Lanthanum Fluoride Electrode Response in Aqueous Chloride Media R. E. Mesmer Reactor Chemistry Diaision, Oak Ridge National Laboratory, Oak Ridge, Tenn. 37830
WITHTHE RECENT INTRODUCTIONof solid-state electrodes, very sensitive analytical tools have become available for several anions. The best characterized of these is the lanthanum fluoride electrode doped with Eu2+( I , 2,3). Of particular interest to the analytical and physical chemists are the interferences which must be taken into account. The sensitivity of the fluoride electrode is at least 1000-fold for fluoride over chloride ( I , 3). Because equilibrium studies in aqueous solutions are commonly conducted in the presence of a relatively high concentration of inert electrolyte to control the ionic strength of the medium, the magnitude of the interferences become limiting factors. As a result of our interest in the use of chloride media at this laboratory we have examined in detail the chloride interference at 25 O C. EXPERIMENTAL The Orion Model 94 09 fluoride electrode was used in this study. The fluoride electrode behavior was examined at 25.00' =t0.02"C using the cells indicated below for measuring fluoride and hydrogen ion concentrations.
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llm NaCl '(1 - y)m NaCl LaF31 !Pt, Hz x m NaFl , y m HC1
I1
Ilm NaCl '(1 - y)m NaCll Pt, H ~ I IPt, H2. jxm NaF y m HC1 In both cells the electrode on the right is a reference or constant potential electrode, and the one on the left, the indicating or measuring electrode. The concentration of y in the reference compartment was to 10-3m. (Concentrations are in molal units.) The cell consists of two Teflon compartments separated by a Teflon liquid junction with a leak rate of 0.001 gram per hour per cm head of water. With this junction experimentally determined liquid junction potentials obey the Henderson equation within a few percent. Under the conditions of the experiments discussed here, the liquid junction potential was below 0.5 mV except in the more acidic HCI solutions. The liquid junction potential was calculated and taken into account in every case. The two electrodes on the left are in a stirred outer compartment in which titrations were conducted; the inner compartment contained the reference hydrogen electrode. Hydrogen purified by a Serfass-Hydrogen Purifier Model C-15D was presaturated with water vapor over lm NaCl before passage through the solutions. The system used to measure cell potentials consisted of a vibrating-reed electrometer and a precision potentiometer described previously (4). It was shown by experiment that neither the Teflon vessel nor the LaF3 electrode gave rise to measurable (lO-'m) fluoride in the most dilute solutions investigated during the time required for these experiments. Corrections were made for small potential drifts especially (1) M. S. Frant and J. W. Ross, Jr., Science, 154, 1553 (1966). 39,881 (1967). (2) J. J. Lingane, ANAL.CHEM., (3) G. A. Rechnitz, Chem. & Eng. News, 147 (1967). (4) C. F. Baes, Jr. and N. J. Meyer, Inorg. Chem., 1, 780 (1962).
I 10-61A0 140
I30
IIO
110
400
90
8;
70 mV
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50
4b
30
1 1 20 10 0
Figure 1. Changes in potential of the fluoride electrode with fluoride concentration in 1.00m chloride solutions prepared from three sources of chloride (The reference cell solution composition was different in each experiment). Open circles-experimental data ; closed circles-data corrected for assumed fluoride impurities of 6.5 X 10-5m (I), 1.5 X (II), and 6.0 X lo-% (111); solid curves calculated Nernst plots. in the most dilute fluoride solutions (a few tenths of a millivolt per hour to a few millivolts per hour) by the method described by Baes for the glass electrode (4). For experiments designed to observe any interference of chloride, 1 .OOm chloride solutions were prepared from three different sources of purified chloride: NaCl (Baker and Adamson Reagent Grade) recrystallized once from water, (I); NaCl prepared from Fisher Certified NaOH and Baker Analyzed Reagent HCI, (11), and KCI prepared as a large single crystal of very high purity at Oak Ridge National Laboratory, (111). RESULTS AND DISCUSSION The pH of 1.OOm solutions of chloride I, 11, and I11 was adjusted to 6-7 and dilute fluoride solutions in lm chloride were added incrementally to obtain the data in Figure 1. The solid lines in the figure represent the theoretical Nernst slope (59.16 mV) and the dashed lines the experimental data (open circles). The data show departure from the Nernst slope which can be due to several possible causes: (a) activity coefficient changes; (b) changes in liquid junction potential; (c) interferences; (d) presence of fluoride impurities in the purified chloride; and (e) fluoride from the electrode or vessel. However, because the ionic strength was maintained, constant activity coefficients are not expected to change appreciably, especially at the low levels of fluoride. The liquid junction potentials for these cells were shown to follow the Henderson equation in this medium and changes do not exceed 0.1 mV for the data of Figure 1. Likewise, it was shown that no measurable fluoride was contributed by LaF3 or the Teflon in the most dilute fluoride solution of Figure 1 . In the pH region 6-7 the interference due to hydroxide is negligible at these concentrations of fluoride. VOL 40, NO. 2, FEBRUARY 1968
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Table I. Calculated Dissociation Constant for HF From Results of Additions of HCI to l m NaCl Solution(1V)” IH+1 TF-1 [F-I 3.8 X IO-’ 9.62 X 5.58 5.20 X 5.58 2.29 X 3.29 X 1W6 8.49 X 1.46 X 10-6 3.47 X 5.56 4.10 X 8.05 X IW7 7.36 X 10-3 5.54 4.74 X l W E 3.90 x 10-7 1.67 x 10-2 5.39 5.09 x 10-6 1.45 X l e 7 4.80 X 5.31 5.16 X 8.79 x 5.08 5.00 X 8.06 X 1 0 - 8 Deviation from Nernst behavior equivalent to 5.58
1.32 1.22 1.24 1.25 1.28 1.35 1.41 X 10-6m.
Chloride interference and a fluoride impurity in the purified chloride cannot be distinguished by a single experiment. If the experimental data in Figure 1 are corrected for an assumed fluoride impurity in the chloride (solid circles), then Nernst behavior is obtained down to about 8 X 10-6m fluoride. The corrections which give straight lines with the Nernst slope are 6.5 X 10-5m, 1.5 X 10-5m, and 6.0 X lO-om, respectively, for I, 11, and 111. Because these corrections are different for the three 1.00m chloride solutions this clearly cannot be due to chloride. Chloride interference must depend on the chloride concentration which is constant in these three experiments. However, one must allow the possibility that the lowest observed correction, 6 X IO-Bm, represents the chloride interference. This possibility was eliminated by the following experiment which identified the cause of the deviation as a fluoride impurity. An HC1 solution was added incrementally to a 1 .OOm chloride solution(1V) for which the Nernst correction was 5.6 X lO-6m. Data for measured fluoride concentrations, [F-1, measured hydrogen ion concentration, [H+], and the total fluoride concentration, [FIT,are listed in Table I. The
[HFI concentration was obtained from the difference in [FIT and [F-1. The concentrations of the fluoride species in the table were calculated on the assumption that the Nernst deviation in a plot of the type shown in Figure 1 is due entirely to fluoride impurity. If the deviation were chloride interference, then no dependence on pH would be expected; however, the fluoride concentration will decrease with decreasing pH below pH of 4 where some H F is produced. The dissociation constant calculated from these data is in excellent agreement with published values (5) (the range of best values in similar media is 1.5 X to 1 . 1 X indicating that the deviation was indeed the result of a fluoride impurity in the chloride. From the data of Table I the upper limit for the chloride interference was set at 2 X 10-8m in view of the constancy of the equilibrium quotient. Also, this enables us to conclude that Nernst behavior (within 10%) is obtained to
