Amide interactions in aqueous and organic medium - American

Lebanon Valley College, Department of Chemistry, Annvllle, Pennsylvania. 17003 (Received: October 6, 1980). Enthalpies of solution of amides in organi...
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J. Phys. Chem. 1981, 85, 1236-1241

1236

Amide Interactions in Aqueous and Organic Mediumt J. N. Spencer,* Department of Chemistty, Franklin and Msrshall College, Lancaster, Pennsylvania 17604

Scott K. Berger, Carla R. Powell, Bruce D. Hennlng, Gary S. Furman, Wllllam M. Loffredo, Elaine M. Rydberg, Robert A. Neubert, Chris E. Shoop, and David N. Blauch Lebanon Valley College, Department of Chemlstry, Annvllle, Pennsylvania 17003 (Received: October 6, 1980)

Enthalpies of solution of amides in organic solvents and water have been used in the pure base method to calculate the enthalpy of hydrogen bond formation of the N-H proton of N-methylformamide (NMF) and N-methylacetamide (NMA) to water. NMA forms a 1.2 kcal mol-’ stronger hydrogen bond with water than does NMF. The hydrogen bond enthalpy of formamide and acetamide to DMA, DMF, and water has also been estimated. Both formamide and acetamide form two N-H hydrogen bonds with DMA and DMF but only one with water. The enthalpies for the formation of the water to amide carbonyl hydrogen bonds with NMF and NMA have also been calculated. Each carbonyl can bond to two water molecules with the NMA carbonyl being the better proton acceptor. Water molecules are found to structure about methyl groups contributing 1.7 kcal mol-’ per methyl group to the interaction enthalpy of the amide and water. The methyl group when attached to the amide carbon also alters the electron density in the CON linkage leading to a greater basicity for the acetamide carbonyl and a greater acidity for the acetamide N-H proton. Analysis of the transfers of the amides from organic to aqueous medium allows calculation of the interaction enthalpies of the amides in these solvents. The relative importance of amide hydrogen bonding to water in determining the enthalpy contribution to stability in water may be seen from these interaction enthalpies. Introduction Experimental data on amide hydrogen bonding in aqueous medium is limited. Several theoretical ~tudiesl-~ on amide hydrogen bonding in vacuo and in aqueous solutions have recently been reported but the work of Klotz et al.5v6and of Gill and Noll’ remains as representative of the limited experimental work in aqueous medium. Other worker^^^^ have determined aqueous heat capacities of amides and given additivity relationships for the calculation of these heat capacities. Thermochemical data for amide hydrogen bonding in aqueous solution are of particular interest because of the extension to protein behavior and stability in aqueous solution. Experimental Section One of the calorimeters and the calorimetric procedures have been previously described.l@l2 A Tronac Model 450 solution calorimeter was also used for enthalpy of solution measurements. Purifications of most of the reagents have been given in earlier works.loJ1J3 Baker grade formamide was fractionally recrystallized by repeated freezings and thawings. The purified formamide was stored over molecular sieves under a nitrogen atmosphere. Baker-analyzed reagent grade acetamide was purified by sublimation. Chloroform was washed with water, dried over PZO5, and fractionally distilled under a nitrogen atmosphere. Fisher “spectranalyzed” cyclohexane was stored over sieves and distilled from P,OB under a nitrogen atmosphere. Aldrich 99.97% methylene chloride was dried over Pz05 and fractionally distilled. Enthalpies of solution are given in Table I. (Additional data are available as supplementary material. See paragraph at end of text regarding supplementary material.) The transfer enthalpy, AHt,(A+B), is the difference in the enthalpy of solution of a given solute, in solvents A and B, AH,,(A-+B) = A&(B) - Ms(A). AA”, is defined as the difference between AHt, for the two solutes in a ‘This work was performed at Lebanon Valley College. 0022-3654/81/2085-1236$01.25/0

given solvent pair, AAH, = AH,,(A-+B) - AH\,(A+B). The term van der Waals interactions as used in this work refers to dipole-dipole, dipole-induced dipole, and dispersion interactions. Results and Discussion Solvation in Aqueous Medium. A M t r for formamide and acetamide between DMF and H 2 0 is 1.86 kcal mol-l while AA”, for the same amides between DMA and H 2 0 is 1.92 kcal mol-l. These two Amt,are expected to be identical even though DMA has been shown to be a better proton acceptor than DMF13 and acetamide is a better proton donor than formamide. Because the DMA carbonyl forms a stronger hydrogen bond to both formamide and acetamide and the acetamide N-H forms a stronger bond to both DMF and DMA, these effects will cancel when comparing Amt,. Am,,for NMF and NMA from DMF or DMA to H 2 0 would be also expected to be identical and experiment shows this is so. AAH,, for NMF and NMA from DMF to H20 or from DMA to H 2 0 is the same as the corresponding transfers for formamide and acetamide. Later considerations will show that formamide and acetamide form one N-H hydrogen bond to water and two N-H ~

~

(1)Rossky, J.; Karplus, M. J. Am. Chem. SOC. 1979,101,1913.

(2)Rossky, J.; Karplus, M.; Rahman, A. Biopolymers, 1979,18,825. 1978,100,1387. (3)Del Bene, J. J. Am. Chem. SOC. McKelvey, J. J . Am. (4)Johansson, A,; Kollman, P.; Rothenberg, S.; Chem. Soc. 1974,96,3794. (5)Kresheck, G.; Klotz, I. M. Biochemistry 1969,8,8. (6)Klotz, I. M.; Franzen, J. S. J. Am. Chem. SOC. 1962,84, 3461. (7) Gill, S.J.; Noll, L. J. Phys. Chem. 1972,76,3065. (8)Nichols, N.;Skold, R.; Spink, C.; Suurkuusk, J.; Wadso, I. J. Chem. Thermodyn. 1976,8,1081. (9)Guthrie, J. P. Can J. Chem. 1977,55, 3700. (10)Spencer, J. N.;Sweigart, J. R.; Brown, M. E.; Bensing, R. L.; Hassinger, T. L.; Kelly, W.; Hovsel, D. L.; Reisinger, G. W.; Reifsnyder, D. S.;Gleim, J. E.; Peiper, J. C. J. Phys. Chem. 1977,81,2237. (11)Spencer, J. N.;Gleim, J. E.; Hackman, M. L.; Blevins, C. H.; Garrett, R. C. J. Phys. Chem. 1978,82,563. (12)Spencer, J. N.; Gleim, J. E.; Blevins, C. H.; Garrett, R. C.; Mayer, F. J. J . Phys. Chem. 1979,83,1249. (13)Spencer, J. N.;Garrett, R. C.; Mayer, F. J.; Merkle, J. E.; Powell, C. R.; Tran, M. T.; Berger, S. K. Can. J. Chem. 1980,58,1372.

0 1981 American Chemical Society

The Journal of Physical Chemistty, Vol. 85, No. 9, 1981

Amide Interactions in Aqueous and Organic Medium

1237

TABLE I : Enthalpies of Solution AW,(kcal mol-', 298 K) solvent CCl, H,O DMF DMA EtOAc

NMF +3.89 t 0.13' -1.67 t 0.04 -0.02 i 0.02' -0.25 t 0.03' +1.11 i 0.06'

DMF

NMA

DMA

formamide

acetamide(s)

+0.85 i 0.04" +4.48 f 0.Oga +0.40 i 0.09' -3.52 t 0.04 -3.24 t O . O l b -4.97 i 0.09 +0.49 0.01' +2.33 i 0.01' 0.00 +0.32 2 0.03' +0.01 t 0.01' -0.95 i 0.03 +2.75 i 0.10 -0.10 i 0.04' +0.04i 0.04' 0.00 -1.49 t 0.07 +2.27t 0.06 +0.32t 0.02' +1.47 t 0.15a +0.27 t 0.03' C6H12 +3.25 i O.Ogd +3.20 t 0.12d CH,Cl, +1.16 i 0.07d -1.49 i 0.06d +1.54 i 0.16d -1.57 i 0.08d -0.09 f 0.14d -3.40 i O.Ogd CHC1, +0.24 t 0.02d -2.93 i 0.15d a Data from ref 13. Data from ref 25 corrected to A G e for liquid with heat of fusion given in ref 26. Data from ref 27. d Data from ref 29.

hydrogen bonds to DMF and DMA. The aqueous heat capacities of acetamide and Nmethylacetamide are identical within experimental error? If acetamide formed two hydrogen bonds with water it would not be expected that the heat capacity would be the same as that of NMA. Nichols et a1.* have developed additivity relationships to calculate aqueous heat capacities for the amides. Their relationships predict identical heat capacities for acetamide and NMA. Previous studies in organic medium have shown that the DMA and the NMA carbonyls are 0.3 kcal mol-l better proton acceptors than the NMF and DMF carb0ny1.l~ Theoretical c a l c ~ l a t i o n s have ~ ~ ~ Jshown ~ that the carbonyl of amides can accommodate two water molecules. If these calculations are accepted and if the results obtained in organic medium may be extended to aqueous medium, the DMA or NMA carbonyl would be expected to be solvated by 0.6 kcal mol-l more than the DMF or NMF carbonyl. If A M t r from any organic solvent to H 2 0 is considered for DMF and DMA, the enthalpy effects of H 2 0 structuring about a methyl group may be determined. AHtrfor DMA from DMF to H20involves the following considerations.12 To remove DMA from DMF requires heat to overcome van der Waals (VDW) forces between DMA and DMF molecules; the closing of the cavity in DMF which originally accommodatedthe DMA molecule releases heat. To place the DMA molecule in H20 requires heat to form a cavity and returns heat due to the VDW attractions and H-bond formation with H20. In addition, in H20, due to the structuring of water molecules about nonpolar groups, heat will be evo1ved.l The transfer of DMF from DMF to H20 involves the same considerations. If AHk for DMA and DMF from DMF to H20 are compared, the difference A M , , should be due to the one additional methyl group of DMA over DMF and the better proton-acceptingability of DMA. The cavity and VDW interactions per CH2have previously been found for most of the organic solvents used in this work and similar calculations were used to provide the additional datal2 given in Table 11. The VDW terms for H20interactions per CH2were calculated as previously shown.12 The cavity term for water was obtained from calculations given by Pierotti15and by using an estimated solubility parameter of 4.8 for H20 in conjunctionwith the solubility parameter theory equations.16 Both calculations gave similar results. If the cavity and VDW interactions per CH2 are considered, the following procedure will allow an evaluation of the heat terms which contribute to AAH,; 2.5 kcal mol-l are released when the cavity in DMF closes, 3.1 kcal mor1 are required to overcome the VDW attractions, thus to remove DMA from DMF requires 0.6 kcal m o r more than (14)Scheiner, S.;Kern, C. W. J. Am. Chem. SOC.1977, 99, 7042. (15)Pierotti, R. A. J. Phys. Chern. 1965, 69, 281. (16) Krishnan, C.V.; Friedman, H. L. J. Phys. Chem. 1971, 75,3598.

TABLE 11: Interaction Enthalpies per CH, (kcal mol-')

solvent

cavity

dispersion

dipoleinduced dipole

structurea

cc1,

1.3 2.1 0 1.5 2.3 0.02 1.8 2.4 0.02 2.5 2.9 0.2 2.3 2.8 0.2 1.4 2.2 0.05 1.2 2.1 0 0.3 0.13 0.04 1.7 H2O For each CH, group in H,O restructuring of H,O releases 1.7 kcal mol-'. When cavity, dispersion, and dipole-induced dipole interactions are included the net interaction increment per CH, group is -1.6kcal mol-'. All interaction terms are calculated as previously indicated (ref 12). CHCl, CH,Cl, DMF DMA EtOAc

the removal of DMF from DMF. To form a cavity in H20 requires 0.3 kcal mol-l more heat for DMA than for DMF. The VDW interactions of H 2 0 with the methyl group of DMA return 0.2 kcal mol-l more heat for DMA than for DMF. Thus 0.1 kcal mol-l more heat is required for the introduction of DMA into H 2 0 than for the placing of DMF in H20. A M t r would be expected to be +0.7 kcal mol. The experimental Amt,is -1.46 kcal mol-l leaving 2.2 kcal mol-l unaccounted for. Of the 2.2 kcal, 0.6 is due to the stronger H bonds formed by the two water molecules to the DMA carbonyl. Therefore 1.6 kcal must be attributed to the restructuring of water about the methyl group. The same procedure may be used for all other organic solvent to H 2 0transfers for the DMF-DMA pair. In certain solvents, CC4, CH2C12,and CHC13,specific interactions of the solvent with the solute must be considered. When all transfers have been considered the heat attributed to water restructuring per methyl group is -1.7 f 0.1 kcal mol-l. The heat released due to the structuring of water about a methyl group may also be found by consideration of the transfers of methanol and ethanol from CCll to H20.17-19 When procedures analogous to those just described is used a value of 1.7 kcal mol-l is found for the heat term due to the structuring of water molecules about a methyl group. Because this value is the same as that obtained for the amides, in which case a 0.6 kcal mol-l stronger hydrogen bond was assumed to be formed by the DMA carbonyl than by the DMF carbonyl, the assumption appears reasonable. Similar analysis of the data for propanol and butanol shows that insertion of successive CH2groups into a chain does not bring about a linear increase in the (17) Krishnan, C. V.; Friedman, H. L. J.Phys. Chem. 1969, 73,1572. (18)Arnett, E.M.;Joris, L.; Mitchell, E.; Murty, T. S. S. R.; Gorrie, T. M.; Schleyer, P. v. R. J.Am. Chem. SOC.1970,92, 2365. (19)Arnett, E. M.;Carter, J. V. J. Am. Chem. SOC.1971, 93,1616.

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The Journal of Physical Chemistty, Vol. 85, No. 9, 1981

number of water molecule contacts with a given hydrocarbon chain. The 1.7 kcal mol-' restructuring enthalpy applies to those cases in which structures differ only by one methyl group and that methyl group cannot be part of a chain longer than two carbon atoms. Additional evidence for the enthalpy effects per methyl group produced by the structuring of water may be found from an analysis of the differences in the free energy for the transfer of hydrocarbons to water. The free energy change for the transfer of pure toluene to water is -5430 cal mol-l while that for pure benzene to water is -4620 cal mol-1.20 Thus toluene is 0.810 kcal mol-' more stable in water than is benzene. If the additional methyl group of toluene is stabilized enthalpically by 1.6 kcal mol-l (Table 11, -1.7 kcal mol-I restructuring, +0.3 cavity, -0.2 VDW), the change in entropy produced by the methyl group may be calculated from AAp = -0.810 = -1.6 - T A A S A A S so calculated is -2.7 cal deg-' mol-l. Similar analysis of the transfers of ethylbenzene and the 0,m, and p-xylenes20 gives an average A A S of -2.7 cal deg-' mol-l per methyl group. The entropy changes for the transfers of benzene, toluene, and ethylbenzene to water are given as 13, 16, and 19 cal deg-' mol-l, respectively,20an increment of 3 cal deg-l mol-l per methyl group. The entropy changes for the transfer of ethanol to water and for propanol to water are reported as 10.7 and 13.4 cal deg-' mol-l, respectively,20a difference of 2.7 cal deg-l mol-l. It should be noted that as the length of the chain increases the entropy change per methyl group tends to fall off. For AAS between pentanol and ethanol the average AAS is 2.1 cal deg-l mol-' per methyl group. It has previously been noted from enthalpy data that a similar change in the number of water contacts with straight chain molecules occurs as the chain length increases. The changes in water contacts which occur due to a lengthening of a chain must be roughly compensated by entropy and enthalpy changes so that a linear free energy relationship is found.20 The calculations presented here have indicated that for a straight chain hydrocarbon the number of water contacts tends to decrease with increasing CH2 insertions. This results in fewer water-nonpolar group contacts and less structuring of water molecules. Hence a less exothermic enthalpy term and a less negative entropy term results, which tend to counterbalance so as to maintain linearity of the free energy relationships. For incremental methyl groups up to a chain length of not more than two carbons the difference in free energy, AAp, for a transfer from pure hydrocarbon to water can be found from AAp = -1.6"298(-0.0027)N', where N' is the difference in number of methyl groups. AAp for the transfer of benzene and ethylbenzene from the pure liquids to water would be calculated as AAp = -3.2 298(2)(-0.0027) = -1.6 kcal mol-l. The experimental AAp is -1.610 kcal mol-1.20 Calculation of Hydrogen Bond Enthalpies. The pure base calculation of the hydrogen bond enthalpies is made by use of eq 1.12 The superscripts A and M refer to acid AHHB

- m S M ) b a s e - ( U s A(AHCA - AHC')

+

msM)ref=

+ (miA - miiM) + (AHcM' -

mcq+ (miM-'m y ) (1)

and model compound, respectively, the subscript S refers to the enthalpy of solution, A H C is a cavitation enthalpy, A H i is an interaction enthalpy, the primes refer to refer(20) Tanford, C."The Hydrophobic Effect"; Wiley, New York, 1973.

Spencer et al.

TABLE 111: Hydrogen Bond Enthalpies (kcal mol-') acid base -AH NMF (N-H)

DMF" DMA" EtOAca

2.9 3.1 2.3 2.0 NMA (N-H) 3.6 DMA" 3.9 EtOAc" 2.9 3.2 formamide (N-H) EgF 5.1 (5.2)b DMA 5.6 (5.6jb H.0 2. oc acetamide (N-H) DMF 6.2 (6.5)b DMA 6.7 ( 7 . 0 ) b 3.2' H,O 5.3d NMF (C=O) H2O 5.9d NMA (C=O) H,O From ref 13. For two H bonds. The value first given was obtained by estimates from the H-bond enthalpies of NMF and NMA with H,O. The value in parentheses was calculated as outlined. The formamide and acetamide N-H bonds to H,O are considered to be the same as the NMF and NMA H bonds to H,O. For two H bonds.

%&

ence solvent, and A H H B is the corrected pure base hydrogen bond enthalpy. The terms grouped in parentheses on the right-hand side of eq 1 correct for any difference between an acid and its model compound. In all cases in this work the model compound differs from the acid only by one methyl group. The correction factors for use in eq 1 are given in Table I1 and the enthalpy of solution data are given in Table I. Equation 1 may be used for the calculation of the NMA-water N-H hydrogen bond enthalpy if a suitable model compound is chosen. When DMA is used as the model compound and CCll as the reference solvent eq 1 gives (-3.24 4.97) - (4.48 - 0.40) = AHHB- 0.3 0.2 1.7 1.3 - 2.1 from which AHHBis found to be -3.2 kcal mol-'. Table I11 lists this enthalpy along with other enthalpies similarly determined, The N-H proton of NMA forms a 1.2 kcal mol-l stronger hydrogen bond with water than does the N-H proton of NMF. In organic medium the NMA proton was found to form a 0.7 kcal mol-I stronger bond than the NMF proton. The enhanced ability of proton donation of NMA in aqueous medium is not surprising. The high dielectric constant of water tends to support increased ionization; H-bond formation is but the first step toward ionization. The pure base method may also be used to calculate the enthalpy of hydrogen bond formation of N-ethylacetamide (NEA) with water. According to the procedures previously outlined,12 if DMA is used as the model compound no corrections are necessary because both model compound and NEA have the same number of methyl groups. I t would be expected that the NEA-H20 hydrogen bond enthalpy would be nearly the same as that of the NMAH 2 0 hydrogen bond. The enthalpies of solution of NEA in water and CC14have been reported as -3.70O2l and +4.99 kcal mol-1,22respectively. These enthalpies of solution in conjunction with those for the model compound DMA (Table I) may be used directly in eq 1 to calculate the N-H hydrogen bond enthalpy. The value so obtained, -3.3 kcal mol-l, is in good agreement with that calculated for the NMA-H20 N-H hydrogen bond, -3.2 kcal mol-l. The near identity of these two hydrogen bond enthalpies, obtained in one case by the use of correction factors and in the other

+

+

+

+

(21) Konicek, J.; Wadso, I. Acta Chern. Scand. 1971, 25, 1541. (22) Lindheimer, M.;fitienne, G.; Brun, B. J. Chirn. Phys. 1974, 71, 135.

The Journal of Physical Chemistry, Vol. 85, No. 9, 198 1

Amide Interactions in Aqueous and Organic Medium

TABLE IV : Specific Interaction Enthalpiesa A f f s s , kcal

solute

solvent

NMF (C= 0) DMF (C=O) NMA (C= 0 ) DMA (C= 0) NMF (N-H) NMA (N-H)

mol-'

CCl,

CHCl,

CH,Cl,

-2.2 -2.2 -2.6 -2.6 0 0

-4.8 -4.8

-2.3 -2.3 -2.3 -2.3 -0.2 -0.7

-5.2 -5.2 tO.l -0.8

Calculated according t o discussion given in ref 29. Although the estimated C= 0 interactions of the formamides and acetamides with CH,Cl, are the same, when comparing differences between the two carbonyls the acetamide interaction is consistently more exothermic by 0.3 kcal mol I , a

without any corrections, lends validity to the procedures used and to the corrections used for the restructuring of water. A M H ,for formamide and NMF from DMF to H20 used with the hydrogen bond enthalpies given in Table I11 shows that Amtr can only be rationalized if two H bonds are formed by formamide to DMF and one formamide N-H hydrogen bond is formed with H20. Similar results are obtained for AmH from , DMA to H20. Acetamide also forms two N-H hydrogen bonds with DMA and one with H20. Formation of the second hydrogen bond is not as favorable as formation of the first due to an electronic rearrangement. Previous results for the anilines'l suggest that the overall enthalpy change for the formation of two bonds may be estimated by assuming that 80% of the strength of the first bond is a good approximation for the second. Table I11 gives the estimates for the enthalpies of hydrogen bond formation with water by formamide and acetarnide obtained in this way and also lists estimates obtained from calculations to be described later. If the parameters of Tables 11,111,and IV are used to calculate Amtr all experimental AAH, may be calculated to within f0.2 kcal mol-l. As an example of how these calculations are made consider A M H ,for NMF and DMA from CC14to H20. DMA has two more methyl groups than NMF thus when the cavity in CC14 closes 2.6 kcal mol-I are released, 4.2 kcal mol-1 more are required to overcome the VDW interactions in DMA than in NMF, both NMF and DMA interact specifically with CC14 through the carbonyl, the DMA interaction is 0.4 kcal mol-l stronger, thus to remove DMA from CC14 requires 2.0 kcal mol-l more than to remove NMF. The combination of cavity, VDW, and water restructuring terms for H20 release 3.2 kcal mol-l more for DMA than for NMF, the watercarbonyl H bond releases 0.6 kcal mol-l more for DMA than for NMF, the N-H hydrogen bond energy of NMF to H20 is 2.0 kcal mol-l. Thus the heat released in placing DMA in H20 is 1.8 kcal mol-l more than for NMF. By this procedure Amtr is calculated to be +0.2 kcal m o F , the experimental A M t r is +0.19 kcal mol-'. The work of Klotz et al.51~ suggests that the zero enthalpy of formation of amide-amide hydrogen bonds in water is due t o the equivalence of a C=O 'water plus a N-H NH plus a water-water bond.2 *water bond and a C-== C=O-..NH WW = C=O-.-W + N-H-W AH = 0 (2) The enthalpy change for the NMA-water N-H hydrogen bond calculated by the pure base method refers to a nearly solvation-free reaction because of the use of model compounds which subtract out solvation. Thus the enthalpy change given in Table I11 refers to the reaction NH + W = N-H... W (3)

+

--.

..

Similarly for the carbonyl-water hydrogen bond c=o + = c=o. ..w

w

1239

(4)

If the enthalpy change previously reported for the formation of the NMA-DMA hydrogen bond13is used (Table 111) NH C=O = NH*-.C=O (5) with eq 2-5 and if the water-water dissociation energy is taken to be 5.2 kcal mol-l W + W = WW AH = -5.2 kcal mol-l AH for eq 4 may be found to be -5.9 kcal mol-l. Because NMA carbonyls were used for this calculation the NMF carbonyl to water hydrogen bond enthalpy may be estimated to be -5.3 kcal mol-l. Thus if two water molecules solvate the carbonyl the average water hydrogen bonding enthalpy to an acetamide carbonyl is about -3.0 kcal mol-' and that to a formamide carbonyl is about -2.7 kcal mol-l. Calculation of Transfer Enthalpies. AHH,may be calculated by considering dlthe interactions terms separately. Equation 6 breaks AHtr into its component parts: AH,(A--B) = NAHc,i AH"(B) - AH"(A) "Co(A) + m c o ( B ) + ~ F (6)G A and B refer to solvents; N is the number of methyl groups attached to the functional group C(=O)N(CON); AHc,iis the cavitation term, VDW, and water structuring factor per CH2for the A B transfer (these data are given in Table 11); AH" is the specific interaction enthalpy including hydrogen bond formation at the N-H site (Tables I11 and IV); AHCOis the specific interaction enthalpy including hydrogen bond formation at the C=O site (Table IV). The functional group CON of the amide is considered separately and terms which account for the interactions of methyl groups (NAHc,i)are added to it. A H F G is the change in enthalpy for the functional group when transferred between solvents A and B. Because eq 6 takes into account all specific interactions separately from A H F G , A H F G gives the differences in interactions in the two solvents with CON which are due to VDW interactions and cavitation enthalpies. All terms on the right-hand side of eq 6 may be found from data given in Tables 11,111,and IV except for A H F , ; this is calculated by using the experimental transfer enthalpy. For any solvent pair A H F G calculated for all possible solute transfers is constant to within an average deviation of f0.1 kcal mol-'. A H F G is given in Table V. By use of eq 6 and A H F G calculated for a given solvent pair the enthalpy of hydrogen bond formation for acetamide and formamide with DMA and DMF may be calculated. These enthalpies, which require the formation of two hydrogen bonds by formamide and acetamide, are given in Table 111along with the estimates obtained as described earlier. From AHW it is possible to determine the contributions to the interaction enthalpy in water due to the methyl groups and H bonds formed with water. For example, for the transfer of NMF from CC14 to H20, A H F G NAHc,i is -0.5 kcal mol-l, and the experimental AHtris -5.6 kcal mol-l. The formation of H bonds by NMF to H 2 0 accounts for a large part of the transfer enthalpy, only 1.6 kcal mol-l are accounted for by the structuring of H20. For DMF transferred from CHC13 to H20, AHFG+ NAHc,i is -0.1kcal mol-', the experimental AHtr is -0.59 kcal mol-l,

+

1,23124

+

-

+

(23) Pullman, B.; Miertus, S.;Perahia, D. Theor. Chim. Acta (Berlin) 1979, 50, 317. (24) Gebbie, H. A.; Burroughs, W. J.; Chamberlain, J.; Harries, J. E.; Jones, R. G. Nature (London) 1969,221, 143.

Spencer et al.

The Journal of Physical Chemistry, Vol. 85, No. 9, 1981

1240

TABLE V : Comparison o f A H F G Calculated b y Different Methods solvent CC1, H,O CHCl, -+ H,O CH,Cl, H,O DMF + H,O DMA H,O EtOAc + H,O cyclohexane -+ H,O CCl,-+ DMF CCl, DMA CC1, -+ EtOAc CCl, cyclohexane CHCl, -+ CH,C1, CHC1, + DMF CHC1, DMA CHC1, -+ EtOAc CHCl, +cyclohexane CH,C1, DMF CH,Cl, + DMA CH,Cl, -+ EtOAc CH,Cl, + cyclohexane DMF -+ DMA -+

-f

--f

-+

-+

-+

-+

AHFG~ +0.3 P 0.1 t 1 . 5 P 0.0 t 3 . 1 f 0.2 t 3 . 8 f 0.1 +3.7 i 0.1 +3.1 t 0.1 -0.2 f 0.1 -3.5 P 0.1 -3.3 f 0.1 -2.7 t 0.0 +0.5 f 0.1 -1.6 t 0.2 -2.4 i 0.1 -2.2 f 0.1 -1.5 t 0.0 +1.7 f 0.1 -0.7 P 0.1 -0.6 P 0.2 -0.08 t 0.1 + 3.2 t 0.3 +0.2 P 0.1

A H F G ~ AHF& +0.4 +0.4 +1.6 +1.4 +3.2 +2.9 +4.0 +3.9 +3.7 +3.6 +3.1 t3.0 0.0 -0.1 -3.6 -3.5 -3.3 -3.2 -2.7 -2.6 +0.4 t 0.5 -1.6 -1.5 -2.4 -2.5 -2.1 -2.2 -1.5 -1.6 +1.6 +1.5 -0.8 -1.0 -0.5 -0.7 +0.1 -0.1 + 3.2 + 3.0 +0.3 +0.3

a Calculated from eq 6. Calculated by taking the appropriate difference in H F G terms given in Table VI for DMA. Calculated by taking the appropriate difference in H F G terms given in Table VI for DMF.

and AHFGis +1.5 kcal mol-'. The methyl group effect on the interaction enthalpy of DMF in water is -1.6 kcal mol-', therefore the H-bond contribution to the enthalpy in HzO is only -0.5 kcal mol-l over that in CHC13. Thus the methyl group structuring effect on water accounts for the greater interaction enthalpy of DMF in water as compared to CHC13. However, in most cases the H-bond formation with water is the principle contributor to the exothermicity of the transfer. Because A H F G refers only to the functional group, the values listed in Table V give a measure of the difference in the dipole-dipole, dipole-induced dipole, cavitation terms, and dispersion force interactions between the functional group and solvent for a given solvent pair. For the transfers of the amides from one organic solvent to another the functional group is energetically stable in the more polar solvent due primarily to dipole-dipole interactions. In those organic solvent pairs with nearly equal dipoles such as CHzClzand EtOAc or cyclohexane and CC14 A H F G is small. For the organic solvent to HzO transfers the functional group is energetically more stable in the organic solvents in all cases with the exception of cyclohexane. This increased interaction enthalpy in the organic solvents is due to larger dipole and/or dispersion interactions in the organic solvent. It is also possible to estimate the absolute interaction due to VDW forces between the solvent and the functional group in each solvent. This calculation would also include the cavity term for the functional group. The transfer of DMA from the gas phase to CC14may be found from the enthalpy of vaporization of DMA, +12.7 kcal mol-'?* and the enthalpy of solution in CC14to be -12.3 kcal mol-l. In C C 4solvent a -2.6 kcal mol-' solvation interaction at the DMA carbonyl occurs (Table IV) and 3 X -2.1 kcal mol-' (25) Ojelund, G.; Skold, R.; Wadso, I. J. Chem. Termodyn. 1976,8,45. (26) Kreis, R. W.; Wood, R. H. J. Chem. Thermodyn. 1969, 1, 523. (27) Skold, R.; Suurkuusk, J.; Wadso, I. J. Chem. Thermodyn. 1976, 8, 1075. (28) Weissberger, A. "Techniques of Chemistry"; 3rd e& Wiley-Interscience: New York, 1970; Vol. 11. (29) Spencer, J. N.; Berger, S. K.; Powell, C. R.; Henning, B. D.; Furman, G. S.; Loffredo, W. M.; Rydberg, E. M.; Neubert, R. A.; Shoop, C. E.; Bluach, D. N. to be published in J. Solution Chem.

TABLE V I : H F G for CON in a Given Solvent

-HFG,(I kcal mol-' solvent DMA DMF DMA 10.6 10.1 cc1, 7.3 6.9 6.9 6.5 H,O 8.5 7.9 CHC1, CH,Cl, 10.1 9.4 10.4 DMF 10.9 EtOAc 10.0 9.5 C," 6.9 6.4 a Calculated from heat of vaporization of DMA and DMF, vide infra.

(Table 11) come from the dispersion interactions of the methyl groups. But when DMA is introduced into CC14, a cavity must be made which requires 3 X 1.3 kcal mol-' (Table 11) for the three methyl groups. Thus if the interaction terms are taken into account, the difference between the gas phase to CC14transfer and the dispersion, specific interaction, and cavity terms should give the VDW interactions with the functional group, HFG;or -12.3 - (-8.9 3.9) = H F G = -7.3. Similarly, for the transfer of DMA from gas phase to water, AH* = -17.6 kcal mol-'; 3 X -1.6 kcal mol-l accounts for the VDW, cavity term, and restructuring of water about the three methyl groups of DMA and -5.9 kcal mol-' is the water-carbonyl interaction. For DMA transferred from gas phase to water, HFG = -17.6 - (-4.8 - 5.9) = -6.9 kcal mol-l. Thus AHFGfor the transfer CC14 HzOis -6.9 - (-7.3) = +0.4 kcal mol-l. AHFG given in Table V obtained by use of eq 6 is +0.3. Similar procedures may be used to find H F G for all solvents and the results are listed in Table VI. From H F G it is possible to calculate directly, with the parameters given in Tables II-IV, the transfer from solvent to gas phase, the enthalpy of solution of the pure liquid solute, and the transfer enthalpies between solvents. Interaction Enthalpies in Aqueous and Organic Medium. The energy required to remove an NMA molecule from DMA is found as follows: Each methyl group of NMA is stabilized through dispersion and dipole-induced dipole forces by 3.0 kcal mol-'; the NMA-DMA hydrogen bond stabilization is 3.9 kcal mol-'; the VDW interactions of DMA with the functional group of NMA, H F G , is 10.6 kcal mol-'; when NMA is removed from DMA, the cavity closes returning 4.6 kcal mol-l (2 X 2.3). The enthalpy required to remove NMA from DMA is then +15.9 kcal mol-l. The negative of this enthalpy, -15.9 kcal mol-', is the interaction enthalpy of NMA in DMA. The interactions of NMA in H20 are found in an analogous fashion: the NMA carbonyl water solvation term is -5.9 kcal mol-'; the NMA-water hydrogen bond is -3.2 kcal mol-'; each methyl group contributes -1.6 kcal mol-' for a total of -3.2 kcal molu1; this term includes cavitation and VDW effects; H F G is -6.9 kcal mol-l. The overall hydration enthalpy for placing NMA in H 2 0 is -19.2 kcal mol-l. The transfer of NMA from DMA to H20 is then -19.2 15.9 = -3.3 kcal mol-I. The experimental transfer enthalpy is -3.28 kcal mol-'. It is also possible to calculate H F G from the enthalpy of vaporization of DMF, 11.4 kcal and the enthalpy of solution by the same means used to find HFGfrom DMA data. H F G calculated from DMF data is given in Table VI. In all cases, H F G calculated from the DMF data is about 0.5 kcal rno1-l less negative than H F G calculated from DMA data. The dipolar terms contributing to H m must be larger when the carbonyl of CON is attached to a methyl group. The methyl group is more polarizable than hydrogen and

+

-

+

1241

J. Phys. Chem. 1981, 85, 1241-1243

stabilizes the amide resonance form -O-C=N+ to a greater extent than does hydrogen. Thus the carbonyl of acetamides is more basic than the carbonyl of formamides. Similarly, the proton-donating ability of the acetamide N-H is enhanced over that of the formamide N-H. When the interaction terms found in this work for the amides are used, the hydration enthalpy of a model peptide unit, NMA, in water may be found.

\N

(-6.9) '

CH

(-1.6)

\ H. -3 l2 .1*0H ,

*

/H

The numbers in parentheses give the NMA-water interaction in kcal mol-l. The enthalpy terms at the methyl groups result principally from the energetically favorable restructuring of the water molecules about the methyl groups. For the water molecules at the carbonyl or N-H, a hydrogen bond can be formed, which may take the place of a water molecule previously hydrogen bonded to another water molecule. Consequently, the rearrangement of water molecules about the carbonyl and N-H groups is not significantly altered from normal structure.l The overall enthalpy of NMA in water is then found by summing all the interaction terms. This enthalpy is -19.2 kcal mol-l as previously given. Similar calculations may be made for the amides in all solvents. It is instructive to compare the hydration enthalpy of NMF in water with that of NMA.

0\H

H/)\H;-5.3y

.. **e

H-C-N

1

(-6.5)

/CH3(-l.6)

\

He.

'0-H

j

(-2.0)

H

The overall enthalpy of NMF is -15.4 kcal mol-l as compared to -19.2 kcal mol-l for NMA in water. The additional methyl group of NMA accounts for 1.6 kcal mol-l of the 3.8 kcal mol-l difference, but the remaining 2.2 kcal mol-' come from the effect of the methyl group attached to the carbon on the functional group. The carbonyl of NMF is not as good a proton acceptor as that of NMA and the NMF proton is not as available for hydrogen bonding. In addition, the dipole of the functional group C=ON< is not as great for NMF as for NMA so that the VDW terms are less for NMF. Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for support of this research and to the Research Corporation for funds for the purchase of the Tronac 450 calorimeter. The authors also acknowledge Project SEED for providing funds for one student (D.B.) to assist in this work.

Supplementary Material Available: Heats of solution and mmol quantities of solute used for NMF, DMF, NMA, and DMA in water; formamide and acetamide in DMF and DMA (3 pages). Ordering information is available on any current masthead page.

Gas-Phase Pyrolysis Kinetics of 5-Acetoxy-2-methylpent-2-ene Gabrlel Chuchani," Ignaclo Martin, and Mlguel E. Alonso Centro de QGmica, Instituto Venezolano de Investlgaclones Clenrjfcas,Apartado 1827, Caracas 10 10-A, Venezuela (Recelved: October 6, 1980)

The kinetics of the gas-phase pyrolysis of 5-acetoxy-2-methylpent-2-ene has been measured over the temperature range 330-380 "C and pressure range 53-210 torr. The reaction, in a static system seasoned with allyl bromide, and in the presence of propene inhibitor, is homogeneous, obeys a first-order law, and is unimolecular. The rate constants are given by the Arrhenius equation log k ( d ) = (13.21f 0.14) - (199.6 f 1.7) kJ m o r (2.303RT)-l. The presence of the (CH&C=CH group at the P-carbon atom of ethyl acetate does not provide anchimeric assistance in the elimination of this ester. A simultaneous effect of both steric acceleration and the allylic weakening of the p hydrogen appears to cause a slight rate enhancement of the Z = (CHJ2C=CH group relative to Z = CH2=CH group in the pyrolysis of ZCH2CH20Ac. TABLE I: Temperature Dependence of the Introduction Rate Constants The neighboring olefinic double bond has been found temp, "C 330.2 340.3 350.1 355.1 360.2 370.1 380.1 to assist anchimerically the gas-phase dehydrohalogenation 104k,,s-'0.84 1.63 2.95 3.99 5.62 9.84 17.62 of CH2=CHCH2CH2C1' and (CHS)2C=CHCH2CH2C1.2 These studies were associated with the a-bond participation during solvolysis of their corresponding t ~ s y l a t e s . ~ , ~ participate in the rate of pyrolysis of this ester; it only increases the elimination rate due to an allylic weakening However, the vinyl group in 3-buten-1-yl acetate does not of the Cp-H bonda5 The loosening of C6-H has also been found to be caused by other .Ir-bonds adjacent to the p(1) Chuchani, G.; Hernhdez A., J. A.; Martin, I. Int. J. Chem. Kinet. 1979, 11, 1279 carbon atom in ethyl acetates, where the sequence in py(2) Chuchani, G.; Martin, I.; Alonso, M. E.; Jano, P. Int. J. Chem. Kinet.,1981, 13, 1. (3) Servis, K. L.; Roberts, J. D. J. Am. Chem. SOC.1964, 86, 3773. ( 4 ) Rogan, J. B. J. Org. Chem. 1962,27, 3910.

( 5 ) Martin, I.; Hernindez A., J. A,; Rotinov, A.; Chuchani, G. J.Phys. Chem. 1979,83, 3070.

0022-3654/81/2085-1241$01.25/00 1981 American Chemical Society