174
V. Schünemann
4.1 Iron Centers in Chemical Complexes
and Biomolecules: Structural Overview, Biological
Relevance and Physical Properties
The trace element iron is indispensable for almost all life forms on earth. The best
known example is the function of iron as an oxygen carrier in the red blood pigment
hemoglobin. However, by exploiting its redox properties, iron also plays an important
role in biological electron transport chains, e.g. in mitochondria or in photosynthetic
reaction centers of thylakoid membranes of green plants [1]. Many enzymes, in turn,
use the catalytic activity of iron, such as the enzyme cytochrome P450, which is
also vital for humans and occurs in liver cells of mammals [2]. Iron is present in
mononuclear form either in heme centers [3, 4] or as non-heme iron [5], which is
directly bound to amino acid chains of the protein. Iron ions can also be linked by
oxygen or carboxyl bridges to dimeric centers such as methane monooxygenase [6]
or ribonucleotide reductase [7].
Iron centers bridged by sulfur can be found in iron-sulfur proteins. The function
of these centers ranges from electron transfer to catalysis [8]. Unbound iron can
produce free radicals in the presence of water, which are a danger to any cell. For
this reason, iron metabolism in any organism is strictly regulated. Ferritins, proteins
containing amorphous iron hydroxide clusters with diameters up to 8 nm [9], are
used for iron storage. Biological complexing agents, transferrins, which contain iron
in monomeric form and bind so strongly that its catalytic activity cannot lead to the
formation of free radicals even in the presence of water, are used for iron transport
[10].
The most common coordination environment of iron ions in biomolecules and
chemical complexes are the sixfold octahedral and the fourfold tetrahedral coordination (Fig. 4.1) [11]. Sixfold coordination is found in heme proteins involved in
electron transfer reactions like the cytochrome c family and also in the iron center of
hemo and myoglobins when O 2 is bound. Iron sulfur proteins which are involved in
electron transfer, but can also have catalytic (enzymatic) functions have tetrahedral
sulfur coordinated iron centers in most cases.
The type of coordination defines the order of the 3d-orbitals of the iron center
due to the electronic interaction with the surrounding ligands. These are in ligand
field theory regarded as negatively charged point charges (Fig. 4.2). This is quite a
simplification of the “real” electronic situation, but describes the electronic properties
of iron centers often quite astonishing well. However, this simple picture fails when
strongly covalent bonds are formed like metalorganic bonds.
Iron in biomolecules occurs in different spin and oxidation states. Divalent iron
with 4 unpaired electrons in the 3d shell often occurs in the S = 2 state, where one
3d orbital is double and the others are single occupied according to Hund’s rule
(Fig. 4.3). Such a state is also often called a ferrous high spin state. If the ligands
exert a strong ligand field, this leads to the diamagnetic S = 0 form of the divalent
iron [12]. The latter state is also often called a ferrous low spin state.
V. Schünemann
4.1 Iron Centers in Chemical Complexes
and Biomolecules: Structural Overview, Biological
Relevance and Physical Properties
The trace element iron is indispensable for almost all life forms on earth. The best
known example is the function of iron as an oxygen carrier in the red blood pigment
hemoglobin. However, by exploiting its redox properties, iron also plays an important
role in biological electron transport chains, e.g. in mitochondria or in photosynthetic
reaction centers of thylakoid membranes of green plants [1]. Many enzymes, in turn,
use the catalytic activity of iron, such as the enzyme cytochrome P450, which is
also vital for humans and occurs in liver cells of mammals [2]. Iron is present in
mononuclear form either in heme centers [3, 4] or as non-heme iron [5], which is
directly bound to amino acid chains of the protein. Iron ions can also be linked by
oxygen or carboxyl bridges to dimeric centers such as methane monooxygenase [6]
or ribonucleotide reductase [7].
Iron centers bridged by sulfur can be found in iron-sulfur proteins. The function
of these centers ranges from electron transfer to catalysis [8]. Unbound iron can
produce free radicals in the presence of water, which are a danger to any cell. For
this reason, iron metabolism in any organism is strictly regulated. Ferritins, proteins
containing amorphous iron hydroxide clusters with diameters up to 8 nm [9], are
used for iron storage. Biological complexing agents, transferrins, which contain iron
in monomeric form and bind so strongly that its catalytic activity cannot lead to the
formation of free radicals even in the presence of water, are used for iron transport
[10].
The most common coordination environment of iron ions in biomolecules and
chemical complexes are the sixfold octahedral and the fourfold tetrahedral coordination (Fig. 4.1) [11]. Sixfold coordination is found in heme proteins involved in
electron transfer reactions like the cytochrome c family and also in the iron center of
hemo and myoglobins when O 2 is bound. Iron sulfur proteins which are involved in
electron transfer, but can also have catalytic (enzymatic) functions have tetrahedral
sulfur coordinated iron centers in most cases.
The type of coordination defines the order of the 3d-orbitals of the iron center
due to the electronic interaction with the surrounding ligands. These are in ligand
field theory regarded as negatively charged point charges (Fig. 4.2). This is quite a
simplification of the “real” electronic situation, but describes the electronic properties
of iron centers often quite astonishing well. However, this simple picture fails when
strongly covalent bonds are formed like metalorganic bonds.
Iron in biomolecules occurs in different spin and oxidation states. Divalent iron
with 4 unpaired electrons in the 3d shell often occurs in the S = 2 state, where one
3d orbital is double and the others are single occupied according to Hund’s rule
(Fig. 4.3). Such a state is also often called a ferrous high spin state. If the ligands
exert a strong ligand field, this leads to the diamagnetic S = 0 form of the divalent
iron [12]. The latter state is also often called a ferrous low spin state.