Kinetic and Equilibrium Studies of the Reactions of Cysteine and

16 Kinetic and Equilibrium Studies Downloaded by PENNSYLVANIA STATE UNIV on September 27, 2013 | http://pubs.acs.org Publication Date: May 5, 1997 | doi: 10.1021/ba-1997-0253.ch016

of the Reactions of Cysteine and Penicillamine with Aqueous Iron (III) M.J.Sisley and R. B. Jordan* Department of Chemistry, University of Alberta, Edmonton, Alberta, Canada T6G2G2

The reactions of cysteine and penicillamine

(H L ) 3

with aqueous

+

iron(III) have been studied in dilute acid solutions (0.01-0.2 M H ) with +

iron(III) concentrations (0.01-0.1 M) in excess over the amino acid (~1.5-8.5 x 10 M) at 25 °C in 1.0 M

LiClO /HClO .

-4

4

Stopped-flow and

4

standard spectrophotometry were used to measure the kinetics of the complexation and oxidation-reduction

reactions, respectively.

The

absorbance immediately after the complexation reaction has been used to determine complexation equilibrium constants, defined as {[(H O) 2

n

FeLH ][H ] }/{[Fe(OH ) ][LH ]}

and

[H ] }/{[Fe (OH ) (µ-OH) ][LH ]},

with values of 0.063 and 0.38 M

2+

+

2

+

2

2

2

2

6

3+

+

4+

8

3

+

2

3

{[(H O) (µ-OH) Fe LH ] 2

m

2

3+

2

for penicillamine, and 0.025 and 0.07 M for cysteine, respectively. The dependence of the complexation rate on iron(III) and H

concentrations

+

indicates that the dominant reaction pathways and their rate constants (M

-1

+ H L : k = 1.4 x 10 (cys); 1.2 x

s ) are (1)for (H O) Fe(OH) -1

2

2+

5

2

2+

5

10

3

2

3

3

2

(pen); and (3) for Fe (OH ) (µ-OH) 2

+

3

+ H L: k = 7.4 x 10 (cys); 7.4 x 10

10 (pen); (2) for (H O) Fe(OH) 3

8

4+

2

3

+ H L: k = 4.25 x l0 (cys); 8.1 x 3

2

(pen). The oxidation of cysteine is second-order in cysteine, with

terms in the rate law first, second, and third order in Fe(OH ) . From 2

the [H ] dependence, these are assigned to reactions of +

H L, 2

H L, 2

(H O) FeLH 2

and

2+

4

+

(H O) FeLH 2

4

Fe (OH ) (µ-OH) (LH) 2

2

6

2

3+

2+

+

or

3+

6

(H O) FeLH

+

2

2+

4

+

Fe (OH ) (µ-OH) (LH) 2

2

(H O) FeLH . 2

4

2+

6

3+

2

Penicillamine

reacts similarly but more slowly, and the first pathway is not observed, but one fourth-order in iron(III) is found. With excess penicillamine the reaction is second-order in penicillamine and is inhibited by H

+

and

iron(II), and the reactive species is suggested to be Fe(LH) . +

2

*Corresponding author.

© 1997 American Chemical Society

In Electron Transfer Reactions; Isied, S.; Advances in Chemistry; American Chemical Society: Washington, DC, 1997.

267

268

E L E C T R O N TRANSFER REACTIONS

IVIERCAPTOCARBOXYLIC ACIDS OF T H E GENERAL FORMULA (HSRC0 H)

Downloaded by PENNSYLVANIA STATE UNIV on September 27, 2013 | http://pubs.acs.org Publication Date: May 5, 1997 | doi: 10.1021/ba-1997-0253.ch016

2

react with aqueous iron(III) to form blue complexes that subsequently undergo oxidation-reduction to the corresponding disulfide and iron(II). Equilibrium studies (1-4) have shown that two protons are released on complexation so that the - C O g and - S ~ groups are coordinated to iron(III) in the blue species. Cys teine (I) and penicillamine (II) also form blue complexes, and the observations of Stadtherr and Martin (5) show that - S ~ and - C O j are coordinated because S-methyl-L-cysteine and cysteine methyl and ethyl esters give no blue color. Furthermore, the amino group is not coordinated i n acidic solution because the 2V-acetyl derivatives of cysteine and penicillamine give the blue color, and penicillamine complexation (4) involves release of only two protons. Tomita et al. (6) characterized 1:1 complexes of cysteine and thioglycolic acid at -78 °C in 90% ethanohwater with absorption maxima at 620 nm (ε « 500-600 M cm ). They also isolated violet tris complexes ^ ~ 590 nm, ε ~ 3 χ ΙΟ M c m , for cysteine) and red complexes (X ~ 490 nm, ε ~ 1 χ 10 M " c m , for cysteine), which they assigned to FeiOHXSRCO^f-. The complexation kinetics of several H S R C 0 H systems have been stud ied by McAuley and co-workers (1-3) and by Baiocchi et al. (4). These studies involved different conditions, with iron(III) i n excess in the former and H S R C 0 H i n excess i n the latter, and the kinetic results are in reasonable agreement and follow the usual pattern for substitution on aqueous iron(III). Cysteine was not studied, but Baiocchi et al. did study penicillamine. More recently, Jameson and co-workers (7, 8) studied the reaction of aque ous iron(III) with excess cysteine at p H 2.7-4.87 and 8.5-11.68. These workers seem to have been unaware of the earlier studies of Stadtherr and Martin, Tomita et al., and Baiocchi et al., and in the lower p H range, Jameson et al. (7) assigned a blue species (À 614 nm, ε 1.03 χ 10 M cm ) to a mono-com plex of fully deprotonated cysteine with S- and N-coordination, while all previ ous work would indicate S- and O-coordination with an uncoordinated - N H J . At p H > 8, the species assigned as Fe(OH)(L) and Fe(OH)(L)f- by Jameson et al. appear to be Fe(OH)(L)f- and Fe(L)^ , respectively, from the observations of Tomita et al. (6). McAuley and co-workers (3) and Baiocchi et al. (4) studied the oxidation step for several H S R C 0 H systems, although neither studied cysteine, and McAuley et al. noted only that penicillamine oxidation is much slower. With excess reductant, Baiocchi et al. proposed formation of a steady-state amount - 1

- 1

-1

3

m a x

- 1

3

max

-1

2

2

3

max

- 1

-1

-

2

H C— CH-NH I I 2

3

+

(H C) Ç-CH-NH 3

2

3

+

3

HS

C0 -

HS

C0 " 2

2

I

II


1

16.

SISLEY AND JORDAN

Fe Reactions with Cysteine and Penicillamine

269

m

H OH H 0^ I I ^OH H0 H OH Η0

2

2

H

2»/, J ^°'' COâ + H +e~ = H S ^ COâ "0

2

2

2

E° = 1.72 V E° = -0.03 V

The E° for the last reaction (at 1.0 M H ) is consistent with measured val ues {10, 11) of about -0.25 V at p H ~ 7 for various species. It is apparent from the first two equations that oxidation to form a radical in these systems is highly unfavorable, although overall two-electron oxidation in the last reaction is possible with quite mild oxidizing agents. The E° for the first reaction can be combined with the known complexation constants {4) and E°(Fe /Fe ) = 0.75 V to estimate the driving force for the following intramolecular electron trans fer reaction. +

III

F e ( S - ^ COâ ) m

> Fe +

COg

11

ζ

11

This reaction has E° of -0.54 V {K = 9 χ 10" M) for cysteine and penicill amine. Clearly the intramolecular electron transfer is quite unfavorable thermodynamically. If the reverse of this reaction is assumed to have a diffu sion-controlled rate constant of ~ 1 0 M " s , then the upper h'mit for the for ward rate constant is F e + O2C^

F e ( S - v - COâ ) < m

S-ST^

11

2

CO2

which has E° of -0.79 V {K = 4 χ 10~ M) for cysteine and penicillamine. This calculation has assumed that the formation constant of the bis complex is onefifth that of the known value of the mono complex, but changes of a factor of 10 in this ratio just change the E° by ±0.06 V. O n the other hand, the following overall bimolecular reaction is quite favorable, 14

2 F e ( S ^ COâ )
2Fe + " 0

m

n

2

C ^ S - S - ^ COg

and has E° of 0.24 V (K = 1 χ 10 M) for cysteine and penicillamine. 8


16.

SISLEY A N D JORDAN

Fe Reactions with Cysteine and FeniciUamine m

271

The present work was undertaken to determine how the biologically important cysteine and penicillamine fit into the general picture from the mercaptocarboxylic acids. The conditions of [iron(III)] > [amino acid] and modest acidities (0.01 to 0.2 M H ) were chosen to determine i f the μ-dihydroxy dimer (IV) of iron(III) might show unusual reactivity. This species could bring two iron(III) and two - S - groups together in a manner somewhat analogous to the intermediate III proposed by McAuley et al. (3), and might impart unusual reactivity. Downloaded by PENNSYLVANIA STATE UNIV on September 27, 2013 | http://pubs.acs.org Publication Date: May 5, 1997 | doi: 10.1021/ba-1997-0253.ch016

+

Experimental Methods Materials. The L-cysteine hydrochloride monohydrate, L-cystine, penicil lamine hydrochloride, DL-penicillamine, and D-penicillaminedisulfide were used as received (Aldrich). Stock solutions of iron(III) perchlorate were prepared by dis solving primary standard grade iron wire (Allied Chemical) in excess 3.5 M HC10 , and oxidizing the iron(II) so produced with H 0 . The concentration of H C 1 0 in the final solution was determined by titrating the H released from Dowex 50 W - X 8 ( H ) cation resin and correcting for the H released by adsorption of iron(III). Solutions of iron(II) perchlorate were prepared by dissolving hydrated Fe(C10 ) (Alfa) in aqueous perchloric acid. The iron(II) content was determined spectrophotometrically as the 1,10-phenanthroline complex, and the solutions were found to contain > Κ/|ΧΗ ] for our conditions. This rate law becomes firstorder i n [ F e ] i f fc [Fe ] > > fc_g [ H ] , and then eq 9 simplifies to eq 10, where the minor second-order terms described i n the previous section have been added for completeness. n

n

t

+

aq

f

2

+

2

m

7

IH

+

rate = -,—^5[Fe(LH) ] " KHFe ] "•-5 2

n

+ s e

t

c o n d - order terms

,. (10) ΛΧ

f 1

[H ] +

Absorbance-time profiles calculated on the basis of this rate law are shown in Figure 2. In Figure 2A, no iron(II) has been added, and these runs show the penicillamine and [ H ] dependence. In Figure 2B, the effect of iron(II) is shown; the dashed curves (calculated with added [iron(II)] = 0) show the magnitude of the inhibition, and that the inhibition decreases at the higher acidity. A l l of the curves in Figure 2 have been calculated from eq 10 with k = 1.25 χ 10" and k^Kçlk^ = 3.5. These parameters provide a satisfac tory description of all the data. Models i n which the dominant reaction is F e ( L H ) + H L (&/') or F e ( L H ) + H L (fc/) are also first-order in iron(III) and second-order in penicillamine, but are not consistent with the [ H ] dependence. A specific rate constant that predicts the rate at 0.02 M H , predicts a rate that is 5 times too fast at 0.05 M H . This analysis leads to the upper limits for k" and fc/ given in Table I. It remains something of a mystery why these pathways, especially do not seem to contribute significantly. +

2

5

2+

2 +

2

+

3

+

+

+

Conclusions The equilibrium constants K and K are both smaller for cysteine than peni cillamine. This difference is typical of other metal ion systems (13) in which formation constants have been determined for both ligands. KQ for complexation by the iron(III) dimer is larger than K by 3 and 6 times for cysteine and penicillamine, respectively. For various α-mercaptocarboxylic acids, McAuley and co-workers (1-3) and Baiocchi et al. (4) are in reasonable agreement on the a

f â

a



16.

SISLEY A N D JORDAN

Fe

m

2000

281

Reactions with Cysteine and PeniciUamine

4000

6000

8000

10000

Time, s

Figure 2. Absorbance-time profiles for the oxidation of excess penicillamine by aqueous iron(III). For A, the concentrations of penicillamine, total Fe(III) and H , respectively, are 2.60 χ ΙΟ- , 2.01 χ 10^, 2.06 χ (o) 3.80 χ 10^, 2.01 χ 10^, +

3

;

2.11 χ 10- (n); 6.00 χ 10~ , 2.01 χ KH, 2.18 χ 10r (O); 6.15 χ ΙΟ" , 4.00 χ 1(H, 2

3

3

2

5.08 χ ΙΟ- (m). For clarity, the curves are offset by 0.20, 0.12, 0, and —0.04 absorbance units, respectively. For B, the concentrations of Fe(II), penicillamine, total Fe(III) and H , respectively, are 1.14 χ ΙΟ" , 2.60 χ ΙΟ- , 2.04 χ lOr , 2.07 χ 2

3

+

3

4

10- (o); 2.12 χ ΙΟ- , 2.41 χ 10~ ,1.99 χ i(H, 2.09 χ 10~ (μ); 4.00 χ ΙΟ- , 4.09 χ ΙΟ- , 2.02 χΙΟ- , 2.17 χ ΙΟ- (Ο); 3.97 χ ΙΟ- , 6.17 x1ο- , 4.03 χΚΗ, 5.13 χΙΟτ 2

3

3

4

3

2

2

3

3

3

2

(Μ). For clarity, the curves are offset by 0.05, 0.06, 0, and -0.08 absorbance units, respectively. The curves are cahuhtedfrom eq 10; dashed curves are cahuhted assuming [Fe(II)] = Ofor comparison to the (•) and (O) data in B.



282

values in the general range of 0.5-3 M . For penicillamine, our value of K = 0.068 M is in excellent agreement with that of Baiocchi et al.. These authors suggested that K is smaller for penicillamine because it forms a six-membered chelate ring, while the α-mercaptocarboxylates form five-membered rings. For the molar absorptivity coefficients of the iron(III)-a-mercaptocarboxylate complexes, McAuley and co-workers found values ranging from 140 to 1400 M c m , while Baiocchi et al. reported 700-1100 M " c m for the same ligands. For penicillamine, our value of 9.7 χ 10 is in only fair agreement with the 1200 M c m reported by Baiocchi et al., and we find a somewhat lower value of 7 χ 10 for cysteine. These values are somewhat difficult to determine because of the transient nature of the complexes. Previous work indicates that K involves coordination of - C 0 and - S ~ groups, and the same seems true for Κ because the same number of protons are released, and the iron(III) monomer and dimer complexes have similar absorbance maxima. The dimer complex could involve chelation at just one iron(III) center, but then one might expect K > Κ because of the higher charge per iron in the monomer. Alternatively, die amino acid could be bridg ing between the two iron centers in the dimer, in which case complexation of one end of the ligand would not adversely affect complexation of the other end at the other iron(III) center. This could explain the observation that Κ > K . The proton loss indicates that the μ-(ΟΗ-) bridge is not lost, unless it is con verted to an μ-(0) bridge. The latter seems unlikely under the acidic condi tions of this study. For the complexation reaction, the specific rate constants can be calcu lated from the composite values determined from the least-squares analysis that are summarized in Table I. The values are quite typical for substitution on aqueous iron(III) and in reasonable agreement with the earlier work of Baioc chi et al. (4) on penicillamine. There is the usual proton ambiguity between the paths involving F e ( O H ) + H L (fc ) and F e ( O H ) + H L (fc ). A choice between these terms is often based on the magnitude of the calculated rate constant compared to values for other ligands. In the present cases, this crite rion does not eliminate either pathway as a reasonable contributor. For the first pathway, the rate constants of 1.2-1.4 χ 10 M " s are high for reaction of a cation, and are more typical of neutral ligand complexation, but the positive charge on the - N H J group is somewhat removed from the reaction center. Baiocchi et al. (4) argued that assignment to k leads to unreasonably high and variable values for several mercaptocarboxylic acids and penicillamine. O n this basis assignment to k seems more appropriate. For the oxidation of cysteine, the dominant reaction pathways have a common feature in that they involve at least two iron(III) centers and two cys teines to reach the transition state. This appears to be a complicated way to proceed, but the disulfide product, cystine, requires that two electrons are released. This feature was also observed by Ellis et al. (I, 3) for the oxidation f l

f t

_ 1

-1

1

- 1

2


_ 1

- 1

2

a

2

β

f l

β

β

2

2

2+

3

+

2

2

3

1

3+

2

3

_1

3

2


f l

16.

SISLEY A N D JORDAN

Fe

m

283

Reactions with Cysteine and Penicillamine

of several mercaptocarboxylic acids by aqueous iron(III). In the latter work, reactions of the iron(III) dimer were not observed because of the high acidity, and the only pathway identified corresponds to k h i present terminology. The value of k = 6.4 χ 10 M s for cysteine is somewhat larger than those of 48-167 M " s- (25 °C, 1.0 M N a C l 0 ) found by Ellis et al. for different reductants. However, it is still unclear how the nature of the reductant affects the reactivity, and the magnitude of k seems reasonable. The alternative assignment of this kinetic term to k gives a five times larger rate constant and implies that k " > k . The latter order does not agree with the order < nor with the expectation that H L should be a better reducing agent than H L . From the earlier suggestion that the reaction of two F e ( L H ) units might react though a sulfide bridged structure (III), one might expect that the dimer complex F e ( O H ) ( L H ) would undergo facile reduction by cysteine as H L . In fact, the k " contribution in eq 7 is too small to be detected, and only an upper limit can be given (Table I). What is revealing is that k " (Fe (OH) ( L H ) + H L ) is smaller than fc ' ( F e ( O H ) ( L H ) + F e ( L H ) ) . This indi cates that it is more favorable to have both cysteines coordinated to iron(III) centers for oxidation to occur. The present results and the earlier work of McAuley and co-workers pro vide a general picture of the oxidation of mercaptocarboxylate systems under conditions of iron(III) in excess. The situation with a deficiency of iron(III) is not so clear. Baiocchi et al. studied several systems under these conditions and published details for mercaptosuccinic acid ( M S A H ) . These data indicate a first-order dependence on [MSAH ] and the proposed mechanism involves the bis complex (Fe(MSAH) ) as a steady-state intermediate. However, the analysis assumed that formation of the mono-complex (Fe(MSAH) ) is complete under all conditions, but this is certainly not the case from the formation constant also determined by Baiocchi et al. Our study with excess penicillamine clearly shows that the reaction is sec ond-order in penicillamine, that the bis-complex forms as an equilibrium species, and that the latter is the kinetically dominant reactant. The penicil lamine reaction also shows iron(II) inhibition, and this and the other kinetic features can be accounted for by the mechanism in Scheme II. It must be acknowledged that this mechanism has some puzzling features; why, for instance, does the bis-complex undergo intramolecular electron transfer, whereas the mono-complex seems to be unreactive? The driving force, although slight in either case, would seem to favor the mono-complex. The thiyl radical, once formed, couples with free penicillamine, rather than intramolecularly with the other penicillamine ligand in the bis-complex pre cursor. This could be rationalized if the S - atoms are trans to each other in the bis-complex and therefore not situated to couple efficiendy. Further studies on analogous systems may clarify these issues. t n e

2

2

2

1

_ 1

_1

1

4

2

2


2

2

2

3

+

2+

2

2

3+

2

2

2

3 +

2

3

2

2

3+

2

2+

3

3

2

+


2

284


Acknowledgments This work was supported by a grant from the Natural Sciences and Engineer ing Research Council of Canada.

References


1. Ellis, K. J.; McAuley, A. J. Chem. Soc. Dalton Trans. 1973, 1533 2. Lappin, A. G.; McAuley, A. J. Chem. Soc. Dalton Trans. 1975, 1560.

3. 4. 5. 6. 7.

Ellis, K. J.; Lappin, A. G.; McAuley,A.J.Chem. Soc. Dalton Trans. 1975, 1930. Baiocchi, C.; Mentasti, E.; Arselli, P. Trans. Met. Chem. 1983, 8, 40. Stadtherr, L. G.; Martin, R. B. Inorg. Chem. 1972, 11, 92. Tomita, Α.; Hirai, H.; Makishima, S. Inorg. Chem. 1968, 7, 760. Jameson, R. F.; Linert, W.; Tschinkowitz, Α.; Gutmann,V.J.Chem. Soc. Dalton Trans. 1988, 943.

8. Jameson, R. F.; Linert, W.; Tschinkowitz, A. J. Chem. Soc. Dalton Trans. 1988, 2109. 9. (a) Surdhar, P. S.; Armstrong, D. A. J. Phys. Chem. 1987, 91, 6532; (b) J. Phys. Chem.

1986, 90, 5915; (c) Mezyk, S. P.; Armstrong, D. A. Can. J. Chem. 1991, 69, 533. 10. Millis, Κ. K.; Weaver, K. H.; Rabenstein,D.L.J.Org. Chem. 1993, 58, 4144. 11. Lees, W. J.; Whitesides,G.M.J.Org. Chem. 1993, 58, 642. 12. Sisley, M . J.; Jordan, R. B. Inorg. Chem. 1991, 30, 2190.

13. Smith, R. M . ; Martell, A. E.; Motekaitis, R. J. NIST Critical Stability Constants of

Metal Complexes Database; U.S. Department of Commerce. National Institute of Standards and Technology: Washington,D.C.,1993. 14. Baes, C. F.; Messmer, R. E. The Hydrolysis of Cations; Wiley: New York, 1976. Mil -burn, R. M.; Vosburgh,W.C.J.Am. Chem. Soc. 1955, 77, 1352.


Kinetic and Equilibrium Studies of the Reactions of Cysteine and

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