1. Introduction: What are Coordination Compounds?

Coordination compounds are a special class of compounds in which a central metal atom or ion is bonded to a surrounding group of molecules or ions, known as ligands. This bonding is a type of Lewis acid-base interaction, where the ligands donate lone pairs of electrons to the central metal. These compounds are also often referred to as complexes or complex compounds.

A key characteristic is that they retain their identity in solution. For example, when $[Cu(NH_3)_4]SO_4$ dissolves in water, the $[Cu(NH_3)_4]^{2+}$ complex ion remains intact.

2. Key Terminology

• Central Metal Atom/Ion: The metal atom or ion (typically a transition metal) to which a fixed number of ligands are attached. In $[Co(NH_3)_6]Cl_3$, the central metal is $Co^{3+}$.
• Ligands: The atoms, ions, or molecules that donate a pair of electrons to the central metal atom to form a coordinate bond. Ligands can be neutral ($H_2O$, $NH_3$, $CO$) or charged ($Cl^{-}$, $CN^{-}$, $OH^{-}$).
• Coordination Number: The number of donor atoms of the ligands that are directly bonded to the central metal atom. In $[Co(NH_3)_6]Cl_3$, the coordination number of cobalt is 6.
• Coordination Sphere: The central metal atom and the ligands directly attached to it. It is usually enclosed in square brackets. In $[Co(NH_3)_6]Cl_3$, the coordination sphere is $[Co(NH_3)_6]^{3+}$.
• Coordination Polyhedron: The spatial arrangement of the ligand donor atoms around the central metal atom. Common geometries include octahedral (coordination number 6), tetrahedral (coordination number 4), and square planar (coordination number 4).
• Oxidation State of the Central Atom: The charge the central metal atom would have if all ligands were removed along with their electron pairs. In $[Co(NH_3)_6]Cl_3$, the ligands ($NH_3$) are neutral, and the three chloride ions outside the coordination sphere have a total charge of $-3$, so the charge of the complex ion must be $+3$. Therefore, the oxidation state of the cobalt is $+3$.

Classification of Ligands

• Monodentate: Ligands with a single donor atom (e.g., $Cl^{-}$, $H_2O$, $NH_3$).
• Polydentate (Chelating): Ligands with two or more donor atoms. These form ring-like structures called chelates.
  • Bidentate: Two donor atoms (e.g., ethylenediamine, $en$).
  • Hexadentate: Six donor atoms (e.g., EDTA, ethylenediaminetetraacetate).
• Ambidentate: Ligands that can coordinate to the central atom through two different atoms (e.g., $NO_2^{-}$, which can bond via nitrogen or oxygen; $SCN^{-}$, which can bond via sulfur or nitrogen).

3. Werner’s Theory of Coordination Compounds

This was the first successful explanation of bonding in coordination compounds. Its main postulates are:

• Metal atoms exhibit two types of valencies:
  • Primary Valency: Corresponds to the oxidation state of the metal. It is ionizable and satisfied by negative ions.
  • Secondary Valency: Corresponds to the coordination number. It is non-ionizable and satisfied by ligands (neutral molecules or negative ions).
• Every metal has a fixed secondary valency.
• The secondary valencies are directed in space, leading to a definite geometry for the complex.

4. Nomenclature of Coordination Compounds

The naming of coordination compounds follows a specific set of IUPAC rules:

1. Name the cation first, followed by the anion.
2. Within the coordination sphere, name the ligands alphabetically before the central metal.
3. Ligand names:
   • Negative ligands end in ‘-o’ (e.g., $Cl^{-}$ is chloro, $CN^{-}$ is cyano, $OH^{-}$ is hydroxo).
   • Neutral ligands have special names (e.g., $H_2O$ is aqua, $NH_3$ is ammine, $CO$ is carbonyl).
4. Use prefixes ‘di-‘, ‘tri-‘, ‘tetra-‘, etc. for multiple simple ligands. For complex ligands (e.g., $en$), use ‘bis-‘, ‘tris-‘, ‘tetrakis-‘.
5. Name the central metal atom:
   • If the complex ion is a cation, use the normal name of the metal (e.g., Cobalt, Platinum).
   • If the complex ion is an anion, the metal name ends in ‘-ate’ (e.g., Cobaltate, Platinate).
6. The oxidation state of the central metal is indicated by a Roman numeral in parentheses.

Example:

• $[Co(NH_3)_6]Cl_3$: Hexaamminecobalt(III) chloride
• $K_4[Fe(CN)_6]$: Potassium hexacyanoferrate(II)

5. Isomerism in Coordination Compounds

Isomers are compounds that have the same chemical formula but different arrangements of atoms.

Structural Isomerism

• Ionization Isomerism: Occurs when the counter ion and a ligand inside the coordination sphere swap places.

  Example: $[Co(NH_3)_5Br]SO_4$ and $[Co(NH_3)_5SO_4]Br$.

• Hydrate Isomerism: A specific type of ionization isomerism where water acts as a ligand or a solvent molecule outside the coordination sphere.

  Example: $[Cr(H_2O)_6]Cl_3$ (violet) and $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$ (green).

• Linkage Isomerism: Arises from ambidentate ligands bonding in different ways.

  Example: $[Co(NH_3)_5(NO_2)]^{2+}$ (nitro) and $[Co(NH_3)_5(ONO)]^{2+}$ (nitrito).

• Coordination Isomerism: When both the cation and anion are complex ions, and the ligands are swapped between them.

  Example: $[Co(NH_3)_6][Cr(CN)_6]$ and $[Cr(NH_3)_6][Co(CN)_6]$.

Stereoisomerism

• Geometrical Isomerism (cis-trans): Arises when the ligands occupy different positions around the central metal atom.
  • cis-isomer: Identical ligands are adjacent to each other.
  • trans-isomer: Identical ligands are opposite to each other.

  Common in square planar ($[Ma_2b_2]$) and octahedral ($[Ma_4b_2]$) complexes.

• Optical Isomerism: Occurs when the isomers are non-superimposable mirror images of each other (enantiomers). This happens in complexes that lack a plane of symmetry.

  Common in octahedral complexes with bidentate ligands, such as $[Co(en)_3]^{3+}$.

6. Bonding Theories

These theories explain the properties (color, magnetism, geometry) of coordination compounds.

a) Valence Bond Theory (VBT)

VBT, proposed by Linus Pauling, explains bonding by considering the overlap of the central metal ion’s atomic orbitals with the ligand’s orbitals.

Key Principles:

1. The central metal ion forms a number of empty hybrid orbitals equal to its coordination number.
2. The ligands have at least one orbital containing a lone pair of electrons.
3. A coordinate bond is formed by the overlap of an empty metal hybrid orbital with a filled ligand orbital.
4. The geometry of the complex is determined by the type of hybridization. For example:
   • Coordination Number 4: $sp^3$ hybridization leads to a tetrahedral geometry. $dsp^2$ hybridization leads to a square planar geometry.
   • Coordination Number 6: $sp^3d^2$ hybridization leads to an octahedral geometry.
5. The magnetic properties are determined by the presence of unpaired electrons in the metal’s d-orbitals.
6. Inner and Outer Orbital Complexes: VBT distinguishes between complexes where the metal uses inner d-orbitals (e.g., $(n-1)d$ orbitals for $dsp^2$ or $d^2sp^3$) and those that use outer d-orbitals (e.g., $nd$ orbitals for $sp^3d^2$).

Limitations:

• It does not explain the characteristic colors of coordination compounds.
• It provides only a qualitative explanation of magnetic properties.
• It does not predict the geometries of all complexes correctly.

b) Crystal Field Theory (CFT)

CFT is a more advanced, purely electrostatic model that explains bonding based on the interaction between the central metal ion’s d-electrons and the ligands’ electrostatic fields.

Key Principles:

1. The bond between the metal and the ligands is assumed to be purely ionic. Ligands are treated as point charges (for anions) or point dipoles (for neutral molecules).
2. In a free metal ion, the five d-orbitals ($d_{xy}$, $d_{yz}$, $d_{xz}$, $d_{x^2-y^2}$, and $d_{z^2}$) are degenerate (have the same energy).
3. When ligands approach the central metal, the electrostatic repulsion between the ligand’s lone pairs and the metal’s d-electrons causes the d-orbitals to split into different energy levels. This is called crystal field splitting.
4. Octahedral Splitting: In an octahedral field, the ligands approach along the axes. The $d_{x^2-y^2}$ and $d_{z^2}$ orbitals point directly at the ligands and are raised to a higher energy level (called the $e_g$ set). The $d_{xy}$, $d_{yz}$, and $d_{xz}$ orbitals lie between the axes and are lowered in energy (called the $t_{2g}$ set). The energy difference is the crystal field splitting energy ($\Delta_o$).
5. Tetrahedral Splitting: In a tetrahedral field, the ligands approach between the axes, so the splitting pattern is inverted. The $t_{2}$ set is at a higher energy, and the $e$ set is at a lower energy.
6. High-Spin vs. Low-Spin: The filling of the split d-orbitals depends on the relative magnitudes of the crystal field splitting energy ($\Delta_o$) and the pairing energy ($P$).
   • Strong-field ligands (e.g., $CN^{-}$, $CO$) cause a large splitting ($\Delta_o > P$). Electrons will fill the lower energy orbitals completely before pairing up, leading to low-spin complexes.
   • Weak-field ligands (e.g., $Cl^{-}$, $F^{-}$, $H_2O$) cause a small splitting ($\Delta_o < P$). Electrons will fill the orbitals singly before pairing up, leading to high-spin complexes.
7. Explanation of Color: The color of a complex is due to the absorption of light, which excites an electron from a lower energy d-orbital to a higher energy d-orbital (d-d transition). The absorbed light corresponds to the energy gap, $\Delta_o$. The color we see is the complementary color of the light absorbed.
8. Spectrochemical Series: The spectrochemical series is an experimentally determined list of ligands ordered by their ability to cause crystal field splitting. A short list is:

   $I^{-} < Br^{-} < Cl^{-} < F^{-} < H_2O < NH_3 < en < CN^{-} < CO$

   (Weak field $\rightarrow$ Strong field)

Limitations:

• CFT is a purely electrostatic model and does not account for the partial covalent character of the metal-ligand bond.