Pekka M. Juusola

Comparison of different methods for calculation of the stoichiometric dissociation constant of acetic acid from results of potentiometric titrations at 298.15 K in aqueous sodium or potassium chloride solutions

Abstract Single-ion activity coefficient equations were determined for the calculation of the molality-scale dissociation constants, Km, for acetic acid in aqueous NaCl or KCl solutions at 298.15 K. The salt alone determines the ionic strength, Im, of the solutions considered in this study. The activity coefficient equations are of the Huckel type, and Km can be calculated by those for a certain ionic strength from the thermodynamic dissociation constant. The data used for the estimation of the parameters for these equations were measured by potentiometric titrations in a glass electrode cell. Three different methods to calculate the experimental Km values from the titration data were considered. The activity parameters for the Huckel equation were determined from the Km values calculated by all of these methods. The parameters obtained by these methods are very consistent with each other and also with the Huckel parameters obtained previously for acetic acid from the literature data measured on Harned cells. The final activity parameters recommended in this study seem to be reliable. Despite the theoretical difficulties associated with the single-ion activity coefficients and the simplicity of the calculation method based on Huckel equations, Km can be obtained by this method almost within experimental error for acetic acid in NaCl and KCl solutions up to Im of about 1 mol kg−1. It was shown in this study, in addition, that the Km values obtained by the Huckel equations agree well also with those calculated by the Pitzer equations.

Determination of the Glass Electrode Parameters by Means of Potentiometric Titration of Acetic Acid in Aqueous Sodium or Potassium Chloride Solutions at 25°C

Single-ion activity coefficient equations are presented for the calculation of stoichiometric (molality scale) dissociation constants Km for acetic acid in aqueous NaCl or KCl solutions at 25°C. These equations are of the Pitzer or Hückel type and apply to the case where the inert electrolyte alone determines the ionic strength of the acetic acid solution considered. Km for a certain ionic strength can be calculated from the thermodynamic dissociation constant Ka by means of the equations for ionic activity coefficients. The data used in the estimation of the parameters for the activity coefficient equations were taken from the literature. In these data were included results of measurements on galvanic cells without a liquid junction (i.e., on cells of the Harned type). Despite the theoretical difficulties associated with the single-ion activity coefficients, Km can be calculated for acetic acid in NaCl or KCl solutions by the Pitzer or Hückel method (the two methods give practically identical Km values) almost within experimental error at least up to ionic strengths of about 1 mol-kg−1. Potentiometric acetic acid titrations with base solutions (NaOH or KOH) were performed in a glass electrode cell at constant ionic strengths adjusted by NaCl or KCl. These titrations were analyzed by equation E = Eo + k(RT/F) ln[m(H+)], where m(H+) is the molality of protons, and E is the electromotive force measured. m(H+) was calculated for each titration point from the volume of the base solution added by using the stoichiometric dissociation constant Km obtained by the Pitzer or Hückel method. During each base titration at a constant ionic strength, Eo and k in this equation were observed to be constants and were determined by linear regression analysis. The use of this equation in the analysis of potentiometric glass electrode data represents an improvement when compared to the common methods in use for two reasons. No activity coefficients are needed and problems associated with liquid junction potentials have been eliminated.

Determination of Stoichiometric Dissociation Constant of Ammonium Ion in Aqueous Potassium Chloride Solutions at 298.15 K

Abstract Equations were developed for the calculation of the stoichiometric (molality scale) dissociation constants, Km, of ammonium ion in aqueous KCl solutions at 298.15 K from the thermodynamic dissociation constant, Ka, of this acid and the ionic strength, Im, of the solutions. Excess KCl was used in the solutions considered so that this salt in practice determined the ionic strength of these solutions. Equations for Km were based on the single-ion activity coefficient equations of the Hückel type. Potentiometric titration data measured with a glass electrode cell were used in the estimation of the parameters for the Hückel equations. By means of the calculation method suggested in this study, Km can be obtained almost within experimental error up to Im of 1.0 mol kg−1 for KCl solutions. The Km values obtained by this method were compared to the literature results obtained from the Harned cell data of Bates and Pinching (1949) and from the concentration cell data of Everett and Wynne-Jones (1938) measured with hydrogen electrodes on cells containing two liquid junctions.

Determination of Stoichiometric Dissociation Constants of Glycolic Acid in Dilute Aqueous Sodium or Potassium Chloride Solutions at 298.15 K

Equations were determined for the calculation of the stoichiometric (molality scale) dissociation constant, Km, of glycolic acid in dilute aqueous NaCl and KCl solutions at 298.15 K from the thermodynamic dissociation constant, Ka, of this acid and from the ionic strength, Im, of the solution. The salt alone determines mostly the ionic strength of the solutions considered in this study, and the equations for Km were based on the single-ion activity coefficient equations of the Hückel type. The data measured by potentiometric titrations in a glass electrode cell and the literature data obtained by Harned cells were used in the estimation of the parameters for the Hückel equations of glycolate ions. By means of the calculation method based on the Hückel equations, Km can be obtained almost within experimental error at least up to Im of about 0.5 mol kg-1 for glycolic acid in NaCl and KCl solutions.