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Anal. Chem. 1981, 53, 2141-2143
AIDS FOR ANALYTICAL CHEMISTS Mlcroprocessor-Controlled Determinatlon of Fluoride in Environmental and Biologlcal Samples by a Method of Standard Additions wlth a Fluoride Hon Selective Electrode K. A. Phiillps" and C. J. Rix Department of Applied Chemistry, Royal Melbourne Institute of Technology, 124 Latrobe Street, Melbourne, 3000, Vlctoria, Australia
This work is a development of the determination of fluoride in seawater by Rix et al. (I). In that work, addition of buffers was unnecessary because the seawater electrolytes maintained a constant background ionic strength and their presence did not affect the calculated result. Various buffers have been used to control the ionic strength and pH for fluoride analyses in water samples with low salinity or in biological samples with complicated and variable background matrices (2-5). In the present study we have found that the general purpose Total Ionic Strength Adjustment Buffer called TISAB (6) is satisfactory. EXPERIMENTAL SECTION Apparatus. A Radiometer titration system comprising a TTT6O titrator controller, a PHM64 research pH meter (which has BCD output of the potential), an ABU13 Autoburette (to which was added a B246 electronic counter module to give BCD output of the volume delivered from the buret), and a TTAGO titration assembly (stand, electrode holder and stirrer) was used. Fluoride ion selective electrodes used were an Orion Model 96-09 (with built-in Ag/AgCl reference electrode) and a Radiometer Model F1052F with a separate Radiometer Model K801 Ag/AgCl reference electrode. All solutions were pipetted with a Macro-Set pipet (Model 8900, Oxford Laboratories Inc., Foster City, CA) equipped with disposable polyethylene tips. A 6502-microprocessorwas used to read the BCD outputs of potential and volume and to control the titrator as follows: (a) wait for an initial preset time to allow the system to reach steady state, (b) read potentials at fixed time intervals until they become constant within operator-definedlimits; (c) stir the solution while adding an amount of titrant; (d) stop stirring and repeat the sequence from step (b) until the required data are collected. Reagents. All chemicals used were of reagent grade purity. Deionized water distilled in a glass-lined still was used to minimize background fluoride levels. Standard fluoride solutions were prepared from sodium fluoride dried in an oven at 110 "C for at least 8 h. Procedure. Surface water samples were filtered on site using a Sartorius SM165-10 polycarbonate filtration unit employing a 0.45-pm Sartorius filter fitted with a Nalgene hand pump and stored in polyethylene containers at ambient temperature. Urine samples were collected in sterile plastic containers and immediately frozen and stored in a domestic freezer. Fluoride analyses were made on 5.00-mL aliquots of sample mixed with 5.00-mL aliquots of TISAB in polyethylene vessels thermostated at 25.0 i 0.1 "C. The initial potential indicated the approximate fluoride concentration in the buffered sample and the standard NaF used as titrant was about 100 times this concentration. A total of 2.5 mL of titrant was added in amounts such that the successive buret readings formed a geometric progression. This gives approximately evenly spaced potentials over a suitable range for the regression analysis. Different titrant concentrations or volumes can cause curve fitting programs to give absurd results. Both the Nernstian slope of the electrode (i.e., the potential change per 10-fold change in concentration) and the fluoride concentration in the buffered sample can be calculated from a standard additions analygis using the program of Brand and Rechnitz (7) which uses a successive approximations method of 0003-2700/81/0353-2 141$0I 25/0
Table I. Background Fluoride Levels In Reagents and Comparison of Found and Expected Concentrations titrant calcd [F-]/ [F-lfound/ Nernstian sample PPm PPm slope (a) TISAB (b) 1%NaCl ( c ) 1%NaCl +. TISAB ( 1 : l ) (d) 10 mL of (c) plus: (i) 0.050 mLlof 10 ppm F' (ii) 0.100 mL of 1 0 ppm F (iii) 0.050 mL of 100 ppm F(iv) 0.060 mL of 1000 ppm F a
10 10 10
0.024 0.044 0.034
59.3 59.0 59.2
10
0.083 (0.084) a 0.132 (0.133) a 0.536 (0.531) a
59.2 59.3
6.01
59.2
10
100 1000
58.4
(6.00) a
Expected value including the background [F-] in (c).
Table 11. Fluoride ]Levels in Surface Waters
sample Murray River (Swan Hill, Victoria) Albert Park Lalke (Melbourne) The Golden Spring (Hepburn Springs, Victoria)
titrant calcd [F-]/ [F-]found/ Nernstian ppm slope ppm 10 0.11 59.5 100
0.47
58.9
100
0.51
59.2
calculation. When the Nernstian slope is known, a linear relationship ( I ) or a Gran function (8) can be used to calculate the concentration. Our procedure takes advantage of both methods of calculation. To get an approximate fluoride concentration we use a linear relationship assuming the Nernstian slope equals Ei9 mV. This provides good initial guesses for the Brand and Rechnitz program. The resulting Nernstian slope is finally used in a Gram function calculation. We consistently obtain straight line regression coefficients better than 0.999 995 without rejecting arly data. In another approach, Frazer et al. (9) use an interactive computer to reject part of the data to select the "best" portions of the Gran plots for extrapolation. RESULTS AND DISCUSSION The fluoride ion selective electrodes used in this study were found to give more &able potential readings in nonstirred solutions. Indeed, in a stirred solution of 1%NaCl containing 2 ppm fluoride, the electrode response tended to drift even after an hour of continuous stirring. However, when the stirring was stopped, the electrode assumed a steady-state potential within about 30 min and the potentials measured a t 1-min intervals over the next 3 h were identical; in our ca9e the electrode gave am output of -72.80 mV with no variation even in the second decimal place. When stirring was reconi0 1981 Amerlcan Chemical Soclety
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ANALYTICAL CHEMISTRY, VOL. 53, NO. 13, NOVEMBER 1981
Table 111. Fluoride Levels in Urine Samples sample child A (duplicate analyses on 2 samples)
note a, c
(4)bIC
child B ( 3 samples)
a, c (10) b , c
(22) bsc
child B (1sample)
C
(50) (140) (254)
titrant [F-l/ppm
[F-]found/ ppm
calcd Nernstian slope
10 10 1000 1000 100 100 1000 100 1000 1000 1000
1.18 1.09 16.5 17.0 0.42 12.7 64.2 12.7 14.6 18.2 21.0
61.5 60.7 59.0 59.3 59.5 59.3 59.0 59.3 59.0 59.4 59.5
Number in parentheses ( n ) denotes urine sample was collected n hours after fluorocarbon anaesa Basal [F-] in urine. thesia, Analysis on freshly thawed sample. Number in parentheses denotes the number of hours sample stood at ambient temperature after thawing and prior to analysis. menced, the electrode potential immediately began to drift. These observations led us to adopt the procedure of stirring after each addition to mix the solutions (typically 2 min) followed by a nonstirring period in which the electrode potential was monitored. As has been noted by others (10-12) we also found that equilibration times decreased as the fluoride concentration increased and indeed for most of the analyses reported in this work it was clear from the electrode output that a steady-state potential was established within about 5 min except for samples containing the lowest levels of fluoride. With samples containing the lowest levels of fluoride, the general practice adopted was to keep the buret silicone rubber delivery tip out of the solution for the initial (no added fluoride) potential measurement. This was done in order to prevent any sample contamination which might arise as a result of diffusion of a trace of titrant from the tip. For subsequent potential measurements,the buret tip was lowered into the solution so that titrant could be added directly to the buffered sample. The results presented in Table I clearly indicate that good agreement is obtained between an expected fluoride concentration and the experimentally determined value using the standard additions method in a relatively simple matrix containing equal volumes of 1% NaCl and TISAE! spiked with various concentrations of standard fluoride solution. In addition, the Nernstian slope value calculated for each analysis using the Brand and Rechnitz program always compares favorably with the expected Nernstian slope of 59.16 mV at 25 "C (see Table I). This agreement with both the expected analytical concentrations and the Nernstian slopes is assumed to validate the analysis procedure we have adopted. Although we have no data with which to compare our surface water analyses, the values obtained (see Table 11) lie within the expected limits for uncontaminated environmental waters (13-15). In addition, it should be noted that the calculated Nernstian slopes are in close accord with the expected value. In the results obtained for the urine samples (see Table 111) it is again clear that the calculated Nernstian slopes are close to the expected value and so again our analytical procedure is validated. In addition, it should be pointed out that Tusl (16) also used TISAB with urine samples and determined inorganic fluoride using an ion-selective electrode, but the method used was to construct a standard calibration curve. The basal concentration of fluoride in urine is known to span the range of approximately 0.1-1.1 ppm, depending on the fluoride content and volume of fluid in the diet (17,18). The basal concentrations reported in our study fall in the expected range and appear to be consistent with values re-
ported elsewhere (19)for children imbibing fluoridated water (Melbourne,Australia, has been fluoridating its water supply to approximately 1 ppm since 1977). The grossly elevated levels of inorganic fluoride in urine presented in Table I11 arose because they were obtained from children who had undergone anaesthesiawith the fluorocarbon inhalation anaesthetic, methoxyfluorane. It is well-known that this anaesthetic gas is rapidly metabolized in humans to produce inorganic fluoride and other fluorine-containing species (20-23). In the present work, only short-term studies have been performed, but in all patients so far examined the concentration of fluoride in urine has steadily increased after exposure to the fluorocarbon anaesthetic reaching a maximum of about 60 ppm 24 h after use, similar to previously reported trends (20,21). The total volume of urine voided by the child for each sample was not determined and so it is not possible to calculate the total amount of fluoride excreted in the urine and compare it with the total inhalation of fluorine in the form of methoxyflurane. From the data in Table I11 for child B it is clear that the concentration of fluoride measured in a single urine sample gradually increases upon storage at ambient temperature. This is particularly noticeable in samples containing abnormally high concentrations of fluoride. It is suggested that the observed increase in inorganic fluoride concentration is due to the slow decomposition at ambient temperatures of some organo-fluorine species, the presence of which has been described (20-22). Thus, we recommend that urine samples be frozen as soon as possible after collection and that analyses be performed on freshly thawed samples. Only in this way will the measured concentration of inorganic fluoride truly represent the physiological level present in the urine.
ACKNOWLEDGMENT The urine samples were obtained from patients under R. Westhorpe of the Royal Childrens' Hospital, Melbourne, Victoria, Australia. LITERATURE CITED Rix, C. J.; Bond, A. M.; Smith, J. D. Anal. Chem. 1978, 48, 1236-1239. Manakova, L. I.; Bausova, N. V. Zavod. Lab. 1978, 42, 635-637; Chem. Abstr. 1977, 86, 3 4 0 7 7 ~ . Galleao. R.: Bernal. J. L.: Pardo. R. Aflnidad 1978, 35, 333-337; Cheni. Abstr. 1976, 90. 1743981. ANay, G. Gyogyszereszet 1977, 21, 171-173; Chem. Abstr. 1977, 87, 95326~. Pantucek, M. H. Cesk. Hyg. 1978, 23, 187-195; Chem. Abstr. 1978, 89, 86648q. Instruction Manual-Fluorlde Electrodes, Model 94-09, Model 94-06; Orion Research Inc.: Cambridge, MA 1977; pp 2-3. Brand, M. J. D.; Rechnitz, G. A. Anal. Chem. 1970, 42, 1172-1177. Gran, G. Analyst(London)1952, 77, 661-671. Frazer, J. W.; Kray, A. M.; Selig, W.; Lim, R. Anal. Chem. 1975, 47, 869-875.
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Anal. Chem. 1981, 53, 2143-2144
Baumann, E. W. Anal. Chim. Acta 1971, 54, 189-197. Durst, R. A. Ed. “Ion Selective Electrodes”; Department of Commerce, National Bureau of Standards: Washington, DC, 1969; NBS Spec. Publ. (US.)No. 314, pp 154-155. MkJgley, D.; Torrance, K. “Potentiometrlc Water Analysts”; Wlley: London, 1978, pp 318-319. Waldbott, G. L.; Burgstahier, A. W.; McKinney, H. L. ”Fluoridation: The Great Dilemma”, Coronado Press: Lawrence, KS, 1978; pp 35-36. Liberti, A.; Masclnl, M. Anal. Chem. 1969, 41, 676-679. Reference 11, p 361.
Tusl, J. Anal. Chem. 1972, 44, 1693-1694. Iriweck, K.;Sorantin, ti. Mikrochim. Acta 1977, 2 , 25-31. &em. Abstr. 1977 87, 149230d. Wassenaar, J. E.; Binnefts, W. 1. Voedins 1976, 39, 18-20. Chem. Absfv. 1978. 89. 194650. Reference 19, pp 51-62.
(20) Holaday, D. A.; Rudofsky, S.; Treuhaft, P. S. Anesthesiology 1970, 33, 579-593. (21) Mazze, R. I.; Trudell, J. R.; Cousins, M. J. Anesthesiology 1971, 35, 247-252. (22) Cousins, M. J.; Mazze, R. I. Anaesfhesia Intensive Care 1973, 1 , 355-373. (23) Yoshimura, N.; Holaday, D. A,; Fiserova-Bergerova, V. Anesthesiology 1976, 44, 372-379.
RECEIVED for review November 26, 1980. Resubmitted June 15,1981. Accepted July 13,1981. We are grateful to the Ro:yal Melbourne InstituteofTechnologyfor the award ofa grant.
Determination of Mercury in the Presence of Iron(1II) by Iodide Ion Selective Electrode Guler Somer Chemistry Department, Hacettepe University, Beytepe, Ankara, Turkey
It has been tried for a long time to develop an electrode for determining mercury. Since it was known that mercury was an interfering ion for the silver/sulfide electrode, we tried to use this electrode for the determination of mercury. However, the Ag/S electrode was not reversible, and even after brief exposure to mercury it became completely unresponsive. It was found later that an iodide electrode could be used for the determination of mercury and that this electrode responds to mercuric ion down to about loW8M (1). The mechanism appears to involve the reaction of mercuric ion with silver iodide on the electrode membrane surface to release silver, which is sensed by the electrode. Because the electrode actually senses silver, a monovalent ion, it exhibits a monovalent (60 mV) slope, in spite of”the fact that mercuric ion is divalent. Since mercuric ion readily forms chloride complexes, this ion should be removed before measurement by passing the sample through an anion exchange column (I). Overman, on the other hand, used potentiometric titration of mercury with sodium iodide using the same electrode as the sensor (2). As he also discusses, ferric ion, peroxide, and chelating agents are interfering ions and have to be eliminated. Therefore for the determination of mercury, especially in natural matrices in which considerable amounts of iron are present, interference of Fe3+should be eliminated. Recently some new electrodes which have a response to mercury have been suggested, one in liquid state (3) and some in solid state ( 4 , 5 ) but they are not commercially available. In this work the determination of mercury with an iodide ion selective electrode and the interference of iron(II1) is investigated. For this purpose a potentiometric titration method is applied with sodium iodide solution as the titrant. EXPERIMENTAL SECTION Apparatus. A Corning Model 12 Research pH meter with Orion (Model 94-53) solid-state iodide ion selective electrode and a Corning fiber junction saturated calomel electrode was used. No difference in results was observed when an Orion double junction reference electrode was used instead of the fiber junction electrode. Reagen1,s. The mercury@) solution was prepared by dissolving Hg(NO& in 0.1 N “0% This stock solution was diluted with triple distilled water and used in concentrations of W--10-* M. Merck pro analysis grade NaI was dissolved in triple distilled water in appropriate concentration with mercury solution and used as the titrant. 0003-2700/81/0353-2143$01.25/0
Table I. Change of Electrode Potential with Hg2+Concentration
10-4 10-3
210 21 2 24 5 29 5 338
140 179 247 302 358
160 175 235 290 349
Procedure. Solutions of mercuric ion (10-2-10-8M) are titrated with NaI using a microburet of 5-10 mL volume. Before ealch new experiment the bottom of the ion-selective electrode is cleaned with a specific polishing paper which is suggested by the Orion Co. Potential readings are taken while stirring with a magnetic stirrer.
RESULTS AND DISCUSSION The amount of mercury can theoretically be determined directly, by measuring the potential of a solution into which an iodide-selective electrode and a reference electrode are immersed. However it is observed that, especially at low concentrations of mercury, the proportionality between potential and the logarithm of the mercury concentration does not exist and therefore the results are not reliable. Tablo I summarizes the data obtained for the change of potential with Hg2+concentration. As it is observed, the change of potential with concentration if3 not proportional since theoretically for a 10-fold change of concentration a potential change of 59 mV is expected. In addition, potential measurements of identical solutions a t different times were not reproducible. They changed depending on the surface of the electrode. The results of three sets of experiments (EI,EII, EIII)which show the nonreproducibility are also given in Table I. According to Overman (2) mercuric ion down to 5 X 104 M can be determined by using a potentiometric titration. Because of the nonreproducible results as shown in Table I, it is preferable to use the NaI titration of mercury(I1) ions. In this method for each mercury(I1) ion, two iodide ions aire consumed. During the titration, mercury ions form Hg12,and after the equivalence point they form HgId2-complexes (I). Figure 1 shows the titration curve of 50 mL of M mercury(I1) solution with IO-$ M NaI. Because of a 200 mV change of potential about the end point, this titration can be 0 1981 American Chemical ,Society
