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Biophysical and Bioinorganic

ChemistryProf. Dr. Tatjana N. Parac-Vogt

Dr. Gregory Absillis

Department of ChemistryKU Leuven, Belgium

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What is Biophysical/Bioinorganic

Chemistry?

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Physical Methods in

Bioinorganic Chemistry

Principles of

Inorganic Chemistry

Related to

Bioinorganic Research


Bioinorganic ChemistryAn Inorganic Perspective of Life

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Bioinorganic chemistry constitutes the discipline at the interface

of the more classical areas of inorganic chemistry and biology

Biophysical chemistryis a physical science that uses the concepts of physics andphysical chemistry for the study of biological systems.

EPR study of the bacteriochlorophyll reaction center (RC) of Rb. spaeroides


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Principles of

Inorganic Chemistry

Related to

Bioinorganic Research


The Hard-Soft-Acid-Base concept

Electronic and geometric stuctures of

metal ions

Tuning of redox potentials

pKa values of coordinated ligand

Ligand exchange kinetics

...

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Physical Methods in



EPR spectrocopy

NMR spectrocopy

Mssbauer spectroscopy

EXAFS spectrocopy

CD spectroscopy

...

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1 The content will be FUNCTIONbased:

Metalloproteins:

O2transport

e-transfer

structural role

metalloenzymes

hydrolytic enzymes

e- reduction

rearrangements

Communication

Interaction with nucleic acids

Metal ion transport and storage

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Crystal structure of ferritin multimer

(24mer) and monomer.


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2 Metal based probes and diagnostic/therapeutical pharmaceuticals

Cardiolyte, Tc(CNR)6heart imaging agent

Auranofin, arthritis drug

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3 Biomimics for catalysis


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Mono- and bimetallic polyoxometalate complexes

as mimics for the active site of cytochrome P-450

and methane monooxygenase

Bimetallic Cu-complexes as artificial phosphatases


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Metals in Biological Systems

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Na, K, and Cl:

Osmotic control, electrolytic equilibria,

Mg:

Phosphate metabolism,

P:

DNA, RNA, ATP,

S:

Amino acids,

Analysis of many bacteria, a few hundred among the known 0.4 million plant

species, and of about 200 among the catalogued 1.1 million animal species as

well as of organs, tissues, and other substances of biological origin, have enabled

us to establish the number and identity of the chemical elements present in

biological systems and to recognize those that are essentialfor bacterial, plant

and animal life:

1. Eleven elements appear to be approximately constant and predominant in all

biological systems (99.9% of the total number of atoms present in the human

body).

Element Atom Percentage

H 62.8

O 25.4

C 9.4

N 1.4

99.0

Due to highH2O

content

Basic elements of organic structures

and metabolites (H, O, C, N)

Na, K, Ca, Mg, P, S, and Cl (0.9%)


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2. Ten metals and non-metals are required by most biological systems, but not

necessarily by all biological systems:

Mn, Fe, Co, Ni, Co, Zn, Mo, B, Si, and Se

3. Eight elements, some of which may be required by plants and animals,

whereas others may be required by just plants or just animals or sometimes by

relatively few species of plants or animals:

W, V, Cr, F, I, As, Br, Sn

4. Cd, Sr, and Ba are known to be important in the chemistry of one or two

particular species.

The limited number of species examined, the difficulties of the analytical work,and the lack of detailed knowledge concerning the role of each element make it

necessary to check hypothetical essentiality by delicate tests, usually by

following the developmental growth of species, animal or plant, while giving

them carefully prepared diets deficient in the particular element considered.


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Definition: In general terms we will consider as essential an element consistently present in

a certain biological species such that its deficiency in the nutritive sources of

that species leads to disease, metabolic anomalies, or perturbations in its

development.

Essential or not essential Evaluation of the effects of its deficiency

Mainly lighter elements (Z < 36)distributed over practically all groups

All organisms require about 20 elements though the precise

requirement differs somewhat within different species


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In some cases essentiality studies have enabled the classification of a certain

element as essential or not essential.

In other cases the results are ambiguous:

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Effects of deficiency not fully tested?

Requirements are so low that trace

amounts in carefully purified diets

satisfy the need?

Uptake essentiality: some species

accumulate elements even though

they may not need them


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Mainly lighter elements (Z < 36)

distributed over practically all groups

Na

K

Li

Rb

Cs

Group 1

Mg

Ca

Be

Sr

Ba

Group 2

Zn

Cd

Hg

Group 12

Cl

Br

F

I

Group 17

B

Si P

Non-redox non-metals

H

N O

Redox non-metals

C

S

Se

Redox metals

Cr Mn Fe Co Ni Cu

Mo

Chemical Nature

Reactivity

Functionality


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Chemical Nature Reactivity Functionality

Metal ions may serve multiple functions depending on their location within

the biological system, so that there classification is somewhat arbitrary and

overlapping:

Group 1 and 2 metals operate as structural elements or in the

maintenance of charge and osmotic balance.

Transition metals that exist in only one oxidation state, such as Zn(II),

function as structural elements

Transition metals that exist in multiple oxidation states serve as:electron carriers

facilitators of oxygen transport

sites at which enzyme catalysis occurs


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Na

K

Li

Rb

Cs

Group 1

1. Charge Carriers (Maintenance of charge and osmotic balance)


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Mg

Ca

Be

Sr

Ba

Group 2

Zn

Cd

Hg

Group 12

2. Structural and Triggers (Structural elements in SOD, zinc fingers, triggers for

protein activity, )


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3. Electron transfer (cytochrome, nitrogenase activity, ) Redox metals

Cr Mn Fe Co Ni Cu

Mo



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4. Oxygen Transfer (transport and storage of O2) Redox metals

Cr Mn Fe Co Ni Cu

Mo



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5. Enzyme Catalysis (hydrolysis, rearrangement, oxido reduction )

Redox metals

Cr Mn Fe Co Ni Cu

Mo



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Inorganic Chemistry Basics

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Hard-Soft Acid-Base Classification

Ligand preference and possible coordination geometries of the metal center are important

bioinorganic principles.

Metal-ligand preference is closely related to the hard-soft acid-base nature of metals and

their preferred ligands:

Hard metal cations form their most stable compounds with hard ligandsSoft metal cations form their most stable compounds with soft ligands

Hard metal cations: hard dense less polarizable cores of positive charge, e.g. Na+, Ca2+,

Co3+, Fe3+,

Hard ligands: small electronegative elements or ligand atoms with a hard polyatomic ion,

e.g. oxygen ligands in (RO)2PO2-, crown ethers,

Soft cations and anions: characterized by highly polarizable large electron clouds, e.g.

Hg2+, sulfur ligands,


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In biological systems these

ligands are provided by

protein side chains, the basesof nucleic acids, small cellular

cytoplasmic constituents,

organic cofactors, water,

e.g. alkali and alkaline earth

metals are like Ca2+ are mostcommonly coordinated by

carboxylate oxygen atoms,

Fe3+ by carboxylate and

phenoxide oxygen donors,

and Cu2+ by histinen

nitrogens.

For species that can have

multiple oxidation states , the

lower oxidation state is softer

than the higher.



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Nucleophilic attack of the C6-OH of glucose on the -phosphate of a Mg2+-ATP complex in

hexokinase.

ATP

Mg2+Glucose

The hard acid Mg2+stabilizes hard bases including ATP and tRNA via its strong interaction

with phosphate groups.


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Nearly 30-35 percent of the amino acids of metallothionein proteins are cysteine residues

containing sulfhydryl groups that bind avidly to soft metal ions such as Cd2+, Hg2+, Pb2+and

Tl+hereby protecting the cell against the toxic effects of these metal ions.

Metal ion

Cysteine residue

Tetrametallic and trimetallic clusters (Cd2+, Hg2+, Pb2+and Tl+) in metallothioneins


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Cu2+in hemocyanin

Hemocyanins are proteins that transport oxygen throughout the bodies of some

invertebrate animals. These metalloproteins contain two copper ions that reversibly bind a

single oxygen molecule. Whereas hemoglobin carries its iron atoms in porphyrin rings, the

copper ions of hemocyanin are bound as prosthetic groups coordinated by histidine

residues.


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The Chelate Effect

Chelation refers to the coordination of two or more donor atoms from a single ligand to a

central metal ion. The resulting metal-chelate complex has an unusual stability derived in

part from the favorable entropic factor accompanying the release of nonchelating ligands,usually water, from the coordination sphere:

[Cu(H2O)6]2++ 2NH3 [Cu(H2O)4(NH3)2]

2++ 2H2O

H = -46 kJ/mol

S = -8.4 J/K mol

[Cu(H2O)6]2++ en [Cu(H2O)4(en)]

2++ 2H2O

H = -54 kJ/mol

S = +23 J/K mol

The reaction enthalpies (H)are fairly similar, because two Cu-N bonds are formed in each

cases, but the reaction enthalpies (S)differ greatly. The difference is understood in terms

of the change in total number of molecules in each reaction. In the ammonia reaction,

there is no net change in the total number of molecules: two ammonia molecules become

coordinated to the copper ion releasing two water molecules. In contrast, for the reaction

with en, the net number of molecules increases: one en molecule becomes coordinated to

the copper, causing the release of two water molecules. Consequently, the disorder or

entropy increases more in the cases of the en reaction making it thermodynamically more

favorable.

Ethylenediamine (en)


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The Chelate Effect

The chelation of Ni2+ by ETDA (ethylenediaminetetraacetate):

[Ni(OH2)6]2++ H2edta

2- [Ni(edta)]2-+4H2O +2H3O+

Applications of EDTA:

Medicine: chelate metal ions that might be present in toxicexcess.

Food industry: Limit the availability of essential elements

to harmful bacteria.

Research: Reduce the concentration of free metal ions that

could promote undesired side reactions.


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Note that the coordination of histidine residues in hemocyanins also represents an

example of a chelating effect in which several histidine residues belonging to the same

polyppetide chain coordinate to various copper centers in the protein.

The Chelate Effect


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The Chelate Effect

An important example of the chelate effect in bioinorganic chemistry is afforded by

porphyrin, corrin and chlorin ligands. These macrocyclic molecules have four nearly planar

pyrrole rings with their nitrogen donor atoms directed towards the central metal ion.

The resulting metallo-porphyrin, -corrin, or -chlorin units are thermodynamically very

stable, accommodating a variety of metal ions in different oxidation states.

Pyrrole ring


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The Chelate Effect

Fe2+ in cytochrome c

The Fe2+ in cytochrome c is

octahedrally coordinated to the 4

equatorial N atoms of the

porphyrine ring, the N atom of a

His residue and the S atom of a

methionine residue.


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The Chelate Effect

The structure of vitamin B12is based on a corrin ring. The central metal ion is Co+

. Four ofthe six coordination sites are provided by the corrin ring, and a fifth by a

dimethylbenzimidazole group. The sixth coordination site, the center of reactivity, is

variable, being a cyano group (cyanocobalamin), a hydroxyl group (hydroxycobalamin), a

methyl group (methylcobalamin) or a 5'-deoxyadenosyl group (adenosylcobalamin). Here

the C5' atom of the deoxyribose forms the covalent bond with Co+.

Co+in cobalamin


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The Chelate Effect

Mg2+in chlorophyll a, b, and d

Chlorophyll is a chlorin pigment, which is structurally similar to and produced through the

same metabolic pathway as other porphyrin pigments. At the center of the chlorin ring is a

magnesium ion.


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The Chelate Effect

Fe2+in hemoglobin

The heme group consists of Fe2+

held in a porphyrin ring. The iron ion, which is the site ofoxygen binding, coordinates with the four nitrogens in the center of this ring, which all lie

in one plane. The iron is bound strongly (covalently) to the globular protein via the

imidazole ring of the histidine residue (also known as the proximal histidine) below the

porphyrin ring. A sixth position can reversibly bind oxygen completing the octahedral

group of six ligands.


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pKaValues of Coordinated Ligands

An acid dissociation constant, Ka, (also known as acidity constant, or acid-ionization

constant) is a quantitative measure of the strength of an acid (HA) in solution:

HA H++ A-

pKa= - log Ka

When a protic ligand is bound to a metal ion, the ligand generally becomes more acidic,

because the positively charged metal ion stabilizes the anionic conjugate base of the

ligand. This effect is best exemplified by coordinated water, but occurs for many other

biological ligands such as thiols; imidazole, phenols, alcohols, phosphoric and carboxylic

acids, and their derivatives.


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The enzyme carbonic anhydrase contains a Zn2+ion at its active site and a watermolecule

molecule bound to it. A free water molecule in bulk water has a pKaof 15. Binding of the

water molecule to Zn2+

lowers the pKa to 7. The Zn2+

bound aqua ligand is thereforedeprotonated to a significant extent at physiological pH, giving a Zn2+ -hydroxy complex.

The hydroxo group acts as a nucleophile and attacks CO2to form HCO3- in the enzymatic

mechanism.


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Trivalent metal ions are better able to lower the pKa values of protic ligands than their

divalent analogs, as expected on the basis of charge considerations.

Coordination of two or more metal ions to a protic ligand lowers the pKaeven more:

Deprotonation of the side chain of histidine

resulting in a 2-imidazolato copper(II)-zinc(II)

moeity of SOD.


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Sometimes both protons dissociate from the aqua ligand to form mononuclear oxo, O2-

complexes:

Generation of a tyrosyl radical in the active site of ribonuclease reductase

-oxo bridge between

the two iron centers

d l


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Redox Potential

Oxidation is the loss of one or more electrons and a species that loses one or more

electrons has been oxidized. The species that accepts this electron(s) is reduced or

undergoes reduction. A reaction in which one reactant is reduced while another is

oxidized is referred to as a redox reaction.

Oxidation reaction can be split into two half-reactions, one for the oxidation, and one for

the reduction that together represent the overall reaction. The reduction of H+by metallic

iron, for example, can be split into two half-reactions:

Overall redox reaction: Fe + 2H+Fe2+ + H2(g) E = E(H)-E(Fe) = +44 V

Oxidation half-reaction: Fe Fe2+ + 2e- E(Fe) = - 44V

Reduction half-reaction: 2H++ 2e-H2(g) E(H) = 0 V

The potential for each redox reaction at standard conditions is E. By convention, half-

reactions are usually written as reduction reactions, with their potentials listed as

reduction potentials. When the reduction half-reaction is written as an oxidation, the sign

of E is reversed.

d l


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Redox Potential

The more positive the potential, the greater the species' affinity for electrons and

tendency to be reduced.

R d P i l


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Redox Potentiale- flow in the mitochondria:

The components of the respiratory chain contain a variety of redox cofactors. Complex I

contains five iron-sulfur clusters and FMN. Complex II contains several iron-sulfur clusters,

FAD, and cytochrome bS68. Complex III contains a [2Fe-2S] iron-sulfur center and cytochromes

bS62, bS66, and c1. Complex IV contains at least two copper ions and cytochromes a and a3.

R d P i l


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Redox Potential

A cascade effect in which the component which has the most positive redox potential gets

reduced. Once reduced that component acts as a reducing agent for next component.

R d P i l


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Redox Potential

Alteration in the ligand donor atom and stereochemistry at the metal center can produce

great differences in the potential at which electron transfer will occur.

Cu+, a closed shell, d10 ion prefers tetrahedral 4-

coordinate or trigonal 3-coordinate geometries. Cu2+

complexes, on the other hand, are typically square

planar with sometimes one or two additional weakly

bound axial ligands.

A ligand environment that produces a tetrahedral

geometry will stabilize Cu+ over Cu2+ rendering the

latter a more powerful oxidizing agent by raising the

redox potential. Adding bulky R groups in Cu(R-sal)2

distorts the geometry from planar to tetrahedral,

making it easier to reduce copper and raising itspotential.

Cu+ is a soft acid preferring to bind to soft ligands such as RS-and R2S. Soft ligands in the

coordination sphere increase the Cu+/Cu2+ potential.

R d P t ti l


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Redox Potential

High redox potentials in copper containing proteins (e.g. azurin E 300-800mV) are

obtained through distortion of the coordination geometry towards trigonal planar ortetrahedral and the use of histidine imidazole and cysteine thiolate side chain as donor

ligands:

Li d E h R t


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Ligand Exchange Rates

M-OH2bonds are very labile, breaking and reforming as fast as a billion times per second.

The labilities of metal-ligand bonds typically follow the trends for aqua complexes:

Ligand exchange rates are faster for less highly charged M2+ions

than for M3+metal ions.

Kinetically

inert

Li d E h R t


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Second and third row transition metal complexes are much more kinetically inert than their

first row counter parts:

Cisplatin binds to DNA through the loss of Cl- ligands. The Pt can not be exchanged even

upon prolonged dialysis of the platinated DNA. Only strong Pt-binding ligands such as

cyanide can displace the Pt-adduct.

Li d E h R t


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The fast metal-ligand exchange rates of first row transition metal ions such as Fe2+ are

remarkably diminished when they are bound to multidentate chelates such as porphyrines:

The axial ligands, which are not part of the chelate ring undergo exchange at rather fast

rates. However, ligands such as CO (carbon monoxide intoxication), CN-en RS- form more

inert M-L bonds.

El t i d G t i St t f M t l I


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Electronic and Geometric Structure of Metal Ions

The d-electron configuration is obtained by

subtracting the formal oxidation state from

the atomic number Z and calculating how

many electrons must be added to the

preceding noble gas element (usually Ar, Z =

18):

Fe3+: 26-3-18 = 5

Mo

4+

: 42-4-36 = 2Cu+: 29-1-18 = 10



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The most common coordination geometries for coordination numbers 3 to 6 for metals

encountered in bioinorganic chemistry. Substantial distortions from these idealized

structures can occur.



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When a metal ion in a given formal oxidation state is placed at the center of a coordination

polyhedron defined by a set of ligands, the energy levels of the d-orbitals housing these the

metal electrons are altered from those found in the free metal ion. This phenomenon iscalled ligand field splitting.



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The Octahedral Field:

The d(z2) and d(x2-y2) orbitals comprise the upper pair of orbitals and are referred to as eg

orbitals. The lower set of orbitals, d(xy), d(xz) and d(yz), are referred to as the t 2gorbitals.

The two levels egand t2gare seperated by an amount, owhich is known as the ligand field

splitting parameter.

The magnitude of odepends on the identity of the metal ion, its charge, and the nature of

the ligands.



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The Tetrahedral Field:

The d(xy), d(xz) and d(yz) orbitals comprise the upper pair of orbitals and are referred to as

t2orbitals. The lower set of orbitals, d(z2) and d(x2-y2) , are referred to as the e orbitals.

The two levels e and t2are seperated by an amount, twhich is known as the ligand field

splitting parameter.

The magnitude of tdepends on the identity of the metal ion, its charge, and the nature of

the ligands.

0.4t

0.6t

e

t2



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It is valuable to compare the magnitude of twith that of o for two complexes, one six

coordinate (octahedral) and one four coordinate (tetrahedral), in which the metal ions, the

ligands and the M-L bond lengths are the same. Intuition suggests that tshould be smallerthan o simply because it is caused by interaction with four rather than six ligands. This

finding is indeed the case, and it is shown that tis in fact 4/9 the value of o, when all else

is equal.



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For a given ligand geometry and specific metal ion, the crystal field parameter is

dependent on the nature of the ligand. For example, coordination of the ligands CN-and

CO always lead to a large relative to the coordination of halide ligands such as I -and Br-.

Measurements of many complexes reveal an ordering of common ligands in terms of their

relative splitting. This ordering is known as the spectrochemical series:

CN- CO > NO2-> 2,2-bipyridine > ethylenediamine > NH3> edta >

NCS-> H2O > OH-> F-> Cl-> Br-> I-

The ligands towards the high end of the series are know as strong-field ligands, and thoseon the low end as weak-field ligands. Strong-field ligands typically result in larger orbital

splittings than do weak-field ligands.



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The coordination geometry and the nature of the ligands determine the magnitude and

complexity of the d-orbital splitting in the transition metal ion. The splitting contributes to

the spectroscopic and magnetic properties of the complex as well as to its stability.

As a result, these ligand field splitting diagrams are extremely useful when attempting to

correlate the physical properties of metal centers in proteins (optical spectra, magnetism,

EPR spectra, ) with their structure and reactivity.

If, for example, one were to encounter a diamagnetic (no unpaired electrons) Ni2+center in

a protein, it most likely would have a square planar geometry, since both tetrahedral andoctahedral d8 complexes would be expected to have two unpaired electrons and be

paramagnetic:



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The splitting of the energy levels of the d-orbitals has major effects on the UV-Vis-NIR

electronic absorption spectra of transition metal complexes. The energy (wavelength) of

spectral bands due to transitions between d-orbitals is strongly influenced by the splittingconstant , while the number of bands is generally determined by the number of levels

that are formed:

Weak-field ligand Strong-field ligand

Cr3+: d3- Octahedral geometry



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Co3+: d3- Octahedral geometry

As we go from weak- to strong field ligands, o increases and hence the absorption

wavelength decreases.



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In general, tetrahedal complexes absorb at lower energy (higher wavelength) than

octahedral complexes as L< o.

Tetrahedral complexes absorb with more intensity than octahedral complexes.A good example of this difference in intensity can be seen in comparing solutions of

[CoCl4]2- (tetrahedral) and [Co(H2O)6]

2+ (octahedral). At the same concentration of cobalt,

the solution of the former complex is intensely blue, while that of the later is pale pink. The

blue color is due to absorption in the orange part of the Vis spectrum, and the pink from

absorption in the green, a shorter wavelength.

[CoCl4]2-[Co(H2O)6]

2+



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In addition to electronic transitions due to metal ion centered d-d transitions, transition

metal complexes also can have electronic transition in which electrons transfer from ligand

to metal and vica versa.

A ligand-to-metal-charge transfer (LMCT) band excitation arises from the excitation of an

electron in a ligand centered orbital into a d-orbital of the metal ion. Such transitions

typically give intense absorptions. The LMCT bands can vary from UV excitations in hard

ligands (O,N) to Vis excitations for soft ligands (S).

For example, the spectra of iron-sulfur

proteins show a broad absorption

envelope in the Vis-NIR range

resulting from several overlapping

absorption bands derived from

transition with predominant S Fe3+

charge transfer character (S = Cys of

inorganic sulfur):



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Electronic and Geometric Structure of Metal IonsThe magnetic properties of transition metal complexes depend on the number of unpaired

electrons that reside in the d-orbitals, which in turn depends on the strength of the field

created by the surrounding ligands (i.e. the magnitude of the d-orbital splitting). If the

splitting is greater than the energy required to pair electrons in a single orbital (i.e. thepairing energy), then the metal exists in a so-called low-spin state. On the other hand, if

the pairing energy is greater than the splitting, then a high-spin state will occur.

Tetrahedral complexes: The small tis always less than the pairing energy. Tetrahedral

complexes are therefore all high-spin.

Octahedral complexes: the general rule of thumb is that strong field ligands lead to low-

spin states, while weak ligands lead to high spin states.


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