Virtual Free Radical School Iron Chelation in Biology
Virtual Free Radical School Iron Chelation in Biology Alvin L. Crumbliss Department of Chemistry Duke University Box 90346 Durham, NC 27708-0346 Telephone: (919) 660-1540 Fax: (919) 660-1605 E-mail: [email protected] Website: http://www.chem.duke.edu/%7Ealc/labgroup/ Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 1 Iron Chelation in Biology Tutorial Guide Introduction: Biological Iron Coordination Chemistry Panels 3, 4 & 5 Chelation and Solubility Chelation and Redox Potential Panel 6 Panel 7 Common Iron Ligands in Biology Chelate Stability Definitions Panel 8 Panel 9 Chelation and Redox Control Oxidation State Influence on Chelate Stability Panels 10, 11 & 12 Panel 13 Iron Chelation and Transport Panels 14, 15 &16 Influence of pH on Chelate Stability Panel 17 Influence of Chelate Stability on E0 Influence of Chelation on Kinetics Panel 18 Panel 19 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 2
Introduction: Biological Iron Coordination Chemistry Iron is the second most abundant metal on the earths surface, falling closely behind aluminum and in near equivalent concentration to calcium and sodium. It is an essential element for virtually every living cell. The biochemistry of iron is controlled to a large extent by its coordination chemistry; i.e. the immediate chemical environment in the first coordination shell. This first coordination shell controls irons biological activity in small molecule storage (e.g.O2), electron transport, and catalysis. Common oxidation states: +2, +3 3 Common coordination numbers: 4, 5, 6 Iron in Biology Fe 1st coordination shell; immediate chemical environment Society For Free Radical Biology and Medicine Crumbliss 3 References [1,2] Introduction: Biological Iron Coordination Chemistry Examples of the extensive use of iron in biological systems, all of which are controlled or mediated by chelation, are as follows: redox chemistry involved in simple electron-transfer reactions; redox chemistry involved in reactions with O2, ranging from O2 transport and storage to O2 reduction by cytochrome oxidase, and O atom insertion catalyzed by cytochrome P450; and substrate activation by the electrophilic behavior of iron; for example, hydrolase enzymes such as purple acid phosphatase. 4 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 4 References [1,2] Introduction: Biological Iron Coordination Chemistry Fe 5 Iron in Biology
The first coordination shell Prevents hydrolysis/precipitation Influences molecular recognition Controls redox potential Controls mobility Society For Free Radical Biology and Medicine Crumbliss 5 Iron Chelation and Solubility Fe insoluble due to hydrolysis H2O 3+ OH2 H2O OH2 Fe H2O OH2 2+ OH2 H+ OH2 Fe OH2 H2O H+ L: OH2 OH H2O OH2 H O Fe H2O OH2 OH2 4+ OH2 Fe
O H OH2 OH2 Fe Higher insoluble polymers [Feaq3+]tot = 10-10 M @ pH 7 Strong chelators prevent hydrolysis and precipitation 6 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 6 Iron Chelation and Redox Potential L L L Fe L + 1.2 n+ L L Easy to reduce Eo volts Iron(II) stabilized Fe(phen)3 +0.8 +0.4 Fe(bipy)3 Fe(terpy)2 Fe(OH2)6 Iron(III) stabilized Fe(CN)6 HEME DERIVATIVES - Fe(salicylate) Fe(III/II) redox potential varies significantly with ligands in 1st coordination shell
0.0 Iron in Biology myoglobin Fe(EDTA) Fe(oxinate)3 -0.4 - 7 Fe(oxalate)3 hemoglobin hydroxamate siderophores Society For Free Radical Biology and Medicine e.g. Desferal Crumbliss 7 Common Iron Ligands in Biology Common iron ligand donor groups in biology include amino acid side chains, such as amine (I), carboxylate (II), imidazole (III), phenol (IV), and thiol (V). Other ligating groups include -hydroxy carboxylate (VI), catecholate (VII), hydroxamate (VIII) and porphyrin (IX). Fe Fe Fe Fe Fe O O S O (II) H2N Fe O (III) (I) (V) (IV)
N Fe N H (VIII) N N Fe N 8 (IX) N Iron in Biology O O Fe O O R O O N 1 (VI) R2 (VII) Iron(III) is a hard Lewis acid and prefers ligation to hard Lewis base donors (e.g. O, amine N) and iron(II) is a borderline soft Lewis acid and prefers ligation to soft Lewis base donors (e.g. S, pyrrole N). Society For Free Radical Biology and Medicine Crumbliss 8 Iron Chelate Stability Definitions Compilations of metal-ligand complex stabilities, such as that edited by Martell and Smith, use pH independent equilibrium constants, FeLH, as defined below for the reaction between Fe(III) and a hexadentate triprotic ligand, LH3, in aqueous solution. Fe(OH2)63+ 3- + L
FeL 110 = [FeL] 3+ 3[Fe(OH2)6 ][L ] However, in an in vivo or in vitro situation protons compete for the Fe(III) binding sites and the degree of complexation of the metal will be influenced by the ligand pK a values and the pH of the medium. + [FeL] [H ] 3+ + Fe(OH2)6 + H3L FeL + 3 H K = [Fe(OH2)63+][H3L] Since stability constants and K are determined as concentration quotients, their units differ on changing the denticity of the ligand. Consequently, 110 for a hexandentate ligand and 130 for a bidentate ligand cannot be directly compared. A pFe scale circumvents this problem and the problem of H+ competition due to different ligand pKa values. The pFe value for a particular ligand is the negative log of the free Fe(III) concentration at a fixed set of conditions: [total ligand] = 10 M, [total Fe(III)] = 1 M, and pH = 7.4. A high pFe value denotes a stable chelate complex. Panel 17 illustrates the influence of pH on Fe(III)-siderophore complex stability, using pFe values to express the stability of the complex. 9 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss References 9 [3,4,5,6] Iron Chelation and Redox Control Why is it important? A mechanism for preventing iron from participating in a catalytic cycle to produce toxic hydroxyl radicals and/or reactive oxygen species (ROS) (e.g. via the Fenton reaction or Haber Weiss cycle) is to control its redox potential by selective chelation. Through chelation, the redox potential for iron may be removed from the region where it can undergo redox cycling and produce hydroxyl radicals and ROS. This is illustrated in Panel 12. From the following thermochemical cycle, Equation (1) can be derived which relates the redox potential of an Fe complex to the chelators ability to discriminate between Fe(III) and Fe(II), as expressed by III and II. This relationship illustrates that the selectivity of a chelator for Fe(III) over Fe(II) increases with decreasing redox potential. III Fe(H2O)63+ + L
Fe3+L E0complex E0aq Fe(H2O)62+ + L 10 Iron in Biology E0aq E0complex = 59 log(III/II)  II Fe2+L Society For Free Radical Biology and Medicine Crumbliss 10 Reference  Iron Chelation and Redox Control Why is it important? From Equation (1) it is evident that the redox potential and stability of an iron complex are inter-related. These inter-relationships are important in characterizing the biological chemistry of iron because controlling the oxidation state of iron is a method of controlling both the thermodynamic and kinetic stability of a coordination compound. This is illustrated in Panel 13. As a result, the redox potential of a complex may be viewed as a measure of the sensitivity of a molecular level switch for changing the chemical environment of the iron (1st coordination shell). Data in Panel 13 show that for high spin complexes, changing the oxidation state of iron from +2 to +3 changes both the kinetic lability and thermodynamic stability of an iron chelate complex. 11 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 11  Reference Iron Chelation and Redox Control Prevent redox cycling & ROS production Fe(III) selectivity Control stability Control ligand exchange kinetics Control "switch" sensitivity . O2 -/H2O2 +770 Fe(H2O)63+/Fe(H2O)62+ +460
. H2O2/HO , HO- RH E (mV vs NHE) Why control E0? +940 HO ROS . HO- H2O2 .- -160 O2/O2 -320 NAD(P)+/NAD(P)H 100 Fe3+ Fe2+ . O2 - O2 1010 Haber-Weiss Cycle Eoaq - Eocomplex = 59 log( II) -480 -500 12 Iron in Biology Ferrioxamine B Transferrin Society For Free Radical Biology and Medicine 1020 Crumbliss 12 Reference 
Oxidation State Influence on Chelate Stability n+ L L FeIII Stable Inert Thermodynamics Illustration of the loss of several orders of magnitude of stability on reduction of high spin Fe(III) complex to Fe(II). Kinetics Illustration of an increase in 1st coordination shell lability on reduction of Fe(III) to Fe(II). 13 Iron in Biology L L +e L - e- L L (n-1)+ L FeII L L L L Less stable Labile Fe(III)transferrin Fe(II)transferrin log K @ pH 7.4 log K @ pH 7.4 Fe(III)ferrioxamine B Fe(II)ferrioxamine B
log 110 = 30.6 log 110 = 10.3 3+ Fe(OH2)6 +e - - L t1/2 = 4 ms + *OH2 = 20 = 3 3+ + OH2 2+ + OH2 Fe(OH2)5(*OH2) -e- 2+ Fe(OH2)6 + *OH2 t1/2 = 0.2 s Society For Free Radical Biology and Medicine Fe(OH2)5(*OH2) Crumbliss 13 References [7,8,9,10,11] Iron Chelation and Transport In humans, the host protein transferrin (Tf) is produced in excess of circulating free iron and sequesters extracellular iron at extremely high -20 affinity (Kd ~10 M). This chelation of iron prevents it from precipitation and also has a bacteriostatic effect by keeping iron as an essential nutrient from being available to bacterial pathogens. Human Transferrin Fe(III) Binding Site Nhistidine Otyrosine III
Fe aspartate O Otyrosine Fe3+ + apo-Tf + CO32Kb ~ 1020 M-1 O C O FeIIITf(CO32-) O 14 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 14 Reference  Iron Chelation and Transport Microbes solubilize environmental iron by a chelation process, whereby the microbe secretes chelators called siderophores which have a high and specific affinity for Fe(III). Siderophore mediated iron acquisition by microbes is illustrated here where the cell synthesizes and releases a polydentate siderophore (i) which solubilizes insoluble iron deposits by chelation (ii). The Fe(III) chelate diffuses back to the cell (iii) where it is recognized by a cell receptor (iv) and the iron is released into the metabolic processes within the cell (v). 15 Iron in Biology Fe Release microbial cell (i) (v) siderophore O O O O
O Molecular Recognition (iv) (iv) O (iii) O O O Fe O (iii) O O Fe3+ Ion Recognition (ii) Complexation Ksp ~ 10-39 Fe2O3.6H2O or Fe(OH)3 environmental iron Ca Al Mn Cr Cu Zn Ni Society For Free Radical Biology and Medicine Pb K Crumbliss 15 References [4,5,6] Iron Chelation and Transport Siderophores, microbially synthesized Fe(III) specific chelators, are low molecular weight molecules that usually incorporate bidentate catechol, hydroxamic acid, and/or -hydroxy carboxylic acid donor groups. These chelators exhibit high Fe(III) complex stabilities (high and pFe) to enhance delivery of iron to the cell, and large negative redox potentials (Panels 7, 12 and 18) for Fe(III) complexing specificity and to prevent redox cycling leading to the production of toxic hydroxyl radicals and ROS
(Panels 10, 11 and 12). Shown below are the structures of two hexadentate siderophores; enterobactin, a tris catecholate (in red), and ferrioxamine B, a tris hydroxamate (in blue). O O O O OH O OH O HO O N H O N O O O O O NH O Fe(III)-enterobactin complex 110 = 1049 ; pFe = 35.5 O O enterobactin 16 O O O NH O O OH HO
N H N N H N O O NH3+ O O N Fe O O O N H N O HO Iron in Biology NH3+ O O O O Fe O H N desferrioxamine B HO O O HN OH O N
H N H N OH ferrioxamine B complex 110 = 1030.2Crumbliss ; pFe = 16 26.6 Society For Free Radical Biology and Medicine References [4,5,6] Influence of pH on Fe(III)-Chelate Stability 40 NH NH O O N H 30 O OH NH2 OH O H N N O N N OH N H O NH2 NH pFe (I) 20 O H N
N HO O OH N O NH3+ O HO N H N O 10 (II) Plot of Fe(III) complex stability, expressed as pFe (Panel 9), as a function of pH for two siderophores, exochelin MN (I) and ferrioxamine B (II). Although they have approximately the same stability at pH 6.0, above this pH exochelin MN has a higher affinity for Fe(III) and below this value ferrioxamine B exhibits a higher affinity. This is due to different levels of competition from H + for the Fe(III) binding sites, due to different pKa values for the donor groups (shown in blue) in these two siderophore chelators. Exochelin MN (I) Ferrioxamine B (II) 0 0 17 2 4 Iron in Biology 6 pH 8 10 12 Society For Free Radical Biology and Medicine Crumbliss 17 [9,12] Reference
Influence of Fe(III)-Chelate Stability on E0 -E1/2 (mV) vs NHE 500 1 450 3 2 4 5 400 1. 2. 3. 4. 5. 6. 7. 8. 6 350 7 300 250 Fe(desferrioxamine B)+ Fe(Desferrioxamine E) Fe2(alcaligin)3 Fe(saccharide-trihydroxamate) Fe2(rhodotorulic acid)3 Fe(N-methylacetohydroxamate)3 Fe(acetohydroxamate)3 Fe(L-lysinehydroxamate)3 8 200 5 10 15 20 25 30 pFe Plot of the reversible Fe(III/II) redox potential (-E1/2) as a function of the stability of the complex, as expressed by pFe values (Panel 9). Data are for hexadentate (1,2,4), tetradentate (3,5) and bidentate (6,7,8) hydroxamic acid siderophores
and siderophore mimics. Note that: as the stability of the Fe(III)-complex increases, the complex becomes more difficult to reduce; and the stability of the complex decreases with decreasing denticity. 18 Iron in Biology Society For Free Radical Biology and Medicine Crumbliss References 18 [6,10,13,14] Influence of Fe(III)-Chelation on Kinetics H3C N Iron chelate formation places a strong electron donor in the H2O first coordination shell, which labilizes the remaining H2O aquated coordination sites. This is illustrated here for the reaction of hexa(aquo)iron(III) with N-methylacetohydroxamic acid, a siderophore mimic. Incorporation of the bidentate hydroxamate group in the first coordination shell labilizes the remaining aquo ligands by a factor of ~500 (k2/k1). OH2 Fe OH2 3+ OH H+ O H3C OH2 OH2 H2O O N CH3 HO k2 = 8.1 x 102 M-1s-1 O H3C
1+ H3C N O O OH2 Fe N CH3 H+ OH2 O O H3C Iron in Biology O H3C H3C 19 OH2 Fe H2O k1 = 1.8 M-1s-1 2+ OH2 Society For Free Radical Biology and Medicine N CH3 Crumbliss 19 [5,15] Reference References 1. 2. 3. 4. 5. 6. 7. 8. 9.
20 Crichton, R. (2001) Inorganic Biochemistry of Iron Metabolism, John Wiley & Sons, Ltd, New York. Harris, W. R. (2002) in Molecular and Cellular Iron Transport (Templeton, D. M., Ed.) pp 1-40, Marcel Dekker, Inc., New York. Martell, A. E. and Smith, R. M., Eds. (1974, 1975, 1976, 1977, 1982, 1989) Critical Stability Constants, Plenum Press, New York. Raymond, K. N. and Stintzi, A. (2002) in Molecular and Cellular Iron Transport (Templeton, D. M., Ed.) pp. 273-320, Marcel Dekker, New York. Albrecht-Gary, A.-M., and Crumbliss, A. L. (1998) in Metal Ions in Biological Systems Vol. 35, Iron Transport and Storage in Microoganisms, Plants and Animals (Sigel, A. and Sigel, H., Ed.) pp. 239-327, Marcel Dekker,New York. Crumbliss, A. L. and Boukhalfa, H. (2002) BioMetals 15, 325-339. Aisen, P. (1998) in Metal Ions in Biological Systems Vol. 35, Iron Transport and Storage in Microoganisms, Plants and Animals (Sigel, A. and Sigel, H., Ed.) pp. 585-632, Marcel Dekker,New York. Harris, W. R. (1986) J. Inorg. Biochem. 27, 41-52. Schwarzenbach, G., and Schwarzenbach, K. (1963) Helv. Chem. Acta 46, 1390-1400. Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 20 References 10. Spasojevic, I., Armstrong, S. K., Brickman, T. J., and Crumbliss, A. L. (1999) Inorg. Chem. 38, 449-454. 11. Helm, L. and Merbach, A. E. (1999) Coord. Chem. Rev. 187, 151-181. 12. Dhungana, S., Miller, M.J., Dong, L., Ratledge, C. and Crumbliss, A. L. (2002) manuscript in preparation. 13. Wirgau, J. I., Spasojevic, I., Boukhalfa, H., Batinic-Haberle, I., and Crumbliss, A. L. (2002) Inorg. Chem. 41, 1464-1473. 14. Dhungana, S., Heggemann, S., H., Gebhardt, P. Mllmann, U. and Crumbliss, A.L. (2002) Inorg. Chem. 41, submitted for publication. 15. Caudle, M. T., and Crumbliss, A. L. (1994) Inorg. Chem. 33, 4077-4085. Acknowledgements I thank my co-workers, some of whose names appear in the References, for their hard work, questions, ideas, and intellectual stimulation. Our work in this area is supported by the National Science Foundation, the National Instutites of Health, and the American Chemical Society Petroleum Research Fund. 21Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 21
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