, hydrated at 110, 170, and 200 °C under PH2O of 57.8 kPa).

Figure S1-10. Hydration behavior of LiOH/MgO at (a) 110, (b) 170, and (c) 200 °C (all samples
were dehydrated at 350 °C, and then, hydrated at 110, 170, and 200 °C under PH2O of 57.8 kPa).
183

Figure S1-11. Hydration behavior of NaCl/MgO at (a) 110, (b) 170, and (c) 200 °C (all samples
were dehydrated at 350 °C, and then, hydrated at 110, 170, and 200 °C under PH2O of 57.8 kPa).

Figure S1-12. Hydration behavior of KCl/MgO at (a) 110, (b) 170, and (c) 200 °C (all samples
were dehydrated at 350 °C, and then, hydrated at 110, 170, and 200 °C under PH2O of 57.8 kPa).
184

A2. SEM images
Here, the SEM images for the samples before and after reaction tests are shown. These images
are taken by scanning electron microscope (JEOL, JSM-6510A) belonged to the center for analytical
instrumentation in Chiba university.

An acceleration voltage was 15 kV.

The samples were

Mg(OH)2-W, L10, LO20, and L10LO10 (the sample composition was written in the main section).

A2.1. SEM images for the prepared samples
Figures A2-1 (a)–(d) show the SEM images of prepared samples. The sample descriptions are
the same as shown in the main section. It is difficult to detect any differences between samples.
Therefore, the appearances of the samples were not different from the authentic Mg(OH)2 even after
adding the Li compounds (LiCl and/or LiOH).

A2.2. SEM images for the reacted samples
Figures A2-2 (a)–(d) show the SEM images of samples heated to 800 °C by thermobalance. The
appearances of Figures A2-2 (a)–(c) (Mg(OH)2-W, L10, and LO20) were not significantly different
from those after preparation (Figures A2-1 (a)–(c)). Therefore, the appearance of Mg(OH)2 wasmaintained and the component of a sample converted from Mg(OH)2 to MgO as reported previous
studies, which is called a pseudomorph.[A4] On the other hand, the surface of L10LO10 was
dramatically dissimilar to the original surface and other samples. That had really smooth surface.
This probably indicates that the Li compounds (LiCl and LiOH) melted during heating. Certainly
the melting temperature of components is decreased in the binary system. When the mixing ratio of
LiCl and LiOH is equal to 1: 1, the melting point of the binary system is 315 °C which is lower thanthat of single LiCl and LiOH (610 and 477 °C, respectively) as shown in Figure A2-3. Therefore,
the lower melting point could make the surface of L10LO10 smooth.
Figures A2-4 (a)–(d), A2-5 (a)–(d), and A2-6 (a)–(d) show the SEM image of the samples before
hydration reaction at 110, 170, and 200 °C. Before the hydration test, the shape and appearance ofthe samples were unchanged compared with after preparation (Figures A2-1 (a)–(d)). Salts added
to Mg(OH)2 were presumably not melt during dehydration at 350 °C.
Figures A2-7 (a)–(d), A2-8 (a)–(d), and A2-9 (a)–(d) show the SEM image of the samples after
hydration reaction at 110, 170, and 200 °C for 80 min. The appearance of the samples was not
dramatically different from the samples before the hydration reaction (Figures A2-4 (a)–(d), A2-5(a)–(d), and A2-6 (a)–(d)). Therefore, the component in a sample changed, on the other hand, theappearance was unchanged during the transition Mg(OH)2 to MgO and MgO to Mg(OH)2.

185

Figure A2-1. The SEM images for prepared samples: (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d)L10/LO10

Figure A2-2. The SEM images for the samples after heating to 800 °C: (a) Mg(OH)2-W, (b) L10, (c)LO20, and (d) L10/LO10186

Figure A2-3 Phase diagram of LiCl and LiOH. The y axis is melting point and the x axis is moleratio (Research center of computational mechanics Inc: http://www.rccm.co.jp/product/thermodynamics/factsage/).

Figure A2-4. The SEM images for the samples before hydration at 110°C (The samples were
dehydrated at 350 °C for 30 min, and then dehydrated at 110 °C for 10 min): (a) Mg(OH)2-W, (b) L10,(c) LO20, and (d) L10/LO10.

187

Figure A2-5. The SEM images for the samples before hydration at 170°C (The samples were
dehydrated at 350 °C for 30 min, and then dehydrated at 170 °C for 10 min): (a) Mg(OH)2-W, (b) L10,(c) LO20, and (d) L10/LO10.

Figure A2-6. The SEM images for the samples before hydration at 200°C (The samples were
dehydrated at 350 °C for 30 min, and then dehydrated at 200 °C for 10 min): (a) Mg(OH)2-W, (b) L10,(c) LO20, and (d) L10/LO10.

188

Figure A2-7. The SEM images for the samples after hydration at 110 °C (The samples were dehydrated
at 350 °C for 30 min, and then dehydrated at 110 °C for 10 min. After that, the samples were hydrated
at 110 °C for 80 min, and finally dried for 30 min): (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d)L10/LO10.

Figure A2-8. The SEM images for the samples after hydration at 170 °C (The samples were dehydrated
at 350 °C for 30 min. After that, the samples were hydrated at 170 °C for 80 min, and finally driedfor 30 min): (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d) L10/LO10.

189

Figure A2-9. The SEM images for the samples after hydration at 200 °C (The samples were dehydrated
at 350 °C for 30 min. After that, the samples were hydrated at 200 °C for 80 min, and finally driedfor 30 min): (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d) L10/LO10.

A3. DSC measurement
In this section, the results of DSC (differential scanning calorimetry) measurement is mentioned.
Mg(OH)2-W, L10, LO20, and L10/LO10 were used as samples (the sample composition was written
in the main section).
The measurement conditions are as follows: The DSC measurement was carried out under N2 flowof 50 mL min−1 by using Thermo Plus EVO2 (Rigaku Corp.). The sample of approximately 1 mg
was placed in an aluminum pan and heated to desired temperature.

Mg(OH)2-W was heated to

450 °C, and L10, LO20, and L10/LO10 were heated to 400 °C with heating rate of 10 °C min−1. αAl2O3 was used as a reference. The measurement was carried out triplicate to reduce the error andthe 90% confidence interval for the measurement was estimated by Student’s t-distribution (the detailis shown in Chapter 3). The dehydration enthalpy of the samples was calculated by the area of DSC
curve as shown in Figure A3-1. The enthalpy in this study was described as it per g magnesiumhydroxide as follows:(A 3-1)𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅𝑅 𝑒𝑒𝑒𝑒𝑒𝑒ℎ𝑎𝑎𝑎𝑎𝑎𝑎𝑎𝑎 = ∆𝐻𝐻𝑟𝑟 /𝑤𝑤𝑟𝑟

where ∆Hr [J g−1] is the reaction enthalpy per g sample, and wr [–] is the weight ratio of Mg(OH)2 in190

the sample.
Figures A3-2–5 show the DSC curve for Mg(OH)2-W, L10, LO20, and L10/LO10. The red, blue,and green line are the DSC curve on the first, second, and third try, respectively. The curve except
for LO20 were relatively sharp while that for LO20 was very broad and showed a “side step”.[A1] A
previous study showed a similar DSC curve with the side step of Ca(OH)2 with added LiOH by ball
milling.[A1]Figure A3-6 shows the reaction enthalpy by dehydration of the samples with an error bar. This
figure shows the reaction enthalpy for Mg(OH)2-W, LO20, and L10/LO10 were relatively similar.
On the other hand, that for L10 was around 1000 J g−1-Mg(OH)2 which is relatively larger than thatfor the other.

However, that is not a significant difference. The value for the reaction enthalpy andthe peak temperature for the dehydration are shown in Table A3-1. These results indicate that thebond energy in the sample did not significantly vary by Li+ ion substitution into Mg2+ ion sites (theeffect of the addition of Li compounds is discussed in detail in Chapters 3–5). It is because thenumber of the substitution was not drastic.

Figure A3-1. The calculation for the dehydration enthalpy.

191

Figure A3-2. DSC curve for Mg(OH)2-W.

Figure A3-3. DSC curve for L10.

192

Figure A3-4. DSC curve for LO20.

Figure A3-5. DSC curve for L10/LO10.

193

Figure A3-6. Reaction enthalpy by dehydration of 4 samples with an error bar.
Table A3-1. Reaction enthalpy by dehydration and peak temperature for dehydration of the samples
SampleReaction enthalpy [J g−1]

Peak temperature [°C]

Mg(OH)2-W−878 ± −162

379.6 ± 19.5

L10−1014 ± −123

325.6 ± 6.1

LO20−859 ± −53348.3 ± 3.2

L10/LO10−868 ± −127304.6 ± 17.4

A4. CO2 TPD measurement
To investigate the basicity of Mg(OH)2-W, L10, LO20, and L10/LO10, CO2 temperature
programmed desorption (CO2-TPD) was carried out by using a thermobalance (TG-DTA2000SE,NETZSCH Japan, Inc., Japan). The condition is as follows; a sample was heated at 400 °C for 30
min with heating rate of 20 °C min−1 under Ar flowing of 100 mL min−1. After that, the sample wascooled down to 100 °C with heating rate of −20°C min−1 and adsorbed by CO2 of 100 mL min−1 at thetemperature for 60 min. Finally, the desorption process was carried out at 200, 400, and 600 °C for
30 min at a constant rate of 20 °C min−1. Actually, there is a difference between the programmed
temperature and measured temperature. This difference is caused by the difference of heated portionand measured portion, and so on. The actual temperature and TG profiles are shown in Figure A41.

194

Figure A4-1. Measured temperature program and TG profile of Mg(OH)2-W for CO2 TPD
measurement.
The CO2 TPD profiles of the samples shown in Figure A4-2 are derived from the TG profiles after
the induction of CO2. The desorption curve means the CO2 amount adsorbed on a sample. The
desorption of CO2 in the temperature range of 100–200, and 400–600 °C were significant for eachsample. This result means that weak and strong basic sites, corresponding to the temperature rangeof 100–200, and 400–600 °C, were presented. The desorption amount of dehydrated Mg(OH)2-W
was larger than those of dehydrated Li compounds-added Mg(OH)2. Therefore, the basicity of pure
MgO could be higher than those of Li compounds-added MgO. However, the CO2 amount adsorbedon the samples was actually quite low. Thus, the author suggest that other analysis methods are moresuitable to investigate the basicity.

195

Figure A4-2. CO2 TPD profiles of (a) Mg(OH)2-W, (b) L10, (c) LO20, and (d) L10/LO10.

A5. Nitrogen gas adsorption measurement
In this section, the results of nitrogen gas adsorption measurement are mentioned. The specificsurface area of the samples was evaluated by BET (Braunauer, Emmett, and Teller) theory. The
measured samples were listed in Table A1-1. The measurement was carried out for the prepared andreacted samples by using BELSORP-mini Ⅱ (MicrotracBEL Corp.).
Experimental conditions for the nitrogen gas adsorption were as follows: The powder sample wascharged into a sample cell and dried at 120 °C for 1 h under N2 flow of 100 mL min−1 before themeasurement. The measurement was carried out at liquid nitrogen temperature (77K). The partialpressure range for the adsorption process was 0.020–0.995, that for the desorption process was 0.050–
0.970. In general, when the value of C parameter for BET analysis is in the range of 100–200, theanalysis result for specific surface area is reasonable.[A5]Mean particle size in nm can be obtained by BET-plot based on the equation below;

𝑙𝑙 =

6

𝜌𝜌𝑠𝑠 ×𝑎𝑎𝑠𝑠,𝐵𝐵𝐵𝐵𝐵𝐵

× 103(A5-1)
196

where l [nm] is the mean particle size, ρs is the sample density [g cm−3], and as,BET is the specificsurface area [m2 g−1] for the sample. Here, the sample density was calculated as follows;𝜌𝜌𝑠𝑠 = 𝜌𝜌𝑀𝑀𝑀𝑀(𝑂𝑂𝑂𝑂)2 × 𝑥𝑥𝑀𝑀𝑀𝑀(𝑂𝑂𝑂𝑂)2 + 𝜌𝜌𝑎𝑎 × 𝑥𝑥𝑎𝑎 + 𝜌𝜌𝑏𝑏 × 𝑥𝑥𝑏𝑏 + ⋯(A 5-2)where ρMg(OH)2 and ρa,b, are the density of Mg(OH)2 and that of additives (LiCl, LiOH, NaCl ,and soon).

x is the weight ratio of the substances in a sample and was calculated from mixing ratio.

Figures A5-1–6 and A5-6–12 show the adsorption/desorption isotherms and BET plots based on
the isotherms, respectively.

Table 5-1 shows the specific surface area and mean particle size of thesamples. The surface area of Mg(OH)2 with an additive tends to be smaller than that of Mg(OH)2W.

Meanwhile, the mean paticle size of Mg(OH)2 with an additive tends to be larger than that ofMg(OH)2-W, meaning that the additive could partially cover the surface of Mg(OH)2.
Figures A5-13 and 5-14 show the relationship between the dehydration peak temperature and thespecific surface area or mean particle size of each sample based on Tables A1-2 and A5-1. The peaktemperature of the samples with low specific surface area (or large mean particle size) tends to belower. On the other hand, a clear relationship could not be observed. In addition, the hydration
reactivity of the samples was not affected by the surface area, and the particle size. Thus, it can beconcluded that these factors do not affect on the dehydration and hydration reactivity of the samples.

Figure A5-1. Adsorption/desorption isotherm for LiCl/Mg(OH)2.

197

Figure A5-2. Adsorption/desorption isotherm for LiOH/Mg(OH)2.

Figure A5-3. Adsorption/desorption isotherm for LiBr/Mg(OH)2.

198

Figure A5-4. Adsorption/desorption isotherm for NaCl/Mg(OH)2.

Figure A5-5. Adsorption/desorption isotherm for NaOH/Mg(OH)2.

199

Figure A5-6. Adsorption/desorption isotherm for KCl/Mg(OH)2.

Figure A5-7. Adsorption/desorption isotherm for K2CO3/Mg(OH)2.

200

Figure A5-8. BET plots for LiCl/Mg(OH)2 based on the adsorption/desorption isotherm (Figure A51).

Figure A5-9. BET plots for LiOH/Mg(OH)2 based on the adsorption/desorption isotherm (Figure A52).

201

Figure A5-10. BET plots for LiOH/Mg(OH)2 based on the adsorption/desorption isotherm (Figure A53).

Figure A5-11. BET plots for NaCl/Mg(OH)2 based on the adsorption/desorption isotherm (Figure A54).

202

Figure A5-12. BET plots for NaOH/Mg(OH)2 based on the adsorption/desorption isotherm (FigureA5-5).

Figure A5-13. BET plots for KCl/Mg(OH)2 based on the adsorption/desorption isotherm (Figure A56)

203

Figure A5-14. BET plots for K2CO3/Mg(OH)2 based on the adsorption/desorption isotherm (Figure
A5-7)
Figures A5-13 (a) and (b) show the relationship between the dehydration peak temperature andspecific surface area or mean particle size of the samples obtained from BET plots (Figurs 5-7–5-12).
These figures present that the specific surface area (mean particle size) of NaOH-, LiBr-, or K2CO3added Mg(OH)2 tended to be small (large). These additives relatively form their hydrates, thus, thehygroscopicity of the additives can affect the sample morphology. The dehydraton peak temperatureof the samples with the surface area of ~30 m2 g−1 (the particle size of 80 nm) were relatively low.
On the other hand, a clear correlation could not be observed.

Figure A5-13 Relationship between the peak temperature and (a) specific surface area or (b) mean
particle size obtained from BET plots (Figurs 5-8–5-14). Purple dot is for Mg(OH)2-W.

204

Table A5-1. Specific surface area [m2 g−1] and mean particle size [nm] of the samples based on BETplots (Figures A5-7–12)
SamplesSpecific surface area

Mean particle size

L5

30.9 ± 2.04

82.84 ± 5.590

L10

29.8 ± 1.77

85.43 ± 5.412

L20

27.5 ± 1.77

93.91 ± 6.124

L30

25.5 ± 3.55

102.4 ± 13.61

LO5

34.5 ± 3.41

73.78 ± 7.505

LO10

33.4 ± 0.90

75.84 ± 2.046

LO20

31.6 ± 2.50

80.02 ± 6.283

LO25

32.1 ± 0.94

78.73 ± 2.320

LO30

32.1 ± 2.59

78.71 ± 6.492

LO50

26.2 ± 1.72

95.93 ± 6.360

LC10

36.8 ± 1.41

69.97 ± 2.677

LB10

18.1 ± 1.40

132.3 ± 10.49

LB20

9.30 ± 0.70

247.3 ± 18.29

LB30

5.30 ± 1.72

419.3 ± 86.69

N5

35.1 ± 1.67

73.09 ± 3.461

N10

30.5 ± 0.86

84.41 ± 2.377

N20

25.8 ± 1.33

100.6 ± 5.220

N30

27.2 ± 2.00

96.50 ± 6.988

NO10

10.6 ± 0.17

241.7 ± 3.93

NO20

2.55 ± 0.61

1024 ± 235.4

NO30

0.894 ± 0.853512 ± 2896

K5

33.8 ± 2.03

75.85 ± 4.517

K10

31.3 ± 1.67

82.32 ± 4.407

K20

29.0 ± 2.13

90.00 ± 6.450

K30

27.1 ± 1.84

97.39 ± 6.769

KC10

16.5 ± 1.19

155.2 ± 11.15

KC20

11.0 ± 3.60

239.5 ± 78.22

KC30

11.6 ± 1.19

223.0 ± 23.57

L5/LO525.8 ± 1.96

99.78 ± 7.409

L10/LO10

22.7 ± 0.66

114.3 ± 3.272

Mg(OH)2-W

34.2 ± 1.65

74.47 ± 3.650

Mg(OH)2 (Wako)

36.4 ± 2.71

69.96 ± 5.351

205

A6. Dehydration reaction kinetics
Chapter 2 discussed the dehydration reaction kinetics based on the integration method. The
dehydration profile of LiOH/Mg(OH)2 with large LiOH amount (LO20, LO25, LO33, and LO50)
could be reproduced by two steps reaction combined zero- and first-order reactions. On the other hand,
the transition from zero-order reaction to first-order reaction and the definition of the mole fraction
range of zero- and first-order reaction were uncleared. In addition, the transition could not be welldescribed (Figure 2-7). Therefore, it can be suitable to assume the reaction to be the sum of zeroand first-order reaction as follows;(A 6-1)𝑥𝑥[−] = 𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴(𝑘𝑘1𝑡𝑡) + 𝑘𝑘0 𝑡𝑡 + 𝐵𝐵

Figure A6-1. Calculated and measured mole fraction of Mg(OH)2 for (a) LO20, (b) LO25, (c) LO33,
and (d) LO50 at 270 °C. Bule, green, and red dotted lines are measured value, calculated value by 1
step, and 2 steps model, respectively (Chapter 2 showed the detaile).
value based on eq. (A6-1).

206

Orange line is the calculatedwhere x is the mole fraction of Mg(OH)2 [-], A and B are the initial mole fraction of Mg(OH)2 forzero- and first-reactions [-], respectively. k0 and k1 are the reaction constants for zero- and first-order
reactions [s−1], respectively.

t is the reaction time [s].

The investigated range of the mole fraction

were 0.2–0.9 for the reaction at 230, 250, 270 °C, and 0.2–0.8 for 300 °C as described in Chapter 2.
The values of A, B, and reaction constants were determined by least squares method using a SOLVER
in EXCEL.
Figure A6-1 shows the comparison between the measured and calculated mole fraction based on
1 step, 2 steps, and eq. (A 6-1). This figure clearly shows that the values calculated by eq. (A6-1)
followed the measured value well. This indicates that the dehydration reaction can be described as
combination of zero- and first-order reactions rather than the 2 steps reaction (transition from zero- tofirst-order reaction).
Figure A6-2 shows the contribution of each reaction on whole reaction. These can be calculated
as follows;

Figure A6-2. Contribution of each reaction on whole reaction for (a) LO20, (b) LO25, (c) LO33, and(d) LO50 at 270 °C. Red, and blue lines are the fraction change of zero-, and first-order reactions,
respectively. These fractions were calculated based on eq. (A 6-2).

207

𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶𝐶 𝑜𝑜𝑜𝑜 𝑒𝑒𝑒𝑒𝑒𝑒ℎ 𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟𝑟[−] =𝑘𝑘0 𝑡𝑡+𝐵𝐵

𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴(𝑘𝑘1 𝑡𝑡)+𝑘𝑘0 𝑡𝑡+𝐵𝐵

𝑜𝑜𝑜𝑜

𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴(𝑘𝑘1 𝑡𝑡)(A 6-2)

𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴(𝑘𝑘1 𝑡𝑡)+𝑘𝑘0 𝑡𝑡+𝐵𝐵

These contrisutions mean the fraction of each reaction (zero- or first-order reaction) on the wholereaction. Therefore, the fraction change during recation can be observed as shown in Figure A6-2.
This figure presents the fraction of zero-order reaction increased while that of first-order reaction
decreased with increasing reaction time at 270 °C.

In this way, the fraction change can be monitored.

This might provide us some information on the dehydration reaction in the future work.
Table A6-1. Reaction rate constants of LO20, LO25, LO33, and LO50 based on the new model
k0 [s−1]

Sample230 °C

250 °C

270 °C

300 °CLO20

2.54 × 10−6

4.97 × 10−6

3.88 × 10−5

8.57 × 10−6LO25

4.07 × 10−6

2.15 × 10−7

7.43 × 10−5

4.82 × 10−4LO33

3.22 × 10−6

3.17 × 10−6

6.85 × 10−5

1.72 × 10−4LO50

2.90 × 10−6

2.61 × 10−5

1.07 × 10−4

2.86 × 10−4

k1 [s−1]LO20

3.20 × 10−5

1.30 × 10−4

6.30 × 10−4

1.01 × 10−3

LO25

5.22 × 10−5

1.45 × 10−4

1.15 × 10−3

9.20 × 10−3

LO33

5.29 × 10−5

1.29 × 10−4

1.04 × 10−3

3.34 × 10−3

LO50

5.09 × 10−5

2.46 × 10−4

1.42 × 10−3

4.46 × 10−3

Table A6-2. Constants A and B of LO20, LO25, LO33, and LO50 based on the new modelA [-]Sample230 °C

250 °C

270 °C

300 °C

LO20

0.45

0.75

0.48

0.70

LO25

0.41

0.77

0.46

0.25

LO33

0.59

0.68

0.42

0.53

LO50

0.54

0.40

0.36

0.49

B [-]
LO20

0.46

0.14

0.43

0.09

LO25

0.49

0.11

0.44

0.56

LO33

0.29

0.20

0.48

0.26

LO50

0.35

0.50

0.54

0.32

208

Table A6-1–3 show the reaction rate constants, constants A and B, and activation energies andfrequency factors of the samples based on the new model. The rate constant k1 show the typicaltemperature dependence (the values of k1 decreased with increasing temperature), while the onstant k0
of LO20 and LO25 did not show the typical trend. In addition, the frequency factor of LO20 by thezero-order reaction was quite small (Table 6-3). Therefore, it is difficult to mention that the newmodel well-described the reaction behavior.
required to be clarified.

Furthermore, the definition of constants A and B is

The new model based on eq. (A6-1) can help us to understand thedehydration reaction mechanism, on the other hand, the model is needed to be investigated in detail in
the future work
Table A6-3. Activation energies and frequency factors of LO20, LO25, LO33, and LO50 based on the
new model
Ea [kJ mol−1]

SampleA [s−1]

Zero-orderFirst-order

Zero-orderFirst-order

LO20

54

122

1.60

2.01 × 108

LO25

210

183

4.56 × 1015

4.27 × 1014

LO33

155

150

2.38 × 1010

1.79 × 1011

LO50

156

156

6.63 × 1010

9.97 × 1011

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210

Achievements
1. 論文①［国際誌 査読有り 筆頭］
1) Ryo Kurosawa, Junichi Ryu, “Effect of LiOH addition on dehydration reaction of Mg(OH)2”,
Journal of Chemical Engineering of Japan, 52, (2019), 152–158. (Chapter 2)2) Ryo Kurosawa, Masato Takeuchi, Junichi Ryu, "Comparison of the effect of co-addition of
Li compounds and addition of a single Li compound on reactivity and structure ofmagnesium hydroxide", ACS Omega, 4, (2019), 17752−17761. (Chapter 3)3) Ryo Kurosawa, Masato Takeuchi, Junichi Ryu, “Fourier-transform infrared analysis of thedehydration mechanism of Mg(OH)2 and chemically modified Mg(OH)2”, Journal of
Physical Chemistry C, 125, (2021), 5559–5571. (Chapter 4)4) Ryo. Kurosawa Masato Takeuchi, Junichi, Ryu, “Fourier-Transform Infrared and X-ray
Diffraction Analyses of the Hydration Reaction of Pure Magnesium Oxide and Chemically
Modified Magnesium Oxide”, RSC Advances, 11, (2021), 24292–24311 (Chapter 5)②［国際誌 査読有り 非筆頭］
1) Aya Maruyama, Ryo Kurosawa, Junichi Ryu, “Eﬀect of Lithium Compound Addition on the
Dehydration and Hydration of Calcium Hydroxide as a Chemical Heat Storage Material”,
ACS Omega, 5, (2020), 9820−9829.
2) Atsushi Kondo, Ryo Kurosawa, Junichi, Ryu, Masaya Matsuoka, Masato Takeuchi,“Investigation on the mechanism for the dehydration of Mg(OH)2 and hydration of MgO by
Near-infrared spectroscopy”, Journal of Physical Chemistry C, 125, (2021), 10937–10947.
③［和文誌 査読有り 非筆頭］
1) 齋藤鴻輝, 黒沢諒, 劉醇一, ”化学蓄熱のためのクエン酸化合物添加水酸化マグネシ
ウムの脱水および水和反応性”, 鉄と鋼 Vol. 106, No. 8, 549–555.
2．口頭発表・ポスター発表①［国際 筆頭または登壇者］
1) Ryo Kurosawa, Junichi Ryu, “Dehydration and Hydration Reactivity of LiOH-modified
Mg(OH)2 for Chemical Heat Storage”, International Sorption Heat Pump Conference 2017,
C-HM-1029, 2017/8
2) Ryo Kurosawa, “The effect of multiple Li compounds addition on the reactivity andstructure of M(OH)2”, Eurotherm Seminar #112, A110, 2019/5
3) Ryo Kurosawa, Junichi Ryu, “Effect of Li compounds addition on reactivity and crystal
structure of Mg(OH)2 for chemical heat storage”, The Fifth International Symposium on
Innovative Materials and Processes in Energy Systems, IMPRES2019, A102, 2019/10② [国際 非筆頭]
1) Aya Maruyama, Ryo Kurosawa, Junichi Ryu, “Dehydration and hydration reactivity of Li–compound added Ca(OH)2 for thermo–chemical energy storage, Grand Renewable Energy
2018, P–St–15, 2018/6
2) Kouki Saitou, Ryo Kurosawa, Junichi Ryu, “Dehydration and hydration reactivity of citrate
added Mg(OH)2 for thermochemical energy storage, Grand Renewable Energy 2018, P–St–
18, 2018/6
3) Junichi Ryu, Aya Maruyama, Ryo Kurosawa, “Reactivity enhancement of Ca(OH)2 withchemical modification for thermo-chemical energy storage”, Eurotherm Seminar #112,
C109 2019/5
4) Aya Maruyama, Ryo Kurosawa, Junichi Ryu, “Dehydration / hydration reactivity and crystal
structure of Li-compound added Ca(OH)2”, The Fifth International Symposium on
Innovative Materials and Processes in Energy Systems, IMPRES2019, A103 2019/10
211

5)

Masato Takeuchi, Atsushi Kondo, Ryo Kurosawa, Junichi Ryu, “Investigation on thereaction mechanism for dehydration of Mg(OH)2 and hydration of MgO by NIR
spectroscopy”, The 4th Aquaphotomics International Conference, 2021/03③［国内 筆頭または登壇者］
1) 黒沢諒, 劉醇一, “水酸化マグネシウムの脱水反応に対する水酸化リチウム添加の影響”, 第 19 回化学工学会学生発表会, B119, 東京, 2017/03
2) 黒沢諒, 劉醇一, “水酸化マグネシウムの脱水反応に対する水酸化リチウム添加の影響”, 化学工学会第 82 年会, PA137, 東京, 2017/03
3) 黒沢 諒, 劉 醇一, “水酸化リチウム添加水酸化マグネシウム化学蓄熱材の脱水反応速度解析”, 化学工学会第 49 回秋季大会, EB104, 愛知, 2017/09
4) 黒沢諒, 劉醇一, “Mg(OH)2 系化学蓄熱材における Li 化合物添加方法の反応性への影響”, 化学工学会第 83 年会, PA139, 大阪, 2018/03
5) 黒沢諒, 劉醇一, ”Li 化合物同時添加による水酸化マグネシウムの反応性とその構造に対する影響”, 化学工学会第 50 回秋季大会, DB119, 鹿児島, 2018/09
6) 黒沢 諒, 劉 醇一, “Li 化合物添加 Mg(OH)2 の反応性に対する構造的解析”, 化学工学会第 84 年会, PA150, 東京, 2019/03
7) 黒沢諒, 竹内雅人, 劉醇一, “Mg(OH)2/MgO 系化学蓄熱材の反応性に対する Li 化合物の影響”, 化学工学会第 51 回秋季大会, V209, 岩手, 2020/09
8) 黒沢諒, 劉醇一, “Mg(OH)2 化学蓄熱材の脱水反応性に対する Li 化合物添加効果の検討”, 第 10 回 CSJ 化学フェスタ, J2-12, 東京, 2020/10
9) 黒沢諒, 劉醇一, “Mg(OH)2 化学蓄熱材の脱水反応性に対する Li 化合物添加効果の検討”, 第 10 回 CSJ 化学フェスタ, P7-014, 東京, 2020/10
10) 黒沢諒, 竹内雅人, 劉醇一, “Mg(OH)2 化学蓄熱材の脱水反応性に対する Li 化合物
添加効果の検討”, 第 127 回触媒討論会, 1B 04 (B1), 2021/03
11) 黒沢諒, 竹内雅人, 劉醇一, “FT-IR と XRD を用いた Mg(OH)2/MgO 系化学蓄熱材の反応性に対する Li 化合物添加効果の検討”, 化学工学会第 86 年会, PD311, 2021/03
12) 黒沢諒, 竹内雅人, 劉醇一, “Mg(OH)2 系化学蓄熱材の反応性に対する化学修飾の影響”, 第 128 回触媒討論会, 2021/09
13) 黒沢諒, 劉醇一, “Mg(OH)2 系化学蓄熱材の脱水反応性と添加剤の性質の関係”, 化学工学会第 52 回秋季大会, VR114, 2021/09④ [国内 非筆頭]
1) 劉 醇一, 黒沢 諒, “水酸化マグネシウムの熱分解における水酸化リチウム添加効
果と化学蓄材への適用”, 第 120 回触媒討論会, 3E15A, 愛媛, 2017/09
2) 丸山 愛矢, 黒沢 諒, 劉 醇一, “水酸化カルシウムの脱水反応におけるリチウム化
合物添加効果”, 第 20 回化学工学会学生発表会, E01, 東京, 2018/03
3) 齋藤 鴻輝, 黒沢 諒, 劉 醇一, “クエン酸塩添加水酸化マグネシウム系化学蓄熱材の性能評価”, 第 20 回化学工学会学生発表会, E04, 東京, 2018/03
4) 劉 醇一, 黒沢 諒, 丸山 愛矢, “アルカリ土類水酸化物の熱分解におけるリチウム
化合物共添加効果と化学蓄熱材への適応”, 第 122 回触媒討論会, 1J01, 北海道,
2018/095) 丸山 愛矢, 黒沢 諒, 劉 醇一, “Li 化合物添加による Ca(OH)2 の脱水反応と結晶構造への影響, 化学工学会第 50 回秋季大会, DB121, 鹿児島, 2018/09
6) 劉 醇一, 丸山 愛矢, 黒沢 諒, “水酸化カルシウム系化学蓄熱材の化学修飾”, 第 8回潜熱工学シンポジウム, 愛知, 2018/12
7) 齋藤 鴻輝, 黒沢 諒, 劉 醇一, “水酸化マグネシウム系化学蓄熱材のクエン酸ナト
リウム添加による反応性評価”, 第 8 回潜熱工学シンポジウム, 愛知, 2018/12
212

8)

丸山 愛矢, 黒沢 諒, 劉 醇一, “Ca(OH)2 の脱水反応性と結晶構造に関する研究”,
化学工学会第 84 年会, PA162, 東京, 2019/03
9) 梅津 拓実, 黒沢 諒, 劉 醇一, “水酸化マグネシウムの脱水反応における遷移金属
複合効果”, 第 21 回化学工学会学生発表大会, C06, 東京, 2019/03
10) 丸山 愛矢, 黒沢 諒, 劉 醇一, “水酸化カルシウム系化学蓄熱材のアルカリ金属化
合物添加効果”, 化学工学会横浜大会, B108, 神奈川, 2019/08
11) 劉 醇一, 丸山 愛矢, 黒沢 諒, “水酸化カルシウムの熱分解におけるリチウム化合物添加効果と化学蓄熱材への適応”, 第 124 回触媒討論会, 2I13, 長崎, 2019/09
12) 齋藤 鴻輝, 黒沢 諒, 劉 醇一, “化学修飾した硫酸マグネシウム 7 水和物の蓄熱操
作性評価”, 第 9 回潜熱工学シンポジウム, 茨城, 2019/12
13) 河野俊介, 黒沢諒, 劉醇一, “水酸化ランタンの脱水/水和反応に関する金属塩添加の影響”, 化学工学会第 51 回秋季大会, V208, 岩手, 2020/09
14) 青山紗也 黒沢諒, 劉醇一, “アルカリ金属化合物を添加した Mg(OH)2 の脱水・水和に関する研究”, 第 23 回化学工学会学生発表会, H21, 2021/03
15) 永野瑞樹, 梅津拓実, 黒沢諒, 劉醇一, “機械学習による水酸化マグネシウム系化学
蓄熱材の組成最適化”, 第 23 回化学工学会学生発表会, H22, 2021/03
16) 河野俊介, 黒沢諒, 劉醇一, “水酸化ランタンの脱水/水和反応に関する Li 化合物添加の影響”, 化学工学会第 86 年会, PD317, 2021/03
17) 河野俊介, 黒沢諒, 劉醇一, “La(OH)3 の脱水/水和反応に対するアルカリ金属塩添加の影響”, 化学工学会第 52 回秋季大会, VR113, 2021/09
3．受賞・表彰等
1) 化学工学会第 83 年会 学生優秀賞受賞, 2018/03
2) 千葉大学大学院融合理工学府都市環境システムコース優秀論文賞受賞, 2019/03
3) 千葉大学大学院融合理工学府長表彰受賞, 2019/03
4) 化学工学会第 86 年会 学生奨励賞受賞, 2021/03
4．特許
1)

岡田翔太, 北垣昌規, 大塚泰弘, 塘啓祐, 劉醇一, 黒沢諒, 丸山愛矢, “化学蓄熱材及びその製造方法”, 国際出願番号 PCT/JP2019/004259 (WO/2019/159791)

5．研究費, 奨学金獲得実績
1) 千葉大学なのはなコンペ 2019 正賞, 特別賞受賞 （研究費 25 万円, 現金 10 万円）,
2018/122) 日本学生支援機構奨学金全額免除 2019/06
3) 公益財団法人双葉電子記念財団, 給付型奨学金支給内定, 120 万円, 2019/10
4) 公益財団法人双葉電子記念財団, 給付型奨学金支給内定, 120 万円, 2020/10

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Acknowledgement
私は千葉大学大学院融合理工学府都市環境システムコース所属の劉研究室で研究活動に従事させていただきました。5 年半もの間本研究室に在籍し, 劉醇一准教授には多大なる指導をいただき, 研究者として必要な素養や思考法を身に着けるための訓練をさせていただ
きました。その中で学会発表や論文執筆などの貴重な経験をさせていただきました。このよ
うな経験は博士課程まで進学しなければなかなかすることはできないことなので, 非常に勉強になりました。これらの機会を与えていただき, 感謝しております。また, 劉先生には私が地道に努力を続けていたことを評価していただき, こうした機会を与えていただいたと考えています。そのため, 自分の強みはそのように努力を続けられることにあるんだと,自信を持つことができました。さらには今後の進路についても相談に乗っていただき, 数多
くの助言をいただきました。
大阪府立大学大学院の竹内雅人准教授にも研究を進めるうえで多くの助言をいただきました。竹内先生の専門分野である赤外分光法に関する知見を応用することで, 本研究のメインである Mg(OH)2/MgO の脱水/水和反応に対する Li 化合物の添加効果に関して詳細な議論を行うことができました。竹内先生に関わっていただいたおかげで, 3 報の論文がアメリカ
化学会の ACS Omega, Journal of Physical Chemistry C さらに王立化学化学会の RSC Advancesに採択されました。私が行ってきた研究の中でもこれらの業績は代表的なものとなります。
大変お世話になりました。
博士論文の審査を引き受けてくださった, 都市環境システムコース小倉裕直教授, 共生
応用化学コース佐藤智司教授, 物質科学コース小林範久教授, 同じく物質科学コース星野
勝義教授に感謝申し上げます。お忙しい中, 私の博士論文審査会にて貴重なご意見をしていただきありがとうございました。
所属していた化学工学会においても多くの先生方にお世話になりました。特に名古屋大学の窪田光宏助教, 金沢大学の汲田幹夫教授, 北海道大学の能村貴宏准教授には学会期間中に気さくに接していただきました。また 2019 年にスペインに学会出張した際は能村先生に案内していただき, 楽しいひと時を過ごせました。今後も皆様には仲良くしていただければと思います。
タテホ化学工業株式会社の小西武様, 大塚康弘様には共同研究及び学会にて大変お世話になりました。定期的に我々の研究についてディスカッションして頂きました。忌憚ない意
見をいただき大変勉強になりました。また, 我々の研究成果を製品にまでしていただいて,非常に感謝しております。特に大塚様は私と同様に博士課程を修了した経緯がありましたので, 私が博士課程に進学しようか迷っていた際に相談に乗っていただきました。その時は親身に話を聞いてくださり, おかげで博士課程への進学を決意することができました。お二人も今後お仕事等でお付き合いが保てていければと思っております。
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株式会社熱い緑営業部の川上英治様には日頃から実験室内の備品購入の際にお世話になりました。研究室にお越しいただいた際には, 気さくに話しかけていただきました。今後も機会がありましたらお話しいただければと思います。また, これからも劉研究室とのお付き合いを続けていっていただけると幸いです。
劉研究室の後輩のみなさん, 特に博士論文発表会前には貴重なコメントをいただけたことが幸いでした。ありがとうございました。何でもいいので一生懸命取り組めること, 夢中になれることを見つけてください。きっと自分の行動がプラスの方向に変化すると思います。特に修士二年の二人に向けてアドバイスを残しておきます。二人のことはこれまで見てきて, 個人的にはもっと頭を使った方がいいかなと感じることが多かったです。修士号というものがただの飾りにならないように, 修士課程修了後どのように行動するのかが大事かと思います。肩書に惑わされないようにしてほしいし, また自分の持つ肩書と能力が見合う人であってほしいと思います。
湯君はこのまま努力を続けてください。ご両親も君が日本で頑張っていることを願っていると思うので, いい知らせができるように祈っています。また, 母国語ではない日本語もしくは英語を使ってのコミュニケーションはきっとストレスがたまると思います。それでも僕の次に努力を続けている姿勢は素晴らしいものがあると思います。多くの研究業績を挙げ, 自分自身のスキルを高めてください。
（You keep doing effort on your research, Tang. I
think your pearents hope you are hardworking. I hope you make good news for your pearents. You
have been stressful because you must mainly use Japanese or English to communicate with us.
Nevertheless, it is wonderful that you keep your effort for the research. Furtehrmore, I wish you
foucused on your work harder and harder. I hope you will make many research achievements andbrush up on your research skills.）
また, 今回の研究で得られた結果（特に 4–6 章）をうまく活用することで, 化学蓄熱材の開発を促進できればと思っています。ですので, 本論文に示されている検討結果を存分に活かしていただければ本望です。さらに僭越ながら, 劉先生に少しばかりご提案をさせていただきます。ぜひ日頃のゼミや研究活動においても, 細かな指導を学生たちにしていただきた
いと思います。論文を書く際の細かな注意点（表現方法や単位のつけ方など）も改めて指導することはなかなかないかもしれませんが, ぜひ学生たちのことを思ってそのような機会を設けてあげてください。そうすれば学会発表や論文執筆の直前にお互い基本的なところに時間を割く必要性もないと思いますので、ご検討いただけますと幸いです。
最後にこれまで私を支えてくれた家族に感謝の意を表します。これまで好きなことをやらせていただいて, 両親には非常に感謝しています。私が博士課程まで進学し, 経済的な心配も無く研究に専念できたのは, まぎれもなく二人のおかげです。特に授業料は父親に負担してもらっていました。また, 母親には修士課程に在籍していたころから毎日お弁当を作っていただきました。日頃の家事に加えて負担をかけてしまいましたが, 嫌がることなく毎日作ってくれました。このように二人には献身的に支えていただきました。一方で就職先がな215

かなか決まらず, 心配をかけてしまいました。これからはできるだけ恩返しができればと思います。また, 姉と兄と会話することで日頃のストレスを忘れることもできました。繰り返
しにはなりますが, これまで大いに支えていただいて感謝しております。これからは一人暮
らしも始まることになるので余計心配をかけることになるかもしれませんが, よろしくお願いします。

令和 3 年 9 月

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