Ligand

In coordination chemistry, a ligand is an ion or molecule containing at least one functional group capable of binding to a central metal atom or ion to form a coordination complex. Ligand–metal interactions typically involve the formal donation of electron pairs from the ligand to the metal centre, most commonly described using the framework of Lewis acid–base chemistry, in which ligands act as Lewis bases and metals as Lewis acids. In most complexes the ligands donate electrons to the metal, though a few unusual systems involve ligands functioning as Lewis acids.
Ligands exert profound influence on the behaviour of metal centres. They determine the reactivity, substitution processes, spectroscopic properties, redox characteristics, and overall geometry of the resulting complexes. The selection of appropriate ligands is therefore crucial in fields including bioinorganic chemistry, medicinal chemistry, homogeneous catalysis, and environmental chemistry.

Historical Background

Although compounds such as Prussian blue and copper(II) sulfate were studied in the early nineteenth century, the conceptual understanding of coordination complexes emerged largely through the work of Alfred Werner. Werner’s coordination theory explained the structures, stoichiometries, and isomerism of cobalt(III) and chromium(III) compounds by proposing the existence of fixed coordination numbers and spatial arrangements of ligands. He introduced the idea that many complexes possess an octahedral geometry with six surrounding ligands. Werner and Carl Somiesky were the first to employ the term “ligand” in relation to silicon chemistry.
Werner’s resolution of the cobalt complex hexol into optical isomers provided evidence that chirality could arise from metal coordination rather than solely from carbon-based compounds, marking a major advance in inorganic stereochemistry.

Nature of Metal–Ligand Bonding

The metal–ligand bond can range from highly covalent to predominantly ionic, with bond orders varying from one to three. Molecular orbital theory provides a detailed picture of these interactions. When a ligand binds to a metal ion, its highest occupied molecular orbital (HOMO) donates electron density into the metal’s lowest unoccupied molecular orbital (LUMO). Conversely, some ligands also accept electron density from the metal via backbonding, whereby a filled metal orbital donates into a low-energy ligand LUMO. Carbon monoxide is the classic π-acceptor ligand exhibiting strong backbonding.
Ligands with filled π orbitals can act as π donors, altering the electronic structure and stability of complexes.

Strong-Field and Weak-Field Ligands

In an octahedral complex, the five metal d orbitals split into two energy sets:

• A lower-energy triplet: dxy,dxz,dyzd_{xy}, d_{xz}, d_{yz}dxy​,dxz​,dyz​
• A higher-energy doublet: dz2,dx2−y2d_{z^2}, d_{x^2-y^2}dz2​,dx2−y2​

The energy gap between these sets is the octahedral splitting parameter, Δo\Delta_oΔo​. Ligands arranged by increasing Δo\Delta_oΔo​ form the spectrochemical series, which is largely invariant across metal ions. Strong-field ligands (e.g., CN⁻, CO) produce a large Δo\Delta_oΔo​, favouring electron pairing (low-spin complexes). Weak-field ligands (e.g., halides) produce a smaller Δo\Delta_oΔo​, often resulting in high-spin configurations governed by Hund’s rule.
In tetrahedral complexes, only four ligands interact with the metal, producing a smaller splitting parameter, Δt\Delta_tΔt​, with reversed orbital ordering.
The distribution of electrons among the split d orbitals influences the optical absorption, colour, magnetic properties, and chemical reactivity of complexes. Tanabe–Sugano diagrams describe these effects by illustrating the relative energies of electronic states under different ligand field strengths.

Classification of Ligands

Ligands are categorised in various ways, including by charge, size, bulk, coordinating atom(s), electron donation, and binding mode.
A widely used classification distinguishes between L and X ligands:

• L ligands (neutral 2-electron donors), typical of Lewis bases such as amines, phosphines, carbon monoxide, alkenes, dihydrogen, and hydrocarbons involved in agostic interactions.
• X ligands (anionic 1-electron donors), including halides, pseudohalides, hydride, and alkyl groups.

In organometallic chemistry, the Covalent Bond Classification (CBC) system refines these categories by defining L, X, and Z ligands (the latter donating 0 electrons).

Denticity and Chelation

Many ligands bind via more than one donor atom and are therefore polydentate.Denticity refers to the number of separate donor atoms bound to a metal centre:

• Monodentate ligands coordinate through a single atom.
• Bidentate ligands bind through two distinct donor atoms—for example ethylenediamine, whose two amine groups are separated by an ethylene linker.
• Tridentate, tetradentate, and higher polydentate ligands coordinate through additional donor sites.

A ligand that binds through multiple sites often forms a chelate, enhancing complex stability through the chelate effect. The classic ligand EDTA is hexadentate, capable of coordinating through six donor atoms to fully encircle certain metal ions.
Denticity is denoted by κn\kappa^nκn, where nnn is the number of coordinated atoms. For example, a cobalt complex with three bidentate ethylenediamine ligands may be written as Co(en)3_33​³⁺, with each en acting as κ2\kappa^2κ2.

Hapticity

Separate from denticity, hapticity refers to ligands such as cyclopentadienyl that bind through a contiguous set of atoms, often interacting with the metal via delocalised π systems. It is denoted by ηn\eta^nηn.

Influence on Complex Properties

Ligands profoundly shape the chemical behaviour of metal centres. They influence:

• Ligand substitution and reaction rates
• Redox reactivity of the metal
• Stability of oxidation states
• Geometric preference (e.g., octahedral, square planar, tetrahedral)
• Spectroscopic signatures, including colour and magnetic properties