Biological Inorganic Chemistry - Coordination Chemistry in Action

When a metal ion forms a bond with a donor, understanding the bond's nature is crucial.
Consider how many bonds form when a metal ion inserts into a biological environment.
Water is a significant ligand in aqueous environments, with multiple water molecules surrounding a metal ion.
Bare metal ions (e.g., $Fe^{2+}$ or $Fe^{3+}$ ) in an aqueous solution attract several water molecules (4-6) to form a polyhedral arrangement.
This arrangement often results in an octahedral geometry, crucial to understand in biological systems.

In laboratory settings, dissolving ferric chloride in water results in ferric ions surrounded by water molecules, forming an octahedral geometry (6 $Fe-O$ bonds).
It's important to grasp the basic structure of an octahedron.

Metal ions in biological systems can act as redox species, facilitating electron transfer reactions essential for survival.
These reactions are vital for acceptor molecules such as hemoglobin and myoglobin.

Metal ion centers can exhibit catalytic potential.
For example, Magnesium (Mg2+) acts as a cofactor in enzymes, contributing to their enzymatic activity.
Mg2+ cannot undergo electron acceptance or donation; hence, it does not function as a redox catalyst.

Electron pairs, which are negatively charged clouds, occupy a definite region of space relative to the nucleus.
These regions are called orbitals and can hold a maximum of two electrons.
Electron pairs with opposite spins can act as donor molecules (e.g., water or ammonia).

Coordinate bonds are a special type of covalent bond where a pair of electrons is donated to a metal ion center.
For example, a hydrogen atom can form an $H-H$ bond by overlapping their orbitals, with the bonding molecular orbital filled with a pair of electrons, which can then be donated to a Lewis acidic metal ion center.

Ligands are classified based on their electronic and structural properties:
- Electronically: based on the number of electrons donated to the central metal ion.
- Structurally: based on the number of connections it makes.
For Example, Six water molecules can surround an $Fe^{3+}$ ion, each donating a pair of electrons.
The D electrons in metal ions occupy orbitals in three-dimensional space (X, Y, and Z axes).
These D orbitals can accept electron pairs if vacant, leading to a polyhedral arrangement and a structural feature with six connections in a three-dimensional space, resulting in an octahedral shape.

Water molecules effectively complex with metal ions.
Adding a ligand like ammonia can displace water molecules if ammonia has a higher coordination power.
For example, adding dilute ammonia to a copper sulfate solution results in a color change to blue due to the formation of new $Cu-N$ bonds instead of $Cu-O$ bonds.
Multidentate ligands offer higher stability.
Monodentate ligands, like ammonia, replace individual water molecules in a 1:1 manner.

Bidentate ligands, such as ethylenediamine ( $NH<em>2CH</em>2CH<em>2NH</em>2$ ), can displace water molecules.
Ethylenediamine (en) can displace ammonia already bound to a copper center.
Using chelating ligands increases entropy because more water molecules are released into the medium, leading to greater stabilization.
The number of particles released in the medium is more -> energy gain from the entropy point of view.

Common donor groups in biological systems include nitrogen, oxygen, and sulfur.
Amino acids like glycine have $NH_2$ and $COOH$ ends that can bind to metal ions, acting as $N, O$ bidentate ligands.
Ethylenediamine (IUPAC name: 1,2-ethanediamine, abbreviated as en) is a bidentate ligand.
Diethylenetriamine (1,4,7-triazaheptane) is a tridentate ligand.
Triazacyclononane (tacn) is a cyclic ligand with different properties compared to linear ligands.

Binap ligand ( $2, 2'$ -bis(diphenylphosphino)- $1, 1'$ -binaphthyl) is a laboratory-prepared ligand.
Nickel DMG (dimethylglyoxime) complex involves a DMG ligand bound to a nickel center, with one of the oxime functions deprotonated.
Hydrotris(pyrazolyl)borate (TP) is a scorpionate ligand: a tris pyrazolyl borate that binds to a metal ion center in a scorpion-like manner.

Solutions often have solvent molecules coordinating with metal ions, resulting in higher coordination numbers.
Coordination of six water molecules around a ferric ion center is a typical example.
It is difficult to achieve solutions where only one ligand is bound to the metal ion center.
Thallium ( $Tl^+$ ) can form a complex with coordination number one is satisfied, binding to a sterically crowded triazapentadienly ligand through a single $Tl-N$ bond.

When tacn coordinates to an iron center with three ethanethiol arms, it forms an octahedral arrangement with two phases: SSS and NNN.
The S3 triangle is bigger because sulfur ends are free and larger than the nitrogen atoms which are part of the macrocyclic ring, resulting in asymmetry within the octahedral geometry.

The nature of the ligand, donor atom, and stereochemistry affects the electron transfer behavior of metal ions.
Walther Nernst developed the Nernst equation in 1881 to relate the cell potential with the standard potential (E0 value): $E = E_0 - [\frac{RT}{nF}]ln[\frac{[Red]}{[Ox]}]$
- Where:
  - $E$ = Cell potential
  - $E_0$ = Standard potential
  - $R$ = Gas constant
  - $T$ = Temperature
  - $n$ = Number of moles of electrons transferred in the electrochemical reaction
  - $F$ = Faraday constant
  - $[Red]$ = Concentration of the reduced species
  - $[Ox]$ = Concentration of the oxidized species
E0 values can be measured experimentally to understand thermodynamic electron transfer behavior.
Reactions can involve electron transfer, protonation, or substrate binding following electron withdrawal or donation.

Reactions in gases, such as oxygen accepted through myoglobin and hemoglobin, and reactions involving NO produced by the body.
Oxidation of organic substrates in solution.
Reactions in solids, such as those in lithium batteries (e.g., $LiCoO_2$ reacting with carbon to form lithium carbide and cobalt oxides).

In the water oxidation center (WOC), four manganese centers of different oxidation states react with water to liberate $O_2$ molecules.
This process involves a four-electron transfer reaction, reducing the oxidation states of the manganese centers.
Photosynthesis is associated with the liberation of dioxygen molecules from water.

The influence of complexes in reactions, such as hydrogen for electron transfer, is significant.
The standard reduction potential for hydrogen is 0.0 V.
At pH 7 (biological pH), the reduction potential of hydrogen is -0.421 V.

Biological molecules, such as hemoproteins, use E0 values.
The E0 values and the environment, coordination environment of hexaaquairon species affect catalytic biochemical processes.
Changing the coordination environment alters the redox potential values.

The interaction between metal ions ( $M^{n+}$ ) and ligands (L) can be assessed through E0 values, stability, and structure.
Electronic and structural effects in the coordination sphere are important.
Secondary interactions, like hydrogen bonding in DMG, play a role.
The concept of electro-protic equilibrium involves both electron and proton transfer in the medium.