Potentiometric Titration of Mercury Using the Iodide-Selective Electrode as Indicator Robert F. Overman Sacannah River Laboratory, E . I . du Pont de Nemours and Co., Aiken, S. C. 29801
MERCURY, even at relatively low concentrations, is an undesirable contaminant in nuclear fuel elements or reactor targets because it catalyzes the corrosion of aluminum. However, mercuric ions are frequently used as a catalyst to dissolve aluminum-clad reactor fuel and target assemblies in separation plants. Therefore, mercury must be removed before those materials that are to be returned to the reactor can be separated and recovered. When the recycle materials are so highly radioactive that they must be handled behind heavy shielding, as is the case with many transplutonium nuclides, analysis of the process solutions for mercury becomes more difficult. Most methods for mercury analysis were unsuitable for remote operation, because color changes are difficult to observe through thick lead glass. The method of Brandenburger and Bader ( I ) , involving amalgamation of mercury onto copper followed by atomic absorption, is limited by the interference of chelating agents and is tedious to perform with highly radioactive solutions. Baumann and Wallace ( 2 ) have described an EDTA potentiometric titration method which is, however, subject to interference by other chelating agents, such as were present in many of our solutions. A more suitable method involving the titration of mercuric ion against iodide with the iodide-selective electrode as an indicator is described. Interfering reactions from ferric ions, peroxide, and chelating agents may be eliminated by appropriate treatment of the sample. The field of ion-selective electrodes has been reviewed by Toren (3) and by Rechnitz (4). Rechnitz ( 5 ) has demonstrated the use of the iodide-selective electrode for the titration of I- by Ag+. The corresponding reaction between Hg2+ and 1- has not been exploited, perhaps because of the complex equilibria in dilute solutions: Hg2+
+ 41-
HgI4'-
E I20
LL-
2 W
80 End Point
40
1.272 ml
1 I
0
E
7
-40
I
1
0.0IOOM Na I , ml Figure 1. Titration of Hg2+ with iodide selective electrode
(1)
As will be demonstrated, however, this titration is quite useful for solutions containing as little as 10+M Hg*+. EXPERIMENTAL
Apparatus. A pH meter with a solid-state iodide ion activity electrode (Model 94-53A, Orion Research Inc., Cambridge, Mass.) and a saturated calomel reference electrode was used. Reagents. The titrant, sodium iodide, was made by dissolving reagent grade NaI in water and then standardizing gravimetrically with AgN03. The iodide was 0.0100M. This stock solution was diluted as required. (1) H. Brandenburger and H. Bader, Helc. Chirn. Acta, 50, 1409 1967. (2) E. W. Baumann and Richard M. Wallace, ANAL.CHEM., 41, 2072 (1969). (3) E. C . Toren, ibid., 40,402R (1968). (4) G. A. Rechnitz, Clzern. Eng. News, 45 (25), 146 (1967). ( 5 ) G. A. Rechnitz, M. R. Kresz, and S . B. Zamochnick, ANAL. CHEM., 38,973 (1966). 616
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ANALYTICAL CHEMISTRY, VOL. 43, NO. 4, APRIL 1971
The mercury solution was made by dissolving a weighed amount of metal in H N 0 3 and diluting with 0.1N "03. The mercury was 0.0100M. This stock solution was diluted as needed. A nonstandardized solution of mercuric nitrate was used in the EMF titration curve shown in Figure 1. A salt solution containing 6M NaN03, 2 M AI(No3)3, and 0.2M Fe(N03)3 was prepared by weighing out reagent Solutions grade crystals and dissolving them in 0.1N "0,. containing ferrous ion were made from FeS04 dissolved in 0.1N HN03. The ferrous solution was made immediately prior to use. These solutions were not standardized. Diethylenetriaminepentaacetic acid (DTPA) was dissolved in H 2 0 and diluted to 0.05M. Procedure. Samples containing from 12 to 1000 kg of mercuric ion were titrated with NaI. An iodide-selective electrode was used to detect the end point. The standard NaI was added from a 1-ml microburet into 50 ml of 0.1N H N 0 3 or H 2 0 . RESULTS
The titration of 10-4M Hg*+ was not affected by adding 1 ml of salt solution containing 6 M NaN03, 2M AI(NO3)3, and 0.2M Fe(NO&. The titration curves in Figure 1
" n i
*t\
L \.
1
3 5 7 Hg2+ in Titration Vessel, M x IO6 1
Figure 2. Stoichiometric ratio of NaI to Hg2+ at apparent end point as a function of Hg2+concentration
Table I. Titration of 1 ml of 5.0 X 10-aM Hg with 0.0100M NaI Titer value, ml 1.010 1.014
Sample composition 50 ml H20, 25 X 3 0 x H201 50 ml HzO, 200 X 30% HtOz 50 ml0.1N "01, 3 ml 1M FeS04 1.002 50 ml 0.1N HN03, 3 ml 1M FeSOa, 200 X salt solutiona 1 . 0 0 8 50 ml 0.1N HN03, 3 ml 1M FeSOa,200 X salt solution, 100 X 30% Hi02 1.011 50 ml 0.1N HNOa, 200 X salt solution 1.009 50 ml 0.1N HNOa, 200 X salt solution, 1 ml 0.5M DTPA* 1.013 50 ml 0.1N "01, 200 X salt solution, 1 rnl 0.5M DTPA 1.017 Av 1.010 Std dev 0.004 a Salt solution: 6 M NaN03 2M Al(NO& 0.2M Fe(No&. b DTPA: diethylenetriaminepentaacetic acid.
+
+
moles of Hg2+ as a function of the Hg2+ concentration are shown on Figure 2. The titrations were made by adding 10-pl aliquots of Hg2+to 50 ml of 0.1N HNO, and titrating with NaI. After the end point was reached, another aliquot of Hg2+ was added, and another titration with NaI was made. No corrections were made for the small changes in volume of the solutions for subsequent titration. The slightly high ratio of I- to Hg2+ (2.04) at the high concentration limit is not considered significant. DTPA decreased the concentration of free Hg2+ in the solution by complexing the mercury. When 1 ml of 0.05M DTPA was added to the 50-ml solution containing 1 ml of 0.005M Hg2+, only 0.500 ml of 0.010M NaI was required to titrate the Hg2+. When Fe(II1) was added to a similar solution, the HgZ+ was released from the DTPA complex by the Fe(III), and 1.013 ml of 0.010M NaI was required to titrate the HgZ+. Other ligands which compete with Ifor HgZ+would be expected to have a similar effect. The presence of Fe(II1) and H202in the sample caused a drifting end point. A stable end point at the proper place was obtained by adding Fe(I1) or sulfamic acid to the solution to decompose the H202. If the titration of the Hg2+was not carried out immediately after the addition of the Fe(II), the presence of excess Fe(I1) in the solution caused a decrease in the amount of NaI required. The amount of H 2 0 2present as a radiolysis product from 1.8 X 1 O I 2 dpm of alpha activity in a 50-ml sample had no effect on the titration. DISCUSSION
The titration of Hg2+ with I- follows the stoichiometry 1:2 down to Hg2+ concentrations of about 5 X 10+M, and useful results can be obtained down to about 1 X 10-6M by calibrating with standard samples. This lower limit corresponds to about 10 pg of mercury in the titration vessel. Complexing agents, if present, can be effectively removed from the reaction by adding a metal with a higher complexing constant than Hgz+, such as Fe(II1). The Hg2+, displaced from the metal complex, can then be titrated. Oxidizing and reducing agents should be absent or decomposed before the titration because any change in the oxidation state of either I- or Hg2+ must be avoided. Some redox reactions require an initiation time for the reaction to start, and it is possible to perform the titration before the redox reaction starts. ACKNOWLEDGMENT
show that the end points are within the experimental precision. As shown in Table I, the relative standard deviation for titrating 1 mg of mercury in each of eight solutions with and without added salts was +0.4%. The number of moles of NaI required to titrate Hgz+ increased as the concentration of the Hg2+ decreased, because of formation of Hg142-. The ratios of moles of NaI to
The author thanks Dr. Elizabeth W. Baumann for helpful discussions of this work. RECEIVED for review October 22, 1970. Accepted January 4, 1971. The information contained in this article was developed during the course of work under Contract AT(07-2)-1 with the U. S. Atomic Energy Commission.
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