1937
ACIDITYMEASUREMENTS AT ELEVATED TEMPERATURES
Acidity Measurements at Elevated Temperatures.
IV.
Apparent
Dissociation Product of Water in 1 m Potassium Chloride up to 292"l
by R. E. Mesmer, C. F. Baes, Jr., and F. H. Sweeton Reactor Chemistry Division, Oak Ridge National Laboratory, Oak Ridge, Tennessee 87830
(Received October 27, 1060)
An apparatus for measurement of acidities at elevated temperatures is described which employs a hydrogenhydrogen ion concentration cell in a Teflon-lined vessel. Nernst behavior to high precision was demonstrated at 150". The apparatus was used to measure the apparent dissociation quotient of water (Qw') assuming complete dissociation of KOH and HC1 in 1 m KC1 from 50 to 292". The results at 50" agree with the best published data to 0.01 log unit and the values range from -13.01 at 50" to -10.04 at 292". The data are represented within experimentalerror by the equation log Q,,.' = - (5909.13/T) 0.007279T - 27.3973 log T 71.6761. Differences between these values and values of log K , reported by Noyes and coworkers from early conductancemeasurementsare discussed in terms of activity coefficients, possible association of KOH and HC1, and the uncertainty in conductance measurements.
+
+
KC1 within the temperature range 60-300". The word apparent is used (and QW'is primed) to indicate that the evaluation of this product is made with the assumption that KOH and HC1 are completely dissociated electrolytes under the conditions of the measurements. The relationship between Qw'and K , is
Introduction
I n previous studies of this series, the hydrolysis behavior of metal ions in aqueous solutions up to 95" was reported.2 The extension of such studies to much higher temperatures is of interest from a fundamental as well as practical view. There are presently no techniques available, however, for measurements with the precision needed for identification of hydrolytic species in solution from potentiometric data at where a, is the activity of water and YH and Y O H are temperatures above 100". the stoichiometric activity coefficients. The utility of glass electrodes is limited to temperaAt the present time the dissociation constant for tures below about 150" because of the rate of attack of water as well as the apparent dissociation product in acids and particularly bases. 2a Hydrogen electrodes, salt media are well defined by results of Harned and his in spite of the chemical reactivity of hydrogen, might coworkers4 up to 60". The only high-temperature provide the best means for accurate measurements of studies carried out between 100 and 300" at the saturaacidity since many of the metal ions are stable in hydrotion pressures were the early conductance work of gen, and hydrogen electrodes have already been s h o w Xoyes, et aZ.,b and the recent similar work by FisheraB to function at high temperatures and pressures. HainsBoth measured the hydrolysis of ammonium acetate. worth, et C L Z . , ~ ~have employed hydrogen electrodes up to 1000 atm at 25" and Lietzke and his c o ~ o r k e r s ~ ~ The - ~ ~uncertainty of this method has not been assessed but is expected to be relatively high judging from the have employed hydrogen electrodes in combination poor reproducibility in the two studies, particularly with silver halide electrodes to determine the thermobelow 200". dynamic properties of HX-NIX, mixtures up t o 275". I n cells of the type Pt,HzlHCl(m)/AgCl,Ag, where (1) Research sponsored by the U. 5. Atomic Energy Commission low pressures of Hz were used, Lietzke, et al., observed under contract with Union Carbide Corp. a small amount of the reaction (2) (a) C. F. Baes, Jr. and N. J. Meyer, I n o r g . Chem., 1 , 780 (1962);
'/zHz
+ AgCl+
Ag
+ HC1
(11
which gave rise to drifts that had to be corrected for in solutions of HC1 from 0.01 nz to 1.0 7n and at temperatures greater than about 150". We have developed a hydrogen electrode concentration cell which is intended for precise measurements of acidity to 300". An initial application of this apparatus has been a determination of the apparent concentration product lor the dissociation of water (Qw')in 1 m
(b) C. F. Baes, Jr., N. J. Meyer, and C. E. Roberts, {bid.,4 , 518 (1965); (c) R. E. Mesmer and C. F. Baes, Jr., ibid., 6 , 1951 (1967). (3) (a) W. R. Hainsworth, H. J. Rowley, and D. A. MacInnes, J . Amer. Chem. SOC.,46, 1437 (1924); (b) M. H. Lietzke, et al., J . Phys. Chem., 64, 652 (1960); (c) ibid., 64, 1445 (1960); (d) ibid., 6 4 , 1861 (1960); (e) ibid., 6 8 , 3043 (1964); (f) ibid., 69, 2395 (1965); ( 8 ) ibid., 70, 756 (1966); (h) ibid.. 71, 662 (1967). (4) H. S. Harned and B. B. Owen, "The Physical Chemistry of
Electrolytic Solutions," Van Nostrand-Reinhold Co., Princeton, N. J., 1958, pp 633-696. (5) A. A. Noyes, Y. Kato, and R. B. Sosman, J. Amer. Chem. SOC., 32, 159 (1910). (6) J. R. Fisher, Thesis, "The Ion-Product Constant of Water to 350°C," Pennsylvania State University, June, 1969.
Volume 743Number 9 April SO, 1970
1938 Recently, Marshall and Quist of this laboratory have estimated the ionization constant of water up to 800" and 4000 bars from conductance data 011 ammonium bromide solutions.' The data do not extend below about 300" but they are in approximate agreement with those of Noyes, et al.
Experimental Section 1Uate~ials. A stock solution of 3.3 M KCl prepared from J. T. Raker analyzed reagent was acidified to pH 3.5 and purged with Nz to remove GOz. The fluoride content of the neutralized stock was 7 X m as determined by the lanthanum fluoride, electrode. The protolytic impurities in a 1 11% solution made from the stock were
