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THEDATIVEOR COORDINATION PEROXIDE THEORY OF AUTO-OXIDATION‘ NICHOLAS .4.MILAS Research Laboratory of Organic Chemistry, Massachusetts Institute of Technology, Cambridge, Massachusetts Receioed November 3, 1939
In a recent paper published in This Journal, Stephens (38) has advanced some arguments against the dative peroxide theory of auto-oxidation (27a, 27b) and concluded somewhat dogmatically that the initial stages postulated by this theory have no justification. It is the purpose of this paper to point out that Stephens’ arguments are not entirely justifiable and in several cases are somewhat inconsistent with facts, and to bring forth additional evidence in favor of the dative peroxide theory. When this theory was first enunciated some years ago, the present author made a careful investigation of the chemical behavior of numerous substances which are capable of auto-oxidation. This study revealed the fact that all substances capable of combining with molecular oxygen are known to be unsaturated towards reagents other than molecular oxygen. The apparent exception of saturated hydrocarbons is easily accounted for by the assumption originally made by Mardles (23) and independently by the present author (reference 27b, p. 347) that combination of molecular oxygen with saturated hydrocarbons occurs either through free radicals (highly unsaturated towards other reagents) or through the exposed electrons resulting from the activation of hydrocarbon molecules prior to their decomposition. Hydrocarbons, especially of the hexaphenylethane type, are known to split very readily into free radicals which add molecular oxygen with considerable avidity. The presence of unsaturation in auto-oxidant molecules is confined entirely to certain atomic nuclei in each molecule known to possess loosely bound molecular valence electrons which can be easily donated to other nuclei by sharing. That these electrons are loosely bound in certain nuclei of polyatomic molecules has been recently amply demonstrated by Mulliken (29). This author gives the electronic structures of NH, and HzO as ls2 2s2 2p ( ~ ) ~ 2 p (and u ) ~ls22.9 2 p ~ ~ 2 p 6 ~ 2respectively, pc~, and states that “the two 2p (u)electrons in NHs avoid the region of the H nuclei, and as a 1 Contribution No. 94 from the Research Laboratory of Organic Chemistry, hlassachusetts Institute of Technology. 411
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result are relatively loosely bound and reduce the energy of formation of the molecule.” “Direct evidence of these electrons in NH3 is given by the low ionization potential of this molecule,which is 11.1volts, Many features of the chemical behavior of NH3, e.g., the formation of NH4+, H,N.BC13, Cu++(NH3)4,etc., are reasonably interpreted as conditioned by the stabilization of the two loosely bound 2p(u) electrons of NH3 under the influence of an additional nucleus.” The ionization of a 2p(a) electron of NH3 would take, according to Mulliken, 15 volts, while that of a 2s electron, 25 volts. Although NH3 reacts with oxygen only at relatively high temperatures, several of its derivatives are known to auto-oxidize quite readily. Sodamide, for example, adds molecular oxygen to form sodium amidoperoxide, NaH2N.02 (35). Several aryl amines, hydrazones, etc., can be isolated only in the form of their salts in which the 2p(a) electrons are stabilized (reference 27a, p. 1211). Furthermore, the instability of free radicals, like the methyl radical, is, according to Mulliken (29), due to the presence in the radical of a 2p(a) electron. I t has already been suggested elsewhere (reference 27a, p. 1211) that the tendency to auto-oxidize of the hydrides and aryl derivatives of the elements of the fifth group of the periodic table increases as the “effective nuclear charge” of the central elements increases. If the auto-oxidation of these substances proceeds through a preliminary addition of molecular oxygen to the unshared pair of electrons present in each molecule, then there ought to exist a relationship between the tendency of these substances to auto-oxidize and their ionization potentials. That such a relationship actually exists not only with the hydrides of the elements of the fifth group, but also with those of the sixth and seventh groups, has been recently pointed out by the author (28). I t has been shown that as we pass down in each group from the lightest to the heaviest element, there is a regular decrease in the ionization potentials of the hydrides, which means an increase of looseness of the molecular valence electrons. With an increase of looseness of these electrons, there ought to be a corresponding increase in the tendency to auto-oxidize. This is made obvious when we compare the first ionization potentials of the alkali and alkaline earth metals with their tendency to oxidize. I t has been found that with a decrease in the ionization potentials there is a corresponding increase in the tendency t o auto-oxidize. These considerations should not be confined to the inorganic substances, since the principles of oxidation and reduction apply equally well to organic substances, When molecular oxygen combines with an organic substance, the latter contributes two electrons to it, and the ease with which these electrons are contributed usually determines the reducing power or the oxidizability of the substance. Furthermore, the ease with which molecu-
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lar oxygen takes up these electrons to form the peroxide group may also depend upon its electron affinity, or the energy required to remove two electrons from the peroxide ion, 0 2 - - . Unfortunately there are no data available at present which will enable us to estimate this electron affinity. Noyes and Beckman (30) have calculated the electron affinity of atomic oxygen for the first electron and found it to be 4.6 volts. Whether that of the molecular oxygen for two electrons is of the same order of magnitude is not known. As an oxygen molecule approaches an easily oxidizable substance, it is easily conceivable that it may attract electrons from the latter with a force equal to its electron affinity, or, in other words, it may reduce the ionization potential of the auto-oxidizable substance. Some evidence in favor of the electronic structure of the reacting oxygen molecule originally proposed by the author, is to be found in the structure of hydrogen peroxide recently proposed by Maas and his students (5,8,21). From a study of the dielectric constant, molecular refractive power, and the parachor of hydrogen peroxide, these investigators proposed the unsymmetrical structure,
“>.
- H
-?.
0
or
.. .. *. ..
H:O:O:
H
which also agrees with the ease with which hydrogen peroxide loses an oxygen atom to form water. This structure is probably in equilibrium with the symmetrical structure. Since hydrogen peroxide is known to form by the addition of hydrogen to molecular oxygen, its unsymmetrical structure provides evidence for our electronic structure of the reacting oxygen molecule. However, as has been emphasized elsewhere by the author (reference 27b, p. 300), the “odd” electronic structure proposed by Lewis should not be excluded in auto-oxidation reactions. Although some of the molecular valence electrons, present in many of the auto-oxidant molecules, are apparently paired, it is not certain that this condition holds when these molecules are activated or excited. Moreover, several of the auto-oxidants, such as nitric oxide, the alkali metals, ferrous and chromous compounds, and free organic radicals, in general, are definitely known to possess unpaired valence electrons through which they interact with other molecules. Auto-oxidation reactions are strongly exothermic and in numerous cases are chemiluminescent. Backstrom has clearly demonstrated that chemiluminescent reactions reveal the presence in the reaction system of molecules of exactly the same kind as are produced by the absorption of light of the same or higher frequency as the emitted light. Furthermore, the energy transferred during auto-oxidation reactions to the reactant molecules, presumably by collisions of the second kind, excites these molecules
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into those quantum states which are responsible for their ultra-violet spectra. That this energy is of chemical origin can be easily shown by the fact that the heat of reaction is usually greater than the energy of excitation which corresponds to the wave length of the emitted light. In the case of benzaldehyde, for example, Backstrom has shown that the region of strong light absorption is approximately 3800 A.U., corresponding to an energy of excitation of 75 kg-cal. The heat of auto-oxidation of this substance is 140 kg-cal., which is more than sufficient to cause the excitation of benzaldehyde molecules. Another notable example is the thermal decomposition of ozone, which shows a heat of reaction of 69 kgcal. whereas the longest wave length in its absorption spectrum corresponds to an energy of excitation of only 47 kg-cal. If the energy of excitation in auto-oxidation reactions were entirely of vibrational origin, as Stephens seems to assert, one would expect infra-red light to have an accelerating effect in these reactions. As far as the present author is aware, there is no evidence that this is the case. On the contrary, the decomposition of hydrogen peroxide which follows a chain mechanism is totally unaffected by the absorbed infra-red light, while it is greatly accelerated by ultra-violet light (18). Moreover, there exists a strong parallelism between the photochemical and the thermal auto-oxidations. Both have a chain mechanism; both are susceptible to the influence of negative and positive catalysts. Another parallelism closely related to this one is that between the effects of peroxides and ultra-violet light on certain chemical reactions. Both are known to accelerate auto-oxidation reactions, to induce polymerization of unsat urated substances and to affect the photographic plate. Furthermore, the recent observations of Kharasch and his students (16, 17) and of Smith (36) seem to indicate clearly that in the presence of oxygen or peroxides the olefinic bond is polarizedin such a way as to influence the direction of the addition of hydrogen bromide to it. Other additional reactions to the double bond cannot be easily understood unless we assume with Lowry (22) and with Carothers (7) that the reacting molecules are polar or electrovalent, in which two of the electrons have been displaced towards one or the other carbon atom. In fact, it is because of this displacement of the electrons that we are justified in speaking of these substances as unsaturated molecules (20). The fact that the electric moment of ethylene and of other unsaturated substances is found to be zero does not disprove the assumption that the reacting unsaturated molecules are polar. This has been clearly stated recently by Smyth (37). Finally, the foregoing considerations lead to the unescapable conclusion that the mechanism of the formation of excited molecules is the same in the dark or in photochemical auto-oxidation reactions, and that these ex-
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cited molecules are produced by the simultaneous supply of vibrational and electronic energy. AUTO-OXIDATION OF HYDROCARBOSS
The fact that saturated hydrocarbons auto-oxidize is not at variance with the general theory proposed by the present author, as Stephens seems to insist, since the theory assumes the addition of molecular oxygen to molecular valence electrons which are exposed either as the result of excitation of the hydrocarbon molecule or of its dissociation to free radicals. Whether the final peroxide is an alkyl hydroperoxide or a dialkyl peroxide depends entirely upon the hydrocarbon in question and the relative strengths of the bonds concerned. This point has been more fully discussed by the author elsewhere. With regard to the auto-oxidation of unsaturated hydrocarbons, there is a displacement of electrons prior to the addition of oxygen, a condition which is fulfilled only by the theory of the present author. The effective strength of the bond broken in the auto-oxidation of an ethylenic substance may be taken as the difference of the strengths of the double and the single bonds. Taking the strength of the C-C bond as 80 kg-cal. (33), and that of the C=C bond as 123 kg-cal. (39), this difference comes out as 43 kg-cal. which is not very far off from the experimentally obtained activation energy for the oxidation of ethylene, namely, 35 to 43 kg-cal. (40). This indicates that in the activated ethylene molecule the double bond is probably broken with subsequent displacement of the electrons to form the polar double bond discussed in the previous section. AUTO-OXIDATIOS O F ALDEHYDES,
ETC.
Pure aldehydes are not known to dehydrogenate easily, as Stephens claims; in fact, we know of no case in which aldehydes dehydrogenate easily, say, to form diketones even in the presence of dehydrogenating catalysts. Furthermore, Eyring (10) has shown recently that in the case of acetaldehyde, for example, the C-H bond is stronger than the C-C bond. Such arguments, therefore, as those proposed by Stephens are entirely irrelevant t o the auto-oxidation of aldehydes in which the seat of unsaturation lies in the carbonyl group. Both Wolf (41) and Smyth (37) have concluded that the electric moment of aldehydes is inherent in the carbonyl group with the oxygen being the negative end of the dipole. Hydrogen peroxide and ozone are stronger oxidizing agents than molecular oxygen, yet in neither case does the hydrogen of the alpha carbon atom of aldehydes oxidize easily to a hydroxyl group. In the case of hydrogen peroxide stable a-hydroxy peroxides and hydroperoxides are formed (34). Ozone, which is known to add on to unsaturated groups very
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much in the same manner as molecular oxygen, forms alkyledine peroxides with aldehydes (13, 12, 11). These peroxides either polymerize or, when warmed up, go over to their isomeric acids.
H
.. + O8
R:C:O:
--f
H .. ,. .. R:C:O:O:
,. ..
H or
O
I
+ O2
RC(i
\I
0
In the case of the auto-oxidation of aldehydes the very recent work of Bodenstein (4),of Backstrom and Beatty (2) and of Kohler and Nygaard (19) confirmed the conclusion of earlier investigators that an active moloxide is first formed which subsequently rearranges to the peracid. Backstrom and Beatty state definitely that the direct formation of a peracid or an activated form of it in one step fails to account for their experimental results; accordingly they were forced to assume the existence of an intermediate moloxide originally proposed by Bach. These and various other considerations confirm our earlier view in regard to the mechanism of the auto-oxidation of these subs$ances. A similar view has been adopted in the case of the addition of oxygen to the ketones and thioketones. That the double bond in thioketones and thioaldehydes is virtually broken under ordinary conditions is shown by the fact that the aliphatic thioketones and thioaldehydes are known to exist only in their dimeric or trimeric modifications (3). This conclusion is strengthened by the high electric moment recently reported for some of these substances by Hunter and Partington (14). Contrary to the assumption of Stephens, no alkyledine peroxides are known to form in the case of the auto-oxidation of thioketones. With regard to the mechanism of the auto-oxidation of ethers, there exists ample chemical evidence on oxonium compounds (31, 32, 24) which supports the correctness of our original conclusion that an unstable oxonium or dative peroxide is initially formed. Finally, the auto-oxidation of triethylphosphine and related substances proceeds, as has already been shown elsewhere, through the initial formation of dative or triethylphosphonium peroxide. Stephens disagreed with this view and proposed an initial peroxide in which the phosphorus atom still retains its original valence of three. This is contrary to our knowledge of the chemical behavior of this class of compounds towards oxidizing agents. The main stable product of this reaction is triethylphosphonium oxide, while secondary products formed by the interaction of this oxide with the peroxide have also been isolated (9, 15, 25). If the peroxide proposed by Stephens were the initial product, one would expect it to hydrolyze in presence of water to yield ethyl hydroperoxide. Actually, no such product is known to form. That the oxide is the main final product is clearly
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shown by the auto-oxidation of triethoxyphosphine, (CnH60)3P,to form first the peroxide and then triethoxyphosphonium oxide, (C2Hs0)3P0(42). Another example of this type is the auto-oxidation of phosphorus trichloride, which yields phosphorus oxycloride (26). Trimethylamine and other alkylarsines are also auto-oxidized to yield through the intermediate peroxides alkylarsonium oxides (6). SUMMARY
1. Additional evidence has been given in favor of the dative peroxide theory of auto-oxidation originally proposed by the author. 2. It has been shown that the recent criticism of this theory is entirely unjustifiable. REFERENCES (1) BACKSTROM: Medd. Vetenskapsakad. Nobelinst. 6; No. 16 (1927). A N D BEATTY: J. Phys. Chem. 36, 2530 (1931). (2) BACKSTROM (3) BAUMANN AND FROMM: Ber. 24, 1419 (1891). (4) BODENSTEIN: Z. physik. Chem. 12B, 155 (1931). (5) BUTTLER AND MAAS:J. Am. Chem. SOC.62, 2184 (1930). (6) CAHOURS: Ann. 122, 205 (1862). (7) CAROTHERS: J. Am. Chem. SOC.46, 2226 (1924). (8) CUTHBERTSON AND MAAS:J. Am. Chem. SOC.62, 489 (1930). (9) ENQLER AND WEISSBERG:Kritische Studien iiber die Vorgangen der Autoxydation, p. 64. Braunschweig (1904). (IO) EYRINQ:Z. physik. Chem. 7B, 244 (1930). (11) See also FONROBERT: Das Ozon, p. 159. Verlag von Ferdinand Enke, Stuttgart (1916). (12) HARRIESAND KOCTSCHAN: Ann. 374, 321 (1910). Ann. 343, 352 (1905). (13) HARRIESAND LANQHELD: (14) HUNTERAND PARTINGTON: J. Chem. SOC.1933, 87. (15) JORISSEN: Z. physik. Chem. 22, 38 (1897). (16) KHARASCH AND MAYO:J. Am. Chem. SOC.66, 2468 (1933). (17) KHARASCH, MCNAB,AND MAYO:J. Am. Chem. SOC.66, 2521, 2531 (1933). (18) KISTIAKOWSKY: Photochemical Processes, p. 173. The Chemical Catalog Co., New York (1928). (19) KOHLER AND NYGAARD: J. Am. Chem. SOC.66, 310 (1933). (20) LEWIS:Chem. Physics 1, 17 (1933). (21) LINTONAND MAAS:Can. J. Research 7, 81 (1932). (22) LOWRY:J. Chem. SOC.123, 822 (1923). Nature 128, 116 (1931). (23) MARDLES: Z. physik. Chem. 19B, 164 (1932). (24) See also MEERAND POLANYI. (25) For the auto-oxidation of other phosphines, see MICHAELIS AND LINK: Ann. 207, 210. (26) MELLOR:A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. VIII, p. 1002. Longmans, Green and Co., London (1928). (27) (a) MILAS:J. Phys. Chem. 33, 1204 (1929). (b) MILAS:Chem. Rev. 10, 295 (1932). (28) MILAS:J. Am. Chem. SOC.66, 486 (1934). (29) MULLIKEN:Phys. Rev. 40, 55 (1932).
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(30) NOYESAND BECBMAN: Chem. Rev. 6 , 89 (1928). (31) For the most recent work in this field see PFEIFFERAND HAACK: Ann. 460,156 (1928), and reference 32. See also reference 24. (32) PPEIFFERAND OCHIAI:J. prakt. Chem. 136, 126 (1933). (33) RICE,JOHKSTON, AXD EVERING: J. Am. Chem. Sac. 64, 3529 (1932). (34) RIECHE : Alkylperoxide und Ozonide. Theodor Steinkopff, Dresden and Leipzig (1931). (35) SCHRADER: Z. anorg. allgem. Chem. 108, 44 (1919). (36) SMITH:Nature 132, 447 (1933). (37) SMYTH: Dielectric Constant and Molecular Structure, p. 137. A. C. S. Monograph No. 55. The Chemical Catalog Go., New York (1931). (38) STEPHENS: J. Phys. Chem. 37, 209 (1933). (39) TAYLOR: Treatise on Physical Chemistry, p. 328. D. Van Nostrand Co., New York (1931). (40) THOMPSON AND HINSHELWOOD: Proc. Roy. Soc. London 126A, 277 (1929). (41) WOLF: 2. physik. Chem. 3B, 128 (1929). (42) ZIMMERMANN:Ber. 7, 289 (1874).
