Electrochemistry of the Manganese(ll1)-(11) Hematoporphyrin IX Couple Donald G . Davis and Joseph G. Montalvo, Jr. Department of Chemistry, Louisiana S t a t e University in New Orleans, N e w Orleans,
The electrochemistry of the manganese(ll1)-manga nese( II) hematop0 r phy r in couple has been investitigated by means of polarography and cyclic voltammetry in aqueous and aqueous-ethanol electrolyte solutions in the pH range 1 to 13. The forms of the complexes undergoing electrode reaction at various pH's have been elucidated. Several chemical reactions coupled with the electron transfer process have been identified. Among these is the displacement of Mnz+ from Mn(ll) hematoporphyrin by hydrogen ion. The effect of adsorption on the electrode process has been investigated. Evidence opposing the possibility of intramolecular bonding between carboxylic acid side chains of the porphyrin ring and the central manganese ion is presented.
MANGANESE PORPHYRINS are of potential interest because they may possibly play a role in human porphyrin metabolism ( I ) and also in photosynthetic energy transfers (2). It has been determined that manganese is an essential element in quantasomes for the photosynthetic production of oxygen (3-5). Although no conclusive evidence has been brought forth regarding the exact nature of the compunds or complexes of manganese involved in photosynthesis, it has been suggested that manganese porphyrin would be a model worthy of study (6)particularly because a variety of oxidation states are available (Mn(II), Mn(III), and Mn(1V)). The solution chemistry of the various oxidation states of manganese porphyrins have been studied by means of spectrophotometry and potentiometry (6-8). Because it has been shown that electrochemical techniques such as polarography and cyclic voltammetry can elucidate chemical details of metalloporphyrin systems to which other methods are not sensitive (9-11), it was decided to apply these techniques to manganese porphyrin systems. This discussion deals specifically with Mn(II1) and Mn(I1) hematoporphyrin compounds in aqueous-ethanol solutions over the pH range 1-13. Experiments involving complexing agents other than OH- and H20 have been reported elsewhere (12). Mn(II1)Hm and Mn(II1)HmdiMe are used for the manganese(II1) hematoporphyrin IX complex and its dimethylester. Complexes of other oxidation states of manganese will be designated with the appropriate Roman numeral.
(1) D. C. Borg and G. C. Catzias, Nature, 182, 1677 (1958). (2) M. Calvin, Rev.Pure Appl. Chem., 15, l(1965). (3) R. B. Park and J. Biggins, Science, 144, 1009 (1964). (4) G . Englesma, A. Yamamoto, E. Markham, and M. Calvin, J . Phys. Chem., 66,2517 (1962). ( 5 ) E. Kessler, W. Arthur, and J. E. Brugger, Arch. Biochem Biophys., 71, 326 (1957). (6) P. A. Loach and M. Calvin, Biochem., 2, 361 (1963). (7) J. F. Taylor, J . Biol. Chem., 135, 569 (1940). (8) P. A. Loach and M. Calvin, Biochem. Biophys. Acta, 79, 379 (1 964). (9) D. G. Davis and D. J. Orgeron, ANAL.CHEM., 38, 179 (1966). (10) D. G. Davis and R. F. Martin, J . Amer. Chem. Soc., 88, 1365 ( 1966). (11) T.M. Bednarski and J. Jordan, ibid., 89, 1552 (1967). (12) D. G. Davis and J. G . Montalvo, Anal. Letters, 1, 641 (1968).
La. 70122
Hm and HmdiMe are used for the metal-free hematoporphyrin IX and its dimethylester. Because the porphyrin ring is essentially planar the four pyrrole nitrogen atoms occupy four coordination positions of the manganese in what might be considered the X-Y plane. Two more coordination positions remain, one above and one below the X-Y plane. They are designated as Z positions and are occupied by either a water molecule or an OH- in the solutions considered here. During the course of former studies (6, 7), it was observed that irreversible changes occurred in Mn(I1) Hm which increased in rate below pH 7.0. We have studied in detail this irreversible change and found it to be the acid displacement of the Mn(I1) ion from the porphyrin ring. The dissociation of a metal ion from the parent porphyrin may be looked upon as a special type of replacement reaction, the metal ion being replaced from the chelate by hydrogen ions as shown below. MP
+ 2H+,
ki k -1
PH2
+ MZf
In solvents containing appreciable amounts of water the rate of the reverse reaction is usually very slow or does not proceed at all. The acid displacement reaction has previously been studied only qualitatively, [except for the Mgz+ reaction investigated by Snellgrove and Plane (1311 the compounds falling broadly into three groups; those dissociated by water, by diluted hydrochloric acid, or by concentrated sulfuric acid. EXPERIMENTAL
Materials. PORPHYRINS. HmdiHCl (Nutritional Biochemical Corp.) was esterified in absolute methanol with dry hydrogen chloride by the method of Loach and Calvin (6), extracted into chloroform, and neutralized with an aqueous ammonia wash. The chloroform was removed by means of a rotary evaporator at 40 "C leaving the crude ester, HmdiMe. The crude product was purified by an elution column chromatography technique which was a modified version of that reported by Nicholas (14). Chromatographic grade magnesium oxide (Curtin Co.) was activated to the proper degree by first deactivating by exposure to the atmosphere for one week and then drying in an oven for two days at 37 "C. The MgO was slurried with chloroform and the 1-m by 5-cm glass column packed under positive pressure. About 2.5 grams of the crude ester were dissolved in chloroform, placed in the column, and the column eluted with a 50:l chloroform methanol mixture. Several bands were visible but the HmdiMe band was identified by infrared spectroscopy. The purified ester was dried at 40 "C with a rotary evaporator. Mn(II1)HmdiMe was prepared in glacial acetic acid according to Loach and Calvin (6). The course of the reaction was followed both by visible spectroscopy and the loss of fluorescence under ultraviolet light. The reaction was found to be 99+ % complete. (13) R. Snellgrove and R. A. Plane, J . Amer. Chem. Soc., 90, 3185 (1968). (14) R. E.H. Nicholas, Biochem. J . , 48,309 (1951). VOL. 41,
NO. 10,AUGUST 1969
1195
Mn(1II)Hm was prepared by the alkaline hydrolysis of Mn(I1I)diMe. Two and one-half grams of the Mn(II1) ester were dissolved in 200 ml of 0.5M potassium hydroxide in methanol followed by the addition of 5 ml of water. The mixture was stirred in the dark for 4 hours. The solution was then extracted with two 20-ml portions of chloroform to remove any unhydrolyzed ester, brought to pH 7 with HCl, and evaporated to dryness. The product was dissolved in 0.5M KOH and precipitated twice by adjusting the pH to 4.2 (the isoelectric point of Mn(II1)Hm). The precipitate was collected by centrifugation, washed three times with water, and dried in a vacuum desiccator over P205. This product, unlike that of Loach and Calvin (6) contained no chloride. Indeed the elemental analysis of our compound agreed more closely with the formulation for Mn(II1)Hm monohydrate-monohydroxide than the Mn(II1)Hm monohydrate with intramolecular bonding through a propionic acid chain of the other authors. The analysis, assuming the monohydrate-monohydroxide were : Found : C, 60.2; H, 5.5; N, 8.1; Mn, 8.1; and calcd: C, 59.6; H, 5.7; N, 8.1 ; Mn, 8.0. Compounds had identical visible spectra to those previously reported. REAGENTS.Reagent or comparable grade chemicals were used throughout this work without further purification. All solutions were prepared with distilled water and, if required, reagent grade absolute ethanol. Triply distilled mercury was used for dropping mercury electrodes (DME) and hanging mercury drop electrodes. All electrochemical experiments were carried out under a nitrogen atmosphere after deaerating the solution. Commercial tank nitrogen was purified by passing it through two vanadous sulfate scrubbers, two chromous sulfate scrubbers, and finally a portion of solution identical with that being studied to minimize evaporation of volative materials such as ethanol from the electrolytic cell. Apparatus and Methods. Absorption spectra were recorded on a Beckman Model B spectrophotometer. Polarograms were recorded with a Sargent Model XV polarograph or a Beckman Electroscan 30 (also used for chronopotentiometry) using a standard dropping mercury electrode and two electrode Sargent H-cell. Current time curves of individual mercury drops were recorded on a Tektronix 564 storage oscilloscope, using a tilted (45 ") capillary, the electrode being polarized by the Electroscan 30. Cyclic voltammetry was performed using a commercially available hanging mercury drop electrode of the Kemula type (Brinkman Instruments). Various apparatus configurations were used depending on the scan rate to be employed. For scans less than 0.25 V/sec the Electroscan 30 was used. For higher scan rates the electrode potential was controlled by a Wenking potentiostat to which a triangle wave was supplied either by an Exact wave function generator, type 250, or a triangle wave generator constructed in this laboratory. The storage oscilloscope was used as a recorder. The basic principles of cyclic voltammetry have been well covered in the literature (15, 16). The three electrode H-cell was of local construction. The auxiliary electrode was of platinum and was placed in a compartment separated from the test solution by a sintered glass disk. The saturated calomel electrode was brought into contact with the test solution by a Luggin capillary terminating within 1 mm of the working electrode. Controlled potential electrolysis, to establish that Mn(II1)Hm was reduced by one electron, to analyze solutions, and to prepare solutions of Mn(1I)Hm was performed with an Analytical Instruments potentiostat and current integrator. The control potential was selected by examining polarograms
(15) R. S. Nicholson and I. Shain, ANAL.CHEM., 36, 706 (1964). (16) R. S. Nicholson, ibid., 37, 1351 (1965). 1196
ANALYTICAL CHEMISTRY
of Mn(II1)Hm in the solution of choice. A mercury pool electrode was used. All solutions used for electrochemical studies were 1.OM in sodium nitrate. An appropriate buffer (0.2M) and ethanol were added as necessary. Test solutions of manganese porphyrins were prepared shortly before use by weighing the appropriate compound into a volumetric flask and adding previously prepared buffer4ectrolyte. All measurements requiring constant temperature were performed at 25 + 0.1 "C. All potentials are reported with reference to the saturated calomel electrode (SCE). The measurement of the ester hydrolysis rates for Mn(1II)HmdiMe were performed by dissolving a small amount of the ester in an appropriate electrolyte-buffer and extracting with chloroform at regular time intervals. The hydrolysis time was taken as the first time that color remained unextracted from the electrolyte-buffer. Mn(II1)Hm is completely insoluble in chloroform. It was found that the ester did not hydrolyze within a period of at least three hours if the pH was below 10. At pH 11.6, however, hydrolysis was complete in 15-20 minutes. A spectrophotometric procedure was developed to measure the rate of the displacement of Mn2f from the porphyrin ring for conditions under which the half-life was greater than 5 minutes. Electrochemical measurements were employed for half-life times between 5 and 95 seconds. When the spectral technique was used, excess dithionite was added to rapidly reduce Mn(I1I)P to Mn(I1)P. The amounts of dithionite used was increased with decreased pH due to the rapid decomposition of dithionite in acid medium. It was observed that dithionite slowly attacked the free porphyrin in acid solutions-even in the absence of light-resulting in a decrease in absorbance of spectral bands. This prevented following the reaction by monitoring the appearance of free P. An aliquot of the reaction mixture was taken at appropriate times during the course of the reaction and immediately pipetted with shaking into a solution buffered at pH 13 (which contained 45% ethanol and 1M sodium nitrate). Oxygen was introduced into the mixture to oxidize the remaining Mn(I1)P back to Mn(II1)P and to precipitate liberated Mn*+ ion as Mn02. The per cent transmittance was measured at 552 mp. This wavelength corresponds to a maximum in the absorbance of Mn(II1)P and a minimum in that of free P. Furthermore, at 552 mg the absorbance of the Mn(II1)P maximum was sufficiently greater than that of the free P minimum that the small decrease in absorbance of the latter by dithionite attack could be ignored within the time scale of the experimental runs. Because this allowed the total analytical concentration of Mn(II1)P and free P in the absorbance cell to be maintained constant, only the reading at 552 mp was taken. A calibration curve of per cent transmittance us. the concentration of Mn(1II)P in the absorbance cell gave directly the concentration of Mn(I1)P present in the reaction mixture during the course of the reaction. The electrochemical method used for following the displacement reaction was the current reversal chronopotentiometry technique (17). If, during the chronopotentiometric reduction of an oxidant to a soluble reductant, the constant current is reversed at or before the transition time for the oxidant, the reductant is reoxidized so that a reverse-transition time is recorded. Provided that the reverse-current density is equal in magnitude to the forward current density, the reverse-transition time n will be equal to one-third of the forward-electrolysis time t,. When the reversible charge transfer reduction is followed by an irreversible first order chemical reaction, the first order rate constant is obtained from a theoretical working plot of rl/taus. kt,. The experimental values of rt/taallow determination of the rate constant (17) A. C . Testa and W. H. Reinmuth, ibid., 32, 1512 (1960).
-030
B -0.40
w
0’ vj
9
-0.50
v)
3 0
’
-0.60
L
-IN
0.oc - 0.1
I
I
I
-0.3
-0.5
-0.7
W
-0.9
-0.70
E, VOLTS vs S.C. E. Figure 1. Polarogram of Mn(I1I)HmdiMe at pH 9.4, 0.80 mM Mn(III)HmdiMe, 47.5% ethanol
0
1
I
I
4 .O
7.0
10.0
1
13
PH k after dividing out t,. Forward transition times varied from 6 to 80 seconds in this work. A constant current of 6 pa was applied to the cell for most of the work. RESULTS AND DISCUSSION In order to elucidate the fundamental solution chemistry of Mn(II1)Hm and Mn(I1I)HmdiMe a series of polarograms were recorded at various pH’s, percentages of ethanol, and concentrations of manganese species. An example of such a polarogram is shown in Figure 1 . The smaller wave ( A ) is a prewave due to adsorption of the product of the electrode reaction (Mn(I1)HmdiMe) at the surfaceof thedropping mercury electrode. The larger wave ( B ) corresponds to the normal reduction of Mn(II1) species to Mn(I1) species. For the present the effect of pH, etc., on wave B will be of interest but the adsorption phenomena will be considered subsequently. Figure 2 depicts the effect of pH on the of waves corresponding to wave B and is analogous to the potentiometric data of Loach and Calvin (6) except that the gathering of more extensive data at low pH’s was possible due to the use of ethanolic solutions. Data such as in Figure 2 may be used to derive information about the species involved in the electrode reaction. As will be discussed, the electrode reaction is not always free of associated chemical reactions and adsorption problems but in alcoholic solutions the effect of these problems is minimized. Numerically the EliZ’s reported here are in excellent agreement with the mid-point potentials (E,) of Loach and Calvin (6). As may be seen’in Figure 2 the Eli2’s vary with pH in the range 1-4 and also between about 11 and 13. This variation indicates that hydrogen or hydroxide ions are involved in the electrode reaction at high and low pH’s but not at intermediate ones. The number of protons may be deduced because at both high and low pH’s the slope of the curves in Figure 2 were found to be: --
- -0.059V APH Because the electrode reaction involves one electron (as proved by controlled potential electrolysis) it is concluded that at low pH’s one hydrogen ion is consumed by the elec-
Figure 2. Variation of half-wave potential with pH 0 = Mn(I1I)Hm in water 0 = Mn(I1I)Hm in 47.5y0 ethanol A = Mn(II1)HmdiMe in 47.5 ethanol trode reaction. In the intermediate range, Hf or OH- are neither produced nor consumed and at high pH’s a hydroxide ion is produced. Other features of Figure 2 may also be used to deduce chemical information. At high pH’s in ethanolic solution the E112’sfor Mn(II1)Hm and Mn(II1)HmdiMe are the same. It was found that ester hydrolysis occurs in basic solution, the rate increasing with pH. For instance, at pH 11.6 in 47.5 ethanol the ester hydrolysis was essentially complete in 18 minutes. Thus by the time a polarogram was taken the original compound had been transformed to Mn(1II)Hm and no difference in Eli2 could be expected. In the pH range 1.5 to 10 hydrolysis required more than five hours and thus meaningful experiments could be performed with fresh solutions. At pH 4 and below the equality of Elj2’s may be ascribed to the fact that the carboxylic side chains are protonated, thus producing a structure similar to that of the ester. That protonation of the side chains takes place in the pH range 4-6 is confirmed by the fact that the isoelectric point of the compound falls in this range, as do the pK,’s for porphyrin carboxylic acids in general (18). At low pH’s some irreversible change took place in Mn(I1)Hm making studies difficult. We have extensively investigated this irreversible change and found it to be the displacement of the Mn2+ion from the porphyrin ring by one or more hydrogen ions. It might seem that this chemical reaction following the polarographic reduction of Mn(1II)Hm would render unmeaningful the Eli2data of Figure 2 at the lower pH’s. This, however, is not the case because of the relative slowness of the displacement reaction. This may be seen by considering the equation (19): (18) J. E. Falk, “Porphyrins and Metalloporphyrins,” Elsevier, New York, N. Y., 1964. (19) J. Heyrovsky and J. Kuta, “Principles of Polarography,” Academic Press, New York, N. Y., 1966. VOL. 41,NO. 10, AUGUST 1969
1197
Table I. Electrode Reactions of Mn(1II)Hrn and Mn(I1I)HmdiMe as a Function of pHa PH range 1.0-4.0
Electrode reaction HzO
Hz0
I Mn(II1)-P I
+ H+ + e-
I I
+ e-
I I
+ HzO + e-
=
I I
Mn(I1)-P-H
(1)
HzO Hz0 Transition region of mixed reaction 6.0-10.6 Hz0 H20 4.0-10.6
Mn(II1)-P
=
Mn(I1)-P
(2)
I
OH OH Transition region of mixed reactions Hz0 12.2-13.0 OH 10.6-12.2
Mn(II1)-P
=
I I
Mn(I1)-P
+ OH-
(3)
OH
OH
a The symbol -P is used for either hematoporphyrin or hematoporphyrin dimethylester. The charges on the species are neglected for simplicity. The symbol -P-H indicates a hydrogen ion bond to one of the pyrrole nitrogens of the porphyrin.
AEi/z
=
0.059 log 0.81 (kt)’”
(2)
where k is the pseudo first-order rate constant for the ejection reaction in sec-l and t is the drop time of the DME in seconds. At pH 2.5 we calculate that the value of Ellz would be influenced by no more than 12 mV indicating that useful results may be derived from the low pH region of Figure 2. On the other hand the half-life of Mn(I1)Hm at pH 2.5 is on the order of the drop time so that kinetic waves appear after the main wave (Figure 5). On the basis of the foregoing discussions we would propose the set of electrode reactions in Table I to explain the polarographic data. The consumption of hydrogen ion at low pH and production of hydroxyl ion at high pH are consistant with the data of Figure 2. The proposed reactions, however, are at variance with those previously proposed (6). It was suggested (6) that, in both Mn(II1) and Mn(I1) species, one of the 2 coordination positions perpendicular to the plane of the porphyrin ring is occupied by an intramolecularly bonded carboxylate group-Le., one of the propionic acid side chains of the porphyrin ring. A different approach is taken here for the following reasons: (1.) Both Mn(II1)Hm and Mn(I1)HmdiMe have very similar Ell2 E‘S. pH curves even though the dimethyl ester presumably could not undergo intramolecular bonding in either Mn(II1) or Mn(I1) complexes. The difference of E I ~ beZ tween Mn(II1)Hm and Mn(1II)HmdiMe at intermediate pH’s is explainable on the basis of the fact that the Mn(I1I)Hm has two more negative charges than the ester. (2.) If El/, values for the dipyridyl complex of Mn(1II)Hm are compared with those for the dipyridyl complex of Mn(II1)HmdiMe at pH9 (7, 12), a difference of 30 mV is found as compared with the 46 mV difference for Mn(II1)Hm and Mn(1II)HmdiMe in Figure 2 for alcoholic solutions and the 20 mV difference reported by Loach and Calvin (6),for aqueous solutions. These slight differences can reasonably be explained by the changes in dielectric constant in going from one solvent to the next. Greater differences would be expected, 1198
ANALYTICAL CHEMISTRY
however, if an intramolecularly bonded carboxylate group were displaced by a pyridine molecule in the Mn(II1)Hm but not in the corresponding ester. Coordination by two molecules of pyridine has been shown (7,12). (3.) The various formulations in Table I are consistant with the known chemistry of metalloporphyrins. (4.) Mn(II1)Hm has been found in these laboratories not to complex with added carboxcylic acids even when they are present in large concentrations. Thus intramolecular bonding of carboxylate groups seems less likely. Table I indicates that at low pH’s one hydrogen ion is consumed in the electrode reaction due to the addition of this hydrogen ion to one of the pyrrole nitrogen atoms. While somewhat unusual, this proposal is justified because at these pH’s the Mnzf is displaced rather rapidly by hydrogen ion. Indeed, similar species have been identified and been found stable under somewhat different conditions (20). Reactions 2 and 3 are also to be expected by analogy with other metalloporphyrins (10, 17). It would seem chemically unusual to have a Mn(1II)Hm species bound to a HzO molecule in one position and a carboxylate group at the other over the pH range 1 to 1 0 as is required by Loach and Calvin (6). At the lower pH’s hydrogen ion should effectively combine with even a bound carboxylate group as it does in EDTA complexes. On the other hand one would expect, as we propose, that Mn(II1)Hm should add an OH- not far from pH 6 and that in strong basic solutions two OH- ions would be coordinated with a $3 metal as is the case with iron porphyrins. Reversibility. In the pH range 5.5 to 13 in aqueous alcohol (or 5.5 to 10.2 in aqueous solution) polarograms exemplified by that shown in Figure 1 are obtained provided the concentration of Mn(II1) species was kept below 1.5 mM. At higher concentrations some maxima appeared and the direct relationship between diffusion, current, and concentration was lost. In addition, some evidence for dimer formation of Mn(I1) species was found (pH’s above 7) with high concentrations. Thus high concentrations were avoided. A careful analysis of the main polarographic wave (B in Figure 1 at pH 9.4) showed that a diffusion controlled process unencumbered by coupled chemical reactions was taking place. This conclusion is based on the following polarographic and other evidence. Measurements of the diffusion current, i d , showed that id/C was constant for variations of the concentration C, between 0.5 and 1.5 m M and throughout the pH range mentioned above. The ratio id/C did vary with the amount of ethanol in solution but the measured decrease in the ratio with increasing ethanol concentration is attributed to the increase in viscosity with ethanol concentration (21). The diffusion coefficient for Mn(II1)Hm was found to be 0.93 x 10-6 cmZ/sec in 1.OM borax (pH 9.0) and 3 0 x by volume ethanol. It is interesting to note that this is somewhat smaller than the 1.35 X 10-6 cmZ/sec found for Fe(II1) protoporphyrin in similar solution conditions (10). In addition it was found that iddh,where h is height of the mercury reservoir, was a constant, indicating diffusion rather than any type of kinetic control. Quantities comparable to the diffusion current for other faster electrochemical techniques were also examined. The chronopotentiometric quantity ir1Iz/C,where T is the transition time, was constant for 7’s between 8 and 65 sec and the stationary electrode peak cur(20) E. B. Fleischer and J. H. Wang, J. Amer. Clzem. Soc., 52, 3486 (1960). (21) 0. D. Shreve and E. Markham, ibid., 71,2993 (1949).
sible (and the rate constant independent of t ) a difference of 0.030 V should have been found.
For a typical reversibility reaction the current time equation i = kr“
I
II I
II
-0.500
I
I
I
-0.700
E,VOLTS vs S.C.E. Figure 3. Cyclic voltammetry peak potentials us. Mn(II1)Hm concentration (first cycle only) pH = 7.10. Scan rate = 0.050V/sec A . 1.500 mM Mn(1II)Hm B. 0.023 m M Mn(1II)Hm
rent, i,, was found to be proportional to the square root of the scan rate up to at least 27 V/sec. Likewise the ratio or cathodic peak current to anodic peak current was 0.91 5 0.7. All of the evidence indicates then that the electrode reaction is diffusion controlled with no contribution from couple chemical reactions in the pH range in question (15). That the reduction giving rise to the wave B in Figure 1 is a one electron change (Mn(III)Hm+Mn(II)Hm) was confirmed by controlled potential electrolysis and the spectra of the product was the same as that in the literature (6). A polarogram of the reduced solution showed that the diffusion current of Mn(I1)Hm was equal to that of Mn(I1I)Hm within experimental error. When log i / i d - i was plotted against potential for polarograms taken at various pH’s the expected straight lines were obtained (19). The slope of the plots depended on pH and on ethanol constant. In solutions which contained 47.5 % ethanol the slopes decreased from 68 mV at pH 6 to the reversible value of 59 mV at pH 12. When alcohol was not present the slope at pH 6 was 90 mV which decreased again to 59 mV at pH 12. These slopes indicate a certain amount of irreversibility especially at the lower pH’s in the absence of ethanol. A series of experiments were performed on an aqueous solution of Mn(I1I)Hm (1.5 mM) at pH 7. The pH and solvent were chosen such that the deviation from the theoretical slope would be large, and a high concentration was chosen to minimize the contribution of adsorption to the total current. It was found that the value of Eliz for such solutions was independent of concentration and of drop time whereas, if the reaction were irreversible, Eliz should vary (19). For instance, it was found that Elizwas -0.578 V with t = 2.30 sec and -0.575V with t = 7.10 sec. If the reaction were irrever-
(3)
has a constant exponent x theoretically having a value of l l 6 . Experimentally x is often found to be somewhat higher than having values between 0.18 and 0.23 for various ions and capillary characteristics (19). Here using a tilted capillary, x was found to have a constant value of 0.22 from 10% up the wave to the diffusion current. For an irreversible process x varies with potential. It is, however, not known how adsorption might effect x . Numerous cyclic voltammograms were recorded at scan rates between 0.2 V/sec and 20.0 V/sec. N o change in the difference between the anodic and’the cathodic peak potential was found with scan rate (AE, did vary with concentration due to adsorption). If an electrode process were a typical irreversible one variations in AE, would be expected (16). These pieces of evidence indicate that the electrode reaction is not irreversible in the conventional sense despite the fact that the slope of the log plot is greater than 59 mV. Possibly the adsorbed reactant tends to slightly block the electrode surface causing the polarographic slope to be greater than expected. Such a phenomenon has been deemed possible under the condition that the extent of blocking is a function of potential (22). Because of this problem of adsorption the data of Figure 2 may be somewhat in error but this error would not be so great as to negate the conclusions as to the participation of hydroxide ions in the electrode reaction. Adsorption. The pre-wave ( A in Figure 1) indicates that the product of the electrode reaction is quite strongly adsorbed under polarographic conditions. Some rough electrocapillary curves taken by measuring drop time as a function of potential also indicate that adsorption is occurring because these curves were suppressed as compared to curves taken with the solvent only (23). Cyclic voltammograms at low and high sweep rates are shown in Figures 3 and 4. In Figure 3 it may be seen that the difference in peak potential, AE,, depends on concentration (but does not vary with sweep rate) with the cathodic peak being displaced most when the concentration is increased. In addition there appears a just noticeable inflection in the 1.5 m M curve shortly before the peak potentials. This change in peak potentials and in shape with concentration indicates the participation of adsorbed reactant in the electrode process (24). The adsorbed reactant is reduced at potentials close to those at which dissolved reactant is reduced so no post peak or post wave is observed. At higher scan rates more information about adsorption phenomena may be obtained. Figure 4 (top) shows the initial and steady state cyclic voltammogram of Mn(II1)Hm at 100 V/sec. Beacuse of the high scan rate the sensitivity is increased and now on the first scan a pre-peak A is visible before the main peak B. This pre-peak is analogous to the pre-wave A in Figure 1 and indicates the reduction of Mn(II1)Hm to adsorbed Mn(I1)Hm. The peak B is sharper and more symmetrical than would be expected for a case without adsorption of reactant, Mn(II1)Hm (24). Thus evidence exists for (22) J. Kuta, “Modern Aspects of Polarography,” T. Kambara, Ed., Plenum Press, New York, N. Y., 1966, p 68. (23) M. Suzuki and P. Elving, CON. Czech. Chem. Comm., 25, 3202 (1960) (24) R. H. Wopschall and I. Shain, ANAL.CHEM., 39, 1514 (1967). VOL. 41, NO. 10,AUGUST 1969
1199
r\”
IL
CATHODIC
I
D
-
3.00
I
0
I
I
I
I
I
-0.50
i
2.25
-
I .50-
- 1.00
E, VOLTS vs S.C.E. Figure 4. Cyclic voltammograms of porphyrins in neutral water pH = 7.06 100 v/sec Current sensitivity = 86 pa/division Part I: 1.2 mM Mn(II1)Hm; solid line is initial cycle; dashed line represents 4 minutes cycling Part 11: 1.2 mM free Hm; solidline is initial cycle; C and C’ decay to small, constant peak heights after 4 minutes cycling the adsorption of both the reactant and the product in a competitive fashion. The peaks B‘ and A’ are the anodic analogies of B and A . On continued cycling for several minutes peaks A and A’ disappear and are replaced by a new set C and C‘. If hematoporphyrin with no manganese is used the cyclic voltammogram at the bottom of Figure 4 is obtained. Apparently repeated cycling causes the porphyrin ring, as opposed to the manganese, to gradually be reduced to a strongly adsorped material which displaces other adsorbed species and thus A and A’ go away as C and C’ appear. No attempt to identify the reaction causing C and C’ was made but the reduction of porphyrins in this potential range has been reported (25). To complete the adsorption picture consideration should be given to the effect of experimental variables on the size of the (25) R. H. Felton and H. Linschitz, J . Amer. Chem. Soc., 88, 1113 ( 1966).
Table 11. Surface Coverage of Adsorbed Mn(I1)Hm as a Function of pHa Em of Surface area adsorption waves, per adsorbed PH V us. SCE molecule, A 2 9.35 9.90
10.50 10.80 11.20 11.60
-0.173 -0.168 -0.155 -0.138 -0.122 -0.113
871 871 726 671 545 513
All solutions 1.2 m M in Mn(II1) Hm and 30% ethanol.
1200
ANALYTICAL CHEMISTRY
E, VOLTS vs. S.C.E. Figure 5. DME polarogram of MnHmdiMe pH 45% ethanol, 0.99 mMMnHmdiMe
=
4.65;
pre-wave, A (Figure 1). It was found that the height of the prewave at a given pH was independent of the total Mn(II1)Hm concentration and ethanol (0-30). The height of the pre-wave varied in direct proportion to the height of the mercury reservoir. This behavior is characteristic of systems in which the product of the reduction (in this case Mn(1I)Hm) is strongly adsorbed at the electrode surface. The height of the adsorption wave is proportional to the number of molecules adsorbed on the mercury drop during its life and the area per adsorbed molecule may be calculated from 13.66nm213t-’13
A =
ia
(4)
where i, is the limiting current of the adsorption wave, r is the drop time, n the number of electrons necessary for the reduction of each molecule, and m the rate of flow of mercury from the DME (26). Table I1 shows some measured values of A and of the Ellz of the pre-wave as a function of pH. From X-ray studies of metalloporphyrins (18) we estimate that if the Mn(1I)Hm were close packed and lying flat on the surface, the area per molecule would be 125A2. From the table it is seen that incomplete coverage is obtained, in contrast to iron porphyrin. In the iron case an essentially complete monolayer is measured (10) and variations with pH are not obtained. It appears likely that the increase in adsorption of Mn(I1)Hm with pH is due to a lessening of adsorption of other species such as Mn(II1)Hm-a possibility which is not inconsistant with the polarographic data. Coupled Chemical Reactions in Base. The height of the polarographic wave for the reduction of Mn(II1)Hm decreases as the pH increases above 10.2. If 3040% alcohol is present, the height of the wave remains constant. Cyclic (26) L. Meites, “Polarographic Technique,” 2nd ed., Interscience, New York, N. Y., 1966.
t
I
-
n
I
0 Q) v)
Y
24
CT
0 -
I
'a .8
-
I
I
I
I
I
I
I
I
I
I
I
PH Figure 6. pH Dependence of the pseudo first-order rate constant for the displacement reaction 45% ethanol T = 25 "C /I = 1.0 Shaded points represent data for MnHmdiMe; the other data is for MnHm. Chronopotentiometry data is given by the triangular points, remaining data obtained by the spectrophotometric technique. Dashed lines indicate extrapolation
voltammetric studies show that at constant concentration the peak current is not proportional to the scan rate and that the ratio of the anodic peak current to cathodic peak current increases above one as the scan rate increases. These observations are indicative of a reversible chemical reaction preceding the electrode reaction (15). Precisely what this reaction may be, has'not been determined. However, because the addition of ethanol eliminates it, the reaction may well involve the breaking up of dimers or other aggregates of Mn(II1)Hm before the electrode reaction. Ethanol has been shown to monomerize Fe(II1) porphyrin complexes (IO). In any case this phenomenon renders potentiometric measurements in basic solutions without added ethanol somewhat unreliable. Displacement of Mn2+. A polarogram of Mn(II1)HmdiMe at pH 4.65 is shown in Figure 5 . Wave B is for the reduction of Mn(III)/Mn(II)P couple. The heights of waves C and D increase as the pH becomes more acidic while that of B is constant. At pH's more basic than about 6.0 only wave B appears on the polarogram. Cyclic voltammetry studies of Mn(II1)Hm in pH 4.65 showed C and D appeared on the initial cycle of the cyclic voltammogram at low scan rates but not at high scan rates. At low scan rates, the ratio of the anodic peak current to the cathodic peak current was less than one for the Mn(III)/Mn(II)P couple and approached unity as the scan rate increased. This behavior is indicative of a reversible electron transfer followed by an irreversible
chemical reaction as determined from cyclic voltammetry diagnostic criteria (19). A polarogram of free Hm gave only two polarographic waves whose wave height and half-wave potentials showed a pH dependence analogous with waves C and D suggesting that these two waves are due to reduction of free Hm generated at the surface of the electrode by the irreversible chemical transformation of Mn(I1)P to free P. Waves C and D are due to a complex reaction possibly involving catalytic hydrogen ion reduction and are not simple diffusion controlled reactions for Hm or Mn(II1)Hm. The number of electrons involved is not known but it is large and thus C and D appear relatively big compared to B although they represent only a small fraction of the decomposition of Mn(I1)P. The pre-wave of Figure 1 is not found in Figure 5 because the adsorption of Mn(I1)Hm decreases as pH decreases. This does not mean, however, that Mn(II1)Hm adsorption decreases. In fact it appears to increase at lower pH's as shown by the variation of polarographic slopes. Controlled potential electrolysis of Mn(1II)Hm using a mercury pool cathode was performed at pH 6.20. It was found that one electron per Mn(II1)Hm was used. The spectrum of the final solution, however, was identical with that of free Hm rather than Mn(I1)Hm (measured at high pH). Thus it seems clear that although Mn(I1)Hm is originally produced by the reduction of Mn(II1)Hm that the Mn(1I)HmdecomposestoHmand Mn2+. A kinetic study of the decomposition of Mn(I1)Hm and Mn(1I)HmdiMe was undertaken. The spectrophotometric method was used for pH's above 5 and the chronopotentiometric method for lower pH's. With both methods the reaction was found to be first order in Mn(1I)P. Data for the chronopotentiometric method is shown in Table I11 for pH 3.79. It can be seen that the rate constant does not vary with concentration of Mn(II1)Hm which is further conformation that the reaction is first order (17). As the table indicates the error is on the order of 1 5 % for the chronopotentiometric measurements but it is on the order of 3 % for the spectrophotometric method.
Table 111. First Order Rate Constants from Chronopotentiometry with Current Reversal Mn(II1)Hm k X lo2 (mM) ta (sec) dta (sec-l) 1.000 1.88 0.215 5.6 2.00 0.218 5.5 2.22 0.198 5.7 2.36 0.195 5.4 1.250 6.32 0.272 3.9 6.48 0.278 3.4 9.58 0.238 4.2 10.34 0.240 3.8 12.48 0.215 4.2 13.56 0.210 4.1 1.500 3.70 0,292 4.6 5.57 0.249 6.1 5.67 0.281 3.7 7.32 0.239 5.4 9.44 0.242 4.0 9.48 0.228 4.8 13.54 0.200 4.6 15.54 0.182 5.0 Average 4 . 7 10.7
VOL. 41, NO. 10,AUGUST 1969
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The dependence of the rate of the reaction on pH is depicted in Figure 6. The equations given below indicate the order of hydrogen ion concentration participating in the reaction: r
=
k[Mn(II)P]
-log k = -log K
+ npH
were r is the rate, k the pseudo first-order rate constant, k the ratio of various protonation constants which might be involved, and n the order of reaction with respect to hydrogen ion. Therefore, the slope of a plot of -log k us. pH should in principle give the order with respect of hydrogen ion. Figure 6 indicates that the reaction appears to be first order in hydrogen ion at both high and low pH’s. In the intermediate range it appears that the order is greater than one. The dependence of the rate of the reaction on the hydrogen ion concentration indicates this reaction takes place through a displacement rather than a dissociation type of mechanism. Phillips, Dempsey, and Lowe (27) suggested that the reaction for the incorporation of zinc ion into solubilized porphyrin esters takes place through a displacement type reaction. Fleischer and Wang (20) detected and studied a reaction intermediate in the combination of metal ions with porphyrins. They showed that in acetone or chloroform solutions, Fe(II), Fe(III), Co(II), Pt(IV), Sn(II), Zn(II), and other metal ions first formed a new type of complex with absorption spectra markedly different from that of the metalloporphyrins. Upon prolonged standing at 60 “C the spectrum changes to the corresponding metalloporphyrins. On the other hand, upon addition of alcohol, pyridine, or water the spectrum changed immediately back to the free porphyrins (which probably explains our failure to detect any reaction type intermediate). The spectrum of the proposed complex was found to bear a striking resemblance to that of the free porphyrin diacid (PHd*+). It was concluded that the reaction between the metal ion and porphyrin to form the metalloporphyrin takes place through a displacement mechanism. The reaction intermediate proposed consisted of the metal ion sitting atop the porphyrin plane bonded to two diagonal nitrogens and with the other two nitrogens bonded to protons from the underside of the porphyrin plane. It is interesting to note that in the mechanisms for the acid displacement of Mn(1I)P given in this study, addition of the first proton to a pyrrole nitrogen resulted in a slow or rate determining step in the reaction, but with addition of the second proton the reaction was fast and the transition state complex immediately decomposed. This is in agreement with the stability of Fleischer and Wang’s reaction intermediate (20). (27) J. N. Phillips, B. Dempsey, and M. B. Lowe, “Symposium on Hematin Enzymes,” Cambena, Australia, Sept 1959.
1202
n
\i,‘
ANALYTICAL CHEMISTRY
Figure 7. Final transition state complex
Thus the final transition state complex leading to rapid decomposition of Mn(1I)P probably has two protons bonded to the pyrrole type nitrogens and has a sitting-atop structure similar to Fleischer and Wang’s model. This is shown in Figure 7. The results of this study indicate that manganese (11) porphyrin is unstable in an aqueous environment at physiological pH’s due to the displacement of MnZ+ion from tl-e porphyrin. Therefore, it seems unlikely that the Mn(I1) state can be important in porphyrin metabolism of cells in man and in the photosynthetic production of oxygen unless the porphyrin is stabilized. Strong electron attracting side chains on the porphyrin ring or a nonaqueous environment could provide this stabilization.
RECEIVED for review January 14, 1969. Accepted May 20, 1969. This investigation was supported by Public Health Service Grant AM-08248 from the National Institute of Arthritis and Metabolic Disease and by National Science Foundation Grant GP-8565.
