Chemistry: Class 121 — Chapter 5: Coordination Compounds

Alfred Werner formulated his theory of coordination compounds based on the concepts of primary and secondary valencies:

• Primary Valency: Ionisable, satisfied by negative ions, and corresponds to the oxidation state of the metal ion.
• Secondary Valency: Non-ionisable, satisfied by neutral molecules or negative ions, and corresponds to the coordination number of the metal. It is fixed for a given metal.
• The ions or groups bound by secondary linkages have characteristic spatial arrangements corresponding to different coordination numbers (coordination polyhedra).

• Double Salts: Dissociate completely into simple constituent ions when dissolved in water. Examples include carnallite ($KCl \cdot MgCl_2 \cdot 6H_2O$), Mohr's salt ($FeSO_4 \cdot (NH_4)_2SO_4 \cdot 6H_2O$), and potash alum ($KAl(SO_4)_2 \cdot 12H_2O$).
• Complex Compounds: Do not dissociate completely into simple constituent ions. The complex entity (e.g., $[Fe(CN)_6]^{4-}$) remains intact in solution.

• Coordination Entity: A central metal atom or ion bonded to a fixed number of ions or molecules. Example: $[CoCl_3(NH_3)_3]$.
• Central Atom/Ion: The atom or ion to which a fixed number of ligands are bound in a definite geometrical arrangement. These act as Lewis acids.

The ions or molecules bound to the central atom/ion are called ligands. They are classified based on the number of donor atoms:

• Unidentate: Bound through a single donor atom. Examples: $Cl^-$, $H_2O$, $NH_3$.
• Didentate: Bound through two donor atoms. Examples: ethane-1,2-diamine ($en$) and oxalate ($C_2O_4^{2-}$).
• Polydentate: Bound through multiple donor atoms. Example: Ethylenediaminetetraacetate ion ($EDTA^{4-}$), which is a hexadentate ligand.
• Chelate Ligand: A di- or polydentate ligand that uses two or more donor atoms simultaneously to bind to a single metal ion, forming a ring structure. Chelate complexes tend to be exceptionally stable (chelate effect).
• Ambidentate Ligand: A ligand that can coordinate through more than one donor atom. Examples: $NO_2^-$ (can bind via N as nitrito-N or via O as nitrito-O) and $SCN^-$ (can bind via S as thiocyanato-S or via N as isothiocyanato-N).

• Coordination Number (CN): The total number of ligand donor atoms directly bonded to the metal ion. Pi bonds are not counted.
• Coordination Sphere: The central atom/ion and its attached ligands enclosed in square brackets. The ionisable species outside the brackets are called counter ions.
• Coordination Polyhedron: The spatial arrangement of ligand atoms directly attached to the central atom/ion (e.g., octahedral, tetrahedral, square planar).

• Homoleptic Complexes: Complexes in which the metal is bound to only one kind of donor group. Example: $[Co(NH_3)_6]^{3+}$.
• Heteroleptic Complexes: Complexes in which the metal is bound to more than one kind of donor group. Example: $[Co(NH_3)_4Cl_2]^+$.

1. The central atom is listed first.
2. Ligands are listed in alphabetical order, irrespective of their charge.
3. Polydentate ligands are also listed alphabetically. For abbreviated ligands, the first letter of the abbreviation is used.
4. The entire coordination entity is enclosed in square brackets.
5. There should be no space between ligands and the metal within the coordination sphere.

1. The cation is named first in both positively and negatively charged coordination entities.
2. Ligands are named in alphabetical order before the name of the central atom/ion.
3. Anionic ligands end in -o (e.g., chlorido, cyanido). Neutral and cationic ligands keep their names, with exceptions: $H_2O$ (aqua), $NH_3$ (ammine), $CO$ (carbonyl), and $NO$ (nitrosyl).
4. Prefixes di-, tri-, tetra- indicate the number of individual ligands. When the ligand name includes a numerical prefix, bis-, tris-, tetrakis- are used in parentheses.
5. The oxidation state of the metal is indicated by a Roman numeral in parentheses.
6. If the complex ion is an anion, the metal name ends with the suffix -ate (e.g., cobaltate, ferrate). For cationic/neutral complexes, the metal is named as the element.

• Geometrical Isomerism: Occurs in heteroleptic complexes due to different spatial arrangements of ligands. Common in coordination numbers 4 (square planar) and 6 (octahedral). Examples include cis and trans isomers, as well as facial (fac) and meridional (mer) isomers in $[Ma_3b_3]$ complexes.
• Optical Isomerism: Occurs when molecules are non-superimposable mirror images (chiral). The two forms are called dextro ($d$) and laevo ($l$). Common in octahedral complexes containing didentate ligands (e.g., $[Co(en)_3]^{3+}$).

• Linkage Isomerism: Arises in complexes containing ambidentate ligands. Example: $[Co(NH_3)_5(NO_2)]Cl_2$ and $[Co(NH_3)_5(ONO)]Cl_2$.
• Coordination Isomerism: Arises from the interchange of ligands between cationic and anionic entities of different metal ions. Example: $[Co(NH_3)_6][Cr(CN)_6]$ and $[Cr(NH_3)_6][Co(CN)_6]$.
• Ionisation Isomerism: Occurs when the counter ion in a complex salt is itself a potential ligand and can displace a ligand. Example: $[Co(NH_3)_5(SO_4)]Br$ and $[Co(NH_3)_5Br]SO_4$.
• Solvate (Hydrate) Isomerism: Differs by whether a solvent molecule (like water) is directly bonded to the metal ion or present as free solvent in the crystal lattice. Example: $[Cr(H_2O)_6]Cl_3$ (violet) and $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$ (grey-green).

According to VBT, the metal atom/ion under the influence of ligands uses its $(n-1)d$, $ns$, $np$, or $nd$ orbitals for hybridisation to yield equivalent orbitals of definite geometry:

• $sp^3$ (Coordination Number 4): Tetrahedral geometry.
• $dsp^2$ (Coordination Number 4): Square planar geometry.
• $sp^3d$ (Coordination Number 5): Trigonal bipyramidal geometry.
• $sp^3d^2$ (Coordination Number 6): Octahedral geometry (outer orbital or high-spin complex).
• $d^2sp^3$ (Coordination Number 6): Octahedral geometry (inner orbital or low-spin complex).

Magnetic properties can be predicted: complexes with unpaired electrons are paramagnetic, while those with all paired electrons are diamagnetic.

CFT is an electrostatic model that treats the metal-ligand bond as purely ionic, considering ligands as point charges. The electrostatic field of ligands lifts the degeneracy of the $d$ orbitals, causing them to split:

• Octahedral Splitting: The $d$ orbitals split into a lower energy triplet ($t_{2g}$ set: $d_{xy}$, $d_{yz}$, $d_{xz}$) and a higher energy doublet ($e_g$ set: $d_{x^2-y^2}$, $d_{z^2}$). The energy separation is denoted as $\Delta_o$.
• Spectrochemical Series: Ligands are arranged in order of increasing field strength:
  $I^- < Br^- < SCN^- < Cl^- < S^{2-} < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < edta^{4-} < NH_3 < en < CN^- < CO$.
• Weak Field Ligands: Cause small splitting ($\Delta_o < P$, where $P$ is pairing energy), resulting in high-spin complexes.
• Strong Field Ligands: Cause large splitting ($\Delta_o > P$), resulting in low-spin complexes.
• Tetrahedral Splitting: The splitting is inverted and smaller compared to octahedral splitting: $\Delta_t = \frac{4}{9}\Delta_o$. Low-spin configurations are rarely observed in tetrahedral complexes.

The characteristic colours of coordination complexes are attributed to $d-d$ transitions. When light is absorbed, an electron is excited from the lower energy $d$ orbital set ($t_{2g}$) to the higher energy set ($e_g$). The observed colour is complementary to the absorbed wavelength. In the absence of ligands, crystal field splitting does not occur, and the substance remains colourless (e.g., anhydrous $CuSO_4$ is white, while $CuSO_4 \cdot 5H_2O$ is blue).