T H E INHIBITORY ACTION O F ANTHRACENE I N THE AUTOXIDATION OF BENZALDEHYDE* BY HANS L. J. BACKSTROM AND HAROLD A. BEATTY
Introduction On the chain-reaction theory of inhibition, the role of the inhibitor consists in the breaking of reaction chains. I n previous papers’ it has been pointed out that, in autoxidation reactions, the mechanism of ’this process appears to involve an induced oxidation of the inhibitor. I n one case, viz. the inhibition of sodium sulfite oxidation by alcohols, this was definitely established, since it was shown that the alcohol is oxidized along with the sulfite, and that the rate of this induced oxidation is such as to link it directly to the breaking of the chains. Further insight into the mechanism of such a chain-breaking reaction would be of considerable interest as it would serve to throw light on the nature of the chain mechariism itself, a problem about which a t present we know little. The present paper gives the results of an attempt in this direction. As, for various reasons, the inhibition of sulfite oxidation by alcohols does not lend itself very readily to such an attempt, a different case of inhibition was selected for this study, viz. the action of anthracene in the autoxidation of benzaldehyde? Induced oxidations were first observed by Schonbein3 in 1845, and have since been the subject of a great many investigations4 It has been found that, under favorable conditions, for each atom of oxygen taken up by the autoxidizable substance (the inductor), one atom is also taken up by another oxidizable substance (the acceptor) which is present but is not itself attacked by molecular oxygen. This was explained by Engler, Bach and others5 on the theory that in autoxidations the primary reaction product of the inductor is a peroxide which possesses strong oxidizing properties and which can give one atom of oxygen to the acceptor. This theory receives strong support from the fact that in many reactions the intermediate formation of such peroxides can be demonstrated. From the photochemical oxidation of benzaldehyde in acetone solution, Jorissen and van der Beek6 isolated the pure peroxide in 6 3 7 , yield and showed it to be benzoperacid,
* Contribution from the Frick Chemical Laboratory, Princeton University. Backstrom: Medd. K. Vet.-Akad. Nobelinstitut, 6, No. 16 (1927);Trans. Faraday Sac., 24,601 (1928);Alyea and Backstrom: J. Am. Chem. SOC.,51,90 (1929). Backstrorn: J. Am. Chern. Sac., 49, 1460 (1927). Schonbein: J. prakt. Chem., 36, 379 (1845). For new examples and a summary of the literature, see Jorissen and Belinfante: Rec. Trav. chim., 48,711 (1929). Engler and Wild: Ber., 30, 1669 (1897); Bach: Compt. rend., 124, 951 (1897);Bodlander: “Uber langsame Verbrennung,” Ahrens Sammlung, vol. 3,385 (1899). e Jorissen and van der Beek: Rec. Trav. chim., 45, 245 (1926).
IXHIBITORY ACTION O F ANTHRACENE
253 I
previously synthesized by Baeyer and Villiger.' In agreement with the theory, this substance is a powerful oxidizing agent. However, if we apply the Engler-Bach theory to the problem of inhibition, we encounter a serious difficulty: it is not easy to see how an induced reaction of this kind can constitute a chain-breaking mechanism. This appears a t once from the fact that, as was shown by one of US,^ the chain reaction in the autoxidation of benzaldehyde is the formation of the peroxide, which reaction, for want of a better theory, we may write as an energy chain (A = benzaldeEnergy = A' (activated): A' 0 2 = -402'; AOP' A = A02 hyde): h A'; A' O 2 = AOZ', etc. As was pointed out in a later paper,g a subsequent reaction of the deactivated A 0 2 , whether with another molecule of the aldehyde or with a molecule of an added substance, would obviously not cause the chain to be broken. Only a reaction of an added substance with A' or AOz' would interfere with the mechanism of chain propagation. To account for the fact of inhibition it would thus be necessary to modify the peroxide theory of induced oxidation, by assuming that the reaction between the peroxide and the inhibitor (acceptor) is not of the ordinary type of thermal reaction but znvolves a n especially activated form of the peroxide.
+
+
+
+
+
As a matter of fact, the literature on induced oxidations has not been lacking in indications that, a t least in certain cases, the theory needed amplification in some such direction as this. Thus Engler'O states as a well-known fact that there is generally a more energetic transfer of oxygen to the acceptor if the latter is present during the oxidation of the inductor, than if it is added after the peroxide has been formed. In the example which he gives, an indigo solution was found to be somewhat more rapidly decolorized when shaken with a benzene solution of benzaldehyde in an oxygen atmosphere, than when it was shaken, under carbon dioxide, with a corresponding solution of benzoperacid. A more striking case was reported by Jorissen and van der Beek11 who found that the photochemical oxidation of benzaldehyde in dilute chloroform solution a t oo gave benzoperacid in good yield, but a t room temperature the solvent was attacked, yielding phosgene, hydrochloric acid, carbon dioxide and unidentified products. Special experiments showed, however, that benzoperacid scarcely attacks chloroform even a t the boiling point. The results presented in this paper show that the inhibitory action of anthracene in the autoxidation of benzaldehyde is correlated with an induced oxidation of the inhibitor, and that the peroxide theory, t o account for this reaction, has to be modified in the direction indicated above. 7Baeyer and Villiger: Ber., 33, 1569 (1900). 8 Backstrorn: Medd. K. Vet.-Akad. Nobelinstitut, 6 , No. I j (1927). 9 Backstrorn: Ibid., 6, No. 16 (1927). 'OEngler: Z. Elektrochernie: 18, 945 (1912). l1 Jorissen and van der Beek: Rec. Trav. chim., 46, 42 (1927); 47, 286 (1928).
2532
HANS L. J. BXCKSTROM AND HAROLD A. BEATTY
Experimental Materials Benzaldehyde. Kahlbaum’s best grade was redistilled in nitrogen and stored in brown, glass-stoppered bottles. Anthracene. This was prepared from pure, synthetic anthraquinone by the excellent method of von Perger,’* thus avoiding the difficult separation of other coal tar hydrocarbons. Fine crystals, snow white, with intense blue fluorescence, melting a t 2 17’(corr.) were obtained. Anthrone. This was obtained by the reduction of anthraquinone according to the directions of K. H. Meger.13 The recrystallized product melted a t 154’ (uncorr.); it contained a small amount of anthraquinone. 0
11
C
C
Anthraquinone. Du Pont Company’s synthetic, crystalline product, “melting point 284.6’,’’ was used without further purification.
0
II
II
0 Anthranol and Anthrahydroquinone.
The preparation of these is described in connection with experiments using them.
OH
OH
(!
c
I
I
H lZ l3
Von Perger: J. prakt. Qhern., ( 2 ) 23, 127 (1881). K. H. Meyer: “Qrganic Syntheses,” 8, 7 (1928).
I
I
OH
INHIBITORY ACTION OF ANTHRACEKE
2533
Anthrahydroquinone dibenzoate. A pure preparation was obtained by the method of K. H. Meyerll; it melted a t 2 9 2 ' (uncorr.). Ozanthrone. An impure specimen was obtained, with considerable difficulty, by a modification of the directions given by K. H. Meyer." 0
1~ C
C
/\
H
OH
Dihydrodianthrone (also known as dianthrone). This was obtained by the oxidation of anthracene according to Barnett and RIatthews.15 The resulting fine white needles were quite free from anthrone and anthraquinone, and had the characteristic indefinite melting point of 2 5 jo-260'. 0
I I>
C
11
0
AZcohoZ. This refers throughout this paper to a good commercial grade of 9j7c ethyl alcohol, not further purified.
Hydrosulfite. This refers to a reducing solution of 50 g. of commercial sodium hydrosulfite and 5 0 g. of sodium hydroxide made up to 500 cc. with water. The solution was allowed to stand in a tightly-stoppered flask until the sediment of impurities had settled; it was then filtered, in an atmosphere of nitrogen, through a Jena glass filter, and kept in the well-stoppered filter flask. The solution suffers a slow auto-decomposition on standing. l4
K. H. Meyer: Ann., 379, 37 (191I ) . Barnett and M a t t h e w : J. Chem. SOC.,123, 387 ( 1 9 2 3 ) .
2.534
HANS L. J . BACKSTROM AND HAROLD A. BEATTY
Apparatus The shaking apparatus used in the photochemical experiments has been described e1sewhere.l It served to keep the aldehyde solutions a t constant temperature and saturated with oxygen, of which the amount and rate of absorption could be measured. The light source was a mercury arc of the Kromayer type, placed immediately in front of the quartz window of the thermostat. As a rule it was used in connection with filters: a nickel oxide glass for 366 mp, and a quartz cell, I cm. thick, filled with O . O O O ~molar potassium chromate solution, for the line a t 313 mp. In experiments on the thermal reaction, the aldehyde was pipetted into constricted test-tubes, which were then evacuated, filled with oxygen, sealed, and wrapped in black paper,* these operations being performed in red light. The tubes were attached to a rotatory shaker, making 180 revolutions per minute, placed in a dark-room, the temperature of which remained fairly constant. Analytical Since preliminary experiments indicated that anthracene was oxidized in the reaction to be studied, and that gravimetric analysis of the reaction products was not feasible, an optical analytical procedure was developed which was capable of determining both anthraquinone and other simple oxidation products of anthracene. This procedure consisted of mixing a sample of the aldehyde to be analyzed with alcohol and hydrosulfite solution and adding water to make a given volume. The proportions of each liquid were such that a clear solution was obtained, except for a precipitate of anthracene which, if present, was filtered out by suction through a Jena glass filter in an atmosphere of nitrogen. Depending on the amounts of oxidation products present, two proportions were used, as follows: Combination S o . I : 40 cc. of alcohol, z cc. of aldehyde, I O O cc. of hydrosulfite (less than 7 days old), plus water to zoo cc.; Combination No. 2 : 2 0 cc. of alcohol, z cc. of aldehyde,Io cc. of water, I S cc. of hydrosulfite (less than 4 days old), plus water to so cc. After mixing, this solution was then subjected to spectrophotometric analysis. Under these conditions, anthraquinone is reduced to anthrahydroquinone of strong red color; anthrone and anthranol appear as the strongly yellow alkali salt of the latter; while dihydrodianthrone is converted to the yellow-orange alkali salt of its tautomer, dianthranol. As these colors indicate, anthrahydroquinone absorbs farther into the visible spectrum than the others, and it may thus be determined separately even in their presence; and a mixture of anthrahydroquinone with only one of the other substances may be analyzed for each component with accuracy by making measurements a t different wave-lengths. However, anthrone and dihydrodianthrone cannot be distinguished from one another with precision, owing to the similarity of their spectra. The spectrophotometer used was of the Xonig-Martens type. The light source was a 150-watt tungsten filament lamp, placed in a fixed position, so *lndly
supplied b y the Eastman Kodak Co.
INHIBITORY ACTION OF ANTHRACENE
2535
that the relative intensities of the two beams could be determined once and for all, and an analysis made by measurements with the absorption cell in one position only. The relative intensities vary slightly with the wave-length. The scale for different settings of the spectrophotometer was calibrated to read in wave-lengths by observation of the strong lines of a mercury and a cadmium arc. 6000
7-j 5000
4000
3000
PO00
7 0 0 0
3
C 460
L
6
FIQ.I Molecular Extinction Coefficients
0
a
Anthrahydroquinone Anthranol
The solution to be analyzed was contained in a plane-parallel glass or quartz cell which was filled almost to the top and was closed by a glass plate sealed on with vaseline to exclude air. The two cells used permitted a choice of four different layer thicknesses, from I O to 26 mm., which combined with the two dilutions employed, allowed measurements to be made under the most favorable conditions of light absorption. As the surfaces of the cells were not optically perfect, a measurement was made in each analysis a t a
2536
HANS L. J. BACKSTROM AND HAROLD A. BEATTY
wave-length in the red (694 mp) where no absorption by the colored products occurred, giving a small correction for reflection. Additional slight corrections were applied for absorption by the solvent, based on a large number of blank determinations carried out in the same manner. By this method, t'he absorption spectra of ant'hrahydroquinone and anthranol were determined, using known quantities of anthraquinone and anthrone. The results are given graphically in Fig. I , where the molecular extinction coefficient K is defined by the formula: log IJI = Kcd, wherein I,, and I are the intensities of incident' and transmitted light', c is the concentration in moles per liter, and d is the layer thickness of the solution in cm. The curve for anthranol has been corrected for the presence of 2.8% of anthraquinone in the anthrone used, as determined from the absorption spectrum at the longer wave-lengths. On the basis of these curves four wave-lengths were selected as being the most suitable for the identification and estimation of anthrahydroquinone and anthranol. The standard values of K at these wave lengths are given in Table I, those for anthrahydroquinone being the mean of a large number of determinations.
TABLE I hlolecular extinction coefficients, II 2.96 0.0179 0.00422 0.0125 ,! 0.00438 O . O I 2 j 2.85 o.o1;8 11 3.18 0.0161 0.00;91 0 . 0 2 5 2 0.00; j 0.01 13 0 . 0 0 1 7 2 4.36 0.0105 >l 0,00283 0.0123 4.35 0.0099 1, 0.00310 0.0139 4.48 0.0098 I, 0.00329 0.0149 4.53 0.0097 7, 0.00293 0.0130 4.44 0.0049 0.00496 0.0023I 0,0260 1 1 . 2 j 0.00380 > 0.00191 0.0181 9.48 o.00400 0.00160 0.0146 9.13 0.00416 ,, 0.0219 10.0s 0.00387 0 , 0 0 2 I8 0 001j; 0.00027 0.0066 2 4 . j 0.00143 ,, 0.00027 0 , 0 0 7 4 27.4 0.00143 >, 0.0003 2 0 . 0 0 8 1 zj.3 0.00141 0.0803
0.0128
0.01;2
Average
0.031 0.035 0.032 0.034 0.032 0,037 0.041
0.038 0.035 0.040 0,033 0.034 0,034 0.035 0.031 0.032
0.033 0.037 0.034 0.036 0.036 0.034 0.03j 0.03j
0.03j 0.039 0.034 0.034 0.033 0.034 0.038 0.034
0.035
I06
5 7 6 4. j 8 180 4 . j 29 58.2 113 8 . 5 2 7 59.2 188 4.j 2 j 5 7 . 5 2oj 4.5 28 73.9 106 8 . 5 26 76.0 I90 4 . 5 2 4 1 1 2 . 4 134 4.5 1 1 112.9 I34 4.5 1 2 IIj.5 9 98 8 . 5 117.8 86 8 . 5 IO 24.8 1 0 7 4.5 2 3 39.2 I11 4.5 2 2 48.6 125 4.5 2 1 49.6 8j 8 . 5 20 j1.6 85 8.; 19 64.4 123 4.j 18 21 . o 82 4 . 5 37 35.0 81 4 . j 36 37.5 83 4.5 35 48.3 68 8.; 34 49.3 59 8.5 33 3 5 . 2 65 4.5 3 35.5 54 8.5 4 36.4 44 8.5 I 2 37.0 59 4 . 5 8.0 34 4.5 16 10.0 2; 8.; 13 11.0 30 8 . j 15
144.5 146.9 148.5 147.8 42.3
89 93 148 161
8.5 8.5 4.j
2558
HAM L. J. B:~CKSTROM AND HAROLD A. BEATTT
reaction between anthracene and oxygen, a t least of any importance, was confirmed by special experiments on solutions in benzene and I ,4-dioxane. A quantitative experiment with a 0.02 molar solution in the latter solvent, using light of wave-length 366 mp, gave a rate of anthraquinone formation of only IO-' mole/min. which, as seen from Table XII, is less than 1.7~ of the corresponding value for a dilute solution in aldehyde. I n both solvents illumination caused dianthracene formation and intense fluorescence. As previously mentioned, this is not the case with solutions in benzaldehyde which shows that the aldehyde molecules have a specific power of deactivating an excited anthracene molecule. It is interesting to note that, in this process, the energy is apparently lost as kinetic energy? if it were taken up by the aldehyde molecule as electronic energy of excitation, for which purpose it is sufficient, it would start a reaction chain and thus indirectly cause the oxidation of an anthracene molecule; but the above results show that this is not the case. The most dilute solution in Table XI11 gave a very low rate of anthraquinone formation as well as an abnormally high B/Q ratio. Both facts indicate that the anthracene concentration was no longer sufficient to break all the chains, some of which were accordingly being broken in other ways, as in the uninhibited reaction. Table XIV shows that the rate of anthraquinone formation in the thermal reaction varies between very wide limits. The same fact had been observed in preliminary experiments, and every attempt was therefore made to keep the experimental conditions as constant as possible. The aldehyde used was all from the same distillation. It was distributed among a number of 50-cc., glass-stoppered, brown-glass bottles that had previously been used a long time for the same purpose. The reaction tubes were unused test-tubes from the same carton. When cleaning them, with hot alkali, hot dilute hydrochloric acid and distilled water, they were all treated uniformly and a t the same time. After being constricted, they were stored away from dust. They were later numbered in the order in which they were used; these numbers are given in the last column of the table, As appears from the table, the rate is invariably higher when the volume shaken was 4.5 cc. than when it was 8.5 cc.; in some cases the ratio between the rates reaches 1.8, which is almost the ratio between the volumes (1.9). This might indicate surface catalysis, but the fact that the rate seems to increase with the time of shaking and the age of the tube points rather to a slow diffusion of a catalyst into the surface layer of the glass and into the solution. The resulting concentration would be inversely proportional to the volume of the aldehyde. As a further illustration we give, in Table XV, the results of some preliminary experiments made a t 25'. I n this case all the runs were started simultaneously, and the test-tubes were not new but had been in use for some time; before being constricted, they were merely cleaned with alcohol and ether. 26
Compare J. Perrin: Compt. rend., 184, 1097 ( 1 9 2 7 ) .
IKHIBITORY ACTION OF ANTHRACENE
2559
TABLE XV Thermal reaction-Average Anthracene concentration, mole/lit. 0.0025 0.0025 0.005
Volume shaken, cc. IO
4.5 4.5 4.5
temperature Time, hours
14 17.5 21
0.01
IO
42.5 44
0.02
IO
i 2
0.01
0.02
4.5
0.04 0.04
IO
0.08
IO
0.023
4.5 4.5
88 I 60
164 309 313
2
so Rate of quinone form. mole/lit./hr.
x
106
44 66 65 36 43 32 49 38 50
43
77
Table XV shows that, although there are considerable variations in the rate of quinone formation, the results are much more satisfactory than in the final experiments. The table represents a 3 2-fold variation in concentration, and the IO-cc. runs, in particular, show very clearly that we are not dealing with a direct oxidation of the anthracene, but with an induced reaction, the rate of which is independent of the concentration. I n the final experiments, this fact is almost completely masked by catalytic influences. Kuhn and M e y e 9 claim that perfectly pure benzaldehyde does not autoxidize in the dark, and that the “thermal” reaction, as ordinarily observed, is caused by the catalytic action of traces of heavy metals. Our results tend to support this view, but they show that the action of the inhibitor does not consist in combining with or in some other way removing the positive catalyst, but in breaking the reaction chains that are started by the catalyst.
Theoretical Part The reaction mechanism. It has been shown in the foregoing that the first step in the induced oxidation of anthracene is probably a reaction with a peroxide of benzaldehyde, giving anthranol and benzoic acid; but it has also been shown that this reaction cannot be ascribed to the stable form of this peroxide, benzoperacid. We are thus forced to the conclusion that in the formation of benzoperacid from benzaldehyde and oxygen, a short-lived, unstable peroxide appears as an intermediate product, and that it is this peroxide which is responsible for the observed reaction. On the other hand, the fact that the formation of benzoperacid is a chain reaction shows that the energy of activation is in some way passed on from the newly-formed peracid molecule t o another molecule of aldehyde. If 28 Kuhn and Meger: Xaturuissenschaften, 16, 1028 (1928);see also Raymond: Compt. rend., 191, 616 (1930).
2560
HANS L. J. BACKSTROM AND HAROLD A. BEATTY
the assumption is made that the the above-mentioned unstable peroxide represents a stage in this process in which the transfer of the activation energy has not yet occurred and that, consequently, this peroxide forms a link in the reaction chain, it is possible to account for the chain-breaking character of the induced oxidation of the anthracene. We thus arrive a t the following reaction scheme : The primary process is the activation of an aldehyde molecule, photocheniically or otherwise, and its subsequent reaction with oxygen :
+02 CsHjCHO ---+ CsH5CHO’ +C s H i c o ~ H ’ . The active peroxide molecule thus formed can do one of two things. ( I ) It can transfer the energy of activation to an aldehyde molecule: CeHbC03H’
+ CsH5CHO = C~HSCOBHiCeHsCHO’.
The probability of this process will be proportional to the concentration of the aldehyde, C1. The deactivated benzoperacid molecule will eventually react with an aldehyde molecule, giving z molecules of benzoic acid. The active peroxide molecule may react with anthracene, giving one (2) molecule of anthranol and one of benzoic acid:
the anthranol thus formed reacts further with oxygen, eventually giving anthraquinone and water. The probability of this process will be proportional to the concentration of the anthracene, CP. This scheme gives directly the ratio between benzoic acid and anthraSince C1 may be requinone formed as - = kzC2+ klC1 = I + 2 k&. Q kzC2 kzC2 garded as a constant, this formula is essentially identical with the one found B k -. experimentally: - = 0 . 9 2 Q C2 The secondary activated aldehyde molecules formed in reaction ( I ) will again react according to the same scheme. This reaction represents the chain mechanism, and its probability, P, determines the efficiency of chain propagation and thus the chain length. Reaction (z), on the other hand, represents the chain-breaking mechanism: it may be assumed that it does not cause further activations. The chain length, L, may be defined as the total number of activated aldehyde molecules formed for every primary activation. It will obviously be given by*’: L = I P P2 , . . which sum is
+
+ + +
These formulae will only apply under conditions where the influence of the anthracene predominates over other possible deactivating processes, Le. a t sufficiently high inhibitor concentrations, Under these conditions, every 27
Compare, Christiansen. Trans. Faraday Sac., 24, 600 (1928).
INHIBITORY ACTION OF AXTHRACENE
2561
primary activation will eventually lead to the formation of an anthraquinone molecule, and the rate of quinone formation therefore measures the rate of the primary reaction. The above reaction scheme does not take into account the formation of the “yellow substance” by a side reaction, nor the fact that part of the anthraquinone apparently is not formed in the way assumed above, but by autoxidation of anthrahydroquinone. In view of this, the agreement with the theoretical formula must be said to be remarkably good, and, indeed, it seems as if neither of these side-reactions had any effect on the B/Q ratio. The experimental results indicate a value of K which is less than unity by an amount which appears to be slightly outside the experimental error. However, this value is the same for the photochemical reaction a t z o and a t 30’ although there is a great difference in the amount of “yellow substance” formed a t these two temperatures; and it is also the same for the thermal reaction a t 29’ in spite of the fact that, as will be shown presently, no anthrahydroquinone is formed in this case. Accordingly it seems probable that the deviation from unity was caused simply by a slight systematic error. As shown in a preceding section, the amount of hydrogen peroxide formed measures the extent of a side-reaction in which anthrahydroquinone appears as an intermediate product. Tables XI1 and XI11 show that, in the photochemical experiments, this side reaction forms a constant fraction of nearly zoYc of the total reaction. Changes in anthracene concentration, temperature, wave-length of the exciting light, ahd effective light intensity, all appear to be without influence on the relative rates of the two reactions. I n the thermal experiments, on the other hand, no hydrogen peroxide was found among the reaction products. It is true that, in these experiments, the time of shaking was so long that any hydrogen peroxide formed would undoubtedly have, in part, disappeared again by reacting with the aldehyde. However, a determination of the rate of this reaction showed that it is by no means great enough to account for the absence of the peroxide, had that been formed in an amount corresponding to the yield in the photochemical experiments. Thus samples of a 0.08 molar anthracene solution containing a small quantity of dissolved 307~ hydrogen peroxide, were shaken in the dark for about 80 hours: the results showed that, in addition to the benzoic acid formed by autoxidation, which could be calculated from the amount of anthraquinone formed, a further quantity had been formed which was equivalent to the hydrogen peroxide which had disappeared. The rate of this reaction amounted to only about
1%
per hour
log - 0.0047 . Then if hydro) (- dt
gen peroxide had been formed in the experiments of Table XIV a t a constant rate, a t least 50% of the amount formed would still have been present a t the time of analysis; but the quantities found by the benzene-water method were never more than 0.01or 0.02 cc. of 0.05iV.thiosulfate per 2 cc. of aldehyde. Water-insoluble peroxide, on the other hand, was always present in measurable quantities.
HANS L. J. BACKSTRGM AND HAROLD A . BEATTY
2562
The reason for this difference between the results of the photochemical and the termal experiments might conceivably be that, in the latter case, the reaction rate was so low that the concentrations of the various intermediate products must have been extremely small; this would prevent the occurrence of such possible side-reactions as: CeHsC03H C14HQOH= C14Hs(OH)2 CBH~COOH.However, the constancy of the figures for hydrogen peroxide in the photochemical experiments can hardly be explained on any such basis, as the concentrations of the intermediate products must have varied considerably from one experiment to another. Thus it appears that we have here a real difference between thermal and photochemical reaction, caused by the difference in the activation process. This might indicate that an excited aldehyde molecule, formed by absorption of a light quantum, can undergo two different transitions, only one of which leads to the formation of an “active” molecule of the kind appearing in the above equations : the relative transition probabilities would have to be independent of the temperature and also of the wave-length of the absorbed light. But if this is true, it should manifest itself in other ways as well and, so far, confirmatory evidence is lacking; and even on these assumptions we have not been able to formulate a plausible reaction mechanism which would combine the formation of 207~ of hydrogen peroxide with a limiting B/Q ratio equal to I , which is the value indicated by the experiments. A reaction such as C6HsCO3H’ Cl4Hl0 = CsHsCHO C14H8(OH)2(by 20% of the total CsH6C08H’)would give a ratio of only 0.8. As regards the formation of the “yellow substance,’’ it may be fitted into the reaction scheme by assuming that sometimes Reaction ( 2 ) does not lead to the formation of anthranol and benzoic acid, but to a condensation product of some kind, and that the latter process is favored by an increase in temperature.
+
+
+
+
General Conclusions As appears from the foregoing discussion, the proposed reaction mechanism is able to account for all the observed facts, except for a relatively unimportant side-reaction which, moreover, does not appear in the thermal reaction. A closer discussion of the consequences of this mechanism therefore seems to be worth while. First of all, it may be concluded from the experimental results that this mechanism is possible only if a very considerable stability may be ascribed to the “active peroxide.” From a comparison of the formula which represents the experiments at zo, viz. B/Q =
+
0.93
+ 0.019 -,C
with the theoretical formula:
kiCi , in which C1, the concentration of the aldehyde, may be k2C2 put = I O moles/lit., i t appears that the ratio between the specific velocities k 20 of the afore-mentioned reactions ( 2 ) and ( I ) is 2 = = 1050. Seglectk, 0.019 ing differences in molecular diameters, we thus arrive a t the conclusion that B/Q =
I
2
~
2563
INHIBITORY ACTION OF ANTHRACENE
an active peroxide molecule suffersut least 1000collisions with aldehyde molecules before giving up its energy. This figure becomes even higher if not every collision between an athracene molecule and an active peroxide molecule is effective in producing reaction, and it probably is not, as there are more effective inhibitors than anthracene. Reaction ( I ) , in which the energy of activation is transferred from the active peroxide to the aldehyde, would accordingly have the character of a chemical reaction rather than an activation by collision. In agreement with this it is found to have a temperature
kz -
20
- - - - 570, which ki 0.035 indicates that reaction ( I ) has an energy of activation which is higher than that of reaction ( 2 ) by 3600 calories. The fact that we have to ascribe such a considerable stability to the active peroxide molecule indicates that there is no difference in kind between the reaction chains in the autoxidation of a pure substance like benzaldehyde and that of a solution, such as sodium sulfite in water. Furthermore, this stability seems to exclude the possibility that the active peroxide might be identified with newly-formed peracid molecules to which the reaction energy is attached in the form of energy of vibration, or with an electronically excited state of benzoperacid; rather it must be regarded as a chemical individual. It should then be noted that, as emphasized by Engler and Weissberg,*Bthe known structure of other addition compounds of the aldehydes indicates that the primary peroxide should not be identical with benzoperacid but should have the following structure, originally suggested by Bach:29 0 coefficient: at 30’
(29’
for the thermal reaction),
/\
CBH6C
0
I\ /
H O
Engler and Weissberg accordingly assume that this represents the primary reaction product from which benzoperacid is only formed secondarily, by molecular rearrangement. This is in very good agreement with our conclusions concerning the “active peroxide,” and it seems extremely probable that we may assign the above structure to this compound. It may thus be said that, up to this point, the proposed mechanism is supported not only by our own experimental results and by the large fund of experimental evidence of a similar nature which has found expression in the Engler-Bach theory of induced oxidations, but also by structural considerations; but the further assumptions which have to be made in order to obtain a chain mechanism must be said to be of a purely hypothetical nature: it must be assumed that every such rearrangement of the primary peroxide
** Engler and Weissberg: “Kritische Studien uber die Vorgange der Autoxidation,” P. 90 (1904). *9 Bach: Monit. scientif., 1897, 479.
2564
HANS L. J. BACKSTROM AND HAROLD A. BEATTY
caujes the activation of another aldehyde molecule and, since the presence of solvent molecules does not influence the process, it must either be assumed that the energy transfer can take place a t a distance, by some kind of resonance effect, or, which seems more probable, that the rearrangement can only occur in a collision with an aldehyde molecule. Under these conditions it may be asked whether the results should not be taken as proof that we are not dealing with energy chains but with material chains of the type of the hydrogen-chlorine combination, where atoms and radicals form the links in the chain. However, aside from the fact that the general character of our results seems incompatible with any such mechanism, benzaldehyde differs from chlorine in that its absorption spectrum, within the spectral region employed in the photochemical experiments, gives no evidence of dissociation.30 I n view of this, and of the large amount of evidence supporting the proposed mechanism, we seem justified in concluding that it is essentially correct. On this mechanism, the inhibitory effect of a substance is determined by its power to react with the primary peroxide; it is not merely a question of the transfer of the energy of activation. This peroxide must be extraordinarily specific in its reactions with oxidizable substances; for instance, the probability of the corresponding reaction with benzaldehyde, to form two molecules of benzoic acid, must be extremely low as shown by the length of the reaction chains-that is, benzaldehyde is a very poor inhibitor for its own autoxidation. The “active peroxide” does not, however, constitute the only vulnerable point in the reaction chain, and one should expect to find a different class of inhibitors breaking the chains by reaction with the active aldehyde molecules. As previously mentioned, the slight inhibitory action of benzoperacid is ascribed to this cause. We have shown above that anthracene is not oxidized by stable benzoperacid. This cannot be expected to hold for all inhibitors as many of these are extremely easily oxidizable substances. For instance, if diphenylamine is added to peroxidized aldehyde a strong brown coloration develops, indicating that a reaction takes place. Consequently we should expect, in the case of many inhibitors and, especially, a t higher temperatures, to find a reaction with benzoperacid going on side by side with the chain-breaking reaction with the primary peroxide. This is presumably the explanation of some results recently reported by Wagner and Brier.S1 They studied the effect of hydroquinone on the rate of oxidation of linseed oil and found that it prolonged the induction period in proportion to the amount added, but that if further quantities of hydro. quinone were added after the end of the induction period, when the normal rate of oxidation had been reached, they were without effect. From this they concluded that “the velocity of the oxidation reaction determines the efso 31
de Hernptinne: J. Phys. Radium, 9, 357 (1928). Wagner and Brier: Ind. Eng. Chern., 23, 40 (1931).
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INHIBITORY ACTION O F ANTHRACENE
fectiveness of hydroquinone as an antioxidant in hydroquinone-linseed oil systems. Apparently hydroquinone is not able to function as an antioxidant after a critical velocity of reaction is exceeded.” We believe the explanation t o be that, when the uninhibited reaction has been going on for a while, the liquid contains large quantities of stable peroxides that immediately oxidize any hydroquinone added at this stage.
Appendix In a paper by Bodenstein,??which appeared after the completion of the above manuscript, an extremely interesting suggestion is made with regard to the mechanism if autoxidation reactions. On the basis of experiments on the oxidation of acetaldehyde in the gas phase he proposes a reaction mechanism, the essential features of which appear from the following equations :
H
H 0
3.
0
CH3CH
\/ 0
I
/
+ CHPC
=
\
H
I1
CH3c-o--
\
H
I11
+ CH~C \
0-OH
I \‘
Thus, like ourselves, he assumes the primary formation of an unstable peroxide of the Bach type, but reaction 3, which correspondes to reaction ( I ) of our reaction scheme, is not interpreted as a rearrangement of this peroxide but as a transfer of two atoms of oxygen: molecule I11 is assumed to be formed from molecule I, and molecule I V from molecule 11. In this way he obtains a chain mechanism without having to assume a transfer of the energy of activation, which was the objectionable feature of our mechanism. On this mechanism it is understandable that, as mentioned above, the “rearrangement” appears to be possible only in a collision with an aldehyde Bodenstein: Z. physik. Chem., B 12, 151 (1931)
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HANS L. J . B A C K S T R ~ MAND HAROLD A . BEATTY
molecule. It would seem that this is the most satisfactory solution which has yet been offered for the problem of the chain mechanism in autoxidation reactions. In one detail, however, we may add to the mechanism of Bodenstein, viz., as regards the mechanism of inhibition. Here Bodenstein, without disK = P K’, (P = cussing the matter further, writes an equation: P’ peroxide, K = inhibitor), thus assuming a transfer of the energy of activation; but our results show that this process, also, is a chemical reaction in which, however, only one atom of oxygen is transferred. This then appears to be the characteristic difference between the chain-breaking process and its opposite, the chain mechanism itself.
+
+
Summary The inhibitory action of anthracene in the autoxidation of benzaldehyde is connected with an induced oxidation of the inhibitor. The primary oxidation product is anthranol, which substance is autoxidizable and reacts with oxygen to give a peroxide which is later slowly decomposed with the formation of the final reaction product, anthraquinone. With increasing anthracene concentration, the relative amounts of oxidation products formed approaches one mole of benzoic acid per mole of anthraquinone. The course of the reaction shows that the induced oxidation of the anthracene is the result of a reaction with a peroxide of benzaldehyde, but it may be shown that this reaction cannot be attributed to the stable form of this peroxide, benzoperacid. We are therefore forced to assume the existence of an unstable primary peroxide, presumably of the structure
O /\ CBHsC
0
I\ /
H O
which must form a link in the reaction chain. The results are best explained on a chain mechanism recently proposed by Bodenstein. In the photochemical reaction, there is a side-reaction, amounting to nearly zocc of the total reaction, in which anthrahydroquinone appears as an intermediate product. This substance reacts with oxygen to form equimolecular quantities of anthraquinone and hydrogen peroxide. The occurrence of this side-reaction has not yet been accounted for. A further side-reaction, common to both thermal and photochemical reaction, leads to the formation of complex, colored products which could not be identified. At 2’ this reaction amounts to only 67, of the anthracene oxidized, but it is favored by an increase in temperature. The absolute rate of the “thermal” reaction appears to be determined entirely by the quantities of positive catalysts accidentally present; but the
INHIBITORY ACTION OF ANTHRACENE
2567
relative rates of autoxidation reaction and induced oxidation depend only on the composition and temperature of the solution. The photochemical reaction is complicated to some extent by a screening effect of the anthracene, due to its strong absorption of ultraviolet light. For the determination of small quantities of anthraquinone and anthranol (anthrone), a spectrophotometric method was developed, based on a determination of the light-extinction curves of alkaline solutions of anthrahydroquinone and anthranol. Methods for the separate determination of benzo peracid, hydrogen peroxide, and water-insoluble peroxides, were also developed. Methods of preparation are given for anthrahydroquinone and anthranol. The solubility of anthraquinone in benzaldehyde has been determined a t oo and 2 5 ' . Princeton, New Jersey.
