Photochemistry of Fluorescein Dyes. - The Journal of Physical

PHOTOCHEMISTRY OF FLUORESCEIN DYES' H. F. BLUM AND C. R. SPEALMAN Division of Phusiology, University of California Medical School, Berkeley, California Received April 16, 1955

It is well-known that the fluorescein dyes photosensitize the oxidation of readily oxidizable substances by molecular oxygen (Gaffron, 1926; Carter, 1928). Several investigators have shown that peroxide is formed under proper conditions when the dyes are irradiated in solutions which do not contain readily oxidizable substances (Weigert, 1912; Gaffron, 1926; Blum, 1930). The nature of the peroxide has not been clearly shown and the question as to whether it represents an intermediate step in the oxidative process has not been definitely settled. In the following study an attempt has been made to determine the nature of the peroxide formed when fluorescein dyes are irradiated in aqueous solutions (Weigert, 1912; Blum, 1930). Four possible sources for the peroxide have been considered: (1) impurities present introduced in the preparation of the dye, (2) the dye itsell", (3) breakdown products of the dye, since the dye is bleached concomitantly with the formation of peroxide, (4)water, resulting in the formation of hydrogen peroxide. EXPERIMENTAL

The first possibility was tested by preparing fluorescein from very pure intermediates, resorcinol and phthalic anhydride, and purifying carefully by recrystallization. The fluorescein was dissolved by the addition of an M and irradiated equivalent amount of sodium hydroxide, made up to in a thin layer with a 200-watt tungsten filament lamp a t a few centimeters distance. After irradiation for about twenty hours, such solutions conN hydrogen peroxide, as determined iodometrically. tained up to Since the concentration of peroxide formed is of the order of that of the dye, it is highly improbable that it should be a peroxide of any impurities present.

Rate studies It was thought that simultaneous studies of the rates of bleaching and of peroxide formation should be of value in differentiating between the other three possibilities mentioned above. Such studies offer considerable diffiThis investigation was assisted by a grant from the Board of Research of the Cniversity of California. 1123

1124

H. F. BLUM AND C. R. SPEALMAN

culties, since it is necessary t o use fairly intense irradiation over considerable periods of time in order to obtain measurable quantities of peroxide. The method finally adopted was to irradiate successive samples with the same source, a given 200-watt tungsten filament lamp, for varying periods, analysis for peroxide and determination of the degree of bleaching being made for each sample at the end of the irradiation period. Figure 1 is a rate curve obtained in this way. The samples consisted of' 100 cc. of a 0.00028 M solution of eosin2 in distilled water, this concentration being

6

/O

20

30

40

Duration o f irradiation , hrs. FIQ. 1. RATEOF BLEACHING AND PEROXIDE FORMATION FOR IRRADIATED EOSIN Dye concentrations are in moles per liter; peroxide concentrations in equivalents per liter. Numbers of points indicate the order of the determinations. A t 2 hours (point 5) a test for peroxide was obtained, but the quantity was too small for accurate titration.

selected for convenience in analysis. These samples were placed in a 1000cc. Erlenmeyer flask, forming a layer approximately 1 cm. thick on the bottom of the flask and supported a t a distance of approximately 7 cm. above the filament of an ordinary 200-watt Mazda lamp; the distance between the lamp and the flask was carefully reproduced for each sample. The system was cooled by a blast of air directed across the bottom of the 2

Eosin Y (sodium salt of tetrabromofluorescein) from Coleman and Bell.

PHOTOCHEMISTRY OF FLUORESCEIN DYES

1125

flask above the lamp bulb; temperatures were maintained in this way between 32°C. and 33OC. after the first half hour of irradiation. The scheme has the advantage that the solution is irradiated in a fairly thin layer, which is of importance since oxygen is a necessary component of the reactions and has a very definite effect on their rate (see Weigert, 1912, and below). Provision for shaking to establish uniform distribution of oxygen was found impracticable because of frothing of the dye solution. It was found difficult to obtain sufficiently intense irradiation for volumes of solution large enough to provide for the removal of aliquots adequate for analysis, and such a method is subject to the objection that the thickness of the layer of solution is altered upon the removal of each aliquot. The scheme has the disadvantage that the rates are subject to variations in the intensity of emission of the source which definitely varies with the age of the lamp, the extent of this variation being different for different lamps. Some of our determinations were found to be untrustworthy because of this, but the curve shown in figure l , all the points of which were obtained with the same lamp, seems to be free from this objection. That the intensity of the irradiation did not greatly alter in this case is indicated by the accuracy with which certain of the determinations could be duplicated, and that the points obtained in random order (indicated in the figure) fall on definite curves. Furthermore, the general shape of the dye-concentration curve, i.e., having a sharp break downward a t about the time maximum peroxide concentration is reached, was found to be typical of all the curves obtained. The break occurred later with less intense irradiation. The dye concentrations were determined spectrophotometrically. Although Beer's law does not hold exactly for the fluorescein dyes (see Pringsheim, 1928, p. 226), it was found by determining the absorption of various dilutions of eosin a t 5200 A. U., approximately the maximum absorption of eosin, that the law could be applied safely for the concentrations a t which our determinations were made. Since some shifting in the absorption spectrum of eosin occurs upon irradiation (see Wood), the absorption spectra of a certain number of irradiated samples were determined; these appear in figure 2. It will be seen from this figure that any shift in the region of maximum absorption is not of great importance in these cases, nor do the curves appear greatly affected by the presence of colored breakdown products. Thus it seems safe to use the absorption a t 5200 A. U. as an index for calculating the concentration of dye in the irradiated samples; it is probably more correct to consider the concentrations so obtained as maximum dye concentrations, since in the cases of extreme bleaching the presence of breakdown products would probably tend to increase the total absorption. The photo-compound described by Wood (1922) on the basis of absorp-

1126


tion spectrum measurements was studied in the case of fluorescein, but does not appear to play an important part in the reactions herein discussed. It is not formed in the absence of oxygen, but is not a reversible oxidation, since there is no return to the original color when reducing substances are added, hence it is not a peroxide of the dye. It may represent a first step in the oxidative bleaching of the dye. Eosin was used in our rat.e determinations because the shift in the absorption spectrum is less pronounced than in the case of fluorescein, and therefore the determinations of concentration should be more accurate.

ii 9

?

0

480

SO0

520

590

m!FIQ.2. ABSORPTION SPECTRA OF IRRADIATED 0.00028 AI EOSIN 0 = non-irradiated dye. Calculations based on absorption of 1 cm. thickness of solution. Higher concentrations diluted 1to 10 for analysis.

The following method was found satisfactory for the determination of peroxide in the small quantities with which we were dealing. Approximately 2 g. of potassium iodide was placed in a 100-cc. volumetric flask with 5 cc. of 6 N hydrochloric acid and about 70 cc. of water. The sample t o be analyzed, usually about 20 cc., was added by pipette and the mixture made up to volume. After standing for 30 minutes,'the contents of the flask were emptied into a 500-cc. Erlenmeyer flask and titrated with 0.001 N sodium thiosulfate against starch indicator. The sodium thiosulfate was prepared frequently by dilution from a standardized 0.1 N solution. The end point is sometimes difficult because of the presence of

PHOTOCHEMISTRY OF FLUORESCEIN DYES

1127

the dye precipitate formed when the acid is added to the dye solution, but the error from this source is not great. The method was checked by titration of solutions of known hydrogen peroxide concentration diluted to the same approximate concentrations as those of our irradiated dye samples, both with and without the addition of non-irradiated dye. The error was estimated as within 5 per cent a t the lowest peroxide concentrations a t which determinations were attempted; a t higher concentrations it was much less. One source of error is the oxidation of iodide ion by atmospheric oxygen, resulting in high values for hydrogen peroxide; this is avoided by keeping the solution in the narrow necked volumetric flask during the 30 minute period before the titration. The rate curves (figure 1) are subject to certain definite criticisms. Hydrogen-ion concentration was not controlled, because of the fact that electrolytes slow the reaction and hence it was found impracticable t o use buffers. The oxygen tension was likewise not controlled for the reasons mentioned above. Both these factors may have an influence on the shape of the rate curves. Nevertheless two facts stand out quite distinctly. After a considerable period of irradiation, the concentration of peroxide may reach a value considerably greater than that of the dye remaining in the solution, e.g., in figure 1 beyond 20 hours the concentration of peroxide is, mole for mole, a t least five times as great as the concentration of dye, and thus at least a great part of the peroxide would have to be present in some other form than dye peroxide. The increase of peroxide does not appear to show any direct relationship to the destruction of the dye, indicating that the peroxide is not a compound of breakdown products of the dye.

Behavior of the peToxide These facts point toward the fourth of the possibilities outlined above, namely, that the peroxide formed upon irradiation of the dyes is hydrogen peroxide. Comparison of the behavior of this peroxide with hydrogen peroxide was therefore made. Separation of the peroxide from the dye. If the peroxide is hydrogen peroxide it should be possible to separate it from the dye; it was found possible to accomplish this by precipitating the dye as its acid or as an insoluble lake by the addition of Ag+. When a strong acid (say 0.1 N hydrochloric acid) is added to a sample of an irradiated fluorescein dye, the dye is precipitated and can be filtered off. If a filtrate thus prepared was analyzed for peroxide, i t was found to contain quant,itativelythe same amount as a sample from which the dye had not been separated. When the precipitate was redissolved with sodium hydroxide and acidified, no trace of peroxide could be detected. When silver nitrate is added in excess to an irradiated fluorescein dye

1128


solution, the dye is precipitated as an insoluble lake. After centrifuging to throw down the lake, removing the Ag+ from the decanted solution by the addition of sodium chloride, and filtering, it was found that over 80 per cent of the peroxide appeared in the filtrate, and no trace in the precipitate. Similarly, when a dye solution to which hydrogen peroxide had been added was treated in the same manner, approximately the same percentage of peroxide was recovered in the filtrate. Fluorescein, eosin, ~ h l o x i n , ~ and erythrosin4 all gave similar results. T h e chromic acid reaction. This test depends upon the formation of the blue perchromic acid, which may be separated from the aqueous solution because of its solubility in ether. The test is subject to certain difficulties which are discussed by Noyes, Bray, and Spear (reference 6, p. 545). The procedure used by us was as follows: 100 cc. of a 0.001 M solution of eosin was irradiated for 15hours, a t the end of which time a 20-cc. sample of the solution was titrated iodometrically for peroxide, a concentration of approximately 0.001 N being found. Five cc. of 6 N nitric acid was added to 50 cc. of this solution. This precipitated the dye, which was filtered off, To the acid filtrate in a 125-cc. Erlenmeyer flask was added 5 cc. of pure ether, forming a layer a few millimeters thick on the surface of the liquid. Two blanks, one consisting of 50 cc. of the unradiated dye solution, and one of 50 cc. of this solution plus 5 cc. of 0.01 N hydrogen peroxide, were prepared in the same way. Four cc. of approximately 0.01 N potassium chromate was added to each of the above mixtures. The irradiated dye and the blank containing hydrogen peroxide showed a distinct decolorization of the yellow chromate ion, when compared with the blank without hydrogen peroxide, with a faint blue color in the ether layer when viewed from the side through the total thickness of the layer. The ether layer in the blank remained colorless. The intensity of the blue color was comparable in the irradiated dye sample and the hydrogen peroxide blank. From this test, we may conclude that the irradiated dye solution, so treated, contained a concentration of hydrogen peroxide approximately the same as that in the blank, namely 0.001 N , which compares favorably with the figure for the iodometric titration. Rate of reduction by hydrogen iodide. The rate of reduction of our peroxide by hydrogen iodide was studied as follows: Titration of five 20-cc. samples from 100 cc. of 0.001 M irradiated eosin were made a t varying intervals after adding 4 cc. of 6 N hydrochloric acid and 4 g. of potassium iodide to Bach sample. The intervals covered a period of 30 minutes and provide data for a rate curve for the reduction of the peroxide with hydrogen iodide. Sodium thiosulfate (0,001 N ) was used for the titration. At the end of 30 minutes the curve had reached a constant value, indicating 3

Dichlorofluorescein. Tetraiodofluoreacein.

PHOTOCHEMISTRY O F FLUORESCEIN DYES

1129

a peroxide concentration of approximately 0.001 N . Samples of 0.001 N hydrogen peroxide and 0.001 N hydrogen peroxide plus 0.001 M eosin were similarly treated. The rate curves for the three solutions were identical within the limits of accuracy. Since the rate of reduction of the peroxide in the irradiated dye sample is the same as that of hydrogen peroxide, there can be little doubt that the peroxide is the same. Decomposition by catalase. In all the above tests, except the separation of the dye by means of Ag+, it was necessary to acidify the solution in the course of the procedure, and it is possible that such treatment might bring about the formation of hydrogen peroxide from an existing organic peroxide. The decomposition by catalase was studied in the attempt to avoid this possibility. Catalase is known to decompose hydrogen peroxide a t a rapid rate, whereas it does not, so far as has been studied, decompose organic peroxides or peroxyacids a t a perceptible rate (Stern, 1932). Catalase was prepared from liver according to the method described by Waksman and Davidson (reference 11,p. 253). When 5 cc. of the catalase mixture was added to 25 cc. of irradiated dye solution containing approximately 10-3 N peroxide, the decomposition of the peroxide occurred rapidly at a rate comparable with the decomposition of a similar solution of hydrogen peroxide. When the catalase was inhibited by adding 5 cc. of 10 per cent sodium cyanide to like samples, decomposition of the peroxide did not occur, indicating that the above decomposition was due to the catalase and not to reduction of hydrogen peroxide by reducing substances which might have been present in the catalase mixture. The above experiments indicate that hydrogen peroxide is formed, and since the quantity bears no direct relation to the quantity of dye or of breakdown products, it is reasonable to assume that it is formed from water and oxygen. The following energy calculations show that this is possible. The principal reaction involved would appear to be: 2H20 (1)

+ 02 -+ 2H202 (as);

F298

=

50,180 calories; H = 46,000 calories6

(1)

This gives a thermal equilibrium concentration of hydrogen peroxide of about M (Lewis and Randall, reference 4,pp. 596-7). I n agreement with this we found that a dye solution showed no measurable peroxide formation in the dark over a two week’s period. On the other hand, the hydrogen peroxide concentrations of the order of 10-8 M formed in the light indicate we are dealing with a purely photochemical reaction, and not, for example, a thermal chain reaction initiated by light. Lacking experimental evidence, no scheme will be offered here for the intermediate steps necessary in the above reaction. It is possible that a The data for the calculation of F and H in this paper were obtained from Lewis and Randall (1923). The values of H are taken t o the nearest 1000 calories, which is sufficient for our purposes.

1130


dye peroxide may be formed, but if so, it must break down very rapidly to hydrogen peroxide and dye. The value of 46,000 calories for A H is equivalent to a quantum with wave length 6195 A.U.,6 and this represents the maximum wave length which may bring about this reaction. The absorption spectra of the fluorescein dyes in the visible region are included below this wave length. It is understood that this maximum wave length is only approximate, since some molecules will contain energy greater than the mean value and may be raised to the necessary level by smaller quanta. The reaction :

+ O2= HzOz(as); AF = -31,470 calories

2Hf

(5)

has a large negative free energy but could not account for the concentration of hydrogen peroxide formed in our experiments, since the concentration of H+ limits the concentration of hydrogen peroxide which can be formed to 10-14 M . The mechanism : D

+ O2= 2 0 + D;AH = 128,000 calories' HzO

+ 0 = HzOz

is, likewise, ruled out since reaction 6 would require 128,000 calories, which represents a maximum wave length of less than 3000 A.U., to which our vessel is opaque. Some energy of activation would be required to forward reaction 1, and if this were high, the maximum wave length must have a considerably lower value. Unfortunately the data for calculating this energy of activation is not available; it might be greatly reduced by intermediate steps.

The mechanism of bleachfizg The bleaching which occurs concomitantly with peroxide formation is not the reduction of the dye to its leuco base, as is shown by the fact that the color does not return when oxidizing agents are added. While this bleaching reaction has been treated as a separate reaction above, we have obt'ained evidence that it is dependent upon peroxide formation: (1) In the absence of oxygen neither bleaching nor peroxide formation occurs, e.g., a solution of 0.000614 M eosin was irradiated for 5 hours in an evacuated flask, evacuation being cmried sufficiently low to boil off some of the water. A t the end of the irradiation the volume of solution was found to be 15 per cent reduced owing to the evacuation, its concentration being 0.00074 M eosin. This agrees within the limits of accuracy with a concentration of 7

See Taylor (reference 9) for calculation. ValuefromMecke (1929).


1131

the dye by a reduction of volume of 15 per cent, i.e., 0.00072 M . (2) When a readily oxidizable substance (sodium thiosulfate) is present which may react with the peroxide, bleaching is greatly inhibited, e.g., a sample of 0.0006 M eosin was bleached over 99 per cent during 12 hours irradiation, a sample containing the same concentration of dye plus 0.025 N sodium thiosulfate was bleached less than 20 per cent during a like period of irradiation. In the latter case sodium thiosulfate was oxidized in quantity corresponding to 0.005 equivalent per liter, a quantity equivalent to approximately four times the peroxide concentration equivalent mole for mole to that of the dye originally present. This indicates that the quantity of peroxide formed has no direct relation to the quantityof dye bleached. These facts indicate that the bleaching is not a direct decomposition of the dye molecule following the absorption of a quantum of energy, as does also the fact (see figure 1) that the rate of bleaching in the absence of oxidizable substance other than the dye, is not a straight line function of the number of molecules present. As mentioned above, a dye solution remaining in the dark for two weeks formed no peroxide; no bleaching had occurred in the same time. A similar N hydrogen peroxide likewise showed no bleaching, solution containing but a decrease in peroxide concentration comparable to that of a similar hydrogen peroxide solution without dye. Similarly, an irradiated dye solution shbws a slow decrease in peroxide concentration if maintained in the dark for some hours after irradiation, but no bleaching of the dye. The behavior of the peroxide in the irradiated dye solution is, thus, similar to that of hydrogen peroxide, i.e., it does not react with the dye in the absence of light. On the other hand, hydrogen peroxide rapidly bleaches the dye in strongly alkaline solution. This suggested the possibility that the reaction between hydrogen peroxide and dye may be activated by light, and it was found that dye solution containing hydrogen peroxide bleaches much more rapidly in light than does a similar dye solution without hydrogen peroxide, e.g., 0.0006 M eosin plus 0.005 N hydrogen peroxide bleached 40 per cent in 4 hours; a similar dye solution without hydrogen peroxide bleached only 15 per cent in the same time with the same light source. The activation of the reaction by hydroxyl ion may be considered as due to the reduction of the energy of activation of this reaction; the activation by light as due to the supply of the necessary energy of activation by a light quantum absorbed by the dye molecule. The latter would be more probable than that the hydrogen peroxide molecule is activated by light, since hydrogen peroxide does not absorb appreciably in the spectral region supplied by our light source (Urey et al, 1929). Such a mechanism might readily explain the shape of our rate curve (figure 1). The formation of hydrogen peroxide and the oxidation of the dye by hydrogen peroxide would both be dependent upon the absorbed radiation. The second reaction, being de-

1132

H. F. BLUM A N D C. R. SPEALMAN

pendent upon the first, would proceed much more rapidly as the hydrogen peroxide concentration was increased, which would account for the S-shaped curve. Such an explanation would also account for the findings of Wood (1922) whose experiments must have been conducted under conditions similar to ours, Le., without control of hydrogen-ion concentration or oxygen tension. Wood found that a solution of fluorescein bleached more rapidly when a small portion of it was subjected to concentrated radiation, than when the same quantity of radiation was distributed over a larger surface. In the first case the break in the rate curve would occur earlier than in the second. We actually found that the break occurred earlier with more intense radiation; and if taken a t the proper time, the solution subjected to the more intense radiation would show the greater bleaching. Wood suggested that his results were due to the fact that each dye molecule must absorb two quanta of energy before the decomposition occurs; this assumption is not necessary to explain the facts if the above scheme is adopted.

Inhibition by oxygen Weigert (1912) has shown that peroxide formation is inhibited by high partial pressures of oxygen. We have found that when the oxygen tension is raised from 0.2 atmosphere (air) to one atmosphere, peroxide formation and bleaching are both inhibited for a time. The same treatment inhibits the oxidation of a readily oxidizable substance (sodium thiosulfate) when photosensitized by eosin. The nature of this inhibition is not clear and no hypothesis will be offered a t the present time. SUMMARY

1. The peroxide formed when fluorescein dyes are irradiated in aqueous solution is probably hydrogen peroxide. 2. It is suggested that the bleaching of the dyes by light is an oxidation of the dye by the hydrogen peroxide, this reaction requiring activation of the dye molecule by light. REFERENCES (1) (2) (3) (4) (5) (6)

(7) (8)

BLUM,H. F . : Biol. Bull. 68, 224 (1930). CARTER, C. W.: Biochem. J. 22,575 (1928). GAFFRON, H . : Biochem. Z. 179, 157 (1926). LEWIS,G. N., AND RANDALL, M . : Thermodynamics. McGraw-Hill Book Co., New York (1923). MECKE,R. : Naturwissenschaften 17,996 (1929). NOYES,A. A., BRAY,W. C., AND SPEAR,E. B . : J. Am, Chem. SOC.30,481 (1908). A System of Qualitative Analysis for the Common Elements, 111. PRINGSHEIM, P. : Fluorescenz u. Phosphorescenz. Berlin (1928). STERN,K. G.: Z. physiol. Chem. 209, 176 (1932).


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(9) TAYLOR, H. S.: A Treatise on Physical Chemistry. D. Van Nostrand Co., New York (1931). (10) UREY, H. C., DAWSEY,L. H., A N D RICE,F. 0. : J. Am. Chem. SOC.61,1371 (1929). (11) WAKSMAN,S. A., A N D DAVISON, W. C.: Enzymes. Williams & Wilkins Co., Baltimore (192G). (12) WEIGERT,F.: Nernst Festschrift, p. 464 (1912). (13) WOOD,R. W.: Phil. Mag. 43, 757 (1922).

Photochemistry of Fluorescein Dyes. - The Journal of Physical

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