J. Phys. Chem. 1984,88, 2857-2862 planation why an increased period length is observed when Ag+ ions are slowly added to a nonoscillating excitable BZR using an M) initial Ce(1V) concentration.16b*c increased (2.1 X In addition, Ag+ ions may also accelerate the conversion of bromo organic compounds (as, for example, -CBr(COOH)2)to the corresponding hydroxy organic intermediates (A8). Reactions J2Br(COOH)2
+ Ag+
-
.C(OH)(COOH),
+ AgBr + H+
(A8) such as the A8 may be of importance in further Br- releasing (radical) reaction steps as indicated in the work of Edelson et ala8 The nearly constant negative phase shift using high concentrations of AgN03 added is also remarkable (Figure 8D). Topologically, this may be explained by interpreting A as an absolute refractory period.16b*cWhen the system is perturbed inside A, a refractory period implies that the system has to wait the time ( A - t’) until the next Ce(1V) spike can be induced (in order to oscillate with the period length A). Thus, according to eq 4, the phase shift is A 4 = (A - t’) - (Po- t’) = ( A - Po) t’ < A (7a)
A 4 = ( 2 ’ - Po) 2’2 A (7b) Using an approximate value of Poof 40 s and a A value of 15 s, we obtain the solid line in Figure 8D. However, as mentioned above, the real system’s response is always subject to a certain delay due to an intermediate not in the model-realized reaction steps. To account for this, we may add a suitable small positive constant to eq 7b. This is indicated by the dashed line in Figure 8D. The mechanistic explanation of the (nearly) constant phase shifts in Figure 8 appears also to be due to the formation of CeBr3+. The effect of CeBr3+ may be considered as a type of Brbuffering, since when Ag+ ions are added, the sum of reactions A7 and Al-A3 balance the Br- removal of AgN03. The experimental data of Figure 8 support this. At the beginning of the cycle, we have a high Ce(1V) concentration while the Br- concentration increases. Therefore, also CeBr3+becomes high and should cause, due to reactions A7 and Al-A3, an increase of the
2857
(positive or negative) phase shift when Ag+ ions are added. However, as both the Ce(1V) and the Br- concentration decrease slightly later in the cycle, the CeBr3+ amount also should be lowered and more negative phase shifts should be the result when Ag+ ions are added. Therefore a local maximum of the phase shifts at the beginning of the cycle should be observed. At a “phase of stimulation” of about 0.1 such a maximum is indeed found (Figure 8B-D). This again demonstrates the important role Brions play in this system.
Concluding Remarks Although our experimental and calculated data are still in semiquantitative agreement, our results clearly demonstrate the Br- control of the free oscillating MA BZR. Even seemingly contradicting results due to a higher initial Ce(1V) concentration confirm the basic FKN mechanism and provide an explanation of the increased period lengths observed when AgNO, is added at low rates of flow into an excitable BZR with same initial concentrations of reagents. The goal of further work is to construct a mechanism which accounts more quantitatively for these effects. Certainly, phase shift experiments may also be applied to other related chemical oscillators. Although similar methods are well-established in the investigation of biochemical or biological rhythms, purely chemical oscillating systems seem especially well-suited to this form of investigation, since here most of the state variables are generally known, provided a realistic chemical mechanism exists. Note Added in Proof. While this paper was in press, the author became aware of two papers using similar methods for studying the BZR28 and the Briggs-Rauscher o s ~ i l l a t o r . ~ ~ Acknowledgment. The author is indebted to Dr. Cathrine Lillo, Dr. Knut Fzgri, and Dr. Bengt Schwitters for discussing the manuscript. This work was supported in part by H~yesterettsadvokatPer Ryghs legat. Registry No. Br-, 24959-67-9; HOBr, 13517-11-8;Ag, 7440-22-4;Ce, 7440-45-1; Br03-, 15541-45-4;CH2(COOH)2,141-82-2.
Reactions of Iron( I I I ) Porphyrins with Peroxyl Radicals Derived from Halothane and Halomethanes D. Brault and P. Neta* Laboratoire de Biophysique, INSERM U.201. CNRS ERA 951, Museum National d’Histoire Naturelle, 75005 Paris, France, and Chemical Kinetics Division, Center for Chemical Physics, National Bureau of Standards, Washington, D.C. 20234 (Received: October 25, 1983)
The reactions of haloalkane-derivedperoxyl radicals with ferric deuteroporphyrins in aerated acidic or alkaline aqueous 2-propanol solutions are investigated by means of pulse radiolysis. CC1302,CHC1202,CH2C102,and CF3CHC102radicals (the latter one being derived from the anesthetic agent halothane, CF3CHC1Br)are found to oxidize the ferric porphyrins with reaction rate constants ranging between 6 X lo7 and 2.6 X lo8 M-’ s-l. In keeping with an electron-transfer mechanism, the spectrum of the oxidized ferric porphyrin does not depend on the nature of the peroxyl radicals. Also, the rate constant for the reaction of CC130, radicals with ferric porphyrins is lowered by a factor 220 when experiments are performed in the less polar solvents neat 2-propanol and neat carbon tetrachloride. The spectrum of the oxidized ferric porphyrin depends on pH with large changes around pH 2.3 which are attributed to the protonation of an alkoxide ligand of the iron ion. Smaller changes are observed with pK, E 7.4. The relevance of these reactions to cytochrome P4somediated oxidative metabolism and toxicity of haloalkanes is outlined.
Introduction Halogenated alkanes are metabolized by cytochrome P450 to oxidized and reduced products. Trifluoroacetic acid,’ chloro(1) K. Rehder, J. Forbes, H. Alter, 0. Hessler, and A. Stier, AneslhesiOZO~Y, 28, 711-5 (1967).
0022-3654/84/2088-2857$01.50/0
trifluoroethane, and chlorodifluoroethylene’ are produced from halothane (CF3CHClBr), a common anesthetic agent. In the same way, the metabolism of carbon tetrachloride leads to phosgene (COC12)>4carbon dioxide; chloroform,b8 and hexachloroethane.6 (2) J. H. Sharp, T. R. Trudell, and E. N. Cohen, Anesthesiology, 50, 2-8
(1979).
0 1984 American Chemical Society
2858 The Journal of Physical Chemistry, Vol. 88, No. 13, 1984
Reactive intermediates formed in the course of the enzymatic degradation are believed to be responsible for the toxicity of these corn pound^.^^'^ Mechanisms can be inferred from in vivo and in vitro studies on biological systems1-I0 as well as kinetic and spectrophotometric studies of the reactions of iron porphyrins with halogenated and their radical^.'^-'^ Although other schemes have been postu1ated,l6-l8it is likely that, in most cases, the first step in the metabolism of halogenated compounds is their reduction to form r a d i c a l ~ . ~These J ~ radicals may be formed by various reaction^,'^^^^ but the most likely is electron transfer from cytochrome P450:I1,l2 PFe"
+ RX
-
PFelI1 + R
+ X-
(1)
In the case of halothaneI5 and carbon tetrachloride,13 the radicals do not react rapidly with PFerrl and can escape from the iron porphyrin pocket. The radicals may then react with cellular components, especially lipid^.^^^^ Actually, radical adducts have been detected during in vivo and in vitro metabolism of carbon tetrachloride21 and halothane.22 However, the radicals are also likely to react with oxygen either after or before they escape the iron porphyrin pocket
R
+ 0,
-
RO,
(2)
with k2 of the order of lo9 M-I s-1,23-26 In the former case, the
(3) M. Shah, S. P. Hartman, and S. Weinhouse, Cancer Res., 39, 3942-7 (1979). (4) V. L. Kubic and M. W. Anders, Life Sci., 26, 2151-5 (1980). (5) D. Rubinstein and L. Kanics, Can. J . Biochem., 42, 1577-85 (1964); B. B. Paul and D. Rubinstein, J . Pharmacol. Exp. Ther., 141, 141-8 (1963). (6) S. J. L. Fowler, Br. J . Pharmacol., 37, 733-7 (1969). (7) T. C. Butler, J . Pharmacol. Exp. Ther., 134, 311-9 (1961). (8) H. Uehleke, K. H. Hellmer, and S . Tabarelli, Xenobiotica, 3, 1-1 1 ( 1973). (9) T. F. Slater, "Free Radical Mechanism in Tissue Injury", Pion, London, 1972. (10) R. 0. Recknagel and E. A. Glende, Jr., CRC Crit. Reo. Toxicol.,2, 263-96 (1973): R. 0. Recknad. E. A. Glende. Jr.. and A. M. Hruszkewvcz in "Freekadicals in Biology", 701.3, A. W. Pryor, Ed., Academic Press, New York, 1977, pp 97-132. (11) D. Mansuy, M. Lange, J. C. Chottard, P. Guerin, P. MorliEre, D. Brault, and M. Rougte, J . Chem. Soc., Chem. Commun., 648-9 (1977). (12) D. Brault, P. Morlitre, M. Rougce, and C. Bizet, Biochimie, 60, 1031-5 (1978). (13) D. Brault, C. Bizet, P. Morlitre, M. RougCe, E. J. Land, R. Santus, and A. J. Swallow, J . Am. Chem. SOC.,102, 1015-20 (1980). (14) D. Brault and P. Neta, J . Am. Chem. SOC.,103, 2705-10 (1981). (15) D. Brault and P. Neta, J . Phys. Chem., 86, 3405-10 (1982). (16) D. Mansuy, W. Nastainczyk, and V. Ullrich, Naunyn-Schmiedeberg's Arch. Pharmacol., 285, 3 15-24 (1974). (17) I. G. Sipes, A. J. Gandolfi, L. R. Pohl, G. Krishna, and B. R. Brown, Jr., J . Pharmacol. Exp. Ther., 214, 716-20 (1980). (18) L. R. Pohl, B. Bhooshan, N. F. Whittakev, and G. Krishna, Biochem. Biophys. Res. Commun., 79, 684-91 (1977); D. Mansuy, P. Beaune, T. Cresteil, M. Lange, and J. P. Leroux, ibid., 79, 513-7 (1977). (19) B. A. Mico, R. V. Branchflower, L. R. Pohl, A. T. Pudzianowski, and G. H. Loew, Life Sci., 30, 131-7 (1982). (20) A. T. Pudzianowski, G. H. Loew, B. A. Mico, R. V. Branchflower, and L. R. Pohl, J . Am. Chem. SOC.,105, 3434-8 (1983). (21) J. L. Poyer, P. B. McCay, E. K. Lai, E. G. Janzen, and E. R. Davis, Biochem. Biophys. Res. Commun., 94, 1154-60 (1980); J. R. Trudell, B. Bosterling, and A. J. Trevor, Proc. Natl. Acad. Sci. U.S.A., 79, 2678-82 (1982). (22) J. R. Trudell, B. Bosterling, and A. Trevor, Biochem. Biophys. Res. Commun., 102, 372-7 (1981); J. L. Poyer, P. B. McCay, C. C. Weddle, and P. E. Downs, Biochem. Pharmacol., 30, 1517-9 (1981). (23) J. E. Packer, R. L. Willson, D. Bahnemann, and K. D. Asmus, J. Chem. SOC.,Perkin Trans. 2, 296-9 (1980). (24) R. Cooper, J. C. Cumming, S. Gordon, and W. A. Mulac, Radiat. Phys. Chem., 16, 169-74 (1980). (25) J. Rabani, M. Pick, and M. Simic, J . Phys. Chem., 76, 1049-51 (1974); J. Rabani, D. Klug-Roth, and A. Henglein, ibid.,78, 2089-93 (1974); B. Maillard, K. U. Ingold, and J. C. Scaiano, J . Am. Chem. SOC.105, 5095-9 (1983).
Brault and Neta R 0 2 radicals, which are much more reative than R radicals toward several c o m p o ~ n d s , ~would ~ * ~be ~ ,the ~ ~toxic intermediates. In the latter case, we must also consider, the reaction of R 0 2 radicals with PFe"', a step which might be important in oxidative metabolism. In order to examine this possibility we have carried out pulse radiolytic studies on the reaction of several halogenated peroxyl radicals with iron(II1) porphyfins. We have already noted that the peroxyl radical (CF3CHC10,) derived from halothane reacts with iron(II1) porphyrins.I5 Now, we present a detailed study of this reaction under more favorable conditions as well as reactions of peroxyl radicals derived from carbon tetrachloride and other halomethanes. All these peroxyl radicals react rapidly with iron(II1) porphyrins.
Experimental Section" The preparations of ferrideuteroporphyrin IX chloride (DPFe"'C1) and of the pox0 dimer of ferrideuteroporphyrin IX dimethyl ester ((DPDMEFe"'),O) were described e l s e ~ h e r e . ' ~ . ~ ~ Halothane (ICI) was destilled before use. Methyl iodide was purified on activated alumina. The remaining halogenated compounds and solvents were of the purest grade commercially available. Water was purified by a Millipore Milli-Q system. Solutions. Stock solutions of DPDMEFe"' were prepared by dissolving (DPDMEFe111)20in 2-propanol and letting it stand overnight. Under these conditions, the pox0 dimer is split.29 These solutions were mixed with 2-propanol, water, and HC104 (70%) to the desired final concentrations (50% 2-propanol, -4 X M PFeII'). For ex0.001-1.2 M HC104, 5 X periments on pH effects, the DPDMEFe"' solutions in 2-propanol were mixed with aqueous buffer solutions (2 mM phosphate or borate) or with NaOH solutions. Alternatively, alkaline solutions were prepared by dissolving DPFe"'C1 in 0.1 N NaOH and then adding 2-propanol to a final concentration of 6.5 M (50%). The halogenated compounds were added to the above solutions to the required concentration (0.2-0.3 M). The coordination state of ferric porphyrins in the aqueous 2-propanol solutions depends on the pH.29330 In slightly alkaline or neutral solutions, the ferric ion is likely coordinated by one 2-propoxo ion and one 2-propanol molecule (PFe"'0CH(CH,),(HOCH(CH,),). The 2-propoxo ion protonates in acidic medium leading to the positively charged species [PFe"'(HOCH(CH3)z)21f. Experiments were also carried out by using neat CC14solutions of DPDMEFeII'Cl. They were prepared by dissolving (DPDMEFe111)20in CC14, adding a small amount of gaseous HCl, and bubbling out the excess HC1. For all the above solutions, PFe"' will be used as an abbreviation without specifying the ligands except when necessary. The absorption spectra of the solutions were recorded by using a Cary 219 spectrophotometer and were in agreement with the published spectra of assigned structure^.^^-^^ Pulse Radiolysis. The pulse radiolysis experiments were carried out with a single 50-11s pulse of 2-MeV electrons from the Febetron 705 accelerator at the National Bureau of Standards. The dose per pulse was usually 560 rd, determined by KSCN dosimetry. A fresh solution was used for every pulse. The kinetic spectrophotometric detection system consisted of a 2 cm optical path length irradiation cell, a 300-W Varian xenon lamp, a Bausch and Lomb monochromator, and a photomultiplier. The signals were digitized with a Tektronix 7612 recorder and processed with a PDP 11/34 computer.32 Cutoff filters were used to minimize (26) J. Monig, K. D. Asmus, M. Schaeffer, T. F. Slater, and R. L. Willson, J . Chem. SOC.,Perkin Trans. 2, 1133-7 (1983). (27) J. E. Packer, T. F. Slater, and R. L. Willson, Life Sci., 23, 2617-20 (1978); J. E. Packer, J. S. Mahocd, R. L. Willson, and B. S. Wolfenden, Int. J . Radiat. Biol. Relat. Stud. Phys., Chem. Med., 39, 135-41 (1981). (28) D. Brault and P. Neta, J . Phys. Chem., 87, 3320-7 (1983). (29) C. Bizet, P. Morlitre, D. Brault, 0. Delgado, M. Bazin, and R. Santus, Photochem. Photobiol., 34, 315-21 (1981). (30) A. C. Maehly, Acta Chem. Scand., 12,1247-58 (1958); A. C. Maehly and A. Akeson, ibid., 12, 1259-73 (1958). (31) J. 0.Alben, W. H. Fuchsman, C. A. Beaudreau, and W. S. Caughey, Biochemistry, 7, 624-35 (1968).
The Journal of Physical Chemistry, Vol. 88, No. 13, 1984 2859
Reactions of Iron(II1) Porphyrins
TABLE k Reaction Rate Constants solventu 2-propanol (6.5 M), NaOH (0.05 M) 2-propanol (6.5 M), HC104 (2.3 X M) 2-propanol (6.5 M), HClO, (0.23 M) 2-propanol (6.5 M), HClO, (1.16 M) neat 2-propanol, HC104 (2.3 X M) neat CCl, 2-propanol (6.5 M), NaOH (0.05 M) 2-propanol (6.5 M), HC104 (2.3 X M) 2-propanol (6.5 M), HClO, (2.3 X lo-) M) 2-propanol (6.5 M), HC104 (2.3 X lo-' M)
reactants ferrioorDhvrin I
I
Deroxvl radical
,
cc1302
k, M-' s-I (2.5 f 0.3) X (2.3 & 0.3) X (2.5 f 0.3) X (2.6 f 0.4) X
cc1302
=io7
I
,
DPFe1110CH(CH3)2(HOCH(CH3)2)b CCl3O2 [DPDMEFe"'(HOCH(CH1)?)21+
cc1102
cc1;o;
CCl3O2 CF3CHC102 CF3CHC102 CHC1202 CH2C102
lo8 los lo8 los
(1.2 f 0.4) x 107 (3.5 A 0.5) x 107 (6.0 f 0.8) x 107 (8.2 f 1.0) x 107 (7.7 1 . 5 ) x 107
uThe solvent is added with the appropriate halogenated compound (2% CC14 or 3% CF3CHC1Br or 2% CHC1, or 2% CH2C12). [DPFe111(OCH(CH3)2)2]-might also be present in alkaline solutions. photolysis of the solution by the analyzing light. Experiments were performed at room temperature (21 f 1 "C) with air-saturated solutions. Peroxyl Radical Production. The main species produced in irradiated aqueous 2-propanol solutions are e; and (CH,),COH radicals. The latter are formed by the radiolysis of 2-propan01~~ and also by scavenging of H and O H produced by water radiolysi~.~~,~~ In the presence of haloalkanes, the solvated electrons react predominantly with these compounds36 RX
+ e;
-
R
+ X-
(3)
with k, 109-1010M-l s-l. The reactions of e; with 0, and with PFe1I1are negligible in our experiment^.^^ The only reaction to compete with reaction 3 is
H+ + e;
-+
H
(4)
(k4 N 2 X 1OloM-I which becomes important at h!gh acid concentration. It should be pointed out however that H atoms are subsequently scavenged mostly by 2-propanol to form the reducing (CH3)2COH radicals.34 In basic solutions, the (CH3)&OH radicals deprotonate (p& = 12.2), leading to the more reducing anionic form ((CH3)2CO-).37 The fate of the a-hydroxyisopropyl radicals depends on the nature of the haloalkane used and on the pH. Carbon tetrachloride and halothane react with the neutral and anionic forms of ahydroxyisopropyl radicals
- -+ +
C C 4 + (CH3)2COH CF3CHCIBr
CC13
C1-
+ (CH3)2C0+ H+
+ (CH3)2COH
CF3CHCl
Br-
+ (CH3)$0 + Ht
(5)
(6)
with k, = (1-7) X lo8 M-Is-l (ref 13, 38, 39) and k6 = 3.5 X lo7 s-' (ref 15). The anionic form gives rise to parallel reactions with higher rates ( k e 8 X los M-' s-l (ref 40) for CC14 and k = 5.8 X lo8 M-I s-l (ref 15) for halothane). Thus, either in acid or in basic solutions, CCI, or CF,CHCI radicals are expected to be produced nearly quantitatively from all the primary radicals (32) M. G. Simic and E. P. L. Hunter in "Radioprotectors and Anticarcinogens", 0. F. Nygaard and M. G. Simic, Eds., Academic Press, New York, 1983. (33) G. R. Freeman in "Actions Chimiques et Biologiques des Radiations", Vol. 14, M. Haissinsky, Ed., Masson, Paris, 1970, pp 73-134. (34) M. Anbar, Farhataziz, and A. B. Ross, Natl. Stand. ReJ Data Ser. (US., Natl. Bur. Stand.), NSRDS-NBS 51 (1975). Natl. (35) Farhataziz and A. B. Ross, Natl. Stand. Ref: Data Ser. (US., Bur. Stand.), NSRDS-NBS 59 (1977). (36) M. Anbar, M. Bambeneck, and A. B. Ross, Natl. Stand. Ref.Data Ser. (US., Natl. Bur. Stand.), NSRDS-NBS 43 (1973). (37) J. Lilie, G. Beck, and A. Henglein, Ber. Bunsenges.Phys. Chem., 75, 458-65 (1971). (38) R. Koster and K. D. Asmus, Z . Naturforsch. B, 26, 1104-8 (1971). (39) R. L. Willson and T. F. Slater in "Fast Processes in Radiation Chemistry and Biology", G. E. Adams, E. M. Fielden, and B. 0. Michael, Eds., Wiley, New York, 1975, pp 147-61. (40) D. Brault and P. Neta, unpublished results.
(total yield G 6 radicals/100 eV). Methyl iodide was found to react with (CH,),CO- with k = 1.1 X IO8 M-' sv1,l4which allows nearly quantitative formation of methyl radicals in alkaline solutions. Parallel reactions in acidic medium were not examined. The reactions of CHCI, and CH2C12with either form of the radical derived from 2-propanol were too slow to take place efficiently in pulse radiolysis experirnenka In these cases, haloalkyl radicals. produced from reaction 3 with a yield G N 3, and (CH3)&OH radicals, are present after the pulse. In the presence of oxygen, haloalkyl, methyl, and related radicals react to form the peroxyl radicals with rate constants of the order of IO9 M-' s-l , e .g. CC13 0 2 cc1,02 (7)
+ CF3CHCl + 0,
--
CF3CHC102
(8)
(k7 2 lo9 M-l s-l (r ef 23 and 24), ks y 1.3 X lo9 M-Is-l (ref
-
26)). Oxygen also reacts with (CHJ2COH radicals (CH3)2COH+ 0, (CH3),COH(02)
(9)
( k , = (4-4.5) X lo9 M-' s-I (r ef 41 and 42) and with H atoms H O2 H 0 2 (10) ( k l oE 2 X 1Olo M-ls-l (ref 34)). Thus, CCl3O2, CF3CHC102,and CH302radicals can be produced almost quantitatively (G = 6) within less than 1 p s after the pulse (in the case of CH3O2 radicals, this has been checked only in alkaline medium). The total yield might be slightly affected by competition of reaction 9 with reactions 5 or 6. In strongly acidic solutions the yield may be slightly lowered by competition of reaction 10 with the reaction of H atoms with 2-propanol, but it may be increased due to more efficient scavenging of e; from the spurs by H+.43 For experiments involving CH2C12 and CHCl,, the corresponding peroxyl radicals are produced with G e 3 along with a similar concentration of (CH3)2COH(02).
+
Results Alkaline Solutions. Pulse radiolysis experiments with airsaturated alkaline aqueous 2-propanol solutions containing either C C 4 or CF3CHCIBr and PFe"' indicate that the peroxyl radicals react rapidly with the porphyrin. The differential spectra recorded after the completion of these reactions are given in Figure 1. The spectra are characterized by absorption maxima around 530 and 700 nm and indicate bleaching of the starting compound at 585 and 480 nm. They are not significantly dependent on the nature (41) G. E. A d a m and R. L. Willson, Trans. Faraday SOC.,65, 2981-7 (1969); R. L. Willson, ibid., 67, 3008-19 (1971). (42) J. Butler, G.Jayson, and A. J. Swallow, J . Chem. SOC.,Faraday Trans. I , 70, 1394-401 (1974). (43) The increase of the total yield calculated according to Balkas et al.44 may attain 30% in solution containing 1 M acid which exceeds a 5-10s decrease arising from reaction 10. (44) T. I. Balkas, J. H. Fendler, and R. H. Schuler, J . Phys. Chem., 74, 4497-505 (1970).
2860
The Journal of Physical Chemistry, Vol. 88, No. 13, 1984
I
I
I
a
Brault and Neta
b
0
P
-
x o
u mc
n p-5 n
4 v 0
-10 t
1
t.I,'
500
700
600
(4
Figure 1. (a) Optical spectrum of DPFe"' in aqueous alkaline 2-propanol (2-propanol: 6.5 M; NaOH: 0.05 M). (b) Differential spectra recorded after irradiation of the above solution containing either 2% C C 4 (0: experimental, -: smoothed) or 3% halothane (A, - -). Spectra were recorded 850 ps or 1.2 ms after the pulse and the porphyrin concentraor 1.1 X M, respectively. tions were 4.9 X
-
of the halogenated compound. Kinetic analysis of the reaction of the CC1302and CF3CHC102radicals with PFemlwas performed by following absorbance changes at various wavelengths. Pseudo-first-order kinetics were obeyed within a porphyrin concentration range of 1 X 104-8 X lo4 M. Second-order rate constants are given in Table 1. The reaction of CH3O2 radicals with PFe"' was also investigated by irradiating alkaline solutions containing CH31. Very small absorbance changes were observed which did not permit characterization of the kinetics and spectra. It appears that CH3Oz radicals react mostly by other routes rather than with PFe"'. The reaction of CHC1202and CHzC102radicals with PFe"' was not investigated in alkaline solutions since the (CH3)zCOH(Oz) radicals, which are also produced in these solutions, decompose to 02-:5leading to more complicated reactions. Acidic Solutions. The reactions of CCl&, CF3CHC10,, CHC1202,and CH2C102radicals were also investigated in acidic media. CC1302and CF3CHC102radicals which are quantitatively produced by irradiating solutions containing CCl, or CF3CHCIBr were found to react rapidly with PFe"'. The differential spectra recorded after the completion of these reactions are given in Figure 2. They present similar characteristics: absorbance buildup around 435,560,580-590, and 660 nm, and absorbance decrease at 500 and 615 nm indicative of PFe"' bleaching. Spectra showing the same features were obtained with acidic solutions containing CHCI, or CH,CI, (see Figure 2), but with about half the. intensity. In these cases, the possible reaction of (CH3)2COH(02)or its decomposition products with the porphyrin has to be considered. This was examined by adding acetone instead of the halogenated compound which leads to the quantitative formation of (CH&20H(02). The differential spectrum obtained in this case was of very low intensity. Thus, it was difficult to characterize the product and to determine whether (CH3),COH(02) or HO, (produced from it)45react with PFelI1. Nevertheless, the contribution of the related absorption changes (45)Y.Ilan, J. Rabani, and A. Henglein, J . Pbys. Cbem., 80, 1558-62 (1976).
I
I
I
500
600
700
I (
4
Figure 2. (a) Optical spectra of [DPDMEFe111(HOCH(CH,)2)2]t in acidic aqueous 2-propanol solutions (2-propanol: 6.5 M; the same spectra are obtained for HCIO, concentrations ranging from to 1.2 M for a porphyrin concentration of M). (b) Differential spectra recorded after irradiation of aerated solutions of [DPDMEFe'1*(HOCH(CH3)2)2]t (1 X lo4 M) in acidic aqueous 2-propanol (2-propanol: 6.5 M; HCIO,: 2.3 X M) containing 2% CCI4(0: experimental, -: smoothed) or 3% halothane (A, - - - )or 2% CHCI, (0,--). Spectra were recorded 170 ps or 1.2 ms or 400 ps after the pulse, respectively.
represents less than 20% of changes observed with solutions containing CHC13 or CH2Clz. Kinetic analyses of the reaction of these various peroxyl radicals with PFe"' were performed by following absorbance changes at various wavelengths and with porphyrin concentrations ranging between 5 X and 4 X lo-, M. Pseudo-first-order kinetics were obeyed in all cases, but the second-order rate constant were derived from experiments in which full scavenging of the radicals was obtained, Le., at the highest porphyrin concentrations. Best fitting of the data led to the second-order rate constants summarized in Table I. They vary between 6 X IO7 and 2.6 X lo8 M-' s - ~ for the various radicals. In certain experiments with CCl, and halothane, a slower process was found to develop leading to small additional absorbance increase after the main reaction. Its contribution was increased at higher acidities and lower porphyrin, concentrations. This slow process might be due to reaction of HOz radicals or to chain reactions. As mentioned above, H 0 2 formation becomes more important in acidic conditions. Also, from the known rate constants involved in this system it appears that competition between 2-propanol and PFe"' for the peroxyl radicals may occur at low porphyrin concentration. If the radical is scavenged by the alcohol, chain reactions develop leading to further formation of the halogenated peroxyl radicals.28 p H Effects. The difference between the spectra reported in Figures 1 and 2 cannot be ascribed solely to differences in the spectra of the parent porphyrins (cf. spectra shown in Figure l a and 2a) or to the slight dependence of the halogenated peroxyl radical yield on acid concentration. Thus, the observed effects should be due to structural changes in the product. These effects were further examined by irradiating solutions of various pH values (buffers were used at pH 4.7-9.1 in order to avoid side reactions of the porphyrins with protons produced by the solvent radiolysis).2* The differential spectra were found to vary considerably with acid concentration even when the spectra of the parent compound
The Journal of Physical Chemistry, Vol. 88, No. 13, 1984 2861
Reactions of Iron(II1) Porphyrins
r
- 1
I
I
I
I
L
10
c
I
I
500
600
1
I
4
Figure 3. Differential spectrum recorded 170 ps after pulse irradiation of aerated solutions of [DPDMEFelll(HOCH(CH,)z)z]+ (1 X lo4) in
acidic aqueous 2-propanol (2-propanol: 6.5 M; HC104: 1.16 M) containing 2% CCI4:.( experimental;-: smoothed). The spectrum obtained with solutions containing 2.3 X M HCIO, (see Figure 2) is recorded in a dashed line.
I
7
700
)I
(4
Figure 5. (a) Optical spectrum of DPDMEFe"'C1 in CCI4. (b) Differential spectrum recorded 3.6 ms after pulse irradiation of aerated solutions of DPDMEFe"'C1 (8.9 X loe5 M) in CC14.
1
~ 0
' 2
~ 4
~ 6
~ 8
* 10
~ 12
' 14
-log C H + l
Figure 4. Absorbance changes at 660 nm recorded 200 ps after pulse irradiation of DPDMEFe"' (1 X lo4 M) solutions in aqueous 2-propanol (6.5 M) containing 2% CC1, vs. acid concentration (see Experimental Section for details): ( 0 ) experimental, (-) best theoretical fits for one-protonexchange. The curve for the neutral-alkaline region is given on an expanded scale (right).
were the same. An example is given in Figure 3 for 1.16 M HC104 as compared with the one obtained at 2.3 X M HC104. It appears that the spectrum of the product does change with acid concentration with increase of the 660- and 590-nm signals and decrease of the 440- and 550-nm signals. Smaller spectral changes were also observed in the neutral region. A plot of the absorption changes at 660 nm vs. acid concentration is given in Figure 4. They suggest that two successive one-proton equilibria take place with pK, = 2.3 and 7.4, respectively. It should be pointed out that, since the activity coefficients are unknown in water-2-propanol mixtures, no activity correction was made. Also, the yield slightly depends on acid concentration. This might explain some discrepancies between the experimental data and the curves (calculated according to the where aA and equation a = (aA+ aBK1/[H'])/(l + K,/[H+]), aBare the absorbances of the acid and basic form, respectively, a is the absorbance at any pH, and Kl is the equilibrium constant). Solvent Effects. In order to elucidate the nature of the reaction between the peroxyl radicals and PFe"', we examined firstly the
effect of solvent and secondly another one-electron oxidizing system. Radiolysis of air-saturated neat 2-propanol containing CC14 (0.2 M) and HC104 (2.3 mM) basically leads to the CC1302radicals28 as in the aqueous systems. As was reported before,28 CC1302 radicals do not react efficiently with PFe"' in this solvent. Nevertheless, increasing the porphyrin concentration to 4 X M permitted us to estimate a rate constant of = lo7 M-' s-l. The differential spectrum obtained by irradiating 1 X lo4 M solutions of PFe"' was weak but had the same general characteristics as those shown in Figure 2. The reaction is not quantitative under decays by ~these conditions ~ 1 because * CC1302 l l J other routes such as reaction with 2-propanolZ8or radical-radical reactions. Experiments were also carried out in air-saturated neat carbon tetrachloride as a solvent. In this case radiolysis produces cationic species such as CC14+ and CC13+ (ref 46) as well as CC1302 radicals. Several compounds were found to be oxidized rapidly by the cationic species and more slowly by CC1302 radical^.^^,^* A similar situation prevailed with the ferric porphyrin. The rate constant for the reaction of the CCl3O2radical with PFe"' is given in Table I. The differential spectrum recorded after the completion of all oxidation steps is given in Figure 5. Discussion Several pathways for the reaction of peroxyl radicals with PFe"' must be considered: hydrogen abstraction, addition, and electron transfer. CC1,02 radicals are known to abstract labile H atom^.^^,^^ However, the porphyrins used do not seem to have sufficiently reactive hydrogens to give a predominant reaction. Addition to PFe"' has been shown to occur on the metal in the case of the methyl radical14 but other halogenated radicals were found to be u n r e a ~ t i v e . Addition ~ ~ ~ ' ~ on the porphyrin ligand at a meso position (46) 0. Brede, J. BOs, and R. Mehnert, Ber. Bunsenges. Phys. Chem., 84, 63-8 (1980); R. E. Biihler, Radiat. Phys. Chem., 21, 139-46 (1983); M. Washio, S. Tagawa, and Y. Tabata, ibid., 21, 239-43 (1983). (47) J. Grodkowski and P. Neta, J . Phys. Chem., 88, 1205-9 (1984). (48) E. P. L. Hunter and M. G. Simic, to be submitted for publication. (49) J. E. Packer, T. F. Slater, and R. L. Willson, Nature (London),278, 737-8 (1979).
2862
The Journal of Physical Chemistry, Vol. 88, No. 13, I984
was also reporteda5Q Finally, halogenated peroxyl radicals were shown to oxidize various c o m p o u r ~ d s . ~ ~ Our , ~ ~ results , ~ ~ , ~favor ~ the last pathway. Firstly, in a given solvent system, the product spectra do not depend on the nature of the peroxyl radicals. Secondly, the reactivity of the peroxyl radicals decreases in the following order: CC1302 > CHC1262 > CH2ClO2 > CF3CHClO2 >> CH362 indicating dependence on the expected ionization potentials of these species (according to the number and location of the halogens). The same trend was found for the oxidation of various organic compounds by a related series of halogenated peroxyl radicals.23 Thirdly, the rate constants of the reaction were found to decrease significantly with decrease in medium polarity. Thus, the observed main processes are attributed to the general reaction PFe"'
+ RO,
-
[PFe"']+
+ RO2-
(11)
One-electron oxidation products of ferric porphyrins have been prepared electrochemically53-55and ~ h e m i c a l l y . ~ ~ -Reversible ~' half-wave potentials were found at 1-1.2 V vs. SCE.53-55*58The actual electronic distribution has long been debated on the basis of UV-visible spectra,53magnetic s ~ s c e p t i b i l i t y , ~NMR,55v56 ~,~~-~~ IR55and Mossbauer ~ p e c t r a . ~ ~ - ~The " ~ most recent result^^^-^^ includig X-ray crystal structures57 are indicative of iron(II1) porphyrin r-cation radical complexes rather than iron(1V) porphyrins as first suggested.53 However, the electronic distribution might somewhat depend on the solvent and on the nature of ligands bound to the iron.59 As discussed below, in our solvent system, changing acid concentration provides a convenient way to modify the iron axial ligands. In 2-propanol or water/2-propanol mixtures, ferric deuteroporphyrins exist as a neutral form, PFe11'OCH(CH3)2HOCH(CH3)2,or as a charged form, [PFe111(HOCH(CH3)2)2]+, depending on acid concentration according to the equilibrium [PFe111(ROH)2]+
[PFer1'(RO-)(ROH)Io
+ H+
(12)
with pK12 4 as determined from spectroscopic studies.60 Very small featureless optical changes were also recorded in the range pH 11-12 hut it was difficult to assign them to the loss of a proton from PFe"'(RO-)(ROH). (50) C. E. Castro, C. Robertson, and H. Davis, Bioorg. Chem., 3, 343-60 (1974). (51) Zinc porphyrins were found to be oxidized by CClst)z radicals. (52) E. P. L. Hunter and M. G. Simic in 'Oxy Radicals and Their Scavenger Systems", Vol. I, G. Cohen and R. Greenwald, Eds., Elsevier, Amsterdam,-l983, p 32. (53) R. H. Felton, G. S. Owen, D. Dolphin, and J. Fajer, J. A m . Chem. Soc., 92,6332-4 (1971); R. H. Felton, G. S. Owen, D. Dolphin, A. Forman, D. C. Borg, and J. Fajer, Ann. N.Y. Acad. Sci., 206, 504-15 (1973). (54) A. Wolberg and .I. Manassen, J. A m . Chem. Soc., 92, 2982-91 (1970); J. H. Fuhrhop, K. M. Kadish, and D. G. Davis, ibid., 95, 5140-7 (1973). (55) M. A. Philippi 104, 6026-34 . _and M. M. Goff, J . A m . Chem. SOC., (198 2). (56) P. Gans, J. C. Marchon, C. A. Reed, and J. R. Regnard, Now. J. Chim., 5,203-4 (1981); W. F. Scholz, C. A. Reed, Y. J. Lee, W. R. Scheidt, and G. Lang, J . A m . Chem. Sac., 104, 6791-3 (1982). (57) G. Buisson, A. Deronzier, E. Dute, P. Gans, J. C. Marchon, and J. R. Regnard, J . A m . Chem. Soc., 104, 6793-6 (1982). (58) M. A. Philippi, E. T. Shimomura, and H. M. Goff, Inorg. Chem., 20, 1322-5 (1981). (59) D. R. English, D. N. Hendrickson, and K. S. Suslick, Inarg. Chem., 22, 367-8 (1983). (60) D. Brault, unpublished results.
Brault and Neta It can be seen from Table I that the rate for the reaction of the peroxyl radicals with PFe"' does not depend much on the nature of the ligand or on the acid concentration. In the same way, one-electron oxidation potentials of ferric porphyrins were found insensitive to the nature of anionic ligands bound to the iron.58 The oxidized species (PFe"')' is expected to give rise to the same equilibrium as PFe"': [(PFe1*r)+(ROH)2]2+s [(PFe"')+(RO-)(ROH>]+ H + (1 3)
+
The additional charge on the (PFelI1)+ species is likely to shift the pKa to lower values. Therefore, the experimental pKa of 2.3 is attributed to equilibrium 13. The observed pKa at 7.4 is difficult to assign to the further loss of a proton from [(PFer")+(R0-)(ROH)]' as no accurate pKa for the related equilibrium involving PFe"' is available. Alternatively, interactions of anions with the positively charged porphyrin ring might also be considered.61 Although discussion of electronic structure is beyond the scope of this paper, it can be pointed out that the protonation of the second 2-propoxo anion (equilibrium 13) results in drastic optical changes. This might reflect some electronic distribution modification caused by the nature of the iron axial ligands. The reactivity of peroxyl radicals toward ferric porphyrins demonstrated in the present study has important bearing on the toxicity of halogenated compounds. Indeed, it appeared that peroxidation of membrane lipids via radical-mediated chain reactions is not the only process leading to t o ~ i c i t y . ~ , ' ~For *~~>~~ instance, reactions of tissue constituents with other reactive molecules such as C0C12 and HOCl, formed through the oxidative metabolism of CCl,, are likely to be i n v ~ l v e d . 'These ~ molecules might be formed from CC1302radicals but the mechanisms of such reactions are unknown. From a stoichiometric p i n t of view, it must be pointed out that decomposition of CC1302to COCl2 and HOCl involves one additional electron and one proton. They can be supplied through H abstraction from lipids (the postulated initial step in lipid peroxidation). The reduction of CC1302by ferric cytochrome P450followed by protonation affords an alternative pathway which does not involve lipids. The contribution of such pathways to oxidative metabolism may depend on the nature of the halogenated compound. Finally, it should be mentioned that reduction of CC1302 may also lead to the hydroperoxide CC1302Hwhich will have a damaging effect, probably catalyzed by metal ions. Further studies on model systems are under way to estimate the biological relevance of these reactions. Acknowledgment. The research described herein was supported by the Office of Basic Energy Sciences of the U.S. Department of Energy. Registry No. CC1,Oz, 69884-58-8; CHCl,O),, 73761-31-6; CH,ClO,, 73761-32-7; CF3CHC102, 88099-83-6; DPFe"'C1, 21007-21-6; (DPDMEFe"'),O, 23208-98-2. (61) J. H. Fuhrhop and D. Mauzerall, J . A m . Chem. SOC.,90, 3875-6 (1968); D. Dolphin, R. H. Felton, D. C. Borg, and J. Fajer, ibid. 92, 743-5 ( 19 70). (62) D. Mansuy, M. Fontecave, and J. C. Chottard, Biochem. Biophys. Res. Commun., 95, 1536-42 (1980). (63) J. L. Stevens and M. W. Anders, Chem.-Bioi. Interact., 37, 207-17 (198 1). (64) Certain commerical equipment, instruments, or materials are identified in this paper in order to specify adequately the experimental procedures. Such identification does not imply recommendation or endorsement by the National Bureau of Standards, nor does it imply that the material or equipment identified are necessarily the best available for the purpose,
