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The structure of cis-2-phenylcyclobutanecarboxylic acid (1) and cis-3-(p-fluorophenyl)cyclobutanecarboxylic acid (2) have been determined by X-ray diffraction methods. Crystals of 1 are monoclinic, space group C2/c with a = 15.420(8), b = 11.687(6), c = 11.226(5) Å, β = 112.45(4)°, and eight molecules in the unit cell. Crystals of 2 are monoclinic, space group P21/n, a = 8.038(5), b = 5.405(4), c = 22.69(1) Å, and β = 97.30(4)°, with four molecules in the unit cell. In both compounds the cyclobutane ring is puckered, with dihedral angles of 27° and 31°. The bond joining the substituted ring atoms in 1 is significantly longer (1.581(2) Å) than the other three (average 1.535(2) Å), due to crowding of substituents. In 2 both the carboxyl and phenyl substituents are close to the bisecting geometry, whereas in 1 both substituents deviate from this conformation, the carboxyl more than the phenyl.

A recently developed gc–calorimetry procedure involving a combination of gas chromatography and heat of solution measurements has been used to measure heats of vaporization of 17 esters, including methyl 2,2-dimethylpropanoate, butyl acetate, glyceryl triacetate, diethyl phthalate, methyl 9-octadecenoate, and 12 methyl esters of C3- to C15-fatty acids. Comparison with measurements of ΔHv reported in the literature indicates that the gc–calorimetry method often has smaller uncertainties in measurements (0.25 to 1.00 kJ mol−1) and more consistent methylene increments for the series of fatty acid methyl esters. Two of the experimental ΔHv measurements which differ appreciably from recent literature values have been verified by vaporization calorimetry.

Heats of solution of 10 nonpolar solutes (carbon tetrachloride, perfluoroheptane, diiodomethane, benzene, hexane, decane, hexadecane, cyclohexane, toluene, and mesitylene) in 11 nonpolar solvents (heptane, cyclohexane, carbon tetrachloride, benzene, carbon disulfide, diiodomethane, toluene, mesitylene, tetrachloroethylene, 1,1,2-trichlorotrifluoroethane, and hexadecane) have been combined with solute heats of vaporization to give enthalpies of transfer from vapor to solvent (ΔH(v → S)). The heat of vaporization of diiodomethane has been measured (11.70 ± 0.01 kcal/mol). Correlations of the form ΔH(v → S) = a + bMMR + cMV describe the solvent–solute interaction enthalpies with high precision (correlation coefficients 0.991–0.999). Earlier correlations with only the molar refraction are limited to solutes of similar refractive index. In each solvent ΔH(v → S) becomes more exothermic with an increase in the solute molar refraction term, and less exothermic with an increase in the solute molar volume term. These terms represent solvent–solute dispersion interaction enthalpies, and solvent cavity formation enthalpies, respectively.

The heats of vaporization of 1-octyne (10.11 ± 0.02 kcal/mol), 2-octyne (10.63 ± 0.03 kcal/mol), and 4-octyne (10.21 ± 0.02 kcal/mol) have been determined. Heats of solution of the liquid octynes and n-octane have been measured in heptane, cyclohexane, 1,2-dichloroethane, n-butyl ether, ethanol, triethylamine, dimethyl sulfoxide, butyrolactone, dimethylformamide, and hexamethylphosphoric triamide. Enthalpies of transfer from vapor to each solvent have been calculated. Enthalpies of hydrogen bond formation, calculated by the pure base method, become more exothermic in the above solvent order. Correlations with the Taft–Kamlet solvent parameters π* and β indicate that other polar interactions (presumably dipole – induced dipole) are appreciably larger for 1-octyne than for 2- and 4-octyne.

Abinitio calculations using STO-3G, 6-31G, 6-31G**, and 6-311G** basis sets and standard geometries, and 6-31G and 6-311G** basis sets with 6-31G optimized geometries have been performed on a series of acetylene derivatives (H—C≡C—R; –R is —CH3, —Li, —C≡N, [Formula: see text] —NH2,—C≡CH, —CH—CH2, and —F). Substituent effects on stability have been studied using isodesmic methyl exchange reactions (H—C≡C—R + CH4 → H—C≡C—H + CH3R). STO-3G and 6-31G calculations have been carried out for methyl exchange reactions involving several large substituents including phenyl, and for the tert-butyl exchange reactions of 10 substituents (H—C≡C—R + Me3CH → H—C≡C—H + Me3CR).Geometry optimization has an appreciable effect only on the π donor substituents OCH3 and NH2. None of the smaller basis sets reproduces the orderof stabilities or the absolute values of the 6-311G**//6-31G calculations. The latter are in satisfactory agreement with the very few experimentally based enthalpies of methyl exchange.

Enthalpies of solution of 51 polar and nonpolar organic solutes have been determined in the solvent perfluoro-1,3-dimethylcyclo-hexane. The values have been combined with solute enthalpies of vaporization to give the enthalpies of transfer from vapor to solvent. These solvation enthalpies are for most solutes substantially less exothermic than those measured in other solvents. Polar interactions appear to be stronger than with hydrocarbon solvents; but this may result, in part, from incomplete dissociation of strongly associated solutes in the perfluoro solvent.

Heats of solution of several aromatic solutes (benzene, toluene, mesitylene, nitrobenzene, α,α,α-trifluorotoluene, anisole) and model compounds (n-butyl methyl ether, cyclohexane) in 17 organic solvents (n-heptane, cyclohexane, carbon tetrachloride, 1,2-dichloroethane, α,α,α-trifluorotoluene, triethylamine, butyl ether, ethyl acetate, dimethylformamide, dimethyl sulfoxide, benzene, toluene, mesitylene, t-butyl alcohol, 1-octanol, methanol, 2,2,2-trifluoroethanol) have been combined with solute heats of vaporization to give enthalpies of transfer from vapor to solvent (ΔH(v → S)). Differences between solute and model values (ΔΔH(v → S) = ΔH(v → S) (aromatic solute)–ΔH(v → S) (model) were used to evaluate aromatic solute–solvent polar interactions. Correlations of ΔΔH(v → S) with solvent dipolarity–polarizability (Taft–Kamlet π* parameter) have been determined.

A recently developed empirical method for measuring beats of vaporization (ΔHv25) by a combination of gas chromatography and calorimetric heats of solution in the stationary phase has been applied to a series of 9 monoalkylcyclohexanes having C1–C12 side chains. The uncertainties in the measurements (0.07 – 0.23 kcal/mol) are critically evaluated. Comparison with literature ΔHv measurements indicates good agreement for low boiling compounds, and for high boiling compounds, values substantially more reliable than were previously obtained from vapor pressure measurements.

Heats of solution of 5 – 8 C5–C16 alkane liquids have been determined in each of 19 organic solvents covering the Taft–Kamlert π* (dipolarity-polarizability) range −0.08 to 1.00. Enthalpies of alkane transfer from vapor to solvents have been calculated (ΔH(v → S) = ΔHs − ΔHv). The methylene increments in (ΔH(v → S) in the 19 solvents and water (−1.18 to −0.67 kcal/mol) are not well correlated with the Hildebrand solubility parameter δ, but an empirical correlation exists with solvent π* and polarizability:[Formula: see text](n = 20, r = 0.977, see = 0.032). By this criterion the small methylene increment in water appears to be associated with the high dipolarity and low polarizability of water.

Enthalpies of solution (ΔHS) of a series of ketones (acetone, 2-butanone, 2-heptanone, 2-nonanone, 5-nonanone, 2,2,4,4-tetramethyl-3-pentanone, cyclohexanone) and alkane model compounds (n-heptane, n-nonane, 2,2,4,4-tetramethylpentane, cyclohexane) have been determined in 17 organic solvents (n-heptane, cyclohexane, CCl4, α,α,α,-trifluorotoluene, 1,2-dichloroethane, triethylamine, butyl ether, ethyl acetate, DMF, DMSO, benzene, toluene, mesitylene, 1-octanol, methanol, t-butyl alcohol, 2,2,2-trifluoroethanol), and combined with heats of vaporization to give enthalpies of transfer from vapor to solvent (ΔH(v → S)). These values have been used to evaluate ketone–solvent polar interactions (ΔΔH(v → S) = ΔH(v → S)(ketone) − ΔH(v → S)(alkane)). The linear relationships between ΔΔH(v → S) and solvent dipolarity-polarizability (Taft-Kamlet π* parameters) are derived. Based on the deviations from these correlations, ketone–CF3CH2OH enthalpies of hydrogen bond formation have been evaluated. The other alcohol solvents show no evidence of exothermic H-bond formation with ketones.

Heats of solution of m-cresol, 1-butanol, 1-pentanol, t-amyl alcohol, and model compounds (toluene, ethyl ether, n-butyl methyl ether, t-butyl methyl ether) in 17 organic solvents (n-heptane, cyclohexane, carbon tetrachloride, 1,2-dichloroethane, α,α,α-trifluorotoluene, triethylamine, butyl ether, ethyl acetate, dimethylformamide, dimethyl sulfoxide, benzene, toluene, mesitylene, t-butyl alcohol, 1-octanol, methanol, 2,2,2-trifluoroethanol) have been combined with solute heats of vaporization to give solvation enthalpies (ΔH(v → S)). Dependencies of solute vs. model solvation enthalpy differences on solvent dipolarity–polarizability and hydrogen-bond-accepting basicity were determined via correlations with Taft–Kamlet solvatochromic parameters (π*, β, ξ).m-Cresol is a substantially stronger H-bond donor than 1-butanol, 1-pentanol, and t-amyl alcohol, and H-bonds to acceptor solvents including alcohols. Cresol acts as an H-bond acceptor with the strong H-bond donor solvent trifluoroethanol.

Enthalpies of solution of representative ketones, phenols, alcohols, and ethers have been determined in α,α,α-trifluorotoluene, benzene, toluene, and mesitylene, and combined with heats of vaporization to give enthalpies of transfer from vapor to the aromatic solvents [ΔH(v→S)]. Comparison of these values with ΔH(v→S) of nonpolar model compounds provides an estimate of the special interactions (dipole – induced dipole or charge transfer, and hydrogen bonding) of the polar solutes with each aromatic solvent. Unless the model compound is perfectly matched, an alternative procedure, the pure base method, is superior for evaluating hydrogen bonding.The special interactions of ketones and ethers with aromatic solvents increase with decreasing π electron density of the solvent. By contrast, m-cresol shows the strongest special interaction with the most electron-rich solvent (mesitylene > toluene > benzene > trifluorotoluene). These results demonstrate that —OH and —OR groups undergo quite different interactions, and that a distinctive interaction occurs between the hydroxyl group and the aromatic ring.Calorimetric heats of vaporization have been measured for the solutes n-butyl ether (10.68 ± 0.02 kcal/mol), 2,2,4,4-tetramethylpentane (9.20 ± 0.03), n-pentane (6.36 ± 0.02), 1-butanol (12.46 ± 0.00), and 2-methyl-2-butanol (11.94 + 0.02).

The structures of cis-2-(m-nitrophenyl)cyclopropanecarboxylic acid (1) and 1-phenyl-2-methylcyclopropene-3-carboxylic acid (2) have been determined by X-ray methods. Crystals of 1 are triclinic, space group [Formula: see text], with a = 7.422, b = 8.615, c = 9.499 Å, α = 60.24, β = 63.76, γ = 71.36°, and two molecules in the unit cell. Crystals of 2 are monoclinic, space group P21/c, with four molecules in a unit cell of dimensions a = 9.345, b = 13.286, c = 8.187 Å, and β = 98.22°. The carboxyl group of 1 approaches the bisecting conformation. The phenyl group is 37° from this conformation, by contrast with the unhindered compound 2-(p-nitrophenyl) cyclopropyl methyl ketone. In 2 the carboxyl group bisects the three-membered ring, indicating that carbonyl-cyclopropene π interactions can occur if sterically allowed. The two rings are essentially coplanar, permitting maximum interaction of the phenyl group and the double bond. In both compounds, the carbonyl oxygen of the carboxyl is the nearer oxygen to the three-membered ring.

The heats of vaporization of 1-methylcyclopentene, 3-methylcyclopentene, ethylidenecyclopentane, 1-ethylcyclopentene, methylenecyclohexane, allylcyclopentane, vinylcyclohexane, ethylidenecyclohexane, allylcyclohexane, 3,3-diethylpentane, 2,2,4,4-tetramethylpentane, and trans-2,2,5,5-tetramethyl-3-hexene have been measured by the gas chromatography – calorimetry method. These values have been combined with previously reported liquid heats of formation to give gaseous values of ΔHf. The results indicate that the internal double bond is favored by about 0.5 kcal over the exo in both 5- and 6-membered rings, but the endo–exo differences are much smaller than previously believed. Several of the liquid heat capacities that were measured were not well predicted by group additivity schemes.

The accelerating effect of added dimethyl sulfoxide on the rate of hydroxide-catalyzed ionization of chloroform in water has been examined in terms of the solvation of reactants and transition state. Enthalpies of solution of chloroform in water, DMSO, and mixtures of the two have been measured. From these, together with enthalpies of activation for the detritiation of chloroform, and enthalpies of transfer of hydroxide ion from water to aqueous DMSO mixtures, transition state enthalpies of transfer have been calculated. The reaction is accelerated by small concentrations of DMSO because of increased transition state solvation, but by large DMSO concentrations due to reduced solvation of hydroxide ion. Transition state solvation is unlike that of the reactants, but resembles that expected for CCl3− ion.

The following heats of vaporization at 25 °C (ΔHv0) have been measured by vaporization calorimetry: cyclopropyl methyl ketone (39.41 ± 0.09 kJ mol−1). methyl cyclopropanecarboxylate (41.27 ± 0.06 kJ mol−1), cyclobutyl cyanide (44.34 ± 0.20 kJ mol−1), methyl cyclobutanecarboxylate (44.72 ± 0.09 kJ mol−1), ethylcyclobutane (31.24 ± 0.03 kJ mol−1), and cyclopentyl cyanide (48.12 ± 0.09 kJ mol−1). Values of ΔHv0 of the last four compounds have been combined with liquid state heats of formation to give new values of the gaseous state heats of formation (ΔHf(g)).When cyclobutyl cyanide, ethylcyclobutane, and cyclobutylamine are involved in hypothetical isodesmic reactions with ethane to give propane and isopropyl derivatives, the enthalpies of reaction are −102.1 ± 1.7 kJ mol−1 (−24.4 ± 0.4 kcal/mol), 8.3 kJ mol−1 (2.0 kcal/mol) less exothermic than the value for cyclobutane. Possible origins of this difference are discussed.

Enthalpies of transfer from the vapor to solution (ΔH(v → S)) in cyclohexane are similar for alkenes and the corresponding alkanes. The transfer from vapor to dimethylformamide is about 0.6 kcal/mol more exothermic for 1-alkenes, which results from the dipole-induced dipole interaction of the solvent dipole with the π system. When the ethylenic bond is hindered by large adjacent groups (trans-1,2-di-tert-butylethylene) to the approach of solvent dipoles, ΔH(v → DMF) becomes similar to that of the corresponding alkane (1,2-di-tert-butylethane). Similarly the interaction of di-tert-butyl ketone with DMF is 0.7 kcal/mol weaker than that of the less hindered isomer, 2-nonanone.

The enthalpies of transfer from methanol to dimethylformamide for a wide variety of aromatic compounds are additive properties of the groups present within the molecules. Values of ΔΔHs fail to be additive for benzaldehydes by an amount of energy which increases with increasing electron withdrawing power of the substituent. Nuclear magnetic resonance studies indicate that this results from hemiacetal formation in methanol, which varies from nil for p-dimethylamino-, p-hydroxy-, and p-methoxybenzaldehyde to 71% for the p-nitro compound. The non-additivity of group ΔΔHs values can indicate that specific interactions are occurring with a solvent, even though no stable product can be isolated.

The reaction of hydrogen chloride with primary alcohols, which occurs slowly and in poor yields in water, has been found to give good yields of alkyl chlorides in 30–45 min at 70–85 °C in the solvent hexamethyiphosphoric triamide. 2-Octanol and neopentyl alcohol required longer reaction periods. Neopentyl chloride was formed without contamination by rearrangement products.

The heat capacities of 25 liquid aliphatic, alicyclic, and aromatic ketones and aldehydes at 298.2 K and atmospheric pressure have been measured calorimetrically with a typical uncertainty of 0.5%. Additive parameters are proposed for estimating Cp(1) of ketones and aldehydes.

Mesure des chaleurs de dissolution des solutes polaires et non polaires dans divers solvants et combinaison de ces resultats avec les chaleurs de vaporisation pour donner les enthalpies de transfert de la vapeur aux solvants