tion, the pH was adjusted to the yellow color of methyl red indicator, and a drop of pyridine was added. The mixture was back-titrated with standard 0.1M La(NO&, using either Arsenazo or Xylenol Orange as the indicator.
Table 11. Comparison of Fluoride and EDTA Titrations in Radioactive Process Solutions AKIII) found, M Q Method Fluoride EDTA 1.47, 1.47 1.48 1.87, 1.87 1.88 1.73, 1.73 1S6, 1.73 1.90 1.88 1.85, 2.06 2.03 2.08 2.08
RESULTS AND DISCUSSION
A typical potentiometric titration curve is shown in Figure 1. A sudden drop in potential on the first addition of fluoride is followed by a plateau until the equivalence point is approached. The potential change at the equivalence point is >20 mV for a 20941 addition of 0.1MNaF. Table I compares determinations by the fluoride and EDTA procedures in a solution of aluminum nitrate and in a simulated Tramex process solution. The results represent the average of four or more determinations. Bias between the two methods is within the precision of the determination. The relative standard deviation is 1 to 2 % for both methods at 10 micromoles of Al(II1). Table I1 compares determinations in Tramex process solutions. General agreement between the methods is obtained for the six samples. The presence of radioactivity does not affect the fluoride titration. Successful potentiometric titration of aluminum with fluoride depends upon removal of other ions that form complexes or insoluble compounds with fluoride. In the present application, where aluminum is a major constituent, NaOH precipitation is practical because the aluminum occluded on the metal hydroxide precipitate is negligible. However, other mixtures may require different purification procedures to avoid losses of aluminum. The method is based on the postulated formation of an aluminum hexafluoride that is insoluble or undissociated in
a
Original solution diluted 1 :100 in 1M NaOH; 500-1000
p1
of supernate titrated.
ethanol. The slight solubility of cryolite in water is reduced in ethanol. In aqueous solution, Tananaev and Lel‘chuk (5) found that the solid phase consists of two double salts: A1F3.3NaF and 4A1F3.11NaF. In both of these salts, the F/Al ratio is very close to six. The agreement between the fluoride and the EDTA methods indicates that this value can be used for the equivalence factor. RECEIVED for review July 10, 1969. Accepted October 20, 1969. 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. (5) I. V. Tananaev and Yu. L. Lel’chuk, Zhur. Anal. Kliim., 2, 93 (1947); C.A., 43, 5695c.
Quinhydrone Electrode Drift George Dahlgren and Melvin J. Goodfriend, Jr. Department of Chemistry, Unioersity of Cincinnati, Cincinnati, Ohio 45221 REPORTS of a slow decay in the potential of quinhydrone OS. silver-silver chloride electrodes are as old as the cell system itself (]). Literature references usually ascribe this emf drift to chlorination of the p-benzoquinone component of quinhydrone ( 2 3 ) . However, others either report no problem with drift (4) or report drift only in cells containing certain buffer ions (5). We have observed a linear emf decay with time in cells containing acetate buffers. Possible sources of the drift such as poor electrode behavior or diffusion of quinhydrone between the half cells were eliminated and the actual source of the problem appears to be a nucleophilic attack by acetate ion on the p-benzoquinone component of quinhydrone. The reaction was followed by changes in the UV spectrum of the p-benzoquinone with time and this rate correlated with changes in the emf of the cells.
EXPERIMENTAL
(1) D. J. G. Ives and G. J. Janz, “Reference Electrodes,” Academic Press, New York, N. Y . , 1961. (2) H. S. Harned and D. D. Wright, J. Amer. Chem. Soc., 55,
Materials. Materials were the best grade available and were used without further purification except in the case of quinhydrone. Several different recrystallization procedures were used with quinhydrone and all gave identical results. Solutions were prepared from deionized-distilled water which had been freed of dissolved carbon dioxide and oxygen by flushing with pure nitrogen. Apparatus. A semi-micro cell similar to the one used in previous studies ( 4 ) was constructed from borosilicate glass. The silver-silver chloride electrodes were of the thermal type described by Bates (6). However, the finished helices were inserted in rubber plugs which were then sealed in 10/30 borosilicate points with silicone rubber. Electrodes prepared in this way had longer life because the channeling in platinumglass seals observed by many investigators (2) could not occur. Prior to each emf run sets of like electrodes were intercompared and were usually found to be within 10.03 mV of each other, even after two months use. Potentials were measured using a Leeds and Northrup K-3 potentiometer and Model 9834 null detector.
4849 (1933). (3) H. S. Hovorka and W. C. Dearing, ibid., 57, 446 (1935). (4) G. Dahlgren and F. A. Long, ibid., 82, 1304 (1960). ( 5 ) K. C. Rule and V. K. LaMer, ibid.,60, 1974 (1938).
(6) R. G. Bates, “Determination of pH, Theory and Practice,” John Wiley and Sons, Inc., New York, N. Y . , 1964, p 281.
~
ANALYTICAL CHEMISTRY, VOL. 42, NO. 1, JANUARY 1970
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Table I. Absorbance and emfchanges with Time for the Acetate-Quinhydrone System at 25 "C
[HOAc] 0.01246 0.02397 0.02378 0.03564 0.03780 0.04800 0,05993 0.05870
Buffer concentrations, m [NaOAc] 0.01168 0.02427 0.02435 0.03643 0.04605 0.04834 0.06108 0.06118
[NaCI] 0.04575 0.02041 0.05621 0.03090 0.03904 0.04101 0.05181 0.05193
UV spectra were taken on a Cary 11 spectrophotometer. The spectrum of quinhydrone is a composite of those for pbenzoquinone and hydroquinone. Both substances obey Beer's law. Freshly prepared solutions of p-benzoquinone a t 246 nm and log a = 4.338 [compared to log gave a, ,A, a = 4.51 at 244 nm in (7) and compared to log a = 4.2 at 250 nm in (S)]. Similar solutions of hydroquinone gave , , ,A at 289 nm with log a = 3.408 and at 222 nm with log a = 3.696 [compared to log a = 3.49 at 294 nm in (9)]. At 246 nm log a for hydroquinone is 2.279. Run Determinations. Buffer solutions were prepared by weight from previously standardized stock solutions of acetic acid, sodium hydroxide, and sodium chloride. The solutions were divided into two portions, one for the determination of the emf and the other for spectrophotometric studies. After preparation, the emf cells were immersed in a water bath at 25.0 f 0.05 "C and the potentials of duplicate silver-silver chloride and duplicate quinhydrone electrodes were measured. Readings were taken at about 15-minute intervals for two to three hours after thermal equilibrium was established, usually 30 minutes after immersion. After approximately three hours the diffusion of quinhydrone into the silver-silver chloride compartment affected the readings. The stopcock separating the two compartments was opened only when readings were taken. The second portion of the buffer solution was saturated with freshly recrystallized quinhydrone and placed in the thermostated bath. UV spectra of these solutions were taken a t regular intervals. RESULTS AND DISCUSSION
We had constructed our emf system to study ionization constants of a certain class of organic acids. It was during our equipment shakedown runs, determining the pKa of acetic acid, that we became aware of a regular decrease in emf with time, In order t o determine whether quinhydrone diffusion into the silver-silver chloride half cell was the cause of the drift, the potentials of duplicate cells were measured. One set of cells was measured in the usual way, that is by momentarily opening the stopcock separating the half cells every 15 minutes for about two hours, while the stopcocks on the other set of cells remained unopened for the same period, thereby preventing quinhydrone diffusion. The potentials of these matched sets were identical after two hours. As a n example of the accuracy of the measurements, extrapolation of the emf back to zero time for runs at an ionic strength of 0.072 and 25 "C gave a pKa for acetic acid of 4.764, in good agree(7) R. A. Freidel and M. Orchin, "Ultraviolet Spectra of Organic Compounds," J. Wiley and Sons, Inc., New York, N. Y.,1951, p 77. (8) L. E. Orgel, Trans. Faraday Soc., 52, 1172 (1956). (9) R. A. Morton and A. L. Stubbs, J. Chem. SOC.,1940,1348. 112
-
103 (hr-1) 1.48 2.18 2.23 2.75 3.36 3.48 4.24 4.28
6bsd. 0.08
0.16 0.10 0.19 0.22 0.21
0.24 0.28
($)T
(mV hr-l) Equation 4 0.09 0.13 0.13 0.16 0.20 0.21 0.25
0.25
ment with the value of 4.760 of Harned and Ehlers (10) The drift in these runs was about 0.2 mV hr-l. Because the UV spectra of acetate buffer solutions (with sodium chloride added) containing quinhydrone showed a decrease in the p-benzoquinone peak at 246 nm with time, other aspects of the emf drift problem were investigated spectrophotometrically. To determine the effect of chloride ion on the emf" decay the concentration of chloride ion in run solutions was varied from zero to 0.04m and the changes in absorbance with time at 246 nm were observed. The decrease in absorbance was independent of the chloride ion concentration, thereby ruling out chlorination as a cause. Importantly, the decrease at 246 nm was accompanied by concurrent increases a t 289 nm and 222 nm, indicating that p-benzoquinone was being converted to hydroquinone or more probably to a substituted hydroquinone. We were not able to isolate a new product because of the attending polymerization reactions under the conditions employed. Companion experiments in both the light and dark showed that the observed spectral changes were light independent. The effect of pH on the rate of decrease in the absorbance at 246 nm was determined by measuring changes in absorbance with time of buffered solutions. Solutions were prepared in the pH range 3.7 to 5.3. which contained identical concentrations of acetate ion and varying concentrations of acetic acid. N o differences in the decrease in absorbance with time were observed. p-Benzoquinone, when used in place of quinhydrone, showed a 25% slower decrease in absorbance than when quinhydrone itself was used. N o explanation for this apparent catalysis by hydroquinone is offered. However, further work on the details of the mechanism of the reaction is contemplated. The reaction responsible for the decrease in the absorbance at 246 nm with time was found to be dependent on the acetate ion concentration as well as on the quinhydrone concentration, to be virtually independent of chloride ion concentration, and to show what appears to be a water or neutral reaction. There remained the question of relating the absorbance decrease to the emf decay. For the cell Pt IQHz,Q, HOAc, NaOAc, NaCl IAgCl, Ag the potential is given by
(IO) H. S. Harned and R. W. Ehlers, J. Amer. Chem. SOC.,54,
ANALYTICAL CHEMISTRY, VOL. 42, NO. 1, JANUARY 1970
1355 (1932).
or
where aQ and aQ& represent the activities of the p-benzoquinone and hydroquinone components of quinhydrone. The other symbols have their usual meaning. In typical applications of the quinhydrone electrode the activities of the quinhydrone components are considered equal and the last term of Equation 2 is dropped. It is this final term which is the link between the absorbance or quinhydrone concentration decrease and the emf decay. Because the activities of hydrogen ion and chloride ion should remain essentially constant, the change in the measured potential with time at constant temperature can be written as
If we assume ideal behavior and Beer's law, and further assume that the decrease in the concentration ofp-benzoquinone is accompanied by a n increase in the concentration of hydroquinone (or a substituted hydroquinone), Equation 3 becomes
where AQ is the absorbance of p-benzoquinone. Equation 4 explicitly relates changes in the potential of the cell to changes in the absorbance ofp-benzoquinone. Plots of log AQ at 246 nm us. time gave good straight lines as did plots of emf us. time. Data from runs covering a range of acetate ion concentrations can be found in Table I. Column 4 is the slope of the absorbance plots while column 5 contains the slopes of the emf plots. Column 6 is the change in emf with time calculated from Equation 4 using column 4 absorbance data. The agreement is reasonable. It should be noted that the results obtained in this work do not preclude the use of the quinhydrone electrode. When a high degree of accuracy is desired and significant emf drift is encountered, an extrapolation back to zero time can and should be made. As an example, at 25 OC acetate ion gives a drift rate of 0.2 mV h r ' which results in an error of 0.004 pKa unit if one hour data is used without extrapolation or 0.002 pKa unit for 30 minute data should thermal equilibrium be established in that time. We do not expect that most organic anions would be as active as acetate ion. Our data on bimaleate, maleate, and on the anions of ethylene tetracarboxylic acid indicate a drift rate comparable to the water or neutral reaction.
RECEIVED for review September 11, 1969. Accepted October 9, 1969. The authors gratefully acknowledge the financial support of the American Cancer Society Grant No. P-272.
Determination of Vanadium in Natural Waters by Neutron Activation Analysis K. Daniel Linstedt' and Paul Kruger Civil Engineering Department, Stanford University, Stanford, Calv. VANADIUM HAS BEEN quantitatively determined by neutron activation analysis in a variety of sample matrices. For example, Brooksbank, Leddicotte, and Mahlman (I) determined the vanadium concentration in crude oil using a method for nondestructive comparator analysis. Kemp and Smales (2) applied activation analysis for determination of the vanadium concentration in rocks and meteorites. Their procedure included a sodium peroxide fusion followed by separation of vanadium from interfering radionuclides by a cupferronchloroform solvent extraction. Lukens, Heydorn, and Choy (3) developed a pre-irradiation concentration procedure for activation analysis for vanadium in blood. This selective concentration of vanadium permitted the post-irradiation radiation measurement to be made by purely instrumental methods. Grimanis, Pantazis, Papadopoulos, and Tsanos (4determined
vanadium and some other trace elements in several Greek Lakes. For the vanadium analyses, 100-ml samples were evaporated to provide pre-irradiation concentration. Following irradiation the vanadium was isolated by a 3.5-minute cupferron-chloroform solvent extraction prior to counting. In the present work, the occurrence and behavior of vanadium in several natural waters has been investigated as part of a study of the fate of this element in domestic supply waters. A generally applicable procedure has been developed for quantitative analyses of dissolved vanadium in natural waters. This procedure includes a rapid pre-irradiation concentration to enhance the sensitivity of analysis with low flux reactors, and a post-irradiation radio-chemical separation to reduce the interference of other radionuclides.
Present address, Civil Engineering Department, University of Colorado, Boulder, Colo. 80302
Reagents. Ammoniacal Buffer Solution, pH 9.4. Dilute 100 ml of 4N ammonium hydroxide and 50 ml of 4N nitric acid to 1 liter with distilled water. Oxine-Chloroform Solution, 1%. Dissolve 10 grams of reagent-grade oxine in 1 liter of chloroform. Procedure. Samples of natural water are filtered through acid washed, hard finish, Whatman No. 42 filter paper to isolate the suspended particulate matter. The filtrate is concentrated prior to activation by passing 1-liter aliquots through a column of hydrogen-form Dowex 50W-X8 cation exchange resin. This type of resin has been shown to quan-
1
(1) W. A. Brooksbank, G. W. Leddicotte, and H. A. Mahlman, J. Phys. Chem., 57, 815 (1953). (2) D. M. Kemp and A. A. Smales, Anal. Chim. Acta., 23, 397 (1960). (3) H.R.Lukens, K. Heydorn, and T. Choy, Trans. Amer. Nuclear Soc., 8, 331 (1905). (4) A. P. Grimanis, G. Pantazis, C. Papadopoulos, and N. Tsanos, Third U.N. Intern. Conf Peaceful Uses A t . Energy, 854 (1964).
EXPERIMENTAL
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