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The H-function and the acidity of aromatic amines. O'Donnell, Joseph Patrick

Abstract

Our present knowledge of the acidity constants of acids which are too weak to "be measured with respect to the pH scale is limited. In this work discrepancies existing among the results published by previous investigators are reconciled by a reinterpretation of their data. No accepted basicity values are available for concentrated aqueous alkali solutions (greater than 0.1 Molar), since measurements of pH are unreliable in such solutions. The extreme insolubility of compounds which might be used as indicators has prevented evaluation of the basicity of these solutions by spectrophotometric measurement of indicator equilibria. The present work uses the Hammett acidity function concept to develop the H- function for the following systems: For water containing benzyltrimethylammonium hydroxide H- varies from 11.98 for 0.01 Molar to 16.20 for 2.38 Molar; For pyridine-water containing 0.011 Molar tetramethyl-ammonium hydroxide H- varies from 12.27 for water which contains 1.0 mole percent pyridine to 15.43 for water containing 64.7 mole percent pyridine; For a solution of 30 mole percent pyridine and 70 mole percent water H-varies from 14.06 when the concentration of benzyltrimethylammonium hydroxide is 0.03 Molar to 17-65 when the concentration of the base is 2.19; For a solution 50 mole percent pyridine and 50 mole percent water H- varies from 15.67 when the base concentration is 0.01 Molar to 18.91 when the base concentration is 2.12 Molar; For sulfolane-water containing 0.011 Molar tetramethylammonium hydroxide H- varies from 12.39 for 1.79 mole percent sulfolane to 19.18 for 93.4 mole percent sulfolane; For dimethylsulfoxide-water containing 0.011 Molar tetramethylammonium hydroxide H- varies from 12.17 for 5 mole percent dimethylsulfoxide to 18.61 for 70 mole percent dimethylsulfoxide: For aqueous solutions of lithium hydroxide H-varies from 13.43 for 1 normal lithium hydroxide to 14.31 for 5 normal. In addition to these results, adjustments are made to previously published scales for hydrazine-water and ethylenediamine water. The corrected values range from 12.76 for 9.2 mole percent ethylenediamine in water to 15.48 for 31.0 mole percent and 13.06 for 23.2 mole percent hydrazine in water to 15.08 for 51.1 mole percent. These results are obtained using the Hammett acidity function concept from the spectroscopic determination of the ionization equilibria of 24 substituted anilines and diphenylamines. pKa values are reported for these compounds ranging from 2.63 for symmetrical hexanitrodiphenylamine to 18.37 for paranitroaniline. All previous H- determinations have been based on paranitrobenzyl cyanide. This indicator is shown to be unsuitable for this purpose. Results which have been reported by other investigators are adjusted by eliminating the use of paranitrobenzyl cyanide and assigning another indicator as a starting point. This procedure results in a good correlation of all reported H- data so that the largest discrepancy between results reported in this work and those reported previously is found to be 0.26 pka units. A limited amount of data is obtained which permits correlation with sigma-zero substituent constants. For diphenylamines carrying nitro groups in one ring, the acidity change resulting from varying a substituent in the second ring follows closely the result expected on the basis of sigma-zero values. A good correlation is obtained between the acidities reported here and the acidities of the conjugate acids of the same indicators which have been determined through the Haramett technique in sulfuric acid systems. A summary is presented of much of the available information on the relationship between the first and second dissociation constants of dibasic acids. An explanation of the enhanced basicity of the solutions studied is presented which states that the competition of added solvent for available water molecules reduces the amount of water available for hydration of the hydroxyl ion and thus greatly increases its thermodynamic activity.

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