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Hydrogen storage in carbon and silicon carbide nanotubes

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Hydrogen storage in carbon and silicon carbide nanotubes

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HYDROGEN STORAGE IN CARBON AND SILICON CARBIDE NANOTUBES

by

Seyedhamed M Barghi

A Dissertation Presented to the
FACULTY OF THE GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(CHEMICAL ENGINEERING)

May 2015

i

Table of Contents
Dedication ...................................................................................................................................... iv
Acknowledgments .......................................................................................................................... v
List of Tables. ...............................................................................................................................vii
List of Figures. ........................................................................................................................... viii
Abstract ......................................................................................................................................... xi
Chapter 1: Introduction ............................................................................................................... 1
1.1. Hydrogen as a Green Fuel ......................................................................................... 1
1.2. Hydrogen Storage Using Solid Adsorbents for on-Board Vehicle Applications ...... 3
1.3. Silicon-Carbide Nanotubes ........................................................................................ 8
1.4. Outline of the Research ........................................................................................... 11
1.5. References ............................................................................................................... 12
Chapter 2: Hydrogen Storage Using Carbon Nanotubes ........................................................ 17
2.1. Introduction ............................................................................................................. 17
2.2. Experimental ............................................................................................................ 22
2.2.1. Characterization of the Carbon Nanotubes .............................................................. 22
2.2.2. Monitoring the Gas Composition ............................................................................ 23
2.2.3. Effect of the Drift .................................................................................................... 24
ii

2.2.4. Stability of the MWCNT Sample ............................................................................ 24
2.2.5. Determining the Volume of the MWCNT Sample ................................................. 25
2.2.6. Hydrogen Adsorption Isotherms ............................................................................. 26
2.3. Results and Discussion ............................................................................................ 27
2.4. Conclusions ............................................................................................................. 37
2.5. References ............................................................................................................... 39
Chapter 3: Hydrogen Storage Using Silicon Carbide Nanotubes ........................................... 43
3.1. Introduction ............................................................................................................. 43
3.2. Experimental ............................................................................................................ 47
3.2.1. Synthesis of SiCNTs ................................................................................................ 47
3.2.2. Purification of the SiCNTs ...................................................................................... 50
3.2.3. Hydrogen Sorption Measurements .......................................................................... 50
3.3. Results and Discussion ............................................................................................ 52
3.4. Conclusions ............................................................................................................. 64
3.5. References ............................................................................................................... 68
Chapter 4: Hydrogen Storage Using Doped Silicon Carbide Nanotubes .............................. 74
4.1. Introduction ............................................................................................................. 74
4.2. Experimental ............................................................................................................ 76
4.2.1. Materials .................................................................................................................. 76
iii

4.2.2. SiCNTs synthesis and purification .......................................................................... 76
4.2.3. Doping SiCNTs with K ........................................................................................... 77
4.2.4. Doping SiCNTs with Ti ........................................................................................... 77
4.2.5. Hydrogen uptake experiment .................................................................................. 78
4.3. Results and Discussion ............................................................................................ 80
4.4. Conclusions ............................................................................................................. 87
4.5. References ............................................................................................................... 89
Chapter 5: Solubility and Diffusivity of H2 and CO2 in the Ionic Liquid [bmim][PF6] ........ 94
5.1. Introduction ............................................................................................................. 94
5.2. Experimental ............................................................................................................ 98
5.2.1. Materials .................................................................................................................. 98
5.2.2. Gas absorption measurements ................................................................................. 98
5.3. Measurement of the Diffusivities ............................................................................ 99
5.4. Results and Discussion .......................................................................................... 101
5.5. Conclusions ........................................................................................................... 112
5.6. References ............................................................................................................. 114
Chapter 6: Future Work ........................................................................................................... 117

iv

Dedication

To my loving parents, Abbas and Iran
for their kindness and support in every stage of my life.

v

Acknowledgments
I would like to express my appreciation to my advisors Professors Muhammad Sahimi and
Theodore T. Tsotsis for their precious guidance in every stage of this work. They were also
absolutely supportive in providing me with the experimental tools, without which it would
have been impossible to successfully finish this Thesis, and for which I am grateful.

My thanks and regards go to Professors John O'Brien, Katherine Shing, and Malancha
Gupta for serving on my qualifying committee. I am especially thankful to Professor John
O'Brien for letting me use the scanning electron microscope in his lab, which was a great help
for my experiments.

I would also like to thank Professor Andrea Armani for generously allowing me to use the
FTIR that belongs to her group.

I am thankful to Tina Silva and Shokry Bastorous for helping me do my experiments and
keeping our labs operating safely. The efforts of Martin Olekszyk and Angeline Fugelso for
smoothly handling our purchases and financial services were very valuable. I want to thank
Heather Alexander for organizing the inspiring social events in our department. The amazing
work of Karen Woo, Heather Alexander, Andy Chen, Jason Ordonez, and Laura Carlos to
handle my administrative tasks is greatly appreciated. They gave me so much positive energy
throughout the four years of my Ph.D. work that I feel truly lucky for working with such
amazing persons.
vi

Financial support of the University of Southern California for a Provost’s Doctoral
Fellowship, the National Science Foundation, and the U.S. Department of Energy is gratefully
acknowledged.

vii

List of Tables:

Table 3. 1. The true (skeleton) density of the as-synthesized and purified SiCNTs and the
corresponding calculated SiO2 content based on the true density values.............. 67

Table 5. 1. Henry's constant values for CO2 in [bmim][PF6]. ................................................ 105
Table 5. 2. The hydrogen Henry's constant values in [bmim][PF6]. ...................................... 108

viii

List of Figures:

Figure 1. 1. Side- and top-view of the most stable SiCNTs with alternating Si and C atoms
[38]. ......................................................................................................................... 9
Figure 1. 2. Doping of Ti atom on the SiC nanotube: (a) top-view, (b) side-view [41]. .......... 10

Figure 2. 1. (a) Analysis of the gas streams entering and leaving the adsorption vessel using
the residual gas analyzer. (b) Weight of the empty sample container and the fresh
sample’s weight vs. time under dynamic vacuum at pressures of 10
-5
bar. (c) Plot
of Eqn. 1 for the MWCNTs for various pressures of Helium gas. ........................ 30
Figure 2. 2. (a) Hydrogen adsorption at the beginning of the 1
st
adsorption/desorption cycle
(step from 0 to 5 bar); (b) Hydrogen desorption at the end of 1
st
hydrogen
adsorption/desorption cycle (from 5 to 0 bar). (c) Desorption of chemisorbed
hydrogen under vacuum at 120
º
C after the end of the 3
rd
hydrogen
adsorption/desorption cycle ................................................................................... 32
Figure 2. 3. The 1
st
and 2
nd
hydrogen adsorption/desorption cycles (note that for the 2
nd
cycle
the adsorption and desorption branches are indistinguishable). ............................ 34
Figure 2. 4. Hydrogen uptake (w) and absolute value of hydrogen uptake change rate (dw/dt).
............................................................................................................................... 35
Figure 2. 5. The calculated sorption isotherms in the MWCNTs at 25
◦
C. ............................... 36
Figure 2. 6. The amount of hydrogen physisorption on the MWCNTs during the first sorption
cycle and the equilibrium hydrogen values during the second and third hydrogen
sorption cycles (the amount of the strongly sorbed hydrogen after the 1
st
cycle is
subtracted). ............................................................................................................ 37

Figure 3. 1. The experimental layout for synthesis of SiCNTs. ............................................... 49
ix

Figure 3. 2. The XRD spectra of CNTs, the as-synthesized SiCNTs produced at various
temperatures in the range 1200-1500
o
C, and the purified SiCNTs. ..................... 53
Figure 3. 3. FTIR results for the as-synthesized and purified SiCNTs at 1200
o
C. ................. 53
Figure 3. 4.(A) SEM image of the as-synthesized SiCNTs at 1200
o
C; (B) purified SiCNTs;
(C) TEM images of the as-synthesized SiCNTs at 1200
o
C, and (D) purified
SiCNTs. ................................................................................................................. 56
Figure 3. 5. Hydrogen adsorption isotherm at 25
o
C for the precursor CNTs, the as-
synthesized SiCNTs produced at 1200-1500
o
C, and the purified SiCNTs prepared
at 1200
o
C and at 1500
o
C. .................................................................................... 58
Figure 3. 6. Hydrogen adsorption/desorption cycle at 25
o
C for the as-synthesized SiCNTs
produced at 1200
o
C, the purified SiCNTs, and the initial CNTs. ........................ 59
Figure 3. 7. The rate of hydrogen uptake in SiCNTs and CNTs at 25
o
C. ............................... 63

Figure 4. 1. SEM image of, (A) pure SiCNTs; (B) K-doped SiCNTs; (C) Ti-doped before
calcination, and (D) Ti-doped after calcination..................................................... 84
Figure 4. 2. Hydrogen uptakes for the SiCNTs doped with Potassium (K-doped) and Titanium
(Ti-doped). The uptakes for the pure SiCNTs (before doping), and CNTs are also
showed for comparison. ........................................................................................ 85
Figure 4. 3. Additive hydrogen storage capacity (AHSC) values for the K-doped and Ti-doped
SiCNTs . The AHSC values for the pure SiCNTs (before doping), and CNTs are
also shown for comparison.. .................................................................................. 86

Figure 5. 1. Schematic of the gas/ionic liquid system used in this work. .............................. 101
Figure 5. 2. The experimental data and their fit to Eq. (3) for CO2 molality in [bmim][PF6] at
0.04 MPa and 35
o
C. ............................................................................................ 102
Figure 5. 3. The experimental data for H2 molality in [bmim][PF6] at 2 MPa and 35
o
C, and
x

their fit to Eq. (3). ................................................................................................ 103
Figure 5. 4. H2 and CO2 diffusivity vs. pressure at 25
o
C. ..................................................... 104
Figure 5. 5. Pressure-dependence of CO2 molality in [bmim][PF6] in the temperature range at
25
o
C- 55
o
C. The lines represent the best fit of the data. ................................... 105
Figure 5. 6. Pressure-dependence of H2 molality in [bmim][PF6] in the temperature range 25
o
C-55
o
C. The lines represent the best fit of the data. ......................................... 105
Figure 5. 7. Time evolution of hydrogen dissolution in [bmim][PF6] at 25
o
C and 55
o
C (P=2
MPa for both temperatures). The solid lines represent the experimental fit
provided by Eq. 3. ............................................................................................... 109
Figure 5. 8. Temperature-dependence of Henry's constants of H2 and CO2 in [bmim][PF6].
The lines represent the fit of the data to Eq. (4). ................................................. 110
Figure 5. 9. Temperature-dependence of diffusivity of H2 and CO2 in [bmim][PF6]. The lines
represent the fit of the data to Eq. (5). ................................................................. 111

xi

Abstract
Hydrogen is considered a promising renewable energy source. Developing a safe and
efficient hydrogen storage medium is the most crucial challenge for the commercial
application of hydrogen as fuel for vehicles. Adsorptive hydrogen storage on solid adsorbents,
including carbon nanotubes (CNTs) and silicon-carbide nanotubes (SiCNTs), can be a novel
storage technique. In the present work, the hydrogen storage behavior of CNTs, pure SiCNTs,
and doped SiCNTs is investigated experimentally using a combination of gravimetric and
mass analysis methods at room temperature for pressures of up to 100 bar. The combination
of the techniques enabled us to obtain experimental data for hydrogen adsorption and
desorption, and the rate by which they occur.
In the first part of this work, the question of chemisorption versus physisorption during
hydrogen storage in carbon nanotubes (CNTs) is addressed experimentally. We utilize a
powerful measurement technique based on a magnetic suspension balance coupled with a
residual gas analyzer, and report new data for hydrogen sorption at pressures of up to 100 bar
at 25 ℃. The measured sorption capacity is less than 0.2 wt.%, and there is hysteresis in the
sorption isotherms when multi-walled CNTs are exposed to hydrogen after pretreatment at
elevated temperatures. The cause of the hysteresis is then studied, and is shown to be due to a
combination of weak sorption - physisorption - and strong sorption – chemisorption - in the
CNTs. Analysis of the experimental data enables us to calculate separately the individual
hydrogen physisorption and chemisorption isotherms in CNTs that, to our knowledge, are
reported for the first time here. The maximum measured hydrogen physisorption and
chemisorption are 0.13 wt.% and 0.058 wt.%, respectively.
xii

In the second part of this work, we report for the first time the results of an extensive
experimental study of hydrogen sorption in SiCNTs, which were synthesized using the
reaction between SiO vapor and CNTs in an argon atmosphere in the temperature range 1200

o
C - 1500
o
C. The as-synthesized SiCNTs were then purified using a sodium hydroxide
solution, in order to remove the side products of the synthesis reaction. The hydrogen sorption
characteristics of the as-synthesized SiCNTs, as well as those of the purified SiCNTs were
then measured at 25
o
C and for pressures of up to 100 bar. The results reveal hysteresis
between the adsorption and desorption isotherms, which we attribute to the presence of metal
impurities and/or the multilayer structure of the nanotubes. The hydrogen storage capacity of
the as-synthesized SiCNTs is similar to that of the CNTs, whereas for the purified SiCNTs it
is 50% higher than that of the CNTs. In addition, the hydrogen uptake rate in the SiCNTs is
about five times faster than that in the CNTs and, in contrast with the CNTs, its desorption
from SiCNTs is completely reversible under vacuum.
In the third part of this work, the SiCNTs were synthesized by the gas/solid reaction
between SiO vapor, as the Si source, and carbon nanotubes as the carbon precursor. The
resulting SiCNTs were then purified with a hot and concentrated NaOH solution in order to
remove the amorphous silica from their surface. The purified nanotubes were then doped with
either K or Ti. The hydrogen adsorption behavior of the nanotubes was gravimetrically
measured with the aid of a magnetic suspension balance. According to the results, the K-
doped silicon carbide nanotubes show more promise as hydrogen adsorption materials among
all the nanotube samples.
xiii

In the last part of this work, we report on the measurement of the solubility and diffusivity
of H2 and CO2 in the ionic liquid 1-butyl-3-methylimidazolium hexafluorophosphate
([bmim][PF6]) over the temperature range 25
o
C - 55
o
C and for pressures of up to 10 MPa.
The data were obtained using a magnetic suspension balance, a gravimetric instrument that is
capable of simultaneously and accurately measuring the gas solubility and diffusivity in such
liquids. To our knowledge, this is the first time that the H2 diffusivity has been measured in an
ionic liquid. While solubility data for H2 have been previously reported, they vary widely. The
cause for such variations is discussed as well. The diffusivity data for both H2 and CO2 follow
an Arrhenius-like dependence on temperature.

1

Chapter 1
Introduction

1.1. Hydrogen as a Green Fuel
The use of non-renewable fossil fuels has brought about increasingly serious
environmental pollution problems and concerns about a looming energy crisis. Extensive
research has recently been carried out on finding alternate renewable energy resources in
order to eliminate the world’s dependence on fossil fuels. If one is to look for a new fuel
source with such characteristics as unlimited supply, zero emissions of greenhouse gases, and
high energy efficiency, then, hydrogen is the leading candidate for the future energy industry
[1]. Hydrogen is considered as a green fuel source because it burns cleanly and its combustion
product is only water. Its other advantage is the fact that it can be produced in all parts of the
world [2]. Unlike natural gas and oil, which are produced after millions of years of organic
decay, hydrogen can be produced in refineries and chemical plants in a short time [3] using
renewable sources, such as biomass and water.
The main sources of hydrogen atoms in nature are fossil fuels and water. Raw materials,
such as petroleum crude, natural gas, biomass or coal, contain hydrogen atoms that are
bonded to carbon atoms. In water, an oxygen atom is bonded to two hydrogen atoms. To
obtain hydrogen as a fuel source, it must be stripped from such sources in a process that uses
2

energy. For example, electrolysis produces hydrogen from water [4], while gasification and
steam reforming produces hydrogen from a variety of sources including renewable biomass.
The current primary method used to generate hydrogen commercially is through steam-
methane reforming (SMR). In the United States, around 95% of the hydrogen is produced by
this method [5]. In this process, steam at around 700 - 1000 ℃ is used to heat-up the methane
in a reaction vessel in the presence of a (metal-based) catalyst under 3 to 25 bar of pressure.
The ratio of the steam to methane is kept around three [5]. The SMR reaction is described as
follows,
CH4 + H2O (+heat) → CO + 3H2 (1.1)
This reaction is typically accompanied by the water-gas-shift reaction that produces additional
hydrogen via the reaction of CO with steam. Other raw materials, including propane,
methanol and renewable ethanol, can also produce hydrogen via the catalytic reforming
reaction with steam [6].
In another common approach, hydrogen is produced directly from water through
electrolysis. In this process, an electrical current provides the energy for a redox reaction [7].
Hydrogen is reduced, while oxygen is oxidized so that both return to their elemental forms
according to the following reactions
Anode Reaction: 2H2O→ O2 + 4H
+
+ 4e- (1.2)
Cathode Reaction: 4H
+
+ 4e- → 2H2 (1.3)
Overall Reaction: 2H2O → 2H2 + O2 (1.4)
3

The reactions take place in an electrolyzer. In a polymer electrolyte membrane (PEM)
electrolyzer, the electrolyte is a solid polymeric material. Water reacts at the anode to produce
elemental oxygen and hydrogen ions. Electrons then move from the PEM to the cathode
where the hydrogen ions combine with the electrons to form hydrogen gas. Hydrogen
production via water electrolysis is considered as the most environmentally-friendly process
to produce hydrogen, as long as the electricity source is “green”, e.g., solar energy, wind
power, hydropower, ocean power, tidal power, wave power, etc. [8]. Unfortunately,
electrolysis is mainly used at small scales with appliance-sized electrolyzers. In addition, the
electrolyzers that are currently available are too expensive to be worth their output [7]. The
electricity used in electrolysis needs to be around ten times cheaper in order for this method of
hydrogen production to be competitive with SMR of natural gas.

1.2. Hydrogen Storage Using Solid Adsorbents for on-Board Vehicle
Applications
Hydrogen is the ideal candidate as an energy carrier for both mobile and stationary
applications, while averting adverse influences on the environment and reducing dependence
on imported oil for countries without such natural resources. For mobile applications,
hydrogen can be produced on-board the vehicle by reforming of liquid fuels [9] (e.g.,
methanol). Alternately, it can be produced centrally (e.g., in a refinery) and then distributed
and stored in the vehicle as, (i) pressurized gas; (ii) cryogenic liquid, and (iii) via its chemical
or physical combination with materials, such as organic hydrides, metal hydrides, and carbon
materials. Each of these options possesses attractive attributes for hydrogen storage [10].
4

Traditional technologies permit direct storage of hydrogen either in a gaseous or in a liquid
state in pressurized or in cryogenic tanks. Complications here result because of the
hydrogen’s low boiling point (−252.87 ℃) and low density in the gaseous state (0.08988 g/L
at 1 atm). Liquid hydrogen requires the addition of a refrigeration unit to maintain a cryogenic
state [11], thus adding weight and energy costs, and resulting in 40% loss in the energy
content [12]. High-pressure storage of hydrogen gas is limited by the weight of the storage
canisters and the potential for developing leaks. Moreover, hydrogen storage either in liquid
or in a gaseous form poses important safety problems for vehicular applications.
Hydrogen storage via sorption in various materials is another approach [14]. Carbon-based
materials and metal hydrides are the two major types of sorbents that have been considered to
date, and have been the subject of many studies over the past 10 years [15, 16]. There are
currently a few commercially available metal hydrides, belonging to the AB, AB2 and AB5
families, which easily absorb and desorb hydrogen at room temperature, but their maximum
hydrogen content is no more than 1.5 wt%, which is too low for on-board storage application.
For example, the United States Department of Energy (DOE) published a long-term vision for
hydrogen-storage applications [13]. According to this report, for commercial viability the
predicted minimum hydrogen-storage capacity should be at least 6.5 wt%, with 65 g/L of
hydrogen available at the decomposition temperature between 60 and 120 ℃.
Magnesium-based materials are a very promising alternative, due to their high hydrogen
storage capacity, of more than 6 wt%, and their low cost. According to Stampfer et al. [17],
one problem with MgH2, however, is the high value of its dissociation enthalpy (H = −74.5
kJ/mol) and its very low plateau pressure at ambient temperature (about 0.1 Pa). Thus, to
5

achieve a plateau pressure higher than the atmospheric pressure in order to be able to release
the stored hydrogen, it is required to raise the temperature to about 573 K. Many efforts have
been done to improve the thermodynamic properties of MgH2. One possibility is to alloy Mg
with other metals. For example, such alloys as Mg2Ni, Mg2Cu, Mg3Al2 have lower enthalpy
of formation and higher plateau pressure at ambient temperature than MgH 2 alone, but a lower
hydrogen storage capacity. Another possibility is to treat MgH2 powders with high energy
ball-milling with a dual purpose of increasing the surface to bulk ratio to enhance the gas–
solid reactivity, and mixing the hydride with some additive having catalytic effects. Liang
[18] claimed a slight increase of plateau pressure and a decrease in the formation enthalpy as
a result of ball-milling MgH2 together with indium or cadmium, but the metastable structures
obtained after milling tended to decompose after a number of hydrogen absorption/desorption
(a/d) cycles. Thus, ball-milling does not modify permanently the thermodynamics of the a/d
process, as confirmed by others [19, 20] as well. Another difficulty is that MgH2, like other
light metal hydrides, has slow kinetics of hydrogen a/d. The a/d kinetics are important,
especially for vehicular use, which requires short times of refueling the hydrogen reservoir
and a rate of hydrogen release which is high enough to fulfill the power requirements.
Among the various types of carbon-based adsorbents proposed for hydrogen storage,
carbon nanotubes (CNTs) have received considerable attention in recent years, due to their
high surface area, nanometer size pores with a narrow pore size distribution, and low mass
density [21]. In recent years, several research groups have reported that CNTs could adsorb a
great deal of hydrogen (>3 wt.% at ambient temperature) [22-24]. Some theoretical studies by
Monte Carlo simulations and other types of molecular calculations supported the findings [25,
6

26], which resulted in CNTs becoming one of the most studied materials for hydrogen storage
applications due to their novel structure characteristics.
Elemental carbon in the sp
2
hybridization state can form a variety of interesting solid
structures, such as graphite, graphene, CNTs and fullerene [27]. CNTs, first discovered by
Iijima [28], are regarded as a unique material with hollow tubular structures of nanometer
dimensions and large length/diameter ratios made from rolled-up graphite sheets, with their
honeycomb arrangement of the carbon atoms [29]. The CNTs can be divided into two
different types: single-walled carbon nanotubes (SWCNTs) and multi-walled carbon
nanotubes (MWCNTs). SWCNTs consist of a seamless graphene sheet rolled-up into a
hollow cylinder of a few nanometers in diameter and several microns in length. Most of them
are aligned and packed together to form “ropes” of 10 to 100 parallel tubes. A MWCNT
consists of several (from tens up to hundreds) concentric tubular graphite sheets, separated
from each other by a distance of ∼0.34 nm. In each tube, the carbon atoms are arranged in a
helical fashion along the tube axis. The outer diameter of the MWCNTs is typically several
tens of nanometers, and they have a length of 10–100 μm [27]. Depending on the type of
rolling, the CNTs can be further classified as armchair, zigzag and chiral type [30].
Since the discovery of CNTs, these new carbon tubular macromolecules have attracted
great attention owing to their important characteristics, such as a unique structure, high
surface area, a highly porous nature, low density, high strength, good electrical conductivity,
extraordinary mechanical and thermal properties, special functional properties, and relatively
good chemical stability [30, 31]. Many physical or chemical properties of nanotubes are
associated with the smooth, straight, one-dimensional channel present in their cores. The
7

inside channels can hold other molecules by either capillarity or via adsorption; this then
opens the unique possibilities of using nanotubes as super-adsorbents, nano-sized containers
and reactors, and as templates for the fabrication of a variety of other novel, one-dimensional
nanomaterials such as nanowires [32].
The CNTs have potential for use as effective high-capacity hydrogen storage media.
Several reviews have already been published dealing with the applications of CNTs for
hydrogen storage [33-35]. Darkrim et al. [33] discussed and compared various results for
hydrogen adsorption in the CNT materials over a wide range of pressures and temperatures.
Both experimental data and simulation results were reported. It was concluded that controlling
the microscopic characteristics of the CNTs enables one to control the gas adsorption.
Moreover, it was noted that both material synthesis and purification will need to be optimized
in order to improve the gas adsorption. Theoretical predictions and experimental results on the
hydrogen uptake of the CNTs and nanofibers were also summarized by Cheng et al. [34].
They pointed out that in order to accelerate the use of the CNTs and nanofibers as a practical
hydrogen storage medium in fuel cell-driven vehicles, efforts must be devoted to reproducing
and verifying the results, and to investigate the CNT cycling characteristics, volumetric
capacity, and release behavior. In a recent review of Yurum et al. [35], it was noted that
despite all the encouraging developments, current nanomaterial technologies still remain far
from meeting the DOE target of 6.5 wt.% hydrogen loading. Furthermore, while hydrogen
storage capacities and kinetics have been satisfactorily quantified in carbonaceous materials,
the mechanisms of hydrogen uptake and release remain to be better elucidated. Developing
future materials for hydrogen storage appears to depend on a better understanding of a variety
of factors, such as the nature of the surface functional groups, the pore and surface
8

microstructure and topology, the adsorption and desorption properties, the thermodynamic
and kinetic behavior of pure materials as well as their metal-doped composites, and the
hydrogen uptake–release mechanism.
Several studies have reported that further modifying the CNTs could improve their
hydrogen uptake capacity. For instance, Lee et al. [36] used CO2 activation in order to prepare
activated multi-walled carbon nanotubes (Acti-MWNTs) with well-developed pore structures,
high specific surface area, and higher hydrogen adsorption capacities. The activation was
performed in the temperature range 500 - 1100 ℃. Lee et al. [36] stated that the hydrogen
storage capacities of the Acti-MWNTs were enhanced to 0.78 wt.% by increasing the
activation temperatures to 900 ℃, while the hydrogen uptake capacity of the unmodified
CNTs was less than 0.3 wt.%. The increase in the hydrogen uptake capacity of the Acti-
MWCNTs was a result of the formation of a defective structure during the CO2 activation
process.

1.3. Silicon-Carbide Nanotubes
Silicon-carbide nanotubes (SiCNTs) are another promising nanomaterial for hydrogen
storage. Compared to carbon, silicon has more electrons in its outer shells, which leads to
higher polarizability and a stronger dispersion force. Accordingly, the SiCNTs may exhibit a
stronger van der Waals (vdW) attraction to hydrogen than the CNTs [37]. A theoretical study
conducted by Mpourmpakis et al. [38] showed that between the two energetically stable
forms of SiCNTs, the one in which the Si and C atoms have alternating positions in the tube
wall is full of point charges, see Figure 1.1. This happens because there is a continuous charge
9

transfer of more than half an electron from Si to C. The existence of point charges on the
surface of SiCNTs, which were first synthesized in 2001 [39], makes them good candidate
materials for hydrogen storage.

Figure 1. 1. Side- and top-view of the most stable SiCNTs with alternating Si and C atoms
[38].

Doping SiCNTs with transition or alkali metals may be a promising modification to
improve their hydrogen adsorption capacity [40-42]. Meng et al. [40] used a first-principles
simulation method to investigate the adsorption of a single Ti atom on single-wall SiCNTs
(SWSiCNTs), and the corresponding adsorption of hydrogen molecules onto this Ti atom.
They studied the favorable sites for a single Ti atom to be adsorbed on the outside surface of
the SWSiCNTs, and concluded that Ti and C atoms form a strong chemical bond, while the
interaction between Ti and Si is rather weak. The simulation results suggested that up to four
hydrogen molecules can be attached onto the Ti atom.
10

Ti doping also proved to be an efficient way to intensify the hydrogen uptake capacity of
SiCNTs in the study of Banerjee et al. [41]. They reported hydrogen adsorption isotherms in
two SiC nanostructures, namely a planar sheet and a nanotube of the type (10,0) of 1 nm in
diameter, decorated by Ti atoms. The lowest energy structure of the Ti-decorated SiC sheet
shows that the Ti atom distorts the sheet in such a way that one of the Si atoms recedes under
the plane and the Ti atom binds with the nearest three C atoms. The interaction of the Ti-
decorated sheet with hydrogen suggests that each Ti atom can bind to up to four hydrogen
molecules (all of which are adsorbed in the molecular form) with an average binding energy
of 0.37 eV. For SiCNTs, the adsorption of Ti favors the hexagonal hollow site. Moreover,
interaction of the Ti-decorated tube with hydrogen leads to the dissociation of the first
hydrogen molecule in the atomic form; thereafter, it adsorbs hydrogen in the molecular form.
The average binding energy of hydrogen molecules on the Ti-decorated tube is estimated to
be 0.65 eV [41]. Based on these results, Banerjee et al. [41] inferred that the Ti-decorated SiC
nanostructures moderately bind with the hydrogen molecules within the energy window for
hydrogen storage materials and, therefore, can be considered as one of the potentially
promising hydrogen storage materials.

Figure 1. 2. Doping of Ti atom on the SiC nanotube: (a) top-view, (b) side-view [41].
11

In another study, Wang et al. [42] investigated the interaction of hydrogen molecules (H2)
with a single lithium (Li)-doped SiCNT, using ab-initio density-functional theory. According
to them, hydrogen molecules physisorb onto a pure SiCNT with a binding energy of ~0.086
eV. However, the binding energy rises to 0.211 eV when H2 binds to a Li-doped SiCNT. The
increase in the binding energy is due to the charge transfer from Li to the nanotube. Up to four
H2 molecules can be attached to a Li-doped SiCNT with an average binding energy of 0.165
eV, which is close to the lowest requirement proposed by the U.S. DOE, and indicates that
this system is potentially a good storage medium for H2 [42].

1.4. Outline of the Research
The discussions so far make it clear that SiCNTs are promising nanostructured materials
for hydrogen storage, and might have the potential to satisfy the requirements set-forth by the
U.S. DOE. What is currently lacking are detailed experimental data for hydrogen uptake in
SiCNTs, their comparison with corresponding data for the CNTs, and the modification and
optimization of the nanotubes for achieving higher hydrogen storage capacity. This is the
primary goal that will be pursued in this research. In addition, although ionic liquids have the
potential to contribute to the hydrogen production and separation areas, the lack of reliable
hydrogen solubility and diffusivity data in such materials hinders their practical applications.
Therefore, the novel experimental method developed in this work is also used to determine
such characteristics in [bmim][PF6] as a representative of the imidazolium-based ionic liquids.
12

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17

Chapter 2
Hydrogen Storage Using Carbon
Nanotubes
2.
2.1. Introduction
1

Among the various types of carbon-based adsorbents proposed for hydrogen storage,
carbon nanotubes (CNTs) have received considerable attention in recent years, due to their
high surface area, nanometer size pores with a narrow pore size distribution, and low mass
density [1-3]. To investigate the performance of CNTs for hydrogen storage, it is necessary to
accurately determine the equilibrium hydrogen adsorption/desorption isotherms on such
materials. In particular, the kinetics of adsorption are highly important in determining the
charging time of any practical hydrogen storage system, as well as the method by which such
charging may be implemented. Moreover, in order to be able to ensure that the storage system
is capable of delivering the adsorbed hydrogen to the vehicle’s engine with an appropriate
rate, the kinetics of hydrogen desorption from the CNTs must also be measured.
To date, three techniques have been used to measure hydrogen adsorption in CNTs,
namely, the Sievert method, thermal desorption spectroscopy (TDS), and the gravimetric

1
The material in this Chapter is part of a paper previously published: Barghi SH, Tsotsis TT, Sahimi M. Chemisorption,
physisorption and hysteresis during hydrogen storage in carbon nanotubes. Int. J. Hydrogen Energy 2014;39:1390-7.
18

method [4].

The Sievert method is a commonly-used volumetric technique based on changes
in the pressure of the measurement vessel caused by adsorption/desorption of the adsorptive
gas. Its simplicity and the fact that it may better approximate the conditions under which
commercial hydrogen storage systems operate are, reportedly, the key advantages of the
method. On the other hand, since in the Sievert method the change in the pressure of the
adsorption vessel is the main means via which adsorption and/or desorption amounts are
measured, omnipresent leaks during the experiments can (and often do) lead to erroneous
results [5]. In addition, since vessel fluid dynamics interfere with measurements at short
times, it is difficult to accurately measure fast sorption kinetics and, thus, to accurately dif-
ferentiate by this method between weak – physisorption - and strong – chemisorption -
isotherms, as we are able to do with the technique used in this study. Another drawback of the
Sievert method is the fact that atmospheric pressure is the lowest pressure at which desorption
experiments may be conveniently carried out [6].
The TDS method is based on measuring the hydrogen that is adsorbed in CNTs via its
desorption in high vacuum using mass spectrometry. The high sensitivity of the technique,
allowing it to detect adsorption even by a small amount of CNTs (as low as 1 mg), is the main
advantage of the method [7].

On the other hand, it is not possible to use this technique for
identifying adsorption hysteresis that may occur, or to differentiate among the various types
of sorption that may take place simultaneously. In addition, since hydrogen desorption
happens under high vacuum, it is difficult to gauge from the isotherms measured by this
method whether the material studied can function as a practical storage medium in the range
of pressures (of up to 100 bar) suggested by the U.S. Department of Energy [8].
19

The gravimetric method is the third technique utilized for the analysis of hydrogen
adsorption. In this method, solid adsorbents are exposed to the adsorptive gases at various
pressures, while the gravimetric balance monitors the weight change of the adsorbent that is
caused by adsorption. The advantage of this method, compared to the TDS, is that one can
independently determine both the adsorption and desorption isotherms, as we do in this study
and, thus, detect the presence of hysteresis, In addition, one is able to study fast dynamics and,
thus, to discriminate among various types of sorption phenomena occurring either
simultaneously or at distinct time scales. Compared to the Sievert method, the advantage of
the gravimetric method is that it is not affected by the often inevitable gas leaks that develop
in the measurement cell [4].

Though the technique cannot, by itself, distinguish the various
gases adsorbed (a key requirement for studying multi-component adsorption), coupling the
gravimetric balance equipment to a sensitive mass analyzer (as is the case with the
experimental system used in this study) overcomes this main drawback [5].
The hydrogen uptake capacities of carbon nanomaterials published in the literature are
remarkably scattered. Geng et al. [9] reported, for example, that CNTs are capable of
adsorbing 0.1 wt.% hydrogen at 293 K and 10 MPa, whereas Chambers et al. [10] claimed
that tubular graphite nanofibers adsorb 11.26 wt.% hydrogen at 298 K and 11.35 MPa. Other
experimental data, which are in variance with the two aforementioned studies, have been
reported as well [11]. Such inconsistent results have been a matter of controversy in the area
of hydrogen storage using carbon nanomaterials. To resolve the controversy, Tibbetts et al.
[12] examined hydrogen sorption in nine carbon materials, including graphite particles,
activated carbon, graphitized PYROGRAF vapor-grown carbon fibers (VGCF), CO and air-
etched PYROGRAF fibers, Showa-Denko VGCF, carbon filaments, and nanotubes from the
20

MER Corp. and Rice University. Sorption experiments were carried out at temperatures
between -80
◦
C and 500
◦
C and at pressures of up to 11 MPa. The maximum hydrogen uptake
for the nine carbon materials at room temperature was 0.1 wt.%. According to Tibbetts et al.
[12], the hydrogen adsorption capacity of a number of the carbon materials at room
temperature is so low that, without doing careful calibration, it is impossible to even detect it
with a reliable accuracy. These findings, potentially, cast doubt on the prior experimental
work that has claimed hydrogen uptake capacities higher than 1 wt.% for carbon materials at
room temperature.
In another work, Zuttel et al. [13] investigated the hydrogen storage capacity of a number
of carbon materials, including multi-walled CNTs (MWCNTs) fabricated by the pyrolysis of
acetylene. Their experiments yielded hydrogen storage capacities, at room temperature and at
a pressure of 10 MPa, not greater than 0.6% (on a per mass basis). Zuttel et al. [13] suggested
that the early reports of extraordinary large hydrogen adsorption capacities of CNTs must,
therefore, be viewed with skepticism. In addition to the aforementioned experimental studies,
molecular simulations have also indicated relatively low hydrogen adsorption capacities for
CNTs. Dodziuk and Dolgonos [14] used molecular mechanics calculations and molecular
dynamics simulations, for example, to study the ability of individual armchair, zigzag, and
chiral CNT, and bundles of nanotubes to function as hydrogen storage media, using the
consistent-valence force field (CVFF) and the extensible systematic force field (ESFF) in their
simulations. Their results suggested that high hydrogen storage in CNTs cannot be achieved
through physisorption alone.
21

Hydrogen adsorption/desorption hysteresis in CNTs has been reported previously in
studies that used the conventional Sievert method to measure the sorption isotherms [6, 15,
16]. It has been hypothesized that the reason for observing the hysteresis may be the
formation of some type of chemical bonding between hydrogen and the metal catalyst
residues or the amorphous carbon in CNTs. Previous studies have not, however, been able to
perform quantitative experiments in order to confirm this hypothesis. For example, one cannot
exclude conclusively that the observed differences between the adsorption and desorption
isotherm branches are not simply due to the experimental difficulties associated with leaks in
the Sievert measurement vessel [4]. Another possible scenario is that the hysteresis may have
been a result of chemisorption of hydrogen on CNTs, which has been reported [17] as an
adsorption mechanism in CNTs. Thus, it is crucial to be able to clearly distinguish between
hydrogen chemisorption from physisorption occurring in CNTs.
In this Chapter, we specifically address the question of hysteresis during hydrogen sorption
in CNTs, as well as whether chemisorption or physisorption occurs in such nanostructured
materials. Another motivation for the present work is the need for high-precision experiments,
so that one can potentially identify the reason for the scattered hydrogen adsorption uptake
data in the literature. We report here new measurements of hydrogen sorption in CNTs, using
a precise measurement technique that enables us to measure separately the contributions of
chemisorption and physisorption to the overall storage capacity of CNTs. In contrast, most
measurement techniques used in the past [16] cannot readily differentiate between
physisorption and chemisorption. To carry out the experiments, we use a high-pressure
gravimetric instrument, coupled with a highly-sensitive mass analyzer system for measuring
hydrogen sorption isotherms of MWCNTs. To our knowledge, the technique we use has not
22

been utilized before in such studies, and there have been no previous studies that could
unequivocally distinguish between chemisorption and physisorption of hydrogen occurring on
MWCNTs at room temperature. This has been made possible in this study by the combination
of the gravimetric and mass spectrometric techniques because:
(i) Hydrogen adsorption and desorption isotherms are measured accurately and
simultaneously with the same experimental method. This is not possible with the TDS
method, whereas the Sievert method is sensitive to system leaks [19-22]. Therefore, since the
two methods are distinct with their own inherent experimental errors, a quantitative
comparison between the results obtained with these two methods is not reliable.
(ii) With our experimental technique, the presence of potential impurities in the adsorptive
gas or desorption of volatile impurities from the solid sample can be detected with the residual
gas analyzer that is connected to the gravimetric equipment. This, then, distinguishes the
method from the gravimetric method used alone, or from the conventional Sievert method.

2.2. Experimental
2.2.1. Characterization of the Carbon Nanotubes
The hydrogen used in the experiments was ultrahigh purity (UHP) grade (99.999%) from
Gilmore. The MWCNTs with a purity of 97.46 wt.% were purchased from Nanostructure and
Amorphous Materials, Inc. (the manufacturer reports the following other elements as
impurities: Al, 0.19 wt.%; Cl, 1.02 wt.%; Co, 1.09 wt.%, and S, 0.24 wt.%). The MWCNTs
are produced by natural gas catalytic decomposition over a Co-based catalyst. The inner and
23

outer diameters of the nanotubes are approximately 5 and 8 nm, respectively. Specific surface
area, average pore diameter, and pore volume of the MWCNTs sample were determined by us
via nitrogen adsorption at its normal boiling temperature using an ASAP 2010 BET
instrument. The crystalline structure of MWCNTs was examined with the aid of a Rigaku
XRD equipment.

2.2.2. Monitoring the Gas Composition
The composition of hydrogen gas used in the experiments was monitored with a residual
gas analyzer (RGA200), manufactured by Stanford Research Systems (SRS), Inc. The RGA
system was equipped with an electron multiplier that makes it possible to measure partial
pressures of species as low as 10
-14
Torr. The effect of the background noise on the mass
analysis results was corrected by running the RGA200 system without injecting any gases. In
the first step, the composition of the UHP hydrogen was determined prior to injecting the gas
into the system, in order to ensure that there were no unexpected impurities in the hydrogen
feed. In the next step, the gas stream leaving the adsorption vessel was analyzed to check that
there was no source of contamination in the system, and that the MWCNTs do not contain any
unstable contaminants. The aim of this step was to guarantee that the data represent solely
hydrogen adsorption and that possible adsorption of other gasses or desorption of impurities
from MWCNTs do not cause experimental errors.

24

2.2.3. Effect of the Drift
The gravimetric hydrogen uptake of MWCNTs was measured using a magnetic suspension
balance (MSB), manufactured by Rubotherm, by monitoring the changes in the weight of the
MWCNTs sample due to hydrogen adsorption. Therefore, in order to make definitive
estimates of the change in the sample’s weight during the experiments, it is necessary to
consider the drift of the MSB over time, which is the error in the weight measurement by the
MSB that may happen during the experiments. Since no adsorption or desorption would occur
when there is no sample in the container, measuring over time the weight of the empty
stainless steel sample container is the most accurate way of observing the effect of the drift on
the accuracy of the data.

2.2.4. Stability of the MWCNT Sample
Before starting the experiments, it is important to remove any potential adsorbed impurities
(e.g., water, CO2, and various hydrocarbons) from the MWCNTs surface. This was done in
this study by heating-up the sample at 120
o
C for 5 hr under dynamic vacuum. The sample’s
weight was subsequently recorded under vacuum at 25
o
C. The goal was to ensure that there is
no desorption of possible adsorbed impurities from the MWCNT sample during the hydrogen
adsorption experiments. This step is important because desorption of any impurities from the
sample during the hydrogen adsorption experiments would lead to misleading estimates of the
hydrogen uptake capacity of MWCNTs.

25

2.2.5. Determining the Volume of the MWCNT Sample
Buoyancy forces have a small but non-negligible effect on the sample’s weight,
particularly at the higher pressures. To account for their impact, one must first measure the
sample’s true solid volume (otherwise known as its skeletal volume). For gravimetric
measurements, this is typically accomplished by measuring the sample’s (apparent) weight mp
at various pressures of Helium and by correlating to the sample’s true weight (and also its
apparent weight under vacuum conditions) m0 according to the simple Archimedes formula
𝑚 0
= 𝑚 𝑃 + 𝑉 𝑠 𝜌 𝑔 (2.1)
where 𝑉 𝑠 is the sample’s true (skeletal) volume, and 𝜌 𝑔 is the density of Helium that was used
as a test gas here because it is considered to be inert, non-adsorbing and the lightest among
the noble gases. In the experiments, its bulk gas density 𝜌 𝑔 was measured directly by
weighing a reference stainless steel insert of known volume at various pressures. Plotting
(𝑚 0
- 𝑚 𝑃 ) vs. the gas density 𝜌 𝑔 yields a straight line [Figure 2.1 (c)] with its slope being the
sample’s true volume, 𝑉 𝑠 . Once the skeletal volume is known, and the hydrogen density is
measured at various pressures, the true sample weight during adsorption/desorption is
calculated by adding to the apparent weight the buoyancy correction term. The measured
hydrogen density is 0.007625 g/cm
3
at 100 bar, 0.000801 g/cm
3
at 10 bar and 0.0000813
g/cm
3
at 1 bar (the calculated densities using the Peng Robinson equation of state [23] at 100,
10, and 1 bar are 0.00783g/cm
3
, 0.00081g/cm
3
, and 0.0000813g/cm
3
, respectively) and, thus,
the total buoyancy term corresponds to 0.354 wt.% at 100 bar, 0.0372 wt.% at 10 bar, and
0.00378 wt.% at 1 bar. During the sorption calculations the assumption was made, due to the
small amounts of hydrogen adsorbed (particularly the strongly adsorbed hydrogen that
26

occupies less than 5% of the BET surface area), that no significant changes in the sample’s
skeletal volume take place. Relaxing that assumption and assuming a 5% volume change
would correspond to an error of 0.0177% at 100 bar, 0.00186% at 10 bar, and 0.000189% at 1
bar, and would, in no way, change any of the conclusions about adsorption hysteresis and the
presence of strong and weak adsorption occurring in the MWCNTs.

2.2.6. Hydrogen Adsorption Isotherms
The hydrogen uptake of the MWCNTs sample was determined by weighing the sample at
various hydrogen pressures in the measurement vessel. The effect of the buoyancy forces on
the results was taken into account in the calculations using Eq. (2.1). After reaching the
equilibrium state at each pressure, a back-pressure controller increased the system’s pressure
to the next higher pressure. After hydrogen adsorption was measured at 100 bar, the system’s
pressure was decreased step by step to measure the desorption isotherm, which represents the
first hydrogen adsorption/desorption cycle.
At the end of the first hydrogen sorption cycle, the hydrogen uptake of the MWCNT
sample was measured again, representing the second hydrogen adsorption/desorption cycle.
The difference between the first and the second cycles -- see further discussion below -- was
in the initial treatment of the MWCNT samples that was employed. For the first cycle, the
MWCNTs were first degased at elevated temperature of 120
o
C. For the second cycle the
MWCNTs were not degased at the elevated temperatures and, therefore, any hydrogen still
remaining on the surface from the first cycle was still adsorbed on the MWCNTs at the
beginning of the second cycle.
27

2.3. Results and Discussion
The surface area and pore structure characteristics of the MWCNT sample were measured
using the BET technique. The specific surface area is 441.3 m
2
/g, the average tube diameter is
7.4 nm, while the pore volume is 0.82 cm
3
/g. The x-ray diffraction (XRD) analysis of the
MWCNTs produces two sharp peaks for the carbon layers that are characteristic of the C(002)
and C(100), which are the typical XRD peaks for carbon reported for MWCNTs, but detects
no other crystalline phases [24]. The mass-spectrometric analyzer (RGA200, coupled with an
electron multiplier), was used to analyze the compositions (in the range of 1-64 amu) of the
gas streams entering and leaving the sorption measurement system. The results of the analysis
are reported in Figure 2.1 (a) and they show only the presence of molecular hydrogen (a sharp
peak at 2 amu); there are no other impurities (analytical detection limit of <0.1 ppm v) in the
gas phase that could have resulted in experimental uncertainties.
We note that there was no drift in the measured weight of the empty container, even after
the measurement had been carried out for 3 days (see Figure 2.1(b)). This means that the drift
of the magnetic suspension balance during the hydrogen adsorption experiments is negligible.
Figure 2.1(b) also shows the MWCNT sample’s weight at 25
o
C and under dynamic vacuum
as a function of time (prior to that the sample was degased at 120
o
C for 6 hr to remove water
and other impurities that might have potentially adsorbed during exposure to laboratory
conditions). The sample’s weight remains constant under dynamic vacuum even after 2 days,
indicating the absence of volatile impurities on the surface of the MWCNTs. Coupled to the
findings that no other species are detected in the gas phase (see Figure 2.1(a)), one concludes
that any subsequent weight changes detected with this particular MWCNT sample is only a
28

result of hydrogen adsorption/desorption, and not due to instrument drift or potential
instability of the MWCNTs themselves.
Figure 2.1(c) illustrates the results for the measurement of the sample’s weight at various
helium pressures, corresponding to various gas densities. The linear relationship between the
sample’s weight and the helium gas density confirms the validity of Eq. (2.1) for determining
the MWCNT sample’s true volume. The true density of the MWCNT sample, determined
using Helium gas, was 2.15 g/cm
3
, which is virtually identical to what has been previously
reported (~2.1 g/cm
3
) by others for such materials [25].
Figure 2.2 shows the hydrogen adsorption in the MWCNTs as a function of time at a
pressure of 5 bar. For this experiment, the MWCNTs were exposed to a hydrogen flow for the
first time after the sample was degased in vacuum at 120
o
C for 6 hr; this also represents the
first hydrogen sorption step on the way to generate the adsorption isotherm. Figure 2.2
demonstrates that it takes around 4 - 5 hr for the system to reach equilibrium, which is in line
with the equilibration times reported in the literature [24] for such materials. Careful
examination of Figure 2.2 reveals, furthermore, that a considerable part of the hydrogen
adsorbs almost instantaneously on the MWCNTs, and that ~ 90% of the total amount is
adsorbed within the first hr. This fast “charging” time is an important consideration in terms
of the eventual practical application of these materials.
29

30

Figure 2. 1. (a) Analysis of the gas streams entering and leaving the adsorption vessel using
the residual gas analyzer. (b) Weight of the empty sample container and the fresh sample’s
weight vs. time under dynamic vacuum at pressures of 10
-5
bar. (c) Plot of Eqn. 1 for the
MWCNTs for various pressures of Helium gas.

After the sample’s weight had equilibrated (upon raising the pressure from vacuum to 5
bar), the pressure was raised again in a step-wise manner in order to generate the adsorption
isotherm, which is indicated as the 1
st
hydrogen adsorption cycle in Figure 2.3. Subsequently,
the pressure was lowered in a step-wise manner in order to obtain the desorption branch of the
cycle, which is also shown in Figure 2.3. When the pressure reached back to 5 bar, the weight
of the sample was allowed to equilibrate and, subsequently, the vacuum pump was turned on.
The change in the sample’s weight is also shown in Figure 2.2. Interestingly, even after the
31

weight of the sample levels off, it does not return to its initial weight, and in fact the sample
retains almost 35% of the total hydrogen adsorbed on the CNTs at the end of the first
adsorption isotherm run. Keeping the sample under 10
-5
bar of dynamic vacuum for an
additional two and a half days (the data not shown here) did not change the sample’s weight
any further beyond what is shown in Figure 2.2. The data in Figure 2.3 clearly show that there
is a substantial hysteresis between the adsorption and the desorption branches of the first
cycle. Subsequently, the pressure of the sample was raised from 10
-5
bar to 5 bar and the
adsorption and desorption isotherms were again generated, indicated as the 2
nd

adsorption/desorption cycle in Figure 2.3. Interestingly, no hysteresis was observed between
the adsorption and desorption branches of the 2
nd
adsorption/desorption cycle, which are
completely indistinguishable from each other, as well as being very close to the desorption
branch from the first adsorption/desorption cycle. Upon completion of the second
adsorption/desorption cycle, a third cycle was carried out and, again, no hysteresis was
observed (the results are not shown here because they are very much indistinguishable from
those of the 2
nd
cycle).
It is clear from the data in Figure 2.3 that evacuation of the sample at 25
o
C does not return
its weight to the original value. Since the weight difference involved is rather small (0.25 mg
or ~0.07% of the original MWCNT sample’s weight of ~300 mg – however, as previously
noted, this is a large fraction ~35% of the total amount of hydrogen adsorbed during the first
adsorption cycle), the possibility exists that it may have resulted from dust particles in the gas
atmosphere or in the apparatus chamber being deposited on the MWCNTs sample. To exclude
such a possibility, upon the completion of the 3rd adsorption/desorption cycle the temperature
of the sample was raised (under 10
-5
bar of dynamic vacuum) to 120
o
C. The sample weight
32

started decreasing (see Figure 2.2) and after ~ 6 hr it returned to its original weight. During
the same period the mass analyzer did not indicate the presence of any other gas species (in
the range 1-64 amu) other than hydrogen.

Figure 2. 2. (a) Hydrogen adsorption at the beginning of the 1
st
adsorption/desorption cycle
(step from 0 to 5 bar); (b) Hydrogen desorption at the end of 1
st
hydrogen
adsorption/desorption cycle (from 5 to 0 bar). (c) Desorption of chemisorbed hydrogen under
vacuum at 120
º
C after the end of the 3
rd
hydrogen adsorption/desorption cycle

In our view, the data in Figure 2.3 indicate two different types of hydrogen species on the
surface, weakly adsorbed, or physisorbed, hydrogen and strongly adsorbed, or chemisorbed,
hydrogen, the latter being defined here as the adsorbed species that will not desorb from the
surface under dynamic vacuum at 25
o
C for a period of more than two days (see Figure 2.2)
33

and which are, thus, unlikely to desorb from the same MWCNTs during the normal charging
and discharging cycles as well. In Figure 2.4 we plot the weight change during the 1
st

adsorption/desorption cycle as we transition from 20 bar to 40 bar, as well as the
corresponding weight change profile as we transition back from 40 to 20 bar during the
desorption branch of the cycle. Shown in the same figure are the rates of change in the
sample’s weight, dW/dt. As already noted above, during the adsorption step one observes a
clearly sharp rise in the weight, followed by a more gradual increase. Interestingly, in the rate
of change in the weight profiles one observes two branches with two very distinct slopes. For
the desorption step again most of the weight loss (98%) occurs in the first 7 min. If one is to
assume that the strongly adsorbed species, defined here as those that do not desorb under
dynamic vacuum of 10
-5
bar, are unlikely to desorb either as one transitions from 40 to 20 bar
in pressure, then the desorbed amount can be fully attributed to weakly-adsorbed or
physisorbed hydrogen. Since no hysteresis exists during adsorption/desorption from the part
of the MWCNT surface that is not covered by the strongly adsorbed hydrogen species (see
Figure 2.3), one can then assume that the amount that is desorbed during the desorption
branch of the first cycle is equal to the amount that is physisorbed during the adsorption part
of the cycle. This way one can generate an adsorption isotherm (for the first experimental
cycle) for the weakly-adsorbed hydrogen that is shown in Figure 2.5. By subtracting the
physisorbed isotherm from the total 1
st
cycle adsorption isotherm one can, in addition,
generate the chemisorption isotherm, which is also shown in Figure 2.5. The data suggest that
for pressures of up to 20 bar the physisorption and chemisorption have relatively the same
magnitude. But, above 20 bar physisorption becomes the primary adsorption mechanism as
the strongly-adsorbing sites on the surface of the MWCNT seem to saturate rather quickly.
34

Figure 2. 3. The 1
st
and 2
nd
hydrogen adsorption/desorption cycles (note that for the 2
nd
cycle
the adsorption and desorption branches are indistinguishable).

The assumption that the amount desorbed at each pressure step during the desorption
branch of the 1
st
cycle is only due to weakly-adsorbed hydrogen and by subtracting that
amount from the total adsorption isotherm one can correctly generate the chemisorption
isotherm, entails the fundamental assumption that there are two distinctly different types of
sites on the MWCNT, one on which exclusively weak adsorption occurs, and another on
which preferential strong adsorption takes place. If that is indeed the case, and one subtracts
from the total adsorption isotherm during the 2
nd
(and 3
rd
) cycle the amount that remains
irreversibly adsorbed under dynamic vacuum at 25
o
C, one should be able to generate the
physisorption isotherm of Figure 2.5. To prove that this is indeed the case, all three (the
35

physisorption isotherm of Figure 2.5 and the estimated physisorption isotherms for cycles 2
and 3) are shown in Figure 2.6. The three lines nearly coincide, thus lending credence to the
above hypothesis. This, then, means that hydrogen sorption hysteresis during the first cycle is
a consequence of hydrogen chemisorption on the MWCNTs. The chemisorbed hydrogen
atoms do not leave the MWCNTs when the system’s pressure decreases during the desorption
experiments. Therefore, the accumulation of chemisorbed hydrogen at various pressures
during adsorption gives rise to the observed hysteresis during desorption. For practical
applications, these results validate the idea that only physisorption is involved during the
charging and discharging parts of the cycle.

Figure 2. 4. Hydrogen uptake (w) and absolute value of hydrogen uptake change rate (dw/dt).

36

Figure 2.5. The calculated sorption isotherms in the MWCNTs at 25
◦
C.

Despite the fast rates of physisorption and the apparent strong chemisorption affinities
between hydrogen and the MWCNT surface, the fraction of the BET surface area of the
MWCNT covered by adsorbed hydrogen at the end of adsorption branch is rather small, ~9.2
% for the physisorbed hydrogen and ~4% for the chemisorbed hydrogen (since it is unclear, at
this point, what kind of surface species are involved in these two different types of adsorption,
or whether it even makes good sense to compare the surface area based on liquid nitrogen
adsorption to the area occupied by hydrogen, these estimates are based on the simple idea that
the adsorbed hydrogen molecule occupies an area on the BET surface of MWCNT equivalent
to a circle with a diameter equal to its Lennard-Jones diameter. This is likely to be a good
assumption for the physisorbed hydrogen, but less so for the chemisorbed one, though any
37

differences that may exist are unlikely to be substantial). Such small coverages are in line
with prior modeling investigations [26].

Figure 2. 6. The amount of hydrogen physisorption on the MWCNTs during the first sorption
cycle and the equilibrium hydrogen values during the second and third hydrogen sorption
cycles (the amount of the strongly sorbed hydrogen after the 1
st
cycle is subtracted).

2.4. Conclusions
Precise measurements by a new technique have demonstrated that hydrogen adsorption in
multi-walled carbon nanotubes at room temperature is a combination of reversible
physisorption and irreversible chemisorption. The adsorption measurements unveil and
confirm that it is the chemisorption part that gives rise to hysteresis in hydrogen
38

adsorption/desorption isotherms. For instance, while the chemisorbed hydrogen was still on
the surface of the MWCNTs, the sample was once more exposed to hydrogen and the
adsorption/desorption isotherms were measured again. No hysteresis was detected for the new
cycle, implying that no chemisorption happened during the second cycle. The equilibrium
values of the new hydrogenation cycle are different from those of the first cycle. Our
calculations confirmed that, by subtracting the chemisorbed values from the equilibrium data
of the first cycle, the adsorption and desorption branches of the first hydrogenation cycle and
the equilibrium data of the second cycle take on the same values. This supports the conclusion
that the reversible hydrogen adsorption on MWCNTs during practical conditions is
completely due to physisorption. The technique used in this work is a practical method for
determining separately hydrogen physisorption and chemisorption in carbon nanotubes and
other nanostructured porous materials.

39

2.5. References
[1] Lee SY, Park SJ. Effect of Temperature on Activated Carbon Nanotubes for Hydrogen
Storage Behaviors. Int. J. Hydrogen Energy 2010; 35: 6757-6762.
[2] Zhu H, Cao A, Li X, Xu C, Mao Z, Ruan D, Liang J, Wu D. Hydrogen Adsorption in
Bundles of Well-Aligned Carbon Nanotubes at Room Temperature. Appl. Surf. Sci. 2001;
178: 50-55.
[3] Ci L, Zhu H, Wei B, Xu C, Wu D. Annealing Amorphous Carbon Nanotubes for Their
Application in Hydrogen Storage. Appl. Surf. Si. 2003; 205: 39-43.
[4] Hirscher M, Becher M, Haluska M, Quintel A, Skakalova V, Choi YM, Dettlaff-
Weglikowska U, Roth S, Stepanek I, Bernier P, Leonhardt A, Fink J. Hydrogen Storage in
Carbon Nanostructures. J. Alloys Compd. 2002; 330-332: 654-658.
[5] Panella B, Hirscher M, Roth S. Hydrogen Adsorption in Different Carbon Nanostructures.
Carbon 2005; 43: 2209-2214.
[6] Hou PX, Xu ST, Ying Z, Yang QH, Liu C, Cheng HM. Hydrogen Adsorption/Desorption
Behavior of Multi-Walled Carbon Nanotubes with Different Diameters. Carbon 2003; 41:
2471-2476.
[7] Sudan P, Zuttel A, Mauron P, Emmenegger Ch, Wenger P, Schlapbach L. Physisorption
of Hydrogen in Single-Walled Carbon Nanotubes. Carbon 2003; 41: 2377-83.
[8] DOE Hydrogen and Fuel Cells Program Plan Draft (2011);
http://www1.eere.energy.gov/hydrogenandfuelcells/
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[9] Geng HZ, Kim TH, Lim SC, Jeong HK, Jin MH, Jo YW, Lee YH. Hydrogen Storage in
Microwave-Treated Multi-Walled Carbon Nanotubes. Int. J. Hydrogen Energy 2010; 35:
2073-2082.
[10] Chambers A, Park C, Terry R, Baker K, Rodriguez NM. Hydrogen Storage in Graphite
Nanofibers. J. Phys. Chem. B 1998; 102: 4253-4256
[11] Luxembourg D, Flamanta G, Bˆeche E, Sans JL, Giral J, Goetz V. Hydrogen Storage
Capacity at High Pressure of Raw and Purified Single-wall Carbon Nanotubes Produced with
a Solar Reactor. Int. J. Hydrogen Energy 2007; 32: 1016-1023.
[12] Tibbetts GG, Meisner GP, Olk CH. Hydrogen Storage Capacity of Carbon Nanotubes,
Filaments, and Vapor-Grown Fibers. Carbon 2001; 39: 2291-2301.
[13] Z¨uttel A, N¨utzenadel C, Sudan P, Mauron P, Emmenegger C, Rentsch S, Schlapbach
L, Weidenkaff A, Kiyobayashi T. Hydrogen Sorption by Carbon Nanotubes and other Cabon
Nanostructures. J. Alloys Compd. 2002; 330: 676-682.
[14] Dodziuk H, Dolgonos G. Molecular Modeling Study of Hydrogen Storage in Carbon
Nanotubes. Chem. Phys. Lett. 2002; 356: 79-83.
[15] Tarasov BP, Maehlen JP, Lototsky MV, Muradyan VE, Yartys VA. Hydrogen Sorption
Properties of Arc-Generated Single-Wall Carbon Nanotubes. J. Alloy Compd. 2003; 356:
510-514
[16] Liu C, Fan YY, Liu M, Cong HT, Cheng HM, Dresselhaus MS. Hydrogen Storage in
Single-Walled Carbon Nanotubes at Room Temperature. Science 1999; 286: 1127-1129.
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[17] Bulyarskii SV, Basaev AS. Chemisorption of Hydrogen by Carbon Nanotubes. Technol.
Phys. 2009; 54: 1612-1617.
[18] Orinakova R, Orinak A. Recent Applications of Carbon Nanotubes in Hydrogen
Production and Storage. Fuel 2011; 90: 3123-3140.
[19] Hirscher M, Becher M, Haluska M, Dettlaff-Weglikowska U, Quintel A, Duesberg GS,
Choi YM, Downes P, Hulman M, Roth S, Stepanek I, Bernier P. Hydrogen Storage in
Sonicated Carbon Materials. Appl. Phys. A 2001; 72: 129-132.
[20] Panella B, Hirscher M, Ludescher B. Low-Temperature Thermal-Desorption Mass
Spectroscopy Applied to Investigate the Hydrogen Adsorption on Porous Materials.
Microporous Mesoporous Mater. 2007; 103: 230-234.
[21] Bianco S, Giorcelli M, Musso S, Castellino M, Agresti F, Khandelwal A, Lo Russo S,
Kumar M, Ando Y, Tagliaferro A. Hydrogen Adsorption in Several Types of Carbon
Nanotubes. J. Nanosci. Nanotechnol. 2010; 10: 3860-3866.
[22] Rather S, Naik M, Hwang SW, Kimb AR, Nahm KS. Room Temperature Hydrogen
Uptake of Carbon Nanotubes Promoted by Silver Metal Catalyst. J. Alloys Compd. 2009;
475: 17-21.
[23] Valderrama JO. Interaction Parameter for Hydrogen-Containing Mixtures in the Peng
Robinson Equation of State. Fluid Phase Equilibria 1986; 31: 209-219.
[24] Stamatin I, Morozan, A, Dumitru A, Ciupina V, Prodan G, Niewolski J, Figiel H. The
Synthesis of Multi-Walled Carbon Nanotubes (MWNTs) by Catalytic Pyrolysis of the
Phenol-Formaldehyde Resins. Physica E 2007; 37: 44-48.
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[25] Lehman JH, Terrones M, Mansfield E, Hurst KE, Meunier V. Evaluating the
Characteristics of Multiwall Carbon Nanotubes. Carbon 2011; 49: 2581 –2602.
[26] Ng TY, Rena YX, Liew KM. Adsorption of hydrogen atoms onto the exterior wall of
carbon nanotubes and their thermodynamics properties. Int. J. Hydrogen Energy 2010; 35:
4543–4553

43

Chapter 3
Hydrogen Storage Using Silicon Carbide
Nanotubes
3.
3.1. Introduction
2

Although hydrogen is recognized as a highly efficient energy carrier, several obstacles
have hindered its broad use as a fuel for mobile applications. Chief among such obstacles is
the fabrication and commercial development of an appropriate material to be used as a high-
capacity hydrogen storage medium [1].

Carbon nanotubes (CNTs) are among the materials
that have been considered for this purpose and numerous experimental and theoretical studies
have been carried out in order to investigate hydrogen adsorption in CNTs and to improve
their storage capacities [2]. Nonetheless, such efforts have not, so far, been successful in
reaching the 6 wt.% (weight of hydrogen/weight of storage material) target for hydrogen
storage capacity set forth by the United States Department of Energy for commercial use of
such materials. Thus, significant efforts have been devoted to date to the development of new
hydrogen storage media [3-5].

Computer simulations, for example, have indicated that
introducing point charges on the storage material’s surface can enhance its storage capacity,

2
The material in this Chapter is part of a paper previously published: Barghi SH, Tsotsis TT, Sahimi M. Hydrogen sorption
hysteresis and superior storage capacity of silicon-carbide nanotubes over their carbon counterparts. International Journal of
Hydrogen Energy. 2014;39:21107-15.
44

because it increases the binding energy of the hydrogen molecules due to the dipole
interaction induced by the surface charges [6].
Among the many alternatives to CNTs for hydrogen storage, silicon carbide nanotubes
(SiCNTs) are particularly intriguing. This is due to the resemblance to and similarity with -
being able to store hydrogen through physisorption - but also important differences between
SiCNTs and CNTs, as the former possesses a charge-induced dipole interaction on the SiC
surface that facilitates interaction between hydrogen and the surface. Silicon carbide is, of
course, an important material in its own right with many excellent properties, such as high
fracture toughness, thermal shock resistance, and the ability to withstand high temperatures
and corrosive environments. It finds many important applications in the electronics industry,
and in recent years has also been used for fabrication of a variety of nanoporous materials,
such as SiC membranes for gas separations in harsh environments, where other types of
organic and inorganic membranes typically fail [7-11].

The applications also include various
one-dimensional (1D) nanomaterials, such as nanowires, nanocables, and nanofibers [12, 13],

and quasi-1D structures, such as SiCNTs [14, 17], prepared using a number of methods that
include chemical vapor deposition (CVD) that converts CNTs into SiCNTs, laser ablation,
arc-discharge, various sol-gel routes, and carbothermal reduction [18-23].
Recent atomistic simulations of hydrogen sorption in SiCNTs [6, 24-26],

and flow of
liquids and their mixtures through them have also yielded quite encouraging results for the
potential utility of such nanotubes [27, 28]. According to Mpourmpakis et al. [6],

point
charges on the outer surface of the SiCNTs that result from the continuous charge transfer of
more than half an electron from Si to C are favorable for the adsorption and storage of
45

hydrogen in SiCNTs. Using extensive molecular dynamics (MD) simulations over wide
ranges of temperature and pressure, Malek and Sahimi [26]

reported that SiCNTs have
superior hydrogen storage capacity over their carbon counterparts. Combined with the ability
to reproducibly fabricate the materials, these results suggest that SiCNTs are potentially
promising hydrogen storage media.
Among the various methods of fabrication of SiCNTs, the reaction between a vapor-phase
precursor containing Si and CNTs for converting them into their SiCNTs counterparts appears
to be the most promising method. This is, indeed, the method used in this Chapter for the
preparation of SiCNTs (see below for further details). The advantage of the technique is its
ability for producing SiCNTs with diameters, lengths, and morphology that closely correlate
with the corresponding properties of the original CNTs that are used as the carbon template
[25, 29, 30]. Silicon monoxide (SiO), produced by the high-temperature reaction
Si (s) + SiO2(s)→ 2 SiO (g) (3.1)
is widely used as a vapor-phase silicon precursor to transform CNTs into SiCNTs via the gas-
solid reaction
SiO (g) + 2 C (s)→ SiC (s) + CO (g) (3.2)
It should be noted, that the nature of reaction (3.1), e.g., whether it is indeed a solid-solid
reaction, is a source of controversy in the scientific literature [31]. The outcome of such
discussions has, however, little bearing for the present study. Reactions (3.1) and (3.2) should,
in principle, produce SiCNTs that are faithful replicas of their precursor CNTs. This is why
the method is known as the shape-memory synthesis. In addition to the two main reactions,
however, the following two unfavorable side reactions may also take place [32]:
46

SiO (g) + CO (g) → SiC (s) + CO2 (g) (3.3)
C (s) + CO2(g)→ 2 CO (g) (3.4)
Reaction (3.3) produces solid SiC via a gas-gas reaction that may or may not be deposited on
the SiCNTs growing via the main gas-solid reaction (3.2), whereas reaction (3.4) results in
etching away the CNTs during the synthesis of SiCNTs. The presence of the two side
reactions results in SiCNTs with diameters that may be different from those of the precursor
CNTs, but also in the formation of nanoparticles rather than SiCNTs. Thus, the two reactions
are unfavorable towards efficient fabrication of SiCNTs with predictable dimensions, and
produce SiCNTs with a broader distribution of diameters. In order to prevent the two side
reactions from happening, it is necessary to carry out the fabrication of the SiCNTs under
conditions that substantially minimize the accumulation of CO, which forms by reaction (3.2),
in the reaction environment (see the experimental section below).
Using the above method, Keller et al. [30],

for example, prepared SiCNTs with surface
areas of 30-60 m
2
/g. The same group had previously synthesized SiCNTs using the same
technique [29],

but employing larger size CNTs. Keller et al. [30]

reported that the surface
area and size of the produced SiCNTs depended on the reaction temperature, the Si/SiO 2
starting molar ratio, the (Si+SiO2)/C weight ratio, and the duration of the reaction. Taguchi et
al. [33]

synthesized SiCNTs using silicon powder and CNTs at 1200
o
C under vacuum (a
pressure of ~5 × 10
−4
pa), and for various reaction periods (1-100 hr). They reported the
formation of pure SiCNTs along with CNTs whose surface was coated with a SiC layer. As
expected, longer reaction times led to higher conversions of C into SiC, and the presence of
partially-converted CNTs. Furthermore, smaller diameter CNTs - those whose diameters were
less than ~10 nm - were easier to react completely and to convert to pure SiCNTs.
47

In the present study, we report on the synthesis of SiCNTs using the reaction between SiO
vapor and multiwall CNTs (MWCNTs) in the temperatures range (1200
o
C-1500
o
C), which is
similar to the range of temperatures used previously by other investigators for the successful
fabrication of SiCNTs [32, 34-36].

The as-synthesized SiCNTs were then purified by
dispersing them in a hot, concentrated sodium hydroxide (NaOH) solution for 24 hr.
Hydrogen adsorption and desorption isotherms on the as-synthesized and in the purified
SiCNTs were then measured in order to test their capacity as hydrogen storage media.

3.2. Experimental
We first describe fabrication of SiCNTs using MWCNTs, and then explain the procedure
for their purification, after which measurement of hydrogen sorption in the SiCNTs will be
described.

3.2.1. Synthesis of SiCNTs
We used MWCNTs (5-8 nm in diameter), produced via methane catalytic decomposition
over a Co-based catalyst, with a reported purity of 97.46 wt.%, purchased from Nanostructure
and Amorphous Materials, Inc. According to the manufacturer, the key impurities are Al (0.19
wt.%), Cl (1.02 wt.%), Co (1.09 wt.%), and S (0.24 wt.%). The Si powder, with mesh size
<325 and a purity of better than 99% (containing 0.31% Al, 0.5% Fe, and 0.043% Al) was
purchased from Stem Chemical, Inc. Fused amorphous SiO2 with mesh size <325 and purity
of 99.8% (containing 0.2% Al2O3 and 0.01% carbon) was supplied by Alfa Aesar Chemical,
48

Inc. NaOH with purity of 98.5% was provided by Mallinckrodt (impurities: 0.002% Pb,
0.00001% Hg, 0.0003% As, and 0.4% Na2CO3).
In order to fabricate the SiCNTs, about 1.5 g of the (Si+SiO2) mixture was placed at the
bottom of an alumina crucible. The mixture was then covered with a porous binder-free
alumina paper (purchased from Zircar Ceramics made of 86% Al2O3, 10% SiO2, and 4% of
other oxides) that acts as a barrier. In the next step, approximately 0.15 g of CNTs was placed
on top of the alumina paper. The alumina paper acts as the barrier and provides a distinct
advantage over what has been used by other groups. In the past, either a carbon felt had been
used as the barrier, or the CNTs were directly placed on top of the (Si+SiO2) powder mixture,
without the use of any barrier [29, 30, 35].

The disadvantage of using the carbon felt (as
opposed to using the inert alumina barrier, as is done in the present study) is its possible
reaction with the SiO vapor, via reaction (3.2), resulting in the loss of SiO, but more
importantly, releasing CO, which, as discussed previously, reacts with the SiO vapor via
reaction (3.3) and alters potentially the size of the SiCNTs. The drawback of not using a solid
barrier at all is the likelihood of solid-solid reactions at the contact points between the CNTs
and the (Si+SiO2) powder mixture, and the potential contamination of SiCNTs with the
unreacted Si and/or SiO2.
To prepare SiCNTs, the aforementioned alumina crucible was placed in the middle of a
tube furnace. Flowing argon gas with a flow rate of 20 scc/min was used in order to remove
the CO, produced by reaction (3.2), for the reasons noted previously. For the synthesis
reactions of SiCNTs to begin, the temperature was raised to the desired level at a rate of 4
o
C/min. Four different synthesis temperatures (1200
o
C, 1300
o
C, 1400
o
C, and 1500
o
C) were
49

investigated. The reaction duration for SiCNTs produced at 1200
o
C was 6 hr, but only 2 hr at
1300
o
C, 1400
o
C, and 1500
o
C. Upon completion of the SiCNTs synthesis, the temperature
was lowered to room temperature in flowing argon at a rate of 8
o
C/min. The schematic of the
experimental lay-out used in the experiments is shown in Figure 3.1.
In addition to the aforementioned experimental procedure, we also synthesized SiCNTs by
loading the CNTs and the (Si+SiO2) mixture into two separate crucibles and placing them
next to each other in the tube furnace. However, the observed conversion of C into SiC using
this approach was rather low and non-uniform. Higher conversion was obtained for the CNTs
for the top layers in the crucible. Thus, we abandoned this experimental approach in favor of
the configuration shown schematically in Figure 3.1, and only SiCNTs fabricated by the
former technique are described in this study.

Figure 3. 1. The experimental layout for synthesis of SiCNTs.

The crystallinity of the structure of the synthesized SiCNTs was studied via X-ray
diffraction (XRD) analysis (with the aid of a Rigaku XRD equipment). Fourier transform
infrared spectroscopy (FTIR, model: Nicolet 5700) was used to investigate the chemical
50

bonds of SiCNTs. The composition of SiCNTs was analyzed by an inductively coupled
plasma mass spectrometer (ICP-MS, ELANN 9000 manufactured by Parkin Elmer). In
addition, scanning electron microscopy (SEM, using a JEOL JSM-7001) and transmission
electron microscopy (TEM, using JEOL JEM-2100F) were used to investigate the
morphology of the produced SiCNTs. To prepare SiCNTs for TEM imaging, they were
dispersed in ethanol. A drop of the dispersion was put on a TEM grid and a UV-cleaner was
then used to clean the grid for imaging.

3.2.2. Purification of the SiCNTs
In order to remove the potential impurities from the as-synthesized SiCNTs, they were
dispersed in a hot and concentrated sodium hydroxide solution. For this purpose, a 100 ml
NaOH solution was first prepared by dissolving around 35 g NaOH pellets in distilled water.
Then, around 600 mg of the as-synthesized SiCNTs were added to the NaOH solution and
sonicated for 10 min. A magnet-stirrer was then used to keep the temperature of the
dispersion at 90
o
C for 1 day. Finally, the purified SiCNTs were filtered from the NaOH
solution, thoroughly washed with distilled water, and dried at 60
o
C for 12 hr.

3.2.3. Hydrogen Sorption Measurements
Hydrogen excess adsorption isotherms in SiCNTs were measured using a magnetic
suspension balance (MSB), manufactured by Rubotherm. The MSB measures hydrogen
uptake based on a gravimetric technique. One may also use a volumetric method, but as
51

discussed previously [37], since in the volumetric method the change in the pressure of the
adsorption vessel is the main mechanism via which the adsorption/desorption amounts are
measured, omnipresent leaks during the experiments can (and often do) lead to erroneous
results (especially for H2, since its adsorption is low). Therefore, the hydrogen uptake
capacities of carbon nanomaterials published in the literature (mostly obtained by volumetric
methods) are remarkably scattered. One report, for example, stated that CNTs are capable of
adsorbing 0.1 wt.% hydrogen at 293 K and 10 MPa, whereas others claimed that tubular
graphite nanofibers adsorb 11.26 wt.% hydrogen at 298 K and 11.35 MPa. Still, other
experimental data, which are in variance with the two aforementioned studies, have been
reported as well. Such inconsistent results have been a matter of controversy in the area of
hydrogen storage using carbon nanomaterials. Thus, we have been using the gravimetric
method in our studies.
In this study, the uptake capacity of SiCNTs was determined by monitoring the
equilibrium weight of the sample at various hydrogen pressures in the measurement vessel. A
residual gas analyzer (RGA200, manufactured by Stanford Research Systems) was coupled to
the aforementioned gravimetric instrument. The RGA measures the gas phase composition of
both the feed gas into and the exit gas from the measurement vessel. This is useful in
determining the presence of potential impurities in the adsorptive gas, or any desorption of
volatile impurities from the solid sample loaded in the gravimetric equipment chamber, and
distinguishes the experimental method used in the present work from the gravimetric method
used by itself alone and from the widely used volumetric methods. Additional details
regarding the experimental set-up and procedure used in this work are described elsewhere
[37] and in Chapter 2.
52

The MSB is also capable of measuring the true (skeletal) density of the SiCNTs. This is
accomplished by weighing the SiCNTs sample at various pressures in the presence of a non-
adsorptive gas (Helium in this study). The Archimedes formula,
mAP = mS -VSρ
g
(3.5)
is then used to determine the true (skeletal) volume of the SiCNTs, where m AP is the apparent
weight of the SiCNTs at the selected pressure of Helium (this is the weight that the balance
records), mS is the sample’s true weight (the weight at zero pressure), VS is the sample’s true
(skeletal volume), and ρ
g
is the density of Helium at a given pressure. Experimentally, mS and
VS are calculated by plotting mAP vs. ρ
g
and determining the slope and intercept of the
experimental line.

3.3. Results and Discussion
Figure 3. 2 shows the XRD spectra of the as-synthesized SiCNTs, the purified SiCNTs,
and also of the precursor CNTs. They indicate the presence of pure 𝛽 -SiC with a perfect
crystalline structure in the (1 1 1), (2 0 0), (2 2 0), (3 1 1), and (2 2 2) lattice planes with no
carbon or Si diffractions line being present. Others [29, 30, 33]

were not able to make such
carbon-free SiCNTs, and had to remove the carbon impurities on the surface of the SiCNTs
with the aid of a calcination process at 600 ℃, whereas our synthesis procedure needs no
further calcination. This is an important result, because calcination can be damaging to the
structure of the final SiCNTs, and may also cause the surface oxidation of the SiC [29, 30].
53

Figure 3. 2. The XRD spectra of CNTs, the as-synthesized SiCNTs produced at various
temperatures in the range 1200-1500
o
C, and the purified SiCNTs.

Figure 3. 3. FTIR results for the as-synthesized and purified SiCNTs at 1200
o
C.
54

The FTIR spectra for the as-synthesized SiCNTs are shown in Figure 3.3 and reveal the
Si–O stretching, Si–O–Si bending, and the transversal optic (TO) Si–C vibrations mode.
Therefore, the as-synthesized SiCNTs are likely to contain XRD-amorphous SiO2, consistent
with the results reported [38] by Baosheng et al. The solidification of SiO2 vapor that is
produced from the vaporization of SiO2 during the production of SiCNTs, which does not
participate in solid-solid reaction (3.1), is the likely source of the SiO2 impurity. The FTIR
data for the SiCNTs purified by NaOH are also shown in Figure 3.3 and show the absence of
the Si–O stretching and the Si–O–Si bending vibrations, indicating that the treatment with the
concentrated NaOH solution is effective for removing the amorphous SiO2 impurity and, thus,
for increasing the purity of the SiCNTs. This was expected, as at elevated temperatures,
NaOH reacts with SiO2 according to the following reaction
SiO2 (S) + 2 NAOH (aq)→ Na2SiO3 (aq) + H2O (l) (3.6)
ICP-MS analysis of the SiCNTs purified by NaOH indicated that the nanotubes contained
0.13 ppm wt sodium, implying that very little sodium residue remains on the treated SiCNTs
and, thus, it is unlikely to impact their hydrogen storage behavior.
Figure 3.4A shows SEM images of the as-synthesized SiCNTs at 1200
o
C. The image
clearly indicates that under the chosen preparation conditions we have, indeed, managed to
prepare SiCNTs with a well-defined tubular structure (although some of the tubes have
closed-ends). This is an important finding, as in prior studies it had been reported that failing
to adjust the ratio of the rates of the SiO vapor generation and of the gas-solid reaction
between the SiO and the CNTs leads to the destruction of the tubular structure of CNTs, in
which case the final product is SiC nanoparticles, rather than SiCNTs [34].

55

56

Figure 3. 4.(A) SEM image of the as-synthesized SiCNTs at 1200
o
C; (B) purified SiCNTs;
(C) TEM images of the as-synthesized SiCNTs at 1200
o
C, and (D) purified SiCNTs.
57

The image SEM of the purified SiC depicted in Figure 3.4B shows that, (i) the tubular
structure of the as-synthesized SiCNTs was preserved during the purification process, and (ii)
the purified SiCNTs seem to have open-ends. Figure 3.4C displays the TEM image of the as-
synthesized SiCNTs. It is evident from Figure 3.4C, but also upon closer inspection from
Figure 3.4A as well, that the end of the nanotube is blocked by a spherical shaped end-cap.
The TEM images of the purified SiCNTs, shown in Figure 3.4D, indicate, on the other hand,
that the purified SiCNTs no longer have end-caps (which is consistent with their SEM image
in Figure 3.4B). This would indicate that the end-caps are made of amorphous SiO2 that
deposits during the SiCNTs synthesis reaction. In addition, a comparison between the
diameter of the as-synthesized SiCNTs (Figure 3.4C) and the purified SiCNTs (Figure 3.4D)
reveals that the latter have a smaller diameter that is almost identical with that of the CNT
precursors, while the diameter of the former is substantially greater. These results demonstrate
the effectiveness of the NaOH-based purification process for removing the amorphous SiO2
layer from the surface of the as-synthesized SiCNTs. They also improve the hydrogen storage
characteristics of the nanotubes, as will be discussed further below, potentially by allowing
hydrogen to have access to the inner structure of the SiCNTs.
In Figure 3.5 we present the hydrogen adsorption isotherm at 25
o
C for the initial
MWCNTs, the as-synthesized SiCNTs (produced at various temperatures between 1200
o
C-
1500
o
C), and the purified SiCNTs (synthesized at 1200
o
C and at 1500
o
C). Figure 3.5
indicates that among the as-synthesized SiCNTs, those that were prepared at lower
temperatures have higher hydrogen adsorption capacities, with the one produced at the lowest
temperature (1200
o
C) possessing the highest hydrogen uptake capacity. Note that
58

temperatures lower than 1200
o
C are not effective to fabricate SiCNTs because of the slower
reaction rates lead to significantly diminished conversions of CNTs into SiCNTs.

Figure 3. 5. Hydrogen adsorption isotherm at 25
o
C for the precursor CNTs, the as-
synthesized SiCNTs produced at 1200-1500
o
C, and the purified SiCNTs prepared at 1200
o
C
and at 1500
o
C.

The differences in adsorption capacities of the as-synthesized nanotubes can be either
attributed to differences in the adsorption properties of the SiCNTs themselves, or to
differences in the SiO2 content of the as-synthesized SiCNTs produced at various
temperatures. Figure 3.5 indicates that the presence of the SiO2 impurity has, indeed, a major
59

impact on the hydrogen storage capacity of the SiCNTs: the nanotubes after the NaOH
treatment that removes the SiO2 impurity have a substantially higher hydrogen storage
capacity than the as-synthesized SiCNTs at the corresponding temperatures. On the other
hand, as Figure 3.5 indicates, there is little difference in the hydrogen storage between the two
purified SiCNTs, one prepared at 1200
o
C and the second one at 1500
o
C. This is to be
expected, as the SiCNTs consist of pure 𝛽 -SiC (Figure 3.2) and their size/shape is very
similar to that of the precursor CNTs (Figure 3.4). That the pure SiCNTs have a higher
adsorption capacity than the precursor CNTs, is consistent with the MD simulation results for
hydrogen sorption in SiCNTs [26].

Figure 3. 6.Hydrogen adsorption/desorption cycle at 25
o
C for the as-synthesized SiCNTs
produced at 1200
o
C, the purified SiCNTs, and the initial CNTs.
60

That the SiO2 present on the surface of the as-synthesized SiCNTs would have such a
major impact on their hydrogen storage properties is not totally unexpected. As the SEM
images of Figure 3.4 would tend to indicate. The SiO2 layers would act as a barrier between
hydrogen and the SiC surface during the hydrogen uptake experiments, and by plugging the
ends of the nanotubes (Figure 3.4C) they would prevent hydrogen from reaching the inner
space of the nanotubes. Removing SiO2 from the surface of the nanotubes (via the NaOH
treatment) would allow the hydrogen molecules to freely access the SiC surface and adsorb.
The differences in the adsorption capacities of the as-synthesized SiCNTs prepared at various
temperatures can be attributed to the varying amounts of SiO2 impurity present. The SiCNTs
prepared at the lower temperatures have lower amounts of SiO2 present as verified
experimentally by FTIR and skeletal density measurements (see Appendix A), as it is to be
expected as the vapor pressure of SiO2 is higher at the higher temperatures.
Figure 3.6 compares the hydrogen adsorption/desorption isotherms at 25
o
C for the
precursor CNTs, the as-synthesized SiCNTs produced at 1200
o
C, and the purified SiCNTs.
The hysteresis indicated by Figure 3.6 is quite an interesting phenomenon, as it takes place
under less than monolayer coverage conditions. For the purified SiCNTs at a pressure of 100
bar, the hydrogen adsorbed corresponds to ~130 % of monolayer coverage based on the
surface area measured by N2 adsorption. However, the precursor MWCNT have interlayer
spacing of 3.5 Å at room temperature [36],

and assuming that the resulting SiCNTs have
structures that faithfully replicate the MWCNT structures, it is likely that H2 with a kinetic
diameter of 2.9 Å [39] penetrates the SiCNTs interlayer space, whereas N2 used in the
conventional BET technique has a kinetic diameter of 3.64 Å [39] and, thus, is unlikely to
penetrate the 3.5 Å interlayer spacing of SiCNTs, implying that BET measurements are likely
61

to underestimate the area available for H2 sorption. Diffusion of hydrogen in MWCNTs has
also been reported by Hou et al. [40] They studied experimentally hydrogen
adsorption/desorption in MWCNTs of various diameters at room temperature, and concluded
that hydrogen does adsorb in the interlayer spacing of the nanotubes. Moreover, they
proposed that the curvature of the nanotubes is a key factor that determines the extent of the
hydrogen mobility. This is consistent with the results reported by Chen et al. [41] who used
the density functional theory (DFT) to investigate hydrogen spillover in various carbon
adsorbents, and demonstrated that the curvature of carbon materials has a pronounced
influence on the hydrogen mobility. Similar hysteresis phenomena for hydrogen adsorption
under supercritical conditions have been previously reported for a variety of microporous
carbons [40-48].
Sorption hysteresis in porous carbon materials has been associated with the polarization of
the carbon surface,

spillover effects,

formation of chemical bounding (chemisorption),
penetration of hydrogen between the adjacent graphene sheets in MWNTs, and the presence
of metal constituents [40-48]. The hysteresis has also been observed for metal-organic
framework (MOF) materials [49, 50],

for which hysteresis has been attributed to changes in
the pore structure induced by the presence of the adsorbed hydrogen.

ICP-MS analysis of the as-synthesized and also purified SiCNTs indicated the presence of
the following metal impurities: Co (~10 ppm wt), Al (~1.2 ppm wt), Na (~ 0.13 ppm wt), and Fe
(~0.07 ppm wt), most likely resulting from the impurities found in the precursor MWCNT [50]

and Si and SiO2 (albeit at much lower levels than what was reported by the MWCNT
manufacturer), and may potentially be contributing to the observed hysteresis. The potential
62

presence of microporosity in the interlayer regions, and the changes that hydrogen adsorption
may bring upon it could also be a contributing factor as well. It should be noted, however, that
as we previously reported [37], the original MWCNTs exhibit no similar closed-loop
hysteresis, despite the fact that they contain similar type of metal impurities and probably
have similar type of interlayer microporosity. Moreover, the observed hysteresis is not an
experimental artifact due to insufficient times for adsorption and desorption, and is
reproducible for subsequent cycles of adsorption and desorption beyond the first cycle shown
in Figure 3.6.
Note that, as Figure 3.6 indicates, the adsorption/desorption isotherms for the SiCNTs
merge as the pressure approaches zero (vacuum), which is not the case for the CNTs, as
previously reported [37]. This is an important potential advantage of the SiCNTs over the
CNTs, because SiCNTs, (i) can release all the adsorbed hydrogen at the end of the
hydrogenation cycle and, thus, (ii) can be used for hydrogen storage for more than one cycle
without diminishing their uptake capacities. It also indicates a more energetically uniform
surface towards hydrogen sorption than the precursor CNTs.
Note also that, as discussed previously, in our view, the data for sorption in CNTs shown
in Figure 3.6 indicate two different types of hydrogen species on the surface, weakly
adsorbed, or physisorbed, hydrogen and strongly adsorbed, or chemisorbed, hydrogen, the
latter being defined here as the adsorbed species that will not desorb from the surface under
dynamic vacuum at 25
o
C for a period of more than two days and, thus, are unlikely to desorb
from the same MWCNTs during the normal charging and discharging cycles as well.

63

Figure 3. 7.The rate of hydrogen uptake in SiCNTs and CNTs at 25
o
C.

Figure 3.7 compares the rate of hydrogen adsorption in the CNTs and SiCNTs as a
function of time, when the pressure is increased from vacuum to 10 bar during the first
adsorption cycle. The relative error of the instrument and, thus, that of the data shown in
Figure 3.7, is 5 percent of the values shown. In terms of the reproducibility of the results, we
measured the hydrogen uptake rate at 5 pressures in 4 different samples for a total of 20 data
points and observed a similar trend. In order to make a better quantitative comparison, the
hydrogen uptakes at every moment were divided by the equilibrium value of hydrogen
adsorption at the same pressure, shown in Figure 3.6. As Figure 3.7 indicates, initially,
hydrogen adsorption in SiCNTs occurs much faster than in CNTs. The dipole interaction
induced by the SiC surface charges is presumably the major cause of this phenomenon,
64

because it intensifies the SiC/H2 interaction [6]. The faster sorption kinetics of SiCNTs is
another important improvement over the CNTs, as it decreases the “charging time” for
hydrogen storage systems.

3.4. Conclusions
SiCNTs were synthesized by a chemical-vapor deposition reaction in the temperature
range 1200-1500
o
C. The as-synthesized SiCNTs were subsequently purified by treatment
with a NaOH solution at elevated temperatures. The hydrogen adsorption in the SiCNTs at
room temperature and pressures between 10
-5
to 100 bar was studied using a gravimetric
instrument, coupled with a mass analyzer. The data indicate that the hydrogen uptake capacity
of the as-synthesized SiCNTs is comparable with that of their CNTs precursors. On the other
hand, hydrogen storage capacity of the SiCNTs purified by NaOH is more than 50% higher
than that of the CNTs precursors. The increase in hydrogen storage capacity is due to the
removal of amorphous SiO2 from the surface and from the pore volume of the as-synthesized
nanotubes which results in an increase in the number of available hydrogen storage sites and
provides more ready access to them for the hydrogen molecules.
The uptake data for the SiCNTs reveal the existence of an adsorption/desorption hysteresis
phenomenon. In contrast to the precursor CNTs, we do not observe any strongly adsorbed
hydrogen at the conclusion of the desorption cycle. The absence of strongly adsorbed
hydrogen at the end of the hydrogenation cycle makes SiCNTs a better storage material, as
they can be used for more than one hydrogenation cycle without any deterioration of their
hydrogen uptake capacity. In addition to the complete reversibility of hydrogen
65

adsorption/desorption, another superior characteristic of the SiCNTs as an adsorption medium
is their faster hydrogen adsorption rate, when compared with that of CNTs. In the experiments
reported here, for example, hydrogen adsorption equilibrium in SiCNTs is reached typically
in less than 1 hr, whereas the corresponding time for CNTs is about 5 hr. The faster uptake
rate is important, because it decreases the charging time necessary to store hydrogen in such
nanotubes.
Despite such superior characteristics of SiCNTs for hydrogen storage, their low adsorption
capacity, which is less than 0.3 wt%, still hinders their commercial application as storage
materials for hydrogen. Doping the surface of the nanotubes with alkali metals and other
materials is the next plausible step for producing solid adsorbents with high hydrogen storage
capacities. High thermal and chemical stability of SiCNTs makes them a promising candidate
for the doping modification, because they can readily resist the harsh oxidative and reductive
environments associated with the doping procedures.

66

Appendix A: Calculation of the SiO2 Content the SiCNTs
The true density of the synthesized nanotubes, 𝜌 𝑠 , is related to the mass and volume of its
pure SiC content and the SiO2 impurity according to Eq. 3.7:
𝜌 𝑠 =
𝑚 𝑠 𝑉 𝑠 =
𝑚 𝑆𝑖𝐶 +𝑚 𝑆𝑖𝑂 2
𝑉 𝑆𝑖𝐶 +𝑉 𝑆𝑖𝑂 2
(3.7)
where 𝑚 𝑠 , 𝑚 𝑆𝑖𝐶 , and 𝑚 𝑆𝑖𝑂 2
are the masses of the nanotubes and the SiC and SiO2
components, respectively, and 𝑉 𝑠 , 𝑉 𝑆𝑖𝐶 , and 𝑉 𝑆𝑖𝑂 2
are their corresponding volumes. Rewritten
in terms of the silicon carbide true density, 𝜌 𝑆𝑖𝐶 , and SiO2 true density, 𝜌 𝑆𝑖𝑂 2
, Eq. 3.7 takes
the form:
𝜌 𝑠 =
𝑚 𝑆𝑖𝐶 +𝑚 𝑆𝑖𝑂 2
𝑚 𝑆𝑖𝐶 𝜌 𝑆𝑖𝐶 +
𝑚 𝑆𝑖𝑂 2
𝜌 𝑆𝑖𝑂 2
(3.8)
Assuming the mass fraction of the SiO2 impurity to be x, we can write:
𝑥 =
𝑚 𝑆𝑖𝑂 2
𝑚 𝑆 (3.9)
Therefore,
𝜌 𝑠 =
(1−𝑥 ) 𝑚 𝑆 +𝑥 𝑚 𝑆 (1−𝑥 ) 𝑚 𝑆 𝜌 𝑆𝑖𝐶 +
𝑥 𝑚 𝑆 𝜌 𝑆𝑖𝑂 2
(3.10)
Rearranging Eq. 3.10 would give the SiO2 impurity mass fraction as:
𝑥 =
1
𝜌 𝑆 +
1
𝜌 𝑆𝑖𝐶 1
𝜌 𝑆𝑖𝑂 2
+
1
𝜌 𝑆𝑖𝐶 (3.11)
67

Therefore, Eq. 3.11 can be used to calculate the SiO2 contents of the produced SiCNTs at
different temperature based on the true density of the nanotubes. According to the literature,
the true density of SiC and amorphous SiO2 are 3.21 g/cm
3
and 2.25 g/cm
3
, respectively [51].
The calculated SiO2 content of the SiCNTs based on their true density are summarized in
Table 3.1. It can be noticed in Table 3.1 that the SiO2 content of the as-synthesized SiCNTs is
higher at higher synthesis temperatures. In addition, the true densities of the purified SiCNTs
are very close to that of the pure SiC, with a slight difference mostly due to the experimental
errors and possibly the presence of some impurities.

Table 3. 1. The true (skeleton) density of the as-synthesized and purified SiCNTs and the
corresponding calculated SiO2 content based on the true density values.
SiCNTs True density (g/cm
3
) Calculated SiO 2 fraction (%)
As-synthesized, 1200
o
C 3.09 9.1
As-synthesized, 1300
o
C 3.07 10.7
As-synthesized, 1400
o
C 3.02 14.7
As-synthesized, 1500
o
C 2.84 30.5
Purified, 1200
o
C 3.19 1.5
Purified, 1500
o
C 3.18 2.2
68

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doped 2D graphene. Int. J. Hydrogen Energy 2011;36:12950-4.
[48] Saha D, Deng S. Hydrogen adsorption on ordered mesoporous carbons doped with Pd,
Pt, Ni, and Ru. Langmuir 2009;25:12550-60.
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[50] Guo Z, Li G, Zhou L, Su S, Lei Y, Dang S, Zhang H. Magnesium-based 3d metal-
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2006.

74

Chapter 4
Hydrogen Storage Using Doped Silicon
Carbide Nanotubes
4. References
4.1. Introduction
3

Over the past decade or so extensive efforts have been devoted to the development of
nanoscale carbonaceous materials for hydrogen storage applications [1-16]. Despite the
promising progress made in using nanostructured materials for this purpose, the availability of
efficient hydrogen storage materials still remains a key barrier to the commercial application
of hydrogen as a fuel. Use of silicon-carbide nanotubes (SiCNTs), instead of carbon
nanotubes (CNTs), is a potential solution to this challenge [17-21]. In our previous work [22,
23], we showed that converting CNTs into SiCNTs does increase their hydrogen storage
capacity by more than 50%; however, the hydrogen uptake of the resulting SiCNTs was still
far below the target set forth by the United States Department of Energy (DOE) for vehicular
applications, which is 6-7 wt.%.

3
The material in this Chapter is part of a paper previously submitted to Int. J. Hydrogen Energy: Barghi SH, Tsotsis TT,
Sahimi M. Experimental investigation of hydrogen adsorption in doped silicon-carbide nanotubes.

75

In order to investigate methods for increasing the hydrogen storage capacity of nanoscale
materials, in addition to experimental efforts, complimentary theoretical and computational
studies have also been undertaken over the past few years. Computer simulation studies of
doped nanomaterials, for example, have indicated considerable improvement in their
hydrogen storage capacity brought forth by doping [24-29]. Dag et al. [25], for example, used
first-principles calculations based on the plane-waves approach to investigate the interaction
between hydrogen and pure and doped CNTs. According to their study, doping of the CNTs
with transition metals, such as titanium (Ti), and alkali metals, such as lithium (Li) and
potassium (K), can increase significantly the capacity of the nanotubes for hydrogen sorption
and storage. In another first-principles study Yildirim and Ciraci [26] showed that a single Ti
atom coated on the outer surface of the CNTs can bind with up to four hydrogen molecules.
The adsorption of the first hydrogen molecule was shown to be dissociative, while the next
three would adsorb in their molecular state. In addition to serving as dopants for the CNTs,
transition and alkali metals have also been recommended as efficient dopants for the SiCNTs
in order to improve their hydrogen storage capacity [18, 27-29]. Wang and Liew [28], for
example, studied alkali-doped SiCNTs with the aid of ab initio density-functional theory.
They showed that doping the SiCNTs would increase their hydrogen binding energy from
0.086 eV to 0.211 eV due to the charge transfer from the dopant to the SiCNTs. Four
hydrogen molecules with an average binding energy of 0.165 eV can consequently be
adsorbed on the doping sites and thus increase the hydrogen uptake capacity of the SiCNTs.
In addition, simulations based on pseudopotential methods [29] have indicated that Ti-
decorated SiCNTs exhibit superior hydrogen storage capacity. According to Banerjee et al.
[29], a Ti atom doped on the SiCNTs’ surface is capable of adsorbing two hydrogen atoms
76

and three hydrogen molecules, which is similar to the observed behavior of the Ti-doped
CNTs [26].
In the present paper, we report on an experimental study of the effect on hydrogen storage
capacity of the SiCNTs doped with K (as an alkali metal) and Ti (as a transition metal). The
goal of the work is to determine whether doped SiCNTs have, indeed, enhanced adsorption
capacities as the aforementioned molecular simulations seem to indicate. If they do, we seek
to understand whether these hydrogen storage capacities approach the target set forth by the
United States Department of Energy (DOE) for such hydrogen storage media.
This paper is organized as follows: In the next Section, we describe the experimental
procedure for doping the SiCNTs, and for measuring the sorption of hydrogen in them. The
results are then presented and discussed in Section 4.3.

4.2. Experimental
4.2.1. Materials
Ultra-high pure hydrogen with 99.999% purity was purchased from the Gilmore Liquid Air
Company. Potassium hydroxide (KOH) with purity better than 99.0% (reported to contain
≤0.01% Cl, ≤0.005% SO4, ≤0.05% (Fe(CN)6)
4-
, and ≤0.02 Pb), was supplied by EMD
Chemicals Inc. Sodium hydroxide (NaOH) with purity of 98.5% (impurities of 0.002% Pb,
0.00001% Hg, 0.0003% As, and 0.4% Na2CO3) was provided by Mallinckrodt. The Si
powder, with mesh size smaller than 325 and a purity better than 99% (containing 0.31% Al,
0.5% Fe, and 0.043% Al) was purchased from Stem Chemical, Inc. Fused amorphous SiO2
77

with mesh size smaller than 325 and purity of 99.8% (containing 0.2% Al2O3 and 0.01%
carbon) was supplied by the Alfa Aesar Chemical, Inc. Titanium (IV) chloride (TiCl4) with a
purity better than 99.0% was provided by Sigma Aldrich. Multiwall CNTs (MWCNTs) with
diameters of 5-8 nm, produced via methane catalytic decomposition over a Co-based catalyst,
with a reported purity of 97.46 wt%, were purchased from Nanostructure and Amorphous
Materials, Inc. According to their manufacturer, the key impurities of the MWCNTs were Al
(0.19 wt.%), Cl (1.02 wt.%), Co (1.09 wt.%), and S (0.24 wt.%). Scanning electron
microscopy (SEM) and energy-dispersive X-ray spectroscopy (EDS), using a JEOL JSM-
7001 instrument were used (see below) to characterize the nanomaterial samples.

4.2.2. SiCNTs Synthesis and Purification
The SiCNTs used in this work were produced via the reaction between SiO vapor and
CNTs at 1200
o
C under an Ar atmosphere. A hot and concentrated NaOH solution was then
used to remove the impurities, as well as any potential side products of the synthesis reaction.
The purification was shown to be efficient in opening the SiCNTs’ ends [23]. Complete
details and extensive discussions of the SiCNTs synthesis and purification are given in our
previous publication [23] and in Chapter 3 and, therefore, will not be repeated here.

4.2.3. Doping SiCNTs with K
Potassium doping of the SiCNTs was carried out by dispersing about 250 mg of purified
SiCNTs in a 3M KOH solution and sonicating the mixture for 2 hr. A centrifuge was then
78

used to separate the SiCNTs from the solution. In the next step, the SiCNTs were first dried at
75
o
C in air, and subsequently were heated up to 800
o
C for 2 hr under an Ar atmosphere.
Heating the KOH-impregnated SiCNTs in Ar at 800
o
C was carried out in order to produce
K2O through KOH dehydration [30] according to the following reaction
2KOH → K2O + H2O (4.1)
and, subsequently, to obtain atomic potassium as a result of the carbothermal K2O reduction
[31, 32]
K2O + C → 2K + CO (4.2)
When reaction (4.2) takes place during K-doping, some of the carbon atoms on the SiCNTs
surface are replaced with potassium atoms.

4.2.4. Doping SiCNTs with Ti
Titanium tetrachloride (TiCl4) was used as the Ti-doping agent. For this purpose, 1 ml of
TiCl4 was added drop-wise to 20 ml of ethanol. During the mixing process, a considerable
quantity of HCl gas was released that is, likely, the result of the following reaction [33]
TiCl4 + C2H5OH → TiClx(OC2H5)4-x + (4-x) HCl (4.3)
Approximately 500 mg of SiCNTs were added to the resulting solution, and were
subsequently sonicated for 5 min. The mixture was then stirred for 24 hr before a centrifuge
was used to separate the SiCNTs from the solution. The SiCNTs were then dried at 80
o
C
overnight under an air atmosphere. The product of the aforementioned procedure turned out to
79

be rigid clumps of nanotubes, most likely due to the polymerization of the Ti precursor [33].
Since the goal of the Ti doping of the SiCNTs is to enhance their hydrogen storage capacity, it
is more desirable to separate the SiCNT clumps into individual nanotubes. To accomplish
this, the clumps were calcined at 450
o
C under an air atmosphere. Previous thermogravimetric
analysis (TGA) studies have shown this calcination procedure to be an effective method for
decomposing the Ti polymer by releasing chlorine from its structure [33].

4.2.5. Hydrogen Uptake Experiment
A magnetic suspension balance (MSB), manufactured by Rubotherm, was used to
gravimetrically measure the hydrogen adsorption capacity of the doped SiCNTs. The MSB
recorded the doped SiCNTs weight under various hydrogen pressures, up to 100 bar. A
residual gas analyzer, RGA200 (manufactured by Stanford Research Systems), was coupled
to the MSB in order to analyze the gas inlet/outlet compositions. The purpose of using the
RGA was to ensure that the sorption experiments would not be interfered with by potential
impurities in the adsorptive gas, or by any desorption of volatile compounds from the solid
sample during the adsorption experiments. It should be noted here, that thermogravimetric (as
well as static) approaches measure excess (rather than absolute) adsorption, which is defined
as the mass of gas adsorbed minus the mass of gas at the temperature and pressure of the
experiment that has the same volume as the adsorbed layer. Since hydrogen is such a light
gas, the differences between excess and absolute adsorption are small (e.g., the excess
adsorption is estimated to range from 98.9% of the absolute adsorption value at 10 bar to
89.2% at 100 bar). From a standpoint of judging whether a material is a good storage medium,
80

it is the excess adsorption that is the important metric. Further details of the hydrogen
adsorption experiments are given in our previous publication [22] and in Chapters 2 and 3.

4.3. Results and Discussion
To illustrate the tubular structure of the SiCNTs, an SEM image of the pure SiCNTs,
before doping, is shown in Figure 4.1(A); it is also used as a reference for determining the
influence of the doping process on the SiCNTs. The EDX analysis of the pure SiCNTs
determined their atomic composition to be Si (51%), C (48%), and O (1%). The small amount
of SiO2 impurity, produced during the SiCNTs synthesis, that remains behind even after the
treatment with the hot NaOH solution, is the likely source of oxygen in their structure [23].
As the SEM image of the K-doped SiCNTs shown in Figure 4.1(B) indicates, the doping
procedure has not damaged the structure of the SiCNTs, as they appear to have retained their
individual tubular structure. The EDX analysis of the K-doped SiCNTs has determined their
atomic composition to be Si (50%), C (45%), K (3%), and O (1%).
The SEM image of the Ti-doped SiCNTs before calcination, shown in Figure 4.1(C),
reveals the production of some particles with considerably larger dimensions than the SiCNTs
during the doping process. In addition, the Ti-doped SiCNTs seem to be sticking together
(forming clumps, as noted previously), which is potentially unfavorable for hydrogen storage
applications. The EDX analysis of the Ti-doped SiCNTs before calcination determined their
atomic composition to be Si (38%), C (36%), Ti (5%), Cl (17%), and O (4%). Figure 4.1(D)
depicts the SEM image of the Ti-doped SiCNTs after calcination at 450
o
C for 2 hr. It
indicates that after the calcination reaction, the large particles seen in Figure 4.1(C) have
81

vanished. Moreover, the EDX analysis of the Ti-doped SiCNTs after calcination did not
detect any chlorine, and determined their atomic composition to be Si (46%), C (43%), Ti
(6%), and O (5%). Therefore, as Figure 4.1(D) and the corresponding EDX results indicate,
the calcination step removes the side products of the Ti-doping reaction, mainly chlorine,
while keeping the SiCNTs intact.
Figure 4.2 shows the hydrogen uptake (the adsorption isotherms) of the K-doped and Ti-
doped SiCNTs at 25
o
C for pressures of up to 100 bar. To provide a better insight, the
hydrogen storage capacity (the sorption isotherms) of the SiCNTs before doping, as well as
the corresponding isotherm for the CNTs are also included in Figure 4.2. As Figure 4.2
illustrates, doping of the SiCNTs with K increases their hydrogen storage capacity over that of
the CNTs by a factor of about 3 at low pressures and about 2 at high pressures, which is
consistent with the previous computational results for the alkali doped-SiCNTs/H2 systems
[28]. However, Ti-doping of the SiCNTs seems to have an adverse influence on their
hydrogen storage capacity, since the adsorption isotherm for the Ti-doped SiCNTs are lower
than the corresponding one for the pure SiCNTs. The experimental results with the doped
SiCNTs are in agreement with the molecular simulations for the alkali (K) – SiCNTs [24]

but
in sharp contrast with the simulations for the Ti-doped SiCNTs [27, 29]. This may be,
potentially, due to the oxygen impurity being present in the structure of Ti-doped SiCNTs (as
shown by EDX), which is likely to be bound to the Ti dopant, while pure Ti was used in the
molecular simulations (The oxygen impurity is present in the SiCNT structure even before the
calcination process; though some of that may have been avoided by carrying out the doping
step in a glove-box under an inert atmosphere, however, the need for eventually using air
calcination to remove the chlorine impurity makes this a rather fruitless precaution).
82

Consideration of the desorption isotherms reveals that the doped-SiCNTs possess a stronger
hydrogen desorption hysteresis than the pure SiCNTs. In addition, and in agreement with the
behavior of the CNTs (but in sharp contrast with the behavior of the pure SiCNTs) the doped-
SiCNTs exhibit substantial sorption irreversibility. This, as in the case of CNTs [22], indicates
the presence of two different types of sites one associated with weak (“physisorption”) and
the other with strong (‘chemisorption”) adsorption behavior. The latter sites, with increased
binding energy, which makes the release of adsorbed hydrogen from the SiCNTs’ surface
more difficult, are likely to be the result of the presence of the dopants [18, 27-29].

83

84

Figure 4. 1. SEM image of, (A) pure SiCNTs; (B) K-doped SiCNTs; (C) Ti-doped before
calcination, and (D) Ti-doped after calcination.

In order to better judge the value of using a nanomaterial as a hydrogen storage system, it
is essential to take its solid volume into account, since if the material was not present that
volume would be occupied by the gas. The sorption capacity corrected for the solid volume is
defined in this work as the additive hydrogen storage capacity (AHSC). AHSC can be
calculated by subtracting from the excess adsorption values measured by the TGA the mass of
hydrogen in the bulk gas phase that has been replaced by the solid volume of the
nanomaterial. AHSC is an important factor in assessing the practical value of the application
of solid porous materials for gas storage purposes, since adding them to a storage vessel
85

would lead to a decrease in the available volume of the system for the bulk gas phase, and the
AHSC directly takes into account that impact [34].

Figure 4. 2. Hydrogen uptakes for the SiCNTs doped with Potassium (K-doped) and Titanium
(Ti-doped). The uptakes for the pure SiCNTs (before doping), and CNTs are also shown for
comparison.

The AHSC for the K-doped and Ti-doped SiCNTs, as well as those for the CNTs and pure
SiCNTs are shown in Figure 4.3. For pressures higher than 10 bar, the AHSC of the CNTs
turns negative. This implies that for pressures larger than 10 bar the total hydrogen stored in a
vessel filled with the CNTs is less than the amount of hydrogen that could be stored in the gas
86

bulk phase in the same vessel in the absence of the CNTs. Since the excess adsorption
measured here with these CNTs is pretty much in line with most of the recent technical
literature [35-39], one may then conclude that using the pure CNTs as a hydrogen storage
medium at room temperature appears not to be promising. On the other hand, the AHSC of
both the pure and the K-doped SiCNTs are positive for the entire 0-100 bar range of
pressures, confirming the superiority of the K-doped SiCNTs relative to the CNTs as a
hydrogen storage material.

Figure 4. 3. Additive hydrogen storage capacity (AHSC) values for the K-doped and Ti-doped
SiCNTs . The AHSC values for the pure SiCNTs (before doping), and CNTs are also shown
for comparison.

87

In addition, the AHSC of the K-doped SiCNTs reaches a maximum at a pressure of 20 bar,
hence suggesting this as the optimal pressure for their hydrogen storage application. However,
the strong hysteresis behavior and sorption irreversibility are a key handicap that must be
overcome prior to the use of these materials in hydrogen storage. The Ti-doped SiCNTs are
less promising than the K-doped SiCNTs, and surprisingly enough (given the previous
molecular simulation studies) less so than the pure SiCNTs themselves. Their AHSC is
positive for pressures up to 80 bar, and becomes slightly negative above that. Therefore, using
Ti-doped SiCNTs as a storage medium is less promising and likely.

4.4. Conclusions
The gravimetric hydrogen uptake measurements of K-doped SiCNTs reveal a significant
increase in their hydrogen adsorption capacity, when compared with pure (un-doped) SiCNTs.
An increase in the binding energy caused by the charge transfer from the metal impurity to the
Si and C atoms is the reason for the superior performance of the doped nanotubes (As a
reminder, the pure SiCNTs, themselves, are also capable of adsorbing a much higher amount
of hydrogen, when compared with the CNTs, mostly due to the charge transfer from the Si
atoms to the C atoms in the SiCNTs’ structure). However, the presence of adsorption sites
with increased binding energy in the K-doped SiCNTs gives rise to desorption irreversibility.
This means a higher temperature would be required to release the adsorbed hydrogen
molecules, which is unfavorable from an energy storage point of view.
Interestingly enough, the Ti-doped SiCNTs show inferior performance than the pure
SiCNTs (though still superior to that of CNTs). This contrasts prior molecular simulation
88

studies that indicate that the doping with Ti should significantly improve the sorption capacity
of the SICNTs. The reasons for this surprising finding are not entirely clear at this point and
time (and should be further investigated), but the behavior is suspected to be due to the
presence of an oxygen impurity introduced during the preparation of the materials.
The AHSC of the pure and K-doped SiCNTs are positive throughout the region of
pressures investigated. This is definitely an important improvement over the CNTs, for which
their AHSC turns negative for pressures higher than 20 bar, implying that a hydrogen storage
system working without any solid adsorbent is capable of storing more hydrogen than a CNT-
based system (at room temperature). It should be noted, however, that though the SiCNTs
exhibit richer behavior than their CNT counterparts, have much higher capacity for hydrogen
storage, and are much more resistant to the reactive environment in which hydrogen is
typically produced, they still are far from being a practical candidate for storing large amounts
of hydrogen, because their storage capacity falls quite short of the target capacity set-forth by
the U.S. DOE. Furthermore, as with the CNTs and other novel nanomaterials, production cost
remains a challenge for their large-volume applications, such as hydrogen storage.

89

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Hydrogen Storage in Decorated Multiwalled Carbon Nanotubes by Ca, Co, Fe, Ni, and Pd
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mechanism in Mg modified multi-walled carbon nanotubes for hydrogen storage.
International Journal of Hydrogen Energy. 2012;37:1919-26.
[4] Silambarasan D, Surya VJ, Vasu V, Iyakutti K. Experimental investigation of hydrogen
storage in single walled carbon nanotubes functionalized with borane. International Journal of
Hydrogen Energy. 2011;36:3574-9.
[5] Liu E, Wang J, Li J, Shi C, He C, Du X, et al. Enhanced electrochemical hydrogen storage
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[8] Jena P. Materials for Hydrogen Storage: Past, Present, and Future. The Journal of Physical
Chemistry Letters. 2011;2:206-11.
[9] Singh P, Kulkarni MV, Gokhale SP, Chikkali SH, Kulkarni CV. Enhancing the hydrogen
storage capacity of Pd-functionalized multi-walled carbon nanotubes. Applied Surface
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[10] Lee S-Y, Park S-J. Effect of temperature on activated carbon nanotubes for hydrogen
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[11] Park S-J, Lee S-Y. Hydrogen storage behaviors of platinum-supported multi-walled
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microwave-treated multi-walled carbon nanotubes. International Journal of Hydrogen Energy.
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[16] Gayathri V, Devi NR, Geetha R. Hydrogen storage in coiled carbon nanotubes.
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material for hydrogen storage. Nano letters. 2006;6:1581-3.
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hydrogen storage in carbon nanotubes. International Journal of Hydrogen Energy.
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[24] Pellenq RJ-M, Marinelli F, Fuhr JD, Fernandez-Alonso F, Refson K. Strong
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silicon carbide nanotubes and the corresponding adsorption of hydrogen molecules to the Ti
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[32] Lillo-Ródenas MA, Cazorla-Amorós D, Linares-Solano A. Understanding chemical
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[33] Zhu Y, Zhang L, Gao C, Cao L. The synthesis of nanosized TiO 2 powder using a sol-gel
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[36] Dodziuk H, Dolgonos G. Molecular modeling study of hydrogen storage in carbon
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94

Chapter 5
Solubility and Diffusivity of H
2
and CO
2

in the Ionic Liquid [bmim][PF
6
]
5. References
5.1. Introduction
4

Hydrogen is among the most promising candidates as a renewable-energy carrier and,
therefore, extensive efforts have been made to date to improve its production, purification,
and storage techniques [1-7]. The use of ionic liquids (IL) is attracting significant attention in
this area, due to their non-flammability, non-toxicity, negligible vapor pressure, wide range of
viscosities (10–1000 cp), and high thermal and chemical stability [8].
Ionic liquids manifest good promise for use with both absorption-[9] and membrane-based
separations [10] of H2-containing gas mixtures. One example is the separation of H2 from CO2
that is of practical importance, because CO2 is a key species involved in hydrogen production
reactions (e.g., steam reforming of methane), and in the subsequent purification steps. For
example, membranes incorporating IL have a high CO2/H2 separation factor [10], a unique
characteristic known as reverse selectivity, whereby larger molecules, such as CO2, permeate
faster through a membrane than smaller molecules, such as H2. This is thought to be due to

4
The material in this Chapter is part of a paper previously submitted to Int. J. Hydrogen Energy: Barghi SH, Tsotsis TT,
Sahimi M. Solubility and diffusivity of H2 and CO2 in the ionic liquid [bmim][PF6].

95

the considerably higher solubility of CO2 in the IL. According to Yokozeki et al. [10], at room
temperature IL solvents possess CO2/H2 sorption selectivities in the range of 30–300,
indicative of the promise of IL for such a separation and for outperforming competitive
polymeric membranes that, under typical operating conditions, have a CO2/H2 selectivity
below 30 [11]. Therefore, in a process equipped with IL-based membranes, hydrogen can be
potentially produced at high pressures without requiring expensive recompression, which is a
weakness of the processes using conventional membranes employing a molecular-sieving
mechanism for separation [10].
Ionic liquids have also been studied as a reaction medium during hydrogen synthesis and
storage applications. For example, in a laboratory study De Souza et al [12], investigated the
application of IL as the electrolyte for the electrochemical production of hydrogen. They
reported that imidazolium-based IL, such as PF6, BMI, and BF4 are superior electrolytes for
electrochemical hydrogen production, as a result of their high chemical stability. Bluhm et al.
[13] studied the effect of using 1-butyl-3-methylimidazolium chloride IL as a carrier (solvent)
for ammonia borane systems during their use as hydrogen storage media. They reported the
positive effect of the IL on both the extent and the rate of hydrogen release in the
temperatures range 85 - 95
o
C.
An important aspect of using IL for gas separations, either as absorption media or in the
form of IL-based membranes, is the ability to accurately determine the solubility and
diffusivity of the various components in the gas mixture to be separated. For example,
knowledge of such properties is essential for estimating the permeability and selectivity of IL-
based membranes [14-21] and, thus, their measurement has become of prime importance in
96

recent years. There are four main methods that have been used to investigate experimentally
the solubility and/or diffusivity of gases in IL systems, namely, (i) the synthetic (bubble-
point) method, used for measuring solubility, which involves visual detection of a phase
transition due to the increase in the system pressure at constant temperature and composition;
(ii) the isochoric saturation method, which involves monitoring the pressure decay of the gas
phase as a result of gas dissolution in the IL phase; (iii) the supported IL-membrane method,
which involves the measurement of the gas permeability through a membrane impregnated
with IL; and (iv) gravimetric methods [22], which involve the detection of changes in the IL
weight after exposure to the gas of interest. Among them, the last three techniques have the
added advantage that one can measure simultaneously the gas solubility and diffusivity.
Nevertheless, the accuracy of the isochoric and the supported IL-based membrane techniques
decrease considerably when one deals with low-solubility gases, such as hydrogen. This is
mostly due to the inherent susceptibility of the two methods to gas leaks and temperature
fluctuations [14, 15].
In our previous studies of the adsorption of hydrogen in carbon and silicon-carbide
nanotubes [23, 24], (see Chapters 3 and 4), we utilized a highly accurate gravimetric
measurement technique involving the use of a magnetic suspension microbalance. A key goal
of this Chapter is to extend the applicability of this technique as a viable experimental tool to
simultaneously measure the solubility and diffusivity of various gases in ionic liquids. As a
model IL, we have selected 1-butyl-3-methylimidazolium hexafluorophosphate
([bmim][PF6]), and as a model mixture the gas pair H2/CO2. The practical importance of
separating H2 from CO2 has already been discussed. The reasons for selecting this particular
IL is because it is a common and widely available material that exhibits high solubility
97

towards CO2, which makes it a good candidate for various CO2 removal applications [15, 25],
(for example, CO2 separation and capture from flue-gas), including the H2/CO2 separation
relevant to hydrogen production under study here. There are, in addition, several papers
reporting data for CO2 solubility and diffusivity in the IL that we use, using the
aforementioned other techniques that could be used to compare the data generated by our
experimental method [15, 25-29].
One of the motivating factors in our search for a new technique to measure hydrogen
diffusivity and solubility in IL is the fact that previous efforts to measure such properties in
IL, including [bmim][PF6], were either unsuccessful when measuring diffusivity, due to the
lower sensitivity of the methods used [9], or inconclusive, when measuring solubilities. For
[bmim][PF6], in particular, the reported hydrogen solubility data are widely scattered. For
example, Jacquamin et al.[26] investigated hydrogen solubility in [bmim][PF6] in the
temperature range 10
o
C - 70
o
C, using the isochoric saturation method. They reported the
Henry's constant for hydrogen, 𝐾 𝐻 2
, to be 188 MPa at 40
o
C (on a mole fraction basis).
However, Kumelan et al.[30] used a similar isochoric technique, but determined 𝐾 𝐻 2
in
[bmim][PF6] to be 422 MPa at 40
o
C, instead. Dyson et al.[31] utilized a high-pressure H-
NMR spectroscopy method to measure hydrogen solubility in [bmim][PF6], and reported 𝐾 𝐻 2

to be 538 MPa at 25
o
C, considerably higher than the value of 210 MPa that was reported by
Jacquamin et al. [26] under the same operating conditions. Therefore, in addition to the
practical importance of establishing the technique as a standard experimental tool for
measuring solubility and transport characteristics of gases (particularly sparingly soluble ones
like H2) in IL, the measurements reported here aim to fill an important need for such reliable
data for the important gas pair (H2/CO2)/IL([bmim][PF6]) system.
98

The rest of this Chapter is organized as follows: In the next Section 5.2 the experimental
procedure is described. Section 5.3 describes the measurements of the solubility and
diffusivity. The results are presented in Section 5.4, while Section 5.5 summarizes the
Chapter.

5.2. Experimental
5.2.1. Materials
The ionic liquid [bmim][PF6] with the chemical composition C8H15F6N2P, with a purity of
better than 97% and a water content of less than 0.1%, was purchased from the Sigma Aldrich
Chemicals Company. Ultrahigh purity hydrogen (99.999% pure) and research-grade carbon
dioxide (99.99% pure) were both provided by Gilmore Liquid Air Company, USA.

5.2.2. Gas Absorption Measurements
In order to measure the hydrogen solubility and diffusivity over a given temperature range,
approximately 1 g of the [bmim][PF6] IL was loaded into the sample holder of the magnetic
suspension microbalance manufactured by Rubotherm. For the CO2 uptake experiments, a
smaller amount (~0.4 g) of [bmim][PF6] was used, due to the high CO2 solubility in the ionic
liquid. We then degased the [bmim][PF6] sample under vacuum at 70
o
C for 24 hr (during
which no change in the weight of the IL sample was observed). After the degasing step was
completed, the system's temperature was set to the value at which gas absorption was to be
investigated, in this study in the range 25
o
C - 55
o
C. The system pressure was then increased
99

stepwise by injecting the absorptive gas into the gravimetric vessel, and the sample's weight
was recorded. More details about the experimental set-up used in this Chapter and the
operating procedures followed are given in publications [23, 24] and in Chapter 2.

5.3. Measurement of the Diffusivities
Figure 5. 1 shows a schematic of the gas/ionic liquid system used in our experiments. The
sample holder is a cylindrical vessel with a diameter of 1.6 cm and a height of 3 cm. Based on
the density of [bmim][PF6], 1 gr of IL sample at 25
o
C and a pressure of 1 MPa occupies a
volume of 0.73 ml (a cylindrical slice with a height of 0.38 cm) at the bottom of the cell. The
density of the IL is measured in situ during the experiment using He, and the sample volume
(thickness) corresponding to the density is used in the analysis of the data (see below). As the
temperature and pressure of the experiment varies, the density of the IL varies slightly (for all
the data reported here the variation was less than ±1%).
Since no transport occurs through the side-walls of the sample-holder, diffusion in the
ionic liquid can be considered to be a 1D process along the height of the sample, governed by
the standard diffusion equation
𝜕 𝐶 𝑖 𝜕𝑡
= 𝐷 𝑖 𝜕 2
𝐶 𝑖 𝜕 𝑥 2
{
𝐶 𝑖 (𝑥 , 0) = 𝐶 𝑖 0

𝐶 𝑖 (0, 𝑡 ) = 𝐶 𝑖 ∗
𝜕 𝐶 𝑖 (𝐿 ,𝑡 )
𝜕𝑥
= 0
(5.1)
where, Ci (mol/cm
3
) is the concentration of the absorbed gas in the ionic liquid sample at
height x and time t, 𝐶 𝑖 0
(mol/cm
3
) is the initial concentration of the gas dissolved in the IL,
100

𝐶 𝑖 ∗
(mol/cm
3
) is the concentration of the gas in the top surface (x=0) of the IL sample in
contact with the flowing gas, and Di (cm
2
/s) and L (cm) are, respectively, the gas diffusivity in
the IL phase and the IL sample’s thickness. The volume of the sample (and the corresponding
thickness L), based on the density of IL measured under the same temperature/pressure
conditions with the gravimetric balance using He, is used in the analysis. We have assumed
that in the presence of H2 and CO2 the density of the IL does not deviate from its value
measured under the same temperature/pressure conditions using He. For H2, for which the
maximum solubility measured in the IL (see Figure 5. 6, below) is less than 0.04% (g H2/g
IL), this by all accounts is quite a safe assumption. For CO 2, for which solubilities in IL are
significantly higher (see Figure 5.5, below), this assumption would appear to be questionable.
If the thickness of IL was to significantly change with CO2 dissolution, however, one would
expect the calculated diffusivities (based on the IL sample thickens from the He-based
densities) to vary significantly as well. However, our own experiments (see Figure 5.4)
indicate otherwise. Others [25, 32-34] have made similar observations in the past.
The solution of Eq. (5.1) is given
𝐶 𝑖 −𝐶 𝑖 0
𝐶 𝑖 ∗
−𝐶 𝑖 0
= 1 −
4
𝜋 ∑
1
2m+1
exp[
−(2𝑚 +1)
2
𝜋 2
4
(
𝐷 𝑖 𝑡 𝐿 2
)] sin[
(2𝑚 +1)𝜋 2
(
𝑥 𝐿 )]
∞
𝑚 =0
(5.2)
Though Eq. (5.2) is an infinite series, higher order terms become progressively smaller in
size. In the analysis of the experimental results, therefore, we have used the first 10 terms in
the series (the value of the 11
th
term is smaller than the 10
th
term by a factor of 10
4
) and have
disregarded all additional higher-order terms. To interpret the experimental values obtained
by the gravimetric experiment, it is necessary to determine the time-dependence of the total
101

number of moles of gas dissolved in the IL. This is done by integrating Eq. (5.2) over the
volume of the IL sample (from x=0 to x=L), which results in the following equation:
𝑛 𝐴 𝑤 𝐼𝐿
=
𝑆𝐿 𝐶 𝑖 ∗
𝑤 𝐼𝐿
−
8𝑆𝐿 (𝐶 𝑖 ∗
−𝐶 𝑖 0
)
𝑤 𝐼𝐿
𝜋 2
∑
exp [−(2𝑚 +1)
2
𝜋 2
(
𝐷 𝑖 𝑡 4𝐿 2
)]
(2𝑚 +1)
2
∞
𝑚 =0
(5.3)
where 𝑛 𝐴 (mol) is the total moles of the gas dissolved, 𝑤 𝐼𝐿
(kg) is the IL mass, and S (cm
2
) is
the cross-sectional area of the sample holder containing the ionic IL. Equation (5.3) is then
used to estimate the diffusivity of the dissolved gas in the ionic liquid with the aid of
MATLAB's nonlinear least-squares curve fitting toolbox.

Figure 5. 1. Schematic of the gas/ionic liquid system used in this work.

5.4. Results and Discussion
Figure 5.2 shows the measured CO2 molality (mol/kg) in the IL as a function of time at 35
o
C and 0.04 MPa of CO2 pressure. The fit of the data using Eq. (5.3) is also shown in Figure
5.2. The values of the fitted parameters are
𝐶 𝑖 ∗
𝑆𝐿
𝑤 𝐼𝐿
= solubility of CO2 in IL at equilibrium =
0.0149 mol/kg and the average diffusivity Di=2.51×10
-6
cm
2
/s. These estimates are consistent
102

with the data reported in the literature (see further discussion below) [15, 18, 25, 28], for CO2
in [bmim][PF6] and, therefore, confirm the accuracy of our experimental method.

Figure 5. 2. The experimental data and their fit to Eq. (5.3) for CO2 molality in [bmim][PF6]
at 0.04 MPa and 35
o
C.

Figure 5.3 presents the corresponding measured H2 molality (mol/kg) in the IL as a
function of time at the same temperature of 35
o
C, but at (a much higher) pressure of 2 MPa.
The fit of the data using Eq. (5.3) is also shown there, and seems to “track” the experimental
data fairly well. The values of the fitted parameters are
𝐶 𝑖 ∗
𝑆𝐿
𝑤 𝐼𝐿
= 0.0332 mol/kg, and 𝐷 𝑖 =
9.49×10
-6
cm
2
/s. To our knowledge, this is the first time that the hydrogen diffusivity in
[bmim][PF6] (but also in any IL for that matter) has been measured experimentally. As
103

pointed out earlier, previous attempts by other groups to directly measure the solubility and
diffusivity of hydrogen and other gases with low solubility (e.g., oxygen and nitrogen) in ILs
had failed, due to the inadequate sensitivity of the experimental methods used [9, 18].
Using the same method, the hydrogen diffusivity was measured at the same temperature
for pressures of up to 10 MPa, and as with CO2 (Figure 5.4), the results are within ± 5% of
the estimated diffusivity at 2 MPa. This apparent insensitivity of hydrogen diffusivity to
pressure is to be expected, since diffusivity in liquids is, generally, a very weak function of
pressure.

Figure 5. 3. The experimental data for H2 molality in [bmim][PF6] at 2 MPa and 35
o
C, and
their fit to Eq. (5.3).
104

Figure 5. 4. H2 and CO2 diffusivity vs. pressure at 25
o
C.

Our experimental data for carbon dioxide molality at equilibrium in [bmim][PF6] as a
function of pressure, for pressures up to 5 MPa and at four different temperatures 25
o
C, 35
o
C, 45
o
C and 55
o
C, are shown in Figure 5.5. As previous authors also have reported [9, 25,
26, 30, 35],

even for such high CO2 pressures the data are still well described by a Henry-like
law. The Henry's constant for CO2, 𝐾 𝐶𝑂
2
(MPa), is calculated from the regression of the
solubility data. Table 5.1 summaries the results for the 𝐾 𝐶𝑂
2
measured in this work, as well as
other corresponding values reported in the literature [9, 25, 26, 30, 35]. It is evident from
Table 5.1 that the 𝐾 𝐶𝑂
2
values obtained by our gravimetric method are consistent with the
literature data for CO2 solubility in [bmim][PF6] systems.
105

Figure 5. 5. Pressure-dependence of CO2 molality in [bmim][PF6] in the temperature range 25
o
C- 55
o
C. The lines represent the best fit of the data.

Figure 5. 6. Pressure-dependence of H2 molality in [bmim][PF6] in the temperature range 25
o
C-55
o
C. The lines represent the best fit of the data.
106

Table 5. 1. Henry's constant values for CO2 in [bmim][PF6].

𝑲 𝑪𝑶
𝟐 (MPa)
20
o
C 25
o
C 30
o
C 35
o
C 40
o
C 45
o
C 50
o
C 55
o
C
Present work - 5.236 - 6.452 - 7.886 - 9.227
Anthony et al.[9] - 5.34 - - - - 8.13 -
Shiflett et al.[25] - 5.572 - - - - 8.367 -
Jacquemin et
al.[26]
4.678 - 5.709 - 6.834 - 8.053 -
Kamps et al.[28] 4.22 - - - 6.65 - - -
Scovazzo et
al.[35]
- - 5.9 - - - - -

Figure 5.6 depicts hydrogen molality at equilibrium in [bmim][PF6] over the temperature
range 25
o
C - 55
o
C for pressures of up to 10 MPa, as well as the linear regression of the data.
The regression coefficient (R
2
) for all four lines is more than 0.995, hence indicating a
perfectly linear dependence of hydrogen solubility in [bmim][PF6] on pressure. The results for
the hydrogen uptake experiment in [bmim][PF6], as well as the corresponding Henry's
constants are summarized in Table 5.2, which also presents the Henry’s constants reported by
107

previous investigators. As Table 5.2 indicates, our data for the Henry's constants are
consistent with those reported by Jacquemin et al. [26], with the differences being less than
15% for the entire temperature range. In their work Jacquemin et al. [26] measured hydrogen
solubilities using an isochoric saturation technique, and reported that the 𝐾 𝐻 2
is, generally, an
increasing function of temperature in the range of temperatures of (40 - 70
o
C), as one would
normally expect. A similar trend is observed (see Table 5.2) by our 𝐾 𝐻 2
values. The 𝐾 𝐻 2

values reported by Kumelan et al. [30] (see Table 5.2) are substantially higher than our results
and those of Jacquemin et al. [26] Moreover, their 𝐾 𝐻 2
values decrease as the temperature
increases from 40
o
C to 100
o
C, which is counterintuitive, and also contrasts the observed
trends in our own data and those of Jacquemin et al. [26].
The challenge one faces using volumetric approaches to measure the solubility of gases
with low solubilities in IL, such as hydrogen, can be seen in Figure 5.7 where we plot the
hydrogen uptake fitted data (mol H2/kg of IL vs. time) in [bmim][PF6] at two different
temperatures (pressure is 2 MPa for both experiments). As reported above (see Figure 5.6 and
Table 5.2), the H2 solubility (mol H2/kg) at equilibrium is higher at 25
o
C than that at 55
o
C,
and that is also obvious in Figure 5.7 from the uptake data at large times. However, the trends
are quite the opposite during the first 3.4 hr of the experiments, where kinetic effects (i.e.,
differences in the diffusivities) dominate. The H2 diffusivity at 25
o
C is lower than that at 55
o
C, which makes the hydrogen uptake a slower process at this lower temperature. The time
required for a given (gas/IL) system to reach equilibrium could be quite substantial, especially
when employing the volumetric technique [24, 30, 31], since larger quantities of IL are
typically utilized in order to improve experimental sensitivity and to overcome the inevitable
gas leaks (the characteristic time for transport is L
2
/D, where L is the characteristic transport
108

dimension of the IL sample and D the diffusivity of the gas in the IL). Waiting for equilibrium
to be established, particularly at lower temperatures, in order to estimate the gas solubility
may require a long time during which the inevitable system leaks may interfere (particularly
for low absorbing gases such as H2) with the experimental accuracy, since in volumetric
systems the amount absorbed is calculated form differences in the system pressure. The
gravimetric instrument used here directly measures the weight of dissolved hydrogen in the IL
phase and is, thus, immune to system leaks. Further, it requires a much smaller IL sample,
which significantly shortens the time needed to reach equilibrium. Such reasons help explain
the widely scattered hydrogen solubility data in the literature for H2 in [bmim][PF6].

Table 5. 2. The hydrogen Henry's constant values in [bmim][PF6].

𝑲 𝑯 𝟐 (MPa)
25
o
C 30
o
C 35
o
C 40
o
C 45
o
C 50
o
C 55
o
C 60
o
C
Present work 183.27 - 206.51 - 221.24 - 238.51 -
Jacquemin et al.
[26]
- 198.8 189.2 185.8 187.8 196.2 - 225.9
Kumelan et al.[
30]
- - - 415.2 - - - 367.7

109

Figure 5. 7. Time evolution of hydrogen dissolution in [bmim][PF6] at 25
o
C and 55
o
C (P=2
MPa for both temperatures). The lines represent the experimental fit provided by Eq. 5.3.

The Henry's constants for H2 and CO2 at different temperatures could be used to determine
the heat of absorption for the aforementioned gases in [bmim][PF6] according to Eq. (5.4)
∆𝐻 𝑖 = 𝑅 (
𝜕 ln 𝐾 𝑖 𝜕 (
1
𝑇 ⁄ )
) (5.4)
where ∆𝐻 𝑎𝑏𝑠 (kJ/mol) is heat of absorption of the gaseous component in the IL, and 𝑅 the gas
constant (0.008314 kJ/mol. K). Figure 5.8 shows the plots of the Henry's constants for H2 and
CO2 (on a natural logarithmic scale) as a function of 1/T. It can be noticed that both curves are
quite linear and, therefore, the regression of the data could be used to determine the heat of
absorption for the two gases in [bmim][PF6]. Based on the curve-fitting results, the heat of
110

absorption for H2 is ∆𝐻 𝐻 2
= -7 kJ/mol and that of CO2 is ∆𝐻 CO
2
= -15.7 kJ/mol. For CO2,
Anthony et al.[ 9] reported ∆𝐻 CO
2
= -16.1 kJ/mol which is in a good agreement (within 2.5%)
with the value determined in our work. For H2, Kumelan et al. [30] found ∆𝐻 𝐻 2
= -4.9 kJ/mol,
30% smaller than the value obtained in this work, mostly due to the difference in the
measurement techniques used in the two studies.

Figure 5. 8. Temperature-dependence of Henry's constants of H2 and CO2 in [bmim][PF6].
The lines represent the fit of the data to Eq. (5.4).
111

Figure 5. 9. Temperature-dependence of diffusivity of H2 and CO2 in [bmim][PF6]. The lines
represent the fit of the data to Eq. (5.5).

The diffusivities for both H2 and CO2, when measured under isothermal conditions,
indicate no dependence on pressure, as noted previously (see Figure 5.4). Shieflet et al. [25]
measured CO2 diffusion in the same IL using a gravimetric technique, and also reported a
similar observation. The diffusivities, on the other hand, depend on temperature. In Figure 5.9
plots of the diffusivity for H2 and CO2 as a function of (1/T) are presented. Both diffusivities
obey, with a great degree of accuracy, an Arrhenius-like relationship,
ln𝐷 𝑖 = ln 𝐴 𝑖 −
𝐸 𝑖 𝑅𝑇
(5.5)
112

where 𝐷 𝑖 (cm
2
/s) is the diffusivity, T (K) the tempearture, 𝐴 𝑖 the pre-exponential factor
(cm
2
/s), 𝐸 𝑖 (kJ/mol) the activation energy, and 𝑅 the gas constant (0.008314 kJ/mol. K). For
H2, 𝐴 𝐻 2
=146 and 𝐸 𝐻 2
=42.43. For CO2, 𝐴 𝐶𝑂
2
=0.64 and 𝐸 𝐶𝑂
2
=32. Similar functional
dependence on temperature has been reported for the diffusivity of carbon dioxide and
methane in a number of ILs [14, 15, 25]. Using a gravimetric method for CO2 diffusivity in
the same IL, Shiflett et al. [25] reported 𝐸 𝐶𝑂
2
=27.2 (kJ/mol), which is consistent (but 15%
less) with the value obtained in our work.
The Arrhenius-like dependence of the H2 and CO2 diffusivities in [bmim][PF6] is
consistent with what is known about gas diffusion in more conventional liquids. Such
diffusivities are generally related to the liquid viscosity, 𝜇 , through an equation of the form,
𝐷 𝑖 ∝ 𝜇 −𝑛 , where for most compounds n lies between 0.5-1 [25]. In addition, for a narrow
range of temperatures, the temperature-dependence of liquid viscosity is often described by
the Andrade equation [25]: 𝜇 = 𝐴 exp (𝐵 /𝑇 ). Combining the gas diffusivity’s known
dependence on viscosity with the Andrade equation results in an Arrhenius-like equation for
the impact of temperature on diffusivity in ILs [25].

5.5. Conclusions
A gravimetric method was used to measure hydrogen and carbon dioxide solubilities and
diffusivities in the ionic liquid [bmim][PF6] over a temperature range between 25
o
C - 55
o
C
and pressures up to 10 MPa. To our knowledge, this is the first time that the H 2 diffusivity is
reported in any ionic liquid, including [bmim][PF6]. The diffusivity data for both gases
113

indicate very little, if any, dependence on the pressure, indicative of Fickian-like diffusion and
the lack of swelling effects. The data for both gases indicate an Arrhenius-like exponential
dependence on temperature. The solubility data for both H2 and CO2 in the same ionic liquid
indicate a Henry-like dependence for the entire ranges of pressures and temperatures studied.
Both Henry’s constants increase with increasing temperature, as one may have expected. This
is the first time that both solubility and diffusivity data are reported for the important
(H2/CO2) gas pair in [bmim][PF6], or in any IL for that matter. Hopefully, this will motivate
the use of the proposed technique for measuring the transport and thermodynamic properties
of other important gas mixtures in this and other ionic liquids.

114

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[9] Anthony JL, Maginn EJ, Brennecke JF. Solubilities and Thermodynamic Properties of
Gases in the Ionic Liquid 1-n-Butyl-3-methylimidazolium Hexafluorophosphate. The Journal
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117

Chapter 6
Future Work

As discussed in Chapters 2 and 3, the hydrogen uptake capacities of carbon nanotubes
(CNTs) and silicon-carbide nanotubes (SiCNTs) are far below the target set by the United
States Department of Energy (DOE), which is 6 wt.% (weight of hydrogen/weight of the solid
adsorbent). Therefore, in Chapter 4, we investigated the effect of doping SiCNTs with Ti and
K because based on the previous theoretical studies, doping the nanotubes is a promising
candidate method to increase their hydrogen storage capacities. Although doping SiCNTs
with K showed some promising improvement for their hydrogen storage characteristics, it
would be worthwhile to also try other possible dopants. Alkali metals, including lithium and
sodium, and also aluminum, and boron are the dopants that have been the subject of previous
simulation studies in the literature.
The most important challenges for doping SiCNTs are: (i) identifying the best doping
material that significantly increases the hydrogen storage capacity of the nanotubes; (ii)
developing a doping technique that does not damage the structure of the nanotubes (this is
important since the destruction of the structure of nanotubes would lead to a decrease in the
surface area of the solid adsorbent which is an unfavorable outcome); (iii) improving the
adsorption/desorption kinetics of the doped nanotubes; (iv) minimizing the desorption
hysteresis of the doped nanotubes, since the hysteresis would make a portion of the stored
118

hydrogen inaccessible to the vehicle’s engine. The development of a technique that satisfies
all these aforementioned requirements would be a worthwhile undertaking.

Abstract (if available)

Abstract Hydrogen is considered a promising renewable energy source. Developing a safe and efficient hydrogen storage medium is the most crucial challenge for the commercial application of hydrogen as fuel for vehicles. Adsorptive hydrogen storage on solid adsorbents, including carbon nanotubes (CNTs) and silicon-carbide nanotubes (SiCNTs), can be a novel storage technique. In the present work, the hydrogen storage behavior of CNTs, pure SiCNTs, and doped SiCNTs is investigated experimentally using a combination of gravimetric and mass analysis methods at room temperature for pressures of up to 100 bar. The combination of the techniques enabled us to obtain experimental data for hydrogen adsorption and desorption, and the rate by which they occur. ❧ In the first part of this work, the question of chemisorption versus physisorption during hydrogen storage in carbon nanotubes (CNTs) is addressed experimentally. We utilize a powerful measurement technique based on a magnetic suspension balance coupled with a residual gas analyzer, and report new data for hydrogen sorption at pressures of up to 100 bar at 25 ℃. The measured sorption capacity is less than 0.2 wt.%, and there is hysteresis in the sorption isotherms when multi-walled CNTs are exposed to hydrogen after pretreatment at elevated temperatures. The cause of the hysteresis is then studied, and is shown to be due to a combination of weak sorption—physisorption—and strong sorption—chemisorption—in the CNTs. Analysis of the experimental data enables us to calculate separately the individual hydrogen physisorption and chemisorption isotherms in CNTs that, to our knowledge, are reported for the first time here. The maximum measured hydrogen physisorption and chemisorption are 0.13 wt.% and 0.058 wt.%, respectively. ❧ In the second part of this work, we report for the first time the results of an extensive experimental study of hydrogen sorption in SiCNTs, which were synthesized using the reaction between SiO vapor and CNTs in an argon atmosphere in the temperature range 1200 ℃ - 1500 ℃. The as-synthesized SiCNTs were then purified using a sodium hydroxide solution, in order to remove the side products of the synthesis reaction. The hydrogen sorption characteristics of the as-synthesized SiCNTs, as well as those of the purified SiCNTs were then measured at 25 ℃ and for pressures of up to 100 bar. The results reveal hysteresis between the adsorption and desorption isotherms, which we attribute to the presence of metal impurities and/or the multilayer structure of the nanotubes. The hydrogen storage capacity of the as-synthesized SiCNTs is similar to that of the CNTs, whereas for the purified SiCNTs it is 50% higher than that of the CNTs. In addition, the hydrogen uptake rate in the SiCNTs is about five times faster than that in the CNTs and, in contrast with the CNTs, its desorption from SiCNTs is completely reversible under vacuum. ❧ In the third part of this work, the SiCNTs were synthesized by the gas/solid reaction between SiO vapor, as the Si source, and carbon nanotubes as the carbon precursor. The resulting SiCNTs were then purified with a hot and concentrated NaOH solution in order to remove the amorphous silica from their surface. The purified nanotubes were then doped with either K or Ti. The hydrogen adsorption behavior of the nanotubes was gravimetrically measured with the aid of a magnetic suspension balance. According to the results, the K-doped silicon carbide nanotubes show more promise as hydrogen adsorption materials among all the nanotube samples. ❧ In the last part of this work, we report on the measurement of the solubility and diffusivity of H₂ and CO₂ in the ionic liquid 1-butyl-3-methylimidazolium hexafluorophosphate ([bmim][PF6]) over the temperature range 25 ℃ - 55 ℃ and for pressures of up to 10 MPa. The data were obtained using a magnetic suspension balance, a gravimetric instrument that is capable of simultaneously and accurately measuring the gas solubility and diffusivity in such liquids. To our knowledge, this is the first time that the H₂ diffusivity has been measured in an ionic liquid. While solubility data for H₂ have been previously reported, they vary widely. The cause for such variations is discussed as well. The diffusivity data for both H₂ and CO₂ follow an Arrhenius-like dependence on temperature.

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Asset Metadata

Creator Barghi, Seyedhamed M. (author)

Core Title Hydrogen storage in carbon and silicon carbide nanotubes

School Viterbi School of Engineering

Degree Doctor of Philosophy

Degree Program Chemical Engineering

Publication Date 04/20/2016

Defense Date 03/03/2015

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag [bmim][PF6],adsorption,carbon dioxide,carbon nanotubes,desorption,diffusivity,hydrogen storage,ionic liquids,OAI-PMH Harvest,SiC nanotubes,solubility

Format application/pdf (imt)

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor Sahimi, Muhammad (committee chair), Tsotsis, Theodore T. (committee chair)

Creator Email hamedm.barghi@gmail.com,mbarghi@usc.edu

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c3-554892

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