Ind. Eng. Chem. Res. 2009, 48, 10015–10020
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Adsorption and Diffusion of Carbon Dioxide on Metal-Organic Framework (MOF-5) Zhenxia Zhao,†,‡ Zhong Li,‡ and Y. S. Lin*,† Department of Chemical Engineering, Arizona State UniVersity, Tempe, Arizona 85287-6006, and Research Institute of Chemical Engineering, South China UniVersity of Technology, Guangzhou, 510640, China
Adsorption equilibrium and diffusion of CO2 on microporous metal-organic frameworks (MOF-5, or IRMOF1) crystals were experimentally studied by the gravimetric method in the pressure range up to 1 atm. The MOF-5 crystal cubes of about 40-60 µm in sizes were synthesized by the solvothermal method. Freundlich adsorption isotherm equation can fit well CO2 adsorption isotherms on MOF-5, with isosteric heat of adsorption of about 34 kJ/mol. Diffusion coefficient of CO2 in the MOF-5 is in the range of 8.1-11.5 × 10-9 cm2/s in 295-331K with activation energy of 7.61 kJ/mol. MOF-5 offers attractive adsorption properties as an adsorbent for separation of CO2 from flue gas. 1. Introduction Mental-organic frameworks (MOFs) represent a fundamentally new class of nanoporous materials.1,2 The crystalline porous materials are constructed from transition metal ions and bridging organic ligands, which have a vast, sturdy and open crystalline structure.3 Owing to their extra-high porosity, ordered, and wellcharacterized porous structures and adjustable chemical functionality, MOFs porous materials are attracting increasing attention on the potential application in adsorption, separation, gas storage, and heterogeneous catalysis.4,5 Carbon dioxide separation is a key step in carbon sequestration for preventing global warming.6 Adsorption separation by a porous material is one of promising methods considered for separation of CO2 from flue gas.7 Identifying an adsorbent which stores and releases CO2 with fast kinetics and high reversibility over multiple cycles is the key to the development of the viable adsorption process for CO2 separation.7 The unique structure of MOFs, make them attractive as adsorbents for CO2 separation with properties better than other porous materials, such as silicates, carbons, and zeolites. For example, Millward and Yaghi8 reported ambient temperature adsorption equilibrium data for CO2 in a large variety of MOFs. The voluminous space enclosed by MOF-5 and MOF-177 results in a CO2 adsorption capacity of 22 and 33.5 mmol/g, respectively, which was substantially greater than that of any other porous material reported. In addition, the ability to synthesize MOFs with various organic linkers and metal joints provides tremendous flexibility in designing this porous material to have specific physical characteristics and chemical functionalities for the application of special gas sorption.9 For design and operation of the novel MOFs materials utilized for CO2 capture and separation, experimental data of CO2 adsorption equilibrium and diffusion coefficients in the crystals of MOFs are of great importance. Several research groups studied CO2 adsorption equilibrium on MOFs, such as those with square channels (MOF-2),7 pores decorated with open metal sites (MOF-505 and Cu3(BTC)2),10 hexagonally packed cylindrical channels (MOF-74),11 interpenetration (IRMOF-11), 2b amino- and alkyl-functionalized pores (IRMOFs-3 and -6)2 * To whom correspondence should be addressed. E-mail: Jerry.Lin@ ASU.edu. † Arizona State University. ‡ South China University of Technology.
and the extra-high porosity frameworks MOF-5 and MOF-177.12 In comparison, research work on diffusion of gases including CO2 in MOFs was only limited to theoretical and molecular dynamics (MD) simulation results on argon diffusion in MOF2, MOF-3, MOF-5, and CuBTC and of other gases (H2, N2, CO2)13 and alkanes (methane, n-pentane, n-hexane) diffusion in MOF-5.10,11 No experimental data on diffusivities of CO2 in MOFs are available in literature. The objective of the present work was to investigate the adsorption and diffusion mechanism of CO2. The data were obtained with carbon dioxide pressures up to 1 atm due to the availability of experimental setup and the interest in exploring the use of MOF-5 as adsorbents for separation of carbon dioxide from flue gas in the atmospheric pressure. 2. Experimental Section 2.1. Synthesis and Characterization of MOF-5 Crystals. MOF-5 crystals of uniform size were synthesized by using terephthalic acid (BCD, +99%, from Aldrich) and zinc nitrate (Zn(NO3)2, 99.5%, from Fluka), and dimethylformamide (DMF, 99%, from Mallinckrodt) as the organic solvent. All chemicals used in this work were purchased from these vendors and used without further purification. The MOF-5 was synthesized following the optimized procedure similar to that reported previously in the literature.1 In synthesis, the solvent DMF was first degassed by argon for 60 min, and then Zn(NO3)2 · 6H2O (1.664 g, 5.60 mmol) and terephthalic acid (H2BDC, 0.352 g, 2.12 mmol) were dissolved in 40 mL of degassed DMF solvent. The mixture was quickly transferred to a 50 mL glass vial and sealed. The vial was then heated to 403 K and held for 2 h under autogenous pressure by solvothermal synthesis. After the reaction, the vial was taken out of the oil bath and cooled down to the room temperature naturally. The cubic-like crystals of colorless powder were isolated by filtration and washed thoroughly with DMF in order to remove the unreacted zinc nitrate. After that, the crystals were immersed in chloroform (50 mL), sealed tightly, and put into the oven at 343 K for 3 days. During the heating process, the solvent was decanted and replenished every day. Finally, the sample was dried under vacuum at 363 K overnight and stored in a desiccator until it was used.14 The surface morphology and particle size of the crystalline MOF-5 samples were observed by a Philips FEI XL-30 scanning
10.1021/ie900665f CCC: $40.75 2009 American Chemical Society Published on Web 09/11/2009
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Figure 1. Schematic diagram of Cahn microbalance system for studying CO2 adsorption on the MOF-5: (A) CO2; (B) helium; (C) flow rate display; (D) furnace; (E) Cahn C-1000 electronic microbalance; (F) temperature controller; (G) computer.
electron microscope at an accelerating voltage of 20 kV after gold deposition. Its crystal phase structure was examined by powder X-ray diffraction (PXRD) with a conventional Bruker D8 diffractometer at 20 kV, 5 mA with a scan speed of 2°/min and a step size of 0.02° in 2θ, using Cu KR (λ ) 0.1543 nm) radiation. The pore textural properties of the MOF-5 samples were measured by Micromeritics ASAP 2020 adsorption porosimeter at liquid nitrogen temperature of 77 K. Their specific Langmuir area and Brunauer-Emmet-Teller (BET) surface area, pore volume, and pore size distribution were obtained from N2 isotherms measured. 2.2. Adsorption and Diffusion Measurements. Adsorption and diffusion experiments were conducted on a Cahn electronic microbalance system (Cahn D-101) as shown in Figure 1. One arm of the microbalance had a stainless-steel pan suspended at its end to hold the crystalline MOF-5 sample. The pan was attached to the microbalance arm by a platinum wire (Gauge 36, Fisher Scientific). The temperature of the sample pan was maintained at the desired temperature with an Omega CN76000 temperature controller connected to a tubular furnace mounted on the outside of the 2-in. i.d. Pyrex balance tube. The furnace was mounted so that the sample pan was exactly at its center. The temperature of sample pan (close to the sample temperature) was monitored by a thermocouple (K-type Omega K-72-SRTC). The gas delivery system consisted of a helium purge gas and carbon dioxide sorption gas. The purge and sorption gas were alternatively delivered to the system by switching a crossover value (Swagelock). The transient and equilibrium weight changes were recorded using a computer-aided data acquisition system. About 10-20 mg of MOF-5 sample was held in the stainlesssteel sample pan and degassed at 393 K for 8 h under helium purge flow at 100 mL/min. After the sample weight became constant, the temperature was cooled to the desired adsorption temperature. After that the adsorption process started by switching from purge to the adsorption gas at 100 mL/min, using the crossover valve. After the adsorption process reached equilibrium, gas desorption was conducted at the same temperature by switching back to purge gas at 100 mL/min. The 100 mL/min flow rate was found optimum where systematic problems of gas dispersion (caused by switching between gases) and weight instability (taking place at high flow rates) were minimized.15 Data were deduced to obtain uptake (mg gas
Figure 2. Scanning electron microscopy (SEM) image of the MOF-5 cubic crystals prepared in this work.
adsorbed/g adsorbent or mmol/g) and fractional uptake rate (uptake/equilibrium uptake) curves. Gas adsorption isotherms were measured at 295.7, 304.7, 313.7, 322.7, and 331.7 K, with the CO2 partial pressures varied between 0 and 1 atm, respectively. The adsorption uptake of CO2 on the MOF-5 was calculated as follows: qe )
We - Wa 1000 WaMCO2
(1)
qt )
Wt - Wa 1000 WaMCO2
(2)
where MCO2 (g/mol) is the molecular weight of CO2 molecule; We (g) and Wt (g) are the amount of adsorbent (the MOF-5) at equilibrium and time t(s); Wa is the initial weight of the MOF5; and qt (mmol/g) and qe (mmol/g) are the amount of adsorbed per gram at time t(s) and at equilibrium, respectively. 3. Results and Discussion 3.1. Characteristics of MOF-5 Crystals. Figure 2 shows the SEM image of the MOF-5 sample prepared in this work. It can be seen that the crystals are of cube-shape with sizes in the
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Figure 3. XRD patterns of MOF-5 sample prepared in this work.
range of 40-60 µm. Most of these MOF-5 particles present in the form of the monocrystal structure, instead of polycrystal structure.1 The crystal size is slightly smaller and more uniform compared with other MOF-5 samples prepared by other research groups.1,16 In this work, the reaction time for MOF-5 synthesis was varied from 1 to 4 h. The MOF-5 structure was formed at 403 K with 2 h reaction time. After 2 h, many crystallites grew together and formed polycrystals with a larger particle size. Therefore, 2 h reaction time was considered as the optimal, and the MOF-5 samples used for adsorption studies in this work were synthesized with 2 h reaction time. Figure 3 shows the powder X-ray diffraction (PXRD) of the synthesized MOF-5 sample. The main peaks of the MOF-5 sample match well with the published XRD pattern for the MOF-5.12 The well resolved peaks shown in Figure 3 imply the high crystallinity of the MOF-5 samples obtained in this work. Figure 4 shows the nitrogen adsorption isotherm of the MOF-5 sample. This N2 isotherm on the MOF-5 sample is of type-I with an initial steep increase in the nitrogen uptake followed by a plateau at high pressures. Its BET surface area, Langmuir surface area, and pore volume are calculated to be 2304 and 2517 m2/g and 1.01 cm3/g, respectively. Its micropore area and micropore volume are 2167 m2/g and 0.86 m3/g, respectively. In density function theory (DFT) differential pore volume plot, only one sharp peak is observed at approximately 14 Å for the MOF-5 sample (Figure 4B). The high surface area as well as the pore size confirms the formation of MOF-5 structure. MOF-5 can be synthesized by several different synthesis methods.17 For fast, large quantity and cost-effective synthesis of MOF-5, many groups adopted N,N′-dimethylformamide (DMF) as the alternative solvent instead of N,N′-diethylformamide (DEF). The average special surface area of the MOF-5 synthesized in DMF were in the range of 600-1300 m2/g.17 In comparison, the surface area of our MOF-5 samples by DMF approach is relatively high (∼2500 m2/g) as compared to that reported previously.18,19 Substantial variation (from 600 to 3400 m2/g) in the surface area values of the synthesized MOF-5 are mainly attributed to (1) the presence of Zn(OH)2 species which partly occupied the cavities of the MOF-5 framework,20 (2) some crystallographic defects, and (3) the minor phase consisting of doubly interpenetrated MOF-5 framework.17 Therefore, some variation in synthesis, including solvent, ratio of Zn/BDC, and synthesis time, can influence the surface area value of the MOF-5 samples. 3.2. CO2 Adsorption Isotherms. Figure 5 shows adsorption isotherms of CO2 on MOF-5 at different temperatures in 295-332 K measured experimentally and fitting results by
Figure 4. (A) Nitrogen sorption isotherm of the MOF-5 sample. (B) Pore size distribution the MOF-5 sample calculated from the density function theory (DFT) model.
Figure 5. CO2 adsorption isotherms on MOF-5 at various temperatures (points, experimental data; solid curves, model fitting by Freundlich equation). Table 1. Values for Freundlich Isotherm Equation for CO2 Adsorption on MOF-5 Freundlich -1
temperature (K)
KF (g )
n
R2
295.7 304.7 313.7 322.7 331.7
2.18 1.54 1.13 0.82 0.54
1.21 1.15 1.11 1.07 1.07
0.999 0.998 0.999 0.999 0.999
Freundlich isotherm equation (qe ) KFpi1/n). The values of Freundlich isotherm parameters and the regression coefficients are listed in Table 1. The experimental data show a modest increase in the amount adsorbed with increasing CO2 partial pressure. There is no plateau in adsorption isotherm in the pressure range investigated, indicating that MOF-5 can adsorb more CO2 at higher CO2 partial pressures. As shown in Figure
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Table 2. Comparison of CO2 Adsorption Capacity for MOF-5 with Other Adsorbents Reported in the Literature
adsorbent
adsorption capacity (mmol/g)
pressure (atm)
temperature (K)
reference
MOF-5 IRMOF-1(solvothermal) IRMOF-1(solvothermal) MOF-5 (sonochemical) MOF-5 (microwave) zeolite 13X all-silica DD3R zeolite Queensland coal activated carbon
2.10 1.92 1.8-1.9 1.13 1.12 3.3 1.7 0.8 2.0
1.0 1.0 1.0 1.0 1.0 1.0 2.0 20 1.0
296 298 298 298 298 323 298 313 298
this work 9 8 22 23 24 25 26 27
5, the maximum adsorption CO2 is found as 2.1 mmol/g (92.5 mg/g) at 295 K and 1 atm. This corresponds to a percentage of monolayer coverage of 1.03% under studied conditions, using the equation Xmonolayer ) qAoNA/SA, where q(mol/g) is the uptake capacity, SA is the Langmuir surface area (2517 m2/g), Ao is the adsorbate cross sectional area estimated from the molecular kinetic diameter assuming a spherical shape (9 Å2 for CO221), and NA is Avogadro’s number (6.02 × 1023 molecule/mol). Several groups have studied CO2 adsorption isotherms of MOF-5 (IRMOF-1) in the low pressure regime and their equilibrium amounts are compared with the result obtained in this work in Table 2. The amount of adsorbed CO2 obtained in this work is comparable to the results reported by Yaghi’s group,9 about twice that value for MOF-5 samples synthesized by sonochemical22 and microwave heating.23 It is safe to conclude that the MOF-5 sample prepared in this work by the solvothermal method exhibits essentially the same CO2 adsorption properties as the MOF-5 samples reported by other research groups. The minor difference may be due to a small amount of entrapped solvent inside MOF-5 prepared by different procedures. Table 2 also shows the CO2 adsorption capacities of other adsorbents, including activated carbon, coals, and zeolites. Obviously, MOF-5 material has higher adsorption capacity than coals and activated carbon, and comparable with 13X zeolite at pressures less than 1 atm. Compared with some commercial activated carbons and char coals, the extra-large surface area, large pore volume, and uniform pore size distribution of MOF-5 offer even higher CO2 adsorption amount. At 1 atm, however, MOF-5 has a lower CO2 sorption capacity than zeolite 13X. The comparison in Figure 5 and the values of regression coefficients (R2) listed in Table 1 show that the Freundlich isotherm equation fits well the experimental data, which is consistent with H2 adsorption isotherm on MOF-5 reported by Deng’s group.28 This empirical model applies only to nonideal sorption on heterogeneous surface as well as multilayer sorption.15 The good fitting of the Freundlich equation with the experimental data suggests that CO2 adsorption on MOF-5 can be described by reversible adsorption and the occurrence for multilayer adsorption. As shown in Table 1, the multilayer saturation capacity calculated from Freundlish equation (KF) is slightly higher than the real adsorbed values with the adsorption intensity (n) greater than unity. From the comparison of the adsorption intensity (n) with H2 (n ) 1.025, at 195 K and low pressure),28 it can be seen that MOF-5 is more favorable for the adsorption of CO2 than H2. Isoteric heat of adsorption can be obtained from the adsorption isotherm by ∆HS RT2
)
(
∂ln(P/PO) ∂T
)
q
(3)
A plot of ln (P/PO) versus (1/T) for CO2 adsorption on the MOF-5 at constant q of 1.5 mmol/g is given in Figure 6. The
Figure 6. Temperature dependence of multilayer saturation capacity for CO2 adsorption on MOF-5 (KF).
Figure 7. CO2 transient adsorption uptakes on MOF-5 at different feed pressures (at 295.7 K).
isosteric heat of adsorption (-∆Hs) obtained from the slope of the plot for MOF-5 is 34.1 kJ/mol. This value is close to the isoteric heat of adsorption for CO2 adsorption on 13X zeolite at similar adsorption capacity (33-34 kJ/mol).24 3.3. CO2 Adsorption Kinetics. Figure 7 shows CO2 adsorption uptakes on the MOF-5 at different CO2 partial pressures. As discussed above, the equilibrium adsorption amount increases with pressure. To show the effects of the pressure (or CO2 loading) on diffusion rate, the uptake curves are plotted as the fractional uptake, the ratio of the uptake (qt) at time t to the equilibrium uptake (qe), versus adsorption in Figure 8. As shown, the fractional transient uptake data for CO2 adsorption at different pressure range in 0.1-1.0 atm are essentially same. This indicates that CO2 loading has negligible effect on diffusivity of CO2 in MOF-5. Figure 9 shows CO2 fractional sorption uptakes on MOF-5 at different temperatures. Assuming diffusion into porous spheres, the transient fraction uptakes can be described by the following equation in the region with uptake qe less than 70%:29 qt 6 = qe rc
D Mt π
(4)
where DM (cm2/s) is the intracystalline diffusion coefficient, and rc (cm) is the crystal radius. In this case, the diffusion time constant (DM/rc2, s-1) was obtained from the slope of qt/qe versus t. Figure 10 plots fractional uptake (qt/qe) of CO2 versus the square root of adsorption time at different temperatures. It can
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Figure 11. Arrhenius plot of CO2 diffusivity in MOF-5 crystals. Figure 8. CO2 fractional dsorption uptakes on MOF-5 at different feed pressures (at 295.7 K).
Table 3. CO2 Diffusion Constant and Diffusivity for MOF-5 Measured in This Work temperature (K)
DM/rc2 (10-3 s-1)
DM (10-9 cm2/s)
R2
295.7 304.7 313.7 322.7 331.7
1.31 1.53 1.64 1.77 1.85
8.17 9.57 10.26 11.04 11.54
0.990 0.993 0.996 0.996 0.996
Table 4. Comparison of Expermental CO2 Diffusion Data for MOF-5 with 13X Zeolite measurement temperature adsorbent method (K)
Figure 9. Fractional CO2 adsorption uptakes on MOF-5 at different temperatures (0.5 atm).
Figure 10. Plots of the fractional adsorption uptake (qt/qe) against the square root of adsorption time at different temperatures (0.5 atm).
be seen from Figure 10 that eq 4 fits the experimental data reasonably well (with R2 > 0.99). Table 3 lists the values of diffusion time coefficient for CO2 on the MOF-5 obtained by regressing the uptake curves with eq 4. The CO2 diffusion time constants (DM/rc2) on the MOF-5 are in range of 1.307-1.847 × 10-3 s-1. The corresponding diffusion coefficient (DM) using the average particle radius rc of 25 µm is in the range of 8.17 × 10-9 to 11.54 × 10-9 cm2/s. The Arrhenius plot of the diffusivity is given in Figure 11. The activation energy Ea for diffusion of CO2 in MOF-5, obtained from the slope of the plot in Figure 11, is 7.61 kJ/mol. CO2 diffusion in MOF-5 follows activated diffusion process. Similar to the diffusion in zeolites, the experimentally measured CO2
MOF-5 13X 13X 13X 13X
gravimetric manometry NMR gravimetric frequency response
296 303 302 493 397
DM (cm2/s) 7.9 × 10-9 6.49 × 10-11 3.4 × 10-6 2.0 × 10-4 6.4 × 10-11
activation energy (kJ/mol) reference 7.6 11.7 10 35.6
this work 33 34 35 36
diffusivity data for MOF-5 is several orders of magnitude smaller than that obtained by the molecular dynamics simulation (4.0 × 10-5 cm2/s at 298 K).30 MOF-5 has cubic cage structure with inner cavities of about 12 and 15 Å in diameter for the ajacient cages37 and aperture openings of 8 Å width in three-dimensional directions.38 The zeolite that has similar pore structure and size is FAU type zeolites (including 13X) defined by 10 solidite cage with cavity of about 12 Å in size and aperture opening of 7.4 Å. Table 4 compares the CO2 diffusion data in MOF-5 with the data in 13X zeolite measured by different experimental methods. As shown in Table 4, there is significant differences in the data of CO2 diffuvisity in 13X measured by the different methods. Bulow31 suggested a complex and yet unknown mechanism of CO2 diffusion in 13X zeolite that might be responsible for such differences and relatively few reliable studies on CO2 diffusion in 13X zeolites. The activation energy for CO2 diffusion in MOF-5 is slightly smaller than that in 13X zeolites. Xiao and Wei32 show that the activation energy for diffusion increases with increasing ratio of the kinetic diameter of diffusing molecule to the pore opening of zeolite (λ ) dm/dp, which is 0.275 for MOF-5 and 0.285 for 13X). The smaller activation energy for CO2 diffusion in MOF-5 is consistent with the theory of Xiao and Wei. 4. Conclusions CO2 adsorption equilibrium and diffusion were studied gravimetrically on homemade metal-organic framework (MOF5) crystals. The MOF-5 sample prepared in this work is of cubic crystals of 40-60 µm in sizes with specific surface area of 2517
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m2/g. MOF-5 has CO2 adsorption capacity of 2.1 mmol/g (or 9.24%) at 295.7 K and 1 atm. CO2 adsorption isotherms for the MOF-5 can be well described by Freundlich adsorption isotherm equation. CO2 diffusion in MOF-5 crystals is an activated process with diffusivity in the range of 8.1-11.5 × 10-9 cm2/s in 295-331 K and activation energy for diffusion of 7.6 kJ/ mol. The CO2 pressure (or loading) has negligible effect on the rate of CO2 diffusion in MOF-5 crystals. Acknowledgment Z.Z. is grateful to China Scholarship Council for a fellowship to support her visit to ASU. The authors acknowledge the financial support of the U.S. National Science Foundation, National Natural Science Foundation of China (No. 20606012) and the Doctorate Foundation of South China University of Technology on this project. Literature Cited (1) Panella, B.; Hirscher, M.; Pu¨tter, H.; Mu¨ller, U. Hydrogen Adsorption in Metal-Organic Frameworks: Cu-MOFs and Zn-MOFs Compared. AdV. Funct. Mater. 2006, 16, 520. (2) Eddaoudi, M.; Kim, J.; Rosi, N.; Vodak, D.; Wachter, J.; O’GKeeffe, M.; Yaghi, O. M. Systematic Design of Pore Size and Functionality in Isoreticular MOFs and Their Application in Methane Storage. Science. 2002, 295, 469. (3) Sanderson, K. Space Invaders. Nature. 2007, 448, 16–746. (4) Chae, H. K.; Siberio-Pe’rez, D. Y.; Kim, J.; Go, Y. B.; Eddaoudi, M.; Matzger, A. J.; O’Keeffe, M.; Yaghi, O. M. A Route to High Surface Area, Porosity, and Inclusion of Large Molecules in Crystals. Nature 2004, 427, 523. (5) Antek, G.; Adam, W. F.; Matzger, J.; Yaghi, O. M. Exceptional H2 Saturation Uptake in Microporous Metal-Organic Frameworks. J. Am. Chem. Soc. 2006, 128, 3494. (6) Loo, S. V.; Elk, E. P. V.; Versteeg, G. F. The Removal of Carbon Dioxide with Activated Solutions of Methyl-Diethanol-Amine. J. Pet. Sci. Eng. 2007, 55, 135. (7) Li, H.; Eddaoudi, M.; Groy, T. L.; Yaghi, O. M. Establishing Microporosity in Open Metal-Organic Frameworks: Gas Sorption Isotherms for Zn(BDC) (BDC ) 1,4-Benzenedicarboxylate). J. Am. Chem. Soc. 1998, 120, 8571. (8) Millward, A. R.; Yaghi, O. M. Metal-Organic Frameworks with Exceptionally High Capacity for Storage of Carbon Dioxide at Room Temperature. J. Am. Chem. Soc. 2005, 127, 17998. (9) Walton, K. S.; Millward, A. R.; Dubbeldam, D.; Frost, H.; Low, J. J.; Yaghi, O. M.; Snurr, R. Q. Understanding Inflections and Steps in Carbon Dioxide Adsorption Isotherms in Metal-Organic Frameworks. J. Am. Chem. Soc. 2008, 130 (2), 406. (10) Chen, B. L.; Ockwig, N. W.; Millward, A. R.; Contreras, D. S.; Yaghi, O. M. High H2 Adsorption in a Microporous Metal-Organic Framework with Open Metal Sites. Angew. Chem., Int. Ed. 2005, 44, 4745. (11) Rosi, N. L.; Kim, J.; Eddaoudi, M.; Chen, B.; O’Keeffe, M.; Yaghi, O. M. Rod Packings and Metal-Organic Frameworks Constructed from RodShaped Secondary Building Units. J. Am. Chem. Soc. 2005, 127, 1504. (12) Li, H.; Eddaoudi, M.; O’Keeffe, M.; Yaghi, O. M. Design and Synthesis of an Exceptionally Stable and Highly Porous Metal-Organic Framework. Nature. 1999, 402, 276. (13) Skoulidas, A. I. Molecular Dynamics Simulations of Gas Diffusion in Metal-Organic Frameworks: Argon in CuBTC. J. Am. Chem. Soc. 2004, 126, 1356. (14) Li, Y. W.; Yang, R. T. Gas Adsorption and Storage in MetalOrganic Framework MOF-177. Langmuir 2007, 23, 12937. (15) Alsyouri, H. M.; Lin, J. Y. S. Gas Diffusion and Microstructural Properties of Ordered Mesoporous Silica Fibers. J. Phys. Chem. B. 2005, 109 (28), 13623.
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ReceiVed for reView April 26, 2009 ReVised manuscript receiVed August 18, 2009 Accepted August 24, 2009 IE900665F
