April 20, 1961
MECHANISM OF TERTIARY AMINEOXIDE REARRANGEMENTS
tion 6 implies that hydrogens are transferred within the molecule. Using steady state assumptions regarding the transient species Rf* and Rf', we obtain for the quantum yield.
where k6(Q) refers to the contributions by the various quenchers (Rf, DfHzand K I ) and where the constant k , is included to account for the acid-base catalytic action of water alone (it?., finite rate for B = 0). From the variation of a with (23) we calculate that ks = 210 k , liter2 mole-2. Using the rate a t zero buffer concentration we calculate that k , = 133 sec.-l. If Q is taken t o be the photochemical product then the concentration of Q is proportional to Eo - E since the dye obeys Beer's law. If the quenching term is dominant in the denominator of the above theoretical equation, then @ = a//3 - E in conformity with the observed result (Fig. 2 ) . From the experimental data i t is found that the quenching for the photoproducts k g = 1.06 X 10lomole-' sec.-l. When K I is added to the system there is competition between iodide ion and photoproduct as quenchers. I t f o l l o ~ sfrom the theoretical expression for @ that for a given degree of retardation the KI concentration must be higher for higher product concentration, thzt is, for greater extent of reaction. From our retardation studies with K I we calculate a quenching constant of 1.5 X lo6 k, liter mole-1 scc.-l when extrapolated to zero time (i.c.,no product present). If every encounter between Rf' and iodide ion led to quenching, then for K I kg = 6.6 X l o o liter mole-' sec.-l as shown froin diffusional argumentsz3 so from this argument k , = 4.4 X lo3set.-'. If i t is assumed that all singly excited riboflavin molecules which do not fluoresce are converted to long-lived species, then k3/(kZ ka) is unity minus the fluorescence yield, namely, 0.26. In pure water CB = 0.006 so the
+
(23) B. Sveshnikov, A c t a physicochim. U.S.S.R., 4, 453 (1936); 755 (1937).
7,
1871
lifetime of the long-lived species k d - l = 1.0 millisec. or about one hundred thousand times longer than the first excited singlet state.24 The fact that the photoproduct is a more efficient quencher than K I suggests that the former quenches statically possibly as a semiquinone dimer. Deviations from the above scheme appear after the reaction has proceeded for some time (points not lying on the straight lines of Fig. 2 ) . This may arise from a dark reaction of leuco deuteroflavin with riboflavin: (7) DfHz Rf Df 4- RfHz followed by the photolysis of deuteroflavin leading, we assume, to lumichrome Lc plus fission products from the side chain :
+
hv
+
(8) Di +Lc fragments. When oxygen is admitted,' leuco deuteroflavin is Oz -+ Df 4- H202. I n a oxidized: (9) DfH2 02 -c similar way RfHz is oxidized: (10) RfHz Rf H2O2. These reactions proceed via a free radical since the polymerization of vinyl monomers are initiated under these conditions.l5 Extraction of the faded and then oxidized solution yields lumichrome formed in step 8. Treatment with alkali yields a second chloroformOHsoluble compound, lumiflavin Lf: (11) Df * Lf fragments. The contribution of this reaction is unimportant unless the pH is high (about 9) and/or the intensity of light is low. The rapid anaerobic fading of the reoxidized solution must be ascribed to a new compound Df since the wave length dependence is different from that of the fading of riboflavin. The leucoflavin can reduce the substrate S (DPIP): (12) DfHz S -+ Df SHz. Unlike true photosensitizers, however, the sensitizer riboflavin is destroyed in the over-all reaction. The fact that HzOz has been detectedlO in the faded solution before oxygen is admitted has been takenz6to be the proof for the water-splitting hypothesis. On the other hand, H202 could be produced in the reactions leading to lumichrome (step 8) or lumiflavin (step 11) without the envolvement of water.
+
+
+
+
+
+
+
(24) G. Weber, BioGhrm. J., 47, 114 (1950). (25) I. Fridovich and P. Handler, J . B i d . Chem., 236, 1835 (1960).
[CONTRIBUTION FROM THE ORGANIC CHEMISTRY DEPARTMENT, UNIVERSITYOF SYDNEY, SYDNEY, N.S.W., AUSTRALIA,THE CURTIKSCHOOL OF MEDICAL RESEARCH, AUSTRALIANNATIONAL UNIVERSITY,CANBERRA, AUSTRALIA,AND THE NATIONAL HEART INSTITUTE, NATIONAL INSTITUTES O F HEALTH, BETHESDA, MARYLAND]
JOHN
Tertiary Amine Oxide Rearrangements. I. Mechanism BY J. CYMERMAN CRAIG,'F. P. DWYER, A. N. GLAZER AND E. C. HORNING RECEIVED AUGUST5, 1960 It is proposed that t h e metal-complex catalyzed rearrangement of trimethylamine N-oxide occurs in two 1-electron steps involving the formation of an intermediate possessing an unpaired electron on a methylene carbon atom which is then a t tacked by an OH radical to yield the methylolamine known t o give dimethylamine and formaldehyde. The N-oxide is bound to the metal through its oxygen, and a hydroxo or aquo group occupies t h e adjacent (cis) coordination position. The necessary ability of the metal complex to exist in a higher oxidation state and to provide the requisite binding sites is satisfied by a number of iron(III), ruthenium(II1) and (IV), osmium(II1) and (IV) and vanadiurn(1V) compounds.
Introduction
(1) Visiting Scientist, National Institutes of Health, 1969.
amounts of these compounds present in biological
5892 (1955).
1872
J. C. CRAIG,F. P. DWYER,A. N. GLAZER AND E. C. HORXIKG
minal products of amine oxidation but instead have a function as intermediates in some biochemical reaction sequence. The Polonovski reaction, whereby tert-amine oxides containing a t least one N-methyl group are converted by acetic anhydride into sec-amines and formaldehyde, suggested that the biological step might be one of oxidative demethylation mediated in the cellular environment by a catalyst capable of lowering the activation energy of the process to a level commensurate with that existing in natural systems. 4 further instance of such a reaction has been reported by Tsuyuki and co-workers5 in the transformation of octamethylpyrophosphoramide ("Schradan") to an N-oxide readily hydrolyzed to the heptamethyl compound and formaldehyde, and it was pointed out2b that the intermediate in this reaction was probably the N-methylolamine or the equivalent anhydronium base.6
at 10" into dimethylamine, formaldehyde and sulfur dioxide. Both the reaction using acetic anhydride and that employing sulfur dioxide may be formulated as closely-related plausible ionic mechanisms X and B However, recently it has been shown3b that the rearrangement of 2-methylpyridine N-oxide with acetic anhydride (which proceeds with a strongly exothermic phase) involves free radical internnediates. (See ref. 3b for this mechanism which we will designate as (C).) Moreover, the rearrangement of triethylamine N-oxide has been found3bto give acetaldehyde and N,N-diethylacetamide, and that of diethylaniline N-oxide in boiling benzene vias effective in causing the polymerization of styrene. 311 The Polonovski reaction, also characterized3"by its exothermic nature, must therefore proceed hy nti analogous free-radical sequence, as shown in (D)
+ -
+
CH,
AcO-
CHz-0 AcOH
1
>NH
I'
+> N+
(1)
I
>N-0
+
-
CH3
(3)
I
>K
CI12. OAC
---+
I
>S
+ HOAc
CHZ=O
'
-H+
-
(B)
+ AcOH
COCHB ----f>NCOCHz -
0
CHI
t
+
-
-+
\-01. s3
(2112-H.
I I + ++ s-0 - + >N-O-$-O
-
.
. .0 I
++
It
>N
-
H?
CHz
+
0-s-0 - ++ -
CH2-0 I I CH>=O >NH 0-S-0 +>NH SOz
+
-++-
+
A similar reaction is to be found in the ready formation7 from trimethylamine oxide and sulfur dioxide of a compound Me3N0.S02 shown to undergo facile decomposition in aqueous solution (3) (a) M. Polonovski and &I. Polonovski, Bull. BOG. chinz., 41, 1190 (1927); (b) V. Boekelheide and D. L. Harrington, Chem. arzd I n d . ( L o n d o n ) , 1423 (1955). (4) (a) J. Axelrod and J . Cochin, J . Pharrn. and E x p . T h e r a p . , 121, 107 (1957); (b) J. Axelrod, ibid., 117, 322 (1956); ( c ) J , R. Gillette, J. V . Dingell and B. B. Brodie, Nature, 181, 898 (1958); (d) B. N. L a D u , L. Gaudette, N.Trousof and B. B. Brodie, J . Biol. Chem., 214, 741 (1955). ( 6 ) H . Tsuyuki, &I. A. Stahmann and J . E. Casida, Biochem. J . , 59, iv (1955). (6) E. W'enkert, E x p e r i e n l i a , 10, 3.36 (1954). (7) (a) M. Polonovski and R I . Polonovski, Bull. S O C . chiin., 39, 1147 (1926); (b) A. B. Burg, THISJOURNAL, 66, 1629 (1943); ( c ) H . 2. Lecher and W. B . Hardy, ibid., 7 0 , 3789 (1948).
It is evident that 2-methylpyridine N-oxide is the vinylog of the aliphatic system in (D) and closely similar behavior is to be expected. In the case of the sulfur dioxide r e a ~ t i o nthe , ~ available evidence does not permit a decision between an ionic and a free radical mechanism. The stability of the intermediate R3N0.S02to solvents other than water suggests that an ionic mechanism may operate, however. Horning and his co-workers foundbathat the rearrangement of N,N-dimethyltryptaniine oxide, a compound isolated from natural sources, took place in aqueous solution a t @H 2--7 in the presence of ferric ion under relatively mild conditions to yield AT-methyltryptamineand formaldehyde and established that similar rearrangements occurrcd with bufotenine oxide, hordenine oxide, X,K-dimethyltyrosine and N,X-dime thylglycine oxide.RC These authors showed the reaction to require Fe3+, no rearrangement taking place with other metal ions. In the case of dimethylglycine oxidesp the reaction was found to proceed in aqueous solution using Fe3+ and either oxalic acid (PH 1 to 5) or tartaric acid (PH 2 to 9) as complexing agent. (8) (a) M. S.Fish, N. AI. Johnson and E. C. Horniug, ibid.,76, 3668 (1958); (b) &I. S. IWn, N . M . Johnson, E . .'1 1,axrence and E. C . Horning, Biorhem. Biophys. A d a , 18,564 (1955); ( c ) M. S . Pish, C .C . Sweeley, K. M . Johnson, E . P . Lawrence and E . C . Hcirning, ibicl , 21, 196 (1958); (d) M. S. Fish, C . C . Sm-ecley and E . C. Horning, C h c m and I n d . ( L o n d o n ) ,R 34 (1956); !e) C. C. Sivecley a n d F;. C. Hiirniiw THISJ O U R N A L , 76, 2620 (1957).
1873
MECHANISM OF TERTIARY AMINEOXIDEREARRANGEMENTS
April 20, 1961
TABLE I EFFECTOF OXALICACIDAND Fe3+ ON
THE
DEMETHYLATION OF TRIMETHYLAMIXE N-OXIDE' After . .....
Ratio oxalic acid :Fe
1.2531
Temp.,
P 11
Chelating agent
Metal
0.8
Oxalic acid, 0.00375 mole
Fe( N03)3.9Hz0,0.003 mole
OC.
M g
+ HCI
100 90 80 75 25
33 39
19,17
0 0
Oxalic acid, 0,030 mole
0.4
100
3: 1
KI[Fe(C20&13-, 0.003 mole
3.6
90 80 70
16 0 0
6: 1
K3[Fe(CZO4)3 13-, 0.003 mole ( COONH4)2~Hz0, 0.01 mole
3.8 1.6
90 80 100
6.0b
90 80 100
0 0 37 8 0 0 0
10: 1
Fe(i';03)~.9HzO, 0,003 mole
Fe( X'0~)3.9HzO,0,003 mole
+
Yield, %
reduction with
Before
~~~~~~
reduction with M g + HCI
38,43 12 0 0 16 0 0 0 (160 min.)
(75 min.) 0
a All solutions contained 0.001 mole of amine oxide and were kept for 40 min. a t the temperature specified, unless otherwise mentioned. * Amine oxide alone.
group could then be effected either without direct binding of the organic moiety or (more probably) COOHCHzNMe20 + [COOHCHZNM~(CHZOH)] with a t least a single-point attachment of the N--+ COOHCH&HMe CHzO oxide to the metal, necessitating the existence of a t least one vacant or labile coordination site in the COOCH2$MeZ(OH) --+ metal chelate. These various possibilities have now been examined using a variety of metal comCOz [CH2=6Mez] plexes and complexing agents.
The successive steps E, F and G occurred
+
+
[bHz-NMez]
+ OH --+ [OHCHpNMe2]--+ CHzO
COOH.CHzNMei0
+ CH2O --+ COOH.CH2NMeZ
+ "Mez
+ H,COOH
The mechanism of the ferric ion-induced demethylation which occurs in aqueous solution is unknown; in order to study the function of the metal ion and of the chelating agent in the demethylation reaction, without the complication of simultaneous decarboxylation, trimethylamine oxide was selected as the test substance. The expected reactions were H and I (H) Me3SO --+ Me2NH (I) Me3N0 CHzO +Me&
+
+ CH20 + HCOOH
Reaction I results from the reduction of some Noxide by formaldehyde to give the tert-amine and formic acid. The reaction may be regarded essentially as an internal 2-electron oxidation of the methyl group, a t the expense of the nitrogen, and the role of the metal ion catalyst may be in mediating the transfer of electrons to the N - 0 portion of the molecule, f
-
possibly by a transitory oxidation of the metal chelate to a suitable higher oxidation state. The complexing agents employed8 indicated the formation of a metal chelate; the oxidation of the methyl
Experimental All reagents were of analytical reagent grade. Trimethylamine oxide was used as the colorless dihydrate of m.p. 96". Reaction conditions similar to those of Sweeley and Homingaewere employed. Determination of Formaldehyde.-Formaldehyde was determined by direct isolation from the reaction mixture as the 2,4-dinitrophenylhydrazone; the reagent was prepared by dissolving 2 g. of 2,4-dinitrophenylhydrazine in 500 ml. of 2 N HCI. The estimation procedure for total formaldehyde and formic acid was: following the heating period, the reaction mixture was placed in an ice-bath and chilled to 0 to 5 " . I t was then acidified by the addition of 5 ml. of 5 N HCI, and then an excess of magnesium shavings. After a reduction period of 25 min. a t 0 to 5 O , the solution was decanted, or if necessary filtered, from any excess magnesium and 50 ml. of the 2,4-dinitrophenylhydrazine reagent was added. The solution now was shaken vigorously for 2 min. and the 2,4-dinitrophenylhydrazone was filtered, washed with cold water and dried overnight at 70 to 80 '. Duplicate determinations gave results reproducible to within 1%. All rn.p.'s were checked and mixed m.p.'s taken wherever possible. In blank experiments i t was found that this procedure gave a 97 to 99% recovery of formaldehyde and a 957G recovery of formic acid. In the presence of vanadium salts and with ruthenium and osmium complexes, formic acid was not reduced to formaldehyde under these conditions. A large amount of formaldehyde present was also lost, since in these cases the yield of the dinitrophenylhydrazone was very much decreased after the reduction. I t seems probable that the reduction can in some circumstances proceed to methanol. For this reason the efficiency of these catalytic agents is likely to be much higher than the yields of the derivative would indicate.
J. C. CRAIG,F. P. DWYBR, A. N. GLAZER AND E. C. HORNING
1874
Catalysis by Iron(II1) Chelates. A. Oxalate System.The reaction mixture containing 0.001 mole of trimethylamine oxide, 0.003 mole of iron(1II) nitrate 9-hydrate and 0.00375 mole of oxalic acid in a total volume of 15 ml. ( p H 0.8) was kept at the desired teriipcrature for 40 n h . A fe\v experiuients involving longer tiines indicated that no further increase in fornialdehyde formatioil took place. Other systems examined under the same coii.ditions contained 0.001 mole N-oxide, In) alone, ( b ) with (1.003 mole Fe(II1) nitrate %hydrate and v:;rtious ratios of osa!ic acid (nil to 0.03 mole) or (c) with 0.003 mole of potassium trisoxalatoferrateiI1Ii. with arid without the nreseiice of 0.01 mole of added oxalate ion. T h e resuits a& presented in Ta!ilc I .
t 60
1 I
Vol. 83
a t pH 6.3, an 18y0yield wa.c obtained against a 48yo yield after reduction, indicating that reactioti I took place t o an appreciable extent. KO rearrangement occurred a t 25'. On addition of the trimethylamine oxide to the iroti(Il1) tartrate syst.ern, color changes dependent on pI1 occurred a t rooni temperat,ure, as: .bH 0.5 t o 2.5: no change observed. pI1 3.i) t o 5.5: dark brown color after 1 t o 2 ruin., faded slightly on standing. pH 5 . 7 t o 6.2: very intense immediate dark brown color, fading slowly to that described for the 3.0 to 5.5 range. pH 6.6 t o 9.8: no color on mixing; a less intense dark brown color developed on standicg. I). Succinic Acid.--.At pH 1.2, 0.03 niole of succinic acid in a reaction mixture contairiing 0.001 mole triniethylamine oxide and 0.003 mole of Fe3+ gave only a trace of formaidehyde. Expts. at higher PH values could not be carried out owing to the precipitation of iron(II1) hydroxide. E. a-Hydroxy Acids. ( a ) Malic Acid.--A series of experiments over the pH range 0.5 to 8.6 was carried out using a reaction mixture containing 0.001 mole of triincthylamine oxide, 0.003 mole of Fe3 * and 0.03 mole of (*)-malic acid. The results are shown in Fig. 2. 7
I
---
---------
i
2
0
ti
4
10
8
pE, Fig. 1.--Effect of pH on conibiiied yield of forniaidehyde and formic acid in t h e denethylation of triniethylaniiue oxide at 80" for 40 inin. with iron(II1)-tartrate conlpltx.
B. Citric Acid.-Solutions coiitain~ng 0.001 mole of trimethylamine oxide were examined (a) with 0.003 mole of Fe3+ and various ratios of citric acid (0.003 to 0.03 mole) at $11 0.6 t o 8.6 and (h) with 9.Cl03 mole of the preformed complex iron(II1) ammonium citrate (pH 6,6). All solutions were kept a t 80" for 40 min. The results are shown in Table 11. TABLE !I EFFECT O F
Fe3+
IN
CITRIC A C I D O N 1 I I K CATALYTIC .kCTlVITS O F DEI\IETIIYI,ATION O F TRIMETHYLAMINE NOSIDEO
THE
Clielating agent
Citric acid (0.03 mole) Citric acid ( .03 mole) Citric acid ( .03 mole) Citric acid ( ,015 inolej Citric acid ( ,006 mole) Citric acid ( 003 mole) Ferric ammonium citrateb (0.003 inole)
Yield (%) (CHz0 f HCOOH)
Ratio of citric arid:Fe
pH
10: 1 10: 1 10: I 5 :1 3: 1 1: 1
9.6 i, 1 8.6 4.4 0.8 0.8
21 23
1:1
6.6
0
0 0 0
0
2
4
G
8
10
PFK. Fig. 2.---Efiect of pII o n coii"oi:ied yield of forinaldehj d e arid formic acid in the demethylation of trimethylamine oxide a t 80" for 40 min. with: iron(II1)-aspartate complex, 0; iron( ITI)-malate complex, 0 ; iron( 111)--glycinate complex,
a.
0
All solutiolis contained 0.001 mole of trimethylamine oxide arid 0.003 mole of ferric nitrate 9-liydrate and were kept at 80' for 40 inin. ii Preformed complex.
C. Tartaric Acid.-.$ solution containing 0.001 mole o f trimethylnmine oxide, 0.0133 mole of iron(II1) nitrate 9hydrate and (?.OS mole L(+-)-tartaric xcid was adjusted to the desired plI with the use of saturated sodium c:irbonate solution. T h e firial volume varied between 25 and 40 m l . The solution was kept a t 80' for 40 min.. and the experiments were run over the PH range 0 . 5 to 10. The results are presented in Fig. 1. When the formaldehyde was determined in the optimum pH range (prior to reduction)
(b) Lactic Acid.-Lactic acid was tested under the optimum tartrate conditions, i . e . , in a reaction mixture containing 0.001 mole of trimethylamine oxide, 0.003 mole of Feat and 0.03 mole of lactic acid, pH 6.3 [sodium-(&)lactate was used]. A yield of 670 was obtained as compared with 48y0for tartaric acid under the s a m e conditions. ( c ) Glycolic Acid.-This acid could not he adequately tested as under the optimum tartaric acid conditions glycolic acid was unable t o keep the iron in solution. On heating, most of the metal precipitated as the hydroxide. F. a-Amino Acids. (a) Aspartic Acid.-The (&)aspartate system was examined over the pH range 3 t o 7. As shown in Fig. 2, the results were closely analogous t o those observed with malic acid. No experiments could be performed below pH 3, owing t o the reduced chelating capacity of aspartic acid in strongly acid solutions. Above PH 6.5 some precipitation of iron( 111) hydroxide was observed on heating. An expt. using L-aspartic acid was performed under the optimum (& )-aspartic acid conditions,
April 20, 1961
MECHANISM OF TERTIARY AMINEOXIDEREARRANGEMENTS
i.e., pH 4.8. The yield was 5870 as compared with 59% )-acid under the same conditions. for the (i ( b ) Glycine.-The Fe3+-glycine system was examined over the pH range 2 t o 6, using a reaction mixture with 0.001 mole of trimethylamine oxide and 0.003 mole of glycine. The results obtained were similar to those described above for aspartic and malic acids (Fig. 2 ) . (c) Glutamic Acid.-Experiments involving 0.015 mole of L( Lglutamic acid, 0.001 mole of trimethylamine oxide arid 0.003 mole of Fea yielded negative results a t p H 4.1 and 7.7. The colors given by the Fe3+-glutamate complex were much paler than those from the corresponding aspartate complex a t each of the pH values examined. Below PH 4.1 the glutamic acid precipitated to a large extent, while a t p H 7.7 precipitation of iron(II1) hydroxide was observed. ( d ) Asparagine.-No reaction occurred in the presence of L-asparagine (0.03 mole) a t PH 3.2. Other Chelating Agents.-No reaction occurred after 40 miri. a t 80' and a t various p H values with the amine oxide (0.001 mole) and 0.003 mole of Fe3+ in the presence of salicylic acid (0.03 mole), 8-hydroxyquinoline (0.015 mole) or ethylenediaminetetraacetic acid (0.03 mole). Catalysis by Complexes of Cu(I1) and Other Metals.--The experiments summarized in Table I11 were carried out a t 80" for 40 min. using Cu2+ (0.003 mole) and chelating acid (0.03 mole). Using the optimum conditions for the Fe(II1)-tartrate system, no rearrangement could be detected with Zn2+,Mn2+,Xi2+and AIss ions.
+
1875
ratio to 6: 1 resulted in complete suppression of rearrangement. The same conclusions could be reached from the citrate system (Table 11). Here again the highest activity was obtained using a citric acid:Fe ratio of 1:1 a t pH Q.8; use of a 2 : 1 ratio reduced the reaction, while no rearrangement occurred with 5 : 1 or 10: 1 ratios. The inactivity of the preformed iron(II1) ammonium citrate complex a t pH 6.6 suggests that, as in the case of oxalic acid. the binding of Fe(II1) by the chelating agent is so strong that under practical conditions the equilibrium Fe(I1I)
+ nAH
Fe(III)A,
+ nH+
lies far to the right except a t low pK values. Results obtained with the tartaric acid system (Fig. 1) showed that substantial rearrangement occurred in the pH range 5 to 9, with a fall in the formic acid yield on combined formaldehyde either side of this range. The color changes observed on addition of the amine oxide to the Fe(111)-tartrato complex suggest the addition of hydroxo- or aquo groups to Fe(II1) complexes, TABLE I11 and again indicate binding of the N-oxide to the EFFECTOF CHELATING ACID ON THE CATALYTIC ACTIVITYOF metal complex. The same color changes occurred Cu2+ I N THE DEVETEIYLATION OF TRIMETHYLAMINE OXIDE" with malic acid, being most intense a t the optimum pH for the reaction. Yield The a-hydroxy- and a-amino-dicarboxylic acids Chelating agent $€I (HCHO -k HCOOH), % gave interesting results. Both aspartic and malic L( +)-Tartrate 6.0 3 acids had marked pH optima for the rearrangement (&)-Malate 4 7 2 a t about pH 5 (Fig. 2). Glycine showed a reduced Oxalate 1.2 Trace efficiency, with an optimum a t a somewhat lower (&))-Aspartate
5.0
3
+
PH.
The lack of reaction with asparagine and glycine under the optimum conditions for aspartic acid may indicate that additional binding by the second carboxyl group was beneficial; the inactivity Catalysis by Complexes of Transitional Metals.of glutamic acid may therefore be ascribed to the The complexes were used in 0.27, aqueous solution adjusted stereochemistry of the latter acid. For bivalent t o the appropriate p H value and containing 50 mg. of the cations, the stability of the 1: 2-glutamic acid comamine oxide. S o catalytic effort was observed with: (1) aquoshlorotetraniminecobalt(II1) chloride, (2) carbona- plex is appreciably less than that of the 1 : 2 tobis-(ethy1ediamine)-cobalt(III)chloride, (3) cis and trans- aspartic acid one and is comparable with the value dichlorobis-(ethylenediamine)-cobalt(II1) chloride, (4) po- for g l y ~ i n e . ~ Thus glutamic acid shows a markedly tassium tris-(oxa1ato)-chromate(III), ( 5 ) potassium trislower affinity for Co(I1) a t all fiH values and does (oxahto)-cobaltate(IIT), (6) potassium dioltetra-(oxa1ato)cobaltate(II1) [Durant's salt] or ( 7 ) acetylacetone-his-( 1,- not bind the metal significantly below PH 7, 10-phenanthro1ine)-ruthenium(11) chloride. compared with pH 5 5 for aspartic acid.l0 The approximate yields (vide supra) with other comThe results of the experiments employing 8plexes are summarized in Table I V . hydroxyquinoline, ethylenediaminetetraacetic acid and salicylic acid confirmed that the mere retention Results of iron in solution was insufficient to produce the The results obtained using the oxalate system rearrangement; of other common metal ions ex(Table I ) show that the rearrangement occurred a t amined, only Cu(I1) gave slight reaction (Table 80 to 100" with an oxalic acid:Fe ratio of 1.25:l. 111) in agreement with the findingssd that this No reaction took place below 80". Iron(II1) ion metal was slightly effective for the rearrangement of showed a small catalytic effect a t loo", while no bufotenine oxide. decomposition occurred with the amine oxide alone The necessity for free binding sites on the comeven a t this temperature. When the oxalic acid: plex having been established, the negative results Fe ratio was 10: 1, the reaction was largely sup- obtained with the complexes of the transitional pressed, as it was also when the tris-oxalato fer- metals listed above indicated further that the mere rate(II1) complex was employed (ratio 3 : 1). presence of a metal complex possessing two free or The need for one or more available binding sites available chelating positions was not in itself on the complex was further indicated by the effect (9) D. 3. Perkins, Biochem. J . , 56, 649 (1963). of adding oxalate ion to the tris-oxalato ferrate(10) J. Z. Hearon, D Burk and A. 1,. Schade, J . ~Vatl.Cancer Ins!.. (111) system; the increase in the oxa1ate:Fe 9, 337 (19191. All solutions contained 0.001 mole of trimethylamine oxide, 0.003 mole of Cu2+and 0.03 mole of chelating acid and were kept a t 80" for 40 min. Q
1876
J. C. CRAIG,F. P. DWYER,A. N. GLAZERA N D E. C. HORNING
Trol. 83
sufficient to induce catalysis of the rearrangement. The positive results summarized in Table I V pointed to the need for the metal complex, in
large amounts of iron (111) hydroxide separated and substantial rearrangement took place. It is not clear whether the reaction was catalyzed by iron(I1) or iron(II1) ions since the latter had some activity (Table I). TABLE 1V Consideration of the complexes with catalytic EFFECT OF COMPLEXESOF TRANSITIONAL METALS IS activity (Table IV) shows that all are derived from CATALYZING THE DEMETHYLATION OF TRIMETHYLAMINE N- metals that can exist in a higher oxidation state. OXIDE This applies also to the iron(II1) and Cu(I1) complexes to be discussed later. However, mere particiYield (%) pation in an oxidation-reduction reaction does not Complexf $H (CH20 + HCOOH) Ref. suffice to promote rearrangement. This is illustrated &Ru( 111)Cl~.HzO 2 12 204 by the chemically stable red substance acetylace1 12 KzRu( I11 )Clb.HzO 30b tone bis-(1,lO-phenanthro1ine)-ruthenium(I1)chlo1 12 K~RU(IV)CI~OH 30 ride“ which is oxidizable to the blue Ru(II1) 45. 3 13 &Ru( CzOa)a complex (Eo = - 0.65 v.), and thence to the brown 40d 11 KRuCLphen 3 Ru(1V) complex (E, = - 1.3 v.). Though the blue 10 11 [Rubipy,Cl.HzO ] C1 5 complex was produced in the reaction, no rear2-4 Trace 14 ( KH&OsBrB rangement occurred. The vanadyl ion exists in 11 4 15 [OsphenzBr2] aqueous solution as various hydrated species such 3 Trace 18 [RhCls.3HaO] as [VO(H20)b]2f and [V(OH)2(H20)4I2+both of 2 15 Trace KsIrCle which can attach the N-oxide (probably rather .. 1.5 .50e \.OClz weakly) by the replacement of a molecule of water. The value of the potential” of the reaction a Black ppt. of RuOz deposited. Color changed from 2H20 = V(OH)d+aq. 2H+ e-, Color VO+aq. red to red-brown due to formation of KzRuClbOH. (Eo = -1.00 v.) also seems especially favorable as changed from green to red brown (KzRu( C~O&(HZO)~) a result of oxidation and hydrolysis. Color changed from for the rearrangement, because it is evident yellow to dark brown as a result of hydrolysis and oxida- that two competing reactions are always involved : tion. e Color changed from blue t o brown-green, slight brown ppt. (VZO~). f Phen = l,10-phenanthroline, bipy one leading merely to oxidation of the catalyst and the other of a methyl group of the N-oxide. = 2,2 ‘-bipyridine. Although the potential for the oxidation of chloroaddition to providing the requisite two binding bipyridineterpyridine-ruthenium(I1) chloride11 (Eo sites, to be able to exist in a suitable higher oxi- = -0.93 v.) is close to the V(1V)-V(V) potential, dation state. These findings are discussed below. and a point of attachment for the N-oxide can be provided by dissociation of a chlorine atom, the substance had no catalytic effect. This suggests Discussion the necessity of an aquo or hydroxy group coIn common with other amine oxides, trimethyl- ordinated to the metal adjacent to the bonding site amine oxide is an oxidizing agent. The oxidizing of the N-oxide. This can be provided with all of action in the cold is slow and irreversible. In the complexes in Table (IV) by dissociationhot acid solution (fiH l), i t rapidly oxidized iron- aquation reactions. There appears to be no (11) to iron(II1) (Eo = -0.772 v.) with pre- functional value in a hydroxy or aquo group in the cipitation of iron(II1) hydroxide and very slowly trans position, and in any event other chemical evidence1‘ suggests that chloroaquobis-(2,2’-bioxidized bromide ion to bromine’? (Eo = -1.065 pyridine) -ruthenium(II) chloride and dibromobisv.) and tris-(2,2’-bipyridine)-iron(II)ion1* (Eo = -1.096 v.), some decomposition of the latter ( 1 , l O - phenanthro1ine)-osmium(I1) have the cis substance occurring simultaneously in the hot structure. A number of cis and trans isomeric acid solution. It can be inferred that the potential forms of ruthenium complexes have been deof the reaction Me3N0 2H+ = CHIN H20 scribedIg and some proportion a t least of the cis 2e is of the order of - 1.1 v. Formaldehyde form can be expected to exist in the solutions of the [Ruwas found only in the presence of iron(I1) ion. aquated [RuC160HI2-, [ R u C l ~ H 2 0 ] ~and Under the standard experimental conditions used phenC14]- ions. The tris-oxalatoruthenate(II1) with the iron(II1) complexes (0.001 mole), traces ion, though spin-paired, cannot be resolved and only could be detected, but a t 100”in the presence yields a precipitate of calcium oxalate if boiled of high concentrations of iron(I1) ion (0.1 mole), with calcium salt solution^.'^ Dissociation of the oxalate group would yield initially the cis diaquo (11) F. P. Dwyer and H. H. Goodwin, THIS JOURNAL, 83, in complex which may rearrange. press With the ruthenium and osmium complexes, (12) R. Charonnat, A n n . C k i m . , 16, 5 (1931). (13) F. P. Dwyer and A. M . Sargeson, J . P h y s . C k e m . , 60, 1331 the higher oxidation states of 4-4 and +6 are (1956). available especially after halogen atoms have been (14) F. P. Dwyer and J. ‘8. Hogarth, Inorg. Syntheses, 5, 204, replaced by aquo and hydroxy groups. (1957) Aquation of trichlorotriaquorhodium(II1) and (15) M. Delepine, Bull. SOL. ckim., 9, 771 (1911). (16) J. Meyer and M. Kawezyk, 2. onorg. Ckem., 228, 297 (1936). hexachloroiridate(II1) ions proceeds rather slowly (17) W. M. Latimer, “The Oxidation States of the Elements and as evidenced by the rate of precipitation of silver their Potentials in Aqueous Solutions,’’ Prentice-Hall Book Co., S e w halide. The potential17 of the system Rh3+ York, N. Y . , 1952, pp. 340-345.
+
+
+
+
+
+
(18) W. W. Brandt, F. P. Dwyer and E . C. GyarfaJ, Ckem. Reus., 64, 959 (1954).
(19) N. V. Sidgwick, “ T h e Chemical Elements and their Compounds,” Vol. 11, Clareqdon Press, Oxford, 1950, pp. 1470, 1471.
MECHANISM OF TERTIARY AMINEOXIDEREARRANGEMENTS
April 20, 1961
+
+
2H+, E , = 1.42 v. (absence HzO = R h o + of halide ion) suggests that the N-oxide would be ineffective as an oxidizing agent. On the other hand a lower value (E, = -1.2 v.) has been adduced17 for the system [RhC13l3- = [RhClsl2e-. It is possible that the potential of some of the chloro-aquo or hydroxy-chloro-aquo species may be more positive. The potential of the [IrC16l3- [IrCl6l2-systemp0is - 1.017 v. Catalysis by Iron(II1) and Copper(I1) Complexes.-Coordinately saturated iron(II1) complexes such as the tris-(oxa1ato)-ferrate(II1) ion in the presence of excess oxalate ion and the ethylenediaminotetraacetatoferrate(II1) ion do not catalyze the rearrangement. Slow rates of equilibration are characteristic of many iron(II1) complexes, and mixtures of iron(II1) ion and a chelating anion, e.g., oxalate ion, even under favorable p H conditions, are transformed to the tris complex fairly slowly. When the iron(II1) salt is partly hydrolyzed (above pH 3) , depolymerization of the complex hydroxo entity markedly reduces the rate of formation of other complexes. For these reasons, i t is probable that during some of the rearrangement experiments a mixture of complex species was present. Since the chelating agents attach in the anion form it is evident that chelation is favored by a relatively high pH value, but on the other hand competition by hydroxyl ions for the coordination positions about the metal may then become serious. Investigations of the stability of iron(II1) complexes with a-amino acids has shown that a t the highest practicable pH value (4.7), stable 1: 1 complexes are formed.21 In less acid solution separation of iron (111) hydroxide obscures the possible existence of higher complexes. Water and/or hydroxyl groups presumably occupy other coordination positions in the 1 : l complexes. The iron(II1) oxygen complexes usually have two or three molecules of the ligand attached, e . g . , [Fe(CZ0&l3and [Fe(C204)2(OH2)2]-. The structures of the complexes and mode of attachment of tartrato and citrato groups is not known with any certainty and hence the high stability in alkaline solution is not understood. The tartrato complex acts as a catalyst up to p H 10. Relatively, the oxalato complex would seem to be a better catalyst than the citrato a t the same p H value. This may well be because the oxalato complex is weakerz2 (log K = 9.4 compared with 11.85), i.e., a higher proportion of the bis complex, with appropriate sites, is present. The evidence for the existence of high oxidation states of iron has been summarized by K l e i r ~ b e r g and , ~ ~ in recent articles on the mode of action of catalase,24peroxidasez4 and other catalytically active h a e m o p r o t e i n ~ . ~The ~ ion di-
+
1877
chlorobis- [o-phenylenebis- (dimethylarsine) 1- ironz6 undoubtedly contains tetravalent iron. None of the iron complexes used in this work have ever been oxidized to a higher oxidation state, possibly because no serious attempt has been made. On general principles, oxidation of the bis-(chelate)diaquo species should be favored since not only is the detachment of electrons facilitated by the negative charge but dissociation of a proton to form a hydroxyl group would tend to stabilize the oxidized form. The catalytic mode of action envisaged requires only a transitory existence of the oxidized form, which might well be more of the nature of a highly polarized state of the iron atom than a discrete oxidation state. The photoactivated state of tris-(2,2’-bipyridine)-iron(II) salts in which negative charge is transferred from metal to ligand constitutes this kind of internal oxidation-reduction system.21 The situation regarding the oxidation of Cu(I1) complexes is not unlike Fe(II1). A complex Cu(111) fluoridez8 is known and also complex periodatesJ2$but other Cu(I1) complexes have not been oxidized. Mechanism of Rearrangement.-Trimethylamine N-oxide is a weak base, (pK, = 4.6), existing in acid solution as the cation [(CH3)3-
(20) F. P . Dwyer, H. A. McKenzie and R. S. Nyholm, 1.PYOC. Roy. Sac. N.S.W., 78, 260 (1946). (21) D.D. Perrin, J . Chem. Soc., 3120,3125 (1958). ( 2 2 ) J. Bjerrum, G. Schwarzenbach and L. G. Sillen, “Stability Constants,” Chem. SOC.Special Publication, No. 6, London, 1957. (23) J. Kleinberg, “Unfamiliar Oxidation States,” University of Kansas Press, Lawrence, K a n . , 1950. (‘24) N. K. King and M. E. Winfield, Aztstral. J. Chem., 12, 47 (19.59). ( 2 5 ) E. C. Slater, Biochemical Society Symposium No. 15, Cambridge C‘niv. Press, 7fi (1958).
(26) R. S.Nyholm, i b i d . , 101 (1958). (27) R.J. P. Williams, Chem. Reus., 16, 299 (1956). (28) W. Klemm and H. Huss, 2. anorg. Chem., 258, 221 (1949). (29) L.Malatesta, Gam. chim. if& 71, 467 (1961).
1878
J. C. Cam,
F. P. Dwmn, A. N. GLlrpn AHD E.C. HaumK,
NOH]+.
COardinatiOn to I metal mwt occuz, with amine oxidu~,"t h u g h the oxysm atom. A#anipg the llwlpI noo-linearity of the oxygen boa& it can be g e c ~ lnadily fmm modela that one bgdr0gCp;atomofamcthylgrou approachcrthe atanofanadjacentcod inatcdOHgmu OT moaculc. When the N-0-M bond a& it about 120°, the atoms are clo# enough to form a dx-membad ring through a hydrogen bond. The marmuganent is then formulated aso). It is proposed that the metal atom (Fe(III)] loam an electron to the oxygen atom of the Noxide, thaeby b m i n g oddized to Fe(IV). The N-O bond, weakened as a d t , ruptures, the oxygen atom taking with it one electmn from the in the nitrogen (step 2). [An wuo group (Ha) Fe(II1) complex would shed a proton on oxidation to Fe(IV).] In this way the oxygen atom, attached now to the iron atom as an OM) group, has been reduced in two l-electron steps. Attack by ti proton detached from the methyl group then transforms it to a hydmxo (OH) gratrp, and the odd el=(u
tg
%
(80) R. 8. N H m , J. CLm
*
1767 (lOal).
Vdm
*
tam on the nibrogrp atom kwxnrrar peked wifh whidl har migrokd from thcsrboa atom (m3). (The proton rtmovad lrotn th methyl gmup need not ncccrwtily be the OBIP! attacking the oxo group, however.) The mdoI atom accepts 1~11 efectton from the aobrdinstsd hyPIP
dmxylgrcwp,nvcrtin inthepnn#rtoitrorigIsat reduced &ate [Fe(IIIfJ (%4), the 0Ol)rdinrrtcd hydroxyl group being there oxidized to a hydroxyl radical which attacks the unpaired dcctron on the methylene carbon atom to give the methylolamine. An q u o group fillrr the wrcant site on the complex, regenerating the original Fe(111) chelak This mechanism is closely &lat to that PO&lated for the radical-matedPokmovski nBcfioLI, (D) above, and to tbe WJbfree-radical rearrangement of 2-methypyridine "p" Naide, (C) above, and is compatible with the experimental facts. An ionic mechanism would not explain the necessity for a valency change of the complex during the d o n . Further examination of thh reactionisinprogress.
