Ion-pair formation by some alkali, alkaline-earth, transition-metal, and lanthanide acetates in water at 25" has been studied by addition of inorganic salts to sodium acetate-acetic acid solutions. The corresponding pHs were obtained to f0.002 unit by using the pK of acetic acid to derive the standard potential of the cell before adding the salts. The answers are expressed as dissociation constants at zero ionic strength.CELLS such as ours,(A = anion, M = cation, Y = chloride, nitrate, or perchlorate) , have been used extensively to derive the ion-pair constants of MA, MA,, etc. The general procedure of using standard buffers to calibrate the pH meter, then measuring the pHs of fixed mixtures of m2 and m3 to which varying amounts of m, are added, may err in that the electrodes must be washed free from buffer before they are applied to the test solutions. This was avoided by obtaining the potential E from m1 + m2 alone and the pH calculated from (Stockholm conventions 1)(1) log K(HA) = -pH + log [A-] + l og! , log [HA] * I n this Paper, concentrations denoted m, and those in the Tables, are molarities. A represents acetate ion.

The dissociation constants of the ion-pairs LiOH, NaOH, CaOHf, SrOHf and BaOH+ in water at several temperatures, and of CsCl at 25" C, have been calculated from e.m.f. measurements of the cell Pt, H2/M, OH-, X-/AgX/Ag.Although the alkali metal hydroxides are generally regarded as completely strong electrolytes, there are indications that the first two members do not fully dissociate in aqueous solution. Thus Kolthoff,l using an indicator method, concluded that LiOH has a dissociation constant K of about 0.5, while Darken and Meier 2 obtained 1.2 from conductivity data. The latter authors also estimated that K -5 for NaOH, but they did not regard the evidence for this as very conclusive. However, Bell and Prue 3 also report K -5 for NaOH from their kinetic studies, and point out that activity coefficients also suggest that NaOH does not dissociate completely since its value is about 8 % lower at 0.5 M than

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A glass electrode-AgC1 cell, measured to kO.01 mV, has been used in evaluations of pK values at 25°C of four acids in the range pK 1-2.2. The practical work consisted of calibrating the cell with dilute HCl followed by additions of stock solutions containing the acid ligands. Three of the acids have been previously studied with Pt, H2-AgC1 cells so critical comparisons are possible.Attempts have been made to reduce the dependence of the pK values at zero ionic strength (when these are -c 2.5) upon the value of the ion-size parameter p used in the Debye-Huckel activity coefficient expression. The present calculations are based upon the use of a second parameter Q. Round values of this at different values of have been calculated from some HCI, salt data. A second but smaller feature which has been taken into account is the observations of Bates that standard electrode potentials (E") vary slightly with p if E" is derived by linear analysis. By using these two factors, (method C of table 3) the pK against p dependence is markedly diminished. Also, by using the p and Q parameters appropriate to HCl solutions, pK values close to the averages of those obtained by method C are obtained.For acids of pK values < 2.5, the values at zero ionic strength are markedly dependent upon the value assigned to p in eqn (2) when Q = 0. The dependence is much greater than is found with other methods such as conductivity 3* and the root

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