trode by products of the electroreductions to yield (C&)zHg in both cases. The transmetallation of (C6H5)2T1Brwith mercury has been observed in a nonelectrochemical system by Gilman and Jones (18). They reported an excellent yield (90%) of (C6Hs)2Hgafter refluxing in pyridine for eight hours. The polarographic results indicate that the intermediate formed during the second electron transfer (Equation 3) is an adsorbed phenylthallium species, C6HsT1, and not a phenylmercury radical, C6HsHg* . In the polarographic reduction of phenylmercury ions, both the formation of C6H5Hg. and its subsequent reduction to C6H6and Hg are irreversible (13, 19). Because the polarographic reduction to C6Hs and T1" (Equation 5) is reversible, it is an adsorbed C6HsT1 species, not C6HsHg., which undergoes rapid disproportionation to
(C6H&Hg (Equation 4) and reduction to C6H6and TI (Equation 5). Gilman and Jones (20) presented evidence for the formation of phenylthallium by pyrolysis of (C6Hs)3TI in xylene. An attempted electron paramagnetic resonance study to determine the nature of the intermediates formed during electrochemical reduction was unsuccessful. The formation of an insulating layer of (C6H5)2Hgprecipitate on the mercury electrode surface contained in an EPR aqueous flat cell inhibited the electrochemical reduction after a few seconds and no signals were observed. Stirring the mercury layer to expose fresh electrode surface to the depolarizer is not possible in such flat cells. It is possible that EPR studies of the reaction intermediates might be more successful in solvents in which diphenylmercury is soluble.
(18) H. Gilman and R. G. Jones, J. h e r . Chem. SOC.,61, 1513 (1939). (19) R. F. Broman and R. W. Murray, ANAL.CHEM.,37, 1408 (1965). (20) H. Gilman and R. G. Jones, J . Amer. Chem. SOC.,62, 2357
RECEIVED for review February 29, 1968. Accepted May 2, 1968. Presented in part, Division of Fuel Chemistry, 153rd ACS Meeting, Miami Beach, Fla., April, 1967. Work supported in part by National Science Foundation Grant GP6164.
(1940).
Potentiometric Determination of Boron as Tetrafluoroborate R. M. Carlson Department of Pomology, Unioersity of California, Dacis, Calif. 95616
J. L. Paul Department of Environmental Horticulture, Unicersity of California, Davis, Calq. 95616 The potentiometric determination of tetrafluoroborate with a liquid ion exchange membrane electrode i s described. The electrode can be used for tetraInterfluoroborate concentrations down to lO--6M. ference from several anions was estimated. Nitrate and iodide cause the greatest interference of those anions studied. Rate of formation of tetrafluoroborate in solution at 24 and 60 OC was studied. Complete formation of tetrafluoroborate was obtained in 0.28M hydrofluoric acid at 60 OC in 5 minutes. At lower temperatures and lower hydrofluoric acid concentrations more than 6 hours were required for complete formation. Using columns of the boron specific resin Amberlite XE-243 boron can be separated from interfering anions and the tetrafluoroborate ion can be formed in less than 15 minutes with a small volume of 10% hydrofluoric acid. The tetrafluoroborate is eluted with sodium hydroxide and measured potentiometrically. Application of the method to determination of boron in water samples is illustrated.
THE RECENT DEVELOPMENT of liquid ion exchange membrane electrodes that are selective for nitrate and perchlorate ions (1) suggests the possibility of a similar electrode that would respond to the tetrafluoroborate ion. The behavior of the three ions is similar in many ways. All three form sparingly soluble salts with nitron. Perchlorate and nitrate extract as tetrabutylammonium complexes into methyl isobutyl ketone along with tetrafluoroborate (2). It seemed plausible that the type of membrane used for the perchlorate and nitrate electrodes would be suitable for tetrafluoroborate. (1) G. A. Rechnitz, Chem. Eng. News,45 (25), 146 (1967). (2) W. J. Maeck, M. E. Kussy, B. E. Ginther, G. V. Wheeler, and J. E. Rein, ANAL.CHEM., 35, 62 (1963).
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ANALYTICAL CHEMISTRY
Current methods for determining boron are time-consuming and subject to interferences. A rapid potentiometric method sensitive to microgram quantities of boron would have application for boron determinations in many materials. This report deals with a liquid ion exchange membrane electrode sensitive to tetrafluoroborate ion. A procedure for isolating boron and rapidly forming the tetrafluoroborate ion on columns of the boron specific resin Amberlite XE-243 is described. EXPERIMENTAL. Apparatus. EMF measurements were made with an Orion Model 801 digital pH meter. All solutions were stirred during measurement of potentials. As the work was performed in a laboratory with a temperature of 24 f 0.5 "C, no further effort was made to control temperature. The tetrafluoroborate-sensitive electrode was prepared by modifying a nitrate-sensitive electrode manufactured by Orion Research, Inc. (Model 92-07). The liquid ion exchanger was converted to the tetrafluoroborate form by twice shaking 2 ml of exchanger with 50 ml of a molar solution of boric acid in 4 molar hydrofluoric acid and decanting the aqueous phase. The composition of the internal filling solution was 10-*M H3B03, 0.28M HF, and 10-*M KCI. The electrode was assembled according to manufacturer's directions. Electrode Sleeve. Because all solutions contained HF, a plastic sleeve was used to protect the saturated calomel reference electrode from attack. The sleeve was prepared by gently heating a '/*-inch i.d. polyethylene tube and drawing out a constriction. A single strand of size 50 mercerized cotton thread was passed through the constriction and sealed in place by reheating and twisting the tubing. The twisted
Table I. Effect of Some Anions on Electrode Response to 4.46 X 10-4M Boron in 0.028M HF Salt added Conc, X 102M EMF, mV none 203.6
+I20 +I60 -
KNOa
> E
+zoo +240
--/
2 2 2 2 1
185.1 207.9 207.3 154.5 207.5
+200 5 c .
10-4 10-3 10-2 BORON CONCENTRATION (MOLES /LITER)
10-5
10-1
Figure 1. Electrode response to tetrafluoroborate Initial HF concentration : upper curve, 0.28M; lower curve, 0.028M
constriction was then cut to expose the cotton thread wick. The sleeve was filled with sufficient saturated KC1 to make contact with an inserted calomel electrode. A plastic sleeve type reference electrode (Orion Research Model 90-01) is now commercially available. This electrode could be used in place of the sleeve arrangement described above. However, the equitransferent filling solution supplied with the electrode should not be used because it contains substantial quantities of nitrate ion. A potassium chloride solution saturated with silver ions should be used as the filling solution. Resin Column Preparation. Polyethylene tubing, 3/1e-in~h ID, was used to prepare columns of Amberlite XE-243 boron specific resin. The tip of the tubing was drawn out and a small cotton wool plug inserted to support the resin bed. Approximately 1 ml of 40- to 80-mesh ground resin was placed in the tube, and a small cotton plug was placed on top to prevent entry of air. The height of the resin bed was 5 cm. An inverted 2-ounce polyethylene bottle with the bottom removed was fitted to the tube to serve as a reservoir. The flow rate over these columns was approximately 1 ml per minute. Cation exchange resin columns were prepared in a similar manner using 1/4-inch i.d. polyethylene tubing. These columns were filled to a depth of 10 cm with Dowex 50W-X8, 50- to 100-mesh resin. Prior to each use they were regenerated with 10 ml of 3M HC1 followed by 20 ml of HzO. Plastic laboratory ware and reagent grade chemicals were used throughout. RESULTS AND DISCUSSION
Electrode Response. The response of the electrode to tetrafluoroborate was determined with mixtures of boric and hydrofluoric acids. These solutions were prepared several weeks prior to measurement t o ensure complete formation of tetrafluoroborate. Figure 1 shows the response of the electrode as a function of boron concentration a t two initial concentrations of hydrofluoric acid. The slope of the linear portion of each curve is 58 mV per decade concentration change which approximates Nernst response. The curvature at the higher boron concentrations is due t o the consumption of hydrofluoric acid by the formation of tetrafluoroborate. Response is masked by excess hydrofluoric acid at the lower concentrations of boron. The response curve is useful down to lO-4M boron in 0.28M HF and to lO-5M boron in 0.028M HF. However, response at the low boron levels depends upon adequate stirring. The response of the electrode t o tetrafluoroborate in the presence of various anions was determined by adding potas-
sium salts to a solution of 5 x lOP4Mboron in 0.028M HF. The concentration of various added salts and the resulting measured potentials are presented in Table 1. There is considerable interference from iodide and nitrate as would be expected from a consideration of nitrate electrode behavior. The other anions tended to cause a small negative error. This negative error was probably caused by decreased activity coefficients owing to increased ionic strength. The selectivity constants of fluoride and bicarbonate ions were estimated by measuring the electrode potentials of decreasing concentrations of potassium tetrafluoroborate in 0.1M ammonium fluoride and 5 X 10-aMsodium bicarbonate. The selectivity constant for fluoride was 9 X 10-5 and that of bicarbonate 2 X The effect of pH was determined by adding hydrochloric acid or sodium hydroxide t o lop2 and 10-3M solutions of potassium tetrafluoroborate. Varying the pH from 3 to 11 had no effect on the electrode potential. Formation of Tetrafluoroborate in Solution. A serious limitation in the application of the electrode is the time required for the formation of tetrafluoroborate in the dilute hydrofluoric acid solutions that must be employed t o avoid interference from excess fluoride. Wamser (3, 4) and Pasztor, Bode, and Fernando (5) have studied the kinetics of formation and hydrolysis of tetrafluoroborate a t room ternperature. Pasztor, Bode, and Fernando found formation complete in 2 hours for a solution containing 0.2 ppm boron in 0.56M HF. With 0.056M HF, 18 hours was required to form tetrafluoroborate. One means to shorten the time required for complete formation of tetrafluoroborate is to carry out the reaction at elevated temperature. The effect of temperature on formation rate was determined by allowing boric acid to react with hydrofluoric acid at 24 and 60 "C. The solutions used were 1 X 10-*M H3B03in 0.28M H F and 1 X 10-3M in 0.028M HF. The 60 "C data were obtained by bringing H3BOa and H F solutions to 60 "C, mixing appropriate volumes to give the above concentrations, and maintaining the solutions at 60 "C for various times. The solutions were cooled in an ice bath to 24 "C and the tetrafluoroborate was determined potentiometrically. The results are presented in Figure 2. Elevating the temperature to 60 "C does indeed increase the reaction rate. For solutions that contain sufficient boron to allow use of 0.28M HF, the complex can be formed in 15 minutes at 60 "C. However, for more dilute solutions of boron, considerably more than 6 hours is required for the formation of tetrafluoroborate even at 60 "C. This seriously (3) C. A. Wamser, J . Amer. Chem. SOC.,70, 1209 (1948). (4) Zbid.,73, 409 (1951). (5) L. Pasztor, J. D. Bode, and Q. Fernando, ANAL.CHEM.,32, 277 (1960). VOL. 40,
NO. 8, JULY 1968
1293
t
I
I
I
I
I
I
1
100 2
80
0
c
80
a 60
I
1,080 pg BORON
E
e 40
70
20
60
s
0
0
I20 240 TIME (MINUTES)
360
-
50
8
90
Figure 2. Formation of tetrafluoroborate as a function of time, temperature, and H F concentration
limits the usefulness of the method for boron solutions more dilute than 10-4M boron. Use of Boron Specific Resin. The slow formation of tetrafluoroborate in dilute hydrofluoric acid and the interference encountered from anions such as nitrate prompted the investigation of the boron specific resin Amberlite XE-243 for isolating and concentrating boron from aqueous solutions. The properties of this resin are described by Kunin and Preuss (6). The active group in the resin is N-rnethylglucamine which complexes boron. The resin has a remarkable ability to remove boron from electrolyte solutions. The resin also contains tertiary amine groups which give it weak base anion exchange properties. These two properties of the resin offer a unique means of isolating and concentrating boron and forming tetrafluoroborate with a small volume of concentrated HF. Boron is removed from an aqueous sample passed over the resin, Interfering anions can be removed by washing with dilute ammonium hydroxide. The boron on the resin can then be reacted with a small volume of strong hydrofluoric acid. As the tetrafluoroborate is formed, it is adsorbed by the anion exchange complex of the resin. Excess hydrofluoric acid can then be removed by washing with water without loss of tetrafluoroborate because of its strong acid character. In this step, some of the adsorbed fluoride hydrolyzes because of the weak acid properties of hydrofluoric acid. This further decreases the concentration of excess hydrofluoric acid in the final solution. Finally, the tetrafluoroborate can be eluted with sodium hydroxide. If the eluate is passed directly over a cation exchanger in the hydrogen form, the resulting solution is a mixture of tetrafluoroboric acid in dilute hydrofluoric acid. Such a solution is ideal for the potentiometric determination of tetrafluoroborate. The breakthrough capacity of the 1-rnl columns described above was determined by passing increasing volumes of 5 x 10-3M H3B03 over a series of columns. Hydrofluoric acid was added to the effluent and, after sufficient time for formation, the tetrafluoroborate was determined potentiometrically. No boron could be detected in the effluent until more than 2.25 x 10-l mmol of boron had been added to the column. With the addition of 2.5 X 10-I mmol approximately 0.1% of the added boron appeared in the effluent. The breakthrough capacity was taken to be 2.25 X IO-' (6) R. Kunin and A. F. Preuss, Ind. Eng. Chem. Prod. Res. Develop., 3, 304 (1964).
1294
0
ANALYTICAL CHEMISTRY
80 70 60
50
0
IO 20 TIME (MINUTES)
30
Figure 3. Formation of tetrafluoroborate in Amberlite
XE-243columns at three H F concentrations mmol of boron, which is approximately one half of the theoretical capacity given by Kunin and Preuss (6). The batch equilibrium data of Kunin and Preuss indicate similar breakthrough capacities can be expected in the presence of electrolytes except for strong acid solutions where boron is not so strongly adsorbed. The rate of formation of tetrafluoroborate in the columns was determined for two levels of boron and three concentrations of hydrofluoric acid. Columns were treated with 1 ml of 5 , 10, or 25% H F after application of boric acid. At intervals after application of hydrofluoric acid, the columns were washed with 2 ml of water to remove the excess hydrofluoric acid. Immediately after the water wash, 10 ml of 0.3M NaOH were applied to elute the tetrafluoroborate. The eluate was passed directly over cation exchange columns to remove excess sodium hydroxide and to convert tetrafluoroborate to the acid. These eluates were diluted to 25 ml. The results are presented in Figure 3. Formation of tetrafluoroborate was complete in 5 minutes with both 0.1 and 2.5 X mmol of boron when 10 or 2 5 % H F was used. With 5 % H F 20 to 30 minutes are required to form the complex. More time is required to form the complex for the lesser amount of boron when 5 % H F is used. Using this procedure, the operating capacity of the column was 2 . 0 x 10-l mmol of boron which is about 90% of the breakthrough capacity. When the columns are loaded to the breakthrough capacity, addition of hydrofluoric acid displaces some boron. As the amount of fluoride in the final solution affects the sensitivity of the electrode at the lowest levels of boron, the
+I60
-
+I80
-
mV +200
-
+220
-
2
5
IO
5 0 100 pg BORON / 25 ml
500
Figure 4. Calibration curve for procedure employing Amberlite XE-243columns retention of fluoride by the columns against water washing was determined. One milliliter of 25 HF was added to each of five columns. The columns were then washed with 1, 2, 3, 5, and 10 ml of HzO. Fluoride was displaced with 0.3M NaOH and converted to hydrofluoric acid with the cation exchange columns. These solutions were diluted to 25 ml, and neutralized, and fluoride was determined with a fluoride electrode (Orion Model 94-09). The concentrations of
fluoride found in the eluate diluted to 25 ml were 0.16, 4.6 and 3.1 X 10+ molar for 4.3 X lo+, 3.4 X X water washes of 1, 2, 3, 5 and 10 ml, respectively. It is possible, then, to reduce effectively the excess fluoride in the final solution. For best results with low boron samples, the volume of water wash and the amount of resin per column should be carefully controlled. Boron was determined in a sample of tap water to demonstrate a possible application of the method. The procedure used was to pass aliquots of the sample over the Amberlite XE-243 columns, apply 1 ml of 10% HF and, after 10 minutes, 2 ml of water. The tetrafluoroborate was eluted with 10 ml of 0.3M NaOH directly onto a cation exchange column. The effluent from the cation exchange column was diluted to 25 ml prior to potentiometric determination of tetrafluoroborate. A calibration curve was prepared by carrying boric acid standards containing from 2.6 to 520 pg of boron through the same procedure, The calibration curve is shown in Figure 4. Results for the water sample were 0.67 and 0.65 ppm boron for each of two 50-ml aliquots and 0.66 and 0.65 ppm for each of two 25-ml aliquots. The curcurmin method gives a value of 0.66 ppm for this sample indicating excellent agreement between the methods. With adjustment of boron specific resin column size and reagent volumes, this approach to boron determination should find application to a wide variety of materials. RECEIVED for review February 28, 1968. Accepted April 22, 1968.
Nitrogen and Oxygen Compound Types in Petroleum A General Separation Scheme L. R. Snyder and B. E. Buell Union Oil Co. of California, Union Research Center, Brea, Calif. An integrated scheme has been developed for the separation of the oxygen and nitrogen compound types present in high boiling petroleum fractions into reasonably clean cut fractions for further analysis by other techniques. Cation and anion exchange are used in conjunction with adsorption chromatography on alumina, silica, and charcoal. The resulting separations are experimentally convenient, repeatable, and predictable, and sample recoveries are satisfactory. Sample reactions during separation appear to be generally minor and do not interfere with sample analysis. The use of this separation scheme-in whole or in part -by different workers should also serve to interrelate their findings on the analysis of different crude oil fractions.
THECOMPLEXITY OF PETROLEUM is such that the complete identification and quantitative determination of the various nitrogen and oxygen compound types (referred to here as heterocompounds) requires extensive separation of the starting sample. Previous attempts at the analysis of these heterocompounds in petroleum distillates and related materialse.g., I-5-have employed a wide variety of separation methods: acid and base extraction, ion exchange, adsorption chromatography, paper chromatography, gas chromatography, complex formation, liquid thermal diffusion, and other techniques. Each of these separation methods has potential
advantages and limitations in the analysis of petroleum heterocompounds, but in the past little attention has been given to the evaluation and optimization of such procedures in terms of actual petroleum separations. Similarly little work has been reported on their optimum integration into an overall separation scheme. Ideally an integrated separation scheme for the analysis of petroleum heterocompounds should meet several requirements. First, the various heterocompound types should be concentrated into a reasonably small number of distinct fractions, each of which contains only a few heterocompound types. In particular it is necessary that most of the heterocompounds be separated from the hydrocarbons and sulfur compounds which constitute the bulk of petroleum distillates. Second, the separation should be both repeatable and predictable. By this we mean that the approximate distribution of ~~
(1) G. U. Dinneen, G. L. Cook, and H. B. Jensen, ANAL.CHEM., 30, 2026 (1958). (2) G. K. Hartung and D. M. Jewell, Anal. Chim. Acta, 27, 219 ( 1962). (3) D. M. Jewell, J. P. Yevich, and R. E. Snyder, Am. SOC.Testing Mater., Spec. Tech. Pub/., 389, 363-383 (1965). (4) H . V . Drushel and A. L. Sommers, ANAL.CHEM., 38, 19 (1966). ( 5 ) D. K. Albert, ibid., 39, 1113 (1967). VOL 40, NO.
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JULY 1968
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