Hydrogen Bond Energies of Hydrogen Chloride−Carbonyl Complexes

J. Phys. Chem. 1996, 100, 2083-2088

2083

Hydrogen Bond Energies of Hydrogen Chloride-Carbonyl Complexes Douglas S. Dudis Polymer Branch, Materials Directorate, Wright-Patterson AFB, Ohio 45433-7750

Jennifer B. Everhart Department of Chemistry, UniVersity of Cincinnati, Cincinnati, Ohio 45221

Todd M. Branch Department of Chemistry, UniVersity of Kentucky, Lexington, Kentucky 40506

Sally S. Hunnicutt* Department of Chemistry, UniVersity of Dayton, 300 College Park, Dayton, Ohio 45469-2357 ReceiVed: July 28, 1995; In Final Form: October 30, 1995X

Gas phase hydrogen bonding was studied via infrared spectroscopy in the following systems: methyl acetatehydrogen chloride, methyl formate-hydrogen chloride, 2-butanone-hydrogen chloride, and acetone-hydrogen chloride. The intensity of the H-Cl stretching vibration was monitored as a function of temperature, and the hydrogen bond energies and enthalpies were determined. The hydrogen bond energy for all four complexes was within 2.0 kJ mol-1 of -18.3 kJ mol-1, indicating that the hydrogen bond energy is only minimally affected by nonlocal molecular structure. The acetone-hydrogen chloride complex was modeled by ab initio methods. The energy of formation of the complex, calculated at the MP2/6-311++G** level of theory, is -20.8 kJ mol-1, in good agreement with the experiment.

Introduction Hydrogen bonds1 play important roles in intermolecular interactions of solids, liquids, and gases. The gross structures of numerous hydrogen-bonded complexes have been characterized via infrared spectroscopy of complexes in cryogenic matrices2,3 and high-resolution infrared and microwave spectroscopy of complexes in bulk gases or molecular beams.4,5 Properties of hydrogen-bonded complexes have also been determined (see, for example, Del Bene,6,7 Floria´n and Johnson,8 and Zheng and Merz9) via semiempirical and ab initio calculations. Many hydrogen-bonded complexes have been studied, especially in solution, but the bond energies (i.e, complexation energies) of comparatively fewer gas phase, hydrogen-bonded complexes have been measured.10 High-resolution rotational vibrational spectroscopic results have been used to assess D0 and De for hydrogen-bonded complexes of small molecules.4,5 For example, Ohashi and Pine11 reported a De of 9.5 ( 1 kJ mol-1 for (HCl)2 based on infrared spectroscopy; Howard and Pine12 reported a D0 of 1.372 cm-1 for Ar‚HCl, also based on infrared spectroscopy. Measured De values for HF complexes include 3.59, 12.7, and 20.73 cm-1 for H2‚HF, OC‚HF, and OCO‚HF, respectively,4 and 26.1 kJ mol-1 for HCN‚HF.5 Just two systematic studies of hydrogen bond energies have been reported for complexes of larger molecules. Clague and Bernstein13 investigated the complexation energies of formic, acetic, and propionic acid dimers using infrared spectroscopy. They concluded that there are no significant differences among the complexation energies for these dimers. Thomas14 investigated the hydrogen bond energies of hydrogen fluoride complexes with dimethyl ether, methyl ethyl ether, and diethyl ether using infrared spectroscopy. His results show dramatic variation in the hydrogen bond enthalpies: -43.1 kJ mol-1 for X

Abstract published in AdVance ACS Abstracts, January 15, 1996.

0022-3654/96/20100-2083$12.00/0

HF-dimethyl ether, -37.0 kJ mol-1 for methyl ethyl ether, and -30.0 kJ mol-1 for diethyl ether. A systematic study was begun to evaluate the hydrogen bond energies of complexes of hydrogen chloride with larger molecules. Infrared spectroscopy of the H-Cl stretching vibration in complexed HCl was used as a measure of the concentration of complex in the sample. Analysis of the temperature dependence of this band leads to the hydrogen bond energy. Infrared spectroscopy permits a direct probe of the complex itself, even though its concentration in the sample may be very small. Gaseous mixtures were studied because solvation effects on the bond energy are eliminated and because the gaseous carbonyl, hydrogen chloride, and resulting complex are at chemical equilibrium. Herein is reported the first systematic study on the hydrogenbonding interactions of hydrogen chloride with carbonylcontaining compounds. The gas phase hydrogen bond energy of only one carbonyl-hydrogen chloride complex, acetonehydrogen chloride, has been reported.15 This paper investigates four carbonyl-hydrogen chloride (M‚HCl) complexes, including (CH3)2CO‚HCl, C2H5COCH3‚HCl, CH3COOCH3‚HCl, and HCOOCH3‚HCl. These carbonyl compounds are particularly amenable for study because they have high vapor pressures. The carbonyl-hydrogen chloride complex is easily observed in an infrared spectrum, but the equilibrium constant for formation of the complex is small. High-level ab initio calculations of the energy of hydrogen bond formation have been completed for the acetone-hydrogen chloride complex. The results from those calculations will be summarized here and compared to the experimentally determined energies. Experimental Section Infrared spectra of a gaseous mixture of a carbonyl compound and hydrogen chloride were recorded at multiple temperatures. © 1996 American Chemical Society

2084 J. Phys. Chem., Vol. 100, No. 6, 1996

Figure 1. Comparison of infrared spectra of (a) pure acetone and (b) acetone mixed with hydrogen chloride. The spectrum of the acetonehydrogen chloride mixture has been offset by two absorbance units. At the pressures of acetone used (60 Torr), most of the carbonyl peaks are saturated and do not obey Beer’s law.

The spectra were recorded on a Mattson Cygnus 100 Fourier transform infrared spectrometer at 2.0 cm-1 resolution. Each spectrum was the sum of 32 scans. A 10 cm path length cell of volume 57.3 cm3 was equipped with 37 mm diameter KBr windows and wrapped with Nichrome wire heating tape for temperature studies. The temperature of the body of the cell was monitored with a thermocouple. The cell windows comprise just 8% of the cell’s total surface area; accounting for the slightly lower temperature of the cell windows had no effect on the value of ∆EC. Mixtures were prepared on a standard glass vacuum line and transferred to the cell. Acetone, 2-butanone, methyl acetate, methyl formate, and acetaldehyde were purchased (Aldrich) and then subjected to several freeze-pump-thaw cycles before use. HCl was made by dripping concentrated H2SO4 onto solid NaCl; the HCl was purified by trap-to-trap distillation to remove water. The total gas concentration, 0.015 mol L-1, was the same for each study. For all studies the M:HCl ratio was 1:1, and the total pressure in the cell was 120 Torr. Spectra of the M + HCl mixture were recorded at temperatures between 25 and 50 °C in (approximatey) 3 °C increments. In the time necessary to record one spectrum (less than 1 min), the temperature varied by less than 0.3 °C. Results Figure 1 shows the infrared spectra of pure acetone and of the (CH3)2CO‚HCl mixture at 27 °C. Figure 2 shows the infrared spectra of pure methyl formate and of the HCOOCH3‚HCl mixture at 27 °C. The discrete rovibrational lines of uncomplexed HCl can be observed in the mix spectra; in the methyl formate-hydrogen chloride mix spectrum, only the P-branch appears due to overlap with methyl formate C-H stretching bands. In the spectra of the mixtures, a new, broad, low-intensity band appears in the HCl stretching region between 2500 and 2850 cm-1; the arrows in the figures point to the center of this band. The band is red-shifted relative to ν0(HCl). It is observed in all the mixture spectra, and it is assigned as a combination band of the H-Cl stretch in complexed HCl

Dudis et al.

Figure 2. Comparison of infrared spectra of (a) pure methyl formate and (b) methyl formate mixed with hydrogen chloride. The spectrum of the methyl formate-hydrogen chloride mixture has been offset by two absorbance units. At the pressures of methyl formate used (60 Torr), most of the carbonyl peaks are saturated and do not obey Beer’s law.

Figure 3. Comparison of infrared spectra of the acetone-hydrogen chloride mixture in the H-Cl stretching region as a function of temperature. The top spectrum was recorded at 52.7 °C; the bottom spectrum was recorded at 25.0 °C. The intermediate spectra were recorded at temperatures between 52.7 and 25.0 °C.

(M‚H-Cl stretch) and the M‚HCl bend (hindered HCl rotation). The M‚H-Cl band overlaps a low-intensity vibrational band of the uncomplexed acetone. The M‚H-Cl band intensity decreases with increasing temperature, as is consistent with the exothermic equilibrium

HCl + M a M‚HCl

(1)

This same trendsdecreasing band area with increasing temperaturesis observed for all mixtures studied with the exception of acetaldehyde-hydrogen chloride. Figure 3 illustrates the decrease in band area with temperature for the acetone-hydrogen chloride complex. No complex formation

H Bond Energies of Hydrogen Chloride-Carbonyl Complexes

J. Phys. Chem., Vol. 100, No. 6, 1996 2085

TABLE 1: Summary of the Hydrogen Bond Energies and Enthalpies for Hydrogen Chloride Complexes with Acetone, 2-Butanone, Methyl Acetate, and Methyl Formate; Previously Reported Hydrogen Bond Energies and Enthalpies for the Acetone-Hydrogen Chloride Complex Are Included for Comparison Lewis base

∆EC, kJ mol-1

∆HC, kJ mol-1

methyl acetate methyl formate 2-butanone acetonea acetoneb acetonec

-18.8 ( 0.8 -16.1 ( 0.6 -17.8 ( 1.3 -18.8 ( 0.6 -20.1 ( 2.1

-21.3 ( 3.2 -18.6 ( 0.7 -20.3 ( 1.5 -21.3 ( 0.6 -22.6 ( 2.1 -18.8

d

a This study. b From Mettee et al.15 c From Nowak et al.18 d One standard deviation.

TABLE 2: Summary of Results for Experimental Test Acetone/Hydrogen Chloride Stoichiometry; See Text for Equations PC/P′C mix ratio

n)1

n)2

〈C〉/〈C〉′

4:1/3:1 4:1/2:1 3:1/2:1

0.856 ( 0.033 0.722 ( 0.025 0.843 ( 0.026

0.910 ( 0.036 0.862 ( 0.032 0.947 ( 0.031

0.845 0.728 0.862

between HCl and acetaldehyde was observed. In the gas phase, a significant fraction of the acetaldehyde is trimerized as is apparent from its infrared spectrum. The fraction of monomeric acetaldehyde decreased to near zero after HCl was added and the temperature was increased. Temperature studies of unmixed HCl were carried out at the same concentration as the M‚HCl experiments. The HCl spectrum was identical at each temperature. Thus, no part of the broad band between 2500 and 2850 cm-1 is due to (HCl)2 or any higher order complexes of HCl. This result is in agreement with previous experiments11,16 which determined ∆E for 2HCl / (HCl)2. These experiments show that long path lengths (150 cm to 64 m) and low temperatures (less that 298 K) are needed to observe (HCl)2. The stoichiometry (1 mol HCl to 1 mol (CH3)2CO) given in eq 1 was verified by determining that the stoichiometric coefficient, n, equals 1.0 in following equation

PC (PHCl)nPAC 〈C〉 ) ) n P′ 〈C〉′ (P′HCl) P′AC C

(2)

where acetone is abbreviated AC and where the primed and unprimed variables refer to different HCl/(CH3)2CO mixtures at the same temperature, volume, and total pressure. Equation 2 arises from the simple fact that the equilibrium constant must be the same for two different mixtures at the same temperature. Three different mixtures, 4:1, 3:1, and 2:1 HCl:(CH3)2CO, were prepared, and the area under the complex peak in the infrared spectrum (〈C〉) was determined. Table 2 gives the results in terms of the measured values for 〈C〉/〈C〉′ and the calculated values for the right-hand side of eq 4 with n ) 1 and 2. The M‚HCl band area is proportional to the concentration of the complex via Beer’s law and thus is proportional to the concentration equilibrium constant, Kc. However, it is the partial pressure equilibrium constant Kp, not Kc, which is related to the bond enthalpy via van’t Hoff’s equation. As shown by Mettee,17 van’t Hoff’s equation can be rewritten in terms of Kc and the bond energy (energy of complex formation), not enthalpy, using the definition of enthalpy and Kp + Kc(RT)∆n. The band area of the complex 〈C〉 is related to the hydrogen bond energy EC by

Figure 4. Plots of ln〈C〉 versus T-1 for (a) acetone plus hydrogen chloride (circles), (b) 2-butanone plus hydrogen chloride (squares), (c) methyl acetate plus hydrogen chloride (triangles), and (d) methyl formate plus hydrogen chloride (diamonds). Each graph uses the same x-axis. The lines are calculated from the linear regression best fit to eq 3.

ln

-∆EC 〈C〉 ) +a RT Cl[M][HCl]

(3)

where C is the complex molar absorptivity, l is the cell length, a is an integration constant, and [M] and [HCl] are uncomplexed carbonyl and HCl concentrations, respectively. We assume that C does not vary with temperature and that K , 1, so [M] and [HCl] also remain constant with temperature. The complex band area and the energy of hydrogen bond formation are then related by

ln 〈C〉 )

-∆EC + a′ RT

(4)

where a′ now contains C, [M], and [HCl]. ∆EC was determined from a graph of ln 〈C〉 versus T-1. To measure 〈C〉, each spectrum was shifted up or down so that the base line at 3300 cm-1 is 0.000 absorbance units. After the shift, the variation in the base line at all frequencies was at most 0.003 absorbance units. By comparison, the variation in the absorbances at a single frequency within the M‚H-Cl stretching band was at least 0.040 absorbance units for spectra recorded at temperatures from 25 to 50 °C. Next, 〈C〉 was measured by integrating the total band area excluding the HCl rovibrational peaks. This method resulted in the best precision. 〈C〉 was also ascertained by measuring the peak height at 2700 cm-1 and by integrating the total band area following subtraction of the HCl spectrum. The latter two methods gave the same energies of hydrogen bond formation (within 95% confidence limits) but were less precise. Finally, the temperature dependence of the overlapping carbonyl band was assessed and found to be negligible. The integrated area of the complexed H-Cl stretching band changes by 70% over a temperature study. In contrast, the integrated area of the same region in the spectrum of pure M changes by just 1.6% over a temperature study. The ln 〈C〉 versus T-1 plots are shown in Figure 4 for (CH3)2CO‚HCl, C2H5COCH3‚HCl, CH3COOCH3‚HCl, and HCOOCH3‚HCl. Several temperature studies were carried out for each M‚HCl mixture, and the resultant averaged ∆EC’s and ∆HC’s are given in Table 1. The ∆EC for (CH3)2CO, C2H5COCH3,

2086 J. Phys. Chem., Vol. 100, No. 6, 1996 and CH3COOCH3 is the same within the experimental uncertainty. The ∆EC for HCOOCH3 differs significantly, by more than 1.7 kJ mol-1, from the ∆EC for (CH3)2CO, C2H5COCH3, or CH3COOCH3. Table 1 also includes the ∆EC’s for hydrogen bond formation in (CH3)2CO‚HCl from two previous studies.15,18 Notably, there is good agreement between our ∆EC for (CH3)2CO‚HCl and the two previously reported values for the same complex despite the differences among the studies with respect to physical state (liquid or gas), concentration, temperature range, and infrared band monitored. Nowak et al.18 used infrared spectroscopy to study (CH3)2CO‚HCl in CCl4 solution and in a solid Ar matrix, and they reported a ∆H of -18.8 kJ mol-1 for (CH3)2CO‚HCl in CCl4 solution. No details were given regarding their method of determining ∆H. Mettee et al. also used infrared spectroscopy to study (CH3)2CO‚HCl in the gas phase. They reported a ∆EC of -20.1 ( 2 kJ mol-1. They used the peak height, not peak area, of the complexed and uncomplexed acetone CdO stretching band to determine ∆EC. The CdO stretching band in complexed acetone appears as a shoulder on the uncomplexed acetone CdO stretching band. In order to observe this shoulder band, they used 760 Torr of HCl and 10 Torr of acetone. Calculations Ab initio calculations were performed to obtain that which could not be extracted from the experiment thermodynamic properties of the complex such as ∆S, ∆G, and K. The results of the calculations also help to assess the molecular contributions to the thermodynamic properties. All calculations were carried out using the Gaussian-92 suite of programs.19 The geometries, energies, and harmonic oscillator vibrational frequencies of HCl, acetone, and the hydrogen chloride-acetone complex were optimized using second-order perturbation theory with the triplezeta 6-311++G** basis set incorporating diffuse and polarization functions on all atoms (hereafter MP2/6-311++G**). Structures were fully optimized with no symmetry constraints. The vibrational analysis assured that true minima were found. The final values of ∆G and K were calculated from the experimental, not theoretically computed, values of ∆E and ∆H and the theoretically computed values of ∆S. Previous studies of hydrogen-bonded complexes containing HCl or HF showed that calculations at a similar level of theory (SCF/MP2 using polarization and diffuse functions) yield structures, vibrational frequencies, and energies with small errors. Hydrogen-bonded complexes involving HCl or HF20 such as (HCl)2, (HF)2, HCl-H2O, HF-HCN, HCl-HCN, the HF-chloromethanes,21 and the HCl-chloromethanes22 were studied successfully by such high-level computations. Agreement with experimental data, where available, was good; also, little to no change in results was observed as the basis set size was increased. Recently, Del Bene et al. compared23 ab initio calculations (MP2/6-31+G(d,p)) to density functional theory (based on a hybrid functional) for small hydrogen-bonded complexes of HCl and HF. They found that overall the MP2/ 6-31+(d,p) calculations agreed more closely with experimental results for geometries, binding energies, and vibrational frequencies than did the density functional theory calculations. Computational studies of hydrogen-bonded complexes containing carbonyls have also been reported. Second-order perturbation level computations using electron-correlated polarized double-ζ basis sets were completed24 for N-methylacetamide-water complexes. The calculated and experimental enthalpies of complexation at 298 K agreed within the uncertainty. Floria´n and Johnson8 studied the cyclic formamide dimer using SCF/MP2(6-31G(d,p)) and density functional theory calculations; they found that the two techniques gave similar

Dudis et al.

Figure 5. Structure of the acetone-hydrogen chloride complex as determined at the MP2/6-311++G** level of theory.

TABLE 3: Moments of Inertia (10-47 kg m2) Determined by the 6-311++G** Calculation HCl

(CH3)2CO

(CH3)2CO‚HCl

2.639

83.59 98.23 171.4

92.2 668.2 749.9

interaction enthalpies, geometries, and force constants. Zhang et al.25 reported structures and vibrational frequencies calculated at the 6-31G** level (no electron correlation) for the acetonewater complex. Their results for the vibrational frequencies of the complex, after scaling, agreed favorably with the corresponding experimental values taken from matrix isolation FTIR data. Mettee et al.15 calculated the structure of CH3CHO‚HCl at the 3-21G level and used this geometry to calculate the enthalpy (at 298 K) and vibrational frequencies of the complex at the MP2/6-31G* level. These results are compared to our results below. The computational results for HCl‚(CH3)2CO are discussed in the three following sections: geometry, vibrational frequencies, and thermodynamic properties. Geometry According to the calculations, complexation only slightly perturbs the geometries of acetone and hydrogen chloride. Figure 5 shows the calculated structures of (CH3)2CO‚HCl. Changes in the structure of acetone are largely limited to CdO functional group, which directly interacts with the HCl. The CdO bond length increases from 1.220 to 1.225 Å. The C-C bond located on the same side of the complex as the HCl decreases from 1.516 to 1.511 Å. The H-Cl bond length also increases; it goes from 1.273 to 1.293 Å. The changes in bond length for CdO, C-C, and H-Cl are consistent with the changes in stretching frequencies (discussed below) for these bonds. As seen in Figure 5, the OHCl bond angle is slightly less than 180° with the chlorine atom pointing toward the inplane methyl hydrogen. The COH bond angle is close to 120°, suggesting that the hybridization of O is unchanged by complexation with HCl. Table 3 gives the calculated moments of inertia for (CH3)2CO, HCl, and (CH3)2CO‚HCl. Uncomplexed acetone has a shape which is nearly oblate, but complexed acetone has a shape which is nearly prolate. The dissimilar moments of inertia for uncomplexed and complexed acetone contribute significantly albeit modestly to ∆S of complexation (discussed below). The structure of acetone has been characterized experimentally based on microwave spectroscopy and electron diffraction.26 The experimental results agree quite well with the computed results presented here. Iijima27 determined the zeropoint structure of acetone by reconciling the results of microwave and electron diffraction studies. He gives a CdO bond length of 1.210 Å and a C-C bond length of 1.517 Å, in close agreement with our calculated bond lengths for acetone. The moments of inertia for acetone determined via microwave

H Bond Energies of Hydrogen Chloride-Carbonyl Complexes TABLE 4: Summary of Calculated and Experimental Vibrational Frequencies, in cm-1, for (CH3)2CO, HCl, and (CH3)2CO‚HCl (CH3)2CO‚HCl calc

expta

30.1 65.6 93.0 120.7 132.0 386.2 436.5 481.0 510.5 561.5 559 813.7 896.2 921.6 1095.4 1120.7 1271.3 1232 1402.5 1412.6 1481.6 1482.1 1491.6 1505.4 1756.8 1708 2805.6 2545, 2496, 2393 3075.4 3079.9 3157.5 3162.2 3206.4 3209.1

(CH3)2CO exptb

calc

HCl calc exptc normal modes ClHO bend, νB CH3 rotation

74.9 130.4 379.2 391 480.3 488

1397.4 1405.5 1481.8 1482.7 1490.3 1504.2 1767.2

530 (529)d 787 900 1066 1095 1000 1221 (1223,1216) 1360 1360 1415 1415 1435 1435 1710 (1720)

3072.6 3077.3 3154.4 3159.5 3201.7 3202.7

2922 2922 2965 2965 3005 3005

534.5 805.9 890.5 908.9 1090.6 1117.5 1255.9

OH stretch, ν CCC bend HCl bend, ν+ CCO bend HCl bend, νb CCO bend CC stretch CH3 rock

CC stretch CH3 deformation

CdO stretch 3086 2991 HCl stretch CH stretch

a Observed vibrational frequencies for the acetone-HCl complex in a solid Ar matrix.16 b Observed vibrational frequencies for pure acetone.30 c From Huber and Herzberg.31 d Numbers in parentheses are the observed vibrational frequencies for acetone in a solid Ar matrix.16

spectroscopy28,29 are 8.36 × 10-46, 9.82 × 10-46 and 17.14 × 10-46 kg m2, in good agreement with the results presented here (see Table 3). Vibrational Frequencies The vibrational frequencies of acetone, hydrogen chloride, and the acetone-hydrogen chloride complex were calculated at the MP2/6-311++G** level. The results of the calculations are shown in Table 4 along with the corresponding experimental frequencies30 where available. The computed frequencies have not been scaled. The experimental frequencies for acetoneHCl are taken from the matrix isolation study. The experimental HCl frequency is calculated from the constants in Huber and Herzberg.31 One of the methyl rocking frequencies given for acetone is estimated from heat capacity studies.30 The calculated frequencies are higher than the observed frequencies for the majority of the normal modes. The percent differences between the calculated and observed frequencies range from 0.4% for a methyl rocking mode to 6.5% for the C-H stretching modes. It is well-known32 that this type of calculation overestimates the vibrational frequencies. The frequencies are overestimated because the calculations assume the vibrations are harmonic and because the potential energy surfaces are too rigid at this level of calculation. Six new vibrational modes result when acetone complexes with hydrogen chloride. All six occur at frequencies outside the range of our spectrometer and all fall within the typical range of frequencies.5 The frequencies are all substantially (25-50%) lower than the analogous modes for the acetaldehyde-hydrogen

J. Phys. Chem., Vol. 100, No. 6, 1996 2087 chloride complex reported by Mettee et al. This difference is most likely due to the larger basis set (6-311++G** versus 3-21G for Mettee et al.) and the inclusion of electron correlation in our calculation. The calculations show that the acetone CdO, the C-C, and the H-Cl stretching modes are altered significantly upon complexation. The CdO stretch shifts 11 cm-1 to the red and the H-Cl stretch 280 cm-1 to the red, consistent with the lengthening and weakening of this bond in the complex. The C-C stretching vibration shifts to higher energy by 7.8 cm-1, consistent with the slight shortening and strengthening of this bond in the complex. The CCO in-plane bending mode of acetone is also shifted upon complexation. This mode shifts to higher energy because attached HCl obstructs this bending motion. Mettee et al. calculated much larger peak shifts for the acetaldehyde-hydrogen chloride complex. Again, the differences between their results and ours are likely due to the differences noted above between the two calculations. In the experimental infrared spectra recorded in this study, the center of the complexed-H-Cl stretching band appears at roughly 2690 cm-1. Using this point as ν0 gives an experimental shift of approximately 300 cm-1, in reasonable agreement with the calculated shifts. No bands corresponding to complexed acetone were observed in the spectra of the (CH3)2CO‚HCl mixtures; at the acetone concentrations used, the acetone bands are fully absorbing and would mask any small bands due to the complex. However, sharp, nonoverlapped vibrational bands are observed for acetone-hydrogen chloride complexes in the lowtemperature Ar matrix; the frequencies of these bands are consistent with the calculated frequencies reported here (see Table 4). Thermodynamic Properties The ab initio calculations of the minimum energies of (CH3)2CO, HCl, and (CH3)2CO‚HCl are used with the calculated vibrational frequencies to determine the energy of complex formation, ∆EC, at 298 K. The calculated value, -20.8 kJ mol-1, is in good agreement (within 2.0 kJ mol-1 or 0.49 kcal mol-1) with the experimentally determined value, -18.8 ( 0.6 kJ mol-1. The electronic energy of complexation has been corrected for basis set superposition error (BSSE) using the counterpoise technique.33 The thermodynamic properties associated with the complexation of acetone and hydrogen chloride are ∆H ) -21.3 kJ mol-1 (from the experiment), ∆S ) -106.9 J K-1 mol-1 (from the ab initio calculation), ∆G ) 10.6 kJ mol-1 (from ∆H and ∆S), and equilibrium constant K ) 0.0139 (from ∆G). Complexation is exothermic, and the negative bond energy is the driving force behind the complexation reaction. This driving force is offset by the net loss in entropy upon complexation. The ∆S of complexation is a sum of the negative entropy changes due to translation and rotation and the positive entropy changes due to vibration. These present results are compared to those reported by Mettee et al. for the complexation of acetone and hydrogen chloride. First, the present results for ∆H and ∆E agree with those of Mettee et al. within experimental uncertainty. Second, value reported here for K (from experiment and theory) is an order of magnitude smaller than that of Mettee et al., who determined K from the ratio of complexed to uncomplexed acetone CdO band peak heights. To determine K, Mettee et al. estimated the ratio of complexed to uncomplexed acetone CdO band molar absorptivities using the square of the ratios of the difference in calculated Mulliken charges on C and O in complexed and uncomplexed acetaldehyde. (For computational reasons, Mettee et al. used acetaldehyde as a reasonable model

2088 J. Phys. Chem., Vol. 100, No. 6, 1996

Dudis et al.

for acetone.) If we substitute our value of this molar absorptivity ratio for their value, their value for K falls to 0.075. This value is closer to, but still significantly different from, our assessment of K. Third, the theoretical values for ∆S calculated in this report are compared to those of Mettee et al., which were actually calculated for acetaldehyde-hydrogen chloride. Their calculation gave a complexation entropy of -112.51 J K-1 mol-1. As noted by Mettee et al., the smaller basis set used in their calculations overestimates the vibrational frequencies. This is especially true for the frequencies associated with complex formation which, in turn, contribute most to the entropy of complexation. Thus, Mettee et al. conclude that their calculation does not accurately reflect the higher vibrational entropy of the complex relative to HCl and (CH3)2CO. They report a complexation entropy of -92.09 J K-1 mol-1 using literature values for the low-frequency modes of the complex. Last, the loss in rotational entropy from the present calculation, ∆Srot ) -18.5 J K-1 mol-1, is nearly twice that given by Mettee et al. (∆Srot ) -10.5 J K-1 mol-1).

to the CdO group. Its complex with HCl could not be stabilized by the methyl hydrogen-chlorine interaction, and thus the HCOOCH3‚HCl complex would have a slightly lower ∆EC. In lieu of experimental results for the gross structures of the carbonyl-hydrogen chloride complexes, we will investigate this hypothesis with further calculations on the C2H5COCH3‚HCl, HCOOCH3‚HCl, and CH3COOCH3‚HCl complexes.

Contributions to ∆EC

References and Notes

The experimental and computational results show that the hydrogen bond formed between HCl and the small carbonyls is weak. The structure of the complexed carbonyl compound is nearly unchanged compared to the uncomplexed carbonyl. The intermolecular interaction appears to be localized largely to the lone pair electrons of the carbonyl oxygen and the hydrogen atom in HCl. This study suggests that CH3, CH2CH3, and OCH3 groups stabilize the hydrogen-bonded complex equally and that a H group stabilizes the complex less. Structural arguments are used to interpret this result. Steric interference from the bulkier ethyl side group in methyl ethyl ketone might be expected to give a lower ∆EC; the lack of this steric interference from the adjacent hydrogen in methyl formate might lead to a higher ∆EC. Our results do not support this interpretation. The steric interactions instead most likely lower the entropies (not the energy) of the complexes, thereby reducing the free energy of complexation and the equilibrium constant. Mesomeric and inductive effects are more important than steric hindrance. The ether oxygens of the two esters can mesomerically stabilize the complex due to oxygen’s lone pair electrons. However, the high electronegativity can inductively destabilize the complex. For the esters, the stabilization and destabilization mechanisms offset each other. The ketones, by comparison, have alkyl groups adjacent to CdO which can neither stabilize nor destabilize the complex mesomerically or inductively. Thus, the ketones and esters would be predicted to form complexes with HCl of similar stability. This model is consistent with the experimental and computational results, at least for acetone, methyl ethyl ketone, and methyl acetate. Alternatively, the hydrogen bonds of the (CH3)2CO, C2H5COCH3, and CH3COOCH3-HCl complexes may be stabilized by a weak interaction of the in-plane methyl hydrogen with the chlorine. In (CH3)2CO and (CH3)2CO‚HCl, the calculations place this hydrogen cis to the CdO. Rotational spectroscopy measurements34 are consistent with the calculation. Certainly, the methyl hydrogens and the hydrogen chloride will rotate within the complex, but the long-range hydrogen interaction may favor the structure drawn in Figure 5. The Mulliken atomic charge for this hydrogen (0.1145) is lower than the two out-of-plane hydrogens (0.1711 and 0.1542) in the complexed acetone. For uncomplexed acetone, the Mulliken atomic charges for the in-plane and two out-of-plane hydrogens are 0.1789, 0.1696, and 0.1592, respectively. Methyl formate, of course, has no in-plane methyl hydrogen adjacent

Conclusion Experimental results have been combined with ab initio calculations, yielding a thermochemical picture of the hydrogen chloride-carbonyl complexes with some molecular detail. This type of study is complementary to high-resolution vibration and rotation spectroscopy studies. It is well-suited to complexes of larger molecules, and it emphasizes macroscopic thermodynamics instead of intermolecular potentials. Acknowledgment. The authors thank Dr. H. C. Knachel for fruitful discussions and for the use a vacuum line. We also thank D. L. Cecotti for her work in the lab.

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