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Notes On the Mechanism of Selective Electroadsorption of Protons in the Pores of Carbon Molecular Sieves Linoam Eliad, Gregory Salitra, Abraham Soffer, and Doron Aurbach* Department of Chemistry, Bar-Ilan University, Ramat-Gan 52900, Israel Received March 23, 2004. In Final Form: November 3, 2004
Introduction Carbon materials are structurally puzzling solids, and numerous articles have been published trying to resolve their morphology.1 Of interest are the nanoporous, high specific surface area carbons which, besides their high ion adsorption capacity, a property shared with other porous materials,2 also show high electroadsorption capacity, which is due to their electronic conductivity. This enables modifying ion adsorption by changing the electrical potential of the carbon electrode.3 Electroadsorption occurs also in nonporous, electrically conductive materials such as metals.4 However, high surface area porous carbons exhibit such large capacity figures that they become of interest for practical implementation in ion separation processes (e.g., acid separation and brackish water desalination)5 and energy storage in electrical double layer capacitors.6 Further unique features can be observed when the pore size falls close to the adsorbate molecular size. In this case, size exclusion of the larger adsorbates exists, so that electroadsorption stereoselectivity takes place. Using globular “molecular probes” of known dimensions, we could calibrate the pore size of the carbon molecular sieve. In turn, the electroadsorption stereoselectivity of these calibrated carbons were employed to establish a “yardstick” for assessing the dimensions of various ions.7 The obvious question that followed our study of neutral solutions was how the proton, which exhibits unique * To whom correspondence should be addressed. E-mail:
[email protected]. (1) (a) Rodriguez-Reinoso, F.; Molina-Sabio, M. Adv. Colloid Interface Sci. 1998, 76-77, 273. (b) Kinoshita, K. Carbon, Electrochemical and physicochemical properties; John Wiley & Sons, Inc.: New York, 1988; Chapter 1. (c) Jenkins, G. M.; Kuwana, K. Polymeric Carbons- Carbon Fibre, Glass and Char; Cambridge University Press: Cambridge, U.K., 1976. (2) (a) Rudzinski, W.; Charmas, R.; Piasecki, W.; Groszek, A. J. Langmuir 1999, 15 (25), 8553. (b) Rudzinski, W.; Charmas, R.; Piasecki, W. Langmuir 1999, 15 (25), 5921. (3) (a) Soffer, A.; Folman, M. J. Electroanal. Chem. 1972, 38, 25. (b) Yang, K.; Ying, T.; Yiacoumi, S.; Tsouris, C.; Vittoratos, E. S. Langmuir 2001, 17, 1961. (4) Jerkiewitz, G. Prog. Surf. Sci. 1989, 57 (2), 137. (5) (a) Oren, Y.; Soffer, A. Electrochim. Acta 1984, 28 (11), 1649. (b) Farmer, J. C.; et al. J. Electrochem. Soc. 1996, 143 (1), 161. (c) Oren, Y.; Soffer, A. J. Appl. Electrochem. 1983, 13, 473; (d) 1983, 13, 489. (6) (a) Burke, A. Proceedings of the 11th International Seminar on Double Layer Capacitors and Similar Energy storage Devices, Deerfield, FL, Dec 3-5, 2001; Florida Educational Seminars, Inc.: Boca Raton, FL, 2001. (b) Mayer, S. T.; Pekala, R. W.; Kaschmitter, J. L. J. Electrochem. Soc. 1993, 140, 446. (c) Bispo-Fonseca, I.; Aggar, J.; Sarrazin, C.; Simon, P.; Fauvarque, J. F. J. Power Sources 1999, 79, 238. (7) Eliad, L.; Salitra, G.; Soffer, A.; Aurbach, D. J. Phys. Chem. B 2001, 105, 6880.
behavior in the bulk solution and in proton exchange membranes, would behave within the pores of carbon molecular sieving electrodes. As a point charge, H+ is strongly hydrated in the bulk because the oxygen lone pairs of water can come very close to the proton nucleus. Therefore, it would be difficult to strip off this hydration sheath while entering very narrow pores. On the other hand, the proton in the carbon ultramicropores can occupy a very small space via hydrogen bonding. This interaction, which occurs within the pores, may compensate, together with the Coulombic interaction with the electronic negative charge for the loss of the bulk solution hydration sheath. When the pores are enlarged (by controlled burnoff), the stereoselectivity should disappear. This effect was, in fact, realized in a previous work.8 What was surprising, however, was that the selective proton electroadsorption was still very significant even when the carbon electrodes were very poor in surface oxygen groups. Although such carbons may slowly acquire some surface oxide groups from the air at ambient temperatures, they are by far less surface oxidized than carbons activated by hot concentrated nitric acid as was shown by thermogravimetric analysis (TGA) analysis.8 In the present work we wish to study further the mechanism of this unexpected type of selective proton electroadsorption. A special attention will be paid to the role of physisorbed water in the electroadsorption of protons. Experimental Section The carbons employed in this work were made by pyrolysis of a polymer precursor under the flow of nitrogen. Three precursor types were used: raw (unpainted cellulosic) Jeans cloth, ashless Whatman paper type 541, and a polyimide Kapton film 200NH (50 µm; a Pruduct of DuPont). The carbonization temperature programming is described elsewhere.12 It was identical for all three carbon precursors. The cellulosic precursors were carbonized while mixing ammonium chloride with the precursor, as a carbonization promoter. The use of this salt, as well as other electrolytes,9 is known to dramatically change the chemical course of the pyrolysis, leading finally to a much greater carbon yield. Carbon dioxide was used for activation at 900 °C at ambient pressure, and concentrated nitric acid was used at room temperature. In this work, the activation time was the only variable parameter. The electrochemical measurements, cyclic voltammetry in the electrical double layer potential range, cells, and equipment were described previously.10 The reference electrodes were saturated Hg/Hg2Cl2, which included the electrolytic bridge (saturated KCl). All scan rates were 1 mV/s unless specified. TGA was carried out under a stream of nitrogen (200 cm3/ min), at a heating rate of 10 °C/min using a Mettler TGA/STDA 851. Regarding the sample treatment before TGA, the sample (8) Eliad, L.; Salitra, G.; Soffer, A.; Aurbach, D. J. Phys. Chem. B 2002, 106, 10128. (9) Dae-Young, K.; Yoshiharu, N.; Masahisa, W.; Shigenori, K. Cellulose 2001, 8, 29-33. (10) Salitra, G.; Soffer, A.; Eliad, L.; Aurbach, D. J. Electrochem. Soc. 2000, 147, 2486. (11) Burch, D. M.; Desjarlais, A. O. National Institute of Standards and Technology-IR 5681; National Institute of Standards and Technology: Gaithersburg, MD, 1995. (12) Kinoshita, K. Carbon, Electrochemical and physicochemical properties; John Wiley & Sons, Inc.: New York, 1988; Chapter 4.
10.1021/la049238h CCC: $30.25 © 2005 American Chemical Society Published on Web 03/01/2005
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was washed with a large amount of water immediately after the HNO3 treatment and then dried for several hours in a drying oven at 80 °C. The samples were then stored for at least several days before being introduced into the thermogravimetric analyzer. We may, thus, assume that all the samples were equilibrated with the ambient humidity and oxygen. Water adsorption isotherms were obtained at room temperature, by purging the sample with a high purity N2 stream that was prewetted to a constant humidity by bubbling through various saturated salt solutions. The saturated salt solutions employed were11 saturated salt solution
% humidity
LiCl MgCl2 K2CO3 NaBr NH4Cl KCl KNO3
11.3 32.8 43.2 57.6 78.6 84.3 93.6
The wetting process required about 24 h. The weights of the carbon samples were about 200 mg, and the nitrogen flow rates were about 30 mL/min. The first point of zero humidity was taken as the sample weight after drying with a nitrogen purge for several hours. The nitrogen with prefixed humidity was then blown through the sample. This procedure was repeated with increasing specific humidity of the salt solutions.
Results and Discussion 1. Characteristics of the Proton-Selective Environment in Porous Carbon Electrodes. In a previous work,8 a proton-selective environment within the pores of a carbon molecular sieve was formed by treatment with nitric acid. A high concentration of oxygen-containing surface groups is produced12 and most likely interacts, via hydrogen bonds, with adsorbed water molecules that fill the pores.8 The resulting carbon indeed exhibited selectivity in proton adsorption at the proper pore size, which disappeared, as expected, when the pores were further enlarged while still keeping the protic envorinment intact (Figure 7 in ref 8). It is of interest to examine this “protic environment” by means of TGA. The thermograms for samples that have been treated with HNO3 to different extents are shown in Figure 1a. Two different temperature regions show an intensive rate of weight loss, presumed to be predominantly due to water desorption: one at 40-100 °C and the other at 100180 °C. We may assume that the first region corresponds to the desorption of relatively free, physisorbed water, while the higher temperature region corresponds to water that is hydrogen bonded to the oxide surface groups. As expected, the amount of physisorbed water in the porous carbons increases with the increase in the degree of activation. The amount of water desorbed in the range 100-180 °C does not increase monotonically upon activation but rather shows a maximum (Figure 1b). This shows that maximum interaction occurs while the adsorbate effective dimension is close to the pore wall separation. This phenomenon was also inspected with hydrated ions7 and with adsorption from the gas phase13 and was treated theoretically by Everett and Powl.14 For the present case of water without ions, it is difficult to state what effective dimension implies, because water clustering may play a role within the carbon pores,15 as will be shown later on. (13) Koresh, J.; Soffer, A. J. Chem. Soc., Faraday Trans. 1 1980, 76, 2457. (14) Everett, D. H.; Powl, J. C. J. Chem. Soc., Faraday Trans. 1 1976, 72, 619. (15) Brennan, J. K.; Bandosz, T. J.; Thomson, K. T.; Gubbins, K. E. Colloids Surf., A 2001, 187-188, 539-568.
Figure 1. (a) Thermogravimetric curves for a cellulose-based carbon cloth, activated by HNO3 to different extents. (b) Weight loss percent in the temperature range 100-180 °C of the strongly bound water.
Figure 2. Comparison of thermogravimetric curves for (a) carbon activated by HNO3 at 80 °C for 30 min. (b) Same as part a but followed by HTT at 300 °C for 30 min. (c) Same as part a but followed by HTT at 1000 °C for 30 min.
It is most interesting that the pristine carbon provides pore space for the physisorbed water but not for strongly interacting water. This is probably because there is no space to build water clusters. There are also only a few or no surface groups to form hydrogen bonds with the adsorbed water. The results shown in Figure 1 are compared with those of a carbon that is free of the protic environment (prepared by high-temperature degassing of the HNO3-treated samples). Figure 2 shows the thermograms of such carbons together with the parent, HNO3-treated sample. The water coming off at 100-180 °C from the HNO3-treated carbons is absent here, while the physisorbed water that desorbs is significantly greater. These results emphasize the existence of a protic environment which is stable, probably
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Figure 3. Cyclic voltammograms (1 mV/s) of carbon paper electrodes gradually activated in CO2 at 900 °C for different times, starting from 10 min up to 4.5 h. (a) 10 min of CO2 activation; (b) about 100-fold expansion of the adsorption ordinate of Figure 3a; and (c-g) 35, 60, 90, and 135 min and 4.5 h of CO2 activations, respectively.
up to 200 °C, in carbons that were activated by nitric acid and, thus, contain surface groups. In summary, the protic environment exists as water molecules adsorbed on the surface, which interact strongly with surface groups of the carbon molecular sieving pores. This is besides the weakly interacting physisorbed water that is present even in the pristine (nonactivated, nonsurface oxidized) carbon. 2. H+ Selective Electroadsorption in the Absence of the Protic Environment. As mentioned above, the carbon that was not treated with nitric acid and yet showed selective H+ electroadsorption was not totally free of surface oxygen groups. Brennan and coauthors had shown15 that a low concentration of oxygen surface groups can
serve as an anchor for adsorbed water clusters within the pores. Therefore, there is a possibility that the effect of specific adsorption of the small amount of surface groups that might be present in a carbon that was not treated with nitric acid will be enhanced by water cluster formation. To test this possibility, we prepared a series of CO2 activated carbons that were free of oxygen surface groups as much as possible, by avoiding exposure to air on the way from the activation process to the electrodes assembly in deaerated aqueous solutions. The results are shown in Figure 3a-g. Figure 3a is the cyclic voltammetry of a pristine carbon (with no activation). Figure 3b shows the same results on a
Notes
Figure 4. Water adsorption isotherms of various carbon samples at room temperature.
20-fold expanded scale. It shows no electrical double layer activity. The fact that a significant electric double layer capacity could be obtained only after activation as seen in Figure 3c indicates that the protons acquire a certain size while electroadsorbed. This was also observed in previous studies.7,8 It is reproduced here at the condition of minimal exposure to air between treatments. The most important result in this series is that shown in Figure 3c for a carbon electrode that was activated by CO2 at 900 °C for 30 min. There is a zero capacity in all electrolyte solutions except the case of a 0.1 M H2SO4 solution at negative potentials in which pronounced electroadsorption currents are observed, corresponding to proton electroadsorption. At this range, the differential capacity reaches values as large as 100 F/g. Further activation as shown in Figures 3d-g leads to pore-size enlargement, as presented in our previous works.7,8,10 Accordingly, the proton electroadsorption selectivity is weakened gradually, as a result of the buildup of capacity for electroadsorption of other ions as well. Thus, Figure 3d,e shows electroadsorption of Na+ ions. Subsequently, the electroadsorption of the sizable sulfate anion takes place, Figure 3f, and finally, the proton adsorption stereoselectivity practically disappears for all ions tested, as shown in Figure 3g. We realize here that (a) the proton electroadsorption selectivity is a molecular sieving effect, which disappears when the pores are sufficiently large, and (b) the carbon employed in the series presented in Figure 3 is, at least initially, free of chemisorbed oxygen and, therefore, the proton specificity is not based on surface oxygen groups. In other words, in narrow pores that cannot accept highly hydrated ions, nonspecific (nonchemical) interactions are responsible for the proton electroadsorption selectivity. Realizing that a protic environment within the pore system is not necessary for selective proton electroadsorption raises the question what compensates for the loss of hydration interaction of the protons which are inserted into the narrow pores from the bulk. We may think of two mechanisms which may act in parallel: (A) Coulombic interactions with a negative space charge at the carbon bulk, as occurs generally during the course of ion electroadsorption at the electrical double layer, and (B) physisobed water, not necessarily in clusters,15 as a hydrogen bond support for the electroadsorbed protons. The following water vapor adsorption isotherms will serve to examine mechanism B. In Figure 4, a series of water vapor isotherms at room temperature is shown for different degrees of activation. These type-V isotherms16 are known to indicate17 (i) weak adsorbate-to-surface
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interaction, leading to very little adsorption at low relative pressures; (ii) relatively strong lateral interaction, leading to enhanced adsorption at higher relative pressure; and (iii) the adsorbent being a microporous or submicroporous solid, which causes the isotherm to level off at high relative pressures. Properties i and ii correspond to the properties of water18 absorbed into the pristine carbon and to the slightly, 30-min activated carbon, thus, exhibiting the lowest isotherms, while the more strongly activated carbon (CO2 activation at 900 °C for 90 min) shows a high isotherm. The initial slopes are an indication of the concentration of initial nuclei of surface groups that serve as anchors for the water adsorption, prior to cluster formation at the higher pressures. Thus, the HNO3 activated carbon has the highest initial slope, corresponding to high surface group concentration.16 Most interestingly, the pristine carbons, both Kaptonand cellulose-based, still exhibit significant water capacity, indicating that there are suitable pores that can accommodate the small water molecule. Because according to Figure 3a,b the pristine carbons show no proton adsorption, we conclude that the water molecules are effectively smaller than the protic moiety. Because, in turn, the proton is the smallest of the ions studied (Figure 3c), it is likely that the proton is a protium H3O+ ion in the smallest pores which can accommodate it. Thus, mechanism B is effective in compensating for the loss in high hydration energy upon entering the pores. Because the hydronium ion assumes a pyramidal symmetry,19 it is conceivable that it may acquire a larger effective size than the water molecule, assuming that the latter is a planar molecule. This assumption contradicts an earlier view introduced by Pauling,20 commonly cited in the literature,21 according which the water molecule, by means of the two lone bonds of the oxygen, acquires a tetrahedral sp3 symmetry. Later theoretical studies suggest that the O-H bond is an sp2 type22 and the lone bond orbitals are too small to compare to the O-H bond length.23 So far there is no single model that predicts all the properties. Nevertheless, adding the van der Waals radius of hydrogen (1.2 Å)24 to the height of the pyramidal hydronium cation leaves no doubt as to its larger effective dimension than that of the water molecule. Neglecting the effect of the lone bonds, the effective size of the water molecule is determined by the size of the large oxygen atom. The van der Waals radius of hydrogen is perpendicular to this dimension and, thus, has no influence on the effective dimension of water. (16) Gregg, S. J. Adsorption, surface area and porosity; Academic Press, Inc.: London, 1982; p 5. (17) (a) Avgul, N. N.; Dzhigit, O. M.; Kiselev, A. V.; Shcherbakova, K. D. Dokl. Akad. Nauk SSSR 1953, 92, 105. (b) Avgul, N. N.; Dzhigit, O. M.; Kiselev, A. V.; Shcherbakova, K. D. Dokl. Akad. Nauk SSSR 1955, 101, 285. (c) Avgul, N. N.; Kiselev, A. V. In Chemistry and Physics of Carbon; Walker, P. L., Jr., Ed.; Marcel Dekker: New York, 1970; Vol. 6, p 1. (18) (a) Avgul, N. N.; Dzhigit, O. M.; Kiselev, A. V.; Shcherbakova, K. D. Dokl. Akad. Nauk SSSR 1953, 92, 105. (b) Avgul, N. N.; Dzhigit, O. M.; Kiselev, A. V.; Shcherbakova, K. D. Dokl. Akad. Nauk SSSR 1955, 101, 285. (c) Avgul, N. N.; Kiselev, A. V. In Chemistry and Physics of Carbon; Walker, P. L., Jr., Ed.; Marcel Dekker: New York, 1970; Vol. 6, p 1. (19) Sobolewski, A. L.; Domcke, W. J. Phys. Chem. A 2002, 106, 4158. (20) Pauling, L. The Nature of Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, 1960; pp 108-111. (21) Mtcalfe, H. C.; Williams, J. E.; Castka, J. F. Modern Chemistry; Holt Rinehart and Winston: New York, 1982; p 126. (22) Chaplin, M. http://www.lsbu.ac.uk/water/models.html (accessed Feb 2005). (23) Laing, M. J. Chem. Educ. 1987, 64, 124. (24) Pauling, L. The Nature of Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, 1960; pp 275-264.
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Summary (i) Even in nonpolar carbons, water is adsorbed in the smallest pores of nonactivated cellulose- and Kapton-based carbons, whereas the pores are too tight to accommodate any ion, including protons. (ii) Porous carbons with an oxide-free surface still support significant electroadsorption of protons, most likely due to hydration by physisorbed water.
Notes
(iii) The hydrated protons within the pores are effectively larger than the water molecules but are much smaller than fully hydrated protons in solution. Acknowledgment. Partial support for this work was obtained from the Israel-USA Binational Science Fund (BSF). LA049238H
