sure measurements as well as to Barbara Kneebone who assisted with some of the temperature-dependence work.

LITERATURE CITED (1) G. J. Moody, R. B. Oke, and J. D. R. Thomas, Analyst, (London),95, 910 (1970!:v (2) J. Ruzicka, E. H. Hansen, and J. Chr. Tjell, Anal. Chim. Acta, 67, 155 (1973). (3) R. W. Cattrall and H. Freiser, Anal. Chem., 43, 1905 (1971). (4) H. J. James, G. D. Carmack, and H. Freiser, Anal. Chem., 44, 853 (1972). (5) G. H. Griffiths, G. J. Moody, and J. D. R . Thomas, Analyst(London), 97, 420 (1973). (6) J. E. W. Davies, G. J. Moody, and J. D. R. Thomas, Analyst (London), 97, 87 (1972). (7) J. E. W. Davies. W. M. Price, G. J. Moody, and J. 0. R. Thomas, Lab. Pract., 22, 20 (1973). (8) R . P. Buck, Anal. Chem., 46, 28R (1974). (9) J. Pick, K. Toth, E. Pungor. M. Vasak, and W. Simon, Anal. Chim. Acta, 64, 477 (1973). (IO) U.Fiedler and J. RuiiEka, Anal. Chlm. Acta, 67, 179 (1973). (11) E. Eyal and G. A. Rechnitz, Anal. Chem., 43, 1090 (1971). (12) G. Baum, M. Lynn, and F. B. Ward, Anal. Chim. Acta, 65, 385 (1973). (13) G. D. Carmack and H. Freiser, Anal. Chem., 45, 1975 (1973). (14) L. Heyne, in "Fast Ion Transport in Solids," W . van Gool, Ed., North Holland, Amsterdam, 1973, p 123.

(15) American Society for Testing and Materials, Philadelphia. Pa., Method D-1673-6 1. (16) L. G. Johnson and J. Chomicz, lnsulation(London), 14, 61 (1968). (17) R . E. Barker and A. H. Sharbaugh, J. Polym. Sci., Part C, 10, 139 (1965). (18) C. B. Monk, "Electrolytic Dissociation," Academic Press, New York, N.Y.. 1961. (19) A. E. Binks and A. Sharples, J. Polym. Sci., Part A, 6 , 407 (1968). (20) S. D. Hamann. Aust. J . Chem., 18, 1 (1965). (21) P. K. Datta. J . Sci. Ind. Res., 30, 222 (1971). (22) H. Sasabe. K. Sawamura, S. Saito, and K. Yoda, Polym. J., 2, 518 (1970). 123) S. Saito. H. Sasabe. T. Nakaiima. and K. Yada, J. Polym. Sci.. Part A, 6 , 1297 (1968). (24) F. Sandrolini. S. Pietra. and D. Manarisi, Chim. lnd. (Milan), 53, 755 (1971). (25) G. King and J. A. Medley, J. Colloid Sci., 8, 148 (1958). (26) J. H. Kallweit, J. Polym. Sci., 4, 337 (1966). (27) J. M. Davies, R. F. Miller. and W. F. Busse. J. Am. Chem. SOC.,63, 361 (1941). (28) L. E. Amborski, J. Polym. Scl., 62, 331 (1962).

RECEIVEDfor review April 10, 1975. Accepted July 10, 1975. This work was supported by a grant from the Office of Naval Research.

Electrochemical Characterization of Iron Porphyrin Complexes in Aprotic Solvents L. A. Constant and D. G. Davis Department of Chemistry, University of New Orleans, New Orleans, La. 70122

Fe( 111) porphyrin complexes of pyridine, substituted pyridines, and other complexing agents have been made in dimethylsulfoxide and N,N-dimethylacetamlde. The mechanisms of the redox reactlons found by cyclic voltammetry are reported. The E112 of a redox couple of Fe(lll) porphyrins depended on the pK, of the ligand. Ion palr formation was found In some cases, especially without added axial ligand.

Iron complexes of porphyrins play vital and diversified functions in biological systems. They are the prosthetic groups of such hemoproteins as hemoglobin, myoglobin, cytochromes, catalases, and some peroxidases. These hemoproteins are responsible for the transportation and storage of oxygen, the transfer of electrons in the "respiratory chain", hydrogen peroxide elimination, and possibly as an intermediate in the formation of high energy bonds in the mitochondria. In regard to these functions, many questions are still unanswered but much insight has been, and may be, gained by the study of model systems under conditions which approach the physiological state. Previous electrochemical studies with iron porphyrins have been performed in aqueous or aqueous-ethanol solutions (1-4). The aqueous solutions produce two problems in regard to such electrochemical work. First, there is a tendency for dimerization ( 5 ) to occur in aqueous solution without the added ethanol. Second, reduction of the porphyrin ring, especially to form the n-dianion, is followed by protonation. Use of nonaqueous solvents has eliminated many of these problems. The reductions of tetraphenylporphyrin and its metallo derivatives have been studied a t a dropping mercury electrode in dimethylformamide (6, 7) and dimethylsulfoxide (8). The oxidation-reduction poten-

tials in dichloromethane have been measured by cyclic voltammetry for tetraphenylporphyrin (9). Both the monomeric and dimeric iron(II1) forms of tetraphenylporphyrin and octaethylporphyrin have been studied (10, 11). Magnesium octaethylporphyrin and zinc tetraphenylporphyrin have been oxidized to the n-cation radical and the n-dication species in dichloromethane, methanol, and butyronitrile a t a platinum gauze anode (12). Oxidations of metalfree and metallo derivatives of tetraphenylporphyrin a t a platinum electrode in butyronitrile (13, 14) and benzonitrile (15) have been reported. Recently, in these laboratories (16, 17), redox characteristics of transition metal complexes of octaethylporphyrin have been studied by cyclic voltammetry a t a platinum electrode in dimethylsulfoxide and butyronitrile. The reduction of iron(II1) porphyrins has been reported in dimethylformamide without added ligand (18) and with added imidazole or histidine (19). But, except for the dimethylformamide work, only limited mention of the iron complexes may be found. In this work, model iron porphyrin systems were studied in aprotic solvents which are believed to approach more closely the protein sheath or lipid environment of living cell membranes than do aqueous systems. Proton magnetic resonance spectra have been reported for iron porphyrin and some axially coordinated ligand complexes in deuterated chloroform-methanol solutions (20). These results have shown that the species in a solution with added pyridine was the bis(pyridine)iron(III)porphyrin complex and that the exchange rate of free and complexed pyridine was slow a t temperatures below 260 K. Further, a linear relationship between the chemical shifts of the porphyrin protons and the basicity of the coordinated ligand was reported. The results also showed that the low-spin form of the complex was favored by greater basicity of the axial ligands. In the present work, this informa-

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

2253

Table I. Electrochemical Oxidation-Reduction Reactions"

-3

-B""""

+

100-

CY

-100-

-

V

-100 l 02

0

t

Z

1

02

,

l

I

I

-08

-0.3

E (volts vs

SCE)

Figure 1. Typical cyclic voltammograms (See text for explanation)

Table 11. Metal-Free Porphyrin Ring Reduction Half-Wave Potentials Potentials, Porphlrin

TPP

DMSO DMA

PrCN Pyridine

Proto

DMSO DMA'

PrCN" Pyridinea a

E (volts i s . SCE)

Solbent

-1.04 -1.04 -1.19 -1.13 -1.16 -1.17 -1.27 -1.24

-1.41 -1.52 -1.62 -1.49 -1.51 -1.59 -1.73 -1.59

Dimethyl ester of protoporphyrin.

tion has been utilized and the effects of such phenomena as ligand exchange and spin state of the central metal have been coordinated with the electrochemical characteristics of the various iron porphyrin complexes.

(9)

j [; J

Fe(1I)P

+ 2L

f

Fe(I1)P

EXPERIMENTAL

+ S + X-

a P = Porphyrin, X = Halogen, S = Solvent Molecule, L = Ligand, and 2 = Unknown (either solvent or unoccupied).

2254

Protoporphyrin and its iron(II1) chloride salt, hemin, were purchased from Sigma Chemical Company and Nutritional Biochemical Company, respectively, and were used as received. Tetraphenylporphyrin was prepared by the method of Adler et al. (21, 22). T h e resulting product was used without purification to prepare the tetraphenylporphyrin iron(II1) chloride ( 2 3 ) .Both the metal-free porphyrin and the iron-porphyrin were purified by column chromatography on neutral alumina using reagent grade dichloromethane as the eluting solvent. Initially, dimethylsulfoxide (DMSO) was treated with molecular sieves and then distilled a t reduced pressure. However, samples treated with molecular sieves alone yielded similar electrochemical background current; thus the distillation process was discontinued. N,N-Dimethylacetamide (DMA) was purified by vacuum distillation at 10-12 mm of ,mercury-the center 60% fraction being utilized. N,N-Dimethylformamide (DMF) was treated with BaO, filtered and distilled (twice) a t 10 to 15 mm of mercury with the middle 60% being kept each time. This solvent decomposed within two days even if stored under nitrogen and over molecular sieves. Butyronitrile (PrCN) was purified by the procedure of Van Duyne and Reilley (24),followed by distillation at 12 to 15 mm of mercury, retaining the middle 60% each time. The PrCN was stored under nitrogen and over alumina which had been activated at 350 OC for 3 hr. Tetrabutylammonium perchlorate, purchased from varied commercial sources, was vacuum dried over P2Oj and used without further purification as the supporting electrolyte for all

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

Table 111. Iron Porphyrin Half-Wave Potentials Potentials, E (\olta v ) . SCE)

Porphyrin

TPP

Solvent

DMSO DMA

PrCN

Pyridine CHZC1,

Proto

DMSO DMA DMF

Pyridine

Ring oxidations

...

... +1.38 ... +1.40 ... ... ... ...

the solutions. All solutions were made 0.1M with this electrolyte and between 0.5 and 1 m M with porphyrin or iron porphyrin. Most of the ligands were of reagent grade and used as received. T h e 4-aminopyridine was recrystallized from acetone, while pyridine, 4-methylpyridine, and piperidine were purified by distillation. All cyclic voltammet.ric measurements were made with a threeelectrode potentiostatic circuit constructed from operational amplifiers in these laboratories ( 2 5 ) and featuring positive feedback t o help minimize uncompensated resistance losses. T h e cell and electrode system have been previously reported (16). The working electrode, a small platinum disk (4.7 mm2), was cleaned, at least daily, by strong chemical oxidation followed by potentiostatic reduction a t or near -0.2 volt vs. SCE. T h e reference electrode (SCE) was placed in contact with the test solution through a lowflow saturated potassium chloride salt bridge and Luggin capillary placed as close to the working electrode as possible without generating oscillations-further minimizing resistance losses. Oxygen was removed from all solutions with prepurified nitrogen.

RESULTS AND DISCUSSION Typical voltammograms are shown in Figure 1 where they are labeled as types I, 11, and 111. The assignments for the voltammetric peaks will be made here for each type and later the various solutions will be matched with the corresponding type for discussion. The peak separation for the cathodic-anodic pair of a type I voltammogram is relatively small (near 60 mV at 0.2 V/sec); while the current function, ip,c/V1/2, and the current ratio, ip,a/ip,c, are constant with changing scan rate. For solutions without added ligand, the couple is assigned to the reversible redox system depicted in reaction 2 (Table I); while with added ligand, the couple is assigned to reactions 3 or 4 depending upon whether the potentials are dependent on the ligand concentration or independent of that concentration, respectively, thus becoming an EC case. The peak separation for the voltammetric peaks of type I1 is large (usually greater than 100 mV at 0.2 V/sec.); while the current function and current ratio are still constant with respect to scan rate. For a solution without added ligand, the cathodic peak is assigned to the reduction of reaction 1. A fast chemical reaction (shown in reaction 7 ) follows the reduction and the anodic peak is due to the oxidation of reaction 2, thus showing an ECE mechanism. In the presence of added ligand which is not strong enough to compete with the solvent for the axial positions of the iron(111) porphyrin, the cathodic peak may be assigned to reduction of reaction 1 for solvents where the ion pair exists or of reaction 2 when the ion pair does not exist. For all the ligands studied, added ligand complexed the iron( 11) porphyrin and, therefore, the reduction process is followed rapidly by reaction 9 or IO. Thus, the anodic peak is due to the oxidation of reaction 4. The type I11 voltammogram shows a cathodic peak, C1, which is at the potential expected for the solvated species and, therefore, the peak is assigned to reduction of reaction

...

... +1.17

...

+1.20

... *.. ... ...

Fe(III), Fe(I1)

-0.11 -0.11 -0.30 +0.17 -0.50 -0.17 -0.21 -0.26 +0.06

R i n g reduct.,nr

-1.17 -1.08 -1.09 -1.48

-1.68 -1.66 -1.72 -1.72

-1.30 -1.24 -1.20 -1.65

-1.80 -1.81

...

...

...

-1.84

1 or 2 as in type 11. The cathodic peak, C1, and the anodic peak, AI, form a reversible one-electron (confirmed by coulometric data) couple and are associated with the redox system of reaction 4. The absence of the anodic peak, Az, indicates the presence of a rapid chemical reaction as shown in reactions 9 or 10. The ratio of peak currents for C1 to C2 is scan rate dependent and decreases to some limiting value as the scan rate is increased. This may be explained by the presence of an equilibrium as depicted in reactions 11 and 12. When there exists a mixture of solvated and complexed iron(II1) porphyrin in the bulk of solution and the potential is scanned cathodically, the complexed species is reduced first, lowering its concentration in the area near the electrode. This causes a shift in the equilibrium to convert more of the solvated species to the complexed form (an CE mechanism). During the time that the potential is between C1 and the foot of Cz, the reduction of the complexed form continues but reduction of the solvated species does not occur. Once the potential reaches the foot of Cz, reduction of both species occurs. Thus, the amount of current observed for reduction of the complexed form (peak Cl) is dependent on its initial concentration in the bulk of solution and the amount formed by the equilibrium shift. In turn, the current of peak C2 will be decreased by the amount of the equilibrium shift from what its initial concentration would produce. At faster scan rates, the time in which the equilibrium shift will occur and the amount of complexed species thus produced decreases. Eventually, a scan rate will be achieved such that no equilibrium sXft has time to occur and a limiting ratio of currents for C1 to Cz will be observed. Thus, the presence of an equilibrium between solvated and complexed iron(111) porphyrin explains the data observed for a type I11 voltammogram. The actual ratio of the two iron(II1) forms in the bulk of the solution is equal to the limiting ratio of cathodic currents observed a t rapid scan rates. Porphyrin Ring Reactions. From previous works (6, 7, 9, 13, 25, 26), the metal-free porphyrins are expected to exhibit four one-electron redox reactions. The first and second ring reductions were observed and the half-wave potentials are listed in Table 11. The only evidence of any electrochemical oxidation of the metal-free porphyrin ring was an anodic peak at +1.12 volts vs. SCE for the dimethyl ester of protoporphyrin in butyronitrile. No corresponding cathodic peak was observed, nor was the anodic shoulder upon multiple scanning. In contrast, the iron porphyrins exhibit the expected five one-electron redox reactions, the half-wave potentials of which are listed in Table 111. Proposals have been made that the redox reactions assigned here as the first ring reduction and first ring oxidation should be assigned to reduction of iron(I1) to iron(1) (18, 27-29) and oxidation of iron(II1) to iron(1V) (IO),respectively. Thus, it must be emphasized that the assignments here and in a former work (17) may not be correct, but the

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

2255

IF-. I

/

I I

POTENTIAL (volt) Figure 2. Multiple scan cyclic voltammograms of iron(ll1)-iron(l1) couple for hemin in NJVdimethylforamide 1mM hemin, 0.1M tetrabutylammonium perchlorate. (A) Slow scan rates (10 V/sec)

reduction and oxidation potentials have values appropriate for ring reactions (17). The reductions for tetraphenylporphyrin are observed at potentials 0.11 f 0.01 volt more positive than the corresponding potentials of protoporphyrin. The absorption maxima in the visible spectrum of tetraphenylporphyrin in chloroform are a t longer wavelengths than in the corresponding spectrum of protoporphyrin dimethyl ester. Falk (29) has shown that for iron porphyrins and their axial complexes, the redox potentials are more positive and the visible absorption maxima are a t longer wavelengths with lower basicity of the porphyrin pyrrole nitrogens. Comparison of the half-wave potentials and visible absorption maxima for the DMSO solutions of the iron(II1) chloride of these porphyrins also shows these same trends. Therefore, it is confirmed that tetraphenylporphyrin is less basic than protoporphyrin. Iron(II1)-Iron(I1) Reactions. Cyclic voltammograms over the potential range of the iron(II1)-iron(I1) couple for the protoporphyrinin(II1) chloride in DMSO and DMA are shown in Figure 1 as type I and 11, respectively. The couple is quasireversible for both solvents and the major difference is the larger peak separation a t a given scan rate for the DMA solution as compared to the DMSO solution. In DMA, the situation may be made more complex by a following chemical reaction (reaction 7 ) . The same behavior was observed when tetraphenylporphyriniron(II1)chloride was studied. The conductivity study of Brown and Lantzke (30) has shown that the protoporphyriniron(II1) chloride dissociates into an iron(II1) porphyrin species and a chloride ion in DMSO, while in DMA and DMF it exists largely as the associated ion pair. The visible spectra of the solutions (even with added electrolyte) are in excellent agreement with those published (30). The porphyrin species in DMSO has solvent molecules in both axial positions and overall charge of +l;whereas, in DMA, it has the chloride ion and a solvent molecule as ligands and is electrically neutral. The structural difference of the porphyriniron(II1) chloride dissolved in DMSO and DMA is due not to the slightly higher dielectric constant but rather to the stronger complexing ability of DMSO. Thus, the larger peak separation for the voltammogram of the DMA solution is related to the ion pair existing in solution and the energy required for the 2256

*

separation of the chloride ion from the iron ion and subsequent solvent reorientation upon reduction of iron(111) to iron(I1) (reactions 1 and 7 ) . According to absolute rate theory, a larger activation energy would result in a slower electron transfer (31) and the larger peak separation indicates just this phenomenon. An electrolyte study in DMA using the perchlorate, fluoroborate, iodide, and bromide salts of tetrabutylammonia indicated that the chloride is the participant in the ion pair except in the presence of bromide. The cyclic voltammogram for hemin in DMA with bromide as electrolyte exhibits a smaller peak separation as compared to the voltammograms with the other electrolytes. Solutions which were electrolytically reduced from iron(II1) to iron(I1) and reoxidized to iron(II1) yielded voltammograms identical to those of the original solutions. The single scan cyclic voltammograms for the DMA solution containing perchlorate as electrolyte are of type 11. The multiple scans exhibited different behavior as shown in Figure 2. The slow scan rates ( l o V/sec.) show two cathodic peaks and a lone anodic peak; and the intermediate scan rates show all four peaks. A mechanism that is consistent with these data is suggested as follows. Initially the iron(II1) porphyrin-chloride ion pair is the solution species and the lone cathodic peak of the single scan, the more negative cathodic peak of multiple scans, is due to reduction of reaction 1. The existence of the equilibrium depicted in reaction 7 is indicated by the presence of the more negative anodic peak upon slow scanning but not upon faster scanning. Obviously, k f >> k , and the true equilibrium concentration of iron(I1) porphyrin-chloride ion pair is small. But upon slower scanning, the iron(I1) porphyrin-chloride pair is oxidized first and its concentration a t the electrode surface is lowered. The equilibrium shifts to produce more iron(I1) porphyrin-chloride ion pair and this adds more anodic current. The current for the more negative anodic peak increases a t the expense of the more positive anodic peak. Thus, oxidations of reactions 1 and 2 are responsible for the more negative and more positive anodic peaks, respectively. The more positive cathodic peak observed a t fast multiple scanning is assigned to the reduction of reaction 2, and the absence of this peak a t slower scan rates indicates the existence of reaction 8. One noteworthy point concerning this mechanism is that both chemical reactions 7 and 8 favor the electrically neutral form. Therefore, the mechanisms presented explain all the data; but, considering the complexity, may not be the only possible mechanisms. The cyclic voltammograms of protoporphyriniron(II1) chloride in DMF were very similar to those with DMA and the major difference between the two solutions is the relative stabilities of the solvents themselves. DMA remained suitable for use for two weeks after distillation if stored under nitrogen with molecular sieves, while even with these precautions DMF decomposed within two days. The decomposition was accompanied by shifts in the half-wave potentials observed for the iron porphyrins, presumably because of complexation with dimethylamine. Brown and Lantzke (30) had also observed rapid changes in the visible spectrum of hemin in DMF and slower changes in DMA. The cyclic voltammetric data for DMSO solutions of iron(II1) porphyrin-chloride are much less complex giving type I voltammograms for slow and fast, single and multiple scans. The electron transfer mechanism is illustrated by reaction 2. Iron(II1) Porphyrins with Pyridine as Axial Ligand. Addition of pyridine as an axial ligand eliminates the iron-

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

(111) porphyrin-chloride ion pair in DMA but still DMA and DMSO exhibit different voltammograms. The difference with added pyridine stems again from the stronger competitive nature of DMSO for axial positions as compared to DMA. The iron(II1)-iron(I1) redox couple of both protoporphyriniron(II1)-chloride and tetraphenylporphyriniron(II1)-chloride in DMSO with added pyridine appears as a single set of peaks and the general shapes of the voltammograms are similar to these without pyridine in DMSO (type I). The data for all voltammograms fit the criteria for a quasireversible couple. Plots of the reduction half-wave potential of the iron(II1)-iron(I1) couple vs. the logarithm of the pyridine concentration in DMSO show straight lines a t concentrations >lo0 mM with slopes of 110 mV and 135 mV for hemin and tetraphenylporphyriniron(II1)-chloride, respectively. The half-wave potential shifts anodically with increase in pyridine concentration and tends toward a limiting value as the concentration is lowered below 100 mM. The peak separation was relatively small (60-70 mV) a t a scan rate of 0.20 V/sec. at both low and high ligand concentrations, but increased to 120-140 mV a t intermediate concentrations (50-100 mM). The visible spectra of iron(II1) porphyrin solutions in DMSO containing as much as 2M pyridine are the same as without added pyridine. These data indicate that the iron(II1) is not complexed by pyridine and that two ligands are added upon reduction as shown in reaction 3. The first porphyrin ring reduction in DMSO containing pyridine is observed as a single set of peaks which meet the criteria for a quasireversible couple but the half-wave potential of this redox couple shifts cathodically with increasing pyridine concentration. The plots of half-wave potential vs. logarithm of pyridine concentration show the same tendency of lower slope at the lowest pyridine concentrations as the iron couple plots. The slopes of the linear portion a t higher ligand concentrations are -107 and -125 mV per log unit of pyridine concentration for hemin and tetraphenylporphyriniron(II1)-chloride, respectively. This signifies that the two axial pyridine ligands are lost upon reduction of the iron(I1) porphyrin to the iron(I1) porphyrin x-anion radical as shown in reaction 5. I t has not been ascertained whether the axial positions are vacant or occupied by solvent molecules, therefore the symbol Z has been used. The second ring reduction half-wave potential is independent of pyridine concentration. The results are consistent with the assignment of the peaks for this couple to reaction 6 with the same uncertainty as to occupation of the axial positions. The behavior of the iron porphyrins in DMA with added pyridine was observed to be distinctly different than its behavior in DMSO solutions. The major difference stems from the fact that the pyridine complex of the iron(II1) porphyrins exists ( a t least in equilibrium concentrations) in DMA whereas it did not in DMSO. Thus, pyridine has less complexing abilit,y for iron(II1) porphyrins than DMSO but is competitive with DMA. Again, both hemin and tetraphenylporphyriniron(II1)-chloride behave in the same manner and the following discussion applies equally to both. The DMA solutions with added pyridine produced type 111 cyclic voltammograms for the iron(II1)-iron(I1) couple of the iron porphyrins. The cathodic peak, Cs, has a potential corresponding to a DMA solution without added pyridine, while the quasireversible couple has more positive potentials as expected for a complex which is more stable for the lower oxidation state. The potentials do not show any dependence on the pyridine concentration but the ratios of cathodic peak currents do. Table IV shows the relationship

Table IV. Cathodic Current Ratio, Hemin-Pyridine-DM A Concintration P)ridine, .?A

icilicz

0.5 0.75 1.o 1.25

0.14 0.22 0.33 0.33 0.44

12.5 17.8 25.0 25.0 31 .O

1.o

49.4

1.5 2 .o

Complexed

between the pyridine concentration, the ratio of cathodic peak currents (limiting values from fast scan rates) and the percentage of iron(111) porphyrin complexed with pyridine as calculated from the limiting values of that ratio. The visible spectra of the iron(II1) protoporphyrin in DMA and in 1.5M pyridine-DMA are similar. The DMA solution shows maxima ( t shown after wavelengths) at 638 (5300), 540 (lOOOO), 509 (10500), and 398 (-1 nm; whereas the pyridine-DMA solution shows maxima a t 641 (4100), 540 (9300), 518 (9500), and 389 (91500) nm. The spectrum in pyridine as solvent shows maxima at 633 (1100), 556 (13000), 526 (13000), and 413 (-) nm. The cyclic voltammetric data have shown that approximately 31% of the iron(II1) porphyrin exists as the pyridine complex at a concentration of 1.5M pyridine in DMA. The spectra reflect this fact, but the changes in the visible spectra for these iron(II1) porphyrin solutions are more subtle than the changes observed by voltammetry. Therefore, it is concluded that the voltammetric data are better indicators of the degree of complexation than the iron(II1) visible spectra. However, the iron(I1) protoporphyrin visible spectra in DMA with 1 M pyridine and pyridine as solvent are very similar and have band maxima at 556 (20000),526 (17000), and 480 (12000) nm. The spectrum of iron(I1) protoporphyrin in DMA without pyridine is different with maxima a t 598 (6600), 570 (76001,and 555 (7800) nm. In pyridine as solvent both the iron(II1) and iron(I1) porphyrin are assumed to be the bis(pyridine) form. On the basis of the visible spectra of the iron(I1) complex in pyridine-DMA solutions, it is concluded that the bis(pyridine) complex is present in these solutions also. Since the peak potentials are independent of pyridine concentration, no pyridine exchange occurs upon electron transfer. Therefore, the iron(II1) porphyrin complex must also be the bis(pyridine1 form. Thus, the couple, Cl-AI, is assigned to reaction 4 and Ca to reduction of reaction 1. The absence of an anodic peak corresponding to oxidation of reaction 2 leads to the conclusion that rapid ligation occurs to form the bis(pyridine)iron(II) porphyrin complex. An attempt was made to determine the number of pyridine ligands from the percent of iron(II1) porphyrin complexed as a function of pyridine concentration. But the range of pyridine concentration was too limited (by the requirement that measurable currents for both species be observable) for any valid conclusions to be drawn. The first ring reduction potential of iron protoporphyrin shifted cathodically some 250 mV upon the addition of 2M pyridine to a DMA solution. The peak separation of the first ring reduction in DMA containing pyridine was much larger than in DMA without pyridine and increased as the pyridine concentration was lowered between 2M and 0.5M. This indicates that the pyridine ligands are lost upon reduction to the iron(I1) porphyrin x-anion radical just as in DMSO solution. The second ring reduction in DMA has the same peak potentials and the same peak separation with or without pyridine.

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

2257

2

a

I

-03;

2

4

PYRIDINE

I

I

I

6

8

10

I

I2

NITROGEN pK,

Figure 3. Dependence of cyclic voltammetric peak potentials for iron(ll1)-iron(ll) couple vs. ligand nitrogen pK, 1mM hemin, 1M ligand, 0.1M tetrabutylammonium perchlorate, Dimethylsulfoxide. Scan rate, 0.2 V/sec. (1) 4-Cyanopyridine, (2) 4-Acetylpyridine, (3) Pyridine, (4) 4-Phenylpyridine, (5) 4-Aminopyridine, (6) Pyrazole, (7) Benzimidazole, (8) 5.6-Dimethylbenzimidazole, (9) Imidazole

A brief study of tetraphenylporphyriniron(II1)-chloride in butyronitrile containing pyridine showed that the iron couple and both porphyrin ring reductions have the same behavior as in DMA. The objective of using butyronitrile was to utilize the anodic potential range of this solvent to study the effect of pyridine concentration on the porphyrin ring oxidations; but success was limited for two reasons. First, only equilibrium concentrations of the bis(pyridine)iron(III) porphyrin complex exists and these concentrations are small except with a larger concentration of pyridine, such as 0.5M to 2M. Second, the pyridine itself is oxidizable in the same potential region as the porphyrin ring. Because of the higher concentrations of pyridine necessary, this resulted in large background currents and coating of the electrode with products of the pyridine oxidation. Even the electrochemistry of iron(II1) porphyrin-chlorides in pyridine as solvent is not trivial. The iron(II1)iron(I1) potential region showed two sets of peaks. The current for C1 was about six times that for Cz and the current for A1 was about twelve times that for A2 and there was no variation of the ratios with scan rate. I t has been reported that one of the axial pyridine molecules is replaced by a water molecule in pyridine with as little as 3.9% water (32). Thus, the extra set of peaks may be due to formation of small amounts of this species produced by traces of water in the pyridine. Another possibility is the existence of some iron(II1)-chloride ion pair, the presence of which is not totally unreasonable in light of the relatively low dielectric constant of pyridine. Peaks for the first porphyrin ring reduction in pyridine are a t more negative values as compared to DMSO and DMA as solvents, and have larger peak separations; while the second ring reduction appears completely normal. This behavior of the ring reductions is completely in accordance with the trends observed for the mixed pyridine-DMSO and pyridine-DMA systems and shows the anion and the dianion are uncomplexed by pyridine. Iron(II1) Porphyrins with Other Nitrogen-Containing Ligands. A series of substituted pyridine compounds and compounds related to imidazole were used as ligands to study the effect of complexation on the electrochemical behavior. The series chosen spanned the range of pK, values for the ligand nitrogen from 1.9 to 11.3. These pK, values refer to aqueous solution data from the literature; while the comparisons concern solutions in DMSO and DMA, both of which have moderately high dielectric constants. Wooten and Hammett (33) have stated that a par2258

ticular acid will have dissociation constants that are determined primarily by the basic properties of the solvent. Therefore, the relative strengths of any two acids of the same charge type are approximately the same in all solvents of similar dielectric constants, as long as the dielectric constant is greater than 20. Thus, it is suggested that although the absolute values change, the relative order within the series holds constant. All substituents for the pyridine series were in the para position in order to minimize steric consideration. The majority of the work utilizing these ligands was performed with protoporphyriniron(II1)-chloride but tests with tetraphenylporphyriniron(II1)-chloride showed it to behave similarly. The cyclic voltammetric peak potentials of the iron(II1)iron(11) couple for the substituted pyridines and imidazole related series in DMSO are plotted in Figure 3 as a function of the pK, of the ligand nitrogen. All these data were ~ ~1~. ligand. The recorded at a Scan rate of 0.2 v / ~and peak separations using the more basic ligands are approximately 60 mV, indicating that the redox reactions of these complexes are diffusion controlled a t this scan rate. For solutions containing ligands of lower basicity, the peak separations are larger with the lower pK, of the ligand. This increased peak separation, indicating charge transfer control, is apparent below a pK, of 5 . The 4-cyanopyridine and pyrazole solutions exhibit such negative cathodic peak potentials because these peaks correspond to reduction of an iron(II1) porphyrin with solvent molecules rather than ligands in the axial positions. In DMSO solution, the species occupying the axial position of the iron(II1) porphyrin is related to the pK, of the ligand nitrogen. Cyanopyridine, the least basic of the substituted pyridine series, does not compete to any large extent with the solvent. For a 0.125M solution of cyanopyridine, the solvated species is evidently the only iron(II1) porphyrin species present and the observed cyclic voltammogram for all scan speeds appears as a type 11. At fast scan rates, a 2M solution exhibits the same type voltammogram. However, a t slow scan rates, the voltammograms begin to approach that of a type I11 by showing C1 as a shoulder on C2. The visible spectra of the iron(II1) state in 0.125M and 2M solutions are the same as without added ligand, while those of the iron(I1) state are similar to solutions with added pyridine except for an additional absorption maximum a t 624 nm. Increasing the ligand basicity, by using acetylpyridine, shows enhanced competition with DMSO for the axial positions. A 0.25M acetylpyridine solution produces a type I1 voltammogram a t all scan rates. But slow scan rates with a 1M solution yield voltammograms of type I; while faster scan rates yield a type I11 voltammogram. The visible spectrum of the iron(II1) porphyrin in a DMSO solution containing 0.25M 4-acetylpyridine is the same as without added ligand, while addition of 1M ligand produces an absorption maximum a t 557 nm as compared to 580 nm without ligand. The iron(I1) porphyrin spectrum is the same as with 4-cyanopyridine except the broad band has its maximum a t 640 nm. Thus, the data indicate that the equilibrium shown in reaction 11 lies farther to the right with 4acetylpyridine than with 4-cyanopyridine. The substituted pyridines of higher basicity, Le., 4-phenylpyridine and 4-aminopyridine, produced cyclic voltammograms of type I for high and low scan rates and ligand concentrations. The visible spectra of the iron(II1) porphyrin indicate also that complexation did occur with these ligands. Further, the spectra indicate that the bis(4-aminopyridine)iron(III) complex is in a low-spin state and that bis(4-phenylpyridine)iron(111) porphyrin complex probably is also. This agrees with a previous report that the low-spin

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

Table VI. Per Centa Iron(II1) Porphyrin Complex, Heminb-1M Ligand-DMA

Table V. Voltammetric P e a k Potentials,a Heminb-1M Ligand-DMA Ligand

4 -Cy anopyr idine 4 - Acet y lpyr idine Pyridine 4 -Phenylpyridine 4 - Methylpyridine 4 - Am inopy r id ine Pyrazole Imidazole 5,6-Dimethylbenziniidazole Piperidine a

C1

+0.12 +0.16 +0.05

+0.09 +0.03 +0.25 -0.09 -0.26 -0.27 -0.21

Potential vs. SCE at, scan rate of 0.20 V/sec.

c2

-0.35 -0.29 -0.32 -0.27 -0.31

.. ... ,

-0.29 , ,

., .,

4

A1

+0.19 +0.22 +0.09 +0.17 +0.10 -0.19 -0.01 -0.21 -0.15 -0.13

Concentration =

ImM.

form is favored by a greater basicity of the ligand (20). The half-wave potential of the 4-aminopyridine complex is more negative than the potential of the solvated species and indicates that the iron(II1) porphyrin complex is actually more stable than the iron(I1) porphyrin complex. This fact is in agreement with Falk’s report (29) that the bis(4aminopyridine)iron(II) porphyrin complex is an order of magnitude weaker than the corresponding bis(4-cyanopyridine) complex. If the equilibrium of reaction 11 does exist, it lies extremely far to the right with both 4-phenyl- and 4aminopyridine. The 4-aminopyridine contains two nitrogens both of which could be capable of coordination. Work with 4-dimethylaminopyridine, which should have strong steric interferences for the amino nitrogen, produced cyclic voltammetric data identical to 4-aminopyridine itself; whereas work with aniline as the added ligand showed no existence for any iron(II1) porphyrin complex. Thus, 4-aminopyridine coordinates through the pyridine nitrogen. Constant potential coulometry showed that solutions of iron( 111) protoporphyrin-chloride dissolved in DMSO, alone, with pyridine or with 4-aminopyridine contained essentially no iron(I1) form; whereas solutions with 4-phenylpyridine contained about 25% of the iron(I1) form. When an attempt was made to reoxidize the iron(I1) to iron(III), the coulometer continued well past the expected number of coulombs in the presence of 4-phenylpyridine but not with the other ligands. Thus, the data indicate that the 4-phenylpyridine caused autoreduction of the iron(I1I) porphyrin. A previous report ( 3 4 ) stated that this type of autoreduction occurred with pyridine but no such tendency was observed here. Addition of piperidine produced almost complete spontaneous reduction of iron(II1) to iron(I1) as had also been noted by others (34, 35) and to date only the bis(piperidine)iron(II) porphyrin has been isolated. Since it was not possible to electrolyze any large amount of the bis(piperidine)iron(II) porphyrin to the iron(II1) state, nothing definite can be said about the structure of the iron(II1) porphyrin species. However, the presence of the cathodic peak in the voltammogram of both DMSO and DMA solutions with added piperidine indicates the existence of an iron(II1) complex. The visible spectrum of the iron(I1) porphyrin complex is low-spin and consistent with the presence of two piperidine ligands. The small peak separation in both DMSO and DMA suggests further that both the iron(II1) and iron(I1) forms are the bis(piperidine) complex under these solution conditions. The facts that neither the current function nor the current ratio show any dependence on scan rate indicate that the chemical reaction is slow on the cyclic voltammetric time scale. The cyclic voltammetric peak potentials (Figure 3) for hemin solutions in DMSO with 1M of added imidazole,

Ligand

PK a

Complexcd

4-Cyanopyridine 1.91 Low 4-Acetylpyr idine 3.3 Low Pyridine 5.45 38 4 - Phenylpyr idine 5.55 48 4-Methylpyridine 6.88 60 4 - Aminopyridine 9.18 100 Pyrazole 2.48 37 Benzimidazole 5.53 100 6.95 100 Imidazole Based on limiting values of C,/(C, + Cz). Concentration = 1mM. benzimidazole, 5,6-dimethylbenzimidazole, or pyrazole showed the same trend as that established with the substituted pyridine series. The voltammograms all showed a single cathodic-anodic set for the iron(II1)-iron(I1) couple and the peak separation was large for pyrazole (the lowest pK,) but small for the others of the series. The spectral data also support the conclusion that the pyrazole complex with iron(II1) porphyrin does not exist in DMSO solution, while the bis(1igand) complex is the solution species when the other ligands are added. The substituted pyridine series in DMA solutions of protoporphyriniron(II1)-chloride yields cyclic voltammograms in accord with that observed for DMA solutions containing pyridine. The voltammograms are of type I11 with all the substituted pyridine series except 4-aminopyridine which exhibited a type I voltammogram. No simple correlation could be found between peak potentials, which are listed in Table V for 1 M ligand concentration, and ligand basicity as has been shown for the DMSO solutions. However, the potential difference between the two cathodic peaks was found to decrease with increases in basicity in the series from 4-cyanopyridine to 4-methylpyridine. The ratio of current for C1 to the total for C1 and CQwas used to determine the percentage of iron(II1) porphyrin existing as the substituted pyridine complex. This percentage increased with increase in basicity of the substituted pyridine as shown in Table VI. The assignments of electrochemical reactions for C1, CQ,and AQare the same as for the DMA solution with pyridine added. Further, it is assumed that when an axial ligand complex exists, it is the bis(1igand) form for both the iron(II1) and iron(I1) porphyrin. The addition of 1 M pyrazole to a DMA solution of 1mM protoporphyriniron(II1)-chloride yielded a type I11 voltammogram; whereas with imidazole and related compounds a type I voltammogram was obtained. The potentials with this series are listed in Table V and the percentages of iron(111) porphyrin complexed with 1M ligand are shown in Table VI. The general trends are the same as with the substituted pyridine series; but the imidazole series results are commensurate with higher pK, values in the pyridine series. CONCLUSION In conclusion, the study has shown that the electrochemical characteristics indicate various chemical equilibria which have direct bearing on the electron transfers from both the central metal ion and the porphyrin ring. Among these are ion-pair formation, solvent competition for axial positions, and complexation with ligands in the axial position. The formation of a complex has been shown to produce a change in spin state of the central metal ion with

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975

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ligands of high basicity and thus additionally affect electron transfer rates. I t has been shown (36) that complexation of imidazole, a biologically significant ligand, with an iron(II1) porphyrin greatly increases the electron transfer rate constant for the iron(II1)-iron(I1) redox system as compared to an iron(II1) porphyrin with solvent in the axial position.

LITERATURE CITED D. G. Davis and R . F. Martin, J. Am. Chem. Soc., 88, 1365 (1966). R. F. Martin and D. G. Davis, Biochemistry, 7, 3906 (1968). D. G. Davis and D. J. Orgeron, Anal. Chem., 38, 179 (1966). K. M. Kadish and J. Jordan, Anal. Lett., 3, 113 (1970). T. M. Bednarski and J. Jordan, J. Am. Chem. SOC., 86, 5690 (196p). D. W. Clark and N. S. Hush, J. Am. Chem. SOC.,87, 4238 (1965). G. Peychai-Heiiingand G. S. Wilson, Anal. Chem., 43, 550 (1971). (8)R . H. Felton and H. Linschitz, J. Am. Chem. SOC.,88, 4238 (1965). (9) N. E. Tokel, C. P. Keszthelyi, and A. J. Bard, J. Am. Chem. SOC.,94, 4872 (1972). (10) R. H. Felton, G. S. Owen, D. Dolphin, and J. Fajer, J. Am. Chem. SOC.. 93, 6332 (1971). (11) K. M. Kadish, G. Larson, D. Lexa, and M. Momenteaux, J. Am. Chem. SOC.,97, 282 (1975). (12) J. Fajer, D. C. Borg, A. Forman, D. Dolphin, and R. H. Felton, J. Am. Chem. SOC.,92, 3451 (1970). (13) A. Stanienda, Naturwissenschaften, 52, 105 (1965). (14) A. Stanienda and G. Biebl, 2.Phys. Chem., 52, 254 (1967). (15) A. Woiberg and J. Manassen, J. Am. Chem. SOC.,92, 2982 (1970). (16) K. M. Kadish and D. G. Davis, Ann. N. Y. Acad. Sci.. 208, 495 (1973). (17) J. H. Fuhrhop. K. M. Kadish, and D. G. Davis, J. Am. Chem. SOC.,95, 5140 (1973). (1) (2) (3) (4) (5) (6) (7)

(18) D. Lexa. M. Momenteau, and J. Mispelter, Biochim. Biophys. Acta, 338, 151 (1974). (19) D. Lexa, M. Momenteau, J. Mispelter. and J. Lhoste, Bioelectrochem. Bioenerg., 1, 108 (1974). (20) H. A. 0. Hill and K. G. Morallee, J. Am. Chem. SOC.,94, 731 (1972). (21) A. D. Adler. F. R. Longo, J. D. Finarelli. J. Goldmacher, J. Assour, and L. Korsakoff, J. Org. Chem., 32, 476 (1967). (22) A. D. Adler, L. Sklar, F. R. Longo, J. D. Finareiii. and M. C. Finarelli, J. Heterocycl. Chem., 5, 669 (1968). (23) A. D. Adier, F. R. Longo, F. Kampas, and J. Kim, J. lnorg. Nucl. Chem., 32, 2443 (1970). (24) R. P.Van Duyne and C. N. Reiiley, Anal. Chem., 44, 145 (1972). (25) J. L. Huntingtonand D. G. Davis, Chem. instrum., 2, 83 (1969). (26) R. H. Felton and H. Linschitz, J. Am. Chem. SOC.,88, 1113 (1966). (27) I. A. Cohen, D. Ostfeid, and B. Lichenstein, J. Am. Chem. SOC., 94, 4522 (1972). (28) G. J. Handschuh, "Spectral Investigations of Iron(1)-Etioporphyrin(1i) and Related Systems", Ph.D. Dissertation, Johns Hopkins University, Baltimore, Md., 1971. (29). J. E. Falk, "Porphyrins and Metalloporphyrins", Elsevier Publishing Company, New York, 1964. (30) S. B. Brown and I. R. Lantzke, Biochem. J., 115, 279 (1969). (31) S. Giasstone, K. J. Laidler, and J. Eyring, "The Theory of Rate Process", McGraw-Hill, New York, 1941. (32) H. A. Degani and D. Fiat, J. Am. Chem. SOC.,93, 4281 (1972). (33) L. A. Wooten and L. P. Hammett, J. Am. Chem. SOC.,57, 2289 (1935). (34) L. M. Epstein. D. K. Straub, and C. Maricondi, lnorg. Chem., 6, 1720 (1967). (35) L. J. Radonovich, A. Bloom, and J. L. Hoard, J. Am. Chem. SOC., 94, 2073 (1972). (36) L. A. Constant and D. G. Davis, in preparation.

RECEIVEDfor review March 28, 1975. Accepted August 6, 1975. The authors thank the National Science Foundation for financial assistance under Grant GP-42479X.

Electrochemistry of Cobalt Tetraphenylporphyrin in Aprotic

L. A. Truxillo and D. G. Davis Department of Chemistry, University of New Orleans, New Orleans, La 70722

The electrochemistry of cobalt tetraphenylporphyrln has been investigated in dimethylsulfoxide, N,Kdimethylacetamide, N,Kdimethylformanlde, n-butyronitrlle, and dichloromethane by cyclic voltammetry and controlled potential coulometry. In addition to the solvent studies, the number and stability constants of axial ligands were determined. Axial ligands included pyrldine, 4-picoline, and plperidine. Generally it was found that one or two axial ligands complex with Co(lll) porphyrin and that ligands were lost on reduction to Co(ll) or Co(l). Some equilibrium constants for complex formation were calculated from the electrochemical data.

This study of cobalt tetraphenylporphyrin (CoTPP) was undertaken because of the biological importance of cobalamins which have a porphyrin-like structure. Thus it was hoped that the materials investigated here would act as model compounds for cobalamins and especially vitamin BIZ.The use of nonaqueous solvents was selected to avoid adsorption problems as previously found ( 1 ) and to better mimic the in vivo environment in which cobalt compounds must function. CoTPP lends itself to the study of model biological systems since it can be both oxidized and reduced electrochemically. Vitamin B12 has been shown to undergo similar reactions (2-4). All electrochemical studies have been done in nonaqueous solvent in this work. 2260

Stanienda and Biebl ( 5 ) first reported some anodic Ellz's for CoTPP. They reported an irreversible oxidation of the metal before the oxidation of the ring in two one-electron steps. The two final steps produced the T cation and the dication (6). Kadish and Davis ( 7 ) investigated the oxidations and reductions and reported a heterogeneous (eleccm/sec. for the trochemical) rate constant of 5 X Co(II1) to Co(I1) reduction. Co(1I)TPP was oxidized in benzonitrile (8) and a reversible reaction a t +0.52 volts vs. SCE was found. Felton and Linschitz (9) recorded two reductions of CoTPP in DMSO, one a t -0.82 volt and the other a t -1.87 volt vs. SCE. The large difference between these as compared to other metal porphyrins (10) implies a difference in the nature of the reduction process and suggests that the site of primary reduction may be the metal rather than the ring in the first reduction step. A similar reduction was also reported ( 7 ) and was found to be the reduction of Co(1I)TPP to Co(I)TPP, the existence of cobalt(I) porphyrins having been discussed previously (2, 11). The reduction of Co(1I)TPP to Co(1)TPP has also been studied on mercury and platinum ( 2 ) in dimethylformamide. The purpose of this research was to apply electrochemical techniques to determine E1/2's and heterogeneous rate constants for CoTPP and its complexes with axial ligands.

EXPERIMENTAL Reagents. Cobalt(I1) tetraphenylporphyrin [Co(II)TPP] was initially prepared from tetraphenylporphyrin (obtained from Mad

ANALYTICAL CHEMISTRY, VOL. 47, NO. 13, NOVEMBER 1975