Anal. Cbern. 1980, 5 2 , 1912-1922
(2) Krauskopf, K. "Source Rocks for Metal Bearing Fluids", in "Geochemistry of Hydrothermal Ore Deposits", Barnes, H. L., Ed.; Rhinehart and Winston: New York, 1967, p 22. (3) Holland, H. D. €con. Geol. 1972, 6 7 , 281. (4) Tooms, J. S. Trans. Inst. Mln. Metal., Sect. B 1970, 79, 116. (5) Dunham, K. C. Trans. Inst. Min. Metal., Sect. B 1970, 7 9 , 127. (6) Roedder, E. U . S . Geol. Surv., Prof. Pap. 1972, 440-JJ. (7) Goldschmidt, V. M. "Geochemistry"; Clarendon Press: Oxford, 1954; pp 588-589. (8) Behne, W. Geochim. Cosmochim. Acta 1953, 3 , 186. (9) Noble, D. C.;Smith, V. C.; Peck, L. C. Geochim. Cosmochim. Acta 1967, 3 1 , 215. (10) Van de Kamp, P. C. J . Geol. 1970, 78, 281. (11) Boyle, R. W. Geol. Surv. Can. Mem. 1961, 370. (12) Huang, W. H.; Johns, W. D. Geochim. Cosmochim. Acta 1967, 37, 597-602. (13) Johansen, 0.; Steinnes, E. Geochlm. Cosmochim. Acta 1967, 3 1 , 1107-1109. ( 1 4 ) Fabbi, B. P.; Espos, L. F. Appl. Spectrosc. 1972, 2 6 - 2 . 293-295. (15) Fabbi. B. P.; Espos, L. F. U . S . Geol. Serv., Prof. Pap. 1976, 840, 89-93.
(16) (17) (18) (19) (20)
Small, H.; Stevens, T. S.; Bauman. W. C. Anal. Chem. 1975, 4 7 , 1801. Abbey, S. X-ray Spectrosc. 1978, 7 - 2 , 99-120. Terashima, S. Bull. Geol. Surv. Jpn. 1974, 2 5 , 175-179. Flanagan, F. J. Geochim. Cosmochim. Acta 1973, 37. 1189-1200. Ando, A . ; Kurasawa, H.; Ohmori, T.; Takeda, E. Geochem. J . 1974 8 , 175- 192. (21) Dreibus. G . ; Spettle, B.; Wanke. H. "Origin and Distribution of the Elements", VoI. 11; Pergamon Press: New York, 1979; pp 33-38. (22) Clark, R . S., Jr.; Jarosewich. E.; Mason, 6.; Nelen, J.; Gomez, M.; Hyde, J. R. Smithson. Contrib. Earth Sci. 1970, 5 , 44-53. (23) Flanagan, F. J.; Chandler, J. C.: Breger, 1. A , ; Moore, C. 8.;Lewis, C. F. U . S . Geol. Surv., Prof. Pap. 1976, 840, 123-126.
RECEIVED for review April 18, 1980. Accepted .June 16, 1980. This research was supported in part by NASA grant NSG03-001-001. This work was presented at the 22nd Rocky Mountain Conference of the Rocky Mountain Chromatography Discussion Group a t Denver, Colo., August 11-14, 1980.
Determination of Atmospheric Sulfur Dioxide without Tetrachloromercurate(I1) and the Mechanism of the Schiff Reaction Purnendu K. Dasgupta," Kymron DeCesare, and James C. Ullrey California Primate Research Center, University of California, Ua vis, California
Formaldehyde ( 7 mM) buffered at pH - 4 is used to stabilize atmospheric SO2 as hydroxymethanesulfonic acid. Equilibrium data for the above reaction are presented. Sulfite, liberated from the compound by base, is added to acidic pararosaniline for color development by the Schiff reaction, and absorbance is measured at 580 nm. The procedure has been optimized with regard to acidity and reagent concentrations. The method is comparable to the West-Gaeke method (Anal. Chern. 1956, 28, 1816) in absorption and recovery efficiency, sensitivity, and precision. No unusual interferences are observed due to 03,NO,, and transition-metal ions, except Mn(I1). A novel ion chromatographic procedure to determine hydroxymethanesulfonate and sulfate in the same sample is also described. Enhancement of sensitivity in the colorimetric method by solvent extraction has been studied. Investigations into the mechanism of the Schifi reaction and structure of the products have established the validity of the alkylsulfonic acid theory. Mechanistically an arninocarbinol seems to be the first intermediate, which undergoes subsequent nucleophilic substitution by bisulfite ion. To account for the high absorptivity of the product, we suggest that the sulfonic acid group is significantly ionized.
I n 1866, Schiff ( I ) reported the color regeneration in a SO2-bleached fuchsin solution upon the addition of an aldehyde. Extensive application of the Schiff reagent (fuchsin-sulfurous acid) to identify carbonyl compounds was pioneered by Schmidt (2) and identification of other functional groups, which can be oxidized to the carbonyl moiety, was introduced by Bauer ( 3 ) and McManus ( 4 ) . Kasten (5) and P e a c e (6) have adequately reviewed the histologic applications of this uniquely important reaction in histochemistry. Steigmann (7) turned the reaction around and utilized acid bleached fuchsin and formaldehyde for the qualitative identification of sulfites; a quantitative procedure by Grant ( 8 ) followed. Kozlyaeva (9) reported the determination of airborne SOz and used evacuated flasks for collection. Atkin ( I O ) used an alkaline glycerol solution, an absorber first described by Haller ( I I ) , to determine relatively high levels of SO?. Urone and Boggs (12) reduced the alkali content of Haller's absorber by a factor of 25 and described the first useful method for measuring ambient levels of SO1 (13, 1 3 ) . It !\a<. 0003-2700/80/0352-1912$01 O O / O
9 5 6 16
however, the work of West and Gaeke ( 1 5 ) ,utilizing sodium tetrachloromercurate(I1) (TCM) as absorber, after Feigl (16), and pararosaniline instead of fuchsin (an impure mixture of pararosaniline and rosaniline) that established the basis for the method of choice in determining ambient SO2. T h e method found other important applications ( 17-19), was automated (20,21),and after relatively minor refinements remains an international reference method (22, 23). In spite of the advent of other colorimetric techniques ( 2 4 , 25), numerous instrumental techniques with various degrees of sophistication (reviewed in (26)) as well as such promising recent entries as ion chromatography (27)and pulsed fluorescence ( 2 8 ) ,the West-Gaeke (WG) method remains in use because of its simplicity, sensitivity, specificity, and marginal equipment cost. Refinements made to the original LVG procedure include studies on the purity of the dye (29),temperature dependence and stability of the chromogen (ZO), interferences due to ozone, oxides of nitrogen, and transition-metal ions and prevention of such (30-33), color enhancement by nonaqueous solvents and equilibrium effects ( 3 4 ) ,and an elaborate optimization of all pertinent parameters (35). However, the major existing deterrent toward the continued use of this technique is the manipulation of relatively high concentrations of toxic and expensive mercuric chloride. Structural studies by Nauman et al. (36) indicated that the mechanism of the WG reaction involves initial formation of a formaldehyde bisulfite addition compound (hydroxymethanesulfonic acid) and culminates in an aminomethanesulfonic acid. During our studies on the dissociation of hydroxymethanesulfonic acid to formaldehyde and S(IV), the possibility of absorbing SO2 directly into an appropriately buffered solution of formaldehyde and then reacting the intermediate adduct with pararosaniline became clear. This paper presents a modified WG method which does not use any mercury compound. Structure of the product and reaction mechanisms are reexamined in light of the new evidence. EXPERIMENTAL SECTION Reagents. Pararosaniline hydrochloride, purified according to Scaringelli et al. (35'), was obtained as a 0.3% solution in 1 M HC1 (Harleco). This solution (113.9 mL) and concentrated HC1 3 mL) were diluted to 1 I, for the working reagent. This is C 1980 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 52,
essentially the same reagent as described by West and Gaeke ( I s ) , designed to produce the same acidity in a final volume of 25 mL instead of 1 2 mL. Sodium sulfite, analytical reagent grade, was purified by repeated recrystallizations in an argon atmosphere (37). The purity of the final product was assayed by iodometric titration by using KI03 as primary standard and found to be better than 99%. Determination of total sulfur by H202oxidation and BaS04 gravimetry agreed satisfactorily. Sodium hydroxymethanesulfonate (Aldrich, 98% ) was purified through precipitation from a saturated aqueous solution by adding ethanol, filtering, and drying at 80 "C. Freedom from sulfate was checked by ion chromatography (38). The sodium content of the purified product was assayed by ion chromatography, and total sulfur was assayed by BaS04gravimetry after alkaline hydrolysis and H202oxidation. The purity of the final product was better than 99%. Buffered formaldehyde adsorber was prepared by diluting formaldehyde solution ( 3 7 % , 530 pL) and potassium hydrogen phthalate (KHP, 204 mg) to 1 L. The reagent (7 mM in HCHO and 1 mM in KHP) exhibits a pH of 4.2 at 25 "C. All other reagents were of analytical reagent grade and used without further purification. Equipment. A Beckman 35 kinetic system, equipped with a 120-~Lquartz flow-through cuvette and a Gilson Minipuls 2 sipper pump, was used for all spectral measurements. The sample did not come in contact with any surfaces other than glass or PTFE prior to optical measurement, and all measurements refer to a temperature of 25 f 1 "C. All gas dilutions and sampling were conducted in Hinners-type environmental chambers (39) of stainless steel and glass construction and ranging in size from 0.44 to 4.5 m3. These chambers allow very large dilutions of test gases with clean, dry, filtered air (40) and provide up to eight equivalent sampling ports for concurrent sampling. All sampling was conducted a t a controlled temperature of 25 "C and a controlled relative humidity of 50% as measured by a Bendix Psychron hygrometer. Sulfur dioxide was obtained as the pure gas or as, 4.3 (&5%)ppm in N2 (Matheson),and chamber concentrations were monitored, independent of the absorbers, by a flame photometric total sulfur analyzer (Meloy SA-285). For high concentrations of SO2,a heated orifice diluter (Meloy Pneumotron:l was connected upstream of the sulfur analyzer and adjusted to give a 1OO:l dilution ratio. The sulfur analyzer was calibrated. with a permeation tube-based Dasibi gas calibrator, Model 1005 CE-2. Nitrogen dioxide was obtained as 0.2% in N, (Matheson:l and chamber concentration was monitored by a chemiluminescent analyzer, TECO Model 14T. The analyzer was calibrated by gas-phase titration methods with the Dasibi gas calibrator. Ozone was generated from medical grade O2 by a silent arc discharge instrument, Sander Model VI, and monitored with a Dasibi 1003AH ultraviolet ozone monitor which in turn was calibrated with a Dasibi 1008-PC absolute ozone photometer. Midget size ( 2 5 mL) fritted glass bubblers (pore size 145-175 pm) were used for sampling. Larger bubblers (500 mL) with similar frit porosities were used to accommodate a larger absorber volume when sampling high concentrations of SO2 or to permit a larger sampling: rate without absorber losses. All flow meters were calibrated with dry test meters and pressure drops across sampling trains were monitored and corrections applied as necessary. Measuremen: of pH was made by an Orion 701A pH meter equipped with a Model 91-02 research electrode and an automatic temperature compensator probe. Dissociation of Hydroxymethanesulfonic Acid to Formaldehyde and S(IV). The equilibrium was studied by ultraviolet spectrophotometry a t every integral pH value from 2 to 9 with the exception of pH 3 for which a suitable optically transparent buffer system was not available. Aqueous solutions of formaldehyde do not absorb in the LTV region due to hydration to tht: gem-diol. Hydroxymethanesulfonic acid, a saturated compound, also displayed no ultraviolet absorption. The U' spectrum of the compound in aqueous solution is solely due to the S(1V) species resulting from dissociations of the type OHCH2S03-= HCHO + HS03-. This is corroborated by the fact that a t any given pH, the spectrum of a solution of the addition compound is superimposable on the spectrum of a S(1V) solution of appropriate concentration at the same pH. Sodium hydroxynethanesulfonate
NO. 12, OCTOBER 1980
solutions ranging in concentration from 1.00 to 100 mM were made up in appropriate buffer solutions of ionic strength 0.25 M and allowed to equilibrate until no further increase in optical absorption was observed. Absorbance values were measured at 10 or more different wavelengths a t which there was significant absorption (>0.1 AU). Measurements were made near the absorption maxima wherever possible, but in most cases this was located beyond measuring range (<200 nm). Sulfur(1V) concentrations in the adduct solution were calculated from similarly measured values of standard sodium sulfite solutions (0.25-1.00 mM) made up in the same buffer. All operations were carried out in an argon atmosphere except for the actual spectral measurements which were made in P T F E stoppered quartz cells. Water used as solvent was deionized, distilled, boiled, and purged with argon for several hours. Collection, Analysis, a n d Calibration Procedure. Midget bubblers, filled with 15 mL of the buffered formaldehyde absorber, were used for sampling. The usual sampling rate was 250 mL/min and did not exceed 400 mL/min, above which absorber loss due to frothing became significant. Antifoaming agents (Dow Antifoam A) could be added to permit a higher sampling rate, within limits, without any interference. Larger bubblers were, however, necessary for flow rates approaching 1 L/min or higher. Subsequent to collection, the sample could be stored a t room temperature for up to 1 month without any appreciable degradation. For analysis, the contents of the fritted tube and the side arm were washed into the bubbler with a small amount (2 mL) of water. Next, NaOH (4.5 M, 1 mL) wi3s added from a dispensing pipet, and the bubbler was stoppered with a ground glass stopper and inverted once. The contents of the bubbler were then transferred into pararosaniline working reagent (5 mL) kept in a 25-mL culture tube. (The reverse process, i.e., addition of acidic pararosaniline to the alkaline solution, leads to lower absorbance values and poor precision.) The contents of the tube were then made up to the neck (25 mL) with water, and the tube was capped and inverted thrice. The optical absorbance of the resulting color was measured at 580 nm after a period of 10 min and not exceeding 15 min. For calibration, purified Na2S03(9.84 mg) and HCl (1 M, 160 PL) were made up to 1 L with the buffered formaldehyde absorber. This solution contained the equivalent of 5 pg/mL SOz, and calibration plots were constructed by diluting x mL of this standard with (15 - x ) mL of the absorber and subjecting to color development as detailed in the previous paragraph. Ion Chromatographic Analysis. Ion chromatographic determinations were carried out with a Dionex Model 10 ion chromatograph. Anionic analysis by standard procedures with this technique involves an alkaline carbonate-bicarbonate eluant (38). Such conditions lead to on-column dissociation of hydroxymethanesulfonate to sulfite. Since the plumbing of the instrument is permeable to air, significant oxidation of sulfite (whether injected as hydroxymethanesulfonate or as such) to sulfate often occurs (41) making it impossible to determine the original amount of sulfate present in such a sample. An analytical method involving 1 mM KHP as eluant was developed to solve this problem. With this eluant, the hydroxymethanesulfonate ion eluted as such and was well separated from sulfate. The conductance detector on the instrument was calibrated with standard solutions of KCl, and all reported data have been corrected accordingly. Product Isolation a n d Identification. Subsequent to color development, the solutions were ext.racted with 1-butanol or 1-pentanol. The organic layer was washed twice with small amounts of 1 M NaCl to remove all 'HgC1, when color was developed according to the WG method. In all cases, the organic layer was next washed thrice with water. The solvent was then removed in vacuo. Macroscale synthetic attempts, designed to isolate milligram quantities of product(s), involved a similar procedure, except that all reagent concentrations were increased by a factor of 10-25, and the isolated crude product was chromatographed repeatedly on a 25 X 1500 mm column packed with Silicar CC-7 (Malinckrodt) using gradient elution with a CHCl,/MeOH solvent system. Thin-layer chromatography was carried out on 0.25 mm thick silica gel 60 F (Merck) plates and eluted with methanol/ethyl acetate (1:l).
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980
Table I. Composite Equilibrium Constant for the Hypothetical Equilibrium HCHO-S(1V) Adduct --L HCHO + S(1V) at 25 "C, I = 0 . 2 5 buffer system PH K' PK ' KCl-HCl phosphate phosphate phosphate Tris
1.54 x 1.56 x 1.17 x 2.53 X 7.57 x 2.16 X 1.83 x
5 6 7
4.81 4.93 4.60 4.12 3.66 2.74
10-5 1 0 - 5
/ 98 8 ppm S O 2 2 5 0 r n L 7 0 r n M H C H G . IOrnM KHP absorber Sampling rates f o r curves from l e f t t o r i g h t 1 1 . 1 , 9.0, 7.2. 5 0 ord 3.0 L/min
Table 11. SO, Absorption Efficiency of the Buffered Formaldehyde Absorber
no. of expts
8 8 10 10 8
6 7 8
9 10 11
0.05 0.11 0.11
0.23 0.23 0.52 0.52 1.08 1.08
4 2 2 4 4 4 4
Sam p 1in g rate, mL/min 250 40 0 250 40 0 25 0 4 00
250 400 250 400 250 400 250 400
duration, min 1000- 2000 1000-15 00
180-1 000 120 180
120 90 60 80
N,N',N"-Pentamethylpamrosaniline was prepared by repeated column chromatography of biological stain grade methyl violet (an impure mixture of tetra-, penta-, and hexamethylpararosaniline). The chromatographic conditions were as described above. The final product was isolated in the carbinol form and recrystallized from ethanol. Anal. Calcd for CZ4Hz9N30: C, 76.76; H, 7.78; N, 11.19. Found: C, 75.82; H, 7.69; N, 11.24. Bis(p-aminopheny1)-p-azomethinophenylcarbinol, the mono(azomethine) Schiff base of pararosaniline in the carbinol form was prepared by adding 10 mM HCHO to a threefold excess of 10 mM pararosaniline in 1 M HCl, heating for 4 h at 80 "C, and neutralizing gradually with NaHC03. The insoluble Schiff base was freed from the pararosaniline leuco base by repeated washings with ethanol. Anal. Calcd for CZ0Hl9K30:C, 75.68; H, 6.04; N, 13.24. Found: C, 75.53; H, 6.10; N, 13.38. RESULTS AND DISCUSSION T h e Equilibrium Constant f o r the Dissociation of the Formaldehyde-Bisulfite Addition Compound. Equilibrium data are presented in Table I. The relative standard deviation of the multiwavelength data set for any given p H was less than 2% in all cases. Close examination of the p H dependence of K' will indicate that the actual operative equilibria are considerably more complex than a system consisting of the hydroxymethanesulfonate ion dissociating into formaldehyde and the bisulfite ion. The pattern of the data qualitatively agrees with that observed for the benzaldehyde-sodium bisulfite addition compound by titrimetric methods ( 4 2 ) , substantiating the presence of other independent equilibria such as HOCH2S03H HCHO + SO2. H 2 0 ,-OCH2S03- HCHO + SOj2-,etc. Detailed implications of these and related results will be presented elsewhere. For present purposes, a composite equilibrium constant ( K ? for the hypothetical single equilibrium expression HCHO-S(1V) adduct HCHO + S(IV), without differentiating between t h e various S(1V) species, may be expressed as K ' = [S(IV)]2/([adduct,,] - [S(IV)]). When pure sodium hydroxymethanesulfonate is allowed to dissociate, the concentration of free HCHO (including the diol) in solution must equal the total concentration of the liberated S(IV) species. I t is seen from Table I that K' has a minimum value of about and reaches this value somewhere between pH 4 and 5 , in reasonable agreement with the value calculable from
Figure 1. The absorption efficiency of a buffered formaldehyde absorber for high levels of SO2 at high sampling rates.
the data given by Arai ( 4 3 ) . At this pH, the bisulfite ion is virtually the only S(1V) species present in significant concentration. From the first and second dissociation constants of S 0 2 . H 2 0 , as given in ( 4 4 ) and ( 4 5 ) , respectively, [HS03-]/[S(IV)],d reaches a maximum value at a p H of 4.1. At p H levels between 4 and 5 , the generally accepted single equilibrium expression involving only the bisulfite ion adequately describes the system. Thus, if any given amount of SO2 is to be absorbed and stabilized as the adduct in an absorber containing formaldehyde at any given concentration, the optimum pH of the absorber solution should be between 4 and 5 since the sum total of the concentrations of the free S(1V) species (available for oxidation or desorption removal) is minimum a t this pH. Absorption Efficiency of t h e Buffered Formaldehyde Absorber. Two midget bubblers connected in series, the first containing 15 mL of the formaldehyde absorber and the second containing 15 mL of 0.1 M TCM, were used to test the efficiency of the absorber in capturing SO2. Table I1 shows the SO2concentrations sampled (FPD),flow rate, and duration of sampling. T h e T C M solution was analyzed according to the WG method, in no case was any significant increase over blank observed. The formaldehyde absorber functions, within the limits of this experiment, with a 100% efficiency for collecting the sampled SOz. T o determine the feasibility of using a formaldehyde absorber to capture SO2 efficiently a t stack concentrations, we sampled sulfur dioxide, 98.8 ppm in concentration (FPD), through 250 mL of a solution 70 mM in HCHO and 10 m M in K H P at flow rates of 3.0, 5.0, 7.2, 9.0, and 11.1 L/min. The SO2 concentration downstream of the bubbler was measured with the total sulfur analyzer. The extent of breakthrough is shown in Figure 1. Even under such adverse conditions, the SO2 absorption efficiency was nearly 99% a t 5.0 L/min for the first 20 min and better than 99.9% a t 3.0 L/min for the first 30 min, at which points the absorber contained -100 pg/mL SO2. Ion Chromatographic a n d Conductance Studies. With 1 mM K H P as eluant, the background conductance was relatively high ( - 175 pS cm-') but still completely suppressible. U'ith a 3 X 150 mm precolumn, 3 X 250 mm separator column and a 6 X 250 mm suppressor and an eluant flow rate of 2.3 mL/min, hydroxymethanesulfonate eluted with relatively little retention, with about the same retention time as that of fluoride under standard conditions. The calibration plots (1OO-pL loop injection) under these conditions were linear for both sulfate and hydroxymethanesulfonate with slopes of -7 /IS cm-' mM S042-and -45 pS cm-' mM OHCH2S03-, respectively. An assessment of the dissociation constant for hydroxymethanesulfonic acid cannot be made from a direct comparison of these two slope values inasmuch as the hy-
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980 ~ _ _ _ _ _ _ ~ ~
Table 111. Stability of SO, Samples Collected in the Buffered Formaldehyde Absorber SO, found
condition original sample 8 h exposure to bright sunlight storage (room temp, fluorescent light) 1 day 3 days 5 days 10 days 15 days 20 days 25 days 30 days refrigerated, 30 days aeration 8 h, 250 mL/min 18 h, 250 mL/min 8 h, 400 mL/min storage, 37 O c a original sample 3 days 5 days 10 days a
as hydroxymethanesulfonate as sulfate (in &/mL (in [email protected]
49.0 48.9 48.9 48.8 48.8 48.9 48.7 48.6 48.8
0.0 0.0 0.1 0.1 0.2 0.3 0.3 0.4 0.0
48.8 48.8 48.5
0.0 0.0 0.0
0.991 0.987 0.972 0.959
By the colorimetric procedure reported in this paper.
droxymethanesulfonate elutes as a very sharp peak and the much more strongly retained sulfate as a broad peak. Data were obtained for Kohlrausch type conductance plots for various concentrations of hydroxymethanesulfonic acid. The independent dissociation of this compound into HCHO and S(IV),aside from its dissociation as an acid, made quantitative interpretation of the conductance vs. concentration data difficult. Qualitatively, the compound is not a weak acid; the acid dissociation constant is certainly substantially higher than the first dissociation constant of S0,.H20. In the absence of an established method for determining hydroxymethanesulfonate, the ion chromatographic procedure permitted the analysis of SO, samples absorbed in the buffered formaldehyde reagent to determine not only the hydroxymethanesulfonate concentration but also the extent to which oxidation to sulfate had taken place in a stored or aerated sample. The analysis is facilitated by the fact that both the sample and the eluant have the same ionic background of 1 mM K H P , thus minimizing any undesirable phthalate peak or solvent dip. Formaldehyde itself does not show up in an ion chromatogram. Stability of Collected SO2 Samples. Sulfur dioxide was absorbed to the extent of -50 Fg of S 0 2 / m L in a buffered formaldehyde absorber. Aliquots were then subjected to various treatment. Results of ion chromatographic analyses for these samples are presented in Table 111. Sulfur dioxide stabilized in 0.1 M TCM is reported to be oxidized at a rate of l % / d a y a t room temperature (35). The formaldehyde absorber, even though much more dilute in concentration, appears to be far more effective in retarding such degradation to no more than 1% /30 days. The lower value for the sample aerated 8 h a t 400 mL/min is probably due to solution loss due to aerosolization since an attendant amount of sulfate was not found. There is no apparent photochemical degradation. Storage at 37 "C leads to a higher degradation rate of approximately 0.3% /day, which we consider acceptable. T h e mechanism and rate of aerobic oxidation of S(1V) solutions are still to be unequivocally established because of t h e high susceptibility of such systems to both positive and
negative catalysis. Within the p H range 4-9, the uncatalyzed rate of oxidation of S(1V) solutions may generally be accounted by assuming that the SO3,- species is oxidized far faster than the HS03- species, and for most practical purposes one may disregard the second process altogether. For the normalized rate expression -d[S(IV)]/dt = k[SO:-], the magnitude of the reported rate constant in units of reciprocal seconds extends from IO+ to lo-,, though most studies report it to be about (46). If we assume that formaldehyde exerts no catalytic effect in either accelerating or retarding the oxidation of the free S(1V) present in a solution containing S(1V) stabilized in formaldehyde, the rate of conversion of the total sulfur originally present to sulfate for a system at pH 4.2 containing 7 m M HCHO may be calculated to be -2.5%/30 days assuming K' (as in Table I) to be the rate constant k to s-l, and the second dissociation constant for SO2-H20 be to be 4.6 X (45). Considering the limitations of the available rate data, the agreement with the observed oxidation rate of 1%130 days is quite goocl. Development of Color: t h e Analytical P r o c e d u r e . A solution of SO, absorbed in the buffered formaldehyde reagent or a solution of sodium hydroxymethanesulfonate did not react with acid-bleached pararosaniline to develop any appreciable amount of color, in direct contradiction to the generally accepted mechanistic hypothesis for the WG reaction (36). Indeed, this perplexing fact prompted us to develop the alternative ion chromatographic procedure to make certain that the absorption of SO2 into a formaldehyde absorber does lead to the stable hydroxymethanesulfonate species. It was found that if HCHO and S(1V) are regenerated from the addition compound in solution by adding strong base, followed by the addition of pararosaniline containing enough acid to lower the final pH to the usual operating region of the WG method, the characteristic WG color developed. T h e initial data, however, showed a remarkable lack of precision. Since the absorber already contained approximately the same amount of HCHO as called for in the WG procedure, additional HCHO was unnecessary. This left only one other alternative in the order of mixing the reagents, namely, addition of pararosaniline to the sample prior to alkaline decomposition, followed by the addition of an appropriate amount of acid. Color developed in this procedure as well, but the sensitivity was lower than that of the first procedure and no improvement in precision resulted; this approach was henceforth discarded. Continuing with the first method, substitution of maleic or oxalic acid solutions of appropriate strength for the hydrochloric acid solution of the pararosaniline reagent, designed to offer better buffering in the final mixture at the p H range of interest, did not improve precision. Any variation of time between the addition of base to the sample and the addition of acidic pararosaniline was not regarded to be an important factor. In strong base, the decomposition of the addition compound is virtually instantaneous as evidenced from spectrophotometric studies; t h e s a m p l e is completely conuerted t o HCHO a n d S ( I V ) prior to t h e a d d i t i o n of t h e pararosaniline. Variation due to any oxidative loss of the S(IV) prior to pararosaniline addition is unlikely inasmuch as oxidation of S(1V) a t pH 12 and above proceeds extremely slowly ( 4 7 , 4 8 ) ;even without stabilizers, such as glycerol, KOH has been reported to stabilize S(1V) to the same extent as TCM ( 4 9 ) . Experiments with intentional variation in time between alkaline decomposition and pararosaniline addition also indicated that this factor did not contribute significantly to the lack of precision. Further experiments showed that the exact mixing conditions of the acidic pararosaniline with the alkaline sample is intimately related with the final color intensity. It appeared that upon acidification with the pararosaniline reagent, two competing reactions were taking place. The first is a reaction
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980
Table IV. Reproducibility of Method
std dev of points,a AU
slope of calibration line,b AU pg-' mL-'
0.558 0.568 0.570 0.567 0.567 0.560 0.590 0.564 0.559
0.007 0.002 0.008 0.010
std dev of slope,C A U p g - l mL-'
blank, AU computedd observed 0.037 0.032 0.040 0.030 0.030 0.039 0.032 0.035 0.029
0.011 0.005 0.012 0.010 0.004 0.012 0.008 0.007 0.006
0.036 0.031 0.041 0.031 0.031 0.038 0.032 0.038 0.028
Recommended regression line: absorbance = 0.567 x crg/mL SO, + 0.030 Percent error (single determination) at 1 pg/mL (25 fig SO, total): 2.7% (at 95% confidence level) Reproducibility of slope: 0.557 t 0.019 (at 95% confidence level) a Deviations from the mean, six observations per set. tions per set. Intercept of regression line.
Deviation from the mean slope, six observa-
Table V. Optimization of Final pH: Effect of Varying Amounts of Basea NaOH added, mmol 1.90 2.85 3.80 3.99 4.18 4.37 4.56 4.75 4.94 5.13 5.32 5.51 5.70
final pH 0.95 0.97 0.99 1.03 1.06 1.09 1.12 1.17 1.21
nm 579.0 578.5 578.5 578.0 578.0 577.5 577.5 577.5 577.5 577.0 576.5 576.5 576.0
blank absorbance, AU (at Amax) 0.016 0.017 0.019 0.021 0.023 0.026 0.028 0.030 0.033 0.038 0.043 0.046 0.050
Amax, A U
0.482 0.522 0.566 0.575 0.584 0.594 0.595 0.598 0.598 0.598 0.599 0.599 0.600
sample-blan k absorbance, A U (at AmELx) 0.466 0.505 0.547 0.554 0.561 0.568 0.567 0.568 0.565 0.560 0.5 56 0.553 0.550
a 25 pg SO, in 1 5 mL 7 mM HCHO, 1 mM KHP absorber, x mmol NaOH, into 5 mL pararosaniline reagent, diluted to 25 The absorption band is more than 5 nm wide at maximum absorbance ( t 1%).Measurement at the recommended mL. wavelength of 580 nm does not significantly change the sample-blank absorbance although the blank value itself may change slightly,
involving pararosaniline and leading to the colored product, and the second is a direct recombination of the HCHO and S(1V) to form the colorless adduct. Thus one would expect to find the highest and the most reproducible color intensity when the alkaline sample contacts the highest available concentration of pararosaniline as neutralization occurs. For any given concentration of pararosaniline (which contributes to the blank and cannot therefore be indefinitely increased), this condition is most easily met if the alkaline sample is added to the acidic pararosaniline and not the other way around. This change in the procedure resulted in excellent precision in the analytical data; typical data for the recommended procedure are presented in Table IV after Scaringelli et al. (35). The calibration plot (not shown) is linear up to at least 30 wg of SO2,corresponding to -0.7 AU, with a correlation coefficient better than 0.999. The method is equivalent t o the WG procedure in precision and sensitivity. One interesting difference in calibration characteristics is that the blank absorbance value from the regression line actually coincides closely with the experimental blank value; for the U'G method the regression value is often significantly different from the experimental blank value. Final acidity was optimized by varying the amount of base added for alkaline decomposition and retaining the other parameters constant. Table V shows the data obtained. The optimum absorbance is reached with 4.4-4.7 mmol of NaOH leading to a final p H of about 1.05. In the recommended procedure, 4.5 mmol of NaOH (1 mL, 4.5 M) is used. T h e amount of pararosaniline was similarly optimized by retaining final p H and other factors constant while varying the pararosaniline concentration in the reagent. The solid
0 600 r
< ° 0 ' 000 0 00 0
...-.. Blonk Absorbance
Pararosaniline Concentralion Equivalent I o Milliliters of Working Reagenl
Flgure 2. Effect of pararosaniline concentration on final color intensity. curve in Figure 2 shows the sample-blank absorbance as a function of the amount of pararosaniline added, plotted in terms of the pararosaniline content of the working reagent. T h e dashed line shows the blank absorbance multiplied by a factor of 10. T h e optimum amount of pararosaniline was determined to be equivalent to that contained in 5 mL of the
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980
Table VI. Effect of Varying Formaldehyde Concentration in Absorber on Color Intensity
mM HCHO in absorber
sample-blan k absorbance at 580 iim ? SD,=
( 1 5 mL)
blank absorbance at 580 nm i SD,= AU
0.578 i 0.005 6.1 0.569 i 0.004 6.5 0.565 t 0.007 7.0b 0.572 i 0.004 7.5 0.573 t 0.007 7.9 0.575 i 0.010 a Five measurements on each set. concentration.
0.030 i 0.001 0.031 i 0.002 0.030 -t 0.001 0.031 t 0.001 0.032 i. 0.002 0.032 i 0.001
TIME ( m i n i
Figure 3. Rate of color development and decay: West-Gaeke vs.
present method. working reagent, leading to -97% of the maximum observed sensitivity and a relatively low blank value of 0.030 AU. Any marginal increase in sensitivity at the expense of the increasing value of t h e blank was considered undesirable. We were unable to account for the decrease in sensitivity for pararosaniline concentrations beyond that contained in 8 mL of the working reagent. In view of the arguments presented for ideal mixing conditions, delivering the same amount of pararosaniline and acid contained in 5 m L of the working reagent from a smaller volume of a more concentrated solution was considered. However, experiments showed a marginal or no improvement from the described procedure, presumably due to decreased precision of delivering smaller amounts and loss of HC1 from considerably more concentrated solutions. Effect of the formaldehyde concentration itself was considered important in view of possible losses from the absorber solution during sampling. Subsequent work, based on an ion chromatographic determination of HCHO as formate (via oxidation with Hz02in an alkaline medium) with borate eluant (50), showed that no significant loss of HCHO occurs from a n absorber solution (15 mL) upon sampling for 8 h a t 250 mL/min a t 25 "C. Presumably, hydration to the diol lowers the vapor pressure. I n any case, varying the formaldehyde concentration, within limits, did not significantly affect the results; the pertinent data are presented in Table VI. Speed of Color Development and Stability of Developed Color. Compared to the WG procedure, this method leads to faster development of color, which also decays more rapidly. Figure 3 shows the respective plots of absorbance vs. time for the WG procedure and this method. The respective absorbances, measured at 560 and 580 nm, for solutions containing 1 pg/mL SO2 in the final volume, are plotted in arbitrary units as a function of time. In this method. maximum color intensity is reached in about 10 min and no significant change occurs until 15 min. Subsequently the
intensity decays approximately with a first-order rate of 0.2% /min. In contrast, the U'est-Gaeke method reaches maximum color intensity in about 30 min, continues undiminished to about 50 min, and then decays approximately with a first-order rate of 0.07 70 /min. Extensive spectrophotometric investigations on the rate of color development and decay were carried out for this method. In essence, these results indicated that within the range of interest, the rate of color development was facilitated by an increasing pararosaniline concentration and, to a lesser extent, by increasing acidity. The rate of decay was almost solely governed by the p H of the final solution, decaying faster with decreasing p H For most laboratories analyzing a large number of samples, the 5-min flexibility in the measurement period was considered adequate, as long as the samples are properly sequenced. Automated applications should present no problems. If desired, measurements over a larger peIiod of time can be made by decreasing the acidity of the final mixture and sacrificing the sensitivity to some degree. Alternatively, for periods up to a t least 1 h after mixing, fairly accurate results can be obtained by applying decay corrections with the listed rate of decay, provided the samples are thermostated at 25 "C. The decay rate is temperature dependent and increases with increasing temperature. The advantage of a faster development, in our opinion, adequately offsets the disadvantage of lower stability. We have considered the possibility that the decay is solely due to the loss of SO, from the solution. the system having reached a state of dynamic equilibrium. Large increases in the decay rate can be brought about by increasing temperatures and acidity or by bubbling a purging gas through the solution. For the WG method, the decay rate decreases with increasing TCM Concentration. Although these observations can be cited in support of the above hypothesis, there is no conclusive proof to support this contention. Effects of Ozone and Nitrogen Dioxide. Ozone, 0.8, 1.5, and 2.4 ppm, was sampled through an absorber (15 mL) containing the equivalent of 25 pg of SO2 at 250 mL/min for periods of 8, 4,and 3 h, respectively Triplicate sample sets were collected and allowed to remain 20-30 min before analysis, after Scaringelli et al. (35).h'o significant interference was found. Nitrogen dioxide, 0.6, 0.9, 2.5, and 5.9 ppm, was sampled through an absorber (15 mL) containing the equivalent of 25 Fg of SOz at 250 mL/min for periods of 2 , 4 , 8 , and 16 h. Six samples were collected in each set; three were analyzed as such and the remaining were analyzed after the addition of 1 mL of 0.6% sulfamic acid (33). For the samples to which no sulfamic acid was added, the correlation between the total amount of the NO, sampled and the decrease in color intensity from the standard value was less than ideal. The largest decrease observed was -30% from the standard value. Sampling 5.9 pprn of NOz at 250 mL/min for 16 h represents nearly 3 mg of NO2 passing through the absorber. Yet, as little as 0.1 mg of nitrite added as NaNOz to the sample containing 25 pg of SO2totally eliminated any color. Evidently, the conversion efficiency of gaseous NO2 to nitrite by this particular absorber is poor; the action of NOz on solutions of S(IV) is known to be quite complex (51-53). For the samples to which sulfaniic acid was added, the results from all the sets were uniformly 556% lower than the standard value. Experiments were conducted to determine effects of sulfamic acid on a standard calibration plot, and it was established that with or without the presence of nitrite, 6 mg of added sulfamic acid decreases the slope of the regression line by 5.3%. Scaringelli et al. (35) have previously argued, for the T C M method, that this suppression of color intensity is due to a decrease in p H as a result of the added acid. This system is already acidic to such an extent that the small amount of sulfamic acid added cannot cause a significant
ANALYTICAL CHEMISTRY, VOL. 52, NO 12, OCTOBER 1980
change in the final p H , and none was discernible by actual measurement. In agreement with Zurlo and Griffini (31) and Pate et al. (33),we find that all compounds containing primary or secondary amino groups, including ammonia, lower the sensitivity of the method, presumably due to competing reactions. We also found that the sulfamic acid addition should be carried out immediately after sampling. If the sample, containing dissolved NO2, is allowed to remain for any length of time prior to analysis, further degradation occurs. The same phenomenon takes place with TCM as absorber and to a greater extent. If the sample must be stored prior to analysis, sulfamic acid should be added before storage and preferably as the sodium salt. This prevents any major changes in the p H of the absorber that may lead to some desorption loss. Recognizing that sampling NO, through an absorber containing SO, is not exactly equivalent to sampling a mixture of both gases, we designed an experiment involving 0.025 ppm of SO, and 2.5 ppm of NOz, which is representative of the most adverse conditions that may possibly be encountered. Parallel sampling was conducted by using both formaldehyde and T C M as absorbers (15 mL, 250 mL/min, 15.5 h). In the absence of NO,, the mean results were in excellent agreement (ppm SO,: 0.025 F P D , 0.0244 WG, and 0.0248 present method). In the presence of 2.5 ppm NO, (2.45 ppm CL) and with sulfamic acid treatment, the formaldehyde absorber showed a 5.2 1.5% decrease which could be attributed to t h e sulfamic acid treatment. Without sulfamic acid, an average decrease of 31% was observed. The TCM samples averaged a 25% decrease, with sulfamic acid treatment. We believe that the T C M absorber is more susceptible to NO2 interference primarily because of its higher p H and the resulting higher solubility of NO2. If t h e present method is applied to samples collected in locations known to have significant concentrations of NO,, sulfamic acid/sodium sulfamate addition should be incorporated into the analytical procedure. An appropriate calibration procedure must, however, be used. T h e method, in general, is considerably less susceptible to NOz interference than the U‘G method. With low levels of NOz therefore, the sulfamic acid addition may be omitted without significantly affecting the results. Effect of Transition-Metal Ions. The effects of Fe(III), Cu(IIj, V(V), and Mn(II), all of which are known to catalyze the aerobic oxidation of S(IV) ( 5 4 ) , were investigated. In separate experiments, to 15 mL of the absorber containing 25 pg of SO2, Fe(III), Cu(II), and Mn(I1) were added as chlorides and V(V) as NaV03. The solution was then aerated for 8 h a t 250 mL/min and analyzed. For Fe, Cu, and V (highest amounts added: 100, 100, and 200 pg as the metal, respectively) no significant interference was discernible. Ion chromatography of the aerated samples also indicated the absence of sulfate. Manganese, however, interfered severely, as little as 10 pg of Mn totally suppressed the formation of the purple compound. A bright yellow color (A,, 426 nm), decaying with time, was formed instead, and experiments showed that the formation of this color was not related to the presence of HCHO, K H P , or S(1V). Ion chromatography of the aerated samples did not show any perceptible amount of sulfate, clearly establishing that the interference was not due to any oxidative degradation of the collected S(1V). As such, this interference has no parallel in the WG method. Mn(I1) in neutral or acidic solutions does not react with acid-bleached pararosaniline. It is only when manganous hydroxide, precipitated from solution by the addition of strong base, is reacidified with the pararosaniline reagent, that the yellow compound forms. The reaction appears to be a redox reaction resulting in the oxidation of pararosaniline to an unstable quinone imine. Alkaline hypochlorite produces the same
[[email protected][email protected]
@$a 0 “2
Figure 4. Structure of pararosaniline and related compounds. yellow compound when added to pararosaniline, the latter may in fact be substituted for o-tolidine in the well-known reaction used for the determination of dissolved Cl,, with a lower sensitivity. Freshly precipitated Mn(OH), quickly absorbs oxygen to produce a compound capable of oxidizing pararosaniline to the yellow quinone imine. Accordingly, no interference due to Mn(I1) is observed when the whole reaction is carried out in the absence of oxygen. Interestingly, if the precipitated Mn(OH)2is fully oxidized to hydrated MnO,, by the addition of H 2 0 2 ,it is no longer capable of oxidizing pararosaniline to the quinone imine. Attempts to eliminate the interference due to Mn(I1) by the addition of EDTA were only partially successful. Color development occurred in the presence of EDTA, but even with a concentration of 1.8 m M of Na,EDTA in the absorber, 10 pg of Mn(I1) produced a negative interference of about 20% in the determination of 25 pg of SO2. (trans-1,2-Cyclohexylenedinitrilojtetraaceticacid, CDTA, which has been reported to be a substantially better chelating agent for Mn(I1) than EDTA ( 5 5 ) ,was studied for masking the interference due to Mn(I1). A concentration of as little as 0.1 mM of the reagent in the absorber totally eliminated any interference due to 10 pg of Mn in the determination of 25 pg of SO,; interference due to 20 pg of Mn was reduced to less than 3%. Studies showed that the results were the same whether the CDTA was incorporated into the absorber itself or added subsequent to sampling, confirming that the interference due to Mn(I1) does not take place via the oxidative degradation of S(1V). While the results reported in this paper do not pertain to an absorber containing CDTA, we have confirmed that the presence of CDTA does not affect the method sensitivity or the blank value. We recommend the inclusion of 0.1 mM CDTA in the absorber solution. The reagent should be added as the disodium salt, in which case the p H of the absorber remains virtually unchanged. Enhancement of Sensitivity by Solvent Extraction. The sensitivity of the present method or that of the WG method can be increased by an order of magnitude by extracting the product into 1-butanol or 1-pentanol. The latter is preferable because of its lower solubility in aqueous media. Some unreacted pararosaniline is also extracted and contributes to the blank. The color of the dye cation (D+j in pararosaniline (Figure 4a) is due to the chromophore in the canonical form shown in Figure 4b and its two other equivalent canonical forms. Similar cations of other triaminotriphenylmethane dyes, e.g., crystal violet (N,N’,N”-hexamethylpararosaniline), display a red-violet to blue-violet color. Protonation of one of the nitrogen atoms in D+ to yield DH’+ results, in the case of crystal violet, in a green color. Only two canonical forms of such a structure can be drawn, and the color is similar to that of a dye cation (D+)of a diaminotriphenylmethane dye, e.g., Malachite Green (N,N’-dimethyl-p,p’-diaminotriphenylmethane). Protonation of a further nitrogen atom to yield DH,?+ causes a further blue shift in absorption and a decrease
ANALYTICAL CHEMISTRY, VOL. 52, NO.
12, OCTOBER 1980
Table VII. Comparative Studies: West-Gaeke vs. Present Method
no. of 1
4 5 6 p 5~
PO^ o r o s o n ,I I p e - HC H O - s ( I V ) r e o c I 10" p r o d u c t ( 1 ), , ,A 560nm @ p H 5
r e d s h i f t s w i t h d e c r e a s i n g p H 1 I n b o f h cases
Flgure 5. Color intensity vs. pH: pararosaniline and pararosanilineHCHO-S(1V)
in absorptivity and the color, for pararosaniline, rosaniline (which contains a methyl substituent ortho to the amino group in one of the p-anilino groups of pararosaniline), crystal violet, and other related dyes become yellow, similar to the color of the dye cation of a monoaminotriphenylmethane dye (56). Often, D+ and DHZ3+are resonance stabilized relative to DH2+ t o such an extent that a distinct stage with the predominance of DH2+and the attendant green color is never observed, and the intensely violet dye gradually becomes a weak yellow with increasing acidity; such is the case for pararosaniline and rosaniline (57). An organic solvent such as 1-pentanol is able to extract the unprotonated dye as the ion pair D+X- but not t h e protonated dye. Since the product of the pararosaniline-HCHO-S(1V) reaction is considerably more acidic than pararosaniline (Figure 5 ) the product is preferentially extracted into the organic solvent. However, a typical reaction mixture in the WG method contains a much higher concentration of unreacted pararosaniline relative to the amount of the product formed (and even more so in the present method) resulting in too high a blank value to be of any practical use. Extraction of the reaction mixture (25 mL) with 1-pentanol (2 mL) subsequent to the recommended color development periods for the respective methods and measurement of absorbance of the organic layer results in an enhancement of sensitivity by well over 1 order of magnitude, b u t the absorbance blank itself under these conditions is nearly 2 AU, making measurements impossible by all but the most sophisticated instruments. However, once extracted into the organic solvent, the color intensity appears to be stable for long periods of time, at least 1 week. The added stability, combined with the enhanced sensitivity, may find use in special cases. While we have not optimized the procedure, we find t h a t the best way to circumvent the problem of high blanks is adding acid to the extracted layer prior to measurement; sensitivity is only marginally sacrificed. Adding solid trichloroacetic acid is ideal, since no dilution is involved. Comparative Studies: West-Gaeke vs. Present Method. Side by side parallel samples were drawn into 0.1 M T C M and the buffered formaldehyde absorbers and analyzed by the respective methods. T h e sampling conditions were as given in Table I1 and the results are shown in Table VII. At low concentrations of SO, there is a consistent trend for the present method to yield values slightly higher than the WG method. Within the limits of accuracy of our data, the difference could be due to a number of possible reasons, namely, calibration error, inaccuracy of flow measuring devices, and a n actual difference in absorption efficiency. We do not believe the last one t o be responsible because, under these sampling conditions, we were never able to detect any SO, in the back-up bubblers for either absorber. However, the
8 8 10
0.02 0.02 0.05 0.05
4 2 2 4 4
SO, concn, ppm this method ( + S D ) ~~
set no. expts FPD
0.23 0.23 0.52 0.52
WG ( r S D ) 0.009
0.008 i: 0.003 0.019 i 0.002 0.018 i 0.003 0.052 r 0.007
0.052 f 0.107 f 0.105 i 0.224 *
0.009 0.012 0.015 0.028
0.545 0.538 t 1.176 i 1.152 t
0.061 0.015 0.029 0.099
0.020 0.020 0.055 0.052 0.105 0.101 0.221 0.229
i t ir t f i
0.002 0.002 0.005 0.007 0.015
0.025 0.029 0 . 5 4 1 i 0.009 0.529 f 0.022 1 . 1 6 2 I0.086 1.148 + 0.053 f
M E C H A N I S M OF THE R E A C T I O N A N D S T R U C T U R E ( S ) OF THE P R O D U C T ( S ) data clearly show that the present method is capable of yielding results comparable in accuracy and precision to that of the WG method. The mechanisms of the Schiff staining reaction as practiced in histochemistry and that of the complementary reaction involved in the determination of SOz have been controversial for years. In the staining reaction, the tissue under investigation is treated with Schiff reagent which consists of S O z / bisulfite/pyrosulfite bleached rosaniline/pararosaniline. When a suitable nucleophile such as the bisulfite ion is present in significant concentration, it adds to the central carbon atom of the dye cation, leading to loss of the chromophoric canonical forms and consequent loss of color (Figure 4c). A similar loss of color occurs a t high p H for this class of dyes due to the formation of the carbinol (Figure 5 ) . Bisulfite addition has been clearly established to result in bonding via sulfur rather than oxygen (58)leading to a sulfonic acid rather than a sulfite ester. Schiff reagent is thus pararosanilinesulfonic acid. Previously it has been shown on the basis of spectral studies that there is no significant spectral difference between t h e products of the pararosaniline-aldehyde-S(1V) reaction and those of the Schiff reagent-aldehyde reaction (59). In this investigation thin-layer chromatography of butanol/pentanol extracts of the reaction mixtures for (a) the WG method, (b) the present method, and (c) the Schiff reagent-formaldehyde reaction showed that aside from the starting reagents, two components with respective R, values of -0.4 and -0.6 were present in all cases. With increasing concentration of formaldehyde, a third compound with a R f value of -0.8 was formed a t the expense of the slower moving component (R, 0.4) in (b) and (c). In subsequent discussions, these compounds are referred to as I , 11, and I11 respectively. This experiment established that the same products are formed in all the reactions, presumably through similar mechanisms. Previous studies with 35S-labeled Schiff reagent (60) also showed only three colored components with which radioactivity was associated. A fourth component, radioactive but colorless, was also present and could be the Schiff reagent itself or the formaldehyde-bisulfite adduct. As many as ten components, however, have been reported to form in the Schiff reagent-formaldehyde reaction when the formaldehyde concentration is increased to large amounts (61),but this study used basic fuchsin as the starting material for t h e Schiff reagent and the conditions had little similarity with the reactions of our interest. The debate on the nature of the product(s) formed in t h e Schiff reaction dates back to the 1860's. T h e idea that the original dye is regenerated was discarded because of an easily discernible difference in hue between the reaction product and the original dye. Schiffs original idea ( I ) that the color is due
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980
to the formation of a Schiff base was discarded because the presence of sulfur in the product could be easily established. An elaborate study by Wieland and Scheuing (62) led to the sulfinic acid theory which invoked that bisulfite attack on the amino nitrogens of pararosaniline resulted in N-sulfinic acid (-NHS02H) formation a t two of the nitrogen atoms. One or two molecules of formaldehyde subsequently reacted with one or both of the N-sulfinic acid groups to yield -"-SO2CH2OH moieties. In the last major review of the chemistry of the Schiff reaction, this theory was held to be valid (5). The alkylsulfonic acid theory originated even earlier (63) and the theme t h a t one or more aminomethanesulonic acid (-NHCHzS03H) groups are formed as a result of the reaction was repeated by a number of workers (64-66) even prior to the work of Nauman et al. (36). Damianovitch (63),Rumpf (65), and Nauman et al. thought that the aldehyde-S(1V) addition compound was first formed followed by reaction with the amine. Bucherer (64) did not elaborate on the mechanism beyond a ternary reaction between the amine, aldehyde, and bisulfite. Hormann e t al. (66) suggested a carbinolamine (-NHCH,OH) as an intermediate, formed by addition of the amine to the aldehyde. Further spectral studies (59) also appeared in the literature in support of the alkylsulfonic acid theory. Secondary arguments involve the charge and the degree of protonation of the product dye cation (i.e., D+ or DH2+,etc.) and the extent to which substitution of the amino groups occurs (34, 35, 66). Of all the studies dealing with the structure of the product(s) of the Schiff reaction, the work of Nauman et al. is generally considered t o be chemically the most elegant (68). What seems to be clearly established is that the reappearance of color is intimately associated with the basicity of the nitrogen atoms in pararosaniline. In the case of the WG reaction or the present method, the pH of the system is sufficiently low to eliminate any color due to the D+ cation and only the DHz3+form is present. Any substitution of the amino hydrogen($ which decreases the basicity of the nitrogen atom(s) will regenerate the D+ type cation and lead to the reappearance of the red-violet/blue-violet color. In the case of the Schiff reaction itself, the decreased basicity of the nitrogen atom(s) is expected to decrease the electrophilic character of the central carbon atom leading to elimination of the nucleophile (bisulfite) attached t o it and reappearance of the D+ type cation. Thus, Stoward (69) has correctly argued that from visible absorption spectra alone it is impossible to conclude whether the product has one or more -NHCH2S03H, - N H S 0 2 H , -NHSO2CH20H, or even -N=CH2 groups. Our efforts to obtain enough pure product(s) for unequivocal structural assignments via NMR studies were unsuccessful primarily due to the fact that macroscale synthetic attempts resulted in compounds other than those formed in the analytical reactions. These compounds had very low R, values in the thin-layer chromatograms, were insoluble in most solvents, and contained little or no sulfur. We believe these compounds t o be the Schiff bases, quite possibly the same compounds observed by Hiraoka (61) with high formaldehyde concentrations. These compounds predominate whenever pararosaniline and formaldehyde are used a t concentrations substantially higher than those used under the analytical conditions; sulfite may play a merely catalytic role under such conditions as observed by Hartough e t al. (70) for aminomethylation of thiophenes. However, we believe that our results combined with those of some recent studies will be sufficient to provide an unambiguous solution to the problem. Structure. The charge or the degree of protonation of the product dye cation can of course be assigned simply from the color of the product to be D+ or unprotonated. Further, N,N',N"-pentamethylpararosaniline can be substituted for
pararosaniline in all three methods of our interest. The color of this dye is blue-violet for D', green for DH2+,and yellow for DH23+as detailed in a previous section. T h e products in all three cases have a blue color and can be made green and yellow with progressive addition of acid. We discard the N-sulfinic acid theory for a number of reasons: (a) Such compounds are unknown in aqueous solution; the UV-visible spectrum of a solution containing microgram amounts of bisulfite and the pararosaniline working reagent is identical with the sum of the spectra of the individual components. (b) Thiols, alkali sulfides, and thiosulfates all give a reaction virtually identical with that of sulfites, when these compounds are treated with pararosaniline and formaldehyde. This is impossible to explain in terms of the sulfinic acid theory unless radically different structures and mechanisms are proposed. (c) It has recently been shown by Mirek and Rachwal (71) that even sulfinylamines (RN=S=O) do not react with aldehydes to yield a N-S-C linkage. In aprotic solvents only slow formation of the Schiff base takes place, accompanied by loss of SO2. With the addition of water and especially catalyzed by acids, crystalline precipitates appear. These compounds have been shown, on the basis of detailed NMR and IR studies, to be amine salts of aminoalkanesulfonic acids. R-N=S=O
H2O + R'CHO 7 RNHCHR'SO3-RNH3'
T h e cation is easily exchangeable with those derived from other amines. (d) If a N-sulfinic acid of pararosaniline were to be formed a t all, one would expect reappearance of color even without the benefit of further reaction with an aldehyde, because of decreased basicity of the nitrogen atom(s). We regard the alkylsulfonic acid theory to be valid because: (a) a-Aminoalkanesulfonic acids have been prepared by the reaction of primary and secondary amines with aldehydes and S(1V) or by heating the amines with the aldehyde-bisulfite addition compound. In the case of primary amines, an identical a-aminoalkanesulfonic acid results if the corresponding Schiff base is treated with bisulfite (72-74). (b) The mono(azomethine) Schiff base of pararosaniline, (p-NHzPh)2+CPhN=CH2,reacts with S(1V) in acid solution to give a single product which thin-layer chromatography shows to be identical with product I isolated from the analytical reactions. (c) The conjugate acid of pararosaniline displays a pK, of 2.7 (Figure 5 ; 2.58 from ref 57) whereas the conjugate acid of product I shows a pK, of less than I, the difference is consistent with the formation of an aminoalkanesulfonic acid (73). Regarding the degree of substitution, if all six amino hydrogens of pararosaniline were replaceable, one can calculate the possible formation of nine different products (one each of mono, penta, and hexa substituted pararosaniline and two each of di, tri, and tetra substituted pararosaniline). Only three compounds are found indicating only one of the hydrogens associated with each amino group of pararosaniline is replaceable. We assign therefore the structure N-(methylsulfonic acid)pararosaniline to I, substantiated by its formation from the mono(azomethine) Schiff base of pararosaniline and S(1V). Correspondingly, I1 must be N,N'bis(methylsu1fonic acid)pararosaniline. I t has classically been assumed that only I is formed in the WG reaction since a large excess of pararosaniline is present (34,36). We find however that a t the S O z level of 1 pg/mL, both I and I1 are formed. From the intensity of the spots on thin-layer chromatograms, it appears that substantially more I than I1 is formed in the WG method, whereas in the present method the relative amounts appear to be nearly equal. This is to be expected
ANALYTICAL CHEMISTRY, VOL. 52, NO. 12, OCTOBER 1980
since the absolute concentration of free S(1V) in solution is much less in the WG method due to the presence of TCM. For either method, variation of SOz concentrations between 0.1 and 1 pg/mL does not result in any easily discernible difference in the relative amounts of I and I1 formed. For both methods, however, a decreased formaldehyde concentration results in the formation of more I relative to I1 and increasing the formaldehyde concentration results in the formation of more I1 relative to I. With the present method, I11 begins to be formed with a further increase in formaldehyde concentration; we consider I11 t o be the trisubstitued compound. Mechanism. Insofar as the mechanism of the sulfomethylation of the amine is concerned, reaction of the amine with the aldehydebisulfite addition product must be rejected. As we have noted, no reaction takes place between sodium hydroxymethanesulfonate and acid-bleached pararosaniline. At the time we undertook this study, we were unaware that Wieland and Scheuing (62) had made the same observation, which in fact was the main impetus behind the N-sulfinic acid theory. We have considered the possible participation of some intermediate formed during bisulfite addition to formaldehyde. This seems unlikely since the electrophilic character of the carbonyl carbon can only decrease subsequent to bisulfite attack. Further, thiols and sulfides, which undergo the same reaction as sulfite, also do not react if formaldehyde is added first t o the above compounds and sufficient time is allowed for the formation of the corresponding addition compound before the addition of pararosaniline. Thiosulfate, which does not add to a carbonyl group, reacts normally; the slow rate of color development and the spectrum of the reaction product indicate that the reaction actually occurs via the acid hydrolysis of thiosulfate to sulfite. In the literature reports of the syntheses of a-aminoalkanesulfonic acids from an amine and an aldehyde-bisulfite addition compound, either the latter was generated in situ or heating was necessary which presumably permitted sufficient dissociation to occur. Accordingly, we found that pararosaniline and sodium hydroxymethanesulfonate react to a small extent upon warming. This leaves the addition of the amine to the aldehyde as the only other logical alternative for the first step of the reaction. I t is obvious from increasing reagent blanks due to increasing concentrations of formaldehyde, even when pH is strictly controlled (see also Figure 3, ref 15),that a reaction takes place between pararosaniline and formaldehyde. The first step in this reaction would be the formation of the carbinolamine as elucidated by Jencks (75). The following steps can now be visualized: (a) elimination of water to form the Schiff base followed by (b) addition of bisulfite to the C=N bond or (c) direct nucleophilic attack of bisulfite on the carbinolamine leading to elimination of OH-. The reaction of our interest takes place only in strongly acid media, virtually no reaction takes place above a p H of 2. However, no definite conclusions can be drawn from this observation since both steps (a) and (c) are acid catalyzed. T h e rate of color development for the reaction between the mono(azomethine) Schiff base of pararosaniline and bisulfite was studied under the same conditions as the analytical reaction and found t o be significantly slower compared to the analytical reaction. T o further determine whether Schiff base formation plays any role in the mechanism of the reaction whatsoever, the following experiments were conducted under argon: (a) 1 mL of the pararosaniline reagent was treated with 1 mL of 70 m M HCHO and allowed to remain for 5 min, followed by 25 pg of SO2 as Na2S03; (b) 1 mL of the pararosaniline reagent was treated with 1 mL of a solution containing 25 pg of SOz as NaZSO3and allowed to remain for 5 min, followed by 1 mL of 70 mM HCHO; (c) 1 mL of the pararosaniline reagent was treated with 1 mL of $0 mM
_ _ _ l o i Parorasaniline
a n d H C H O . S u l f i l e 5 m i n loler l b i P a r a r o s o n l l i r l e a n d s u l f i t e , H C ~ O5 m n l o t e r and I C ) P o r a r o s a n ~ l i r i ea n d HCHO, S u l f i t e i m m e d i a t e l y
Figure 6. Rate of color development and the order of reagent addition.
HCHO immediately followed by 1 mL of a solution containing 25 pg of SO2 as Na2S03. In each case the volume was made up to 25 mL and rate of color development was followed a t 560 nm. The results are shown in Figure 6. There was no difference between (b) and (c), shown therefore as a single curve, while (a) was decidedly slower during the initial part of the reaction and made up the difference to some degree toward the end. Results of these experiments clearly favor participation of the carbinolamine rather than the Schiff base. The reaction of thiols and sulfides provides additional evidence, these reactions are 2 to 3 times faster than the corresponding reactions with sulfite a t the same pH; we attribute this to the greater nucleophilicity of RS- and HS- compared to HSO,-. Conclusive evidence in favor of nucleophilic substitution is given by the compound N,N’,N”-pentamethylpararosaniline which cannot form a Schiff base but yet undergoes the Schiff-type reaction to give a single product. We suggest therefore the following mechanisms: (A) For the West-Gaeke reaction and the present method (NH2Ph),C+
(B) For the Schiff staining reaction (NH2Ph),CS03H
(NH2Ph),C ( S O , H ) P h N H C H ( O H ) R (NH,Ph)ZC(SO,H)PhNHCH(SO,H)R
A second attack a t the same nitrogen does not take place apparently due to its decreased basicity. If strongly basic aliphatic amines are used, a second addition reaction occurs easily (73). Absorptivity of the Reaction Product. The absorptivity of tris(paminopheny1)methane dyes appears to be directly related to the basicity of the amino groups. The more basic the amino group, the greater is the stability of the chromogenic canonical form(s) shown in Figure 4b. Accordingly, N,N’,N”-hexamethylpararosaniline (crysta1 violet) has an tmaxmore than 1 order of magnitude higher than the weaker base pararosaniline. The emax for the methylsulfonic acid derivatives of pararosaniline should be lower than pararosaniline inasmuch as these compounds must have basicities considerably lower than pararosaniline and should in fact decrease with progressive substitution. Corresponding reaction products with thiols and sulfides should display values of t,, between those of pararosaniline and its methylsulfonic acid derivatives. The puzzling fact is that if the different reactions are assumed
ANALYTICAL CHEMISTRY, VOL. 5 2 , NO. 12, OCTOBER 1980
to be completed to the same extent, the emax values for the products obtained in the reactions with thiols and sulfides are a t least 1 order of magnitude lower than that obtained with sulfites. Even more puzzling are the results of the calculations made by Huitt and Lodge ( 3 4 ) which state that the e, for the tris(methylsu1fonic acid) is nearly 50% larger than that of the bis(methylsu1fonic acid), more than 4 times that of the methylsulfonic acid and is, in fact, much higher than pararosaniline itself. The only way to reconcile these observations is to assume that the methylsulfonic acid derivatives actually exist in solution to a large degree as the methylsulfonate anion and increasingly so with increasing substitution. In such a system, in addition to the canonical form =+NHCH2S03H, the canonical zwitterion =+NHCH2S03 may exist in another chromophoric form =NCH2S03H leading to a significant increase in the absorptivity. This hypothesis also explains the hitherto unexplained increase in sensitivity of the WG reaction upon the addition of dimethylformamide ( 3 4 ) . The basic solvent promotes a greater degree of ionization for the methylsulfonic acid, thus increasing absorptivity. CONCLUSIONS Substitution of a nontoxic buffered formaldehyde absorber for T C M should result in an impetus toward the continued use of the already established West-Gaeke method, in this modified version, for the determination of atmospheric SOz. Results from the new method are comparable in every respect to those from the established procedure. Investigation into the mechanism of the Schiff reaction and the structure of the products establishes t h e validity of the alkylsulfonic acid theory. An aminocarbinol appears to be the first intermediate which undergoes nucleophilic substitution by bisulfite. I t appears that substitution of other related dyes for pararosaniline could substantially increase the sensitivity of the method. Studies with dyes such as N,N',N"-pentamethylpararosaniline which display a n absorptivity much greater than pararosaniline or bis(paminopheny1)formaldimine hydrochloride, (p-NH2Ph)2C=+NH2C1 (Auramine 0),which has been reported in cytochemical studies to produce a fluorescent product ( 5 ) ,should prove particularly interesting. ACKNOWLEDGMENT T h e technical help of B. K. Tarkington and T. K. Duvall is gratefully acknowledged. Discussions with G. K. Newkome, Louisiana State University, Baton Rouge, were helpful. LITERATURE CITED Schiff, H. Ann. Chem. Pharmacol. 1866, 740, 92. Schmidt, J. G. Ber. Dtsch. Chem. Ges. 1880, 13, 2342. Bauer, H. 2. Mikrosk.-Anat. Forsch. 1933, 33, 143. McManus. J. F. A. Nature (London), 1946, 158, 202. Kasten, F . H. I n t . Rev. Cytol. 1960, 10, 1. Pearce, A. G. E. "Histochemistry", 3rd ed.; Little Brown: Boston, 1968: Vol. 1, Chapter 13. Steigmann, A. J . SOC. Chem. Ind., London 1942, 6 1 , 18. Grant, W. M. Anal. Chem. 1947, 19, 345. Kozlyaeva. T. N. Zh. Anal. Khim. 1949, 4 , 75. Atkin, S. Anal. Chem. 1950, 22, 947. Haller, P. J . Soc. Chem. Ind.. London 1919, 38, 52T. Urone, P. F.: Boggs, W. E. Anal. Chem. 1951, 23, 1517. Pauius, H. J.;Floyd, E. P.; Byers, D. H. Am. Ind. Hyg. Assoc., 0.1954, 15, 4. Moore.. G. E.:. Cole.. A. F. W.: Katz. M. J . Air Pollut. Control Assoc. 1957. 7, 25. West, P. W.; Gaeke, G. C. Anal. Chem. 1956, 28, 1816. Feigl. F. "Chemistry of Specific, Selective and Sensitive Reactions"; Academic Press: New York, 1949; p 75. Burke, K. E.; Davis, C. M. Anal. Chem. 1962, 34, 1747. Barabas, S.;Kaminski, J. Anal. Chem. 1963, 35, 1702.
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RECEIVED for review May 29, 1980. Accepted July 25, 1980. This research was supported by a contract from the Electric Power Research Institute (RP 1112-2).