RESULTS AND DISCUSSION Previous work on the determination of mercury in biological materials culminated with a volatile element separation of mercury (23). Because of the volatile nature of selenium compounds, the method was investigated to see if the simultaneous collection and determination of selenium and mercury could be accomplished. The quantitative recovery of selenium during the separation procedure was determined using selenium-75 tracer and non-irradiated biological materials. Because of the higher melting point of sodium and potassium selenides (-875 “C), a higher temperature than was used for mercury was necessary for quantitative recovery of selenium. The method was tested for complete recovery of selenium by burning non-irradiated samples of biological material with known quantities of selenium-75 nitrate solutions added. The recovery selenium-75 was brought to 50-ml volume and compared to the activity of a like quantity of tracer diluted directly to 50-ml volume and counted in a standard configuration. The results of these tests indicated a complete recovery of selenium-75 tracer. A total of five separations was carried out with a mean recovery of 99.1 % and a relative standard deviation of 1.1 %. Using the analytical procedure developed, the selenium concentration in three new Standard Reference Materials (SRM) being issued by the National Bureau of Standards was determined. These materials are SRM 1630 SMR 1577 SRM 1571
Mercury in Coal Bovine Liver Orchard Leaves
The method described was used to determine the selenium contents in these materials with an accuracy in the mean results of better than *lo% relative at the 95% confidence level. The results were in excellent agreement with independent analyses using isotope-dilution spark-source mass spectrometry. The individual results are given in Table I. The results indicate that the described separation affords a degree of confidence in the h a 1 results not obtainable using the nondestructive activation analysis technique. Yet with the simple one-step combustion separation, samples may be processed in a minimum of time (-15 minutes each), maintaining the process-time advantage of nondestructive activation analysis when large numbers of samples must be analyzed. The technique has the added advantage that the mercury content, which is often of great interest in biological materials, may be determined simultaneously. ACKNOWLEDGMENT The cooperation of P. J. Paulsen and R. Alvarez is gratefully acknowledged. Their investigations and analyses gave special credence to this work. The authors would also like to thank H. Y. Yule for his suggestions and critique. RECEIVED for review December 3, 1971. Accepted January 31, 1972. In order to specify procedures adequately, it has been necessary to identify commercial materials in this report. In no case does such identification imply recommendation or endorsement by the National Bureau of Standards, nor does it imply that the material identified is necessarily the best available for the purpose.
Perchlorate Determination by Thermometric Enthalpy Titration Peter W. Carr’ and Joseph Jordan2 Department of Chemistry, The Pennsyhania State University, University Park, Pa. 16802 SATISFACTORY INSTRUMENTAL METHODS for the determination of the common oxy-anions SO4*-and c104-are conspicuous by their absence. Paradoxically, this problem is a specter of yesteryear which still haunts analytical chemistry in this day and age, A simple and convenient procedure, described and discussed in this paper, is ideally suited for the rapid and convenient quantitative analysis of perchlorate at millimolar concentration levels ; relying on judicious fundamental considerations, it involves a precipitation titration to a thermometric end point using an organo-arsenic reagent, and is readily amenable to automation. The cations tetraphenylphosphonium, tetraphenylarsonium, and tetraphenylantimoniumyield sparingly soluble perchlorate salts (1-5). The relevant precipitation reactions are of the type c104- A~44+= As44C104 ( S ) (1)
+
Present address, Department of Chemistry, University of Georgia, Athens, Ga. 30601 * To whom requests for reprints should be addressed. (1) (2) (3) (4) (5)
H. Nezu, Bunsaki Kaguku, 10, 561 (1961). R. J. Baczuk and W. T. Belleter, ANAL.CHEM., 39, 93 (1967). R. J. Baczuk and R. J. Dubois, ibid., 40,685 (1968). H. H. Willard and L. R. Perkins, ibid.,25, 1634 (1953). M. D. Morris, ibid., 37, 977 (1965).
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ANALYTICAL CHEMISTRY, VOL. 44, NO. 7, JUNE 1972
which is an analog of the process on which a classical-albeit unsatisfactory-gravimetric determination in the form of potassium perchlorate is based (4 = CsH& However, while the solubility of potassium perchlorate (6) in water at 25 “C is as high as 0.14 mole liter1, the molar solubilities of A S ~ ~ C ~ O ~ and Sb&C1O4 are 2 to 3 orders of magnitude smaller ( 2 , # , 5). The solubility product of P44C1o4 has apparently never been actually determined : pertinent literature (1-5) suggests strongly that it is comparable to that of the analogous tetraphenylantimonium and tetraphenylarsonium perchlorates. Indeed, various amperometric, conductometric, and potentiometric precipitation titrations of perchlorate with P+4+,Sb+4+, salts have been reported in recent years (1-3, 5. 7). and These were handicapped by limitations inherent in the methodologies used : thus, amperometric titrations are prone to interferences by numerous electroreactive species; conductometric titrations are interfered within the presence of excess electrolytes or in strongly acid solutions (where the high mobility of the hydrogen ion accounts for special difficulties). (6) I. M. Kolthoff, E. B. Sandell, E. J. Meehan, and S. Brucken-
stein, “Quantitative Chemical Analysis,” 4th ed., Macmillan, New York, N.Y., 1969, p 146. (7) M. J. Smith and S. E. Manahan, Anal. Chim. Acta, 48, 315 (1969).
The most interesting possibility, viz., potentiometry with a perchlorate ion selective electrode, has been explored and found wanting apparently because of a “selectivity coefficient” which was unsatisfactory with respect to nitrate (3, 8). Our rationale for investigating thermometric end points was predicated by the obvious a priori consideration that none of the aforementioned “electrochemical handicaps” could conceivably affect a calorimetric procedure.
tl
T
col .
EXPERIMENTAL Chemicals. Tetraphenylphosphonium bromide, tetraphenylarsonium chloride, and tetraphenylantimonium bromide were obtained from two sources, City Chemical Corporation, New York, N.Y., and Aldrich Chemical Company, Cedar Knolls, N.J. The phosphonium and arsonium salts were purified by double recrystallization from a water-ethanol mixed solvent. A solution of tetraphenylantimonium sulfate was prepared by the method of Morris ( 5 ) except that neutralization of tetraphenylantimonium hydroxide with sulfuric acid was carried out in boiling water. This modification of Morris’ technique yielded a 0.12M solution of the reagent. Stock solutions of tetraphenylarsonium chloride and tetraphenylphosphonium bromide were prepared as follows. The salt was dissolved in a minimum of boiling water; a gram of activated charcoal was added; the mixture was stirred for several minutes, filtered, cooled slowly to room temperature, and refiltered. The solution of tetraphenylarsonium chloride obtained in this manner was 0.25M in AS+^+ and very slightly yellow. This solution was undoubtedly supersaturated but no observable precipitation or change in titer occurred while the reagent was in use. However, precipitation did occur after six days. All other solutions employed in this investigation were prepared from reagent grade chemicals and standarized by classical procedures. Apparatus. Thermochemical measurements were performed in a simple adiabatic calorimeter consisting of a Dewar flask which was immersed in a water bath whose temperature was controlled to 25.00 i 0.01 “C with the aid of a thermistor thermoregulator (Model 71, supplied by Yellow Springs Instrument Co., Yellow Springs, Ohio). Titrant was delivered by a constant rate pump (Sage Instruments, White Plains, N . Y . ,Model 234-3) from a 10-ml ground glass syringe. Before entering the titration cell, the titrant flowed through a length of glass tubing which was submerged in the water bath. At no time did the volume of titrant added exceed the volume of the thermostated glass tubing. A Teflon (Du Pont) needle (Hamilton Company, Whittier, Calif., Catalog item KF24TF) of very small volume (30 pl) was affixed to the glass thermostating tube and immersed in the titration cell. Solutions in the Dewar were homogenized with the aid of a large fourbladed glass stirrer powered by a Sargent synchronous rotator (600 rpm). For monitoring temperature a recording thermistor bridge was employed which has been described elsewhere (9, 10). Electrical equipment for calorimetric calibrations included a heater and power supply built after a design by Stern, Whitnell, and Raffa (11). Procedure. The determination of end points by extrapolation of linear branches of thermometric titration curves is contingent on the virtual invariance of heat capacity. This requires that volume changes be minimized-i.e., that titrants be concentrated relative to titrands (9). Because of this, Sb&+ could not be used as a “direct titrant” in the conven(8) T. M. Hseu and G. A. Rechnitz, Anal. Left., 1, 629 (1968). (9) J. Jordan, in “Treatise on Analytical Chemistry,” I. M. Kolthoff and P. J. Elving, Ed., Part 1, Vol. 8, Interscience, New York, N.Y., 1968, pp 5175-5242. (10) R. A. Henry, Thesis, The Pennsylvania State University, University Park, Pa.. 1967. (11) M. J. Stern, R. Whitnell, and R. J. Raffa, ANAL.CHEM.,38, 1275 (1966).
5 sec H
Time Lapsed After
lnjectim
Figure 1. Typical “injection enthalpograms.” Excess of perchloric acid added to 100 ml of approximately 3 mM solutions I. Shape of curve obtained with tetraphenylphosphorusand tetraphenylarsenic 11. Shape obtained with tertraphenylantimony Arrow indicates time of HCIOl injection. Curves are shifted arbitrarily along horizontal axis tional sense, its solubility being on the order of 0.1M. Instead, a “back titration” technique was employed in this instance: an excess of Sb+4+ was added to the relevant perchlorate sample and the unreacted Sb+4+ was titrated to a thermometric end point (cide infra) with 1 M perchloric acid. It was feasible to prepare sufficientlyconcentrated solutions of As+4+(0.25M) for use as direct titrants. Volumes during any one titration were augmented by 4% or less which still permitted accurate extrapolation of end points. RESULTS Direct Injection Enthalpimetry (DIE) provides an exploratory tool, unmatched in rapidity and convenience, for determining hcats of reaction and discriminating between slow and fast kinetics (9, I O , 12-16); it involves the rapid injection (in 0.3 sec or less) of one reagent solution into an isothermal solution of another reagent in the adiabatic calorimeter and instantaneous mixing (to attain homogenization within 0.01 sec). Temperature-time curves (injection enthalpograms) are subsequently recorded. Ordinates of injection enthalpograms can conveniently be converted to “heat evolved” (expressed in calories) with the aid of suitable calorimetric calibrations. Injection enthalpograms obtained upon adding perchloric acid to solutions of P+4+, As&+, and Sb+4+ are illustrated in Figure 1. It is apparent from the figure that the perchlorates of P44+and As+4+were precipitated rapidly. In contradistinction, the precipitation of the Sb+4C104phase was sluggish. Reaction heats estimated from injection enthalpograms were combined with Gibbs free energy assignments calculated from solubility products available in the literature to yield corresponding entropies. The relevant (12) J. C. Wasilewski, P. T-S Pei, and J. Jordan, ANAL.CHEM., 36. 213 (1964). (13) J.-Jordan, R. A. Henry, and J. C . Wasilewski, ibficrochem. J., 10, 260 (1966). (14) J. Jordan and P. W. Carr, in “Analytical Calorimetry,” R. S. Porter and J. E. Johnson Ed., Plenum Press, New York, N.Y., 1968, pp 203-208. (15) J. Jordan, in “Topics in Chemical Instrumentation,” G. W. Ewing, Ed., Chemical Education Publishing Co., Easton, Pa., 1971, pp 193-199. (16) N. D. Jespersen and J. Jordan, Proceedings, International Symposium on Thermochemistry, Marseille, France, in press (1972). ANALYTICAL CHEMISTRY, VOL. 44, NO. 7 , JUNE 1972
1279
Table I.
Thermodynamicso of Precipitation Reactions c10,- = Z4JrClOc(s) Solubility product at AGO at 25 “C AS” cal Z A H o kcal.mole-1 25 “C of Z@rClOa kcal. mole-’ (mole. degree-’) Phosphorus -10.5 zk 0 . 3 Not available 1.2 x 10-8 - 11.o O?c3 Arsenic -10.9 i 0 . 4 3.98 x Antimony -11.5 zt 0.5 3.5 x 10-8 -10.0 -5 Zt 3 Heats of reaction estimated by DIE in this investigation; solubility products and free energies based on values in the literature ( 2 , 4, 5 ) ; entropies computed from columns 2 and 4.
a, + +
Table 11. Analysis of Perchlorate by Precipitation Titration with Tetraphenylarsonium Chloride. Concentration of “perchlorate unknown” Micromoles Precision,c Error,d (mmoles/l.) Taken Foundb 72 72 0.7 7 240.3 257.3 1.7 480.6 485.6 3.4 0.5 $1 624.6 620.6 0.5 +0.7 4.4 720.9 721.5 0.6 $0.08 5.1 6.85 961.2 955.5 0.6 -0.6 1201.5 1201.5 0.1 Zero 8.6 616.2 613.2 0.4 -0.5 4.3 Average precision in 0.003-0.008M range: 0 . 5 a Volume of samples titrated = 140 mi. ; titrant, 0 . 2 5 M A ~ $ ~ C 1 0 ~ . Based on the stoichiometry of Reaction 1. Expressed as the standard deviation of the mean of three replicates. Titrant was standardized against an accurately known amount of perchloric acid determined by classical alkalimetry. Everything else being equal, similar amounts of perchlorate yielded identical results irrespective of whether the starting material was KCIOa or HClOa. Table 111. Test of Anionic Interferences in the Analysis of Perchlorate Substance Perchlorate addeda found, mmolesb Error, 72 486 Reference None 498 KCI +2 486 KI Zero 486 NaF Zero 960 KMnOa 98 487 KBr03 +0.2 487 +0.2 KN03 548 KClOa $13 489 KHgPO4 +0.6 No end point KFluoroborate detected a In an amount equimolar to perchlorate present in solution analyzed. In sample volumes on the order of 140 ml.
+
thermodynamic data are listed in Table I. They reveal that the precipitation of all three perchlorates was comparably exothermic. On considerations of optimum kinetics and general convenience, tetraphenylarsonium chloride was investigated as a .preferred precipitant for determining perchlorate by enthalpy titration to a thermometric end point. Typical titration curves are shown in Figure 2. Representative precision and accuracy estimates, and the results of a n interference study, are presented in Tables I1 and 111, respectively. 1280
ANALYTICAL CHEMISTRY, VOL. 44, NO. 7, JUNE 1972
I
I 3
1.3 a/.
e d
,,
I
0 Tetrophmylorsoniurn CMoride Added
Figure 2. Thermometric titration curves of perchloric acid solutions Sample size: 130 ml. Perchlorate concentrations: I. 0.0017 M: 11. 0.0051 M. EP = end point. Arrow indicates start of titrant addition. Curves are shifted arbitrarily along horizontal axis It is apparent from Table I1 that in a range of concentrations between 0.003 and O.OOSM, perchlorate was determined with a precision of ‘ 1 2 % by precipitation titration with tetraphenylarsonium chloride t o a thermometric end point. The last column of Table I1 suggests strongly that the accuracy was comparable (note the balanced distribution of positive and negative errors). Below this range of concentrations, the precision remained the same but replicate end points were high by as much as 7%. This is readily accounted for by slow precipitation kinetics which are expected t o yield spurious end points which “lag” behind the stoichiometry of Reaction 1 (14). At concentrations higher than 0.008M, the precision and accuracy must necessarily be at least as good as between 0.003 and 0.008M (and probably better). The “dynamic range” of the method’s applicability has a n upper limit of 0.03M perchlorate, predicated by the maximum attainable concentration of the As+4+ titrant (0.25M) and the requirement that this exceed the titrand’s concentration by about a n order magnitude at least (9, 17). Naturally, the upper limit is further restricted by the solubility product of potassium perchlorate: however, in 0.03M perchlorate, this would be exceeded only when [IC+]> 4.7M. No other common cations form sparingly soluble perchlorates; thus, the presence of cations is not expected to interfere. It is evident from Table 111 that common anions, including halides and nitrate did not interfere. Permanganate and chlorate yielded large positive errors. The magnitude of the (17) H. J. V. Tyrrell and A. E. Beezer, “Thermometric Titrimetry,” Chapman and Hall, London, 1968.
error in the presence of permanganate was due to the fact that MnOa- was precipitated via the reaction: Mn04-
+ AsC#14+= AsC#14Mn04(s)
(2)
The reason for the interference by chlorate (while bromate had no adverse effect) has not been ascertained; possibly, tetraphenylarsonium chlorate (which is soluble per se) might have coprecipitated with tetraphenylarsonium perchlorate under our experimental conditions. Fluoroborate interfered drastically yielding sigmoid titration curves without any welldefined end point. While the effect of fluoroborate was not studied further, it appears consistent with the reasonable assumption that a kinetically slow process of the type BF4-
+ AS+^+ = BF~AsC#I~ (s)
(3)
might have prevailed.
The method described in this paper is remarkably catholic; it represents ZL preferred novel instrumental approach to the rapid and convenient determination of perchlorate in the presence of most other common anions and cations. In the present work, titrant addition was automated and titration curves were-likewise-automatically recorded. Even more complete automation, i.e., digital readout of the end point, is readily feasible using appropriate derivative electronic circuits (18).
RECEIVED for review September 17, 1971. Accepted January 6,1972. Based on a Ph.D. thesis by Peter W. Carr. Supported by Research Grant GP-11386 from the National Science Foundation. (18) P. T. Priestley, Analyst, 88, 194 (1963).
Determination of Sulfinic Acids in the Presence of Thiols by Titration with Aqueous Sodium Nitrite James P. Danehy and Victor J. Elia Department of Chemistry, Unicersity of Notre Dame, Notre Dame, Ind. 46556 SPECIFIC METHODS for the quantitative determination of sulfinic acids based on the reaction between the latter and nitrous acid have been previously published (1-3). The reaction is free from interference by the presence of other species which do not react with nitrous acid (e.g., sulfonic acids, thiolsulhates, thiolsulfonates). 2RS02H
+ HONO
+
(RS02)zNOH
+ H20
(1)
It has been proposed that, in the case of organic disulfides possessing certain structural features, alkaline decomposition occurs by a direct nucleophilic attack of hydroxide ion on one of the sulfur atoms to give thiol and sulfinic acid as the ultimate products (4). While the relatively stable sulfinic acids have actually been isolated subsequent to the alkaline decomposition of aromatic disulfides (5), 2RSSR
+ 40H-
+
3RS-
+ RS02- + 2Hz0
(2)
in the corresponding cases of aliphatic disulfides only sulfonic acids and thiols have been determined as ultimate products (4). Therefore, it was considered worthwhile to attempt to establish whether or not sulfinic acids are formed in these latter cases by using the method of titration with nitrous acid. For this method to be feasible, it had first to be established that titration with nitrous acid can be used in the presence of thiol and disulfide, both of which would be present in significant amounts. Kresze and Winkler (6) have reported that tert-butyl mercaptan reacts quantitatively with nitrous acid to form stable thionitrites. Ashworth and Keller (7) have extended this (1) C . S. Marvel and R. S. Johnson, J . Org. Chem., 13, 822 (1948). (2) J. L. Kice and K. W. Bowers, J . Amer. Chem. Soc., 84, 605 (1962). (3) B. Lindberg, Acta Chem. Scnnd., 17,383 (1963). (4) J. P. Danehy and W. E. Hunter, J. Org. Chem., 32,2047 (1967). (5) J. P. Danehy and K. N. Parameswaran, ibid., 33, 568 (1968). (6) G. Kresze and J. Winkler, Chem. Ber., 96,1203 (1963). 39, 373 (1967). (7) G. W. Ashworth and R. E. Keller, ANAL CHEM.,
observation to other thiols and have developed a spectrophotometric method based thereon for the quantitative determination of thiols. RSH
+ HONO
--*
RSNO
+ HzO
(3)
Since the concentration of thiol in a mixture of thiol and sulfinic acid can be determined without interference from the latter by the use of Folin’s reagent (8), it should be possible to determine the concentration of sulfinic acid in such a mixture by titration with nitrous acid provided that: the thiol under investigation is titrated quantitatively by nitrous acid; the disulfide does not interfere; and sulfinic acid reacts quantitatively in the presence of thiol. EXPERIMENTAL
Materials and Reagents. The thiols and disulfides were purchased, received as gifts, or synthesized by us, as related in earlier reports ( 4 , 5). Folin’s reagent (a phosphotungstic acid) was prepared and used as previously described (8). The sodium nitrite and other chemicals were all of ACS reagent grade. Titration of Aliphatic Thiols and of p-Chlorobenzenesulfinic Acid by Nitrous Acid. An accurately weighed sample of thiol or of p-chlorobenzenesulfinic acid was dissolved in 10 ml of water, stirred magnetically, and 5 ml of 5N H2S04 was added. The solution was then titrated with continuous stirring by addition of standardized sodium nitrite solution until an external end point was obtained: when one drop of the solution gave an instantaneous blue color when added to a starch-KI solution. Titration of Aromatic Thiols by Nitrous Acid. The titrations of p-chloruthiophenol, thiophenol, and m-mercaptobenzoic acid were carried out in the following manner. A sample of the corresponding crystalline disulfide was dissolved in 10-15 ml of dimethylformamide (DMF), and 3 ml of (8) J. P. Danehy and J. A. Kreuz, J . Amer. Chem. SOC.,83, 1109 (1961). ANALYTICAL CHEMISTRY, VOL. 44, NO. 7, JUNE 1972
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