The enthalpies of dissociation of the N-O bonds in two quinoxaline derivatives

Received: 5 May 2008, Revised: 2 June 2008, Accepted: 9 June 2008, Published online in Wiley InterScience: 10 July 2008 The enthalpies of dissociation of the N—O bonds in two quinoxaline derivatives Jose ´R. B. Gomes a,b *, Andre ´R. Monteiro a , Bruno B. Campos c , Paula Gomes a and Maria D. M. C. Ribeiro da Silva a * The present work reports the first experimental thermochemical study of mono-N-oxides derived from quinoxaline, namely, 3-methoxycarbonyl-2-methyl-quinoxaline N-oxide and 3-ethoxycarbonyl-2-methyl-quinoxaline N-oxide. The values of the enthalpies of formation, in the condensed state, and of the enthalpies of sublimation, derived from static bomb calorimetry and Calvet microcalorimetry measurements, respectively, are combined to derive the standard molar enthalpies of formation in the gaseous phase for these two compounds. From the latter values, the first and second N—O bond dissociation enthalpies for the corresponding di-N-oxides have been obtained. The gas-phase experimental results are also compared with calculated data obtained with a density functional theory approach. Copyright ß2008 John Wiley & Sons, Ltd. Supporting information may be found in the online version of this article. Keywords: enthalpy of formation; enthalpy of dissociation; density functional theory; calorimetry; synthesis INTRODUCTION Quinoxaline derivatives play an important role in the development of new therapeutic agents, due to their relevant antibacterial, antiviral or antifungal activities. The oxidation of one or both nitrogen atoms of the heterocyclic ring increases the diversity of their biological properties, since their potential to act as oxidizing agents leads to a large variety of therapeutic applications. These compounds are used as medicinal feed additives [1,2] and they are also used as bioreductive cytotoxic agents/species, [3,4] as a consequence of the hypoxia-selective activity evidenced by some quinoxaline di-N-oxides. [5–8] Over the last decade, experimental and theoretical energetic studies have been expanded to several quinoxaline 1,4-di-Noxides, in order to evaluate the influence of the chemical vicinity on the dissociation enthalpies of the N—O bonds. [9–16] However, experimental studies for quinoxaline 1-N-oxide derivatives have not been described so far, mainly due to the difficulty in obtaining convenient amounts of highly pure samples of those compounds. Recently, we have been able to synthesize very pure samples of two quinoxaline derivatives containing only a single dative N—O bond, by selective reduction of the corresponding 1,4-di-N-oxides previously prepared from benzofuroxan and the appropriate b-ketoester, allowing the experimental study of its molecular energetics by standard thermochemical approaches. Thus, the present work reports the first experimental thermochemical study of mono-N-oxide quinoxalines, namely, 3-methoxycarbonyl2-methyl-quinoxaline N-oxide, 1, and 3-ethoxycarbonyl-2-methylquinoxaline N-oxide, 2, with structures shown in Scheme 1. A static bomb calorimeter was used to determine the standard molar energies of combustion, in oxygen and at T¼298.15 K, of the two title compounds from which the values of the standard molar enthalpies of formation, in the condensed phase, were derived. Furthermore, the standard molar enthalpies of sublimation, at T¼298.15 K, were obtained by high temperature Calvet microcalorimetry using the vacuum sublimation drop method. Finally, combining the standard molar enthalpies of formation in the condensed phase and the corresponding enthalpies of sublimation allowed us to derive the standard (p o ¼0.1 MPa) molar enthalpies of formation in the gas phase, at T¼298.15 K, for both compounds. The latter values were used to obtain the experimental values for the first N—O bond dissociation enthalpy in the parent di-N-oxide quinoxalines since their corresponding standard molar enthalpies of formation in the gas phase were previously reported in the literature. [9,10] The experimental results for first N—O bond dissociation enthalpies were used to calibrate a computational approach that was used to calculate either the enthalpies of formation of the quinoxaline derivatives or of dissociation of the second N—O bond. (www.interscience.wiley.com) DOI 10.1002/poc.1419 Research Article *Correspondence to: J. R. B. Gomes or M. D. M. C. R. da Silva, Centro de Investigac¸a ˜oemQuı´mica, Departamento de Quı´mica, Faculdade de Cie ˆncias, Universidade do Porto, R. do Campo Alegre, 687, P-4169-007 Porto, Portugal. E-mail: [email protected]; E-mail: [email protected] aJ. R. B. Gomes, A. R. Monteiro, P. Gomes, M. D. M. C. R. da Silva Centro de Investigac¸a ˜oemQuı´mica, Departamento de Quı´mica, Faculdade de Cie ˆncias, Universidade do Porto, R. do Campo Alegre, 687, P-4169-007 Porto, Portugal bJ. R. B. Gomes CICECO, Departamento de Quı´mica, Universidade de Aveiro, 3810-193 Aveiro, Portugal cB. B. Campos Departamento de Quı´mica, Faculdade de Cie ˆncias, Universidade do Porto, R. do Campo Alegre, 687, P-4169-007 Porto, Portugal J. Phys. Org. Chem. 2009, 22 17–23 Copyright ß2008 John Wiley & Sons, Ltd. 17 EXPERIMENTAL DETAILS Synthesis and purification The quinoxaline di-N-dioxides (3and 4, Scheme 1) were prepared from benzofuroxan and the appropriate b-ketoester, following the method described by Robertson and Kasubick. [17] The di-N-oxides were then selectively monodeoxygenated with trimethyl phosphite, as described by Dirlam and McFarland, [18] to yield the corresponding mono-N-oxides (1and 2, Scheme 1). The products precipitated from the reaction mixture were isolated by suction filtration and recrystallized from methanol. The structure and purity of both 1and 2were confirmed by 1 H-NMR, 13 C-NMR and elemental analysis, as detailed in the Supplementary Material. Prior to calorimetric measurements, the compounds were further purified by vacuum sublimation. The fusion temperatures for compounds 1and 2, as determined by differential scanning calorimetry (DSC), are 382 K and 360–361 K, respectively. A redetermination of the temperatures of fusion on the cooled samples, already melted in the DSC, leads to the same values, showing that there was no decomposition during fusion. Theaverageratiosofthemassofcarbondioxiderecoveredtothat calculated from the mass of sample, together with the standard deviation of the mean, were: (0.9991 0.0005) and (1.0005  0.0007) for compounds 1and 2, respectively. The densities of both samples were assumed to be r¼1.0 g cm 3 (estimated from the weight and volume of a pellet for each compound). Combustion calorimetry The combustion experiments were performed with a static-bomb calorimeter originally assembled in the National Physical Laboratory, Teddington, U.K. [19] and in the University of Manchester, [20] and now installed in the Thermochemistry Laboratory in the University of Porto, Portugal. [21] The twin valve bomb, with an internal volume of 0.290 dm 3 and wall thickness of 1 cm, is made of stainless steel. Both the apparatus and the operating technique have been described already in the literature. [19–21] The energy equivalent of the calorimeter was determined from the combustion of benzoic acid (NIST Standard Reference Material 39j) with a massic energy of combustion, under standard bomb conditions, of (26434 3) J g 1 . From 11 calibration experiments, e(cal) ¼(15553.3 0.9) J K 1 , for an average mass of water added to the calorimeter of 2900.0 g. In all combustion experiments, samples in pellet form were ignited at T¼(298.150 0.001) K, in oxygen at a pressure p¼3.04 MPa, with a volume of 1.00 cm 3 of water added to the bomb. In the experiments made for 2,n-hexadecane (99þ%, Aldrich, r¼0.773 g cm 3 [22] ) was used as an auxiliary combustion material (standard massic energy of combustion, D c u o ¼(47141.4 3.6) J g 1 ) in order to produce an appropriate increase of temperature. The massic energy of combustion for the cotton thread fuse (r¼1.50 g cm 3 ,c p ¼1.48 J K 1 g 1 , M¼27.700 g mol 1 [23] ) of empirical formula CH 1.686 O 0.843 is D c u o ¼16250 J g 1 . [24] Corrections for nitric acid formation were based on 59.7 kJ g 1 for the molar energy of formation of 0.1 mol dm 3 HNO 3 (aq) from N 2 (g), O 2 (g), and H 2 O(l). [25] At T¼298.15 K, (@u/@p) T for the solid was assumed to be 0.2 J g 1 MPa 1 , a typical value for organic solids. The amount of substance used in each combustion experiment was determined from the total mass of carbon dioxide produced after allowance for that formed from the cotton thread fuse and hexadecane. For each experiment, the value of D c u o was Scheme 1. Synthetic route to quinoxaline mono-N-oxides www.interscience.wiley.com/journal/poc Copyright ß2008 John Wiley & Sons, Ltd. J. Phys. Org. Chem. 2009, 22 17–23 J. R. B. GOMES ET AL. 18 calculated by using the procedure given by Hubbard et al. [26] The relative atomic masses used were those recommended by the IUPAC Commission in 2005. [27] Calvet microcalorimetry The standard molar enthalpies of sublimation of the two N-oxide quinoxalines were measured using the vacuum sublimation drop-microcalorimetric technique. [28] Samples of about 3 mg of the crystalline compounds contained in a small thin glass capillary tube sealed at one end, and a blank capillary with similar mass, were simultaneously dropped at room temperature into the hot reaction vessel in the Calvet high-temperature microcalorimeter (Setaram HT 1000), held at T¼385 K in the case of 1 and at T¼378 K in the case of 2, and then removed from the hot zone by sublimation under reduced pressure. The thermal corrections for the glass capillary tubes were determined in separate experiments, and were minimized, as far as possible, by dropping tubes of nearly equal mass, to within 10 mg, into each of the twin calorimeter cells. The observed enthalpies of sublimation {Ho m(g,T)–Ho m(cr, 298.15 K)} were corrected to T¼298.15 K using D T298.15 Ho m(g) estimated by a group method of energy contributions based on data of Stull et al. [29] and Equ. 1. Finally, the microcalorimeter was calibrated in situ for these measurements by using naphthalene, Dg crHo m¼(72.5 0.1) kJ mol 1 . [30] Computational details The molecular structures of all compounds considered in this work have been initially optimized with the B3LYP method together with the 6-31G(d) basis set. [31,32] The geometries obtained have been characterized as true minima after the computation of the vibrational frequencies at the same level of theory. This set of calculations also yield the thermal corrections for T¼298.15 K that was added to the energy obtained with the same DFT approach but using a larger basis set. Thus, starting from the previously optimized structures, the B3LYP/6-311þ G(2d,2p) approach was used to further optimize the geometry and to obtain the energies of all compounds that were used throughout this work. The absolute enthalpies of all compounds which were used to estimate the standard molar enthalpies of formation and enthalpies of N—O bond dissociation are B3LYP/ 6-311 þG(2d,2p)//B3LYP/6-31G(d) values. All the calculations have been performed by means of the Gaussian 98 computer code. [33] RESULTS The results of a typical static bomb combustion experiment for each compound studied are given in Table 1. In this Table, DU S is the energy correction to the standard state and the rest of the terms have the same meaning as those described in reference [26]. The energy associated with the isothermal bomb process, DU(IBP), after ignition of the samples at the reference temperature was obtained from Equ. 2. In this equation Dm(H 2 O) is the deviation of the mass of water added to the calorimeter from 2900.0 g, i.e., mass assigned to e(cal), "fis the energy of the bomb contents after ignition, DT ad is the adiabatic temperature rise, and DU(ign) is the ignition energy: DUðIPBÞ¼f"ðcalÞþDmðH2OÞcpðH2O;lÞþ"fgDTad þDUðignÞ(2) The individual values of Dcuotogether with the mean value, hDcuoiand its standard deviations are given, for each compound, in Table 2, where Dcuorefers to the idealized combustion reaction yielding N 2 (g), CO 2 (g), and H 2 O(l). Unfortunately, only a small amount of 2was available for the combustion experiments and, therefore, only four experiments were possible. The standard molar energies, DcUo mðcrÞ, and enthalpies, DcHo mðcrÞ, of combustion at T¼298.15 K and in the condensed phase are given in Table 3. The standard molar enthalpies of formation in crystalline state, DfHo mðcrÞ, were derived from DcHo mðcrÞby using the standard molar enthalpies of formation of CO 2 (g) and H 2 O(l), at T¼298.15 K, (393.51 0.13) kJ mol 1 , [34] and (285.830 0.040) kJ mol 1 , [34] respectively. The DfHo mðcrÞ values are (262.4 4.2) kJ mol 1 and (302.7 7.3) kJ mol 1 for compounds 1and 2, respectively. In accordance with normal thermochemical practice, [35,36] the uncertainties assigned to the standard molar enthalpies of combustion are, in each case, twice ð1Þ Table 1. Typical combustion experiments at T¼298.15 K Compound 12 m(CO 2 , total)/g 1.16332 1.36218 m(compound)/g 0.52583 0.39024 m(fuse)/g 0.00275 0.00229 m(hexadecane)/g — 0.15130 DT ad /K 0.85617 1.12572 e f /(J K 1 ) 13.87 14.48 Dm(H 2 O)/g 0.7 0.0 DU(IBP)/J 13325.93 17524.93 DU(hexadecane)/J — 7141.34 DU(fuse)/J 44.66 37.19 DU(HNO 3 )/J 32.57 30.81 (U(ign)/J 0.73 1.19 (U S /J 10.21 9.21 ( c u8/DJg 1 ) 25174.98 26407.31 J. Phys. Org. Chem. 2009, 22 17–23 Copyright ß2008 John Wiley & Sons, Ltd. www.interscience.wiley.com/journal/poc N–O BOND DISSOCIATION ENTHALPIES 19 the overall standard deviation of the mean and include the uncertainties in calibration and in the values of auxiliary quantities used. The results from the Calvet microcalorimetry experiments for compounds 1and 2are reported in Table 4. The standard molar enthalpies of sublimation are, respectively, (118.1 3.3) kJ mol 1 and (129.2 4.1) kJ mol 1 . The standard molar enthalpies of formation in the gas phase and at T¼298.15 K for the two compounds under study are obtained by summing the data given in the last columns of Tables 3 and 4. Thus, the DfHo mðgÞvalues are (144.3  5.3) kJ mol 1 and (174 8) kJ mol 1 for compounds 1and 2, respectively, and appear in Table 5. Note that the last value is anchored on a standard enthalpy of formation in the crystalline state obtained from only four experiments. The first standard molar N—O bond dissociation enthalpy in the gas phase, DH 1 (N—O), for the quinoxaline di-N-oxides 3 and 4are obtained from the standard molar enthalpies of formation of the species appearing in the gaseous reaction described by Equ. 3. ð3Þ The results for compounds 3and 4were determined in previous works and their DfHo mðgÞvalues are (148.7 3.2) kJ mol 1 , [9] and (178.0 4.3) kJ mol 1 , [10] respectively. Those for compounds 1and 2are given in Table 5. The enthalpy of formation for atomic oxygen was taken from the literature, DfHo mðgÞ¼ 249.18 0.10 kJ mol 1 . [34] From the values given above it comes that DH 1 (N—O) ¼253.6 6.2 kJ mol 1 for compound 3and DH 1 (N—O) ¼253 9kJmol 1 for compound 4. These results are 10 kJ mol 1 higher than the value of DH 1 (N—O) calculated previously for compound 4with the B3LYP/6-311 þG(2d,2p)// B3LYP/6-31G(d) calculations approach, i.e., DH 1 (N—O) ¼ 242.9 kJ mol 1 . [10] Coincidentally, the experimental results reported above are identical to the calculated DH 1 (N—O) value when the N—O bond cleaved is that closer to the methyl group in 4. [10] The second (N—O) bond dissociation enthalpies, DH 2 (N—O), computed in that work, Equ 4, for compound 4 were 257.9 kJ mol 1 and 267.0 kJ mol 1 when the second N—O bond broken was near the methyl or the ethoxycarboxyl groups, respectively. Experimental results for the DH 2 (N—O) could not be obtained since it was not possible for us to prepare pure samples of compounds 5and 6, hence preventing the determination of their standard molar enthalpies of formation in the gas phase. ð4Þ Table 2. Individual values of the massic energy of combustion, D c u8,atT¼298.15 K Compound 12 25165.40 26399.52 25174.98 26407.31 25199.58 26450.17 25155.43 26378.19 25204.76 25202.32 (<( c u8> a /(J g 1 )) 25183.8 8.7 26409 15 a Mean values and standard deviation of the mean. Table 3. Derived standard (p8¼0.1 MPa) molar values (kJ mol 1 ) in the crystalline phases, at T¼298.15 K a Compound DcUo mðcrÞDcHo mðcrÞDfHo mðcrÞ 15495.4 3.9 5495.4 3.9 262.4 4.2 26133.1 7.1 6134.4 7.1 302.7 7.3 a The uncertainties are twice the overall standard deviation of the mean. Table 4. Calorimetric values for the standard (p8¼0.1 MPa) molar enthalpies of sublimation, at T¼298.15 K a Compound No. of expts T. Dg;T cr;298:15KHo mDT 298:15KHo mðgÞDg crHo mðT¼298:15KÞ KkJmol 1 kJ mol 1 kJ mol 1 17 385 140.4 3.3 22.3 118.1 3.3 26 378 151.4 4.1 22.2 129.2 4.1 a The uncertainties are twice the overall standard deviation of the mean. www.interscience.wiley.com/journal/poc Copyright ß2008 John Wiley & Sons, Ltd. J. Phys. Org. Chem. 2009, 22 17–23 J. R. B. GOMES ET AL. 20 In the present work, we have computed the DH 1 (N—O) and DH 2 (N—O) values for compound 3at the B3LYP/6-311 þ G(2d,2p)//B3LYP/6-31G(d) level of theory. The optimized structures, energies, and thermal corrections for T¼298.15 K are given as Supplementary Material. The calculated results are 242.9 kJ mol 1 and 266.8 kJ mol 1 , respectively, i.e., identical to those calculated for compound 4. Again, the calculated enthalpies for the dissociation of the first N—O bond in compounds 3and 4are in satisfactory agreement with the experimental results. In fact, suggesting that computed values have an uncertainty of 5.0 kJ mol 1 , the experimental and theoretical values do overlap. It is possible to estimate the DfHo mðgÞvalue for compound 5by computing the enthalpy of the working reaction described by Equ. 5: and by using the experimental DfHo mðgÞvalues for toluene, DfHo mðgÞ¼50.0 0.6 kJ mol 1 , [37] for 2,3-dimethylquinoxaline, DfHo mðgÞ¼172.9 3.0 kJ mol 1 , [9] and for methylbenzoate, DfHo mðgÞ¼(287.9 2.4) kJ mol 1 . [37] Combining all these values with the computed enthalpy of the reaction described by Equ. 5, which is 30.8 kJ mol 1 , it comes that the DfHo mðgÞfor compound 5is 134.2 kJ mol 1 . The accuracy of this value can be tested if we use a similar strategy to compute the enthalpy of formation of 3based on the experimental enthalpy of 2,3-dimethylquinoxaline di-N-oxide, DfHo mðgÞ¼149.4  4.5 kJ mol 1 . [10] The computed enthalpy of formation for 3is DfHo mðgÞ¼145.7 kJ mol 1 , in excellent agreement with the experimental value, DfHo mðgÞ¼(148.7 3.2) kJ mol 1 , [9] giving further support to the result estimated for 5. Using the experimental enthalpy of formation for compound 1and the calculated value for compound 5, both obtained in this work, the standard molar N—O bond dissociation enthalpy for compound 3is DH 2 (N—O) ¼259.3 [7.3] kJ mol 1 (Table 5). ð6Þ Using a similar strategy to that employed above, one can also estimate the DfHo mðgÞfor compounds 4and 6. The working reaction used to obtain the enthalpy of formation of compound 6 is that described by Equ. 6 while a similar reaction but with quinoxaline di-N-oxides was used to calculate that for compound 4. The enthalpies of reaction calculated with the B3LYP/6-311 þG(2d,2p)//B3LYP/6-31G(d) approach were 28.5 kJ mol 1 and 17.1 kJ mol 1 for the reaction described by Equ. 5 containing or not quinoxaline di-N-oxide derivatives, respectively. Using the experimental results introduced above for toluene, methylbenzoate, 2,3-dimethylquinoxaline and 2,3-dimethylquinoxaline di-N-oxide and those for ethane, DfHo mðgÞ¼(83.8 0.3) kJ mol 1 , and propane, DfHo mðgÞ¼ (104.7 0.5) kJ mol 1 , also taken from the compilation of thermochemical data due to Pedley, [37] the estimated enthalpies of formation for compounds 4and 6are, respectively, 180.9 kJ mol 1 and 168.8 kJ mol 1 . The result estimated for compound 4is in excellent agreement with the Table 5. Derived standard (p8¼0.1 MPa) molar enthalpies of formation, enthalpies of sublimation and enthalpies of N–O bond dissociation (kJ mol 1 )atT¼298.15 K a Compound DfHo mðcrÞDg cr:Ho mDfHo mðgÞDH 1 DH 2 1262.4 4.2 b 118.1 3.3 b 144.3 5.3 b 259.3 [7.3] f 2302.7 7.3 b 129.2 4.1 b 174 8 b 254 [9] f 3148.7 3.2 c 253.6 6.2 b 259.3 [7.3] f 4178.0 4.3 d 253 9 b 254 [9] f 5134.2 [5.0] e 6168.8 [5.0] e a The uncertainties are twice the overall standard deviation of the mean. b Experimental result from this work. c Experimental result from reference. [9] d Experimental result from reference. [10] e Calculated value from this work. Deviation was estimated. f Suggested value from this work. ð5Þ J. Phys. Org. Chem. 2009, 22 17–23 Copyright ß2008 John Wiley & Sons, Ltd. www.interscience.wiley.com/journal/poc N–O BOND DISSOCIATION ENTHALPIES 21 experimental result published previously, i.e., DfHo mðgÞ¼ (178.0 4.3) kJ mol 1 . [10] That for compound 6differs by 9kJmol 1 from that previously estimated with another working reaction, c.f (159.2 [5.0]) kJ mol 1 . [10] Nevertheless, it is possible to suggest a value for the standard molar enthalpy of dissociation of the second (N—O) in compound 4using the experimental DfHo mðgÞresult for compound 2and the estimated result for compound 6obtained in this work. The suggested value is also reported in Table 5 and it is DH 2 (N—O) ¼254 [9] kJ mol 1 . Importantly, the present work clearly shows that the energy required to cleave the N—O bonds is almost unchanged with R ¼CH 3 or CH 2 CH 3 confirming previous assumptions. [13] Finally, it must be stressed out here that the differences between calculated and experimental DH 1 (N—O) and DH 2 (N—O) are of about 10 kJ mol 1 , almost included in the uncertainties associated with the experimental results (Table 5). Nevertheless, these discrepancies have obvious influences on the experimental or calculated DH 2-1 ¼DH 2 (N—O)DH 1 (N—O) differences for compounds 3and 4. The experimental DH 2-1 values are 5.7  [9.6] kJ mol 1 and 1.0 [12.7] kJ mol 1 for compounds 3and 4, respectively. The calculated DH 2-1 values are much higher, i.e., 23.9 kJ mol 1 and 24.1 kJ mol 1 for 3and 4, respectively. Nevertheless, one is able to retrieve a clear conclusion from the values above, i.e., the DH 2-1 values for compounds 3and 4are similar. CONCLUSIONS Experimental thermochemical work involving static bomb calorimetry and Calvet microcalorimetry has been performed for the first time on quinoxaline derivatives containing a single dative N—O bond. These quinoxaline N-oxide derivatives have a methyl and methoxycarbonyl or ethoxycarbonyl groups attached to positions 2 and 3. The standard molar enthalpies of formation in the gas phase at T¼298.15 K were used to obtain novel standard molar enthalpies of dissociation of the N—O bond in the parent quinoxaline di-N-oxides based solely on experimental results. In fact, previous experimental enthalpies of dissociation for this class of compounds were a mean of the first and second dissociation enthalpic data. The enthalpies of dissociation of the first N—O bond in 3-methoxycarbonyl-2-methyl-quinoxaline di-N-oxide and 3ethoxycarbonyl-2-methyl-quinoxaline di-N-oxide are 253.6 kJ mol 1 and 253 kJ mol 1 , respectively. These identical values confirm previous assumptions stating that different substituents in the carbonyl group attached to the quinoxaline ring had identical effects on the strength of the N—O bonds since the carbonyl group acts like a firewall that blocks the effects of the different substituents on the quinoxaline moiety. Complementary computational study supported all the experimental results obtained in the present work and permitted also to estimate enthalpies of formation for 3-methoxycarbonyl-2-methyl-quinoxaline and 3-ethoxycarbonyl2-methyl-quinoxaline compounds as well as of the enthalpies of dissociation of the second N—O bond in the parent di-N-oxides. SUPPLEMENTARY MATERIAL Spectroscopic (NMR) and analytical (elemental analysis) data of the title compounds, 1and 2. Table SM1 with B3LYP/ 6-311 þG(2d,2p) selected optimized parameters for compounds 1,3, and 5. Table SM2 with the energies and thermal corrections (T¼298.15 K) for all species considered in the present work. This material is available free of charge in Wiley Interscience. Acknowledgements Thanks are due to FCT and to FEDER for financial support to the research project (POCTI/44471/QUI/2002). J.R.B.G. and A.R.M. thank FCT for the award of their research scholarships. 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