Hydrogen-bonded complex formation. III. Thermodynamics of complexing by infrared spectroscopy and calorimetry

Arnett, Edward M.; Joris, L.; Mitchell, Edward; Murty, T. S. S. R.; Gorrie, Thomas M.; Schleyer, Paul v. R.

doi:10.1021/ja00711a029

Cited by 283 publications

(138 citation statements)

References 6 publications

(11 reference statements)

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“…Usually the errors of the heats of solution of alcohols in alkanes were up to ±(1-2) kJ mol −1 and for other solutions up to ±(0.1-0.3) kJ mol −1 . is range of errors of calculated heats of solution is in agreement with the precise data of direct calorimetric measurements [6][7][8][9][10][11]. Some conclusions can be made from the consideration of the obtained data on Δ soln (Dataset Item 1 (Table)).…”

Section: Methodssupporting

confidence: 79%

“…At the present, there are a lot of data on the heats of mixing of binary liquid systems [1][2][3][4][5] and considerably less data on the heats of solution from direct calorimetric measurements [6][7][8][9][10][11]. In two handbooks [1,2] about 2500 tables of data on the heats of mixing (Δ mix ) of the different liquid solutes in organic and inorganic solvents have been collected, whence the values of the enthalpies of solution were calculated.…”

Section: Introductionmentioning

confidence: 99%

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Heats of Solution of Liquid Solutes in Various Solvents

Kiselev

Shakirova

Potapova

et al. 2013

Dataset Papers in Chemistry

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e values of the heats of solution (2131 solutions) of different liquid solutes in organic and inorganic solvents were obtained from the literature data on the heat of mixing (Δ mix ) in the wide range of concentrations. e limit values of the heat of solution of a solute ( ) in a solvent ( ) (Δ soln ) were calculated from the limit data of the dependence Δ mix / versus at → 0 and the values of that of a solute ( ) in a solvent ( ) (Δ soln ) from the limit data of the dependence Δ mix / versus at → 0, respectively.

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Section: Methodssupporting

confidence: 79%

Section: Introductionmentioning

confidence: 99%

Heats of Solution of Liquid Solutes in Various Solvents

Kiselev

Shakirova

Potapova

et al. 2013

Dataset Papers in Chemistry

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“…Calorimetric measurements of these weak interactions at high dilution in an inert solvent are essentially useless because the H-bond donor and acceptor are incompletely complexed, but this is not a problem in the pure base method (1,2). Ordinarily, hydrogen bonding occurs to a nonbonding (n) electron pair of an electronegative element, but with benzene the n electron system is involved (3).…”

mentioning

confidence: 99%

“…In studies by the pure base method, Arnett et al (1) found the enthalpy of H-bond formation (AHf) of phenol and p-fluorophenol with aromatic hydrocarbons becomes more exothermic (favorable) in the order benzene < toluene < mesitylene. Here AH, represents the interactions of phenol with the aromatic base, less the interactions of the nonhydrogen-bonding model compound, anisole, with the same base.…”

mentioning

confidence: 99%

Enthalpies of interaction of polar and nonpolar molecules with aromatic solvents

Fuchs¹,

Peacock²,

Stephenson³

1982

Can. J. Chem.

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. Can. J. Chem. 60, 1953Chem. 60, (1982. Enthalpies of solution of representative ketones, phenols, alcohols, and ethers have been determined in a,a,a-trifluorotoluene, benzene, toluene, and mesitylene, and combined with heats of vaporization to give enthalpies of transfer from vapor to the aromatic solvents [AH(v+S)]. Comparison of these values with AH(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 x 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 + 0.02kcal/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). On a mesure par calorimttrie les chaleurs de vaporisation des solutes: ether 11-butylique (10,68 + 0,02 kcal/mol), tetramethyl-2,2,4,4 pentane (9,20 + 0,03), 11-pentane (6,36 + 0,02), butanol-1 (12,46 + 0,OO) et methyl-2 butanol-2 (1 1.94 +_ 0,02).[Traduit par le journal]

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“…Over the past 70 years, a large number of experimental studies 13,15,[18][19][20][21][22][23][24][25][26][27][28][29][30][31][32][33][34][35] focused on determining ΔH and Δν. Most of these studies were carried out to validate Badger-Bauer relation for acid-base interactions.…”

Section: Introductionmentioning

confidence: 99%

A brief review of Badger–Bauer rule and its validation from a first-principles approach

et al. 2014

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Understanding the acid-base interactions is important in chemistry, biology and material science as it helps to rationalize materials properties such as interfacial properties, wetting, adhesion and adsorption. Quantitative relation between changes in enthalpy (ΔH) and frequency shift (Δν) during the acid-base complexation is particularly important. We investigate ΔH and Δν of twenty-five complexes of acids (methanol, ethanol, propanol, butanol and phenol) with bases (benzene, pyridine, DMSO, Et 2 O and THF) in CCl 4 using intermolecular perturbation theory calculations. ΔH and Δν of complexes of all alcohols with bases except benzene fall in the range from −14 kJ mol −1 to −30 kJ mol −1 and 215 cm −1 to 523 cm −1 , respectively. Smaller values of ΔH (−2 kJ mol −1 to −6 kJ mol −1 ) and Δν (23 cm −1 to 70 cm −1 ) are estimated for benzene. Linear correlations are found between theoretical and experimental values of ΔH as well as Δν. For all the studied complexes, ΔH varies linearly (R 2 ≥ 0.97) with Δν concurrent with the Badger-Bauer rule yielding the average slope and intercept of 0.053(± 0.002) kJ mol −1 cm and 2.15(± 0.56) kJ mol −1 , respectively.

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Hydrogen-bonded complex formation. III. Thermodynamics of complexing by infrared spectroscopy and calorimetry

Cited by 283 publications

References 6 publications

Heats of Solution of Liquid Solutes in Various Solvents

Heats of Solution of Liquid Solutes in Various Solvents

Enthalpies of interaction of polar and nonpolar molecules with aromatic solvents

A brief review of Badger–Bauer rule and its validation from a first-principles approach

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