Errors in Coulometric Chloride Titrations Due to Photodecomposition of Silver Chloride Charles E. Champion and George Marinenko Institute,for Materials Research, National Bureau of Standards, Washington, D. C. 20234 ERRORSRESULTING FROM PHOTODECOMPOSITION of silver halides have been of concern in precise argentimetry. These errors are particularly important in the determination of halides with coulometrically generated argentous ion ( I ) , since, in high preiision coulometry, in order to reduce errors in time measurement due to switching to a negligible level, titration times of the order of 1000 s or more are preferred. Time intervals of this order of magnitude, as will be shown later, are significant with regard to errors produced by photodecomposition of AgCl. Although Lundell and Hoffman have made a study to show the direction and magnitude of the errors produced by photodecomposition of silver chloride ( 2 ) as applied to gravimetry, the actual rate of photodecomposition of silver chloride is still unknown. Their data showed that the exposure of silver chloride to light causes positive errors when silver is in excess, and negative errors when hydrochloric acid is in excess in the mother liquor. In the present studies, coulometric titrations of chloride are conducted in perchloric acid medium. The photodecomposition of silver chloride in contact with solution liberates free chloride ions (on both sides of the equivalence point) by a mechanism which may be represented as follows (3, 4 ) :
AgCl Cln
+ hv
+ H?O
HClO
--
-+
HC104
Ag
+
’/?
C12
+ C1- + HClO H+ + C1- + [O]
H+
(2)
(3)
Accordingly, positive errors are produced since the chloride ion regenerated by the above mechanism is retitrated. In principle, the rate of photodecomposition of AgCl in contact with solution can be determined by observing changes of potential of a silver wire indicator electrode in a system titrated to the equivalence point. The cool white fluorescent lamps of the laboratory were the source in this case. The illumination on a horizontal plane at the level of the titration cell was 80 foot candles. The measurements were made with a light meter which is accurate to 2 per cent. The change of chloride ion concentration due to photodecomposition of silver chloride in contact with solution, expressed as the number of microequivalents of chloride ions produced in a given interval of time (in a constant volume of 100 ml), was fitted to a cubic equation ( R = at bt2 ct3). The coefficients were determined by the evaluation of simultaneous equations. This equation permits a quantitative correction to be applied to the results for the photodecomposition of silver chloride.
+
+
(1) G. Marinenko and J. K. Taylor, J. Res. Nat. Bur. Std., 67A, (1) 31 (1963). (2) G. E. F. Lundell and J. I. Hoffman, ibid., 4, 109 (1930). (3) I. M. Kolthoff and E. B. Sandell, “Quantitative Inorganic Analysis,” MacMillan, New York, 1952. (4) R. K. McAlpine and B. A. Soule, “Qualitative Chemical Analysis,” D. VanNostrand, New York, 1933.
190
I60
I30
.mv IO0
70
40
0
IO
20
30
40
50
i/eq
Figure 1. Potentiometric end point of coulometric chloride titration using Ag-glass electrode indicator system. The titration curve represents the final 1 % of the titration EXPERIMENTAL
Apparatus. The coulometric circuit and the titration cell are described in previous communications ( I , 5). The potentiometric indicator system consists of a silver-silver chloride indicator electrode, a glass reference electrode (6), and an expanded scale pH meter. The pH meter is used to measure the potential of the indicator electrode. All weighings are made by the substitution method on a 20 g capacity microbalance. The standard deviation for the weighing process is 0.003 mg. Reagents. The supporting electrolyte (1M HC101-1M NaC104) is prepared from reagent-grade perchloric acid and sodium hydroxide. A silver rod (99.995 Ag) serves as the anode for the generation of argentous ions. Single-crystal KCl serves as the analytical chloride sample. Water-pumped nitrogen gas is used for purging the solution to remove the oxygen. Procedure. The supporting electrolyte, consisting of 100 ml of 1 M HC104-1MNaC104 solution, is added to the titration cell, a small amount of HC1 is added and a pretitration is carried out to eliminate ions reacting with Ag+ in the electrolyte and also to saturate the electrolyte with AgCl. A weighed amount of chloride (2 meq sample) is added to the cell and a typical chloride titration (Figure 1) is carried out in the absence of light to the equivalence point (-97.5 mV). At this point the light is turned on and the emf of the indicator electrode is measured at one-hour intervals of time. Finally, the number of microequivalents of chloride, produced by photodecomposition of AgCl in a given time interval, is calculated, making use of a calibration curve prepared under identical conditions.
(5) G. Marinenko and J. K. Taylor, ANAL.CHEM., 39, 1568 (1967). (6) G. Marinenko and C. E. Champion, NBS Tech. Note 425, J. K. Taylor, Ed., p. 45, October 1967. VOL. 41, NO. 1, JANUARY 1969
205
Table I. Effect of Light on Argentimetric Titration of KCI
Sample number 1 2 3 4 5 6
Average :
Light0 99.9831 (1 hr)
Assay, per cent (corrected for exposure) 99. 9653b
...
...
99.9766 (0.5 hr)
99. 9646b
...
99.9764 (1 hr)
... 99.9787 s = 0.0038%
:
99 9589b
...
L’u
I
I
I
I
I
I
I
Without light
... 99.9619
... 99.9640
... 99.9623
99.9629 99.9627 s = 0.0035% s = 0.0012%
I TIME, hrs.
a
The exposure time is indicated in parentheses.
* Value corrected for photodecomposition.
Figure 2. Effects of light on silver chloride in contact with solution Region ab-AgCI precipitateis exposed to light bc-AgC1 precipitate is not exposed to light cd-AgCI precipitate is exposed to light bd‘-represents cd region neglecting the unexposed time interval
RESULTS AND DISCUSSION
Approximately 0.25 t o 2.5 meq-size samples of KCl were titrated under these conditions and measurements were made for the photodecomposition rate of AgCI. A typical potentiometric chloride titration curve is shown in Figure 1. Such a curve is sensitive in the vicinity of the equivalence point t o any change of chloride ion which might be produced by the photodecomposition of AgCI. The change in emf for each one-hour time interval was followed along the potentiometric curve and is represented by the open squares in the expanded rectangle in Figure 1. Thus, the number of microequivalents of chloride produced by photodecomposition of AgCl in a given time interval is determined from these emf differences. The change of emf of the indicator electrode due t o exposure to light when the system is very close to the equivalence point is shown in Figure 2. The number of microequivalents of chloride produced by photodecomposition of AgCl is found t o be a function of exposure time. In the figure, the regions ab and cd represent the sections of the curve where AgCl in contact with solution was exposed t o light. Region bc represents the section of the curve where the AgCl precipitate was not exposed to light. Examination of the “light exposed” and “light excluded” regions of the curve shows that the change in chloride concentration occurs only when the AgCl is exposed to light. By neglecting the region bc (where AgCl was not exposed to light) and displacing the region cd in such a manner that point c of cd coincides with point b of ab, one finds that the segment cd constitutes a continuation of the graphically fitted equation for ab, as illustrated by the dashed curve bd’. Thus one may conclude that the change in chloride concentration near the equivalence point, when the system is exposed t o light, is primarily due to photodecomposition of AgCI. Most coulometric titrations require less than four hours t o complete. Therefore, it was decided t o study photodecomposition of AgCl in the zero to four hour time interval. Five samples of KCI, ranging in size from 0.25 t o 2.5 meq, were titrated in the dark t o the equivalence point. The system was exposed t o light and the amount of chloride produced by photodecomposition was measured at hourly intervals. It was found that the rate of photodecomposition is independent of the amount of AgCl within the indicated limits. The data can be represented by the cubic equation R = at btZ cta, where a, b, and c are fitted parameters, t is the exposure time
+ +
206
ANALYTICAL CHEMISTRY
for the titration (in hours), and R is the number of microequivalents of chloride produced by photodecomposition. The standard deviation for R is 10 per cent. The values of a, b, and c, determined from the graphically fitted curves, are 0.490, -0.090, and 0.010, respectively. The evaluation of R for a specific analysis permits a quantitative correction to be made for the photodecomposition of AgCl occurring during a titration. Six 2-meq samples of KCI were titrated with electrogenerated silver ions. Three of these titrations were conducted in light and the others were excluded from light. The results for the two sets are summarized in Table I. Column 1 contains results for the three titrations that were exposed t o light for a definite time interval. The standard deviation obtained in this set is 0.0038 per cent. The second column represents the results shown in column 1 after the correction was applied for the photodecomposition of AgCI. The third column shows the results of analysis of the same material for the three titrations in which light was excluded. The standard deviation obtained in this set is 0.0012 per cent. From this table it can be seen that the results of analysis for the same material that was exposed t o light are in excellent agreement (0.0002 per cent difference) after the quantitative correction had been applied for photodecomposition with the results of analysis that were not exposed t o light. The error, caused by photodecomposition of AgCl in contact with solution, for a 2 meq-size sample is about 0.02 per cent when the precipitate is exposed to light for a period of one hour. The fitted equation eliminates this bias and provides a significant improvement in the accuracy. The precision for the samples titrated in the presence of light is not significantly changed after the correction for photodecomposition is applied. This larger imprecision results from the propagated uncertainties in photodecomposition and in the analytical method. For example, at one hour of exposure time the uncertainty for the photodecomposition rate is about 0.04 Meq (in terms of 2 meq-size sample, this represents 0.0020 per cent) and the uncertainty in the method is 0.0012 per cent, therefore the combined error is of the same order of magnitude as the standard deviation of the set which was exposed to light.
CONCLUSION
Photodecomposition of silver chloride is a significant source error in high accuracy argentimetric coulometry’ A potentiometric method is described to measure the rate of condiphotodecomposition^ Under tions, the rate Of photodecomposition is about Oe5 peq/hr. In the case of 2-meq chloride sample, this produces a positive error of the order of 0.02 during a titration period of one hour. Using the fitted cubic equation, corrections were applied to titrations of potassium chloride in the presence Of
of light. The corrected results are in good agreement with the results of titrations conducted in the absence of light. For titrations where precision and accuracy of the order of 0.005 are satisfactory, it is sufficient to use the appropriate correction based on the fitted equation when the titration is not protected from light. However, for ultimate precision and accuracy (of the order of 0,0012 per cent), it is necessary to titrate in the absence of light. RECEIVED for review August 16, 1968. Accepted October 2, 1968.
Coulometric Titration of Mercury Fulminate with Electrogenerated lodine A. J. van der Hulst’ TechnologicalLaboratory, National Defence Research Organisation TNO, Rijswijk (Z.H.), The Netherlands
RELATIVELY few methods have been reported for the determination of mercury fulminate, especially when present in small quantities. Mercury fulminate is a substance used as a primary explosive in priming compositions. The amount of active mercury fulminate present is a deciding factor in the effectiveness of the initiation of these mixtures. The content of mercury fulminate is usually determined by a titrimetric procedure known as the Bofors’ method, which is based o n the reaction with sodium thiosulfate ( I , 2). Small amounts of mercury fulminate have been determined by the spectrophotometric determination of mercury (3), and by a cathode-ray polarographic method using a base electrolyte consisting of pyridine and a potassium nitrate solution ( 4 ) . The titrimetric Bofors’ method proceeds according to the following reaction ( I , 5 ) : Hg(0NC)Z
+ 2 NazSzOa+ HzO HgS40a
+
+ 2 NaOH + NaCN + NaOCN
(1)
The amount of thiosulfate consumed is equivalent to the amount of fulminate present. At the same time it appears from this reaction that four hydroxyl ions are produced for each molecule of mercury fulminate. However, this reaction is followed by ( I , 5 ) : HgS40s
-
+ 2 NaOH + NaCN + NaOCN HgS04 + Na2S04+ 2 NaSCN + HzO
(2)
The effect of the second reaction is that the alkalinity of the solution decreases with time. This reaction, however, can be retarded by using a moderate excess of sodium thiosulfate and adding a large excess of potassium iodide. By dissolving mercury fulminate in a known amount of a neutral standard 1 Present address, Central Research Institute of Algemene Kunstzijde Unie N. V., Arnhem, the Netherlands.
(1) T. L. Davis, “The Chemistry of Powders and Explosives,” Wiley, New York, 1943, p 408. (2) “Analytical Methods for Powders and Explosives,” AB Bofors, Nobelkrut, Bofors, Sweden, 1960, p 216. (3) F. H. Go1dman.J. Ind. Hyg. Toxicol., 24,121 (1942). (4) J. Hetman, Tulunta, 3, 127 (1959). (5) R. Philip, Z . Ges. Schiess- u. Sprengsfofw., 7, 109, 156, 180, 198, 221 (1912).
sodium thiosulfate solution and adding excess potassium iodide, the content of mercury fulminate can be determined both acidimetrically and iodometrically using the same solution. After the additionof the reagents anacidimetric titration is performed with standard acid solution using methyl orange as an indicator. It is essential that this titration is carried out immediately after the addition of the reagents, because of the occurrence of Reaction 2. The excess sodium thiosulfate is then determined by back-titration with a standard iodine solution to a starch end point, Speed is less essential in the iodometric titration, because the iodine value of the solution does not change much with time. This iodometric procedure can also be used for the coulometric determination of mercury fulminate. A slight excess of sodium thiosulfate solution, a large excess of solid potassium iodide, and acetate buffer solution are added to the sample, which is either mercury fulminate or a mixture containing this compound. The excess thiosulfate is determined by titrating with electrolytically generated iodine. The thiosulfate titer is also determined coulometrically under the same conditions, omitting the sample. The conditions under which iodine is generated with 100% current efficiency at a platinum anode are well known (6). Rowley and Swift studied the coulometric determination of thiosulfate with electrogenerated iodine and employed it for the indirect determination of small quantities of substances which oxidize iodide (7). The subject of the present work has been the development of a coulometric titration for the rapid and accurate determination of small amounts of mercury fulminate. When mercury fulminate is present in priming compositions, for example, as used in percussion caps, this method is especially suitable for analyzing the contents of one cap. EXPERIMENTAL
Apparatus. Currents accurate to 0.2 were obtained from the Metrohm coulometer E211. The time can be read, accurate to 0.1 second, from a measuring device which is automatically coupled to the generating current circuit. In (6) J. J. Lingane “Electroanalytical Chemistry,” 2nd ed., Interscience, New York, 1958, p 551. (7) K. Rowley and E. H. Swift, ANAL.CHEM., 26,373 (1954). VOL. 41, NO. 1, JANUARY 1969
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