ACTIVE GLUCOSE* BT CHARLES E . CLlFTON AXD JOHN M, ORT
I n the complex equilibrium of solutions of sugar, there seem to be certain forms of the sugars which, although usually present only in traces, are, nevertheless, reactive. Levenea has reviewed the evidence for “active glucose.” Additional facts, discovered by oxidation-potential technic, have been reported recently by one of us12 (Ort). These data can hardly be understood without the assumption of the existence of such an active sugar. Since this technic makes possible the detection of exceedingly slight chemical changes, it appeared possible to determine by it the actual amount of active glucose present. These determinations, carried out with t,he same apparatus and general procedure as previously published,” are herein described. Solutions were first investigated which had stood for fifteen hours while being constantly stirred by pure nitrogen gas at’ the rate of two bubbles a second, as in the previous experiments.** At this time drifts in reduction potentials were quite slow, indicating that something in the nature of preliminary equilibrium had been established. Much longer standing would seem to risk the appreciable accumulation of the products of secondary and side reactions which go on slowly during exposure of the sugar even to weak alkali. Solutions of glucose in sodium hydroxide-sodium phosphate buffers a t pH I O were used in all experiments. They were made up by adding the designated weight of sugar to IOO C.C. of the appropriate buffer. To get an idea of the order of magnitude of the amount of the active form involved, the following preliminary experiments were done. A constant amount of oxidant, 0.j C.C. of a 0.03 percent hydrogen peroxide solution, was added to different amounts of glucose. The drifts of oxidation potentials with time are represented by the curves in Fig. I . As the amount of sugar increases, three things happen: ( I ) the reduction potentials acquired, after standing fifteen hours, drop steadily toward a limiting value; ( 2 ) the highest or peak oxidation potentials reached after additions of oxidant are also lower; and (3) if the amount of sugar is large enough there may not be a great’ change in potential following the addition of this amount of oxidant a t this rate of stirring. The first of these facts suggests that, as the concentration of the sugar increases, the concentration of the electromotively active reductant also increases. Since the peak potentials are later replaced by the original reduction potentials, the total glucose present does not exist in a simple oxidationreduction system in reversible equilibrium with its oxidation products. The * Division of Physics and Biophysical Research, The Mayo Foundation, Rochester, Minnesota. The data presented in this paper are taken from a thesis submitted by Charles
E. Clifton to the Faculty of the University of Minnesota in partial fulfillment of the re-
quirements for the degree of Ph.D., 1928. Read before the American Chemical Society, Swampscott, Masaachusetts, September, 1928.
856
CHARLES E. CLIFTON A S D JOHN M. ORT
step from the active reductant to its first oxidation product may be reversible but the system as a whole under these conditions is controlled preponderantly by a force entirely from outside this reversible part. The addition of a small amount of an oxidant to a simple reversible system would be followed by an almost immediate rise in potential to a thermodynamically calculable value where it would remain. The second and third facts also suggest that there is some quantitative reaction between the amount of sugar present and the oxidant added. This conclusion is also strengthened by the fact that if varying amounts of the peroxide are added to solutions containing 60 gm. of glucose, a similar set
t
"r cc
3/90 %
Time In r n ~ s ~ t e s HrOe added
FIG.I
A constant amount of oxidant added to increasing amounts of glucose.
of curves could be plotted. That is, 0.5 C.C.and 1.0 C.C. of peroxide produced relatively small rises in potentials at this rate of stirring, whereas 2.0 C.C.and 3.0 C.C. brought about much larger ones. It was evident from these experiments that the amount of active reductant present even in concentrated solutions of glucose is very small. I n order to determine it by potentiometric titration an exceedingly dilute standard-oxidizing solution was required. Dilute hydrogen peroxide solutions seemed unsuitable, since part of their total oxidizing capacity is due to dissolved oxygen and part to hydrogen peroxide. It seemed undesirable to titrate with a solution containing two different oxidants, regardless of how similar their reactions might be, especially when one of them, the peroxide must go through intermediate reactions to furnish its active oxidant, and we were titrating a reductant that was being constantly replaced when removed by oxidation. Fortunately it happens that the solubility of air which is free of carbon dioxide in distilled water under ordinary laboratory conditions is just such as to make the solution 0.001N with respect to oxygen. Hence I C.C. =
85;
ACTIVE GLUCOSE
O.OOOOOI gm. equivalent of oxidant. Such distilled water was therefore used in the titrations described here. I n this connection it should be noted that when so C.C. of boiled distilled water was added to the solutions of sugar, a change in reduction potentials was not observed. Likewise a hydrogen electrode in these solutions remained at a constant potential after such an addition, indicating that this amount of dilution with oxygen free water did not appreciably affect the pH. Hence we conclude that all effects herein reported are due to the dissolved oxygen. Solutions containing 60 gm. of glucose were selected for titration. When the sugar was dissolved in the buffer the volume was approximately 139 C.C. and the pH was 10.00 0 . 0 2 . Because these solutions were more viscous
*
I
I
I
I
I
1 2 3 4 5 ficlcrog.ram equivalents of oxidant added
FIG.2 Potentiometric titration curve of active glucose in,6o gm. of the sugar.
t h a n the dilute solutions hitherto used in these studies, the rate of bubbles of nitrogen gas was changed from two to five bubbles a second during the fifteen hours' standing. Even a t this faster rate, it was found that the oxidant was not mixed rapidly enough throughout the entire solution. Therefore, just after an addition, this rate was increased for a quarter of a minute to twenty bubbles a second and then again put back to five bubbles a second. This procedure gave much more constant results than were observed when the rate was left unchanged. Apparently, a t the slower rates the oxygen in the distilled water reacted with the sugar on top and recovery could take place before this part of the liquid penetrated the denser and more viscous layer of solution around the electrodes a t the bottom of the tubes. It has been reported for much more dilute solutions,'* that sudden changes in the rate of bubbles produced sudden changes in the reduction potentials
858
CHARLES E . CLIFTOS A S D J O H S AI. OR?
of those solutions. However, for the 60 gni. .elutions, it was observed in control experiments that such a temporary shift did not produce any appreciable change in the potential readings. I n the dilute solutions formerly used there could have been hardly more than infinitesimal amounts of the active reductant present. Any sudden change of experimental conditions might be expected to produce noticeable fluctuations in reduction pot,entials in such solutions. But, whatever the reason, the temporary violent bubbling did not of itself affect the potentials of the 60 gm. solutions. I t did give rapid mixing throughout the entire solution and much more constant results, and was therefore used in experiments reported here. Fig. z shows the potentiometric titration curve plotted after a consideration of the results of several hundred experiments. The point of inflection of this curve is at 2.j c.c., which fixes the amount of the most active reductant at o.oooooz5 gm. equivalent. The potential for zero added oxidant is shown as -0.406 volt. This is the potential acquired by platinum electrodes after they had stood in these solutions for fifteen hours under the conditions described. This value is an average obtained from several hundred experiments of which more than half gave potentials within j millivolts of this figure and over 80 percent within I O millivolts. This value did not seem to be affected much by the previous treatment of the electrodes, that is, whether they were heated to incandescence just previous to the run, treated with sulfuric and chromic acids or aqua regia, or simply rinsed in distilled water between runs. The behavior of the electrodes after an addition of oxidant was found to be somewhat affected by this previous treatment, although for any series of runs in which this treatment was kept constant, the inflection of the titration curve came always at about 2.j C.C. The curve shown here is from experiments in which the electrodes were simply rinsed in distilled water between runs. The potentials for I c.c., z c.c., and so forth, of added oxidant are the peak voltages reached after the additions of these amounts of aerated, distilled water. In a few moments after reaching these peaks the potentials would begin to drift back toward the - 0.406 value which they would reach in several hours. After this recovery, the subsequent addition of an identical amount of oxidant would usually result in a slightly higher peak value and therefore for the cmve here shown a new glucose solution was used for each addition. However, if after constant preliminary additions of this order of magnitude and the subsequent recoveries, these I c.c., z c.c., and so forth additions were made and the somewhat higher peaks resulting were plotted, the inflection of the curve still came near 2 . 5 C.C. I n this connection it should be noted that these solutions were commonly made up exposed to the laboratory air. Since the stirring and exposure during solution could not be exactly uniform it might be thought that the resultant fifteen-hour voltages or the peaks would be affected. ‘This was found not to be the case for even when compressed air or pure oxygen gas was
ACTIVE GLUCOSE
859
bubbled vigorously through the liquid during and after the solution of the sugar, for a total period of ten minutes, the results on standing and titration were not found to be appreciably affected. Certain solutions were prepared by placing the dry glucose in the electrode chamber and removing air from the space above the sugar by flow of nitrogen gas for an hour. The buffer which had been deoxidized by nitrogen was then added from a special pipette, care being taken to exclude air from both buffer and electrode chamber during the addition. Although the immediate electrode potentials of such solutions, measured just after the last bit of solid dissolved, were somewhat more negative than the ones prepared in air, the voltages acquired on standing and titration were practically the same as the usual values. Therefore we con+ 100 c.c NaOH-NaeHPO4 buffer: ph -10 0.0001 gram equivalent hsscdved Oe added at t 60 grams glucose
Recovery to high reducing intensities after additions of oxidant. All effects are repeatable.
clude that exposure to the air for the short time necessary to dissolve the sugar and to fit the chambers into the apparatus before the solutions are deoxidized by nitrogen gas does not affect our results. The values for the peak voltages were determined as follows: Excluding freak results which occasionally occur, the actual points for the voltages as observed in about a hundred separate runs were plotted. Although the scattering of these values was greater than that observed for the potentials at the end of the fifteen-hour period, it could be seen that the points clustered together around the average values given in Fig. 2 . On drawkg a curve through these clusters in a way fairly to represent the trend of events, it was obvious that the point of inflection of the curve could be tixed t o within 0 . 2 5 c.c., that is, 0.00000025 gm. equivalent of oxidant or I O per cent of the amount needed to combine with the active sugar. If the first oxidation change of glucose be assumed to be bivalent (loss of 2 H+ions) then the total sugar present is 0.67 gm. equivalent and the active sugar is present I part in 266,000 parts f I O percent. If the time drifts of potentials in these concentrated glucose solutions are observed continuously, beginning as soon after starting the deaeration with
860
CHARLES E . CLIFTON AND JOHN M . ORT
nitrogen gas as the salt bridges and other connections can be made, drifts will be noticed (Fig. 3 ) . The variations immediately after starting are, of course, rapid and erratic. The first peak in the curve will not always be observed. From then on, however, the curve represents the consistent course of events. The slow slope downward after about three hours represents the preliminary equilibrium which has been mentioned. Aubel, Genevois and Wurmser,*and Wurmser’4also determined the potentials which reducing sugars, buffered with phosphates, reach in the absence of oxygen at different temperatures and hydrogen-ion concentrations. Their results, when allowance is made for differences in concentration, and so forth, seem to be in fair agreement with our potentials as observed on standing, although their experiments were carried on for a much longer time. At the points marked with arrows O.OOOOI gm. equivalent of oxidant ( I O C.C.distilled water) was added. This is, of course, a considerable excess of oxidant over active sugar and the rise in oxidation potential is considerable. Kevertheless the recovery is complete and rather rapid a t first. The ability to repeat is proof of the existence of a dynamic equilibrium furnishing a fresh supply of active glucose when that present is destroyed. I n attempting to get an idea of the rate a t which the active glucose appears, solutions were next investigated at six hours after the nitrogen gas was turned on. The starting potential3 were, of course, slightly higher than at fifteen hours. The peaks were also found to be somewhat higher but when plotted gave a curve similar to that in Fig. 2 with the inflection at 2.5 C.C. Therefore the amount of active reductant a t this time is also about I part in 266,000 parts of total glucose. The slow progress of secondary and side reactions have doubtless resulted in sufficient byproducts at fifteen hours slightly to affect the actual potentials observed, although the fact that their relative values after additions of oxidant are the same shows that the amount of the most intense reductant is practically the same a t six hours as at fifteen hours. Similar observations were made a t three and a half hours. Therefore we conclude that a t least as soon as the first fluctuations are over there exists in these glucose solutions I part in 266,000 parts of a special or active form. Whether these facts and figures also hold true for fresh solutions immediately after dissolving in the open air cannot of course be decided a t present. This cannot be tested by our technic until the voltage drifts are a t least fairly slow.
A consideration of Fig. z shows that its steep part and point of inflection come in the middle of the curve. A titration curve of a typical reversible oxidation-reduction system also has its point of inflection in the middle but its steep vertical parts are a t the ends. This fact and the fact that the POtentials acquired after repeated identical additions of oxidant are only slightly different show that the first oxidation product of glucose is not electromotively active at the range of potentials belonging to active glucose. Hence we have, in this respect, a condition similar to acid-base titration in which the product, water, does not appreciably affect the electrode potentials since H+ or OH- ions from other sources are so much more numerous. There-
ACTIVE GLUCOSE
861
fore the titration curve of this reductant is of a form similar to acid-base titration curves. Since it is well known that traces of iron have considerable catalytic effect on the oxidation of sugars by air, the influence of small amounts of iron on these titrations was investigated. For the point might be raised that the titrations were simply determining the amount of iron present undergoing oxidation and reduction. Quantitative determinations of iron in these solutions revealed its presence in amounts of the same order of magnitude as the value noted for active sugar. Solutions were, therefore, prepared in which one and two times the amount of iron as FeC13.6H20,equivalent to the amount of active glucose present, was added. The reduction potential of these solutions after fifteen hours was also -0.406 volt. When they were titrated potentiometrically it was found that the curve for one equivalent of iron falls slightly under the curve in Fig. 2 for the sample of Merck’s CP dextrose used and that the curve for two equivalents is slightly lower still. However, their points of inflection both came a t approximately 2 . 5 C.C. and not a t 5.0 and 7.5 c.c., as might be expected if valence changes in iron were alone being titrated. It was noted that the presence of iron seemed to make the potentials a little more definite and more closely reproducible. Hence we conclude that the glucose equilibrium has present in it a form which is chemically much more active than the main bulk and that although its activity in connection with oxidants may be somewhat increased by small amounts of iron, its amount is not; that is inherent within the equilibrium of the sugar itself. The chemical identity of the active form or the mechanism by which this form is produced from ordinary glucose is of course not revealed by these experiments. Since sugar oxidation is favored by increasing alkali whereas in general an oxidant exhibits a higher oxidation intensity in acid solutions, it is evident that the preponderant effect of the OH- ions is on the sugar molecule. The simplest way in which they could act would be by the removal of H+ ions. However, the process is obviously not one of simple neutralization, for the glucose is not appreciably converted into sodium salt even in the presence of a large excess of sodium hydroxide. Yet the active reductant may be simply the negatively charged residues or ions which are left after a splitting off of H+ ions which can only occur when the configuration of the glucose molecule is temporarily strained so as to favor this. The minute amount of the active form present shows that the probability of any given molecule existing in a state to favor this change is small and also that the average life of such an ion is short. The tendency to recombine and revert to the inactive state even though comparatively few H+ ions are present in alkaline solutions must be great; otherwise more of the active form would accumulate and sugar would be oxidized more easily. If this picture of the active form is true, it is evident that the excess electrons of these ions must be in a moreexposedposition on the molecule than those on the common anions because these latter do not exert so intense a force of reduction.
862
CHARLES E. CLIFTOK AND JOHN M. ORT
A few preliminary determinations made on glucose solutions containing only 30 gm. of glucose instead of 60 gm. seem to be in accord with this postulate. That is, they gave a proportion slightly greater than I part in 266,000 parts, although the difference was near the limit of experimental error in determining such small amounts of chemical change. I n a more dilute solution the degree of ionization, if the equilibrium point may be so named, would be expected to be greater. From the point of view of the organic chemistry involved, other easily oxidizable forms of glucose might be postulated. Originally the aldehydic form was supposed to be the “active glucose.” If ordinary glucose exists in the form of a lactal ring this might break and form a free radical. Or an enediol, or its methylenol cleavage product, may be present and active in this respect. For a consideration of these and other theories of the possible constitution of active glucose the works of Nef,gJo Evans and his associa t e ~ ,4 ,~ , Amick, Hudson and Dale, Stieglitz, and especially the review by Levene*should be consulted. The purpose of this paper is merely t o present further quantitative evidence that there is an active glucose. Whatever its form, it exists in a definite although minute amount in equilibrium with the main bulk of ordinary and comparatively inactive glucose. 5v
summary
Further evidence is submitted that in solutions of glucose there exists a small but definite amount of a very powerful reductant or “active glucose,” which alone is responsible for the reduction intensities developed in these solutions. It is shown that this active form is almost instantly destroyed by even a mild oxidant and exists in dynamic equilibrium with the main bulk of inactive or ordinary glucose, since it is in time replaced when so destroyed. Under the conditions of the experiments herein reported the amount of active glucose was found to be I part in 266,000 parts I O percent.
+
Bibliography C. A. Amick: J. Phys. Chem., 31, 1441-1477 (1927). The copper number for glucose. % E .Aubel, L. Genevois and R. Wurmser: Compt. rend. Acad. d. Sc., 184, 407-409 (1927). Sur le potential apparent des solutions de sucres rbducteurs. W. L. Evans, C. $. Buehler, C. D. Looker, R. A. Crawford and C. W. Hall: J. Am. Chem. SOC., 47, 3085-3098 (1925). The mechanism of carbohydrate oxidation. I. W. L. Evans and C. A. Buehler: J. Am. Chem. SOC.,47, 3298-3101 (1925). The me.~ chanism of carbohydrate oxidation. 11. W. L. Evans, R. H. Edgar and J. P. Hoff: J. Am. Chem. Soc., 48, 2665-2677 (1926). The mechanism of carbohydrate oxidation. IV. 6 W. L. Evans and J. E. Hutchman: J. Am. Chem SOC.,50, 1496-1503 (1928). The mechanism of carbohydrate oxidation. VIII. ’ C. S. Hudson and J. K. Dale: J. Am. Chem. Soc., 39, 320.328 (1917 ) . Studies on the forms of d- glucose and their mutarotation. P. A. Levene: Chem. Rev., 5, 1-16 (1928). Active Glucose. 9 J. U. Kef: Ann.d.Chem.,357-358, 214-312 (1907). Veber dasVerhaltenderZuckerarten gegen die Fehlingsche Losung, sowie gegen andere Oxydations mitrel. Io J. U. Nef: Ann. d. Chem., 903, 204-383 (1914). Dissoziationsvorgange in der Zuckerg w w I1 J. M. Ort: J. Opt. Soc. America and Rev. Sci. Inst., 13, 603-608 (1926). Apparatus for determining oxidation-reduction potentials. J. M. Ort: J. Phys. Chem., 33, 825-841 (1929). Vltraviolet light, insulin, and amino acid catalysis. I 3 Julius Stieglite: Proc. Inst. Med. Chicago, 1, 41-50 (1916). The oxidation of carbohydrates. I4 Ren6 Wurmser: Compt. rend. Acad. d. Sc., 185, 1038-1041 (1927). Sur le potentiel apparent des solution de glucose.
