J . Phys. Chem. 1991, 95,4915-4918
dipolarity/polarizability parameter. In terms of eq 6 the correlation equation becomes In K = In K+ (a’- ii)B (6’- 6 ) ~ * (b*’ - b*)r*
+
+
or, as written by Kamlet et al. In K = In KO
+
+ bB +
SA*
If, as previously suggested and demonstrated for the acidTPPO/TEPO systems, the electrostatic interactions in the adduct are relatively unimportant when compared to solvent effects, (b’ - b) may be small compared to (b*’ - b*). For all solvents studied by Kamlet et al. the coefficient of F1@)was found to be positive as would be expected if intrinsic enthalpy effects are important. In some solvents the coefficient of FZ(r*)was negative while in the other solvents the coefficient was positive. The negative coefficient of F2(r*)was attributed to solut-olvent interactions which tended to stabilize the free base relative to the adduct. The positive coefficient for F2(r*)was attributed to dipolar attractions between the acid and base in the adduct, Le., to intrinsic effects. This interpretation would mean that there exists a competition, as indicated by eq 7,between intrinsic and solvent effects and that in some cases the dipolar attractions in the adduct outweigh the solutesolvent interactions. It is not possible to resolve this question without additional data. Later investigations by Spencer et al.’ on the hydrogen-bonded adducts of m-cresol with various bases gave In K correlations similar to those obtained by Kamlet et a1.I6 However, a more detailed interpretation of these In K correlations could be made because the enthalpy and entropy components of the free energy were determined. These data7 show that the negative coefficient enters into the enthalpy and entropy correlations in R way that is explicable only in terms of solvent effects. The enthalpy correlation for organometallic adducts with TPPO requires a similar conclusion. Therefore, both the reversal of sign noted by Kamlet et for the coefficient of A* in their correlation equation and the positive coefficient for the A* coefficient are also likely due to solvation
4915
effects rather than to intrinsic effects as originally proposed by these investigators.
Conclusions NMR and calorimetric data have been obtained for a study of solvent effects on donorjacceptor reactions in a variety of solvents with several different acid-base pairs. Enthalpy of transfer data and thermodynamic parameters for these reactions have been shown to be consistent with interpretations suggested by LFE relationships for these systems. A model has been presented which shows that, when entropy and enthalpy data are known, it is possible to interpret solvent effects in terms of specific and nonspecific interactions. It has further been demonstrated that solvent effects on a wide variety of reactions in organic medium, ranging from hydrogen-bonded species to electron pair donorjacceptor adducts to apolar guesthost complexes, can be more fully evaluated by the use of simple concepts that emphasize the solvation entropy and enthalpy. Even in the absence of entropy and enthalpy data it is sometimes possible to rationalize solvent effects if the effects of solvation on the components of free energy are considered. Although the entropy change due to desolvation in aqueous solution has long been considered to be positive, the entropy changes associated with desolvation in organic media have received much less attention. The results of this study and others cited in this work suggest that the entropy effects which accompany the destruction of a solvent cage may be as important in organic as in aqueous media.17 Acknowledgment. The authors are indebted to the National Science Foundation, the donors of the Petroleum Research Fund, administered by the American Chemical Society, the Howard Hughes Medical Institute, the Franklin & Marshall Hackman program, and the Merck Co. for support of this research. (17) Reed, A. E.; Curtis, L. A.; Weinhold, F. Chem. Rev. 1988, 88, 899-926.
Thermochemistry of Acetylenes and Polyacetylenes Sidney W. Benson* and Leslie J. Garland Loker Hydrocarbon Research Institute, Department of Chemistry, University of Southern California, Los Angeles. California 90089-1661 (Received: November 5. 1990; In Final Form: January 15, 1991)
Recent updates on the thermochemistry of acetylenes and conjugated acetylenes together with empirical estimates of heats of vaporization of liquids have made it possible to extend and make new estimates of group increments for heats of formation, AfHo298, of acetylene derivatives. Theses are generally close to earlier values, but the mean precision of agreement with the data is appreciably improved (f0.3 kcaljmol), and a number of new values are estimated for butadiyne, vinylacetylene, 3,5-cctadiyne, 1,3-hexadiyn-5-ene, and phenyl acetylenes.
Introduction The high-temperature chemistry of acetylene and acetylene derivatives has been the object of extensive investigation in the past decade. Much of the interest has centered on the very important phenomenon of soot formation. Soot formation has been of direct and historic interest to engineers concerned with the production of charcoals and active carbons. It has also been an important problem to chemical engineers using high-temperature pyrolysis to produce useful chemicals. Most recently it has become a critical problem in high-temperature combustion, particularly in diesel engines where soot in the exhaust is considered a potentially hazardous pollutant. The conversion of methane to useful organic compounds (C-1 chemistry) had added
to the interest in acetylene since soot is a product or by-product of the high-temperature pyrolysis of CH,.4 In order to understand the quantitative kinetics of the hightemperature chemistry of acetylene and soot formation it is necessary to know the thermochemistry of the intermediates produced. This article will review the thermochemical data of these compounds and present the “best values” currently available. In 1969, Benson et ale1made a comprehensive review of the literature on the thermochemistry of organic and metalloorganic ( 1 ) Benson, S. W.; Cruikshank, F. R.; Golden, D. M.; Haugen, G. R.; ONeal, H. E.; Rodgers, A. S.;Shaw, R.; Walsh, R. Chem. Rev. 1969, 69, 279.
0022-3654/91/2095-4915%02.50/0 0 1991 American Chemical Society
4916 The Journal of Physical Chemistry, Vol. 95, No. 12, 1991
Benson and Garland
TABLE I: h t r Used To Calculate C u r - P k Heats of Formtion compound 1-butyne 2-butyne 1-pentyne 2-pentyne 3-methyl-l -butyne 3,3-dimethyl- 1 -butyne trans-3-penten- 1-yne cis-3-penten- I -yne 1,s-hexadiyne phenylacetylene 1,7-octadiyne I-wten-3-yne trans-3-decen- 1-yne cis-3-decen- 1-yne 5,7-dodecadiyne 3,9-dodecadiyne
3,3,6,6-tetramethyl-1,7-octadiyne 3,3-dimethyl-l,4-pentadiyne
Tb/ K 280.9b*' 299.8"' 3 136.7 328.g6*' 301.86*7 3 10.9699 323.5'O 3 1 71° 35979'0 4 16's8 408' 4058 45 1I I 441" 490' 488' 43710
U T , ?
kcal/mol 5.90 6.29 6.57 6.91 6.34 6.53 6.79 6.66 7.54 8.74 8.57 8.51
9.47 9.26 10.29 10.25 9.18 7.27
hVHm kcal/mol 5.74 6.31 6.74 7.28 6.38 6.67 7.10 6.88 8.36 10.56 10.24 10.12 1 1.99 11.57 13.62 13.54 1 1.41 7.89
-4Cp.298, cal/(mol K) 9.43 10.60 11.34 12.14 10.72 11.23 11.88 11.55
13.47 15.46 15.22 15.12 16.44 16.17 17.36 17.31 16.07 12.93
4H(W kcal/mol 33.91 j: 0.21 28.51 0.29 27.73 22.53 26.33 18.75 0.58 54.51 0.45 54.1 1.1 91.78 A 1.03 62.7 f 1.0 79.9 1.2 33.6 1.5 24.01 f 0.38 23.70 0.60 43.36 f 0.86 47.25 0.72 50.45 f 1.3 83.34 0.83
** * ** *
"Numbers in superscript are references.
gas-phase compounds. They selected *bestn values from which to deduce group contributions to enthalpies of formation, entropies, and heat capacities of organic compounds. The following year, Cox and Pilcher2 in their extensive text on standard enthalpies covered much the same material, but exof formation, AfH0298, tended it to include data on liquids and solids whose heats of vaporization or sublimation had not been measured. A recent book by Pedley, Naylor, and Kirby3 has reviewed this same material and added the relatively few data which have been reported over the past 20 years. This last effort has changed by small amounts some of the data reported in refs 1 and 2. In the present article we shall look at some of these changes as they affect the thermochemistry of acetylenes. We shall also estimate some heats of vaporization of olefin-acetylenes and diacetylenes. This will permit us to deduce group contributions for conjugated acetylenes for acetylenes and polyunand so permit us to estimate AfHo298 saturated acetylene compounds. Most recently, Benson and Cohen,I3 reviewing the thermochemistry of alkanes, have made use of some of the changes incorporated in ref 3 and arrived at new estimates of the basic groups for the saturated hydrocarbons, both gas phase and liquid. In our current review we have made use of these new group values which again differ only slightly from those used earlier.'J Some recent data on the heats of hydrogenation of a series of phenylacetylenesIs in cyclohexane solution can be used to estimate gas-phase heats of formation. These differ markedly from earlier values2 but appear to be more reliable and self-consistent with other data and we have used them here in place of the earlier data.
Experimental Values Experimental values for the gas-phase 4P of alkynes at room temperature are available only for the smaller species (up to 3-methyl-I -butyne).) The changes from previous valuesI2 are on the order of tenths of kcal/mol. For the rest of the compounds of interest, this value must be obtained from the available liquid-phase AfHovalues and heats of vaporization: AfHo298k) AfH0298(liq) + 4 H 0 2 9 8 (1) Heats of vaporization at the boiling point ArHoTk can be. estimated from the normal boiling point, Tb, of the species by using Trouton's rule: 4H077, = (0.021 kcal mol-' K-')Tb (2) For hydrocarbons this is usually reliable to within 5%.5 Except (2) Cox, J. D.; Pilcher, G. Thermochemistry of Organic and Organometallic Compounds; Academic Press: London, 1970. (3) Pedley, J. B.;Naylor, R. D.; Kirby, S. P. Thermochemical Data o/ Organic Compounds, 2nd ed.; University Press: Cambridge, U.K., 1986. (4) Weissman, M.;Benson, S. W. Prog. Eng. Combust. Sci. 1989, IS, 273.
for the C-10 and C-12 compounds, all of the relevant boiling points have been reported."7r9J0 For cis- and tranr-3decen-l-yne, limited vapor pressure data have been reported," and the normal boiling points were calculated to be 172 and 182 OC, respectively, using the Clausius-Clapeyron relation' T Tb = (3') 1 + - RT In (P/atm) 4H0 and with 4 H 0 = 21TB, Tb = T(l - (21/R) In Pam). These calculated boiling points are comparable to those of similar C-10 straight-chain hydrocarbon compounds: compound decane 1-decene 5-decene 1-decyne
Tb/OC 174.1 170.5 170 174
compound 3-decyne 5-decyne 1,3-decadiene
Tb/OC 176 177 170
Note that a 10 K uncertainty in the absolute boiling point would contribute an uncertainty of 0.2 kcal/mol to 4P.The boiling points of 5,7-dodecadiyne and 3,9-dodecadiyne were estimated to be approximately 21 7 OC from those of other C-12 compounds.' compound dodecane 1-dodecene I-dodecync
Tb/OC 216 213 215
compound 3-dodecyne 6-dodecyne I-dodecen-3-yne
Tb/oc
216 209 224
In order to calculate AfHo298(g),the temperature dependence of the heat of vaporization may be found from the heat of capacity of vaporization. = AvHoTb + (&cT)av(T- Tb) (4) The difference in heat capacities of the liquid and the saturated vapor at the normal boiling point (-ACp,TJ is nearly constant at 10.5 cal mol-' K-' for a wide variety of hydrocarbons. By definition, ACp,Tc= 0 since at the critical point there is no difference between liquid and vapor properties. From the empirical Guye-Guldberg rule, Tb/Tc = 0.625 for most regular liquids.16 ( 5 ) Reid, R. C.; Prausnitz, J. M.;Sherwood,T. K. The Properties of Gam and Liquids, 3rd ed.;McGraw-Hill: New York, 1977. (6) Thermodynamics Research Center. T R C Thermodynamic TablesHydrocarbons; The Texas ABM University System, 1988. (7) CRC Hatuilnwk of Chemistry and Physics, 67th ai.;CRC Ress: Boa
Raton. FL, 1986. (8) Benson, S. W.Thermochemical Kinetics, 2nd ed.;Wiley: New York, 1976. (9) Flitcroft. T. L.; Skinner, H. A. Trans. Faraday Soc. 1958, 51, 47. (10) Kupreev, A. 1.; Shimannaev, G. S.Russ, J . Phys. Chem. 1977, 51, 1403. ( I 1) Skinner, H. A.; Snelson, A. Trans. Faraday Soc. 1959. 55, 404. (1 2) Shaw, R. Thermochemistry of Acetylenes. In The Chemfstry ofthe Carbon-Carbon Triple Bond, Part I ; Wiley: New York, 1978; Chapter 3.
Thermochemistry of Acetylenes and Polyacetylenes TABLE U: G ~ s - P ~ M w o mV ~ W W S Iktr"" (in kerl/mol)
compound
otl E
AfPW(g)
~ ~ d ~ t e ~ ~ t r l TABLE III: Ape, Croup Values 4n-+S)
from liquid data
54.51 f 0.17 44.17 f 0.19 34.50
39.65 f 0.3 34.82 f 0.3 34.47
29.88
29.81
3-methyl-1 -butyric
32.87
32.71 25.42 f 0.6 61.61 f 0.5
3,3-dimethyl-I-butyne trum-3-penten-I -yne 60.98 f 1.1 cis-3-pcnten-1-yne 100.14 f 1.0 1,5-hexadiyne pheny lacetylene 73.3 f 0.6 1 ,7-octadiync 90.14 f 1.2 I-octen-3-yne 43.72 f 1.5 36.16 f 0.4 truns-3-decen-l-ync cis-3-decen-l-yne 35.44 f 0.6 5,7-dodecadiyne 56.98 f 0.9 3,9-dodecadiyne 60.19 f 0.7 3,3,6,6-tctramethyl-1 ,7-octadiyne 61.86 f 1.3 91.23 f 0.8 3,3-dimethyl-1,4-pcntadiyne 1-phenyl-I-propync 64.1 0.7' 1-phenyl-1-butyne 59.1 f 0.5' 1-phenyl-1-hexyne 49.7 f 0.5' diphenylacety lene 93.8 f 0.8' diphenylbutadiyne 142.5 f 1.5' 'These data have been taken from heats of hydrogenation in cyclohexane solution and 4P of the saturated phenylalkanes assuming that heats of hydrogenation are the same in liquid and gas phases. The behavior of AC' as a function of temperature is approximately linear, SO from these two points it is possible to derive the relation A,,Cp,~ (cal/(mol K)) +(17.5/Tb)(T- 1.6Tb) ( 5 )
-4c~ (cal/(mol K)) z 28.0 - 17.5(T/Tb)
(6) The data used to calculate AfHo298(g)are shown in Table I. The Clausius-Clapeyron equation makes the assumption that 4 H o is a constant. Equation 4 makes a first correction in assuming that a constant, average value of ( 4 C ,T),v may be used between T and Tb This leads to an appreciabfy modified vapor pressure equation of the form In P = - ( A / T ) + B + D l n T (7)
where D = A&/R = 6 if Tis considerably below the boiling point. Using eq 6 for 4 C T leads to a further correction In P = - ( A / T ) B D In T + ET (8) For regular liquids the coefficients A, B, D, and E can all be evaluated from the Trouton's relation and eq 6 if Tb the boiling point is known. Both eqs 7 and 8 should be fairly accurate below the boiling point of the liquid. ~ H ' T- 28.9 ~ eU Tb A= R 4H0Tb- 10.50 eu Tb B x - 22.89 ln Tb
+ +
Tb
D = 22.89 E = -17.5 eU/RTb The experimental values of AfHo298(g)are given in Table I1 along with those values calculated from AfHo298(liq) data!" For five species, both gas- and liquid-phase data were available; both the direct experimental gas-phase value and a value calculated (13) Cohen, N.;Bonron, S.W. Chemistry of Alkane Hydrocarbons; htai, Ed.;Wilcy: New York. 1992; Chapter I. (14) Stein, S.E.;Fahr, A. J . Phys. Chem. 1985,89, 3714. (I 5) Davis, H.E.;Allinger, N.L.;Rogers, D.W.J . Org. Chem. 1985,50,
3601,
(in
kul/md)
Drevious value'
acetylene ProPYnc 1-butyne 2-butyne I-pentyne 2-pentyne
39.46 f 0.21 34.81 f 0.29
The Journal of Physical Chemistry, Vol. 95, No. 12, 1991 4917
(16) Partington, J. R. Tncrrise on Physical Chemistry; Longmans, Green: London, 1949; Vol. I.
C,-H Cic C,-G C;C,
26.93
27.55 29.20 (29.20)' -10.20 -4.93 -4.16 -4.73 -1.72 -1 -90
+OS0
new value +27.2 +21.4 +28.5 +24.3 +25.8 -10.014
-5.014 -4.7
-1.7 -2.4" +o. I 14 +0.6 +2.0
+6.26 +8.59
(6.78)O
+3.30 (5.68)' 0.80
0.8" -0.6
via the method outlined here are shown in the table. In all cases these values agree to
