2340
Anderson, Rytting, Lindenbaum, and Higuchi
or oxygen. Especially we would like to point out that even bivalent sulfur atoms like those in the present study have very small ionization potentials for the lone pair and u electrons, and lower levels of virtual orbitals, e.g., the u*c-s bond. Acknowledgment. We wish to thank Professor Hiroshi Kat0 for his helpful discussion and also Dr. T. Kobayashi and Mr. T. Utsunomiya for the measurement of the photoelectron spectra. We are also indebted to Mr. K. Akagi for the synthesis of ETA. References and Notes (1) P. H. Laur, “Sulfur in Organic and Inorganic Chemistry”, Vol. 3, A. Senning, Ed., Marcel Dekker, New York, N.Y. 1972, Chapter 2 4 J. I. Musher, Angew. Chem., 81, 68 (1969). (2) J. R. Sabin, J. Am. Chem. Sac., 93, 3613 (1971); J. S. Rosenfield and A. Moscowitz, ibid., 94, 4797 (1972): A. L. Companion, Theor. Chim. Acta (Berlin), 25, 268 (1972); V. B. Koutecky and J. I. Musher, /bid., 33, 227 (1974). (3) P. W. Sadler. Chem. Rev., 60, 575 (1960); J. D. Watson, “Molecular Biology of the Gene”, W. A. Benjamin, New York, N.Y., 1965. (4) T. C. Bruice, “Organlc Sulfur Compounds”, Vol. 1, N. Kharasch, Ed., Pergamon Press, New York, N.Y., 1961, Chapter 35. (5) M. W. Cronyn, M. P. Chang, and R. A. Wall, J. Am. Chem. SOC.,77, 3031 (1955). (6) H. Yamabe, H. Kato, and T. Yonezawa, Bull. Chem. SOC.Jpn., 44, 604 (1971). (7) J. E. Gano and H. G. Corkins, Chem. Commun., 294 (1973). (8) J. R. Grunweli and H. S. Baker, J. Chem. SOC.,ferkin Trans. 2, 1542 (1973). (9) J. A. Pople and G. A. Segal, J. Chem. fhys., 44, 3289 (1966); D. P. Santry and G. A. Segal, ibid., 47, 158 (1967). (IO) T. Yonezawa, H. Konishi, and H. Kato, Bull. Chem. SOC.Jpn., 42, 933 (1969). (11) D.W. Turner, C. Baker, A. D. Baker, and C. R. Brundle, “Molecular Photoelectron Spectroscopy”, Wiley, London, 1970. (12) (a) T. Wieland and W. Bartmann, Chem. Ber., 89, 946 (1956): (b) Y. Hirabayashi and T. Mazume. Bull. Chem. SOC.Jpn., 38, 171 (1965). (13) In the case of organic compounds with a bivalent sulfur atom, the contribution of the sulfur 3d orbitals to the absorption bands cannot be nec-
(14) (15) (16) (17) (18) (19) (20) (21) (22) (23) (24) (25) (26) (27)
(28) (29)
essarily neglected in the near-ultraviolet region. The present case, however, deals mainly with the absorption bands of comparatively small excitation energy in the ultraviolet region, so that we may neglect the contribution of the sulfur 3d orbitals. For example, see ref 2. W. Gordy, J. Chem. fhys., 14, 560(1946). M. Hayashi and L. Pierce, Spectrochim. Acta, 16, 1272 (1960); L. Pierce and M. Hayashi, J. Chem. fhys., 35,479 (1961). W. D. Closson and P. Haug, J. Am. Chem. SOC.,66, 2384 (1964). H. H. Jaffe and M. Orchin, ”Theory and Applications of Ultraviolet Spectroscopy”, Wiley, New York, N.Y., 1966. a a* transition energy of CH~CONHZ is substituted by that of HCONHz. J. Hinze, M. A. Whitehead, and H. H. Jaffe, J. Am. Chem. SOC.,85, 148 (1963). D. A. Steigart and D. W. Turner, J. Am. Chem. SOC.,94, 5592 (1972). J. A. Pople, Roc. fhys. SOC.,A68, 81 (1955). L. C. Snyder and H. Basch, “Molecular Wavefunctions and Properties”, Wiley, New York, N.Y., 1973. J. N. Murrell, “The Theory of the Electronic Spectra of Organic Molecules”, Wiley, New York, N.Y., 1963. H. Suzuki, “Electronic Absorption Spectra and Geometry of Organic Molecules”, Academic Press, New York, N.Y., 1967. E. E. Barnes and W. T. Simpson, J. Chem. fhys., 39, 670 (1963). C. R. Brundle, D. W. Turner, M. B. Robin, and H. Basch, Chem. fhys. Lett., 3, 292 (1969). T. Kobayashi, K. Yokota, and S. Nagakura. J. Electron Specfrosc., 3, 449 (1973). in previous studles, CNDO/2 has been shown to be very successfui, and superior to the ASMO-SCF method, for the evaluation of dipole moments, rotational barriers, charge densities, etc. Hence we also used this method for TAA and ETA. The contribution of the sulfur 3d orbitals Is neglected for the reasons mentioned above. We used electron diffraction datai4 for the geometry of TAA, and for ETA we used the acyl part of TAA and the S-CY fragment of CH3SCH3.i5 For the calculations of electronic transition energies, CNDO/2 does not seem appropriate: for example, the HOMO of ethylene is not predicted to be a a orbital, but a u orbital of the C-H bond, lnconslstent wlth exa* transition energy obtained is too large. perimental results: the ?r Hence, for the present calculations of transition energies, we used an ASMO-SCF method, which not only includes one-center exchange integrals, but also estimates effectively core integrals, and which then may be appropriate to explain the transition energies of some conjugated molecules with lone pair electrons. 4 ns’ indicates the orbltal of the lone pair electrons of the sulfur atom, but it is different from as in that nst does not partlcipates in a conjugation with the carbonyl group.
-
-
(30)
A Calorimetric Study of the Self-Association of Primary Alcohols in Isooctane Bradley D. Anderson, J. Howard Rytting,* Siegfried Llndenbaum, and Takeru Higuchl Department of Pharmaceutical Chemistry, University of Kansas, Lawrence, Kansas 66045 (Received June IO, 1975)
The heats of dilution of ethanol, propanol, butanol, pentanol, hexanol, heptanol, and octanol have been measured calorimetrically in isooctane as a function of concentration a t 25OC. The resulting AH values are interpreted in terms of the degree of association and nature of the possible associated species assuming that the heat of dilution is essentially due to the breakdown of the various associated species present. The results indicate that dimer formation is not important except in dilute solutions. In solutions having concentrations greater than 0.1 M , tetramers dominate. Equilibrium constants and standard enthalpies are calculated for the major association reactions observed. Ethanol has a greater enthalpy of dilution than the other alcohols studied.
Although the tendency for molecules capable of hydrogen bonding to self-associate even at high dilution in “inert” solvents has long been recognized, an understanding of the nature of self-association interactions has been elusive. Due to the transient character of hydrogenbonded multimers, no direct methods are available for obThe Journal of Physical Chemistry, Vol. 79, No. 22, 1975
taining information. Typically, the change in a solution property with solute concentration due to the self-association of the solute or a change in the spectroscopic behavior of the solute when it is involved in hydrogen bonding have been employed.’ The self-association of primary alcohols has been exam-
2341
Self-Association of Primary Alcohols in Isooctane ined by several investigators using a variety of techniques but the conclusions are often in disagreement. Van Ness et aL2 compared infrared data with heat of mixing data for ethanol-heptane and ethanol-toluene mixtures, concluding that a model consisting of monomers, cyclic dimers, and linear polymer chains having lengths of 20 or more monomeric units best fit the data. Determinations based on vapor pressure measurements for methanol in n-hexadecane and on PVT measurements for methanol vapor have been explained with a monomer-trimer-octamer model.3 Fletcher and Heller4 have shown that their infrared data for 1-octanol and 1-butanol in n-decane are consistent with monomers and tetramers with no other species existing in amounts sufficient to affect the material balance equations. Fletcher5 later concluded that the extent of self-association of primary alcohols in saturated hydrocarbons was in the order 1-octanol > 1-butanol > ethanol-dl > methanol. Using NMR, Dixon6 also found that the monomer-tetramer model gave the best fit when applied to the hydroxyl proton shift data of methanol in cyclohexane. More recent studies have suggested that at very low concentrations more than one associated species must be considered. Aveyard et aL7 concluded that both dimer and tetramer must be considered up to 0.13 M to account for the ir and vapor pressure osmometry data for 1-octanol and dodecanol in octane. Also, a recent dielectric study of l-octano1 solutions8 has indicated that a small, high dipole moment polymer is important at very low concentrations followed by a low dipole moment cyclic polymer and then by larger, high dipole moment polymers at higher concentrations. Although equilibrium constants can generally be directly determined from data obtained using the techniques mentioned, bond enthalpies are usually found indirectly and are therefore regarded as approximate values. Normally the temperature dependence of the equilibrium constant, K , is determined and AHo values are estimated from a van't Hoff plot. Calorimetric studies of association, however, measure heat directly and provide more accurate enthalpies. Calorimetry has been applied to the study of phenol association by Woolley and Heplerg and Woolley et a1.lO In the present study, heats of dilution of primary alcohols in isooctane have been measured to determine the effect of chain length on the bond strength and extent of selfassociation. Data were obtained for solutions of aliphatic alcohols from ethanol through 1-octanol in isooctane in the concentration range of approximately 0.003-1.0 M with particular emphasis placed on concentrations below 0.1 M. Methanol is partially immiscible in isooctane and could not be studied over the same concentration range using the same techniques. The results were interpreted in terms of several models of association. Experimental Section Materials. The chemicals used were of the highest purity available from the supplier and were used without further purification. The ethanol was reagent quality, 200 proof, from U.S. Industrial Chemicals; the 1-propanol was analytical reagent from Mallinckrodt; 1-butanol, 1-pentanol, and 1-heptanol were 99+ mol 96 obtained from Matheson Coleman and Bell; 1-hexanol and 1-octanol were from Fluka AG with purities better than 99%; and the isooctane used was Fisher Scientific Co. certified reagent. The alcohols and isooctane were stored over molecular sieves to remove trace amounts of water.
Procedure. Calorimetric measurements were made using the Tronac Model 550 titration calorimeter in the adiabatic mode. The bath temperature and starting temperature of all runs was 25.00 f O.0loC. The system was tested by measuring the heat of protonation of aqueous Tham, tris(hydroxymethyl)aminomethane, with HC1. The value of AH found for this reaction at 25OC was -11.33 f 0.04 kcal/mol which agrees well with the value of Hill, Ojelund, and Wadsoll of -11.35 kcal/mol. The alcoholic solutions used in the calorimetric titrations were prepared by weight. Densities of these solutions in isooctane were measured a t several concentrations and were employed in converting solution concentrations to molarities. Densities were measured a t 25OC using either a 25-ml pycnometer or the Sodev Inc. densimeter Model 01D based on the oscillating tube principle.12 The titrations were carried out by diluting solutions of ca. 0.6 and 6-10 M (pure alcohol or a concentrated solution) into 40 ml of anhydrous isooctane. At least three runs were made for each solution. Using the continuous titration calorimetric method which is described fully elsewhere,13 several values could be obtained from each run. Results of the measurements were calculated as the heat in calories (1 cal = 4.1840 J) consumed in diluting 1 mol of solution of the initial molarity to the molarity in the reaction vessel a t a specified time. These values are reported in Table I.14 Results The density data for solutions of alcohols in isooctane which were used to convert alcohol concentrations to molarities were fitted by a least-squares method to a quadratic equation p =a
+ bM' + cMr2
(1)
where M' is in units of moles/kilogram of solution. The coefficients for eq 1are reported in Table I1 for solutions of ethanol through 1-octanol in isooctane. The change in enthalpy when an alcoholic solution of specified molarity is diluted to a lower molarity is interpreted as being due to a breaking of bonds as an associated n-mer dissociates to monomer. Considering the case of a single equilibrium, association of an alcohol can be represented by nA1* A,
(2)
with an equilibrium constant which can be expressed as K1,n = (An)/(Al)"
(3)
where (AI) is the molar concentration of the alcohol monomer and (A,) is the molar concentration of associated n mer. The enthalpy, AH, for dissociation from an initial total molarity, Mi, to a final molarity, Mf, is equal to the negative of the standard molar enthalpy for association, AHa I,,, multiplied by the fraction of n-mer dissociating (4) or substituting from eq 3
Using the material mass balance equation and substituting for A, found from eq 3 enables one to calculate the value of A1 a t any concentration for a given Kl,, as follows
The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
Anderson, Rytting, Lindenbaum, and Higuchi
2342
TABLE 11: Density D a t a for Solutions of Alcohols in Isooctanea a
Ethanolb 1-Propanolb 1-BUtanol 1-Pentanol 1- Hex anol 1-Heptanol 1-Oct anol 0
p
=a
+ bM' +
b
C
0.6877 0.002366 0.000502 0.6877 0.00437 0.0006936 0.6872 0.007191 0.000118 0.6870 0.009576 0.000127 0.6870 0.011396 0.000174 0.6873 0.013246 0.000242 0.6872 0.015357 0.000274 CM'2. * Measured with the Sodev densimeter.
Using the above relations one can optimize the variables K and AHo when several sets of AH, M,, and Mf are known. We have used a computer program employing the simplex method of least-squares described by Deming and Morgan15 to fit the data, The values of AH entered were in most cases obtained by determining the difference in two measured enthalpy changes in diluting from a single high concentration to two concentrations below 1M. In this way only the concentration region from 0 to 1.0 M was fitted even though the experimental data were often based on a high initial molarity. By entering a large number of data of low concentration (below 0.1 M ) and fewer values at higher concentrations the computer was forced to fit the regions of low molarity most accurately. The computer program was constructed so that a single equilibrium involving monomer with either dimer, trimer, tetramer, or pentamer could be used to fit the data. Additionally, it was hoped that more information could be gained by looking at various competitive equilibria operating simultaneously. Thus, several combinations of the above equilibria also were considered with as many as eight parameters being optimized. The results of fitting a number of models to the data for 1-butanol are reported in Table 111. The model with the lowest relative standard deviation for all alcohols considered was a monomer-dimertetramer scheme. It was generally found that fits using more than four parameters (more than two competing equilibria) also converged to solutions in which the relative standard deviation was small, but in every case the monomer-tetramer equilibrium was dominant, as is evident from the last two examples in Table 111. The monomer-tetramer model gave good correlation for all concentrations except in very dilute regions. The table shows that the monomer-tetramer model results in a much lower standard deviation and, hence, a much better fit than the monomertrimer or monomer-pentamer models. It has been observed in using the simplex method that the initial guesses for each of the parameters sometimes can affect the final result. In this study, we found that for a single equilibrium the solution was unique, independent of the initial guesses. This was not the case when two or more equilibria were treated simultaneously, but the standard deviation was always significantly higher when the computer converged to solutions other than those reported in Table 111. One of the difficulties in most of the previous association studies and in the present one is the failure to maintain constant temperature. NMR techniques6 have been utilized with a temperature variation of f0.5OC while ir data2Js7 have been obtained with temperature controlled between ranges of &O.l°C at best to floc.Although some The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
calorimetric studied6 of alcohol-alkane systems have been done a t constant temperature the data a t low concentrations are generally sparse. A calorimetric study similar to the present work of 1-octanol in n-decane17 was carried out a t an initial temperature controlled to within f0.2OC and with a net exothermic heat change occurring during the dilution. Although the endothermic heat changes occurring during our measurements were small with greater than half of the data being taken a t temperatures between 24.7 and 25.0°C, the temperature dependence of the equilibrium constant is very large for the tetramer formation. We therefore felt it necessary to apply a first-order correction to the data before attaching any significance to observed similarities or differences. Such a correction was possible because the temperature at each point was known. Since the tetramer formation is the dominant equilibrium (except perhaps at very high dilution in which case the temperature of the determination was also very close to 25.OoC) eq 5 can be used to solve for a K1,4 for each value of AH by entering the measured AH, Mi, Mf, and the A H o 1 , 4 which was previously optimized by the simplex program. Each K1,4 can be corrected individually to 25.0°C by first converting from an equilibrium constant based on the temperature dependent units of molarity, K M ,to an equilibrium constant based on temperature independent units, Km, by the relation Km = KM[P- (O.OO~)(MB)(WB)]~
(7)
where p is the solution density a t a given temperature, and MB and WB are the molarity and molecular weight of solute, respectively. An approximate form of eq 7, K, = K~p~(SOlVent), was sufficiently accurate in most cases and was therefore used in calculating the temperature correction. Applying the van't Hoff equation to K, d In K,IdT = AHo1,4/RT2
(8)
converts Km at any temperature to K, a t 25.0°C. Converting back to a new K1,4 based on molarities and substituting it back into eq 5 allows a AH corrected to 25.0°C to be calculated for each data point. A simplex regression can now be applied to the temperature corrected data set to obtain equilibrium constants and standard enthalpies applicable to 25.0°C. Table IV compares the magnitudes of parameters found after temperature correction to those based on the raw data and demonstrates that the overall results are not altered greatly. The value for AHo1,4 was assumed to be constant with small temperature changes and does not appear to change when the new values for AH are fit to the monomer-tetramer model. The sensitivity of the equilibrium constant to temperature, however, is clearly evident from the table. A useful technique in evaluating the closeness of fit for a particular model is a comparison of computer generated +L values using the best fitting model vs. the experimental data extrapolated to infinite dilution. The relative apparent molar enthalpy represented by t $ ~is equal to the negative enthalpy of dilution of 1mol of alcohol from a solution of given molarity to infinite dilution and can be expressed mathematically as $L
2 AHo 1,,Xl,j((Al)i'/Mi)
j=2
(9)
A comparison of calculated vs. experimental results for 1octanol in Figure 1 shows the excellent agreement obtained with the monomer-dimer-tetramer model. The absolute
2343
Self-Association of Primary Alcohols in Isooctane TABLE 111: Optimization of Parameters Using Various Models to Fit 1-Butanol Enthalpy Dataa
-15.2
3.98 -17.5
-425 -123 -5.3c -100 -120 -15.8 -67.3
0.017 0.059 0.38 0.0001 0.102 0.337 0.107
48.2
-17.5
48.2
-38.8 -19.9
0.003 0.093
-21.4
587
-21.3 -21.3 -21.6
395 411 63 2
-21.6 -21.4
530 422
-25.5
7442
-25.8
3251
-192
0.071
0.20 0.076 0.030 0.061 0.020 0.020 0.027 0.076 0.033 0.024 0.021
AH(kca1) = 2.p - Awl, (kcal/rn~l)K~,~I[(A~),n/M,] - [(Al)p/Mf]l. * u = relative standard deviation between experimental and calculated AH when the given parameters are used in the regression equation. A constant value of -5.3 kcal/mol was entered into the equation, to determine a K for dimer formation based on a reasonable AH" value. As is shown, K1.z remains small in this treatment. a
TABLE IV: Comparison of Temperature Corrected vs. Uncorrected Values for K1,4 a n d AH01,4 (kcal/mol) AH"1,4
Ethanol 1-Propanol 1-ButanOl
1-Pentanol 1-Hexanol 1-Heptanol 1-Octanol
K1,d ( T 5 25°C)
(T 5 25°C)
(25°C)
(25°C)
674 555 587 657 673 633 69 1
-22.6 -21.4 -21.4 -21.1 -21.1 -21.2 -21.2
651 533 555 623 651 599 662
-22.6 -21.4 -21.4 -21.2 -21.2 -21.2 -21.2
4000
-
3000
-
;2 0 0 0
-
K1,4
- 4 ~values reported are based on an extrapolation of the best model to infinite dilution and may be subject to adjustment for a different extrapolation to infinite dilution. Thus, the two theoretical curves shown in Figure 1 differ by approximately 90 cal a t infinite dilution even though both models fit the data well a t higher concentrations. Differences in -& a t a concentration of 1M reflect differences in bonding enthalpies since the alcohol is largely associated at this concentration. The following values of - 4 ~(kcal/mol) a t 1 M concentration were calculated using the monomer-dimer-tetramer model: ethanol = 4.95, propanol = 4.65, butanol = 4.68, pentanol = 4.67, hexanol = 4.65, heptanol = 4.66, and octanol = 4.68. The treatment of the calorimetric data assumes that (i) the heat of dilution is due only to the dissociation of polymers and (ii) activity coefficients of all species are equal to one. Although neither of these assumptions is absolutely true, narrowing the concentration range of interest to 1 M and below minimizes the deviation from these assumptions. Other investigators have recognized that enthalpies of mixing of alcohols and alkanes are due not only to hydrogen bonding but to a dipolar enthalpy contribution as well. Smith and Brownls have estimated that the dipolar energy is only about 5% of that due to hydrogen bonding for butanol in n-decane. Furthermore, the dipolar enthalpy changes more rapidly than the hydrogen bonding enthalpy a t high concentrations so a t a concentration of 1 M roughly 10% of the total apparent dipolar enthalpy of alcohol remains while the extent of hydrogen bond breaking is less
r" (Y
0
V 0
'L
Figure 1. A comparison of the experimental vs. theoretical values of -& for I-octanol: (-) theoretical curve using the monomer-
dimer-tetramer model; (- - - -) theoretical curve using the monomer tetramer model; (I) range of experimental data.
than 20% complete. We assume, then, that the nonhydrogen bonding component to heats of dilution of alcohols is less than 1%in the molarity region considered. The assumption that activity coefficients are unity is also a much better assumption when the dilute region is considered since we are using the infinitely dilute standard state. This assumption allows us to treat activity as concentration in the equilibria equations. Discussion The calorimetric technique alone cannot provide enough The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
2344
information to completely identify the number, size, and configuration of associated species in alcohol-alkane solutions but when the data presented here are combined with the results of other methods, a strong case can be made for the dominance of tetramer. Fletcher and Heller4 proposed that the tetramer predominance is consistent with ir data and reported an equilibrium constant of 774 M-3 for tetramer formation of 1-octanol in n-decane which compares favorably with our calculated value of 660 M-3 for 1-octanol in isooctane. They further interpreted two overlapping peaks as representing a cyclic and a linear structured tetramer and calculated the bond enthalpy of a single hydrogen bond to be between -5.1 and -5.5 kcal/mol depending on whether the linear or cyclic structure was involved. Van Ness et a1.2 determined the hydrogen bond enthalpy to be -5.2 kcal/mol in a study of ethanol-n-heptane solutions. Since the AH value for tetramer formation was approximately -21.2 to -21.4 kcal/mol for all alcohols observed (except ethanol) we can divide this by three for a linear or four for a cyclic structure to determine the enthalpy per bond. Doing this results in a hydrogen bond enthalpy of -7.1 kcal/mol for the linear configuration and -5.3 kcal/ mol for a cyclic structure. Obviously the cyclic bond enthalpy compares much more favorably with the literature results. A cyclic tetramer as the dominant species in dilute alcohol-alkane solutions would not be anticipated from a priori consideration. However, the conclusion that a cyclic tetramer was the more stable polymer was rationalized by Fletcher and Heller4 on the basis of evidence submitted by Bellamy and Pacelg that the free OH groups of linear polymers have a greater tendency to hydrogen bond than the original monomers. Thus a linear tetramer would tend to form the more stable cyclic tetramer or add monomer to form a linear pentamer. Further evidence for a large cyclic polymer has recently been published by Campbell et aLS in which the apparent dipole moment vs. concentration curves for 1-octanol in cyclohexane, carbon tetrachloride, and benzene were determined. The results were consistent with the formation of a small amount of open dimer which causes a slight increase in the apparent dipole moment at very low concentrations followed a t higher concentrations by a cyclic species as reflected by a decrease in the dipole moment with concentration. The low dipole moment cyclic polymer is the primary species until a concentration of about 0.5-1.0 M after which the dipole moment increases rapidly due to the emergence of a linear polymer with a high dipole moment. Although our results are consistent with cyclic tetramer, the formation of linear tetramer at higher concentrations would not be easily detected in a calorimetric technique. The existence of a small amount of dimer has been stressed by several investigators including Aveyard et al.7 who suggested that dimer may reach a maximum percentage of 5% a t approximately 0.1 M . Although allowing dimer to exist in our model calculations does improve the fit considerably, the value of K1,z is always very small relative to K1,4. Additionally, the AH1,z is often unreasonable, being as much as two orders of magnitude higher than expected for the enthalpy of a single hydrogen bond. Since dimers are predicted only a t very low concentrations, where the error in calorimetric measurements is highest, a calorimetric study of the alcohol dimer is also limited. Discussions of monomer-tetramer vs. monomer-dimertetramer in the literature apparently concern a species comprising less than 5% of the formal alcohol concentraThe Journal of Physical Chemistry, Vol. 79, No. 22, 1975
Anderson, Rytting, Llndenbaum, and Higuchi tion. With this in mind it seems reasonable to propose that the monomer-tetramer equilibrium is adequate in describing the behavior of alcohols in alkanes in concentrations less than 1 M and possibly up to neat alcohol as Fletcher5 has suggested. For this equilibrium our data show that the difference between the primary alcohols in their tendency to associate is not large and do not necessarily support the conclusion by Fletcher that the order of association is 1octanol > 1-butanol > ethanol-d1 > methanol. Our values for K1,4 vary by f 1 0 % but there is no general trend. The enthalpies of tetramer formation were nearly identical (fl%)for all the alcohols studies except ethanol. The enthalpy of the ethanol hydrogen bond is greater than the hydrogen bond enthalpies of the other alcohols being about -5.65 kcal/mol if a cyclic tetramer is assumed, as compared to -5.3 kcal/mol for the cyclic tetramer of other alcohols. These values, which are based on the monomer-cyclic tetramer model are in remarkably close agreement with the values determined by Savini et a1.16 of -5.6 kcal/mol for the ethanol hydrogen bond and -5.2 kcal/mol for the octano1 hydrogen bond a t 3OoC in hydrocarbon solvents. The similarity of bond enthalpies for alcohols larger than ethanol has also been observed by Smith and Brown.18 Acknowledgments. This work was supported in part by grants from the University of Kansas General Research Fund and by a Health Science Advancement Award (FR06147) granted to the University of Kansas by the National Institutes of Health. Supplementary Material Available. The calorimetric data contained in Table I will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of this material from this paper only, or microfiche (105 X 148 mm, 24X reduction, negatives) containing all material for the papers in this issue, may be obtained from the Business Office, Books and Journals Division, American Chemical Society, 1155 Sixteenth Street, N.W., Washington, D.C. 20036. Remit check or money order for $4.50 for photocopy or $2.50 for microfiche, referring to code number JPC-75-2340.
References and Notes (1) G. C. Pimentel and A. L. McClellan, "The Hydrogen Bond", W. H. Freeman, San Francisco, Calif., 1960. (2)H. C. Van Ness, T. Van Winkle, H. H. Richtol, and H. B. Hollinger, J. Phys. Chem., 71, 1483 (1967). (3)E. E. Tucker, S. B. Farnham, and S. D. Christian, J. Phys. Chem., 73, 3820 (1969). (4)A. N. Fletcher and C. A. Heller, J. Phys. Chem., 71, 3742 (1967). (5) A. N. Fletcher, J. Phys. Chem., 76, 2562 (1972). (6)W. B. Dixon, J. Phys. Chem., 74, 1396 (1970). (7)R. Aveyard, B. J. Briscoe, and J. ChaDman, J. Chem. SOC., Faraday Trans.,-I, 69,1772 (1973). (8)c. Campbell, G. Brink, and L. Glasser, J. Phys. Chem., 79, 660 (1975). (9)E. M. Woolley and L. G. Hepler, J. Phys. Chem.. 76, 3058 (1972). (IO) E. M. Woolley, J. G. Travers, B. P. Erno, and L. G. Hepler, J. Phys. Chem.. 75. 3591 (19711. (11) J. 0.Hill, G.Ojelund, and I. Wadso, J. Chem. Thermodyn. 1, 1 1 1 (1969). (12)P. Picker, E. Tremblay, and C. Jolicoeur, J. Solution Chem., 3 , 377 (1974). (13)D. .IEatough, . J. J. Christensen, and R. M. Izatt, "Experiments in Thermometric Titrimetry and Titration Calorlmetry", Brigham Young University Press, Provo, Utah, 1973,p 75. (14)See paragraph at end of text regarding supplementary material. (15) S. N. Deming and S. L. Morgan, Anal. Chem., 45, 278A (1973):see also S.L. Morgan and S. N. Deming, bid., 46, 1170 (1974). (16) See, for example, C. G. Savini, D. R. Winterhalter, and H. C. Van Ness, J. Chem. Eng. Data, 10, 168 (1965). (17)R. Aveyard and R. W. Mitchell, Trans. Faraday SOC.,65, 2645 (1969). (18)F. Smith and I. Brown, Aust. J. Chem., 26, 691-721 (1973). (19)L. J. Beilamy and R. J. Pace, Spectrochim. Acta, 22, 525 (1966).
