Anal. Chem. 1985, 57,555-558 (7) Headridge, J. B.; Riddington. J. M. Scanning Electron Microsc. 1981, I I , 357-367. (8) Flanagan, F. J. Geol. Sum. Prof. Pap. ( U . S . ) 1975, No. 8 4 0 , 131-183. (9) Gladney, E. S.;Burns, C. E.; Perrin, D. R.; Roehndts, I.: and Gills, T. E. NBS Spec. Publ. ( U S . ) 1984, No. 260-08. (10) Gladney, E. S.; Burns, C. E. Geostd. Newsl. 1983, 7 , 3-226. (11) Wilson, A. D. Analyst(London) 1964, 89, 18. (12) Benedetti-Pichler, A. A. I n "Physical Methods of Chemical Analysis"; Berl, W. M., Ed.; Academic Press: New York, 1956, Vol. 3, pp 183-21 7. (13) Kratochvil, B.; Taylor, J. K. Anal. Chem. 1981, 53, 925A-938A. (14) Ng, K. C.; Zerezghi, M.; Caruso, J. M. A n d . Chem. 1984, 56, 4 17-42 1. (15) Korotev, R . L.; Lindstrom, D. J. Trans. Am. Nucl. SOC. 1982, 4 1 , 187. (16) Perlman. I.; Asaro, F. Archaeometry 1969, 7 7 , 21-52. (17) Jacobs, F. S.; Filby, R. H. Anal. Chem. 1983, 55, 74-77. (18) Filby, R. H.; Nguyen, S. N.; Grimm, C. A.; Markowski. G. M.; Ekambaram, V.; Tanaka, T.; Grossman, L. Proceedings of the 5th International Conference on Nuclear Methods in Environmental Energy -. Research, Mayaguez, PR, April 1984, in press. (19) Hoffman, B. W.; Van Camerik, S. B. Anal. Chem. 1967, 39, 1198-1199. (20) Davis, A. M.; Tanaka, T.; Grossman, L.; Lee, T.; Wasserburg, G. J.
555
Geochim. Cosmochim. Acta 1982, 4 6 , 1627-1651. (21) Finkelman, R. B. Ph.D. Thesis, Department of Chemistry, University of Maryland, 1980. (22) Ingamells, C. 0.; Switzer, P. Talanta 1973, 2 0 , 547-568. (23) Ondov, J. M.; Zoller, W. H.; Olmez, I.; Aras, N. K.; Gordon, G. E.: Rancitelli, L. A.; Abel, K. H.; Filby, R. H.; Shah, K. R.; Ragaini, R. C. Anal. Chem. 1975, 4 7 , 1102-1109. (24) Fisher, G. L.; Prentice, B. A,; Silberrnan, D.: Ondov, J. M.; Biermann, A. H.; Ragaini, R. C.; McFarland, A. R. €nviron. Sci. Techno/. 1978, 72,447-451. (25) LyOn, W. S. I n Proceedings of the Nuclear Methods in Environmental and Energy Research, Columbia, MO. April 1980, U.S. Department of Energy CONF-800433. pp 587-596.
RECEIVED for review August 6,1984. Accepted November 8, 1984. Supported in part by Electric Power Research Institute under Contract FP 1259-1, under direction of M. McElroy (MRI and WSU) and by the National Science Foundation under Grant EAR-821-8154 and the National Aeronautical and Space Administration under Grant NAG-9-54 (University of Chicago).
Irreversible Reaction Kinetics of the Aerobic Oxidation of Ascorbic Acid David Emlyn Hughes Norwich Eaton Pharmaceuticals, Inc.,' Norwich, New York 13815
Aerobic oxidation of ascorbic acid (AA) is studied at 25, 62, 75, and 86 'C. The AA is determined by 2,6-dichioroindophenol titratlon and dehydroascorbic acid (DHA) is determined simultaneously by continuous-flow derivatlzatlon with o -phenylenediamine using fluorescence detection. The pseudo-first-order, reversible rate constants for the formation of DHA and diketoguionic acid are discussed. The activation energy for AA degradation and the irreversible pseudo-firstorder rate constant for AA and DHA loss are presented. A degradation pathway from AA to products without the formation of DHA is postulated.
Several mechanistic studies on the oxidation of AA have been performed in the last 50 years. An excellent review of the literature of aerobic and catalyzed oxidation is presented by Mushran and Agrawal ( I ) . Definitive kinetic results may have been unobtainable for several reasons: (1)the presence of metal ions as impurities in laboratory water sufficient to catalyze the reaction, (2) the absence of specific methods for the determination of AA, and (3) the absence of precise and sensitive methods for the determination of DHA. Whereas under anaerobic conditions AA degrades to furfural and carbon dioxide, the aerobic mechanistic path is postulated to be oxidation to DHA followed by hydrolysis (2). COOH
Ascorbic acid (AA) is an unsaturated lactone which is a strong reducing agent. It is converted to dehydroascorbic acid (DHA) according to the reaction HO
0
OH
I /
0
+ O=C
C-H
0 '1 '
CHOH
O=C,
+
2e-
(1)
C ,H -
I
h20H
I
OH
\/
I
U
I
CH2OH
CHOH
dehydroascorbic acid
CHOH
I I CHPH
- FI
AOH
(2)
0
oxalic acid -I- other species
diketogulonlc acid
0 1
CH20H ascorbic acid
2Ht
0
dehydroascorbic acid
The reaction is apparently not reversible and is a function of at least the temperature, pH, and oxygen content of the aqueous sample. The autooxidation is strongly catalyzed by several metal ions, notably Cu(I1) and Fe(II1). ' A Procter & Gamble Company. 0003-2700/85/0357-0555$0 1.50/0
The degradation is apparently a function of at least the concentrations of the metal ions present, the pH, and the available light ( 3 , 4 ) . The different kinetic routes apparently all produce DHA. Hence, the oxidation of AA is assumed to occur by the initial formation of DHA followed by conversion of that species to diketogulonic acid (DKA) (5). The reactions are apparently pH dependent. In a more recent study, Blaug and Hajratwala (6) determined the unreacted AA vs. time using the 2,6-dichloroindophenol volumetric determination. An apparent first-order rate of degradation was achieved at 67 "C and a 0 1985 American Chemical Society
556
ANALYTICAL CHEMISTRY, VOL.
57, NO. 2, FEBRUARY 1985
pH range of 3.5-7.2 a t an ionic strength of 0.4. The authors postulated a mechanism in which undissociated AA, monodissociated AA, and a complex formed by the two irreversibly formed products. Some studies ( 1 ) have been performed using the classical 2,4-dinitrophenylhydrazine derivative of DHA [Roe's (7) method] which absorbs a t 500-550 nm after treatment with 85% sulfuric acid. The method does determine AA and DHA but requires a 3-h incubation period. Okamura (8) used a reduction of DHA with dithiothreitol followed by AA determination by the cup'-dipyridyl method which offers more sensitivity than Roe's method. The classical 2,6-dichloroindophenol titration of AA has also been used and is quite sensitive to metal ions but lacks specificity in the absence of a suitable masking agent. Methods relying on ultraviolet detection are hindered by the sensitivity of the AA ultraviolet maximum to pH changes and the very low absorptivity of DHA. High-pressure liquid chromatographic (HPLC) methods (9) combine the above difficulties with kinetic questions concerning the effects of residence of the species in the mobile phase. Several spectrophotometric and titrimetric procedures are available (10, 1 1 ) for the determination of AA including the colorimetric procedure of Schmall modified to use p-nitroaniline (12) where sensitivity is not a problem. In this paper, a procedure for simultaneously determining AA and DHA will be presented. Kinetic data a t several temperatures will be presented and a mechanism proposed. This mechanism suggests that some traditional vitamin C analyses may require careful interpretation.
EXPERIMENTAL SECTION Materials. All reagents used in this study were of an analytically pure grade. Methods. AA was determined by the 2,6-dichloroindophenol volumetric method as modified by Hughes (13) for solutions containing physically and chemically interfering species. The AA standard solution, maintained at the cited temperatures, was sampled periodically with a 2-mL pipet and titrated immediately. Simultaneously, the DHA was monitored by continuous-flow analysis. The sample line was combined with a 1mg/mL aqueous solution of o-phenylenediamine and passed through a delay coil. The resulting fluorescence due to the condensation product 3(1,2-dihydroxyethyl)furo[3,4-b]quinoxalin-l-one was then measured using a Technicon Autoanalyzer I1 Fluoronephelometer equipped with a Corning No. 58 secondary filter and a Corning No. 7-60 primary filter. Continuous flow was provided by a Technicon Autoanalyzer proportioning pump. The general method has been used previously to determine the total vitamin C content by prior oxidation with Norit (14), n-bromosuccinimide, and other oxidizing agents. The availability of automated procedures has made this method more precise (15). However, since some problems in the oxidation step have been reported in the literature, the o-phenylenediamine procedure may not be satisfactory for all determinations of AA. The procedure used here has proven successful for the determination of DHA. This appears to be the first time that this reaction scheme for the determination of DHA alone has been reported in the literature. Five spikes (100-1000 fig) of DHA to 1 mg/mL ascorbic acid solutions at a pH of 2.6 and room temperature gave an average recovery of 101%. DHA standards were linear (RSD f 2 7 0 ) with a correlation coefficient of 0.999 from 100 to 1000 Fg of DHA/mL.
KINETIC STUDY Method of Analysis. An accurately weighed quantity of AA (ca. 50 mg) was dissolved in water equilibrated to the water bath temperature in a 100-mL volumetric flask. The contents were immediately poured into a 125-mL Erlenmeyer flask in contact with a water bath thermostated at the desired temperature. Two-milliliter aliquots were removed periodically for titration and a continuous-flow injection line was placed in the flask to sample the dehydroascorbic acid. The initial pH of the samples was measured. A bubbler was used to
maintain saturation of the solution with oxygen.
DISCUSSION The implicit assumption of existing kinetic schemes is that the mechanism of degradation of AA to DHA and DKA is products
products
kI
AA
y
products
Re T-
DHA
DKA
(3)
characterized by comparatively large rate constants k , , k,, kz, and k2; k , is assumed to be small; and the paths with rate constants k , and k4 are not considered. As long as ascorbic acid is the only species determined (as in previous studies), the rate of disappearance of AA from reversible paths and irreversible paths cannot be distinguished since ( k , + k 3 ) is the apparent pseudo-first-order rate constant
-[AA] = k,[AA]
+ k,[AA]
= (k,
+ k3)[AA]
(4)
where [AA] = d[AA]/dt and the brackets indicate concentration in bg/mL. With the addition of DHA degradation concentration-time profiles, the role of the irreversible reactions of ascorbic acid becomes clearer. If the DHA appearance rate is lower than the ascorbic acid disappearance rate, one or more of the kinetic routes with rate constants k,, k4, or k , must be considered.significant. In the most straightforward case, [AA] [DHA], that is, the absolute rate of disappearance of AA approximates the rate of appearance of DHA and rate constants k,, k,, k,, and k , are small compared with the k l equilibration step. T o test this simple case, the analytically correct form for the concentrations according to Rodiguin (16) was employed, i.e., for eq 3 with k,, k,, and k , = 0
-
PlPZ
Co = [AA] = Co'o'---YlYZ
+
C1 = [DHA] =
Cz = [DKG] = 1 k1hZC0(O'-+ Y1Y2
e-ylt
YI(Y1-
YZ)
+
e-?*t
(7)
Y d Y 2 - Yl)
where y1and y2 are the roots of the equation
with reversed signs. The rate constants were then determined on a Hewlett-Packard 3357 with the experimental AA and DHA concentrations as input. Using Rodiguin's nomenclature, A. is AA, A , is DHA, A2 is DKA, k-, is PI, and k-? is p2, as described in eq 3. The kinetic data (Figures 1-3 and Table 11) were collected a t 25 "C and pH 9.0, 62 "C and pH 2.6, 75 "C and pH 2.6, and limited data at 86 "C and pH 2.6. The rate constants were calculated assuming this simple mechanism. The results appear on Table I. The activation energy for the reaction was determined to be 12.0 kcal/mol at a pH 2.6 by the standard Arrhenius plot of log k , vs. 1 / T , where T is the absolute temperature. The activation energy is consistent with the value of 12.2 kcal/mol, pH 3.52, given by Rogers (17).
ANALYTICAL CHEMISTRY, VOL. 57, NO. 2, FEBRUARY 1985
500
557
r
;JG/ML
250
0
60
30
90
TiME, MINUTES
30
0
Flgure 1. Kinetic data for the oxidation of ascorbic acid to dehydroascorbic acid at 25 "C and pH 9.0: 0 represents ascorbic acid and 0 represents dehydroascorbic acid.
k,, min-'
&, min-'
62 75 86
8X 1.2 x 10-4 3 x 10-4
2X 1 x 10-5
o
90
TIME, MINUTES
Figure 2. Kinetic data for the oxidation of ascorbic acid to dehydroascorbic acid at 62 "C and pH 2.6: 0 represents ascorbic acid and 0 represents dehydroascorbic acid.
Table I. Rate Constants for the Oxidation of Ascorbic Acid" temp, "C
60
&, min-I
k,, min-' 1X
0
7 x 10-5 1 x 10-5
0 0
The exact solution for the mechanism eq 6 with k,, k4, and k , = 0 was utilized. k , and are the forward and reverse rate constants for the AA to DHA conversion and k, and (3, are the forward and reverse rate constants for the DHA to DKA conversion, respectively.
250
1
Table 11. Concentration Profiles for Ascorbic Acid [AA] and Dehydroascorbic Acid [DHA] at 86 O C and pH 2.6 time, min
[AAI, d m L
[DHAI, PLg/mL
10 20 30 40 50 60 70 80 90
479 440 403 369 339 311 285 261 240
66 137 200 250 261 324
_L-_-2-----J
0
(8)
(9)
where R is some "global" rate of formation of products. Note that this rate might itself contain a reversible reaction, e.g., DHA might be reversibly converted to DKA and then irreversibly to products, as in the accepted mechanism. Integrating with respect to time, eq 9, becomes
[AA]
+ [DHA] = I
R dt
90
Figure 3. Kinetic data for the oxidation of ascorbic acid to dehydroascorbic acid at 75 "C and pH 2.6: 0 represents ascorbic acid and 0 represents dehydroascorbic acid.
to the extent that AA and DHA are lost, then
[AA] + [DHA] = R
60
TiME, MiNUTES
Although a reasonable set of rate constants was obtained with the reversible reactions mechanism, the computer fit of DHA-time data was not satisfactory, since the predicted DHA concentrations were higher than the experimental values. The mechanism was therefore extended to include irreversible paths to products. If k l and k - , are much larger than the other rate constants
[AA] + [DHA] = 0
30
(10)
In the mathematically simplest form, R is not a function of time and eq 10 becomes
[AA] + [DHA] = Rt f Ro
(11)
where Ro is the initial sum of the concentrations of AA and DHA. Hence, a graph of ([AA] + [DHA]) vs. time would be expected to be a straight line. For the AA data, the extrapolation formula
[A] = was used since all previous AA degradation studies suggest that the decomposition may be treated as a first-order reaction. The sum of [AA] plus [DHA] was then found and the resulting function plotted vs. time of reaction. After an initial maximum occurring in each case at typically 30-40 min, the ([AA] + [DHA]) then decreases monotonically with time. Linear regression of the data reveals that the relative standard deviation for the 40- to 90-min data points is 1.4% for 25 "C, 0.81?G for 62 "C, 0.41% for 75 "C, and 3.2% for 86 "C. The data are seen to be linear with an average relative standard deviation of 1.5%,which compares well with the expected rel-
558
ANALYTICAL CHEMISTRY, VOL. 57, NO. 2, FEBRUARY 1985
T a b l e 111. R a t e C o n s t a n t s Equation 3
temp,
a
k, and k4 for the I r r e v e r s i b l e Loss of A s c o r b i c A c i d [ A A ] a n d D e h y d r o a s c o r b i c A c i d [DHA] by slope,
"C
PH
unitless
25 62 75 86"
9.0 2.6 2.6 2.6
-0.852 -1.03 -0.489 -0.444
!.G/mL
A([AAl + [DHAl)/ A t , fig mL-' min-'
436 544 351 300
--0.785 -0.201 -0.772 -0.840
intercept,
k,, min-'
k,, min-'
1.8 x 3.7 x 2.2 x 2.8 x
1.5 x 3.8 x 1.1 x 1.2 x
10-3 10-4 10-3 10-3
10-3 10-4 10-3 10-3
Limited data a t this temperature.
ative standard deviation of f3-5% due to instrumental imprecision. If we now consider eq 3 with k 2 and k-, = 0, Le., two irreversible paths
+ k3)[AA] k-,[DHA] = (k-1 + k,)[DHA] - k,[AA]
-[AA] = -[DHA]
(121
-
(13) (14)
combining eq 13 and 14 -([AA] = [DHA]) = k,[AA]
+ k,[DHA]
(15)
and hence k4 [AA] = - -[DHA] k3
-
[AA]
+ [DHA] k3
(16)
A([AA] + [DHA])/At is a constant in the range t 2 40 min, hence eq 16 takes the form [AA] = m[DHA]
+b
(17)
where m = -k,/k3 and b = A([AA] + [DHA])/k3At. A graph of the AA concentration vs. DHA concentration allows the calculation of irreversible rate constants k, and k,. Table I11 summarizes results of the calculations for k3 and k,. Rate constants k , and k., represent total first-order loss of AA and DHA. Rate constant k 3 represents the loss of AA by a mechanism that does not have DHA as an intermediate. Although the kinetic data presented here fit reasonably well with the accepted AA-to-DHA-to-DKA mechanism (eq 6-9), further analysis presented here suggests that an irreversible path from AA to products appears to exist. The rate constant k, appears to be large enough that it may be possible to degrade measurable amounts of AA without any DHA being formed. The concept that AA may react significantly with several species is not new; in fact, Rogers (17) postulated reaction with monodissociated AA, ascorbate ions and a monodissociated AA-AA complex and presented rate constants for the processes. The presence of an irreversible path from ascorbic acid to products without intermediate DHA formations requires some examination of accepted analytical methodology. Vitamin C assays in which the AA is oxidized titrimetrically, e.g., by the classical 2,6-dichloroindophenol (13) procedure, are unaffected by the mechanistic details since these assays only require that the AA present be oxidized to
some other species. Vitamin C assays ( 1 4 , 15) which rely on the quantitative oxidation of ascorbic acid to dehydroascorbic acid may be affected. Assays of this type require careful evaluation since the AA to DHA concentration ratio in samples is not under the control of the analyst. Hence, such a method may perform well with AA standards and samples that contain a high percent of total vitamin C in the AA form. Assays of predominately DHA samples may yield erroneous results. With real samples, the analyst is not aware of the AA-to-DHA concentration ratio and is therefore unable to evaluate the results correctly. The irreversible paths involving the combined AA and DHA pool are of biological interest since only AA and DHA show antiscorbutic activity. Hence, if there are significant irreversible paths from AA and DHA to products, the "total vitamin C", or sum of the AA and DHA, will decrease markedly with time. If the oxidation of AA may be envisioned as solely the conversion to DHA, the total vitamin C activity will be maintained. An irreversible path from AA to products without intermediate DHA formation has a significant effect on the interpretation of the kinetics of the oxidative mechanism and analytical methods that rely on quantitative conversion.
LITERATURE CITED (1) Mushran, S. P.; Agrawal, M. C. J. Sci. Ind. Res. 1977, 36 (6), 274. (2) Connors. Kenneth A.; Amidon, Gordon L.; Kennon, Lloyd "Chemical Stability of Pharmaceuticals"; Wiley: New York, 1979; p 138. (3) Czuros. Z.; Petro, J. Acta Chim. Acad. Sci. Hung. 1955, 7 , 199. (4) Rao, B. S. N. Radiat. Res. 1962, 77,683. (5) Grochmalicka. J. Zesz. f r o b / . Postepow Nauk R o h . 1965, 53, 57; Chem. Absrr. 1986, 64,5813f. (6) Blaug, S. M.; Hajratwala, H. J. fharm. Sci. 1972, 61 (4), 556. (7) Roe, J. H.; Kuether, C. A . J . Biol. Chem. 1943, 774,399. (8) Okamura, M. Clin. Chim. Acta 1980, 703.259. (9) Rose, R. C.; Narhwold, D. L. Anal. Biochern. 1981, 774,140. (IO) Pandey, N. J. Anal. Chem. 1982, 54,793. (11) Verma, K. K.; Glutai, A. K. Anal. Chem. 1980, 52. 2336. (12) Weeks, C. A.; Deutsch. M. J. J. Assoc. Off. Anal. Chem. 1965, 48, 1245. (13) Hughes, David Emlyn J. Pharm. Sci. 1983, 72,126. (14) Deutsch, M. J.; Weeks, Cora E. J . Assoc. Off. Anal. Chem. 1965, 48, 1248. (15) Roy, R. E.; Conetta, A.; Salpeter, J. J. ASSOC.Off. Anal. Chem. 1976, 59, 1244. (16) Rodiguin, N. M.; Rodiguina, E. N. "Consecutive Chemical Reactions"; Van Nostrand: Princeton, NJ, 1964: p 42. (17) Rogers, A. R.; Yacomeni, J. A. J. fharm. fharmacol. 1971, 23. 218.
RECEIVED for review July 30, 1984. Resubmitted November 5 , 1984. Accepted November 16, 1984.
