Spectrophotometric Determination of Mkro Amounts of Ascorbic Acid in Citrus Fruits Bruno Jaselskis and Joseph Nelapaty, S. J. Department of Chemistry, Loyola University, Chicago, Ill. 60626
DETERMINATION OF ASCORBIC ACID in vitamins and natural products is of great importance. Macro and semi-micro amounts have been determined titrimetrically by iodometric methods such as N-bromosuccinimide ( I ) , iodate-iodide (2), bromate-bromide (3),iodine (4,copper(I1) sulfate (5),mercury (11) chloride (6) and others. Among the micro methods may be mentioned the titrimetric methods involving 2,6-dichlorophenolindophenol (7-9), and phenothiazin-3-one (IO), the gas chromatographic method (11) using the trimethylsilyl derivative of ascorbic acid, poloarographic (12), and several spectrophotometric methods. The latter involve a wide variety of reagents based either on the reactivity of the dicarbonyl function of dehydroascorbic acid as illustrated in the method using 2,4-dinitrophenyl hydrazine (13), or on the reductive character of ascorbic acid in the methods involving diazotization reaction with 4-methoxy-2-nitroaniline (14), and the reduction of iron(II1) followed by the determination of iron(I1) as phenanthroline (15) or bipyridyl chelate (16), which have molar absorptivities of 11000 and 8650, respectively. Recently Stookey (17) reported a very sensitive method for the determination of iron(I1) using ferrozine : 3-(2-pyridyl)5,6-bis(4-phenylsulfonic acid)-1,2,4-triazine disodium salt. The reagent forms a very stable chelate with iron(I1) and has its absorption maximum at 562 nm and a molar absorptivity of 27900. This reagent is not only soluble in water, but also reacts readily with iron(I1) in a wide pH range. Thus, it appears to be a suitable reagent for the determination of reducing species in aqueous solutions. In this paper we present the optimum conditions necessary for the quantitative determination of ascorbic acid in citrus fruits. EXPERIMENTAL
Apparatus. Spectral measurements were carried out using Cary XIV spectrophotometer. Corning Model XI1 research pH meter was used for determination of pH. All glassware was washed with 6 M HC1 and deionized water. (1) D. F. Evered, Analyst, 85, 515 (1960). (2) G. Spacu and P. Spacu, 2.Anal. Chem., 128,233 (1948). (3) P. Spanyar and 0. Petro, Elelmiszervizsgalati Kozlem., 13, 130 (1968). (4) J. W. Stevens, IND.ENG.CHEM., ANAL.ED., 10, 269 (1938). (5) L. G. Gein, V. K. Rubtsov, and Z . A. Sumbalkina, Farmafsiya (Moscow) l 8 , 6 5 (1969). (6) L. Kum-Tatt and P. C . Leong, Analyst, 89,674 (1964). (7) W. B. Robinson and E. Stotz, J. Biol. Chem., 160, 217 (1945). (8) H. Vallant, Mikrochim. Acta 1969, 436. (9) H. Liebman and A. D. Ayers, Analyst, 70,411 (1945). (10) M. Raileanu and E. Dobre, Rev. Roum. Cliim., 14, 1053 (1969). (11) L. T. Sennello and C. J. Argoudelis, ANAL.CHEM., 41, 171 (1969). (12) J. E. Page and J. G. Waller, Analyst, 71, 65 (1946). (13) J. H. Roe and C. A. Kuether, J. Biol. Chem., 147, 399 (1943). (14) M. Schmall, C. W. Piefer, and E. G. Wollish, ANAL.CHEM., 25, 1486 (1953). (15) I. Onishi and T. Hara, Bull. Chem. SOC.Japan, 37, 1314 (1964). (16) R. Aragones-Apodaca, Inform. Quim. Anal. Pura Apl. Ind., 21, 230 (1967). (17) L. Stookey, ANAL.CIiEM., 42,779 (1970).
Reagents. Ferrozine, 3-(2-pyridyl)-5,6-bis(4-phenyl sulfonic acid)-l,2,4-triazine disodium salt, was purchased from Hach Chemical Company. Unfortunately the reagent was grossly impure and had to be recrystallized at least four times from distilled water or water-alcohol solutions. A stock solution of approximately 0.01M was used. Stock solutions of acetic acid-sodium acetate buffers were prepared from the ACS reagent grade chemicals. The total concentrations of sodium acetate and acetic acid in all buffer solutions was held at 0.3M. A stock solution of 0.002M of iron(II1) reagent was prepared either from reagent grade ammonium ferric sulfate by dissolving it in dilute perchloric acid or by dissolving iron wire in perchloric acid containing a few drops of 6 M HNOI. Properly prepared reagent could be kept for several weeks before a significant appearance of iron(I1). Reagent grade ascorbic acid was used to prepare approximately 0.01M stock solution containing also 0.02 mole of oxalic acid or sodium oxalate. Water was boiled and deaerated before the addition of ascorbic acid. A 10-ml aliquot of the ascorbic acid solution was then diluted to 1 liter with deaerated water, the contents were transferred to a Machlett buret and were kept under a nitrogen atmosphere. Concentration of ascorbic acid in the stock solution was determined iodimetrically. Only fresh solutions were used for the spectrophotometric measurements. A 0.1M aluminum acetate solution was prepared by dissolving reagent grade aluminum acetate in dilute perchloric acid. Interference studies were carried out using reagent grade chemicals such as fructose, sucrose, mannose, tartaric, citric, phosphoric, and oxalic acids. Procedure. A solution for the colorimetric analysis was prepared by the addition of the following stock reagents to a 25-ml volumetric flask in the following order: 5 ml of buffer solution, 1 ml of 0.1M aluminum acetate, 1 ml of 0.002M iron(1II) sulfate or perchlorate, a suitable amount of ascorbic acid or the unknown, and, finally, 1 ml of ferrozine. The contents were diluted to the mark and the absorbance of iron(I1) ferrozine chelate was measured at 562 nm after 2 or 3 minutes. The absorbance stays constant for hours in the presence of pure ascorbic acid and orange, lemon, or grapefruit juices. For the colorimetric procedure involving fresh fruit, after making a hole in it, a few drops were squeezed out and samples of the juice were withdrawn directly by means of a 11 pipet and the analyses were carried out in the usual way. Commercially available juices were analyzed directly. Comparison of the results between the iodimetric and colorimetric methods was carried out using freshly squeezed and centrifuged citrus fruit juice. In the presence of cobalt(II), copper(II), and nickel(II), the amount of ferrozine had to be increased until the absorbance of iron(I1)ferrozine produced remained constant. RESULTS AND DISCUSSION
Reduction of iron(II1) by ascorbic acid and the resulting absorbance of iron(I1)-ferrozine chelate depends on variables such as the molar ratio of iron(I1) to ferrozine, the pH, the time necessary to reach maximum intensity and the presence of various interferences. Since iron(II1) does not form a stable ferrozine chelate, the amount of ferrozine to be added must be at least in 3 to 1 molar ratio with respect to iron(I1)
ANALYTICAL CHEMISTRY, VOL. 44, NO. 2, FEBRUARY 1972
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I
O 7
3.0
-
Ob 6.0
u 1 W
’*O
z
da
IO
TIME
TIME MIN. Figure 1. Absorbance of iron(I1)-ferrozine at 562 nm produced as a function of time at various pH’s
0
21)
4.0
8.0
Figure 3. Absorbance of orange and grapefuit juice in the presence of aluminurn(III), -and and in the absence of aluminum(III),- . - . - . ,as a function of time
-
produced in the reaction. In the presence of copper(II), cobalt(II), and nickel(II), sufficient amount of ferrozine must be present to complex these ions. In actual determinations, we have used ferrozine in 10 to 1 ratio with respect to iron(I1) formed in solution. The reduction of iron(II1) by ascorbic acid and the formation of iron(I1)-ferrozine chelate depends on pH. The development of color is almost instantaneous in acetic acidacetate buffer solutions in the p H range of 3 to 6. At pH’s below 3, the formation of the iron(I1)-ferrozine chelate, Fe(I1) (ferz>a4-,is hindered by the hydrogen ion competition, while at pH’s higher than 7, iron(II1) ion as a hydroxo complex is not reduced readily. The effect of pH on the formation of Fe(II)(fer~)~~is shown in Figures 1 and 2. The actual optimum p H range for the analysis of citrus fruits is also affected by other substances present. Because of these interferences, color does not develop readily and intensity continues to increase with time as shown in Figure 3. These observations suggest that iron(II1) is removed from the reaction by a chelation with some chemical species present in citrus fruit and, furthermore, that the chelate is slowly 380
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Table 1. Analysis of Ascorbic Acid in Citrus Fruits. AbsorbAmount Amount taken ance found,& Sample ml pgb observed pg Blank 0.00 0.00 0.045 0.00 Ascorbic acid 0.50 7.95 7.88 0.145 1.00 16.16 15.90 0,250 31.80 2.00 0.450 31.93 36.2 0.495 35.5 Orange juice 0.050 0.100 70.1 72.4 0.939 23.7 24.3 0.345 Lemon juice 0,050 0.100 0.645 47.4 48.6 24.6 Grapefruit juice 0.050 24.0 0.350 24.6 24.9 Hi-C 0.100 0.360 37.9 Cranberry juiced 0.100 38.9 0,475 a Ascorbic acid was determined at pH 3.0 in the presence of aluminum(II1) using 1-cm cell. * Amount of ascorbic acid taken was determined iodimetrically using macro samples. Amount found was calculated by the Beer’s law expression
PH
Figure 2. Absorbance of iron(I1)-ferrozine produced as a function of pH after 6 minutes
MIN
pg = ( A X
176 X 1000 X 25)/55800
where A is the absorbance corrected for the blank, 176 is the molecular weight of ascorbic acid, 55800 is the effective molar absorptivity for the ascorbic acid in the reaction with iron(III), and 25 is the final volume of the solution. Absorbance was measured after a reaction time of 2 minutes.
oxidized by iron(III), -iron(II)-ferrozine couple. Addition of citric or tartaric acids to iron(II1)-ferrozine solution at p H 4 produced a very slow formation of F e ( I I ) ( f e r ~ ) ~ ~ species as a result of the direct oxidation of these acids, an oxidation facilitated by light. In the presence of citric and tartaric acids in amounts greater than iron(III), the reaction of ascorbic acid is retarded. This interference can be diminished by lowering the pH to 3 and can be eliminated by the addition of aluminum or lanthanum in amounts approximately 100 times that of iron(II1). Aluminum chelates with citric, tartaric, or similar hydroxy acids leave iron(II1) available for the reaction. Actual analyses of ascorbic acid and citrus fruit juices are summarized in Table I. Analysis of pure ascorbic acid indicates that the oxidation of ascorbic acid by iron(II1) and the formation of iron(I1)-ferrozine chelate proceeds as follows :
ANALYTICAL CHEMISTRY, VOL. 44, NO. 2, FEBRUARY 1972
+
C6HsOs 2Fe(III) Fe(I1)
+
+ 2H+ + 2Fe(II)
CeHeOs
+ 3 ferz2-
+ Fe(II)(fer~)~~-
Thus for each ascorbic acid oxidized there are two chelates of Fe(II)(ferz), 4- produced yielding effective molar absorptivity of 55800 which is the determining factor in the limit of detection. Pure ascorbic acid in the amounts as low as 5 pg/25 ml and in the amounts greater than 10 pg/25 ml can easily be determined with a relative precision of *2% and *l%, respectively. In general the calculated absorbance values using iodimetric method and macro samples are higher by approximately 2 to 4 % as compared to the values obtained by the iron(II1)-ferrozine method. This discrepancy may be due to the inhomogeneity of the sample, since in both cases fresh citrus fruit after centrifugation contains some cellular fragments. Sucrose, glucose, mannose, fructose, and formaldehyde do not interfere. Tartaric and citric acids chelate with iron(II1) and slow down the reaction of ascorbic acid. The interference of these acids is overcome by the addition
of aluminum ion. The analysis of commercial cranberry juice is unsatisfactory by the iron(II1)-ferrozine method, since the developed color changes with time. However, the absorbance values taken after a reaction time of 2 to 3 minutes seem to correspond to the ascorbic acid present in the cranberry juice. Oxalic acid in the amounts as high as 10 times that of iron(II1) has no effect. Phosphate in the amounts approximately 100 times that of iron(II1) does not interfere. Copper(II), cobalt(II), and nickel(I1) form chelates with ferrozine. Thus, in the presence of these ions the amount of ferrozine should be increased until there is no further change of the iron(I1)-ferrozine color intensity. In actual analysis, it is advisable to use an excess of ferrozine. Determination of ascorbic acid in the presence of iron(I1) can be achieved by passing the solution through a cation exchange resin and analyzing the eluent for the ascorbic acid.
RECEIVED for review May 24, 1971. Accepted August 20, 1971.
Cyanide Extraction and Electrodisposition of Trace Amounts of Radioactive Silver from Large Biological Samples V. F. Hodge and T.R. Folsom Scripps Institution of Oceanography, Uniuersity of California, San Diego, La Jolla, Calif. 92037
THESTUDY OF RADIOACTIVE silver in the marine biosphere was recently given new impetus when the long-lived nuclide losmAg(tliz= 127 yr) was found in the livers of several species of marine animals caught in 1964-1965 (1). It became apparent that if llomAg(fl/* = 253 d) could be measured along with losmAg,the ratio of these two nuclides might be useful in some cases in identifying the origin of the radiosilver. These nuclides may enter the biosphere as fallout from nuclear weapons or as pollution from nuclear generating stations or nuclear powered vessels. As part of a continuing study of radioactivity in the marine environment, attempts were made to measure, without chemical pretreatment, losmAgand l1ornAgin the livers of albacore tuna collected during the summers of 1968 and 1970. Although the available multiparameter NaI spectrometer can detect concentrations of the two silver nuclides as low as 1.0 =t 0.2 pCi/sample, other fallout activities such as 6oCoand 65Zn,which were generally present in the livers in relatively high amounts, interfered with the identification of the two silver isotopes in the low concentrations that were encountered. Therefore, a method was sought that would concentrate trace amounts of radiosilver from kilogram quantities of wet tissue without loss, and that would at the same time eliminate interferences from 6OCo and 6jZn. Wet ashing of large organic samples has been difficult because the process requires many gallons of acid and often weeks of regular attention. Dry ashing also had to be rejected since there are reports that silver can be lost at temperatures
as low as 100 “C (2). Therefore, we chose to utilize the complexing action of the cyanide ion that has long been exploited by the mining and electroplating arts. For example, very small amounts of fallout silver (on the order of 10-l6 mole) are readily recovered from kilogram amounts of fresh liver tissue by simply stirring the comminuted liver into a basic cyanide solution which contains a gram of silver ion, and then allowing about two days for the added silver to exchange with the radioactive traces. The silver is then plated directly from the slurry onto platinum strips. The silver may be stripped from the platinum with nitric acid and finally collected as its chloride precipitate. Thus the radiosilver in a large organism can be readily concentrated into a small volume and brought close to a small detector such as a solid-state Ge(Li) detector. Moreover, this procedure is equally effective with slurries of fresh liver, liver long preserved in formalin, and also biological samples reduced at 250 “C to a black char.
(1) T. R. Folsom, R. Grismore, and D. R. Young, Nature, 227, 941 (1970).
(2) T. T. Gorsuch, “The Destruction of Organic Matter,” Pergamon Press Ltd., New York, N.Y., 1970, p 66.
EXPERIMENTAL
Cyanide Extraction. All samples were ground in a Waring Blendor with sufficient water to achieve efficient blending action. For example, 2 kg of wet albacore liver tissue yielded approximately 5 liters of blended material. To this slurry, which was contained in an 11-liter polyethylene bucket, approximately 300 ml of 5M NaOH were added to bring the pH to 9. Then, 2 liters of a cyanide mixture were added. (Caution: All procedures involving cyanide were carried out
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