Anal. Chem. 1980, 52, 2177-2184 (18) (19) (20) (21) (22j
KovBts, E. Helv. Chim. Acta 1958, 4 1 , 1915-1932. McGill, J. R.; Kowalski, 8.R. J . Chem. Inf. Comput. Sci. 1978, 78, 5 2 . Justice, J. B.; Isenhour, T. L. Anal. Chem. 1974, 4 6 , 223. Schechter. J.: Jurs. P. C. ADD/. Soectrosc. 1973. 27. 223. Kowalski, B. R.; Jurs, P. C.; isenhour, T. L.;Reilley, C. N. Anal. Chem. 1989, 4 7 , 695. (23) Ritter, G. L.; Lowry. S.R.; Wiikins, C. L.; Isenhour, T. L. Anal. Chem. 1975, 4 7 , 1951.
2177
(24) Jurs, Peter C.;Isenhour, T. L. "Chemical Application of Pattern Recognition"; Wiley: New York, 1975.
RECEIVED for review March 24, 1980. Accepted August 11, lg80* The authors express gratitude for the support Of the Office of Naval Research.
Fluorometric Reaction Rate Method for the Determination of Thiamine Mary Andrieu Ryan' and J. D. Ingle, Jr." Department of Chemistty, Oregon State University, Corvallis, Oregon 9733 1
The reaction involving the oxidation of thiamine by Hg2+ in basic solutions to fluorescent thiochrome was studied. The initial rate of the reaction, measured as a change in fluorescence signal over 10 s, is iineariiy related to the thiamine concentration from a detection iimlt of 2 X lo-* M up to 1 X lo-" M. The applicability of the new method for determinatlon of thalmine in vitamln-mineral preparations is demonstrated by investigation of the effect of potential Interferences and by analysis of commercial and synthetic preparations.
In the years since the discovery and isolation of thiamine (vitamin BJ, voluminous amounts of literature on its determination cite the use of microbial, chromatographic, direct molecular absorption, and equilibrium-based colorimetric and fluorescence methods (1-3). No reaction rate method for the specific determination of thiamine has been reported even though reaction rate methods may provide improved precision, speed of analysis, and specificity compared to equilibriumbased methods. T h e general advantages of reaction rate methods, in particular with fluorescence monitoring, have been discussed (4, 5). Several of these advantages are especially important with respect to the thiamine determination. Specificity and precision due to sample to sample variability may be improved because only those species which react and cause a change in the monitored fluorescence signal are measured so that steady-state signals like scattering or background fluorescence do not interfere. Increased specificity will make some separations and extractions unnecessary thus providing considerable time savings. Since nonreproducibility, contamination, and loss of analyte are a problem with separations, fewer separations may also result in improved precision. The new rate method presented here has the additional advantages of using fewer and more stable reagents and a simpler analytical procedure. The standard method for the assay of thiamine is an equilibrium-based method called the thiochrome method and involves the oxidation of thiamine to fluorescent thiochrome (TC) (6). In the analysis of pharmaceutical preparations, an alkaline potassium ferricyanide oxidant solution is added to the thiamine sample solution (7). T h e sample solution is shaken with 2-butanol, the 2-butanol layer is withdrawn, and
the fluorescence of TC in the butanol is measured. The blank is the sample treated with all of the reagents except the oxidizing agent. Samples other than pharmaceuticals may require enzymatic hydrolysis of thiamine pyrophosphate to the free thiamine before the oxidation step. The sample must then be filtered and the thiamine removed by an ion exchange column, in order to separate the thiamine from substances which would also consume the K,Fe(CN)6 and from fluorescent interferences. The thiamine is eluted with a hot acidic potassium chloride solution, and the oxidation and extraction are carried out as above. There are several critical points in the standard procedure. The initial purification ion exchange step has been thoroughly studied (8, 9). The temperature and the volume of the potassium chloride solution affect the percentage of thiamine recovered, and usually some of the thiamine is lost. T h e amount of potassium ferricyanide must be sufficient to oxidize the thiamine, but a large excess is undesirable since it may result in the decomposition of TC (I). The presence of other oxidizable materials in the sample complicates the choice of the proper amount of oxidant (10). The yield of the reaction is about 67% ( 3 ) . The extraction of the TC with 2-butanol is necessary to separate the TC from other fluorescent substances and from potassium ferricyanide which quenches the thiochrome fluorescence (11). The addition of the oxidizing reagent, the mixing, and the extraction with 2-butanol must be performed quickly and be carefully standardized to be reproducible (12). This paper discusses the development of a rate method of analysis in which thiamine is oxidized by Hg(1I) in a pseudo-first-order reaction which produces fluorescent TC. The reactions of thiamine in basic solutions and in the presence of Hg(I1) are discussed and related to the empirical results. T h e variables important in any new rate method (13) were studied and are presented here. The equilibrium and kinetics behavior of the reaction is examined in terms of reagent concentrations and other experimental variables. The nature of the compromises involved in optimizing the reaction for a kinetics-based method are presented since different applications may dictate a set of optimum conditions different from those obtained here. T h e results of an extensive interference study and the application of the method to determination of thiamine in multivitamin-mineral preparations are reported.
EXPERIMENTAL SECTION The instrument used in this study was a sensitive spectrofluorometer (5)with a fixed-time digital ratemeter (14) specifically
'Present address: Tektronix, Inc., P.O. Box 500, Beaverton, OR
97077.
'0
0 1980 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 52, NO. 13, NOVEMBER 1980
designed for kinetics-based measurements and has been previously described. Recent modifications include an automatic injecting and filling syringe (1mL) (Hamilton Model 77000) to introduce the last reagent which initiates the reaction. The pneumatically operated syringe was modified to allow triggering by an electrical rather than manual signal. Details of this modification are available from the authors. The syringe is now synchronously triggered by the same switch which triggers the measurement cycle of the ratemeter (14).Thus the measurement time is much more precisely synchronized to the time of initiation of the reactions than previously described ( 5 ) . Reproducibility is also improved because the relative standard deviation of the sample volume delivered is about 0.05%, the rate of sample delivery remains constant from run to run, and the dispenser-reagent reservoir system is closed and minimizes contamination. The plastic tube from the automatic syringe which delivers the last reagent is positioned directly over the cell by an arm which swivels out of the way to allow sample removal, cell cleaning, and addition of other reagents. The excitation monochromator was set at 365 nm and the emission monochromator at 444 nm. The slits were 2 and 1.4 mm, respectively, giving a spectral band-pass of 17 and 12 nm, respectively, for excitation and emission. All solutions were prepared with deionized distilled water (Millipore Milli-Q system). All chemicals were reagent grade except where noted. The glassware was cleaned as previously M), described ( 5 ) . Thiamine (J. T. Baker) stock solutions adjusted to pH 4 with HCl and kept refrigerated, were stable for months. Hg(I1) solutions were prepared from HgClz (J. T. Baker) and made pH 4 with HC1. The buffer was 0.2 M sodium or potassium monohydrogen phosphate (Mallinckrodt) and sufficient 5 M KOH to give the desired pH. The TC (Phaltz and Bauer), of unknown purity, was very slow to dissolve in alkaline aqueous solution. A 200 ppm (7.6 X lo4 M) TC solution retained its initial fluorescent intensity for over 1 month when stored in the dark. The vitamins were from a vitamin test kit (ICN) and most were of unknown purity. When the solutions were made, the chemicals were assumed to be 100% pure unless the purity was known. The substances for the interference study were dissolved in water adjusted to pH 4 with HCl. The biotin, talc, stearate, and starch solutions contained particles due to solubility limitations at the 200 ppm level in water a t pH 4. These solutions were filtered, as were samples, with a 1 . 2 - ~ mmembrane filter. The order and method of addition of the reagents into the sample cell were as follows: 1 mL of Hg(I1) solution, 1 mL of sample, both from Eppendorf automatic pipets, stirrer turned on, and 1mL of buffer solution from the precision liquid dispenser. The importance of this mixing order will be discussed later. For the blank, 1 mL of water at the same pH as thiamine standards was substituted for the sample. For the interference study, 0.5 mL of thiamine solution and 0.5 mL of the potential interferent, both from 0.5 mL Eppendorf pipets, were substituted for the sample solution. To determine if the potential interferent affected the blank signal, we substituted 0.5 mL of water for the above thiamine solution. All concentrations reported are initial concentrations before mixing unless otherwise stated. The pH is always reported as the measured pH of the reaction mixture. Between each run the sample cell was rinsed with pH 2 HCl once and Millipore water three times. The reagent and analyte solution bottles were immersed in the same temperature controlled water bath from which water was circulated through the temperature controlled cell holder. All measurements were made at 22.8 "C. The initial rate was calculated over a 16-9 period following a 16-s delay period which allowed time for the automatic addition of the last reagent, for mixing, and for the reaction to reach its maximum rate. The sample addition, rate measurement, and cell cleaning procedures required about 2 min per run. The rates that are reported in this study are the mean values resulting from at least three runs per test solution. For the equilibrium studies, solutions composed of equal volumes of Hg(II),thiamine, and buffer solutions were mixed in covered containers and allowed to react at room temperature. Samples were withdrawn periodically, and the fluorescence was measured until the reaction reached completion. When reactions were allowed to go to completion in the sample cell under continuous observation, photodecomposition caused lower yields and
a decrease in the time to equilibrium. Multivitamin-mineral tablets were prepared for analysis by grinding to a powder and dissolving in pH 2 HC1. The solution was filtered with a 1.2-pm membrane fdter (Millipore)and diluted to 1L. Aliquots of this were diluted by another factor of 10 giving a final concentration of about lo4 M thiamine. Absorbance measurements were made with an unmodified Cary 118C spectrophotometer. All absorbance data were extracted manually from recorder tracings.
RESULTS AND DISCUSSION Fluorescence is an ideal way to monitor the thiochrome reaction because of all the reactants, intermediates, and products, only TC is fluorescent. In addition, TC is the only product or reactant which absorbs significantly above 300 nm (15) and it has a molar absorptivity of 9.33 X lo3 a t 367 nm a t pH 7 (16). This is fortuitous because the Xe-Hg excitation source used in this work provides a strong line a t 365 nm. Since the other reagents do not absorb strongly a t this wavelength they do not cause a prefilter effect even a t high concentrations. It was determined that on a weight basis the maximum fluorescence of TC is greater than that of quinine sulfate and that its fluorescence is pH dependent in agreement with the literature (10, 17). The emission at 444 nm is at a maximum between p H 8 and 13 and is negligible below pH 4. Since T C is commercially available, a fluorescence Calibration curve was prepared and was linear from a detection limit of about 25 pptr to about 1 ppm TC. The amount of TC formed a t any time during the reaction can be determined from the fluorescence signal if quenching or absorption by other species is not significant. Thus the yield of the reaction can be determined and the measured rate in counts per second or volts per second can be converted into moles per liter seconds of TC produced. The oxidation of thiamine (TM) to fluorescent TC is always accompanied by the simultaneous formation of thiamine disulfide (TDS), a condensation product of two thiamine molecules. The ratio of T C t o T D S is affected by p H ( I @ , solvent (19),and oxidizing agent (20). Potassium ferricyanide is the oxidant in the standard equilibrium method but many other oxidizing agents have been used to oxidize TM such as KMnO,, M n 0 2 (15),cyanogen bromide (21),H202,and Hg(I1) (22-25). Cyanogen bromide is specific and quantitative for TM, but it is stable for only 3 h at r m m temperature and must be used in a fume hood (23). Oxidizing agents like K M n 0 4 (261,H202,12,and O2 favor the production of nonfluorescent T D S (20). Our investigation of various oxidizing agents showed that Hz02and K2Cr207produced very little TC. Low T C yield and a very fast rate made K M n 0 4 unsuitable for a kinetic method. Potassium ferricyanide was also found to be unsuitable for rate measurements with our instrument because the reaction reaches completion too rapidly compared to the mixing time. I t has the additional disadvantages of absorbing strongly at high concentrations and of instability at low concentrations. As part of this study is was determined that 3 X M K3Fe(CN), reduced the T C fluorescence signal by more than 75% but that quenching by 5 X M Hg(I1) did not occur. Data from an equilibrium method for mercury (27) and previous T M studies (22,25) indicated that the Hg(I1)-TM reaction would be appropriate for a kinetics based method. A reaction rate method for Hg(I1) using T M as a reagent (28) indicated the reaction was pseudo first order in TM. Hg(I1) has the additional advantages that i t has been used in the quantitative analysis of TM (22,23) and has been found to provide linear results over a much wider T M concentration range and provide greater specificity than K3Fe(CN)6in certain types of samples (23, 24). The reactions of T M in basic solutions (29) have been studied in detail. The oxidation by K3Fe(CN), has been
Yellow Thiol
Colorless Thiol Thiamine Disulfide
Figure 1. Reaction pathways in thiamine reaction. 008
1
007 0 06
-0 0 0 6 4
r
u 0 0 5 1
'
0 047 0.04
,
,
I
--
'-
0 00
230
250
270
290
310
0 001 230
330
350
Wavelength, nm
Figure 2. Spectra of species in the mercury-thiamine reaction. Each M TM, pH 10.8 buffer, sample consists of 1 mL each of (A) 1 X H,O; (B) 1 X M thiamine disulfide, pH 12.2 buffer, H20; (C) 1 X M TM, pH 12.2 buffer, HO , (after 10 rnin), this is the spectrum M Hg(II), pH 12.2 buffer, H,O. of the colorless thiol; (D) 5 X studied primarily from the point of view of reaction pathways rather than kinetics (20, 30). We have found no mechanistic or kinetic studies of the oxidation of T M by Hg(I1) in the literature. The reaction pathways and intermediates which have previously been identified to be important above pH 11 in aqueous solutions are shown in Figure 1 where it is assumed that Hg(I1) is the oxidant in reactions 2 and 5. Thiamine reacts immediately to form a tricyclic intermediate (CI) (reaction 1) which can be oxidized to T C (reaction 2) and converted to a yellow thiol (reaction 3), which is oxidized to TDS (reactions 4 and 5). The yellow thiol (YT) decays in a few minutes to the colorless thiol (reaction 4) when no oxidizing agent is present. Absorption Studies. Our study of the reaction by fluorescence and UV absorption provided valuable information concerning the Hg(I1)-TM reaction and confirmed some of the results obtained by Maier and Metzler (29). The formation and decay of several of the species identified in Figure 1 can be followed by molecular absorption, but since a great deal of spectral overlap occurs, quantitation is very difficult under many conditions. Strong absorption by the more concentrated Hg(I1) solutions as shown in Figure 2, curve D, and Figure 3, curve C, prevented use of spectral information below 250 nm. The T M and CI spectra are not observable with our spectrophotometer at p H 12.2 due to very rapid reactions (1 and 3) and the relatively low concentrations. The absorption
1I
\ \-*
II, \
, 250
\
.___---- ,
270
290
310
330
350
W a v e l e n g t h , nrn
Figure 3. Spectra of species in the mercury-thiamine reaction. Each M Hg(II), 1 X sample consists of 1 mL each of (A) 1.25 X M TM, pH 12.2 buffer (within 2 min after mixing); (B) 1.25 X M
HdII), pH 12.2 buffer, H20; (C) 9.5 X 10" M TC, pH 12.2 buffer, H,O.
maxima for TDS (26), spectra of TM, CT, and YT, which agree with those presented here, and a spectrum of the CI have been reported (29). When Hg(I1) is present in pH 12.2 aqueous thiamine solutions, several unexpected phenomena occur. The oxidation by Hg(I1) is slow and the increase in T C fluorescence is obM Hg(I1). This is served for over 30 min with 2.5 X especially surprising because a t this pH, T M solutions not containing Hg(I1) are almost completely converted to CT in only 4 min as indicated by the increase in absorption a t 265 nm (curve C, Figure 2). After the CT concentration reaches a maximum, it remains nearly constant for over a half hour (29). It is known (29) and we have experimentally verified that C T is not oxidized to TC. The formation and decay of YT, as observed between 310 and 350 nm (curve A in Figure 3), is also severely altered. For 1 X 10" M thiamine solutions, a t pH 12.2, and between 2.5 X and 2.5 x M Hg(I1) the concentration of YT initially formed is inversely related to the concentration of Hg(I1) and the rate of decay of YT is much less than when no Hg(I1) is present. At concentrations below 5 X N Hg(I1) the decay rate of the YT increases with decreasing Hg(I1) concentrations and approaches the decay rate seen without Hg(I1). Besides the changes in YT, a peak at 268 nm (curve A in Figure 3) is observed which is more intense at higher Hg(I1) concentrations. This unusual behavior can be explained by expanding on the proposed reaction of ref 26 and 29 shown in Figure 1.
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Hg(I1) apparently stabilizes T M or CI by the formation of a complex which can later be oxidized to TC. Addition of a large excess of Hg(I1) to the T M solution drives most of the T M or CI to the metal complex form which will be denoted Hg-TM. This species is observed as the peak a t 268 nm which is most intense a t highest Hg(I1) concentrations and the peak absorbance is directly proportional to T M concentration. This peak decreases with time while the absorption of T C observed a t 350 nm (curve B in Figure 3) and the fluorescence of T C a t 444 nm increase. The rate of decrease is first order with respect to T M as is the rate of formation of T C which is monitored by fluorescence. When Hg(I1) is not in great excess, and in alkaline T M solutions not containing Hg(II), the rapid increase in colorless thiol (CT) can be observed in the spectrum as an increase at 265 nm. At intermediate or low Hg(I1) concentrations the decrease at 268 nm due to the Hg-TM complex is difficult to measure due to the simultaneous formation of CT. Since Hg(I1) solutions begin to absorb below 250 nm, the CT and Hg-TM must both be observed in the 260-275-nm region making quantitative measurements of Hg-TM very uncertain. When Hg(I1) is in sufficient excess such that little Y T is produced and most of the neutral thiamine is tied up as Hg-TM, CT is not a predominant species in the solution. T h e dependence of initial Y T absorbance on Hg(I1) concentration can be explained by the Hg-TM complex. A t high Hg(1I) concentrations only a small amount of the T M is available to be converted to YT. A t lower Hg(I1) concentrations, less Hg-TM is formed leaving more of the T M to react to form YT. Addition of Hg(I1) to alkaline T M solutions which contain Y T have shown t h a t Y T also appears to be stabilized by formation of a Hg complex (Hg-YT). At less than a 1:l ratio of Hg(I1) to thiamine the decay rate of YT approaches that of solutions not containing Hg(I1). The behavior of the mercury(I1)-thiamine reaction below p H 11 also indicates a stabilization of YT by Hg(I1). Maier and Metzler state that below pH 10.6 the YT does not exist to an appreciable extent. But we find that a t pH as low as pH 8.9, more than 10% of the T M in a 1 X M solution is converted to YT in the presence of Hg(II), and it persists for more than 2 min. Without Hg(I1) in the solution no YT was observed. Although the formation of complexes does not appear to be important in the ferricyanide oxidation, a thiazolium ferricyanide complex has been reported (30). Thiamine complexes with other metals do form (31, 32) and Hg(I1) is known to favor complexes with sulfur or nitrogen compounds (33). Hg(I1) complexes with the amino acids, cysteine, and penicillamine, which contain nitrogen and sulfur and have some structural similarities to T M and YT, have been studied and have formation constants of 1014.21 and 1016.15, respectively (34). Penicillamine has even been used as an antidote in Hg poisoning (34). Although YT has been proposed as the TC precursor in the ferricyanide oxidation (20, 29, 30), Risinger and Pel1 (26) suggested that the cyclic intermediate is the TC precursor and this agrees with our findings. We determined that solutions containing higher levels of Y T produced less TC at slower rates upon oxidation by Hg(I1) and that solutions containing an identical amount of Hg(I1) but much less Y T produced more TC at a faster rate. In solutions high in Hg(I1) and YT, much more YT was lost than T C was produced after 5 min of reaction. Thus conditions which favor the formation of YT (i.e., low Hg(I1) concentration or base added before Hg(I1)) produce less TC, indicating t h a t YT is not the major T C precursor. T h e importance of the mixing order, i.e., adding the buffer last, given in the Experimental Section also becomes evident in light of these reactions. If the T M is mixed with base first,
L 1
01 8
I
1
IC
I
I
I2
1 I4
PH
Figure 4. Fluorescence signal from the mercury-thiamine reaction at M Hg(II), 2 X lo-' M TM. equilibrium vs. pH, 2.5 X
/_____________________ 1501
0 c
50 0 0 '
1
0 5
1
2 5 H p ( l l l Concentrafion, M x I O 3
1
5.0
Figure 5. Fluorescence signal from the mercury-thiamine reaction at equilibrium vs. Hg(I1) concentration, pH 12.2, 2 X lo-' M TM. The dashed line represents the theoretical maximum signal. YT and CT form, neither of which produce a significant amount of TC. Any nonreproducibility in the time between base and Hg(I1) addition would affect the amount of thiamine able to be oxidized to TC. Experimentally it is found t h a t p H 12.2 with a 15-s delay after the addition of base before the addition of Hg(II), a 20 times smaller rate of formation of TC was obtained. After a 3-min delay most of the thiamine is in the form of the colorless thiol and the rate was 200 times smaller. Those using ferricyanide have frequently pointed out that the oxidant should be added before or a t the same time as the base (35, 36). A more detailed discussion and a presentation of the experimental data obtained in the investigation of the reactions of T M and Hg(I1) have been made elsewhere (28). Equilibrium Fluorescence Studies. Equilibrium studies were undertaken to gain a knowledge of the yield of the reaction and the time to reach completion. The effect of p H on the reaction was studied by using the same phosphate buffer for all samples and using KOH or HCl to adjust the pH of the buffer. Figure 4 shows the maximum fluorescence signal measured for a particular reaction mixture vs. pH. The signal is influenced by variation with pH of both the T C fluorescence quantum efficiency and the reaction yield (percent of thiamine converted to TC), although the latter effect is dominant. The signal is optimum between pH 10 and 13. The time to equilibrium decreases from many hours a t pH 8 to less than 5 min at pH 14. Studies were not done below pH 7 because of the decreasing quantum efficiency of TC and because T M is much less reactive under acidic conditions (15). The fluorescence of the reagent blank is insignificant a t all pHs. The decreased yield below p H 10.7 is consistent with the reaction sequence we have proposed. The abstraction of H + from the neutral T M to form the tricyclic intermediate is much less likely below pH 10.6 (29). Above pH 13 the yield could be reduced by the effect of OH- on Hg(I1) which will be discussed later.
ANALYTICAL CHEMISTRY, VOL. 52, NO. 13, NOVEMBER 1980
2181
0
0
\
Precipitation
\
\
=
No Precipifation 2 5
a\ \
*.*e, -4
Figure 6. Conditions for precipitation of Hg(I1) for stirred solutions containing 0 . 2 M phosphate buffer: (0)precipitation occurred (a) precipitation did not occur. The effect of Hg(I1) concentration on the reaction is shown in Figure 5. The maximum fluorescence signal was obtained with the highest Hg(I1) concentration. A plot of the UV absorption at 268 nm due to the Hg-TM complex vs. [Hg(II)] has the same shape as Figure 5 . Since higher Hg(I1) concentrations inhibited the formation of the YT and CT, less of the T M would be oxidized to TDS and more to TC. It is also interesting to note that a curve similar to Figure 5 has been published as a calibration curve for an equilibrium method for Hg(I1) using T M as the reagent (27). T h e time to equilibrium decreased with increasing Hg(I1) concentration. A sample containing 5 x M Hg(I1) required about 30 min to reach equilibrium and a reaction mixture containing 3.8 X M Hg(I1) required about 15 min. This is an indication that the reaction is not first order in Hg(I1). The theoretical maximum yield plotted is the signal that would be obtained if all the thiamine were converted to TC. T h e solubility of Hg(I1) a t alkaline pHs was also studied. T h e formation of a precipitate in a fluorescence sample is unacceptable because of the nonreproducibility in the amount and particle size of the precipitate and the increased noise in the total apparent fluorescence signal. In addition, if the precipitate formed during the rate measurement, the increase in scattering would be measured as an apparent rate. Since calculations for the solubility of HgO or Hg(OH)2and the formation of several mercury(I1) hydroxo complexes indicate that precipitation should occur much sooner than was actually observed, it was necessary to determine the precipitation threshold empirically. It was found that for stirred solution in the sample cell with the 0.2 M phosphate buffer, precipitation occurred in 5 min or less for conditions at or above the dashed line shown in Figure 6. That a precipitate had formed was determined by an apparent increase in the apparent fluorescence signal of a blank solution caused by the scattering of excitation radiation by the particles in the solution. Precipitate occurred a t a given p H whether the buffer or just KOH were in the solution with Hg(I1). Kinetic F l u o r e s c e n c e Studies. The effect of the concentration of each of the reagents on the initial rate and on the standard deviation of the initial rate of the reaction was studied. The results of the kinetics p H study are shown in Figure 7 (curve A). The rate of the reaction increases steadily with [OH-]. Since the formation of T C requires the abstraction of hydrogen ions from the neutral TM, it is expected that the rate would increase with [OH-]. The relative standard deviation (RSD) in the initial rate was 1-2% from p H 11.9 to 13.5.Since slower rates were obtained at lower pH and the standard deviation of the blank and dark signals began to be significant, the precision decreased at lower [OH-]. At higher [OH-] poorer precision was also obtained due to the precip-
-2 Log [OH-], M
0
Figure 7. log-log plot of initial rate vs. [OH-], 2.5 X lo3 M Hg(II),2 X lo-' M thiamine: (A) measured rate; (B) rate compensated for changes in fluorescence quantum efficiency.
I
I
-4
-3
I
-2
Log [ H g ( I I ) ] , M
Figure 8. log-log plot of initial rate vs. [Hg(II)],pH 12.12: (A) 5 X M TM; (B) 1 X loT5M TM; (C) 2 X lO-'M TM; (D) 5 X lo-' M TM.
itation of HgO which would eventually occur a t this pH, and because TC was observed a t form without the addition of an oxidant. The presence of dissolved O2in thiamine solutions is known to lead to the formation of TDS (24) and T C (37). The oxidation of T M by dissolved oxygen could have contributed to the decrease in reproducibility. For curve B in Figure 7, the rate has been normalized for the pH dependency of the TC fluorescence quantum efficiency which was evaluated in a separate experiment. This correction provides a plot proportional to initial rate of moles per liter per minute formation of TC. The conversion factor from a rate in counts to a rate in moles per liter per second for 16-9 measurement period, such as those in Figures 7 and 8, is approximately 7.4 X mol L-l counts-'. This conversion factor must be determined for the specific instrument and instrumental conditions used. Since the rate is strongly p H dependent, the reaction mixture must be well buffered. The applicability of phosphate and carbonate buffers was studied. The carbonate buffer gave slower rates but precipitation with Hg(I1) did not occur any more readily than in the phosphate buffer. The phosphate buffer was selected for use in the reaction mixture because it provided better buffering capacity at p H 12 and because better precision in measured rates was obtained. The slope of the plot in Figure 7 changes with pH, but in the region of analytical interest, p H 11 to 13, the slope is approximately +0.9. The fractional order is not suprising in
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such a complex reaction. The relative amounts of CI vary with p H and the rate of formation of C T increases with pH (29). Hg(I1) also interacts with hydroxide and although the concentration of Hg(I1) added to the reaction mixtures is constant, the effective concentration is changing with pH. A decrease in effective Hg(I1) concentration with increasing p H makes the slope of the plot in Figure 7 smaller than if only the effect of the p H on the thiamine could be measured. Precipitation of Hg(I1) occurred above p H 13, although usually not during the measurement period. The dependence of the rate on Hg(I1) concentration at constant p H follows the same trend as the equilibrium study and is shown in Figure 8. Curve C was obtained several months earlier than the other three curves. That these curves nearly parallel indicates that the reaction is pseudo first order in thiamine a t all these Hg(I1) concentrations. I t is known from the absorption studies that the concentration of Hg-TM varies with Hg(I1) concentration and from fluorescence studies that the yield of T C varies with Hg(I1). Since Hg(I1) affects the reaction in several ways, only an apparent order with respect to Hg(I1) can be determined from these plots. The slope a t 2.5 X M Hg(I1) is approximately 0.6. The RSD in the initial rates was 1-2% for Hg(I1) concentrations between 5X and 3.8 X M. The RSD for 5 X M Hg(I1) was 4% at p H 12.2 where precipitation occurs within 5 min (Figure 6) but was 1-3% between p H 11 and 12. T h e reaction rate approximately doubled with a 5 "C increase in temperature. Since the rate could be adequately manipulated by varying p H and Hg(I1) concentration, the convenience of working a t room temperature did not have to be sacrificed. Sixteen seconds was found to be a convenient measurement time for routine analyses. However, this measurement time is short enough that it could be increased to provide a better S / N (38) for very low T M concentrations without suffering excessive photodecomposition. Continuous irradiation of the reaction mixture caused photodecomposition amounting to about 1% min-'. At the optimum conditions selected, 22.8 "C, p H 12.2, 0.2 M phosphate buffer, and 2.5 X M Hg(II), the reaction is pseudo first order in thiamine, 2% of the reaction occurs in 16 s, an 80% conversion of T M to T C is obtained, the precipitation of Hg(I1) is never a problem, and the precision is 1-2% at concentrations well above the detection limit (DL). The apparent orders with respect to hydroxide and to Hg(I1) are 0.9 and 0.6, respectively. These optimum conditions were used for all further work unless it is explicitly stated otherwise. Under these conditions, the pseudo-first-order rate constant determined from the first half-life a t several thiamine concentrations and by other methods (39, 40) was found to be 0.002 f 0.0002 s-*. I t should be noted that the reaction did not appear to remain first order in T M over the whole reaction curve, i.e., successive half-lives decreased in length. Over long periods of time and under continuous irradiation, the reaction curve could have been complicated by side reactions and photodecomposition. These factors also suggest that measurement times be kept reasonably short. The T M calibration curve is linear from the DL of 2 x lo-' to 1 X M. This large dynamic range of almost 4 orders of magnitude is much greater than recommended for the standard method. The RSD is 1% a t and above 1 x lo* M M. The DL is defined as the concentration and 2% at 5 X yielding a rate two times the standard deviation of the blank rate. With the emission monochromator replaced by a 440-nm interference filter (bandwidth 8 nm a t half-height) ( 5 ) ,the fluorescence signals were increased about a factor of 100. The detection limit was improved by a factor of 10 to 2 X M TM. This indicates that shot noise in the blank fluorescence
Table I. Species Which D o Not Interfere at Highest Concentration Tested highest concn tested, ppm Excipients lactose fructose sucrose talc starch
100 100 100
satd soln satd soln Anions
acetate sulfate
103 100
Vitamins A (acetate)a
000.8 (2500 IU/L)
E
satd soln
B6
calciferol calcium pantothenate choline chloride biotin
10.3 111 103
52
satd soln
Metals Ca( 11)
100
Intrinsic fluorescence decays under reaction conditions but vitamin A decomposes readily to a nonfluorescent compound during sample preparation utilized, a
signal is limiting as previously noted (5). The precision in rate measurements at higher thiamine concentrations was not as good as with the emission monochromator system. Interference Studies. An extensive interference study, aimed at the analysis of vitamin tablets, was performed with the optimum reagent concentrations described earlier. A substance was considered to not interfere a t a given concentration if the mean T M rate measured in the presence of the substance was within 1 standard deviation of the rate with T M alone. Each potential interferent was also mixed with the blank to determine if it reacted with the reagents or exhibited an intrinsic fluorescence under the reaction conditions. Compounds which appeared to depress the rate of the T M reaction were further tested with T C to determine if they quenched the TC fluorescence or caused a prefilter effect rather than affecting the rate of the reaction. In most cases the highest concentration tested was 100 ppm. All concentrations are normalized to a sample volume of 1 mL containing the interfering species. The results of the interference study are summarized in Tables I and 11. The effects of Fe(I1) and Mn(I1) were further studied a t p H 11.5. Precipitation is less of a problem at this pH, and the noninterfering levels are 2.5 and 100 ppm for Fe(I1) and Mn(II), respectively. If samples are known to contain considerable amounts of these ions, other reaction conditions, i.e., lower pH and higher Hg(I1) concentration, may provide optimal results. All of the metal ions that interfered tended to depress the rate of the mercury-thiamine reaction. Fe(I1) is potentially the most serious metal interference because i t has one of the lowest interference levels and is a t sufficient concentration in many vitamin-mineral preparations to be near the interference level after the dilution described for the sample preparation procedure. Fe(I1) and other metals have also been reported to produce low results for T M determined by the standard method ( I , 41) in some cases because of decomposition of the vitamin rather than interference with the TC reaction (42). The interferences from vitamin A and B2 have been discussed previously with respect to spectra taken with an intensified diode array system during the course of the T M reaction (43). At typical ratios of T M to A or B2in vitamin
ANALYTICAL CHEMISTRY, VOL. 52, NO. 13, NOVEMBER 1980
Table 11. Species Which Interfere level of
noninterference, PPm bromide iodide B, C
Bl2 nicotinamide folic acid
50 76
Vitamins 1.8 F (1.8 ppm, -2%) 3.5 P, B ( 3 5 ppm, -45%) 0.5 D, A (2.6 ppm, -22%) 25 D (50 ppm, -5%) 5 D, A, B (10 ppm, -3%)
Metals D (100 ppm, -15%) D ( 5 0 ppm, -4%) D, R (50 ppm, -5%) Fe(11) 1 D, P, A ( 5 ppm, -11%) Mn( 11) 1 D, P, A, R (10 ppm, -7%) a A, substance interferes by absorbing excitation or emission radiation or by quenching fluorescence; B, substance reacts with reagents; D, substance depressed apparent reaction rate; F, intrinsic fluorescence of substance decays under reaction conditions; P, precipitation occurs; R, substance reacts with thiamine, values in parentheses (level of substance tested which produced given percentage effect on rate measured for 1 X M (0.34 ppm) or 5 X lo-' M (0.17 ppm) thiamine). Mg(II) Zn( 11) Cu( 11)
Table 111. Formulation of Synthetic Multivitamin-Mineral Preparation Similar to the Parke-Davis Myadec Formula quantity in equiv of 1 tablet, mg Vitamins
type of interferencea
Anions D (100 ppm, -2%) D (760 ppm, -14%)
25 5 5
preparations, there would be no interference problems. Although the interference level for vitamin B12is the lowest, there would be no problem with typical vitamin preparations because the ratio of T M to BI2is always very large (e.g., 100-1000). Vitamin C could cause some problems for preparations which contain excessive amounts of vitamin C. Vitamins A and C may decompose in the sample solution or in the sample preparation steps. Multivitamin-Mineral P r e p a r a t i o n Analysis. The applicability of the method to multivitamin-mineral tablets was tested on two types of samples: commercial multivitamin mineral tablets and a synthetic mixture composed of all the listed ingredients, in about the same proportions, as Myadec multivitamin-mineral tablet. The nominal formula in an 0.8-g pill and the nominal concentrations after dilution are listed in Table 111. For the synthetic sample, each constituent was ground to a powder, weighed, and mixed thoroughly with the other constituents. Samples taken from this synthetic preparation were dissolved in p H 2 HCl, filtered, and diluted to 1 L. Aliquots of this were diluted by another factor of 10 for the analysis, giving an approximate T M concentration of 3 x lo4 M. The solution was p H 4. The results of the analyses are shown in Table IV. The values obtained for the commercial preparations agree reasonably well with the nominal values. In the commercial tablets, age and tablet variations could cause the amount of T M present to be different from the amount claimed by the manufacturer. Analysis of the synthetic preparation provides a better test of the method because the amount of T M in each sample is known. The agreement between the known and determined amounts is much better. It was found that in these samples the T M concentration remained nearly constant for a longer period of time if the final sample solution was p H 2. Even at p H 2, 10% of the T M was lost overnight. Under these analysis conditions, the standards should also be a t p H 2. Unless noted, all determinations were made by the method of standard additions by spiking the final dilutions with a standard T M solution to an added concentration of 2 x lo4 or 4 X lo* M. The standard addition method should com-
2183
vitamin A (acetate) vitamin D
concn in sample, ppm
3 (10000 I U )
0.3
1
0.1
(calciferol)
ascorbic acid thiamine vitamin B, vitamin B, vitamin B,, nicotinamide vitamin E iodide (KI) manganese (MnSO,) iron (FeSO,)
250
25
10 10
1
l(2.97 X
5
0.5 0.01
0.1 100
10
30
(insoluble)
Mineralsa 0.25
0.025
1
0.1
20
copper
M)
2 0.2
2
(CUSO,)
zinc 1.5 0.15 (ZnSO,) magnesium 25 2.5' (MgO) a Masses and concentrations refer to the element of inMgO would not completely dissolve under samterest. ple preparation conditions; therefore, actual concentration would be less than this. Table IV. Determination of Thiamine in Multivitamin-Mineral Preparations
sample Parke-Davis Myadeca Bi-Marta Bi-Mart
amt thiamine amt insam- found, ple, mg mg 10 2.5 2.5
error, %
11.1 2.5 2.1
+11 Ob
-16 av
synthetic prepnC synthetic prepnC synthetic prepnC synthetic prepnC
10.12 10.17 9.92 9.92
10.45 10.21 9.79 9.38
-2.5
+3.3 +0.4
av
-2.1 -5.5 -1.0
Commercial multivitamin-mineral tablet. Determined from calibration curve. Synthetic multivitaminmineral preparation as described in text. a
pensate for changes in the slope of the calibration curve (Le., the reaction rate) and for the gradual decomposition of T M that occurred in dilute solutions containing metals. Determination of the T M concentration in the samples with a calibration curve from external TM standards gave reasonable but less accurate results. This is partially due to the difference in the rate of decomposition of T M in pure standard solutions as opposed to thiamine in solutions containing metals. Samples consisting of thiamine plus all the metals salts, MgO, and KI in the same proportions as the vitamin-mineral mixture were also analyzed. T h e slope obtained from standard additions and external thiamine standards did not differ significantly demonstrating that the thiamine reaction rate is not appreciably affected by the metals. A direct comparison of the reaction rate method and the USP standard method (7) was also performed by using a synthetic vitamin-mineral preparation which did not contain
2184
Anal. Chem. 1980, 52, 2184-2189 Ohnesorge. W. E.; Rogers, L. B. Anal. Chem. 1958, 26, 1017-1021. Ellinger, P.; Holden, M. Biochem. J . 1944, 38, 147-150. Ziporin, Z.Z.;Beier, E.; Holland, D. C.; Bierman. E. L. Anal. Biochem. 1982, 3, 1. Mark, H. B., Jr. Talanfa 1973, 20, 257-268. Wilson, R. L.;Ingle, J. D., Jr. Anal. Chim. Acta 1976, 83, 203-214. Dyke, S.F. "The Chemistry of the Vitamins"; Interscience: New York, 1965, Chapter 2. Weast, R. C., Ed. "Handbook of Chemistry and Physics", 53rd ed.; Chemical Rubber Co.: Cleveland, OH, 1967; p C-508. Kavanagh, F.; %&win, R. H. Arch. Biochem. 1949, 20, 315-324. Kawasaki, C. Modified Thiamine Compounds, in Vitamins and Hormones"; Harris, R . S., Ed.; Academic Press: New York, 1963; Vol. 21. Wostman, B. S.;Knight, P. L. €xperientia 1980, 16, 500. Sykes, P.; Todd, A. R. J . Chem. SOC. 1951, 534-544. Fujiwara, M.; Matsui, K. Anal. Chem. 1953, 25, 810-812. Holman, W. I.M. Biochem. J . 1944, 38, 388-394. Edwin, E. E.; Jackman, R.; Hebert. N. AnaVst(London) 1975, 700, 689-695. Morita, M.; Kanaya, T.; Mlneska, T. J . Vitaminol. 1989, 75, 118-125. Prokhovnik. S. J. Analyst (London) 1952, 77,257-259. Risinger, G.E.; Pell, F. E. 8iochim Biophys. Acta 1985, 707, 374-379. Holzbecher, J.; Ryan, D. E. Anal. Chim. Acta 1973, 64, 333-338. Ryan, M. S.Ph.D. Thesis, Oregon State University, Corvallls, OR, 1980. Maier, G. D.; Metzler, D. E. J. Am. Chem. SOC.1957, 79. 4388-4391. Nesbitt, P.; Sykes, P. J . Chem. SOC. 1954, 4585-4587. Gero, E. C . R.Hebd. Seances Acad. Scl. 1952, 235, 397-399. Tanaka, A. Bhamin 1988, 33,497-502; Chem. Abstr. 1986, 65.3869d; 1988, 6 4 , 8178d; 1988, 6 4 , 97190. Cotton, F. A.; Wilkinson, G. "Advanced Inorganic Chemistry, A Comprehensive Text"; Interscience: New York, 1972; pp 517-519. Lenz. G. R.; Martell, A. E. Biochemistry 1984, 3. 745-760. Penttinen, H. K. Acta Chem. Scand., Ser. 6 1978, 830, 659-663. Waston, H. A. Cereal Chem. 1948, 23, 186-187. Satiidarma, K. S u r a Pharm. 1966, 79, 1-10; Chem. Abstr. 1966, 65, 15 164c. Ingle, J. D., Jr.; Crouch, S . R. Anal. Chem. 1971, 43, 697-701. Laldler. K. J. "Chemical Kinetics"; McGraw-Hill: New Y d , 1965; pp 13, 19. Hall, K. J.; Quickenden, T. I.; Watts, D. W. J . Chem. Educ. 1976, 53. Matsui. K.; Honda. M. Bitamin 1958, 15, 654-659; Chem. Abstr. 1982 57, 754h. De Ritter, E.; Rubin, S. H. Anal. Chem. 1947, 19, 243-248. Ryan, M. A.; Miller, R. J.; Ingle, J. D., Jr. Anal. Chem. 1978, 5 0 , 1772-1777.
Table V. Determination of Thiamine in Synthetic Multivitamin-Mineral Preparation Solutions by the Standard Method and the Kinetic MethodQ
method U.S.P.
kinetic
a
thiamine concn obtained, no. of PM runs 2.40 2.40
RDS,
%
7c
error
3
3
3
1
5 2.5 2.35 5 2 2.45 3 2 The solutions were prepared to contain 2.5
-4 -4
2.47
TM .
-1
-6 -2 X
M
TM. This sample was prepared as before except that the p H 2 sample solution was spiked with a standard T M solution a t the final dilution such that its concentration was 2.5 X lo4 M T M . The results for both methods were obtained from a calibration curve. Table V shows that both methods gave accurate results. Variations in turbidity of the butanol and the formation of a very small amount of TC in the USP method blank gave a DL of 7 X M T M for the standard method. Since the rate method does not involve extractions, the rate method is faster than the standard method in analysis and glassware cleanup time and less expensive due to the use of less chemicals and glassware.
LITERATURE CITED "Methods of Vitamin Assay"; Association of Vitamin Chemists, Inc.; Interscience: New York, 1966; Chapter 6. Strohecker. Rolf; Henning, Heinz M. "Vitamin Assay-Tested Method"; Verlag Chemie: Weinheim/Bergstr., Germany, 1966. Nicheison, 0.; Yamamoto, R. S. I n "Methods of Biochemical Analysis"; Glick, D., Ed.; Interscience: New York, 1958; Vol. 6, pp 191-257. Wilson, R. L.; Ingle, J. D., Jr. Anal. Chem. 1977, 4 9 , 1066-1070. Wilson, R. L.; Ingle, J. D., Jr. Anal. Chem. 1977, 4 9 , 1060-1065. Official Methods of Analysis of the Association of Official Analytical Chemists"; Horwitz, W., Ed., AOAC: Washington, DC, 1970. United States Pharmacopeia", 19th ed.; United States Pharmacopeia Convention, Inc.: Rockville. MD, 1975; p 629. Pippin, E. L.; Potter; A. L. J . Agric. Food Chem. 1975, 23, 523. Gassmann, G.;Janicki, J.; Kaminski, E. Int. Z . Vitaminlorsch. 1963, 33, 1-17.
RECEIVED for review January 26,1979. Resubmitted July 14, 1980. Accepted August 26,1980. Acknowledgment is made to the NSF (Grant No. CHE-76-16711 and CHE-79-21293) for partial support of this research, and M.A.R. gratefully acknowledges an NSF graduate fellowship. Presented in part a t the 1979 Pittsburgh Conference, Cleveland, OH.
Luminescent Quantum Counters Based on Organic Dyes in Polymer Matrices Krishnagopal Mandal,
1.D. L. Pearson, and J. N. Demas"
Department of Chemistty, University of Virginia, Charlottesville, Virginia 22903
New luminescent quantum counters composed of laser dyes dispersed in either poly(viny1 alcohol) (PVA) or poly(vinylpyrrolidone) (PVP) solid matrices are described. The PVA matrix is more compatible wlth water-soluble dyes and the fllms detach easily from the supporting glass substrate. The PVP fllms are more compatible with dyes soluble in organic solvents and form extremely tenacious films which can only be removed by dissolution. A variety of xanthene and coumarin dye quantum counters are reported. Relative spectral flatness and sensitivity as well as usable wavelength range are dlscussed. Several new systems wlth superior properties are presented as well as the considerations involved In the proper selection of a quantum counter.
The measurement of light intensities plays a crucial role for many spectroscopists, photochemists, analytical chemists, photobiologists, and physicists. Accurate measurements are usually made with thermal detectors or quantum counters (1-5).
Thermal detectors (e.g., thermopiles, bolometers, and pyroelectrics) are energy flat detectors that give equal output for equal incident energies a t different wavelengths. They generally suffer from low sensitivities, slow response, and/or small areas ( 4 ) . Most spectroscopic measurements are concerned with photon fluxes rather than beam powers. Only if the spectral distribution is known, however, can the photon flux be calculated directly from the power.
0003-2700/80/0352-2184$01.00/0 0 1960 American Chemical Society
