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  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 1

    Virtual Free Radical School

    Iron Chelation in BiologyIron Chelation in Biology

    Alvin L. CrumblissDepartment of Chemistry

    Duke UniversityBox 90346

    Durham, NC 27708-0346

    Telephone: (919) 660-1540Fax: (919) 660-1605

    E-mail: alvin.crumbliss@duke.eduWebsite: http://www.chem.duke.edu/%7Ealc/labgroup/

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 2

    Iron Chelation in BiologyIron Chelation in BiologyTutorial Guide

    Introduction: Biological Iron Coordination ChemistryPanels 3, 4 & 5

    Chelation and SolubilityPanel 6

    Chelation and Redox PotentialPanel 7

    Common Iron Ligands in BiologyPanel 8

    Chelate Stability DefinitionsPanel 9

    Chelation and Redox ControlPanels 10, 11 & 12

    Oxidation State Influence on Chelate StabilityPanel 13

    Iron Chelation and TransportPanels 14, 15 &16

    Influence of pH on Chelate StabilityPanel 17

    Influence of Chelate Stability on E0Panel 18

    Influence of Chelation on KineticsPanel 19

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 3

    Introduction: Introduction: Biological Iron Coordination ChemistryBiological Iron Coordination Chemistry Iron is the second most abundant metal on the earths surface, falling closelybehind aluminum and in near equivalent concentration to calcium andsodium. 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 coordinationshell. This first coordination shell controls irons biological activity in smallmolecule storage (e.g.O2), electron transport, and catalysis.

    Fe

    1st coordination shell;immediate chemical environment

    Common oxidation states: +2, +3

    Common coordination numbers: 4, 5, 6

    3 References [1,2]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 4

    Introduction: Introduction: Biological Iron Coordination ChemistryBiological 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.

    References [1,2]4

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 5

    Introduction: Introduction: Biological Iron Coordination ChemistryBiological Iron Coordination Chemistry

    The first coordination shellPrevents hydrolysis/precipitationInfluences molecular recognitionControls redox potentialControls mobility

    Fe

    5

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 6

    Iron Chelation and SolubilityIron Chelation and SolubilityFe insoluble due to hydrolysis

    FeH2O OH2

    OH2H2O

    OH2

    OH2

    FeH2O OH

    OH2H2O

    OH2

    OH2

    3+ 2+

    FeH2O O

    OH2OOH2

    OH2

    FeOH2

    OH2

    OH2

    OH2H

    H 4+

    H+

    H+

    Higher insoluble polymers

    Fe

    L:

    [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 7

    -0.4

    0.0

    0.4

    0.8

    1.2

    Fe(OH2)6

    Iron(II) stabilized

    Iron(III) stabilized

    Fe(terpy)2

    Fe(phen)3

    Fe(bipy)3

    Fe(salicylate)

    Fe(CN)6 -

    Fe(EDTA)Fe(oxinate)3

    HEMEDERIVATIVES

    hemoglobin

    myoglobin

    +

    +

    +

    -

    Easy to

    reduce

    Eovolts

    Fe(oxalate)3

    hydroxamate siderophores

    FeL L

    LLL

    L

    n+

    Iron Chelation and Redox PotentialIron Chelation and Redox Potential

    Fe(III/II) redox potential varies significantly with ligands in 1stcoordination shell

    7e.g. Desferal

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 8

    Common Iron Ligands in BiologyCommon Iron Ligands in Biology

    Iron(III) is a hard Lewis acid and prefers ligation to hardLewis base donors (e.g. O, amine N) and iron(II) is aborderline soft Lewis acid and prefers ligation to soft Lewisbase donors (e.g. S, pyrrole N).

    N

    N

    N

    N

    Fe

    Fe

    OO

    Fe

    OO

    R1

    NR2

    Fe

    OO

    O

    Fe

    O

    O

    Fe

    OFe

    H2N

    Fe

    N

    NH

    Fe

    S

    Common iron ligand donor groups in biology include amino acid side chains, suchas amine (I), carboxylate (II), imidazole (III), phenol (IV), and thiol (V). Otherligating groups include -hydroxy carboxylate (VI), catecholate (VII), hydroxamate(VIII) and porphyrin (IX).

    (I)

    (II)(III)

    (IV) (V) (VI)

    (VII)(VIII)

    (IX)8

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 9

    Iron Chelate Stability DefinitionsIron Chelate Stability DefinitionsCompilations of metal-ligand complex stabilities, such as that edited by Martell andSmith, 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+ + L3- FeL 110 = However, in an in vivo or in vitro situation protons compete for the Fe(III) binding sitesand the degree of complexation of the metal will be influenced by the ligand pKa values and the pH of the medium. Fe(OH2)63+ + H3L FeL + 3 H+ K = Since stability constants and K are determined as concentration quotients, their unitsdiffer 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 scalecircumvents this problem and the problem of H+ competition due to different ligand pKavalues. 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, usingpFe values to express the stability of the complex.

    [FeL][Fe(OH2)63+][L3-]

    [FeL] [H+] [Fe(OH2)63+][H3L]

    9 References [3,4,5,6]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 10

    Iron Chelation and RedoxIron Chelation and Redox Control Control Why is it important?Why is it important?A mechanism for preventing iron from participating in a catalytic cycle to producetoxic 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 selectivechelation. Through chelation, the redox potential for iron may be removed fromthe region where it can undergo redox cycling and produce hydroxyl radicals andROS. This is illustrated in Panel 12. From the following thermochemical cycle, Equation (1) can be derived whichrelates the redox potential of an Fe complex to the chelators ability to discriminatebetween Fe(III) and Fe(II), as expressed by III and II. This relationship illustratesthat the selectivity of a chelator for Fe(III) over Fe(II) increases with decreasingredox potential. Fe(H2O)63+ + L Fe3+L

    Fe(H2O)62+ + L Fe2+L

    III

    IIE0aq E

    0complex E0aq E0complex = 59 log(III/II) [1]

    10 Reference [6]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 11

    Iron Chelation and RedoxIron Chelation and Redox ControlControlWhy is it important?

    From Equation (1) it is evident that the redox potential and stability of an ironcomplex are inter-related. These inter-relationships are important in characterizing the biologicalchemistry of iron because controlling the oxidation state of iron is a method ofcontrolling both the thermodynamic and kinetic stability of a coordinationcompound. This is illustrated in Panel 13. As a result, the redox potential of acomplex may be viewed as a measure of the sensitivity of a molecular levelswitch for changing the chemical environment of the iron (1st coordination shell). Data in Panel 13 show that for high spin complexes, changing theoxidation state of iron from +2 to +3 changes both the kinetic lability andthermodynamic stability of an iron chelate complex.

    Why is it important?

    11 Reference [6]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 12

    Iron Chelation and RedoxIron Chelation and Redox ControlControl

    -480

    -320E

    (m

    V v

    s NH

    E)

    NAD(P)+/NAD(P)H

    -160 O2/O2. -

    +940 O2. -/H2O2

    +460 H2O2/HO., HO-

    Fe3+

    Fe2+ O2

    O2. -

    H2O2

    HO.

    HO-

    ROSRH

    Haber-Weiss Cycle

    Eoaq - Eocomplex = 59 log(

    /II)

    Fe(H2O)63+/Fe(H2O)6

    2++770

    1020

    1010

    100

    /

    -500Ferrioxamine BTransferrin

    Why control EWhy control E00??

    Prevent redox cycling & ROS productionFe(III) selectivityControl stabilityControl ligand exchange kineticsControl "switch" sensitivity

    12Reference [6]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 13

    FeIIIL L

    LLL

    L L L

    LLL

    L

    L

    n+ (n-1)+

    + e-

    - e-FeII

    StableInert

    Less stableLabile

    Fe(III)transferrin log K @ pH 7.4 = 20 Fe(II)transferrin log K @ pH 7.4 = 3 Fe(III)ferrioxamine B log 110 = 30.6 Fe(II)ferrioxamine B log 110 = 10.3

    ThermodynamicsThermodynamicsIllustration of the loss of several orders of magnitude of stability on reduction of high spin Fe(III) complex to Fe(II).

    KineticsKineticsIllustration of an increase in 1st coordination shell lability on reduction of Fe(III) to Fe(II).

    Fe(OH2)63+ + *OH2 Fe(OH2)5(*OH2)3+ + OH2

    Fe(OH2)62+ + *OH2 Fe(OH2)5(*OH2)2+ + OH2

    -e-+e -

    t1/2 = 4 ms

    t1/2 = 0.2 s

    13 References [7,8,9,10,11]

    Oxidation State Influence on Chelate StabilityOxidation State Influence on Chelate Stability

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 14

    Iron Chelation and TransportIron Chelation and TransportIn humans, the host protein transferrin(Tf) is produced in excess ofcirculating free iron and sequestersextracellular iron at extremely highaffinity (Kd ~10-20 M). This chelationof iron prevents it from precipitationand also has a bacteriostatic effect by keeping iron as an essential nutrientfrom being available to bacterialpathogens.

    FeIII

    O

    Otyrosine

    OtyrosineO

    Nhistidine

    O

    aspartate

    C

    O

    Human Transferrin Fe(III) Binding Site

    Fe3+ + apo-Tf + CO32-

    Kb ~ 1020 M-1

    FeIIITf(CO32-)

    14Reference [7]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 15

    Iron Chelation and TransportIron Chelation and TransportMicrobes 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).

    OOOO

    O

    O

    Ion Recognition

    microbial cell

    O

    O

    OO

    OO

    Fe2O3.6H2O or Fe(OH)3

    Fe3+

    Molecular Recognition

    Ksp ~ 10-39

    environmental iron

    siderophore

    Fe Release

    Complexation 1030 1050

    AlZnCr CuMn Ni

    KCa Pb

    Fe

    (i)

    (ii)

    (iii)(iii)

    (iv)

    (iv)(v)

    15 References [4,5,6]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 16

    Iron Chelation and TransportIron Chelation and TransportSiderophores, microbially synthesized Fe(III) specific chelators, are low molecularweight molecules that usually incorporate bidentate catechol, hydroxamic acid, and/or-hydroxy carboxylic acid donor groups. These chelators exhibit high Fe(III) complexstabilities (high and pFe) to enhance delivery of iron to the cell, and large negativeredox potentials (Panels 7, 12 and 18) for Fe(III) complexing specificity and toprevent 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 trishydroxamate (in blue).

    enterobactin

    Fe(III)-enterobactin complex110 = 1049 ; pFe = 35.5

    HOO

    OH HOO

    NHN

    O

    N

    HN

    O

    N

    NH3+

    O

    NH

    O

    O

    O

    OHN

    O O

    NH

    O

    O

    OH

    OH

    O

    OH

    HOOH

    HO

    O

    O

    O

    OOO

    NHN

    O

    N

    HN

    O

    N

    NH3+

    Fe

    desferrioxamine B

    ferrioxamine B complex110 = 1030.2 ; pFe = 26.6

    O

    O

    O

    O

    O O

    Fe

    O

    OOO

    O OHN

    O

    OHN

    O

    NH

    References [4,5,6] 16

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 17pH

    0 2 4 6 8 10 12

    pFe

    0

    10

    20

    30

    40

    Exochelin MN (I)Ferrioxamine B (II)

    Plot of Fe(III) complex stability,expressed as pFe (Panel 9), as afunction of pH for two siderophores,exochelin MN (I) and ferrioxamine B (II).Although they have approximately thesame stability at pH 6.0, above this pHexochelin MN has a higher affinity forFe(III) and below this value ferrioxamineB exhibits a higher affinity. This is dueto different levels of competition from H+for the Fe(III) binding sites, due todifferent pKa values for the donor groups (shown in blue) in these two siderophorechelators.

    NNH

    O

    NHN

    NH

    ONH

    NH2

    ON

    O

    OHOH

    NHO

    NH

    OOHNH2

    (I)

    (II)

    HOO

    OH HOO

    NHN

    O

    NHN

    O

    N

    NH3+

    O

    17

    Influence of pH on Fe(III)Influence of pH on Fe(III)--Chelate StabilityChelate Stability

    Reference [9,12]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 18

    pFe5 10 15 20 25 30

    -E1/

    2 (m

    V) v

    s NH

    E

    200

    250

    300

    350

    400

    450

    500

    8

    6

    53

    12

    7

    4

    Influence of Fe(III)Influence of Fe(III)--Chelate Stability on EChelate Stability on E00

    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 siderophoresand siderophore mimics. Note that:

    as the stability of the Fe(III)-complex increases, the complex becomes moredifficult to reduce; and

    the stability of the complex decreases with decreasing denticity.

    1. Fe(desferrioxamine B)+ 2. Fe(Desferrioxamine E) 3. Fe2(alcaligin)3 4. Fe(saccharide-trihydroxamate) 5. Fe2(rhodotorulic acid)3 6. Fe(N-methylacetohydroxamate)37. Fe(acetohydroxamate)3 8. Fe(L-lysinehydroxamate)3

    18 References [6,10,13,14]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 19

    Influence of Fe(III)Influence of Fe(III)--Chelation on KineticsChelation on Kinetics

    Fe

    OH2

    OOO

    O OH2

    H3C

    NCH3

    Fe

    OH2

    OH2

    OH2H2O

    H2O OH2Fe

    OH2

    OOH2O

    H2O OH2

    H3C

    NCH3

    N

    H3C

    H3C

    OHO

    H3C

    NCH3

    O

    OHN

    H3C

    H3C

    H+

    H+

    3+ 2+

    1+

    k1 = 1.8 M-1s-1

    k2 = 8.1 x 102 M-1s-1

    Iron chelate formation places a strong electron donor in the first coordination shell, which labilizes the remaining 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).

    19 Reference [5,15]

  • Iron in Biology Society For Free Radical Biology and Medicine Crumbliss 20

    ReferencesReferences1. Crichton, R. (2001) Inorganic Biochemistry of Iron Metabolism, John

    Wiley & Sons, Ltd, New...

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