Coordination Chemistry

Introduction and Scope

Coordination chemistry deals with compounds in which a central atom or ion (usually a metal from the d‑block, but also main group or f‑block elements) is surrounded by other atoms, ions, or molecules called ligands. These ligands donate electron pairs to the central atom to form coordinate (dative) bonds and build up well‑defined three‑dimensional structures.

Within “Inorganic and Coordination Chemistry” you have already met main group and transition elements and some of their simple compounds. Coordination chemistry focuses specifically on:

• How and why metals bind ligands
• The structural patterns that result
• How these structures influence reactivity and properties

Details of naming, synthesis, bonding, stability, and applications of complexes are treated in the subsequent subchapters. Here the emphasis is on what broadly characterizes coordination compounds, why they are important, and which fundamental ideas are specific to this area.

What Makes a Compound a Coordination Compound?

Coordination compounds (or complexes) are characterized by:

• A central atom or ion, often a transition metal cation such as $ \text{Fe}^{2+}, \text{Cu}^{2+}, \text{Co}^{3+}$.
• One or more ligands, each providing at least one lone pair of electrons that can be donated to the metal.
• Coordinate bonds (also called dative covalent bonds) in which both bonding electrons come from the ligand.
• A well‑defined coordination sphere, described by the coordination number and geometry.

A typical example is the hexaaquacopper(II) complex:
$$
[\text{Cu(H}_2\text{O)}_6]^{2+}
$$
here:

• Central ion: $ \text{Cu}^{2+} $
• Ligands: six water molecules
• Coordination number: 6
• Geometry: approximately octahedral

The entire bracketed part with its charge is the complex ion; if it is associated with counterions (e.g. $ \text{SO}_4^{2-} $), the neutral whole is a coordination compound.

Central Metal Atoms and Ions

Although any atom capable of accepting electron pairs can in principle form complexes, coordination chemistry is dominated by:

• Transition metals (d‑block): e.g. Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn
• Lanthanides and actinides (f‑block)
• Some heavier main group elements (e.g. Al, Sn, Pb, Sb, Bi)

They are especially suited because:

1. They have vacant d (or f) orbitals able to accept electron pairs.
2. They often exist in several oxidation states, which allows a rich variety of complexes.
3. Their ionic radii and charges are such that they strongly attract and hold ligands.

In a given oxidation state, a metal’s size and charge density influence:

• How many ligands it can accommodate (coordination number)
• How strongly it binds those ligands
• Its preference for certain ligand types and geometries

Ligands: Types and Binding Modes

Ligands are ions or neutral molecules with at least one lone pair of electrons that they can donate to a metal center.

Donor Atoms

The atom of the ligand that directly bonds to the metal is the donor atom. Common donor atoms in coordination chemistry include:

• Nitrogen: e.g. $ \text{NH}_3$, $ \text{NH}_2\text{CH}_2\text{CH}_2\text{NH}_2$, pyridine
• Oxygen: e.g. $ \text{H}_2\text{O}$, $ \text{OH}^-$, $ \text{CO}_3^{2-}$, carboxylates
• Halogens: e.g. $ \text{Cl}^-, \text{Br}^-, \text{I}^- $
• Sulfur and other chalcogens: e.g. $ \text{HS}^-, \text{RS}^-, \text{S}^{2-} $
• Carbon: in ligands like carbon monoxide ($\text{CO}$) or cyanide ($\text{CN}^-$)

The nature of the donor atom has a strong impact on the electronic and sometimes magnetic properties of the metal center.

Denticity

The denticity of a ligand tells you how many donor atoms from the same ligand bind to the same metal ion.

• Monodentate ligands: bind through a single donor atom
• Examples: $ \text{H}_2\text{O}$, $ \text{NH}_3$, $ \text{Cl}^-, \text{CN}^- $
• Bidentate ligands: bind through two donor atoms
• Example: ethylenediamine (en, $ \text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2$)
• Polydentate ligands (tridentate, tetradentate, etc.): bind through three or more donor atoms
• Example: EDTA$^{4-}$ is hexadentate (up to six donor atoms)

Ligands that attach through more than one donor atom from the same molecule are also called chelate ligands, and the complexes formed are chelate complexes.

Chelation

When a polydentate ligand binds to a metal at more than one site, it forms a chelate ring with the metal ion at the center. For example, in
$$
[\text{Ni(en)}_3]^{2+}
$$
each en wraps around the nickel and binds through two nitrogens, giving three five‑membered chelate rings.

Chelation typically:

• Increases the stability of complexes compared to similar complexes with only monodentate ligands.
• Can greatly influence solubility, biological activity, and reactivity.

The preference for chelate formation and its consequences are explored further in the subchapter on stability.

Bridging and Terminal Ligands

Depending on how they connect metals, ligands can be:

• Terminal: bound to only one metal center, e.g. in $[\text{Co(NH}_3)_6]^{3+}$ all six NH3 are terminal.
• Bridging: connecting two or more metal centers.
• Common bridging ligands: $ \mu$‑$ \text{OH}^-, \mu$‑$ \text{O}^{2-}, \mu$‑$ \text{Cl}^-, \mu$‑carboxylates, $ \mu$‑cyanide, etc.
• Bridging ligands are indicated with the prefix μ (mu) in structural descriptions and names.

Bridging ligands are important because they give rise to polynuclear complexes, in which several metal centers are linked through shared ligands, leading to extended structures and cooperative properties.

Coordination Number and Geometry

The coordination number (CN) is the number of donor atoms directly bonded to the central metal ion.

Typical coordination numbers and their most common geometries include:

• CN = 2: linear
• Example geometry: $ \text{M(L)}_2 $ with a 180° bond angle
• CN = 4: tetrahedral or square planar
• Many $ \text{Zn}^{2+}, \text{Cu}^{2+}$ complexes are tetrahedral
• Many $d^8$ complexes of $ \text{Ni}^{2+}, \text{Pd}^{2+}, \text{Pt}^{2+}$ are square planar
• CN = 6: octahedral (very common)
• Example: $ [\text{Fe(H}_2\text{O)}_6]^{2+} $
• Higher CN (7, 8, 9...): less common, with more varied geometries such as pentagonal bipyramidal, square antiprismatic, etc., seen more often in lanthanide and actinide chemistry.

The actual geometry results from a balance between:

• The metal ion size and electronic structure
• The size and steric demands of the ligands
• The ligand field preferences (explored in more detail under bonding in complexes)

Coordination number and geometry largely determine:

• The shape of the complex
• Possible isomerism (different spatial arrangements)
• Many physical properties, such as color and magnetism

Inner and Outer Coordination Sphere

In a typical crystalline coordination compound, one distinguishes:

• The inner (primary) coordination sphere: the metal center plus all ligands directly bonded to it.
• The outer (secondary) sphere: counterions and sometimes solvent molecules that balance charge and interact electrostatically or via hydrogen bonding, but are not directly coordinated.

For example, in the salt
$$
[\text{Co(NH}_3)_6]\text{Cl}_3
$$

• Inner sphere: $[\text{Co(NH}_3)_6]^{3+}$ (Co and six NH$_3$ ligands)
• Outer sphere: three $ \text{Cl}^- $ ions

This distinction is useful when discussing:

• Conductivity in solution (which ions dissociate)
• Reactivity patterns (which ligands are labile and which are firmly bound)
• Ion exchange and precipitation reactions used in analysis and synthesis

Special Role of Transition Metals in Coordination Chemistry

Although main group elements do form complexes, coordination chemistry is particularly rich for transition metals because:

• Their partially filled d orbitals enable a variety of bonding interactions.
• They can adopt multiple oxidation states, allowing complexes of the same metal with different charges, colors, and reactivities.
• They frequently exhibit paramagnetism (unpaired electrons), leading to interesting magnetic behavior.
• Their complexes often show intense colors and characteristic spectra due to electronic transitions within the d orbitals and between metal and ligands.

These aspects underpin the importance of coordination complexes in:

• Catalysis (both industrial and biological, e.g. metalloenzymes)
• Materials science (magnetic and optical materials)
• Bioinorganic chemistry (hemoglobin, chlorophyll, vitamin B$_{12}$)
• Analytical chemistry (colorimetric determination, complexometric titrations)

How the metal and ligand orbitals combine, and how this leads to color, magnetism, and reactivity, is considered in detail later in the bonding and properties subchapters.

Fundamental Notions of Reactivity in Coordination Chemistry

Even before looking in detail at reaction mechanisms, some basic patterns of reactivity are characteristic for coordination compounds:

• Ligand substitution: one ligand in the coordination sphere is replaced by another.
• Example (simplified):
  $$
  [\text{Co(H}_2\text{O)}_6]^{2+} + 6\text{NH}_3 \rightarrow [\text{Co(NH}_3)_6]^{2+} + 6\text{H}_2\text{O}
  $$
• Redox changes at the metal center: the oxidation state of the metal changes, often accompanied by ligand changes or structural rearrangements.
• Inner‑sphere vs outer‑sphere processes: depending on whether the ligand bridges between metal centers or not.
• Association and dissociation: changes in coordination number when ligands add to or leave the inner sphere.

The rate at which these changes occur (lability vs inertness) and the thermodynamic stability of complexes are key themes in coordination chemistry, addressed in the subchapters on stability and bonding.

Coordination Chemistry in Nature and Technology

Coordination compounds are omnipresent in natural and artificial systems. A broad overview includes:

• Bioinorganic systems:
• Heme in hemoglobin and myoglobin: an Fe–porphyrin complex for O$_2$ transport and storage.
• Chlorophyll: a Mg–porphyrin complex involved in photosynthesis.
• Vitamin B$_{12}$: a Co complex involved in enzyme catalysis.
• Industrial catalysts:
• Many homogeneous catalysts are transition metal complexes (e.g. for polymerization of alkenes, hydroformylation, hydrogenation).
• Pigments and dyes:
• Complexes of Fe, Cu, Co, and others give intense and stable colors.
• Medicinal chemistry:
• Platinum complexes used in cancer therapy.
• Chelating agents like EDTA used for heavy metal detoxification and in analytical applications.

These applications all rely on the ability of metal centers to bind, transform, and release ligands in a controlled way—core ideas of coordination chemistry.

Summary of Key Concepts in Coordination Chemistry

To orient yourself for the more detailed subchapters, the essential features of coordination chemistry are:

• A central metal atom or ion bound to one or more ligands via coordinate bonds, forming a coordination sphere with a specific coordination number and geometry.
• Ligands are classified by donor atoms, denticity, and binding mode (terminal or bridging), with chelation playing a central role in stability.
• The coordination number and geometry of complexes are determined by the metal and ligands and underlie their structures and isomerism.
• The distinction between inner and outer coordination spheres is fundamental to understanding properties and reactions.
• Transition metals dominate coordination chemistry, providing diverse structures, colors, magnetic properties, and reactivities that are exploited in nature and technology.

The following subchapters build on these foundations to discuss systematically how complexes are named, how they are synthesized, what makes them more or less stable, how bonding in them is understood, and why they are so important in many areas of chemistry and everyday life.