Chapter 7: Hydrogen Storage on Carbon Adsorbents

In recent years, global energy demand and global warming due to greenhouse gas emissions have spurred intensive research for alternative fuel. Hydrogen (H₂), an ideal energy carrier, is obviously a promising candidate. Its utilization efficiency is high; for instance, it is 2.75 times greater than gasoline for the same weight. Moreover, it is environmentally benign with negligible pollutant emissions [1–7]. H₂ can be produced through the conversion of various resources, including different fossil fuels, such as coal, oil, and natural gas, as well as a variety of biomass. The energy required for this conversion can be generated from such resources themselves or from other energy sources, such as wind, solar, and nuclear. As an example, H₂ can be produced indirectly from coal gasification and reforming processes for which advanced technologies are available; these processes, combined with carbon dioxide (CO₂) separation and sequestration, have the potential to manufacture substantial quantities of H₂ with minimum greenhouse gas emissions [7, 8].

Synthesized H₂ can be deployed to generate electricity from fuel cells; alternatively, it can be combusted for providing energy for space heating, replacing natural gas in industry, and fueling aircraft [9]. For such deployments, H₂ needs to be effectively stored for stationary or mobile application. The former entails the installation of storage facilities in a relatively large area, execution of multistep chemical charging/recharging cycles at low temperatures and under high pressures, and provision of extra capacity to compensate for slow kinetics. In contrast, the latter demands the operation with minimum volume and weight specifications, adequate supply of H₂ for the acceptable driving range, i.e., over 380 km (300 miles) on a full tank, with the charge/recharge near room temperature, and supply of H₂ at sufficiently high rates for fuel-cell locomotion of various vehicles [10, 11]. Nevertheless, barriers to gassing up these vehicles with H₂ are substantial. At present, it is impossible to store H₂ as compactly and simply as the conventional liquid fuel. According to George J. Thomas, formerly with Sandia National Laboratories, “…one of the key enabling technologies on which a future hydrogen economy rests is hydrogen storage. The topic is central to all aspects of hydrogen usage…” [12]. Thus, one of the most challenging technical issues remaining unresolved in developing a H₂-based energy system is the efficient and safe storage of a sufficient quantity of H₂ on board without a concomitant, appreciable increase in the weight or volume to the vehicle.

To meet the requirements for H₂ as a transportation fuel, the US Department of Energy (DOE) has mandated that the target energy densities for the development of onboard H₂-storage systems be 5.5 wt% (1.8 kWh/kg system) and 40-kg H₂/m³ (1300 kWh/m³) by year 2015 and 7.5 wt% (2.5 kWh/kg system) and 70-kg H₂/m³ (2300 kWh/m³) ultimately [6]. Currently, three major approaches are available for H₂ storage: (1) pressurization of H₂ in high-pressure vessels, (2) liquefaction of H₂ at the cryogenic temperature, and (3) materials-based H₂ storage [6].

For most of the current prototype fuel-cell vehicles, numbering hundreds, H₂ is stored in high-pressure cylinders at a pressure ranging from 5000 psi (~35 MPa) to 10 000 psi (~70 MPa), which can be readily tapped for use [13]. Strong- and light-weight tanks, such as those fabricated from carbon-fiber reinforced composites, render it possible to safely store H₂ at high pressures. Nevertheless, the density of stored H₂ does not proportionally increase with the increase in the pressure. Even at a pressure of 10 000 psi, the energy content of H₂ is 4.4 MJ/L, significantly less than that for the same volume of gasoline, which is 31.6 MJ/L. Available high-pressure tanks can contain merely about 3.5–4.5 wt% of H₂, which is far from attaining the aforementioned target energy density.

Naturally, the second approach, i.e., liquefaction of H₂ at the cryogenic temperature, can increase the energy density of H₂ to yield a system with 14 wt% and 51-kg H₂/m³, corresponding to the energy content of 8.4 MJ/L, which approximate the aforementioned target energy densities. Transporting or handling liquefied H₂, however, tends to be exceedingly costly. In fact, the cost of liquefied H₂ as a transportation fuel is nearly twice that of gaseous H₂, thereby hindering its deployment. In addition, the extremely low boiling point of H₂ necessitates special precautions for safe handling. Furthermore, any substantial improvement can only be achieved at the expense of severe operating conditions and/or high fuel consumption [14]. Consequently, much effort is needed to reduce the cost of liquefaction for it to be economically viable.

The most promising method for H₂ storage appears to be through materials-based storage technologies, which are inherently safe and should be more energy efficient than pressurization or liquefaction. Materials for such technologies include carbon-based materials, metal-organic frameworks, metal hydrides, conducting polymers, and clathrates [6]. Among them, it is highly likely that carbon-based materials, or carbon adsorbents, have the greatest potential. They are light, inexpensive, and ideal from the standpoint of the gravimetric requirements for deploying fuel cells of automobiles. Moreover, the adsorbent-manufacturing industry has substantial expertise in the production of carbon adsorbents capable of effectively controlling their micro- and nano-structures for large-scale production. Intensive investigations have been performed on H₂ storage on carbon adsorbents, including activated carbons, carbon nanofibers (CNFs), single-walled carbon nanotubes (SWCNTs), and multiwalled carbon nanotubes (MWCNTs).

2. FUNDAMENTALS OF ADSORPTION

Adsorption is a process, which increases in the density of a gas or solute (adsorbate) in the vicinity of the surface of a substrate (adsorbent) due to molecular interactions between the adsorbate and the adsorbent. Gases can undergo supercritical and subcritical adsorption. The former takes place above the gas critical temperature, thereby exhibiting no gas–liquid phase transition. The latter is observed below the gas critical temperature, and its isotherm is characterized by a steep increase in the adsorbed density close to the saturation pressure associated with the formation of a liquid layer on the surface of the adsorbate. Because of the extremely low critical temperature of H₂, i.e., 33.18 K, only supercritical adsorption needs to be taken into account for H₂ storage [15].

2.1. Volumetric and Gravimetric Densities

The H₂-storage system is characterized by its volumetric and gravimetric densities. The amount of H₂ in the adsorbent-filled container includes the H₂ adsorbed by the adsorbent and the bulk H₂ in the nonadsorbing volume of the container. To facilitate the delineation of the role of the adsorbent on H₂ storage, a thermodynamically significant notion of the “excess amount” has been proposed; it is defined as “…the excess material present in the pores over that which would be present under the normal density at the equilibrium pressure” [16, 17]. A plot of the excess amount versus pressure often exhibits a maximum; beyond this maximum, the bulk gas density increases more quickly with increasing pressure than does the adsorbed density. It is highly plausible that H₂-storage capacity can be elevated in a given container by removing the carbon adsorbent [17].

Naturally, the excess number of adsorbate molecules, N_ex, is defined as the difference between the total amount of adsorbate molecules, N_a, within the volume, V_void, of a measurement cell containing the adsorbent at a given temperature and pressure, and the number of molecules N_g that would have been found in the void volume in the absence of adsorbate–adsorbent interactions; thus,

Obviously,

where ρ_g is the bulk gas density (obtained from the equation of state of the adsorbate) at the same temperature and pressure, and V_void is the sum of the pore volume of the adsorbent, V_pore, and the volume of the section of the cell containing no adsorbent, V_empty, i.e., V_void = V_empty + V_pore. Combining Eqs (1) and (2) gives rise to

Dividing N_ex by the mass of the adsorbent, M, in the cell results in the excess adsorbed density, n_ex, i.e.,

Note that N_ex/M is not equal to the absolute adsorbed density, n_a, since it also includes molecules in the volume, V_empty. In terms of the local density, Eq (4) can be rewritten as

Note that Eq (3) can be expressed in terms of the absolute adsorption, without any reference to the empty volume of the measurement cell, as

where

The quantity of H₂ (in moles) stored in an adsorption-based storage unit of total volume V_sys can be estimated from the following expression:

Thus, the stored volumetric density is

and the stored gravimetric density is

In fact, the gravimetric density of the system is often estimated by multiplying the excess density, n_ex, with the molar mass of the adsorbate (H₂): V_Sys and M_Sys are known only after the system is designed. Nevertheless, such estimation assumes that M_Sys = M and also neglects the contribution of the bulk phase in the pore volume to the total amount of gas in the storage unit as well as the contribution of the stored H₂ to the total mass of the filled storage system.

2.2. Physisorption

The binding energy for physisorption via the van der Waals bonds ranges from 0.04 to 0.12 eV with the largest value corresponding to the regime of low H₂ coverage, where the H₂−H₂ repulsion is insignificant [18, 19]. For conventional physisorption, the gas-adsorption performance of a porous solid is maximized when the pore size is less than a few molecular diameters [20]. As such, the potential fields from the walls of the so-called micropores overlap to produce a stronger interaction than would be possible for adsorption on a semi-infinite plane. If the escaping tendency of the gas is much less than the adsorption potential, the entire micropore may be filled with a condensed adsorbate phase. For the case of H₂, with a kinetic diameter of approximately 2.9 Å, pores would have to be significantly smaller than 40 Å to begin to condense H₂ by a nanocapillary filling mechanism.

Any of the suitable or appropriate models for the excess-adsorption isotherms of microporous adsorbents should contain a minimum number of parameters with physical interpretation. Such models currently available include those based on viral expansion, Langmuir isotherm, self-consistent Ono–Kondo approach, and Dubinin pore-filling approach. The details of these models are available [15, 21].

2.3. Chemisorption

In addition to physisorption, H₂ can be adsorbed onto the carbon adsorbents through chemical adsorption, i.e., chemisorption. The binding energy for chemisorption is between 2 and 3 eV, much stronger than the van der Waals interactions [22, 23]. The usual ab initio methods, e.g., density functional theory (DFT), are well established to study the nature and energetics of chemical bonding. The DFT methods include the local density approximations and the generalized gradient approximations. The former tends to overestimate the binding energies; the latter usually reproduces experimental values to within 0.1 or 0.2 eV, thereby providing more accurate information [24, 25]. In fact, chemisorption of H₂ in graphitic nanocarbons is inapplicable for reversible H₂ storage. It is difficult to control, and the adsorbed H₂ can only be released at high temperatures, e.g., 400 K [26]. Thus, the focus of the research has been on physisorption, which gives rise to a much closer energy range for practical H₂ storage than chemisorption.

3. CARBON ADSORBENTS

Carbon-based adsorbents, including activated carbons, CNFs, and carbon nanotubes (CNTs), have been widely recognized as profoundly effective adsorbents. The attributes of ideal carbon adsorbents for H₂ storage should be such that they have uniform and small micropores in high density, minimal macroporosity, and high thermal conductivity. The first attribute is required for increasing heat of adsorption, which might render possible adsorption at the ambient temperature, while the first two attributes collectively insure that the internal volume of the adsorbent is not wasted. The third attribute is required for an enhanced heat flux, which tends to be relatively large for CNTs. Besides the adsorption capacity, the viability of a H₂-storage device hinges heavily on its capability for releasing the adsorbed H₂ [12]. Note that the properties of carbon-based adsorbents are strongly dependent on those of precursors and on processing conditions, none of which is sufficiently well understood at the molecular level.

3.1. Activated Carbons

The term, activated carbons, defines a group of materials with highly developed internal surface areas and porosities. Activated carbons constitute the only variety of carbon adsorbents that are manufactured in abundance and thus inexpensive.

3.1.1. Structure

Carbonized materials, including activated carbons and chars, have structures and properties similar to those of graphite [27, 28]. The density of graphite is 2.267 g/cm³, and its atoms are bonded together in hexagonal two-dimensional networks. The layers of these atoms are held loosely together by van der Waals force. They are principally in the form of A-B-A-B, i.e., the hexagonal form; a small fraction (<10%) in natural graphite stacks in the form of A-B-C-A-B-C, i.e., the rhombohedral form. Note that the latter can convert irreversibly to the former at approximately 2400 K; see Fig. 7.1.

The structure of activated carbons is less perfectly ordered than that of graphite, generally representing an intermediate between the organic precursor and the single graphite crystal [30]. Activated carbons are either graphitic or nongraphitic depending on the degree of crystallographic ordering. The former includes all varieties of carbonized materials in which elemental carbon is in the allotropic form of graphite, irrespective of the presence of structural defects. The latter includes all varieties of carbonized materials in which elemental carbon has two-dimensional long-rangeordered carbon atoms in planar hexagonal networks; nevertheless, in the third dimension, apart from more or less parallel stacking, no crystallographic order can be detected [28].

Figure 7.1 Two structures of graphite: (A) hexagonal unit cell and (B) rhombohedral unit cell [29].

The nongraphitic carbons can be further divided into the graphitizable and nongraphitizable carbons. Graphitizable carbons can be converted into graphitic carbons by heat treatment; the degree of graphitization depends on the temperature of heat treatment and the time allowed to anneal the structure [30]. Such carbons are essentially cokes resulting from the carbonization of pitch materials derived from coal and petroleum [28]. On the other hand, a nongraphitizable carbon cannot be transformed into a graphitic carbon solely by heat treatment under inert conditions even up to a temperature of 3500 K [30]. The majority of such carbons are generated from pyrolysing cellulose and lignin within wood and nut shells or from nonfusing coals, e.g., anthracite and subbituminous coals. These materials lose molecules of water (H₂O), CO₂, and organics of low molecular weights on pyrolysis, with only the cross-linking and basic back-bone structure remaining. The voids created by the loss of small molecules, including H₂ and CO₂, at higher temperatures (500 ~ 1000 °C) constitute the pores, 0.5–2.0 nm in diameter and high surface areas of about or greater than 1000 m²/g, within the carbons [27, 28, 31].

A distinct difference in the structures of graphitizable and nongraphitizable carbons has been identified by Byne and Marsh [30]. According to them, all carbonized materials, including activated carbons, have structures comprising imperfect sections of graphitic lamellae. It is plausible that these lamellae can be sliced into sections, 1.0 nm–500 µm in thickness. The smaller sections of 1.0 ~ 10 nm thick can be crumpled and bonded together to create a three-dimensional network. Such a structure, with defects and irregularities, approximates that of porous activated carbons. Because of the crumpled, defective nature of the constituent lamellarbased molecules (LCMs), their packing density is low, approximately 1.0 g/cm³, and it is the spaces or voids between the LCM that constitute porosity. The graphitizable cokes, however, comprise LCMs that are larger, contain fewer defects, and are less crumpled than those in the porous carbons. Accordingly, they stack parallel to each other over relatively larger dimensions of micrometers, thus increasing their densities, which approach 2.0 g/cm³ (the density of graphite is 2.2 g/cm³); they have little or no pores with dimensions comparable to those of molecules, i.e., nanometers [30]. Figure 7.2 shows a schematic representation of the structures of nongraphitizable and graphitizable carbons [27].

Figure 7.2 Schematic representation of activated-carbon structures: (A) nongraphitizing and (B) graphitizing [27].

The structure of an activated-carbon particle is formed by pores wide ranging in size. According to the classification of the International Union of Pure and Applied Chemistry (IUPAC), the pores of activated carbons can be divided into the following categories based on their diameters; see Fig. 7.3. The diameters of “macropores” for admission and diffusion are greater than 50 nm, those of “mesopores” range between 2.0 and 50 nm, and those of “micropores” are less than 2.0 nm. Pores having diameters of the same order of magnitude as the diameters of molecules, i.e., <0.8 nm, are called “submicropores.”

The macropores serve as transport conduits enabling the molecules of the adsorbate to reach smaller pores situated in the interior of a carbon particle. Thus, although the macropores do not contribute appreciably to the total surface area, they significantly affect the admission rate of adsorbates into the carbon. The mesopores, branching off from the macropores, serve as secondary passages to the micropores for the adsorbate. In mesopores, capillary condensation may take place with the formation of a meniscus of the liquid adsorbate. This phenomenon usually produces the hysteresis loop on adsorption isotherms [27]. The mesopore volume is usually between 0.02 and 0.1 cc/g. In most activated carbons, the surface area contributed by mesopores is relatively low but not insignificant. The micropores contribute most to the total surface area of activated carbons, thereby providing a high adsorptive capacity for molecules of small dimensions; the majority of the adsorption takes place within them. Their small size prevents capillary condensation. They are occupied by the adsorbate at low relative pressures.

Figure 7.3 Schematic representation of the various types of pores in a particle of carbonaceous adsorbent with different diameters: macropores, >50 nm; mesopores, between 2 and 50 nm; micropores, between 0.8 and 2 nm; and submicropores, <0.8 nm [32].

3.1.2. Preparation

Activated carbons have been prepared from a number of precursors, such as lignite, coconut shell, wood, and grains, through one-stage or two-stage processes [33–45]. In the one-stage process, activated carbons are generated directly from raw materials via physical or chemical activation. Physical activation is carried out most frequently by burning off some of the raw carbon in an environment of oxidizing gas to create micropores. The usual choices of oxidizing gas are steam, CO₂, air, or their mixtures. Activation normally takes place at temperatures between 700 and 1000 °C in steam and CO₂ and lower temperatures in air. In chemical activation, a carbon precursor is impregnated with a dehydrating agent before heated between 500 and 800 °C. The most widely used dehydrating agents include zinc chloride, potassium sulfide, potassium thiocyanate, phosphoric acid, sulfuric acid, hydroxides of the alkali metals, magnesium chloride, and calcium chloride. The resultant activated carbons from the one-stage process have small surface areas and high mesoporosities.

The two-stage process comprises carbonization followed by physical or chemical activation. Carbonization is performed by pyrolyzing raw materials at a temperature less than 700 °C to produce char. The char obtained is physically activated through oxidation in the presence of an oxidizing gas or chemically activated through pyrolysis in an inert gas after mixing with the dehydrating agent. The resultant activated carbons usually have high surface areas and high microporosities.

3.1.3. Hydrogen Adsorption

It is entirely plausible that microporous activated carbons will be effective for H₂ storage in light of their high porosities and low cost. Nevertheless, in one of the first investigations [46], the maximum excess H₂ amounting to merely 20.2-g H₂/kg activated carbon was reported at 25 atm and 76 K, corresponding to a gravimetric storage density of approximately 2.0 wt%, which is far from reaching the target H₂ storage density set by DOE [6].

Intensive efforts have been made to increase the H₂-adsorption capacity of microporous activated carbons. One attempt is through metal modification of activated carbons: H₂ is well known to be inordinately adsorptive to adsorbents doped with a transition metal, i.e., a metal in group VIII of the Periodic Table [47–52]. Upon contact with a transition metal in elemental form, H₂ molecules are preferentially adsorbed onto the metal’s surface; some of these H₂ molecules dissociate into H atoms, which are more efficient at filling the available active sites on the activated carbons than H₂ molecules [47, 48, 50]. The effect of this phenomenon, known as H₂ spillover, on catalytic reactions has been extensively documented, although the chemical form of H₂ that spills over is still uncertain [52]. According to Schwarz [50], the transition metals are themselves specifically active chemically. Moreover, they do not affect the hydroxyl, aldehyde, carbonyl, or other active groups on the surface of activated carbons. Such active groups possibly play a significant role in the adsorption of the dissociated H. Apparently, the minute amount of a transition metal can appreciably enhance the adsorbability of H₂ by activated carbons. Nevertheless, the metal-impregnated activated carbons adsorb only approximately 4.8 wt% H₂ at a temperature of 87 K and a pressure of 59 atm [51]. A simple calculation by adopting the parameters specified by Schwarz [51] has revealed that 4.8 wt%, or 16.5 kg/m³, of H₂ could be stored in the tank at the same pressure and temperature without the activated carbons [53].

X-ray diffraction of the nanocrystalline graphite resulting from mechanical ball milling under an H₂ atmosphere has indicated that the graphite inter-layer spacing expands accompanied by the increasing disappearance of long-range order with milling time [54]. Up to 7.4 wt% of H₂ adsorption has been reported after milling for 80 h. Neutron-diffraction measurements have revealed that both molecular H₂ and covalently bonded H₂ are present [54]. Although the carbon adsorbents fabricated via intensive milling appear quite promising for H₂ storage, the adsorbed H₂ is extremely difficult to be desorbed from them.

Inclusion of activated carbons in a storage tank can increase the overall H₂ energy storage density under certain pressure and temperature. For any given activated carbons, this increase is maximized at low temperatures and high pressures, corresponding to the situation where the adsorbent surface is populated, but the density in the gas phase above the adsorbent is not appreciable [14]. Although the overall storage capacity of a tank is maximized at the high-pressure and low-temperature conditions, the effectiveness of the adsorbent is minimized under these conditions [51, 55]. The magnitude of the maximum gain depends on the surface area and microporous volume. Ideal adsorbents should be such that they are highly microporous with relatively high density, i.e., low in macroporous volume. In a recent study, various carbon adsorbents, including activated carbons, carbon black, carbon aerogels, and carbon molecular sieves, were tested to determine if they could augment the capacities of compressed H₂ gas storage systems at the ambient temperature (298 K), acetone and dry-ice temperature (190 K), and liquid-nitrogen temperature (80 K) [17]. The study has revealed that at pressures typical of vehicular compressed H₂-storage systems (approximately 200 bars), only one of 10 carbon sorbents tested can marginally augment the capacity of the storage vessel at 190 and 298 K; such augmentation is nill at 80 K [17]. Relatively little has been published on adsorbents of the activated-carbon type, which are capable of increasing the storage capacities of a high-pressure tank at 298 K [56, 57]. In fact, the H₂-adsorption capacity of activated carbons is low at room temperature and moderate pressure. For instance, according to Jin et al. [58], H₂ adsorption is less than 1 wt% at 100 bars and 298 K, despite the surface area of 2800 m²/g. Although H₂ is mainly adsorbed by physical interaction, H₂ adsorption can be promoted by incorporating surface groups and other reactive sites into the activated carbons. Such measures, however, tend to operate mostly via chemisorption and are ineffective in enhancing the reversible storage of H₂.

As mentioned earlier, activated carbons inevitably contain macropores, mesopores, and micropores. In certain types of activated carbons, approximately 50% of the total pore volume comprises macropores, which hardly contribute to H₂ storage. The poor H₂-adsorption capacity of activated carbons can be largely attributed to the nonuniformity of their pores. The development of the specifically designed carbon adsorbents with uniform pores, therefore, has been extensively pursued for H₂ storage [53].

3.2. Carbon Nanofibers

In an attempt to elevate H₂ storage densities beyond those attainable by activated carbons, various carbon nanostructures have been investigated. Among them are CNFs.

3.2.1. Structure

CNFs, or graphite carbon nanofibers, are usually nongraphitic. They consist of graphene planes arranged in platelet stacks, either in parallel or in angled arrangements, resulting in a conical fishbone (herringbone) structure. Three distinct structures exist: tubular (90°), platelet (~ 0°), and herringbone (45°), where the angle in parentheses indicates the direction of the nanofiber axis relative to the vector normal to the graphene sheets. The spacing between graphite layers in each case is the same as that found in conventional graphitic carbon, i.e., about 3.4 Å. From the physical point of view, CNFs vary from 5 to 100 µ in length and from 5 to 100 nm in diameter. One distinct feature of CNFs is the abundance of open edges, thereby favoring H₂ sorption [59]. Figure 7.4 shows the SEM photographs of CNFs prepared by decomposing a mixture of C₂H₄ and H₂ with a ratio of 1 to 4 over Ni_0.5Cu_0.5 alloy powder [60].

Figure 7.4 SEM photographs of CNFs [60].

3.2.2. Preparation

CNFs have been catalytically synthesized. Specifically, they can be produced by decomposing mixtures of hydrocarbon and H₂ on selected metal and alloy catalysts at temperatures ranging from 400 to 800 °C [61]. Figure 7.5 shows a schematic diagram of catalytically grown CNFs [62]. When a hydrocarbon is adsorbed on a metal surface and conditions exist favoring the scission of a carbon–carbon bond in the molecules, the resulting atomic species may dissolve in the catalyst particle, diffuse to the rear faces, and ultimately precipitate at the interface to form CNFs. The degree of crystalline perfection of the deposited fiber depends on three main factors. They are the chemical nature of the catalyst particle, composition of the reactant gas, and reaction temperature. Surface science studies have revealed that certain faces favor precipitation of carbon in the form of graphite, whereas less ordered carbon will be deposited from other faces [63–65]. By judicious choice of the catalyst, the ratio of the hydrocarbon to H₂ in the reactant mixture, and reaction conditions, it is possible to tailor the morphological characteristics, the degree of crystallinity, and the orientation of the precipitated graphite crystallites relative to the fiber axis [62].

Figure 7.5 Schematic diagram of catalytically grown CNFs: (A) metal surface, (B) metal catalyst particle, (C) interface, and (D) deposited CNFs [62].

3.2.3. Hydrogen Adsorption

The H₂ adsorption on CNFs has been studied extensively at temperatures ranging from 77 to 300 K. The H₂-storage capacities exceeding 60 wt%, i.e., 2-g H₂/g carbon, were achieved for certain CNFs at the ambient temperature under the pressure between 50 and 120 bars [66]. Nevertheless, these results have never been replicated by others. For instance, H₂ adsorption of only approximately 0.08 wt% was attained with similarly fabricated CNFs under similar experimental conditions [67]. Another maximum H₂-storage capacity reported was 1.52 wt% at the ambient temperature under the pressure of 125 atm [68].

To rationalize the results of Chambers and his collaborators [69], they have stipulated that the presence of water vapor would have significantly deteriorated the CNFs, thereby resulting in the expansion of their graphitic layers, which, in turn, gives rise to the CNFs’ high H₂-adsorption capacities; unfortunately, they were unable to probe in situ such an expansion. X-ray data of a pristine sample and the same sample following both adsorption and desorption steps showed the lattice expansion of only 0.007 nm, which is considerably less than 0.03 nm observed for H₂ adsorbed in potassium-intercalated graphite [70].

To enhance the H₂-storage capacities of CNFs, they were grown on a Pd-based catalyst system. Pd has a high potential to dissociate an H₂ molecule into H atoms to be adsorbed at the Pb–C interface [71]. The H₂ uptake experiments were performed volumetrically in a Sievert-type installation by varying the mole ratios of Pd to C from 0.05 to 0.9. The quantity of desorbed H₂ was between 0.04 and 0.33 wt% under the pressure ranging from 1 to 100 bars. This study has indicated that a direct correlation exists between the ratio of Pd and C and the quantity of H₂ absorbed. Unfortunately, a saturation value of approximately 1.5 wt% reached at a high ratio of about 1:1 of Pd to C is far from meeting the energy-density criterion set by DOE [6].

3.3. Single-Walled Carbon Nanotubes

CNTs are allotropes of carbon, representing highly ordered, hexagonal-packed nanostructures that can have a length-to-diameter ratio of up to 28 000 000:1 [72]. CNTs comprise SWCNTs and MWCNTs; the former is covered in what follows.

3.3.1. Structure

An individual SWCNT can be visualized as a single graphene sheet rolled up in an elongated and seamless tube with a diameter in the order of 1 nm and a length in hundreds of micrometers. The pattern of the graphemesheet wrapping is represented by a pair of indices (n,m) termed the chiral vector. The integers, n and m, denote the numbers of unit vectors along two directions in the honeycomb crystal lattice of graphene. If m = 0, the nanotubes are called “zigzag.” If n = m, they are called “armchair.” Otherwise, they are called “chiral”; see Fig. 7.6 [7]. During synthesis, SWCNTs tend to self-organize and form thick nanoropes, consisting of bundles of hundreds of units, via van der Waals interactions, generally in hexagonl arrangement [73].

3.3.2. Preparation

SWCNTs can be synthesized through an electric arc technique. It involves vaporizing carbon in an electric arc discharge or laser vaporization in the presence of cobalt (Co) or nickel (Ni) catalyst. SWCNTs were first prepared by coevaporating a Co catalyst and graphite in an electric arc. The resultant fibers typically consisted of 7–14 bundled, highly impure SWCNTs with 10–15 Å in diameter [74, 75].

Figure 7.6 Structural relationship between graphene sheet and SWCNT: arrows point to two alternative directions of rolling; circles inside rings reflect on aromatic property of carbon rings and delocalization of resonant electrons inside them [7].

SWCNTs can be produced at a much higher yield at approximately 20–50 wt% by laser vaporization [76]. Crystalline ropes containing hundreds of individual SWCNTs, microns in length, were easily obtained [73]. The soot generated via laser vaporization also contained metal–nanoparticle catalysts and various carbon components. As a result, various purification techniques involving filed flow filtration, acid washing, and annealing steps have been developed to prepare SWCNTs of much higher quality [77–82]. In fact, SWCNTs with purity greater than 98 wt% can be fabricated if micron-sized graphite particles are not generated in a vaporization regime during synthesis [83].

3.3.3. Hydrogen Adsorption

The adsorption processes on bundled SWCNTs proceed at four different types of adsorption sites: the internal sites within the individual SWCNTs, the interstices among the nanotubes within the bundle, the grooves between pairs of SWCNTs at their surfaces, and the external surface sites [15]. It is expected that the cylindrical geometry will deepen adsorbate–adsorbent interaction potential wells inside SWCNTs with small diameters. Nevertheless, the H₂ entering into the interior of SWCNTs is not facilitated; instead, H₂ is preferentially adsorbed on the exterior wall at the interstitial sites in a bundle of SWCNTs due to the overlap of the molecular force fields of an individual SWCNT. In fact, the small pore volume of SWCNTs associated with such sites and the small diameter of each SWCNT reduce the overall density of adsorbed H₂ [15]. Thus, the theoretical maximum amount of adsorbed H₂ is merely 3.0 wt% for SWCNTs with a specific surface area of 1315 m²/g at 77 K [10].

SWCNTs were first proposed as H₂-storage media by Dillon et al. [84]. In their study, the soot contained only approximately 0.1–0.2 wt% of SWCNT bundles with 7–14 nanotubes of approximately 12 Å in diameter. Physisorption was identified as the predominant mechanism for H₂ adsorption; the heat of adsorption was estimated to be 19.6 kJ/mol. The study has demonstrated an enhanced interaction between H₂ and SWCNTs compared to planar graphite; for the latter, the heat of adsorption has been reported to be only approximately 4 kJ/mol [18].

The first study of H₂ adsorption on purified laser-generated SWCNTs attained H₂ adsorption of 8 wt% at approximately 40 atm and 80 K [85]; nevertheless, this study failed to elevate high adsorption capacities at 300 K and pressures below 1 atm. It was determined that H₂ was physisorbed on the exposed surfaces of the tubes [85]. In another study, the quantum rotation of H₂ adsorbed on laser-generated SWCNTs was observed via inelastic neutron scattering, demonstrating that H₂ was physisorbed at 25 K under 11 MPa [86].

High H₂-storage capacities on a total sample weight basis were achieved on SWCNTs with a large mean pore diameter of about 1.85 nm produced by a semicontinuous arc-discharge method; the yield from this method tends to be high [87]. The purity of the nanotube soot was estimated to be approximately 50–60%. A sample, first soaked in HCl and subsequently heat-treated in vacuum, was shown to adsorb 4.2 wt% of H₂ at room temperature and 10 MPa; approximately 80% of the adsorbed H₂ could also be released at room temperature. Note that high-capacity adsorption at room temperature was first demonstrated for arc-generated SWCNTs and not for laser-produced nanotubes. This difference might be attributable to a much smaller number of ends or defects in the laser-produced tubes and/or to an enhancement in their stability during the procedures for opening or cutting. A newly developed high-power ultrasonic cutting procedure incorporating a TiAl_0.1V_0.04 alloy has been applied to purified, laser-generated SWCNT materials, thus rendering high-capacity H₂ adsorption possible under ambient conditions [88]. The maximum adsorption capacity of the resultant SWCNTs is approximately 7 wt%, and H₂ adsorption occurs in two separate sites. Approximately 2.5 wt% of the H₂ evolves at 300 K, while the remainder desorbs at 475–850 K. In addition to cutting, the procedure generates an alloy of composition TiAl_0.1V_0.04 due to decomposition of the ultrasonic probe. The presence of the alloy might stimulate H₂ adsorption and desorption. Several control experiments have revealed that the observed uptake is not solely due to the presence of the alloy; some of the experimental results indicate that partial electron transfer is responsible for the stability of the adsorbed H₂ [88].

Electrochemical H₂ storage has also been demonstrated for carbon SWCNTs. Electrodes were fabricated by mixing arc-generated SWCNT soot, containing a few percent of SWCNTs whose diameters range from 0.7 to 1.2 nm, with either copper or gold as a compacting powder in a ratio of 1:4. The kinetics of the SWCNT electrode was found to be relatively poor. The H₂-storage capacity at low discharge currents, however, was as high as 110 mAh/g, corresponding to approximately 0.39 wt% of H₂ [89]. On the assumption that the stored H₂ is totally contained in the small fraction of SWCNTs in the sample, these results indicate that electrochemical H₂ storage on highly pure SWCNTs should not be ignored. In a more recent study [90], SWCNT-composite electrodes, fabricated by mixing SWCNTs, Ni powders, and an organic polytetrafluoroethylene in a ratio of 40:50:10, demonstrated charge/discharge capacities of 160 mAh/g. Unfortunately, no information was provided on either the synthesis or the purity of the SWCNT material.

One approach to enhance H₂ storage resorts to a supported metallic catalyst to promote the dissociation of H₂ molecules into H atoms, which can bind directly with the carbon nanostructures. For instance, the enhancement by a factor of 1.6 was observed for the palladium-catalystdoped SWCNTs [91].

Nondissociative enhancement has been achieved by modifying the nanostructure to intensify average adsorbate–adsorbent interaction [92, 93]. In some of the approaches, SWCNTs are doped with a metal through the formation of Kubas complexes. This leads to enhanced nondissociatively binding of H₂ molecules to transition metal complexes; such binding is much stronger than van der Waals forces [93]. Calculations based on ab initio quantum chemistry have indicated that a titanium atom embedded SWCNT can store about 7 wt% of H₂ at high titanium coverage under ambient conditions [94]. Such enhancement hinges on the existence of a synthesis path leading to a meta-stable unclustered titanium-covered nanostructure separated from the Ti-clustered configuration by a large energy barrier. It has been speculated that metals, such as Sc, Cr, Pd, and Pt, also enhance H₂ adsorption [92].

3.4. Multiwalled Carbon Nanotubes

MWCNTs constitute another major class of CNTs. They are macromolecular arrangements of two or more interpenetrating SWCNTs.

3.4.1. Structure

MWCNTs consist of layers of nested concentric cylinders of graphite with a hollow center; see Fig. 7.7. The spacing between any pair of adjacent concentric cylinders is similar to the interplanar spacing in graphite, and the number of shells varies from 2 up to about 50. MWCNTs have inner and outer diameters that are typically 2–10 and 15–30 nm, respectively, and are generally microns in length. Large bundles of MWCNTs with diameters up to 200 µm and very long individual tubes, whose length exceeds 2 mm [95], have been observed.

Figure 7.7 Structure of MWCNTs [96].

3.4.2. Preparation

Similar to SWCNTs, MWCNTs can be produced via arc-generation; in the generation process, the carbon contained in the negative electrode sublimates due to the high discharge temperature [97–99]. In fact, MWCNTs were observed in the carbon soot of graphite electrodes during an arc discharge with a current of merely 100 amps, which was intended to produce fullerenes [97]. Arc-generated MWCNTs have been purified with air oxidation. To exhaustively remove the nanocrystalline graphite particles present as impurities, however, entails oxidizing 99% of the total material [100]. It is extremely difficult, if not impossible, to completely remove the impurities from arc-generated MWCNTs. For instance, the nanocrystalline graphite again could not be removed even when oxidized with permanganate [101].

MWCNTs have also been produced by the catalytic decomposition of hydrocarbons, such as benzene pyrolysis [102] and acetylene decomposition with metal catalysts [101, 103]. The crystallinity and stability of the tube caps may be compromised when MWCNTs are produced by the catalytic decomposition of hydrocarbons [101]. CNTs produced by the catalytic decomposition of acetylene over a Co-incorporated zeolite were purified with hydrofluoric acid to remove the catalyst support followed by a permanganate or air oxidation to remove amorphous carbon impurities [101].

3.4.3. Hydrogen Adsorption

H₂-storage capacities were reported to be appreciable for alkali-metaldoped MWCNTs formed by the catalytic decomposition of CH₄ [104]. Li and K were incorporated through solid-state reactions with the metal carbonates or nitrates. H₂ desorption and adsorption were measured by thermogravimetric analysis (TGA) and temperature-programmed desorption. The H₂ uptake was 20 wt% for Li-doped MWCNTs at 653 K and 14 wt% for K-doped MWCNTs at room temperature; the latter was reported to combust upon exposure to air. The H₂ adsorption was believed to proceed by a dissociative mechanism. An infrared spectrum indicated the presence of both Li–H and C–H species [104]; nevertheless, the results have not been confirmed. For instance, in an investigation, Li-doped MWCNTs formed under identical conditions only exhibited a weight increase of 12 wt% when exposed to “wet H₂” and only 2.5 wt% when exposed to “dry H₂” [105]. An infrared spectrum of LiOH:H₂O was strikingly similar to the spectrum acquired by Chen et al. [104]. When K-doped nanotubes were exposed to “dry H₂,” a weight increase of only 1.8% was observed [105]. In another investigation, the observation of large H₂ adsorption on Li-doped MWCNTs, similar to those observed by Chen et al. [104], was attributed to the presence of water as an impurity in the TGA atmosphere. In fact, H₂ adsorption by the Li-doped MWCNTs was not evident [106].

Wu et al. [107] studied H₂ adsorption on MWCNTs synthesized by the catalytic decomposition of CO or CH₄ on powdered La₂O₃ catalysts. The CO-generated tubes comprised concentric cylinders, while the CH₄-generated tubes contained graphite layers that were tilted with respect to the tube axis, thus forming cones. In both cases, the catalyst was removed by stirring in dilute nitric acid. The purified nanotubes were then annealed to 1100 °C in vacuum to increase their crystallinity. TGA in flowing H₂ revealed that the CO-generated tubes were capable of adsorbing only a small quantity of H₂ (0.25 wt%) when the sample was cooled at a temperature between 200 °C and room temperature. Another study indicated that MWCNTs could also be charged with H₂ by electrochemical methods [89]. Materials deployed were synthesized by an arc process and contained 10–40 wt% of MWCNTs with diameters ranging from 2 to 15 nm. Stable electrodes were formed by pressing the MWCNT material with palladium powder in a 1:4 ratio. The equilibrium curve of the MWCNT/palladium electrode indicated the occurrence of two separate electrochemical reactions, one of which was not observed for pure palladium electrodes. Although the overall capacity of the palladium/MWCNT electrode was less than that anticipated for pure palladium, the study showed that electrochemical H₂ storage in MWCNTs might be possible. One of the most recent studies of H₂ adsorption on MWCNTs, grown on a Fe:Co:CaCO₃ catalytic system and purified by acid cleaning and air oxidation, gave rise to H₂ uptake value of 0.1–0.2 wt% [71].

To summarize, CNTs have been intensively investigated as H₂-storage media; however, their viability remains uncertain. It appears that the CNTs fabricated to date are far from meeting the requirements of high-capacity, room-temperature H₂ storage for automotive application. Furthermore, the cost of scaling-up the process for manufacturing CNTs is formidable; it is reported that “…making a 20-gal hydrogen fuel tank that uses carbon nanotubes could cost $5.5 million; …” [108]. Economically, therefore, disordered activated carbons might be a more desirable choice than CNTs.

4. CONCLUDING REMARKS

The onboard storage of H₂ is obviously one of the most, if not the most, critical issues in developing an H₂-based energy system. As such, it has lately been a major focus of academic and industrial researches. The storage of H₂ in carbon adsorbents, especially in carbon nanostructures, is still at its infancy and thus is far from mature for commercial application. Many of complexities involved in its large-scale production remain unresolved. They include scale-up, system design, optimization of operating-conditions, and cost estimation.

Adsorption is exothermic, while desorption is endothermic. Hence, the uptake of H₂ into a storage vessel during charging will heat up the bed, thereby reducing its capacity, and the bed will cool on discharging, thus decreasing the deliverable quantity of H₂. Moreover, any impurities in the H₂ gas will accumulate in the storage vessel over repeated charge/discharge cycles; this also will reduce the capacity for H₂. Naturally, it is expected that further extensive research and development need to be undertaken in the near or foreseeable future for the ultimate deployment of carbon adsorbents in the practical onboard storage systems for H₂.

ACKNOWLEDGMENTS

This is contribution No. 10-073-B, Department of Chemical Engineering, Kansas Agricultural Experiment Station, Kansas State University (Manhattan, KS), from which the second and third authors received financial support.

References

[1] Veziroglu TN, Barbir F. Hydrogen: the wonder fuel. Int J Hydrogen Energy 1992; 17:391–401.

[2] Choi B, Nam G, Choi D, Lee B, Kim S, Lee C, et al. Adsorption and regeneration dynamic characteristic of methane and hydrogen binary system. Korean J Chem Eng 1994;21:821–8.

[3] Das LM. Onboard hydrogen storage systems for automotive application. Int J Hydrogen Energy 1996;21:789–800.

[4] Veziroglu TN. Hydrogen energy system as a permanent solution to global energy environmental problems. Chem Ind 1999;53:383–93.

[5] Walker G. Hydrogen storage technologies. In: Walker G, editor. Solid-state hydrogen storage: materials and chemistry (Woodhead Publishing in Materials). New York: CRC Press; 2008. p. 3–17.

[6] DOE. Targets for onboard hydrogen storage systems for light-duty vehicles, http://www1.eere.energy.gov/hydrogenandfuelcells/storage/
current_technology.html, Sept. 2009.

[7] Varin RA, Czujko T, Wronski ZS. Nanomaterials for solid state hydrogen storage (fuel cells and hydrogen energy). New York: Springer; 2009.

[8] Hashimoto K, Habazaki H, Yamasaki M, Meguro S, Sasaki T, Katagiri H, et al. Advanced materials for global carbon dioxide recycling. Mater Sci Eng A 2001; 304–306:88–96.

[9] Bockris JO’M. The origin of ideas on a hydrogen economy and its solution to the decay of the environment. Proceedings – Electrochemical Society. 2000–20 (Global Climate Change). 2001. p. 1–24.

[10] Sandi G. Hydrogen storage and its limitations. Electrochem Soc Interface 2004;13:40–5.

[11] Gutowski M, Autrey T. Hydrogen gets on board. Chem World 2006;3:44–8.

[12] Jacoby M., Filling up with hydrogen. Chem Eng News 2005;83:42–7.

[13] Satyapal S, Petrovic J, Thomas G. Gassing up with hydrogen. Sci Am 2007;296:80–7.

[14] Chahine R, Bénard P. Adsorption storage of gaseous hydrogen at cryogenic temperature. In: Kittel P, editor. Advance in cryogenic engineering, vol. 43. New York: Plenum Press; 1998. p. 1257–64.

[15] Benard P, Chahine R. Carbon nanostructures for hydrogen storage. In: Walker G, editor. Solid-state hydrogen storage: materials and chemistry (Woodhead Publishing in Materials). New York: CRC Press; 2008. p. 261–87.

[16] Coolidge AS. Adsorption at high pressures. I. J Am Chem Soc 1934;56:554–61.

[17] Hynek S, Fuller W, Bentley J. Hydrogen storage by carbon sorption. Int J Hydrogen Energy 1997;22:601–10.

[18] Pace EL, Siebert AR. Heat of adsorption of parahydrogen and orthodeuterium on Graphon. J Phys Chem 1959;63:1398–400.

[19] Okamoto Y, Miyamoto Y. Ab initio Investigation of Physisorption of Molecular Hydrogen on Planar and Curved Graphens. J Phys Chem B 2001;105:3470.

[20] Gregg SJ, Sing KSW. Adsorption, surface and porosity. London: Academic; 1982.

[21] Ruthven DM. Principles of adsorption and adsorption processes. New York: John Wiley & Sons; 1984.

[22] Lee SM, Lee YH. Hydrogen storage in single-walled carbon nanotubes. Appl Phys Lett 2000;76:2877.

[23] Gulseren O, Yildirim T, Ciraci S. Tunable adsorption on carbon nanotubes. Phys Rev Lett 2001;87:116–802.

[24] Kong XJ, Chan CT, Ho KM, Ye YY. Cohesive Properties of Crystalline Solids by the General Gradient Approximation. Phys Rev B: Condens Matter 1990;42:9357–64.

[25] Perdew JP, Burke K, Ernzerhof M. Generalized gradient approximation made simple. Phys Rev Lett 1996;77:3865.

[26] Lee SM, An KH, Lee YH, Seifert G, Frauenheim T. A hydrogen storage mechanism in single-walled carbon nanotubes. J Am Chem Soc 2001;123:5059–63.

[27] Smisek M, Cerny S. Activated carbon: manufacture, properties, and applications. Amsterdam: Elsevier; 1970.

[28] Marsh H. Structure in carbons. In: Figueiredo JL, Moulijn JA, editors. Carbon and coal gasifi cation. Dordrecht: Martinus Nijhoff Publishes; 1986. p. 27–57.

[29] Edwards IAS. Structure in carbon and carbon forms. In: Marsh H, editor. Introduction to carbon science. Oxford: Butterworth-Heinemann; 1989. p. 2–36.

[30] Byrne JF, Marsh H. Introduction Overviews. In: Patrick JW, editor. Porosity in carbons: characterization and applications. London: Edwards Arnold; 1995. p. 1–48.

[31] Capelle A. Adsorption of Cationic Copolymers from Dilute Aqueous Solutions on Powdered Activated Carbon. In Capelle A, De Vooys F, editors. Activated Carbon. A Fascinating Material: Some Thoughts on Activated Carbon. Amersfoort, The Netherlands: Norit, 1983. p. 191–204.

[32] Jüntgen H. New applications for carbonaceous adsorbents. Carbon 1977;15:273–83.

[33] Spiro CL, McKee DW, Kosky PG, Lamby EJ. Catalytic CO₂-gasiflcation of graphite versus coal char. Fuel 1983;62:180–4.

[34] Laine J, Calafat A, Labady M. Preparation and characterization of activated carbons from coconut shell impregnated with phosphoric acid. Carbon 1989;27:191–5.

[35] Hall RC, Holmes RJ. The preparation and properties of some activated carbons modified by treatment with phosgene or chlorine. Carbon 1992;30:173–6.

[36] Solum MS, Pugmire RJ, Jagtoyen M, Debyshire F. Evolution of carbon structure in chemically activated wood. Carbon 1995;33:1247–54.

[37] Toles C, Rimmer S, Hower JC. Production of activated carbons from a Washington lignite using phosphoric acid activation. Carbon 1996;34:1419–26.

[38] Venkatraman A, Walawender WP, Fan LT. Production and characterization of activated carbon from cereal grains. Fuel chemistry division, ACS 1996;41:260–4.

[39] Wu C-C, Walawender WP, Fan LT. Chemical agents for production of activated carbons from extrusion cooked grain products. Extended abstracts and program, 23rd biennial conference on carbon. Penn State University; 1997. p. 116–7.

[40] Diao Y, Walawender WP, Fan LT. Production of activated carbons from wheat using phosphoric acid activation. Adv Environ Res 1999;3:333–42.

[41] Diao Y, Walawender WP, Fan LT. Activated carbons prepared from phosphoric acid activation of grain sorghum. Bioresour Technol 2002;81:45–52.

[42] Zhang T. Preparation and characterization of carbon molecular sieves and activated carbon. Ph.D. Dissertation, Kansas State University; 2004.

[43] Zhang T, Walawender WP, Fan LT, Fan M, Daugaard D, Brown RC. Preparation of activated carbon from forest and agricultural residues through CO₂ activation. Chem Eng J 2004;105:53–9.

[44] Zhang T, Walawender WP, Fan LT. Preparation of carbon molecular sieves by carbon deposition from methane. Bioresour Technol 2005;96:1929–35.

[45] Zhang T, Walawender WP, Fan LT. Enhancing the microporosities of activated carbons. Sep Sci Technol 2005;44:247–9.

[46] Kidnay AJ, Hiza MJ. High pressure adsorption isotherms of neon, hydrogen, and helium at 76 K. Adv Cryog Eng 1967;12:730–40.

[47] Sinfelt JH, Yates DJC. Catalytic hydrogenolysis of ethane over the noble metals group VIII. J Catal 1967;8:82–90.

[48] Kubicka H. The specific activity of technetium, rhenium, ruthenium, platinum, and palladium in catalytic reactions of benzene with hydrogen. J Catal 1968;12:223–37.

[49] Sermon PA, Bond GC. Hydrogen spillover. Cat Rev 1973;8:211–39.

[50] Schwarz JA. Method and apparatus for cold storage of hydrogen. EU Patent 0230384, 1987.

[51] Schwarz JA, Amankwah KAG. The effect of impurities on hydrogen storage capacity on activated carbons at refrigeration temperature. World hydrogen energy conference, 8th. Honolulu and Wukoloa, Hawaii, U.S.A., 22–27 July 1990. p. 973–81.

[52] Kasaini H, Goto M, Furusaki S. Selective separation of Pd(II), Rh(III), and Ru(III) ions from a mixed chloride solution using activated carbon pellets. Sep Sci Technol 2000;35:1307–27.

[53] Dillon AC, Heben MJ. Hydrogen storage using carbon adsorbents: past, present and future. Appl Phys A 2001;72:133–42.

[54] Orimo S, Majer G, Fukunaga T, Zuttel A, Schlapbach L, Fujii H. Hydrogen in the mechanically prepared nanostructured graphite. Appl Phys Lett 1999;75:3093–95.

[55] Carpetis C, Peschka W. A study on hydtrogen storage by use of cryoadsorbents. Int J Hydrogen Energy 1980;5:539–54.

[56] Hynek S, Fuller W, Bentley J, McCullough J. Hydrogen storage by carbon sorption. Hydrogen energy progress: world hydrogen energy conference. Coral Gables, FL; 1994. p. 985–1000.

[57] Chahine R, Bose TK. Low-pressure adsorption storage of hydrogen. Int J Hydrogen Energy 1993;19:161–4.

[58] Jin H, Lee YS, Hong I. Hydrogen adsorption characteristics of activated carbon. Catal Today 2007;120:399–406.

[59] Schur DV, Tarasov BP, Yu. Zaginaichenko S, Pishuk VK, Veziroglu TN, Shul’ga YM, et al. The prospects for using carbon nanomaterials as hydrogen storage systems. Int J Hydrogen Energy 2002;27:1063–69.

[60] Xu W-C, Takahashi K, Matsuo Y, Hattori Y, Kumagai M, Ishiyama S, et al. Investigation of hydrogen storage capacity of various carbon materials. Int J Hydrogen Energy 2007;32:2504–12.

[61] Chambers A, Rodriguez NM, Baker RTK. Catalytic engineering of carbon nanostructures. Langmuir 1995;11:3862–6.

[62] Baker RTK. http://www.wtec.org/loyola/nano/us_r_n_d/09_03.htm; Jan 1998.

[63] Goodman DW, Kelley RD, Madey TE, White JM. Kinetics of carbon deposition from CO on Cu(110) and Ni(100)1980. J Vac Sci Technol 1980;17:143.

[64] Chen JP, Yang RT. Chemisorption of hydrogen on different planes of graphite-a semi-empirical molecular orbital calculation. Surf Sci 1989;216:481–8.

[65] Nakamura J, Hirano H, Xie M, Matsuo I, Yamada T, Tanaka K. Formation of a hybrid surface of carbide and graphite layers on Ni(100) but no hybrid surface on Ni(111). Surf Sci 1989;222: L809–17.

[66] Chambers A, Park C, Baker RTK, Rodriguez NM. Hydrogen storage in graphite nanofibers. J Phys Chem B 1998;102:4253–6.

[67] Ahn CC, Ye Y, Ratnakumar BV, Whitham C, Bowman Jr RC, Fultz B. Hydrogen desorption and adsorption measurements on graphite nanofibers. Appl Phys Lett 1998;73:3378–80.

[68] Ströbel R, Jörissen L, Schliermann T, Trapp V, Schütz W, Bohmhammel K, et al. Hydrogen adsorption on carbon materials. J Power Sci 1999;84:221–4.

[69] Park C, Anderson PE, Chambers A, Tan CD, Hidalgo R, Roderiguez NM. Further studies of the interaction of hydrogen with graphite nanofibers. J Phys Chem B 1999;103:10572–81.

[70] Doll GL, Eklund PC, Senatore G. Elastic neutron scattering studies of H₂ and D₂ physisorbed stage 2 graphite-potassium. In: Dresselhaus MS, editor. NATO ASI Series, v. 148, Intercalation in layered materials. New York: Plenum Press; 1986.p. 309–11.

[71] Biris AR, Lupu D, Dervishi E, Li Z, Saini V, Saini D, et al. Hydrogen storage in carbonbased nanostructured materials. Part Sci Technol 2008;26:297–305.

[72] Pant KK, Gupta RB. Hydrogen Storage in Carbon Materials. In: Gupta RB, editor, Hydrogen fuel: production, transport, and storage. New York: CRC Press; 2009. p. 409–36.

[73] Thess A, Lee R, Nikolaev P, Dai H, Pitit P, Robert J, et al. Crystalline ropes of metallic carbon nanotubes. Science 1996;273:483–7.

[74] Iijima S, Ichihashi T. Single-shell carbon nanotubes of 1-nm diameter. Nature 1993; 363:603–5.

[75] Bethune DS, Kiang C-H, de Vries MS, Gorman G, Savoy R, Vasquez J, et al. Cobaltcatalysed growth of carbon nanotubes with single-atomic-layer walls. Nature 1993; 363:605–7.

[76] Guo T, Nikolaev P, Thess A, Colbert DT, Smalley RE. Catalytic growth of single-walled nanotubes by laser vaporization. Chem Phys Lett 1995;243:49–54.

[77] Tohji K, Goto T, Takahashi H, Shinoda Y, Shimizu N, Jeyadevan B, et al. Purifying single-walled nanotubes. Nature 1996;383:679.

[78] Bandow S, Rao AM, Williams KA, Thess A, Smalley RE, Eklund PC. Purification of single-wall carbon nanotubes by microfiltration. J Phys Chem B 1997;101:8839–42.

[79] Dujardin E, Ebbesen TW, Krishnan A, Treacy MMJ. Purification of single-shell nanotubes. Adv Mater 1998;10:611–3.

[80] Rinzler AG, Lui J, Dai H, Nikolaev P, Huffman CB, Roderiguez-Macias FJ, et al. Large-scale purification of single-wall carbon nanotubes: process, product, and characterization. Appl Phys A 1998;67:29–37.

[81] Shelimov KB, Esenaliev RO, Rinzler AG, Huffman CB, Smalley RE. Purification of single-wall carbon nanotubes by ultrasonically assisted filtration. Chem Phys Lett 1998;282:429–34.

[82] Dillon AC, Gennett T, Jones KM, Alleman JL, Parilla PA, Heben MJ. A simple and complete purification of single-walled carbon nanotube materials. Adv Mater 1999;11:1354–8.

[83] Dillon AC, Parilla PA, Alleman JL, Perkins JD, Heben MJ. Controlling single-wall nanotube diameters with variation in laser pulse power. Chem Phys Lett 2000;316:13–8.

[84] Dillon AC, Jones KM, Bekkedahl TA, Kiang CH, Bethune DS, Heben MJ. Storage of hydrogen in single-walled carbon nanotubes. Nature 1997;386:377–9.

[85] Ye Y, Ahn CC, Witham C, Fultz B, Liu J, Rinzler AG, et al. Hydrogen adsorption and cohesive energy of single-walled carbon nanotube. Appl Phys Lett 1999;74:2307–9.

[86] Brown CM, Yildirim T, Neumann DA, Heben MJ, Gennett T, Dillon AC, et al. Quantum rotation of hydrogen in single-wall carbon nanotubes. Chem Phys Lett 2000;329:311–6.

[87] Liu C, Fan YY, Liu M, Cong HT, Cheng HM, Dresselhaus MS. Hydrogen storage in single-walled carbon nanotubes at room temperature. Science 1999;286:1127–9.

[88] Dillon AC, Gennett T, Parilla PA, Alleman JL, Jones KM, Heben MJ. Nanotubes and related materials. Mater Res Soc Symposium Proceedings, vol. 633. 2001. P. A5.2.1–6.

[89] Nutzenadel C, Zuttel A, Chartouni D, Schlapbach L. Electrochemical storage of hydrogen in nanotube materials. Electrochem Solid-State Lett 1999;2:30–2.

[90] Lee SM, Park KS, Choi YC, Park YS, Bok JM, Bae DJ, et al. Hydrogen adsorption and storage in carbon nanotubes. Synth Met 2000;113:209–16.

[91] Lachawiec Jr AJ, Qi G, Yang RT. Hydrogen storage in nanostructured carbons by spillover: bridge-building enhancement. Langmuir 2005;21:11418–24.

[92] Dag S, Ciraci S. Coverage and strain dependent magnetization of titanium-coated carbon nanotubes. Phys Rev B: Condens Matter 2005;71:165414–1–6.

[93] Kim Y-H, Zhao Y, Williamson A, Heben MJ, Zhang SB. Nondissociative adsorption of H₂ molecules in light-element-doped fullerenes. Phys Rev Lett 2006;96:1–4.

[94] Yildirim T, Ciraci S. Titanium-decorated carbon nanotubes as a potential high-capacity hydrogen storage medium. Phys Rev Lett 2005;94:1–4.

[95] Wang XK, Lin XW, Dravid VP, Ketterson JB, Chang RPH. Stable glow discharge for synthesis of carbon nanotubes. Appl Phys Lett 1995;66:2430–2.

[96] Rochefort A. http://www.nanotech-now.com/nanotube-buckyball-sites.htm. 2009.

[97] Iijima S. Helical microtubules of grahpitic carbon. Nature 1991;354:56–8.

[98] Ebbesen TW, Ajayan PM. Large-scale synthesis of carbon nanotubes. Nature 1992; 358:220–2.

[99] Ishigami M, Cumings J, Zettl A, Chen S. A simple method for the continuous production of carbon nanotubes. Chem Phys Lett 2000;319:457–9.

[100] Ebbesen TW, Ajayan PM, Hiura H, Tanigaki K. Purification of nanotubes. Nature 1994;367:519.

[101] Colomer J-F, Piedigrosso P, Wilems I, Journet C, Bernier P, Tendeloo GV, et al. The production and structure of pyrolytic carbon nanotubes. (PCNTs) J Chem Soc Faraday Trans 1998;94:3753–8.

[102] Endo M, Takeuchi K, Igarashi S, Kobori K, Shiraishi M, Kroto HW. The production and structure of pyrolytic carbon nanotubes (PCNTs). J Phys Chem Solids 1993; 54:1841–8.

[103] Ivanov V, Nagy JB, Lambin P, Lucas A, Zhang XB, Zhang XF, et al. The study of carbon nanotubules produced by catalytic method. Chem Phys Lett 1994;223:329–35.

[104] Chen P, Wu X, Lin J, Tan KL. High H₂ uptake by alkali-doped carbon nanotubes under ambient pressure and moderate temperature. Science 1999;285:91–3.

[105] Fan Y-Y, Liao B, Liu M, Wei Y-L, Lu M-Q, Cheng H-M. Hydrogen uptake in vapor-grown carbon nanofibers. Carbon 1999;37:1649–52.

[106] Pinkerton FE, Wicke BG, Olk CH, Tibbetts GG, Meisner GP, Meyer MS, et al. Thermogravimetric measurement of hydrogen absorption in alkali-modified carbon materials. J Phys Chem B 2000;104:9460–7.

[107] Wu XB, Chen P, Lin J, Tan KL. Hydrogen uptake by carbon nanotubes. Int J Hydrogen Energy 2000;25:261–5.

[108] Feather fibers fluff up hydrogen storage capacity. CEP. August 2009; p.14–5.

..................Content has been hidden....................

You can't read the all page of ebook, please click here login for view all page.

Table of Contents for Chapter 7: Hydrogen Storage on Carbon Adsorbents

Create new playlist

Sign In

Sign Up

CHAPTER 7

1. INTRODUCTION