H-Bonding of Furan and Its Hydrogenated Derivatives with the Isolated

Oct 5, 2010 - ... Abruzzi 24, I-10129 Torino, Italy, and Dipartimento di Meccanica e Materiali, ... The values measured for ΔH0 are compared to the e...
0 downloads 0 Views 1MB Size
Download PDF
J. Phys. Chem. C 2010, 114, 18233–18239

18233

H-Bonding of Furan and Its Hydrogenated Derivatives with the Isolated Hydroxyl of Amorphous Silica: An IR Spectroscopic and Thermodynamic Study F. Mauriello,†,‡ M. Armandi,† B. Bonelli,† B. Onida,† and E. Garrone*,† Dipartimento di Scienza dei Materiali ed Ingegneria Chimica, Politecnico di Torino and INSTM Unit Torino Politecnico, Corso Duca degli Abruzzi 24, I-10129 Torino, Italy, and Dipartimento di Meccanica e Materiali, Facolta` di Ingegneria, UniVersita` Mediterranea di Reggio Calabria, Localita Feo di Vito, I-89122 Reggio Calabria, Italy ReceiVed: July 20, 2010; ReVised Manuscript ReceiVed: September 14, 2010

The interactions of furan and its hydrogenated derivatives (2,3-dihydrofuran, 2,5-dihydrofuran, tetrahydrofuran, and diethyl ether) with the isolated silanol groups at the surface of a highly dehydrated amorphous silica have been investigated by means of IR spectroscopy, at both constant and variable temperature. The adsorbate molecules studied form H-bonds of different strengths with surface silanols. With hydrogenated derivatives, the interaction involves the basic O atom, whereas with the aromatic furan molecule, the interaction probably occurs at the CdC double bond. The spectroscopic features of the stretching mode of H-bonded silanols change with the interacting molecule: classical correlations are observed. Because isolated silanols act as an ideal thermodynamic ensemble, the Langmuir model applies to all interactions. The fraction of silanols engaged in H-bonding can be measured from the decrement in intensity of the 3745 cm-1 band, so that corresponding optical isotherms can readily be drawn. The Langmuir model allows the calculation of the corresponding equilibrium constants and related standard free energies of interaction (∆G0) at ambient temperature. IR measurements run at variable temperature allow the evaluation of the corresponding standard enthalpies, ∆H0, and entropies, ∆S0. The values measured for ∆H0 are compared to the enthalpies of interaction due to the sole H-bond, as calculated from the ∆ν(OH) values: H-bonding was found to account for about twothirds of the total ∆H0 value, the difference being due to the contribution of dispersive forces. The interaction of 2,5-dihydrofuran was found to be slightly more energetic than expected, probably because this molecule can establish dispersive interactions with the silica surface more efficiently than the other molecules. 1. Introduction The surface chemistry of amorphous silica is a very popular topic, and the related literature is immense.1-3 The structure of hydrated surfaces is rather complex, in part because of the pliability of the SisOsSi bond angle, which gives rise to a plethora of possible configurations. As a consequence, ab initio descriptions of the silica surface have become available only very recently.4 The situation becomes much simpler when severely outgassed (dehydrated) samples are considered: a single hydroxyl species (hereafter referred to as tSiOH) is found at the surface of all types of amorphous silica with a nearly constant concentration of ca. 1 tSiOH nm-2 whatever the type of silica, and the corresponding OsH stretching mode is observed in the narrow range of 3747-3745 cm-1.5-7 The isolated tSiOH species is probably the most studied surface object ever: because of its simple structure, the isolated silanol group and its interactions with several molecules were modeled ab initio long ago by cluster approaches.8 The isolated tSiOH groups of amorphous silica offer yet another reason of interest. Being basically all equal and noninteracting, they constitute an ideal ensemble and, therefore, allow simple thermodynamic considerations to be drawn.9 For instance, in an adsorption experiment, it is straightforward to calculate the fraction of free (noninteracting) silanol groups, 1 * Corresponding author. E-mail: [email protected]. Tel.: +39011-5644661. Fax: +39-011-5644699. † Politecnico di Torino and INSTM Unit Torino Politecnico. ‡ Universita` Mediterranea di Reggio Calabria.

- θ, by comparing the intensity of the 3745 cm-1 band, under given circumstances, to the intensity of the same band in the absence of interactions. (Note that θ is the fraction of silanol species engaged.) In the case of carbonylic compounds,9 acetylenic molecules,10 and olefins,11 adsorption of light molecules involves only the tSiOH species, and the process strictly follows a Langmuirtype behavior. With heavier molecules, in addition to the formation of H-bonds with tSiOH groups, physisorption takes place on the rest of the surface, as well as solvation of the already H-bonded molecules, thus causing deviations from the Langmuir-type behavior. From a thermodynamic point of view, Langmuir-type isotherms yield the equilibrium constant at room temperature and, hence, the value of ∆G0. In the past several years, however, variable-temperature infrared (VTIR) spectroscopy has become available, both below and above the ambient temperature, which allows the determination of both the adsorption enthalpy, ∆H0, and entropy, ∆S0, still in the case of adsorption phenomena on an ideal ensemble of sites.12-18 In the present work, both ambient and variable-temperature IR spectroscopic techniques were used to study the interactions of tSiOH groups with a series of closely related molecules, differing by the addition of two hydrogen atoms from term to term. These molecules are (i) furan, (ii) 2,3-dihydrofuran (2,3DHF) and 2,5-dihydrofuran (2,5-DHF), (iii) tetrahydrofuran (THF), and (iv) diethyl ether (DEE). The interaction of such compounds with tSiOH groups has not yet been reported. This set was chosen because their different

10.1021/jp106746a  2010 American Chemical Society Published on Web 10/05/2010

18234

J. Phys. Chem. C, Vol. 114, No. 42, 2010

Mauriello et al.

TABLE 1: Molecular Properties of Adsorbates adsorbate

BPa (K)

IEb (eV)

PAc (kJ mol-1)

pKH-Bd

DEE THF 2,5-DHF 2,3-DHF furan

307.7 339.1 340.5 327.7 304.5

9.51 9.41 9.14 8.3 8.88

828.4 822.1 823.4 866.9 803.4

1.01 1.28 0.53 -0.40

a Boiling point. b Ionization energy. c Proton affinity. Equilibrium constant for H-bond formation with 4-fluorophenol (see text). d

structures allow the basicity of the O atom to be tuned, without significantly affecting the role of dispersive interactions. Indeed, their boiling points are rather close (Table 1). A hot topic in computational chemistry concerns the recent availability of density functional theory (DFT) kernels able to deal with dispersion forces,19 in contrast with those adopted so far, which are able to account for H-bonding only when dealing with intermolecular interactions. For this reason, experimental data on weak adsorption interactions, such as those under consideration, estimating the contributions of both H-bonding and dispersion to the total energy should be of interest to the chemical community.

Figure 1. IR spectra in the OsH stretching region (3800-3100 cm-1) for the adsorption of 2,3-DHF at constant (ambient) temperature on amorphous silica outgassed at 1073 K. Dotted line, spectrum of the bare sample; solid lines: spectra recorded at 2,3-DHF equilibrium pressures in the range of 0.1-10 mbar; arrows, spectral changes with increasing 2,3-DHF equilibrium pressure.

2. Experimental Section All organic chemicals were purchased from Sigma-Aldrich and used without further purification. Some relevant molecular properties are gathered in Table 1. The amorphous silica sample was purchased from Aldrich (specific surface area: 390 m2g-1, pores volume: 0.75 cm3 g-1): for IR measurements, the powder was pressed into a thin selfsupporting wafer (surface density of ca. 5 mg cm-2) and activated under high vacuum (residual pressure below 10-3 mbar) at 1073 K for 60 min in an IR cell allowing to carry out thermal treatments prior to any measurement. IR spectra at constant (ambient) temperature were recorded, at 2 cm-1 resolution on a Bruker FTIR Equinox 55 spectrometer, equipped with a MCT cryodetector. Equilibrium pressures were measured with a Thermovac TM 20 pressure gauge: experiments were usually run by increasing stepwise the pressure up to the vapor pressure of the organic molecule at ambient temperature. VTIR spectroscopy measurements were carried out on a commercial IR cell (AABSPEC)16 equipped with a capacitance pressure gauge (CTR100, Oerlikon-Leybold) and an electronically controlled heating element. The temperature inside the cell was monitored by a K-thermocouple (in contact with the sample wafer) connected to a digital thermometer (CHY 502 A, Tersid). The accuracy of the pressure and temperature measurements was (0.20% and (0.05% (of the read value), respectively. After activation of the sample, the cell was allowed to cool down to ambient temperature and dosed with the appropriate amount of the adsorbate considered. Subsequently, the cell was closed and IR spectra were recorded at several fixed temperature values within the range 300-400 K. 3. Results and Discussion The surface hydroxyl group of silica is a rather weak acid, the pKa of which is estimated to be around 7,20 so that proton transfer is never observed in gas-solid interactions, nor are the IR intricacies of really strong H-bonds, such as the presence of Evans windows.21 Instead, plain formation of H-bonds is expected.

Figure 2. Comparison of IR spectra in the 3800-2600 cm-1 region recorded at constant (ambient) temperature after adsorption of different adsorbates on amorphous silica outgassed at 1073 K. Selected spectra are presented to allow comparison.

3.1. IR Spectroscopic Features. Figure 1 reports IR spectra of the adsorption of 2,3-DHF taken at constant (ambient) temperature. Upon increase of the equilibrium pressure, the 3745 cm-1 band decreases because of the increase of a broad absorption centered at 3385 cm-1 due to H-bonded hydroxyls. The band at 3385 cm-1 has a constant shape and position all along the experiment, indicating the absence of any solvation phenomena. Its half-width (W1/2) is 235 cm-1, which is 24 times larger than that of the pristine band at 3745 cm-1. Figure 2 compares the bands due to H-bonded silanol species in the five cases under study. Spectra were chosen so as to have bands of comparable intensity. Bands in the region 3000-2800 cm-1 are due to stretching modes of saturated CsH species, whereas the weak band at ca. 3100 cm-1 is the corresponding mode of unsaturated CsH species. All bands concerning CsH motions will not be considered further here. Similar features are observed in all cases: The half-width, W1/2, and intensification ratio, Ai/A0, for each case are reported in Table 2. The values

H-Bonding of Furan and Derivatives with Amorphous Silica TABLE 2: IR Spectroscopic Features of the Stretching Mode of H-Bonded Silanols adsorbate

∆ν(OH)a (cm-1)

W1/2b (cm-1)

Ai/A0c

DEE THF 2,5-DHF 2,3-DHF Furan

494 503 467 360 118

320 333 288 235 126

12.4 12.4 11.1 8.7 2.3

a

Bathochromic shift of the OsH stretching mode with respect to the unperturbed mode at 3745 cm-1. b Width at half-height. c Relative increase in specific intensity.

Figure 3. Relationships between the spectral features of the OsH stretching mode of H-bonded silanols (half-width, W1/2, and intensity ratio, Ai/A0) and the shift of the 3745 cm-1 band, ∆ν(OH) (cm-1).

of Ai/A0 were calculated, at low equilibrium pressures, as the ratio between the increment in integrated area of the band formed after adsorption (Ai) and the decrement in integrated area of the free silanol band at 3745 cm-1 (A0). Figure 3 shows that both of these features are steady functions of the main descriptor of the interaction, the frequency shift ∆ν(OH), in agreement with the literature.21,22 3.2. Nature of the Interaction and Correlation with Molecular Properties. In the case of saturated molecules (both DEE and THF), interaction surely occurs at the basic oxygen atom. Shifts of similar magnitude are indeed observed with sister molecules, such as dioxane (unpublished results). With either unsaturated or aromatic molecules, interaction could in principle occur with other electron-rich parts of the molecule, such as the double bond. Previous work11 has shown, however, that the ∆ν(OH) value due to the interaction of isolated silanols with CdC double bonds does not exceed 200 cm-1. It is inferred for the two DHFs that interaction also involves the O atom and that different values of ∆ν(OH) reflect different basicities. The values of ∆ν(OH) are basically the same for DEE and THF (494 and 503 cm-1, respectively), as expected. That of 2,5-DHF (467 cm-1) is somewhat lower, probably because of some strain in the ring caused by the presence of the shorter CdC double bond with respect to the single CsC bond. That of 2,3-DHF (360 cm-1) is instead substantially lower, as the adjacent double bond allows delocalization of the oxygen lone pair. For furan, the observed value of ∆ν(OH) ) 136 cm-1 is compatible with the interaction with the electronic π system,

J. Phys. Chem. C, Vol. 114, No. 42, 2010 18235 instead of the O atom. Comparison with previous work11 suggests interaction with a CsC bond, the order of which is between 1.5 (benzene) and 2 (cyclohexene), in agreement with the CsC bond order in furan. Furan appears to behave rather as a π H-bonding acceptor than an oxygen base, in agreement with what has been reported in the literature for five-membered N-heterocycles:23 The introduction of one (or two) double bonds has a marked effect on the H-bonding basicity, turning the molecules from strong nitrogen bases to very weak π-bases. Although an intermolecular interaction dominated by electrostatics,24,25 H-bonding is known to imply some elongation of the OsH bond (describable as an incipient protonation of the H-acceptor molecule), as well as some transfer of electron charge from the molecule to the acidic hydroxyl group (describable as an incipient ionization of the former). The values of ∆ν(OH) are therefore traditionally related to either the proton affinity (PA, kJ mol-1) or the ionization energy (IE, eV) of the H-acceptor molecule. Accordingly, Cusumano and Low26 long ago tested the charge-transfer-no-bond theory of the H-bond proposed by Mulliken,27 by showing that ∆ν(OH) is linearly correlated with the ionization energy within a series of parent molecules. Figure 4a shows that, indeed, a linear relationship exists between the value of ∆ν(OH) and the ionization energies of DEE, THF, 2,5DHF, and 2,3-DHF. The corresponding value for furan does not fall on the same line, in agreement with the different nature of this molecule. On the other hand, Paukshtis et al.28 reported a correlation (though somewhat coarse) between ∆ν(OH) and the gas-phase proton affinity of the basic molecule, for both the tSiOH species and the Brønsted site in zeolites. An analogous plot for the present case is reported in Figure 4b: As for the correlation with ionization energy, a linear behavior is seen for all of the molecules studied in the present work, with the exception of furan. Another type of correlation is as follows: By means of IR spectroscopy, Berthelot and co-workers29 long ago studied the H-bonding equilibrium in solution for 39 ethers with 4-fluorophenol, the latter being assumed as a reference H-bond donor. The reported values of pKH-B (Table 1) constitute a reference scale for similar interactions in solution. Figure 4c shows that the pKH-B values correlate well with the values measured for ∆ν(OH), even though the former refer to gas-phase interactions. 3.3. Standard Thermodynamic Quantities of Adsorption. The reversible H-bonding interaction can be written as

tSiOH + M(g) a tSiOH · · · M

(1)

where M stands for the adsorbate molecule. Let this process be characterized by the equilibrium constant K. The assumption of the ideality of the ensemble of tSiOH groups leads to the corresponding Langmuir equation

K ) θ/[(1 - θ)(p/p*)]

(2)

with θ being the fraction of silanol species engaged in the interaction and p* being the reference pressure (1 mbar in the present case). The quantity (1 - θ) is the main IR spectroscopic observable, estimated, as described above, as A/Amax, where A is the intensity of the free hydroxyl peak at any given coverage and Amax is that in the bare sample. For all of the compounds considered here, the Langmuir equation is obeyed in the whole range of pressures explored

18236

J. Phys. Chem. C, Vol. 114, No. 42, 2010

Figure 4. Relationships between some relevant molecular properties of the adsorbates studied and the shift of the 3745 cm-1 band, ∆ν(OH) (cm-1): (a) ionization energy (eV); (b) proton affinity (kJ mol-1); (c) pKHB, as defined in the text.

(0.1-10 mbar, corresponding to 10-4-10-3 p/p0 values for all of the examined compounds). Indeed, a plot of the quantity θ/(1 - θ) as a function of the equilibrium pressure yields a straight line in each case (Figure 5). The slopes of these straight lines give the values of the equilibrium constants and, hence, the corresponding ∆G0 values, because ∆G0 ) -RT ln K (Table 3). Note that the temperature is actually not known with accuracy, because of the heating effect of the IR beam, and is expected to be slightly higher than the ambient. For the evaluation of ∆G0, the temperature was assumed to be 300 K; this point, however, will be considered further below. To measure both enthalpic and entropic contributions to ∆G0, the VTIR method was employed. Let us consider a set of IR

Mauriello et al.

Figure 5. Langmuir tests, (1 - θ0)/θ0 versus equilibrium pressure, of the optical isotherms obtained at ambient temperature by dosing the adsorbates studied on amorphous silica outgassed at 1073 K. (a) Furan (solid circles), 2,3-DHF (open squares), 2,5-DHF (gray triangles), DEE (solid squares), THF (open circles). (b) Magnification of the data for furan (black circles).

spectra, such as those in Figure 6, taken over a temperature range in a closed IR cell.13 First, the physical effects of changing temperature must be taken into account. As shown in Figure 6, these are remarkable for the band of H-bonded silanol species. In contrast, they are rather limited for free tSiOH species: A linear hypsochromic shift is observed30 with increasing temperature, with a temperature coefficient (dν/dT) of only 0.0176 cm-1 K-1. The thermal shift is accompanied by a broadening of the tSiOH band. Fortunately, the corresponding integrated intensity, A, remains essentially constant with temperature, so that the evaluation of the fraction of free tSiOH species (1 - θ) is not affected by physical effects. If the equilibrium pressure, p, is known at a given temperature, T, then the equilibrium constant, K, can be determined

H-Bonding of Furan and Derivatives with Amorphous Silica

J. Phys. Chem. C, Vol. 114, No. 42, 2010 18237

TABLE 3: Energetic Features of the H-Bonding Process adsorbate

Ka (mbar-1)

∆G0 b (kJ mol-1)

-∆H0 c (kJ mol-1)

-∆S0 d (J mol-1 K-1)

∆ν(OH) (cm-1)

-∆H* e (kJ mol-1)

T* f (K)

-∆Hdispg (kJ mol-1)

DEE THF 2,5-DHF 2,3-DHF furan

1.62 3.62 0.97 0.24 0.023

-1.20 -3.21 0.076 3.56 9.40

46.9 47.2 49.3 39.6 -

153 150 163 146 -

494 503 467 360 118

32.0 32.4 30.9 26.1 11.5

299 298 302 296 -

14.9 14.8 18.4 13.5 -

a Room-temperature equilibrium constant (reference pressure ) 1 mbar). b Standard free energy of interaction. c Standard adsorption enthalpy, as measured by VTIR spectroscopy. d Corresponding standard adsorption entropy. e Contribution of H-bonding to the adsorption enthalpy, as calculated from the measured bathochromic shift of the OsH stretching band with respect to the unperturbed band at 3745 cm-1. f Estimated temperature of IR measurements carried out at nominally constant (ambient) temperature (see text). g Estimated contribution of dispersion forces to the adsorption enthalpy.

measure, T*, in the IR measurements carried out nominally at ambient temperature, by writing

∆G0 ) -RT* ln K ) ∆H0 - T*∆S0

(5)

from which it results that

T* ) ∆H0 /(∆S0 - R ln K)

Figure 6. IR spectra recorded at variable temperature after dosing ca. 8.810 mbar of 2,5-DHF on amorphous silica outgassed at 1073 K. From bottom to top: spectra recorded at temperatures in the 298-373 K range and at equilibrium pressures increasing from 8.810 to 11.530 mbar; dotted curve, spectrum of bare silica before 2,5-DHF dosage. Inset: Magnification of the free silanol absorption band at 3745 cm-1, showing how it changes with both temperature and equilibrium pressure.

through eq 2. The variation of K with temperature leads to the corresponding values of adsorption enthalpy and entropy, through the well-known van’t Hoff equation (eq 3). Combining eqs 2 and 3 yields eq 4

K(T) ) exp(-∆H0 /RT) exp(∆S0 /R)

(3)

ln[θ/(1 - θ)p] ) -∆H0 /RT + ∆S0 /R(c)

(4)

from which both ∆H0 and ∆S0 can be derived, under the usual assumption of their being constant with temperature. VTIR data plotted according to eq 4 for DEE, the two DHFs, and THF are reported in Figure 7. In all four cases, a satisfactory linearity is observed. Note that no undefined parameter enters the calculation. The corresponding thermodynamic results are reported in Table 3. The interaction of furan was not studied by VTIR spectroscopy, because it is exceedingly weak and thus probably requires the use of a cell operating at temperatures below ambient. Table 3 shows that the energy of the interaction is in the order 2,5-DHF > DEE ≈ THF > 2,3-DHF . furan. The interaction of 2,5-DHF is (slightly) more energetic than expected. A possible reason is discussed below. To check the consistency of the VTIR data with those obtained at constant (ambient) temperature, the values of ∆G0, ∆H0, and ∆S0 were used to calculate the actual temperature of

(6)

Calculated values of T* for the four cases are reported in Table 3. Satisfactory agreement is observed among them, and all result just slightly higher than room temperature, as expected. We assume this finding to be evidence for the consistency of the obtained thermodynamic results. 3.3. Contribution to the Enthalpy of Interaction. It is commonly accepted that ∆ν(OH), the shift in the stretching mode of isolated silanol species, is a measure of the strength of the H-bonding interaction. The quantitative evaluation of such enthalpy of interaction from this spectroscopic finding was a popular topic in the early 1970s. Two approaches have been used: Hair and Hertl31 measured isosteric heats of interaction, qst, with tSiOH groups by the use of the classical relationships

(d ln p/dT)θ ) -qst/RT2 and -qst ) ∆H0 + RT

(7)

where the constancy of coverage θ was monitored through the intensity of the 3745 cm-1 peak. A different behavior was found with different families of compounds, according to the sp2 or sp3 nature of the atom engaged in H bonding. For instance, with adsorbates in which the H-acceptor oxygen atom is sp2-hybridized, the proposed correlation is

∆H0 ) -2.3 + 0.455[∆ν(OH)]1/2

(8)

Curthoys et al.32 instead used direct calorimetry. To discriminate the actual process of interaction with tSiOH groups from plain physisorption, ∆H0 was defined as the difference between the heats of adsorption on a sample fully covered with hydroxyls and the same sample after dehydration. For all substances investigated, the following relationship was proposed

∆H0 ) -1.6 + 0.41[∆ν(OH)]1/2

(9)

Fortunately, eqs 8 and 9 essentially agree and yield, in the range of ∆ν(OH) values observed, results coinciding within few percent. From the ∆ν(OH) values in Table 2, a set of enthalpy

18238

J. Phys. Chem. C, Vol. 114, No. 42, 2010

Mauriello et al.

Figure 7. Plot of the left-hand side of eq 4 against reciprocal temperature for the interactions of (a) 2,5-DHF, (b) THF, (c) 2,3-DHF, and (d) DEE.

changes were calculated (Table 3). These values, termed ∆H*, are representative of the contribution related solely to H-bonding and indeed were found to be smaller than the measured ∆H0 values. The difference between ∆H0 and ∆H*, reported in the last column of Table 3, represents the contribution from dispersive (van der Waals) interactions. It was found that, in the cases of DEE, THF, and 2,3-DHF, these contributions are very close, as expected because of the very similar chemical formulas. Only in the case of 2,5-DHF does the contribution of dispersive interactions appear larger. This can be probably understood by considering the cartoon in Scheme 1, which shows that this molecule (and only this one) can assume a conformation on the surface by which the double bond lies parallel to the surface itself, thus displaying full interaction. Note that the OsH atoms of silanol and the O atom of 2,5-DHF probably do not line up: A partially bent conformation has been shown by Legon and co-workers33,34 to occur in similar adducts between hydrogenated furan derivatives and hydrogen fluoride by microwave spectroscopy. 4. Conclusions H-bonding between the isolated silanol species at the surface of amorphous silica and furan derivatives was observed in all

SCHEME 1: Cartoon Depicting a 2,5-DHF Molecule in Interaction with a Silanol Group at the Surface of Silica

cases examined in this work. With hydrogenated derivatives, interaction involves the basic O atom, whereas with the aromatic furan molecule, interaction probably occurs at the CdC double bond. Classical correlations among spectroscopic features of the stretching mode of H-bonded silanol groups were observed. The Langmuir model applies for all interactions. This allowed the calculation of the equilibrium constants and the corresponding standard free energies of interaction (∆G0) at ambient temperature, whereas IR measurements run at variable temper-

H-Bonding of Furan and Derivatives with Amorphous Silica ature yielded the corresponding standard enthalpies, ∆H0, and entropies, ∆S0. For the former, the contribution due solely to H-bonding accounts for about two-thirds of the total value, the difference being due to dispersive contributions. The interaction of 2,5-DHF is slightly more energetic than expected, probably because this molecule can establish a dispersive interaction with the surface more efficiently than the other molecules. The present results yielding information on the roles of both dispersive and H-bonding contributions to the interaction energies in silanol/furan derivative adducts could afford a benchmark for computational work using newly available DFT kernels. Acknowledgment. Financial support from INSTM (Istituto Nazionale di Scienza e Tecnologia dei Materiali) Consortium is gratefully acknowledged. References and Notes (1) Papirer, E. Adsorption on Silica Surfaces; Surface Science Series; Marcel Dekker AG: New York, 2000; Vol. 90. (2) Iler, R. K. The Chemistry of Silica: Solubility, Polymerization, Colloid and Surface Properties and Biochemistry of Silica; John Wiley & Sons: New York, 1979. (3) Zhuravlev, L. T. Colloids Surf. A 2000, 173, 1. (4) Ugliengo, P.; Sodupe, M.; Musso, F.; Bush, I. J.; Orlando, R.; Dovesi, R. AdV. Mater. 2008, 20, 1. (5) Hoffman, P.; Kno¨zinger, E. Surf. Sci. 1987, 188, 181. (6) Morrow, B. A.; Cody, I. A. J. Phys. Chem. 1976, 80, 1995. (7) Morrow, B. A.; Cody, I. A. J. Phys. Chem. 1976, 80, 1998. (8) Civalleri, B.; Garrone, E.; Ugliengo, P. Chem. Phys. Lett. 1998, 294, 103. (9) Allian, M.; Borello, E.; Ugliengo, P.; Spano`, G.; Garrone, E. Langmuir 1995, 11, 4811. (10) Onida, B.; Allian, M.; Borello, E.; Ugliengo, P.; Garrone, E. Langmuir 1997, 13, 5107. (11) Garrone, E.; Barbaglia, A.; Onida, B.; Civalleri, B.; Ugliengo, P. Phys. Chem. Chem. Phys. 1999, 1, 4649.

J. Phys. Chem. C, Vol. 114, No. 42, 2010 18239 (12) Tsyganenko, A. A.; Storozhev, P. Y.; Area´n, C. O. Kinet. Catal. 2004, 45, 530. (13) Storozhev, P. Y.; Area´n, C. O.; Garrone, E.; Ugliengo, P.; Ermoshin, V. A.; Tsiganenko, A. A. Chem. Phys. Lett. 2003, 374, 439. (14) Area´n, C. O.; Manoilova, O. V.; Bonelli, B.; Delgado, M. R.; Palomino, G. T.; Garrone, E. Chem. Phys. Lett. 2003, 370, 631. (15) Garrone, E.; Area´n, C. O. Chem. Soc. ReV. 2005, 34, 846. (16) Area´n, C. O.; Tsyganenko, A. A.; Platero, E. E.; Garrone, E.; Zecchina, A. Angew. Chem., Int. Ed. 1998, 37, 3161. (17) Area´n, C. O.; Manoilova, O. V.; Palomino, G. T.; Delgado, M. R.; Tsyganenko, A. A.; Bonelli, B.; Garrone, E. Phys. Chem. Chem. Phys. 2002, 4, 5713. (18) Armandi, M.; Garrone, E.; Area´n, C. O.; Bonelli, B. ChemPhysChem 2009, 10, 3316. (19) Grimme, S. J. Comput. Chem. 2006, 27, 1787. (20) Hair, M. L.; Hertl, W. J. Phys. Chem. 1970, 74, 91. (21) Pimentel, G. C.; McLellan, A. L. The Hydrogen Bond; W. H. Freeman and Co.: San Francisco, CA, 1960. (22) Kno¨zinger, H. In The H-Bond: Recent AdVances in Theory and Experiment; Schuster, P., Zundel, G., Sandorfy, C., Eds.; North Holland: Amsterdam, 1976; Vol. 3, p 1269. (23) Arnaud, V.; Berthelot, M.; Felpin, F. X.; Lebreton, J.; Le Questel, J.-Y.; Graton, J. Eur. J. Org. Chem. 2009, 29, 4939. (24) Buckingham, A. D.; Fowler, P. W.; Hutson, J. M. Chem. ReV. 1988, 88, 963. (25) Buckingham, A. D.; Legon, A. C.; Roberts, S. M. Principles of Molecular Recognition; Blackie Academic and Professional: Glasgow, Scotland, 1993. (26) Cusumano, J. A.; Low, M. J. D. J. Catal. 1971, 23, 214. (27) Mulliken, R. J. Am. Chem. Soc. 1952, 74, 811. (28) Paukshtis, E. A.; Yurchenko, E. N. React. Kinet. Catal. Lett. 1981, 16, 131. (29) Berthelot, M.; Besseau, F.; Laurence, C. Eur. J. Org. Chem. 1998, 5, 925. (30) Ryason, P. R.; Russell, B. G. J. Phys. Chem. 1975, 79, 1276. (31) Hertl, W.; Hair, M. L. J. Chem. Phys. 1968, 72, 4676. (32) Curthoys, G.; Davydov, V. Ya.; Kiselev, A. V.; Kiselev, S. A.; Kuznetsov, B. V. J. Colloid Interface Sci. 1974, 48, 58. (33) Cooke, S. A.; Corlett, G. K.; Evans, C. M.; Holloway, J. H.; Legon, A. C. Chem. Phys. Lett. 1997, 275, 269. (34) Cooke, S. A.; Corlett, G. K.; Evans, C. M.; Legon, A. C. J. Chem. Soc., Faraday Trans. 1997, 17, 93.

JP106746A