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Transcript of Bic Parti Vn
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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 Chemistry
Bioinorganic ChemistryAn Inorganic Perspective of Life
3
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
Bioinorganic ChemistryAn Inorganic Perspective of Life
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
Bioinorganic Chemistry
Bioinorganic ChemistryAn Inorganic Perspective of Life
EPR spectrocopy
NMR spectrocopy
Mssbauer spectroscopy
EXAFS spectrocopy
CD spectroscopy
...
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Bioinorganic ChemistryAn Inorganic Perspective of Life
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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Bioinorganic ChemistryAn Inorganic Perspective of Life
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
Bioinorganic ChemistryAn Inorganic Perspective of Life
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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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Chemical Nature Reactivity Functionality
Na
K
Li
Rb
Cs
Group 1
1. Charge Carriers (Maintenance of charge and osmotic balance)
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Chemical Nature Reactivity Functionality
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
Chemical Nature Reactivity Functionality
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4. Oxygen Transfer (transport and storage of O2) Redox metals
Cr Mn Fe Co Ni Cu
Mo
Chemical Nature Reactivity Functionality
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5. Enzyme Catalysis (hydrolysis, rearrangement, oxido reduction )
Redox metals
Cr Mn Fe Co Ni Cu
Mo
Chemical Nature Reactivity Functionality
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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.
Hard-Soft Acid-Base Classification
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Hard-Soft Acid-Base Classification
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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Hard-Soft Acid-Base Classification
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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Hard-Soft Acid-Base Classification
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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pKaValues of Coordinated Ligands
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pKaValues of Coordinated Ligands
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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pKaValues of Coordinated Ligands
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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pKaValues of Coordinated Ligands
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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Ligand Exchange Rates
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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Ligand Exchange Rates
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
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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:
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
Co3+: d3- Octahedral geometry
As we go from weak- to strong field ligands, o increases and hence the absorption
wavelength decreases.
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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+
Electronic and Geometric Structure of Metal Ions
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Electronic and Geometric Structure of Metal Ions
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):
Electronic and Geometric Structure of Metal Ions
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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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