T H E ATOMIC MASS OF POTASSIUM I. The End-point of The Potassium Chloride-Silver Titration * BY CLYDE R. JOHNSOS** AND GDORGE W. LOW, JR.
This article describes a method by which the potentiometer may be substituted for the nephelometer in finding the end-point of certain precise titrations. The method is a general one for a type of titration widely applicable in atomic mass measurements. The technique required is simple, yet it completely avoids “liquid junction” and “activity” difficulties which have hitherto offered an obstacle to the use of the potentiometer in analyzing the systems obtained in work of this nature. The use and scope of the new method may be inferred from the following record of potentiometric and nephelometric analyses of potassium nitrate-nitric acid solutions saturated with silver chloride. The liquids analyzed were similar in every respect to the solution in equilibrium with the precipitated silver chloride at the end-point of the potassium chloride-silver titration. The equilibrium point of this system at o°C. is to be used as a reference end-point in a determination of the potassium chloride-silver ratio. Experimental The experiments consisted in analyzing the unknown saturated solutions for silver and chloride by comparison with standard solutions having almost the same composition. The standard and test systems were prepared from carefully purified reagents with all of the necessary precautions’ to obtain solutions containing equivalent amounts of silver and chloride. Typical Analytical Systems. Each of the two test systems was 0.600 molar in potassium nitrate, 0.308 molar in nitric acid, and contained 13.5 grams of fine-grained silver chloride. Before making up the systems, silver chloride KO.I was allowed to stand in contact with dilute nitric acid containing a small amount of silver nitrate; silver chloride No. z in contact with dilute nitric acid containing a small amount of potassium chloride. These solutions were later washed out with portions of the nitric acid-potassium nitrate mixture. The supernatant liquids of the typical analytical systems were brought to equilibrium at o°C. and analyzed for silver and chloride. While still in contact with the respective precipitates, they were next titrated away from the end-point in opposite directions by known additions of silver nitrate and potassium chloride, and again analyzed. Standard Solutzons. Each standard solution was also 0.600 molar in potassium nitrate, and 0.308 molar in nitric acid, but contained chloride and silver equivalent to 0.600 milligrams of silver per liter. Suitable correction was
* Contribution from the R i c k Chemical Laboratory, ** National Research Fellow in Chemistry. J. Phys. Chem., 36, 1942 (1932).
Princeton Cniversity.
THE ATOMIC MASS OF POTASSIUM
2391
made for the nitrate and potassium added to the standards with the silver and chloride. For the sake of uniformity standard solutions of the same composition were used in all of the nephelometric and potentiometric analyses. It was recognized that this procedure would decrease to some extent the accuracy of the nephelometric analyses of the solutions not a t the end-point. The Analyses. The 4-liter glass-stoppered Pyrex bottles containing the test systems were packed in ice during the entire experiment, and a calibrated thermometer immersed in the solutions invariably registered 0 . 2 O C . The
c W
FIQ.I Cells for E.M.F. Measurements
bottles were shaken once or twice only after the removal of test portions, otherwise twice each day. For each set of six analyses an 80 to IOO milliliter portion of the cold supernatant liquid was withdrawn and filtered through a sintered glass mat. Two 20.00 ml. samples were analyzed nephelometrically for silver and chloride by the standard solution method.' Two I O ml. samples from the same portion were analyzed potentiometrically for silver and chloride with the two cells: Ag/Analytical Solution/Standard Solution/Ag Ag/AgCl/Analytical Solution/Standard Solution/AgCl/Ag. Two forms of cell suitable for the analyses are shown in Fig. I. An essential feature of the design is that it permits one to join the solutions just before the E. M. F. measurements are made. The silver and silver chloride electrodes were prepared in the manner described by MacInnes and Parker,2 with the modifications suggested by Carmody.s However, the customary washing of the silver chloride electrodes with pure water was carefully avoided. For the measurements made a t o°C. the electrodes were placed in the cell only l J. Phys. Chem., 35, 830 (1931); 36, 1942 (1932). 2 8
MacInnes and Parker: J. Am. Chem. SOC., 37, 1 4 5 (1915). Carmody: J. Am. Chem. SOC., 51, 2901 (1929).
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CLYDE R. JOHNSON AND GEORGE W. LOW, JR.
after it had been packed in ice for some time. In general, the electrodes were stored and rinsed in solutions of the composition in which they were to be used, and long exposure to the air was avoided. A uniform procedure was adopted for the analyses. Each pair of electrodes was placed in a portion of the standard solution and short-circuited overnight. Following an initial E. M. F. measurement with the electrodes in the same solution, the electrodes were rinsed and allowed to stand in solutions of the composition in which they were to be used. The electrodes were then placed in position in the cell, and the analytical E. ICI. F. measurements were made, first at 0°C. and then a t 25OC. Immediately following these observations, the contents of the cell were mixed and a final set of readings was made. Each silver and each chloride analysis consisted of four E. M. F. readings taken over a period of thirty minutes after the cell had reached temperature equilibrium. Measuremcrits made with Leeds and Northrup Type K and Students’ potentiometers were equally satisfactory. The standard unsaturated Weston cell was compared with the laboratory standards. All of the work was done in the light of Series OA Wratten Safelights.
Discussion of Results The initial E. M. F. readings were never greater than 0.15millivolt and usually very much less. The final E. &I. F. readings were generally about 0 . 2 or 0.3 millivolts. Silver analysis No. 2 , System No. 2 , in which the final value was 2 . 5 millivolts, was the only marked exception to this rule. Even in this case no correction was applied on the basis of the initial and final readings. However, there is every indication that it would be possible to make significant corrections of this nature which would improve the accuracy of the potentiometric analyses. The results of the analyses are summarized in the following table. For the purpose of comparison, the chloride and silver concentrations in Table I are both expressed as milligrams of silver per liter. That is, the chloride concentrations have been multiplied by the factor Ag/Cl. The “silver added” in the third column gives the total silver which had been added to the system as silver nitrate, or “subtracted” as potassium chloride, at the time the analysis was made. The corresponding additions of these materials were both made on the zzd day of cooling. The analytical E. M. F. readings were remarkably constant and reproducible. The four observations made over a period of thirty minutes rarely covered a range of more than 0.1or 0 . 2 millivolts. Each value given in the appropriate section of the table is based on the average of four such readings, and is calculated from the expression: log C = E/o.ooo1984T 0.778-1. When the silver or chloride concentration, C, is expressed in milligrams per liter, as silver, E is in volts, and T is the absolute temperature. In either cell the electrode in the test solution containing excess chloride was negative, and the electrode in the test solution with excess silver was positive. The limitations of the above expression in interpreting the E. bf. F. date are recognized. The eight final chloride analyses made at 25°C. show the tendency of the silver chloride electrodes to dissolve and saturate‘the solutions
+
THE ATOMIC MASS OF POTASSIUM
2393
TABLE I Summary of Analyses Time Silver No. cooled added Days Mg/Liter:
Nephelometric Analyses Chloride Silver
0.64 0.61 0.63 0.63 0.59 0.62 0.54
0.63 0.61 0.62 0.64 0.63 0.62
26
none none none none none none fo.300
28 29
+0.300 +0.300
0.55
34
+0.300
25
-0.300 -0.300 -0.300 -0,300
16 21 22
18 20
21
28
29 33
Potentiometric A n a l y s e s : At 0°C. At 25°C. Chloride Silver Chloride Silver
0.62 0.64 0.61 0.60 0.67 0.62
1.02
0.62 0.62 0.62 0.60 0.61 0.63 0.49
1.31
0.51
0.79
0.52
0.84 1.02
0.92 0.94
0.62
0.72
0.48 0.49
0.49 0.49 0.79 0.77 0.77
0.75
0.55
0.80
0.57
0.78
0.63 0.66 0.59
0.77
0.62 0.65 0.63 0.77 0.71 0.64 0.77
0.50
0.51
0.49 0.49 0.49
0.49 0~47
0.78
0.50
in the cell. It appears that this source of error does not appreciably affect the potentiometric analyses made a t o°C. Evidence for this belief may be found in the agreement of the first six nephelometric and potentiometric chloride analyses, and also in the symmetry of the sixteen values obtained by the potentiometric analysis, a t o°C., of the solutions containing excess silver and chloride. The fact that the potentiometric chloride analyses apparently give correct values at either o°C. or 2s°C.,when both the standard and test solution contain equivalent amounts of silver and chloride, may mean only that the solution of the electrodes is accompanied by no electrical effect in this unique case. Nevertheless, the observations made by the method described in this article may be used in an actual titration as evidence that the end-point has been reached. Any possibility of serious error due to the solubility of silver electrodes in the 0.3 molar nitric acid is ruled out by the agreement of the first six sets of nephelometric and potentiometric silver analyses. Equal-opalescence “analyses” made by the method of Richards and Willard,l incidental to the first six sets of analyses given in the table, yielded ratios from 1.35 to 1.65,all indicating that the solutions contained excess chloride. It seems that the method leads to a pseudo-end-point in the case of the potassium chloride-silver titration. Investigators who have attempted to determine the potassium chloride-silver ratio have generally used either the above equal-opalescence method or the still more unsatisfactory procedure of Richards and Wells.2 This observation may have some bearing on the fact that authorities have been unable to agree consistently upon an accurate value for the atomic mass of potassium. Thus, in summarizing the results of the various determinations of the potassium chloride-silver ratio made over the 1 2
Richards and Willard: J. Am. Chem. Soc., 32, 32 (1910). Richards and Wells: J. Am. Chem. Soc., 27, 459 (1905).
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CLYDE R. JOHNSON AND GEORGE W. LOW, JR.
past quarter of a century by the classical Harvard methods, Baxterl concludes that “the atomic weight of potassium seems to be in some doubt.” It may be inferred from the data given in the above table that the standard solution method of analysis permits the excess or deficiency of silver in the analytical solutions to be determined to about 0 . 0 2 milligram for every ten grams of silver used, in the case of the potassium chloride-silver titration. The potentiometer may be substituted for the nephelometer in the analyses without loss of accuracy. That is, in so far as the determination of the end-point limits the experimental accuracy, it is possible to obtain analyses correct to one part in 500,ooo by the use of the nephelometric or potentiometric procedures outlined in this report. The present experiments give incidental information which is of interest in connection with the problem of making a precise adjustment to the end-point. The limitation placed on the accuracy of the eight final pairs of nephelometric analyses has been mentioned. In spite of this limitation, it seems very likely that the irregularities in the results of these analyses are in part due to the presence of colloidal silver chloride in the test systems. The additions of silver nitrate and potassium chloride must have resulted in the formation of small amounts of colloidal silver chloride under conditions in which it would be quite stable. The behavior of the colloidal material in the respective systems, on this interpretation of the analytical results, in is accord with Lottermoser’s conclusion2that the positive silver chloride sol is more stable than the negative sol. On the other hand, it is of interest to note that the potentiometric silver analyses, at z~OC., of the same unsaturated solutions used in the nephelometric analyses, show no evidence of the presence of silver having its source in any such colloidal material. The potentiometric analyses a t o°C. would not be expected to reveal the presence of a colloid. Whatever interpretation i s placed on the data, it is evidently advisable to minimize the removals of test portions in the titration. I t is also sound practice actually to adjust the analytical solutions to the correct end-point, where the coagulation of the colloidal material is comparatively rapid. Conclusive evidence that the correct endpoint has been reached may be obtained in any case by comparing the results of the nephelometric or potentiometric analyses with the predetermined solubility of the precipitated compound. The problem of adsorption by the precipitated silver chloride has deliberately been avoided in the present work, as it is one of sufficient importance and difficulty to warrant separate treatment. While the analyses may seem to show that no measurable adsorption of silver nitrate or potassium chloride OCcurs, they are not necessarily conclusive, because of the preliminary treatment of the silver chloride samples. The analytical method used in these experiments would be very suitable for the study of adsorption effects by the examination of successive washings from the precipitated material, in an actual titration. The authors take this occasion to express their thanks to Prof. N. H. Furman for his interest and co-operation in this work. Princeton, New Jersey. Baxter: J. Am. Chem. SOC., 50, 617 (1928). 2 Lottermoser: Alexander’s “Colloid Chemistry,” I, 673 (1926). 1
