8.1 Nomenclature and Structural Types of Coordination Compounds

Coordination Compounds and Ligands

Coordination compounds consist of a central metal atom or ion surrounded by molecules or ions called ligands. The metal center acts as a Lewis acid (electron pair acceptor), while each ligand acts as a Lewis base (electron pair donor). This donor-acceptor interaction is the fundamental bonding concept behind all coordination chemistry.

Components of Coordination Compounds

• The central metal atom/ion is typically a transition metal with empty or partially filled d orbitals available to accept electron pairs.
• A ligand donates one or more lone pairs to the metal center.
• The metal plus its directly attached ligands form the coordination sphere, written inside square brackets (e.g., $[\text{Co(NH}_3\text{)}_6]^{3+}$).
• A chelate forms when a single polydentate ligand binds to the metal through two or more donor atoms, creating a ring structure. Chelates are more thermodynamically stable than comparable complexes with monodentate ligands, a phenomenon known as the chelate effect. This extra stability arises primarily from a favorable entropy change: replacing several monodentate ligands with one polydentate ligand increases the total number of free particles in solution.

Types of Ligands

Ligands are classified by denticity, the number of donor atoms they use to bind the metal.

Nomenclature and Oxidation State

IUPAC Nomenclature Rules

Naming coordination compounds follows a specific set of conventions. Here's the process:

1. Cation before anion. For ionic coordination compounds, name the cation first, then the anion, regardless of which one contains the metal complex.

2. Ligands before metal. Within the coordination sphere, name all ligands before the metal.

3. Alphabetical order. List ligands in alphabetical order by ligand name. The alphabetical order ignores the multiplying prefix (so "dichloro" is alphabetized under "c," not "d").

4. Multiplying prefixes. Use Greek prefixes (di-, tri-, tetra-, penta-, hexa-) to indicate the number of each simple ligand. For more complex or already-prefixed ligand names (like ethylenediamine), use bis-, tris-, tetrakis- instead, and enclose the ligand name in parentheses.

5. Ligand name endings. Anionic ligands get an -o ending (chloro, bromo, cyano, oxalato). Neutral ligands keep their usual name, with two important exceptions: $\text{H}_2\text{O}$ is called aqua and $\text{NH}_3$ is called ammine (two m's).

6. Metal name. If the complex is an anion, add -ate to the metal name (sometimes using the Latin root: ferrate for iron, cuprate for copper, argentate for silver). If the complex is a cation or neutral, use the normal metal name.

7. Oxidation state. Indicate the metal's oxidation state in Roman numerals in parentheses immediately after the metal name, with no space.

Example: $[\text{CoCl}_2\text{(en)}_2]\text{Cl}$

• Coordination sphere: $[\text{CoCl}_2\text{(en)}_2]^+$, counter ion: $\text{Cl}^-$
• Ligands alphabetically: chloro (×2), ethylenediamine (×2)
• Cobalt oxidation state: overall charge is +1, two $\text{Cl}^-$ contribute −2, en is neutral → Co is +3
• Name: dichlorobis(ethylenediamine)cobalt(III) chloride

Determining Oxidation State and Coordination Number

The oxidation state of the metal is found by balancing charges:

Neutral ligands ($\text{NH}_3$, $\text{H}_2\text{O}$, en) contribute 0. Anionic ligands contribute their ionic charge ($\text{Cl}^-$ = −1, $\text{C}_2\text{O}_4^{2-}$ = −2).

The coordination number is the total number of donor atoms directly bonded to the metal. Each monodentate ligand contributes 1, each bidentate ligand contributes 2, and so on. For example, $[\text{Co(en)}_3]^{3+}$ has three bidentate ligands, giving a coordination number of 6.

Isomers

Isomers are compounds with the same chemical formula but different arrangements of atoms. In coordination chemistry, the main categories are structural isomers (different connectivity) and stereoisomers (same connectivity, different spatial arrangement).

Geometric Isomers

Geometric isomers differ in the relative positions of ligands around the metal center. They occur in complexes where ligand positions are not all equivalent.

Square planar complexes (e.g., $[\text{Pt(NH}_3\text{)}_2\text{Cl}_2]$):

• Cis: the two identical ligands sit adjacent (90° apart)
• Trans: the two identical ligands sit opposite (180° apart)

These two isomers can have very different properties. Cisplatin (cis-$[\text{Pt(NH}_3\text{)}_2\text{Cl}_2]$) is a widely used anticancer drug, while transplatin (the trans isomer) is therapeutically inactive.

Octahedral complexes of the type $[\text{MA}_2\text{B}_4]$ also show cis/trans isomerism. For complexes of the type $[\text{MA}_3\text{B}_3]$, the relevant terms are:

• Fac (facial): three identical ligands occupy one triangular face of the octahedron
• Mer (meridional): three identical ligands span a plane that passes through the metal center, so they don't all sit on the same face

Optical and Linkage Isomers

Optical isomers (enantiomers) are non-superimposable mirror images of each other. They rotate plane-polarized light in equal but opposite directions. In coordination chemistry, optical isomerism commonly arises in octahedral complexes with bidentate or polydentate ligands, such as $[\text{Co(en)}_3]^{3+}$. The two mirror-image forms are labeled $\Delta$ and $\Lambda$. Tetrahedral complexes can also be chiral if they have four different ligands, though this is less common.

Linkage isomers arise from ambidentate ligands, which can coordinate through more than one donor atom. The nitrite ion ($\text{NO}_2^-$) is the classic example: it can bind through nitrogen (nitro, $\kappa\text{N}$) or through oxygen (nitrito, $\kappa\text{O}$). Another example is thiocyanate ($\text{SCN}^-$), which can bind through S (thiocyanato) or N (isothiocyanato).

Structural Types

Octahedral Complexes

Octahedral geometry is the most common arrangement for transition metal coordination compounds. Six ligands sit at the vertices of a regular octahedron, with the metal at the center and all metal-ligand bond angles at 90°. The coordination number is 6.

Examples include $[\text{Co(NH}_3\text{)}_6]^{3+}$, $[\text{Fe(CN)}_6]^{4-}$, and $[\text{Cr(H}_2\text{O)}_6]^{3+}$. Octahedral complexes can display both geometric isomerism (cis/trans or fac/mer) and optical isomerism (when chelating ligands create a chiral arrangement).

Square Planar and Tetrahedral Complexes

Square planar complexes have four ligands arranged in a plane around the metal, with 90° bond angles. This geometry is strongly favored by $d^8$ metal ions such as $\text{Pt}^{2+}$, $\text{Pd}^{2+}$, $\text{Ni}^{2+}$ (in strong-field cases), and $\text{Au}^{3+}$. Square planar complexes can show geometric (cis/trans) isomerism. They generally do not exhibit optical isomerism because the plane of symmetry in the square plane makes mirror images superimposable.

Tetrahedral complexes also have four ligands, but arranged at the corners of a tetrahedron with approximately 109.5° bond angles. This geometry is common for metals with $d^0$ or $d^{10}$ configurations (e.g., $\text{Zn}^{2+}$, $\text{Cu}^+$) and for complexes with bulky or weak-field ligands. Tetrahedral complexes do not show geometric isomerism because all four positions are equivalent. However, a tetrahedral complex with four different ligands is chiral and can exhibit optical isomerism.