IUPAC Nomenclature of Coordination Compounds

Coordination compounds are compounds that contain a central metal atom or ion bonded to a surrounding array of molecules or ions called ligands. Understanding their nomenclature is crucial for systematic study and communication in chemistry. The International Union of Pure and Applied Chemistry (IUPAC) has established a set of rules for naming these compounds.

Basic Principles of IUPAC Nomenclature

The naming of coordination compounds follows a sequence similar to simple ionic compounds. The cation is named before the anion. Within the coordination sphere (the central metal and its ligands), ligands are named before the central metal atom. Ligands are listed in alphabetical order, regardless of their charge or the number of ligands.

Ligand Naming

Most ligands have common names, but for IUPAC nomenclature, they are modified.

• Anionic ligands ending in 'o' replace the 'e' with 'o'. For example, 'chloride' becomes 'chlorido', 'cyanide' becomes 'cyanido', 'sulfate' becomes 'sulfato', 'nitrate' becomes 'nitrato', and 'oxide' becomes 'oxido'.
• Neutral ligands are usually named as the molecule itself, such as 'water' (aqua), 'ammonia' (ammine), 'carbon monoxide' (carbonyl), and 'nitrogen monoxide' (nitrosyl).
• Cationic ligands are rare but are named with a suffix 'ium', e.g., 'nitrosonium' (NO⁺).

Prefixes for Number of Ligands

To indicate the number of ligands of a particular type, specific prefixes are used:

• For simple ligands, the prefixes 'di-', 'tri-', 'tetra-', 'penta-', 'hexa-' are used. For example, two chloride ligands would be 'dichlorido'.
• For complex ligands that already contain prefixes or are named as units (e.g., ethylenediamine), the prefixes 'bis-', 'tris-', 'tetrakis-' are used. For example, two ethylenediamine ligands would be 'bis(ethylenediamine)'.

Naming the Central Metal Atom

The name of the central metal atom is written after the names of the ligands.

• If the coordination compound is a cation or neutral, the metal name is used as is (e.g., 'iron', 'copper', 'nickel').
• If the coordination compound is an anion, the metal name is given the suffix '-ate' (e.g., 'ferrate' for iron, 'cuprate' for copper, 'nickelate' for nickel).

Naming Complex Ions and Compounds

The rules for naming complex ions and neutral coordination compounds are as follows:

• For a complex cation, name the ligands alphabetically, followed by the metal name, and then its oxidation state in Roman numerals. Example: [Co(NH₃)₆]³⁺ is hexaamminecobalt(III).
• For a complex anion, name the ligands alphabetically, followed by the metal name with the suffix '-ate', and then its oxidation state in Roman numerals. Example: [Fe(CN)₆]^4- is hexacyanidoferrate(II).
• For neutral coordination compounds, the rules are the same as for complex cations. Example: [Ni(CO)₄] is tetracarbonylnickel(0).
• When both cation and anion are complex, name the cation first, then the anion, following all the above rules. Example: [Co(NH₃)₆][Cr(CN)₆] is hexaamminecobalt(III) hexacyanidochromate(III).

Special Cases

Bridging ligands are indicated by the prefix 'μ-' before the ligand name. If there are multiple bridging ligands, 'μ-' is repeated. For example, [Co(NH₃)₄(OH)₂Co]³⁺ can be named as tetraammine-di-μ-hydroxodicobalt(III). In cases where a ligand is attached to the metal through a specific atom, the point of attachment is indicated by a symbol. For example, 'N' for nitrogen in a nitro group (NO₂^-) attached via nitrogen, and 'O' for oxygen if attached via oxygen (nitrito-O). However, for common ligands like thiocyanato (SCN^-), 'S' is assumed if not specified, and 'isothiocyanato' is used if attached via nitrogen. The IUPAC nomenclature has evolved, and simpler forms are often preferred. For instance, hexacyanidoferrate(II) is now commonly written as ferrocyanide(II).

Isomerism in Coordination Compounds

Isomerism is the phenomenon where two or more compounds have the same molecular formula but differ in the arrangement of atoms or ions. Coordination compounds exhibit various types of isomerism, which can be broadly classified into stereoisomerism and structural isomerism.

Structural Isomerism

Structural isomerism arises due to differences in the connectivity of atoms within the coordination compound. This includes:

1. Linkage Isomerism

This type of isomerism occurs when a ligand can coordinate to the central metal atom through different donor atoms. Such ligands are called ambidentate ligands. Common examples include the nitro ligand (NO₂^-), which can coordinate through nitrogen (nitro-, -NO₂) or oxygen (nitrito-, -ONO), and the thiocyanate ligand (SCN^-), which can coordinate through sulfur (thiocyanato-, -SCN) or nitrogen (isothiocyanato-, -NCS).

Example: [Co(NH₃)₅(NO₂)]Cl₂ can exist as two isomers:

• Pentaamminenitrocobalt(III) chloride (coordination through N)
• Pentaamminenitritocobalt(III) chloride (coordination through O)

2. Solvate Isomerism (Hydrate Isomerism)

Solvate isomerism is a special case of ionization isomerism where the solvent molecule (often water) is involved in the isomerism. If the solvent is water, it is called hydrate isomerism. The solvent molecules can be coordinated to the metal ion or present as lattice water molecules.

Example: The compound CrCl₃·6H₂O can exist in three isomeric forms:

• [Cr(H₂O)₆]Cl₃ (Hexaaquachromium(III) chloride) - no free water
• [Cr(H₂O)₅Cl]Cl₂·H₂O (Pentaaquachloridochromium(III) chloride monohydrate) - one molecule of water in the lattice
• [Cr(H₂O)₄Cl₂]Cl·2H₂O (Tetraaquodichloridochromium(III) chloride dihydrate) - two molecules of water in the lattice

3. Ionization Isomerism

In ionization isomerism, the ligands and the counter ions are exchanged between the coordination sphere and the ionization sphere. The isomers differ in the ions they precipitate when treated with a reagent that reacts with the counter ion.

Example: [Co(NH₃)₄Br₂]SO₄ and [Co(NH₃)₄SO₄]Br₂ are ionization isomers. The first compound will precipitate BaSO₄ with BaCl₂ solution (indicating the presence of SO₄^2- outside the coordination sphere), while the second will precipitate AgBr with AgNO₃ solution (indicating the presence of Br^- outside the coordination sphere).

4. Coordination Isomerism

Coordination isomerism occurs in compounds that contain complex cations and complex anions. The isomers arise from the interchange of ligands between the complex cation and the complex anion.

Example: [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆] are coordination isomers. In the first compound, cobalt is the central metal in the cation, and chromium is in the anion. In the second, this is reversed.

Stereoisomerism

Stereoisomerism arises due to the different spatial arrangements of ligands around the central metal atom. It includes geometric isomerism and optical isomerism.

1. Geometric Isomerism

Geometric isomerism occurs in coordination compounds where ligands have different positions in space relative to each other. It is particularly common in square planar and octahedral complexes.

Octahedral Complexes

In octahedral complexes (coordination number 6), ligands can be in adjacent positions (cis) or opposite positions (trans).

• MA₄B₂ type: Two isomers are possible: cis (B ligands adjacent) and trans (B ligands opposite). Example: [Co(NH₃)₄Cl₂]⁺.
• MA₃B₃ type: Two isomers are possible: fac (facial - three identical ligands on one face) and mer (meridional - three identical ligands along a meridian). Example: [Co(NH₃)₃Cl₃].
• MA₂B₂ type: Three isomers are possible: cis-cis, cis-trans, and trans-trans. However, typically only two distinct geometric isomers exist: cis (both pairs of identical ligands are adjacent) and trans (both pairs of identical ligands are opposite). Example: [Pt(NH₃)₂Cl₂] (square planar). For octahedral, consider [Co(en)₂Cl₂]⁺, which has cis and trans isomers.
• MABCDEF type: 15 geometric isomers are possible.

Square Planar Complexes

In square planar complexes (coordination number 4), geometric isomerism is common, especially in MA₂B₂ and MAB types.

• MA₂B₂ type: Two isomers are possible: cis (A ligands adjacent) and trans (A ligands opposite). Example: [Pt(NH₃)₂Cl₂]. The cis isomer is known as 'cisplatin' and is used as an anticancer drug.
• MAB type: Two isomers are possible: cis and trans. Example: [Pt(NH₃)(Py)ClBr] where Py is pyridine.

2. Optical Isomerism

Optical isomerism occurs when a compound is chiral, meaning it is non-superimposable on its mirror image. Such compounds rotate the plane of polarized light. Optical isomers are called enantiomers.

Optical isomerism is observed in coordination compounds that lack a plane of symmetry. This is often seen in complexes with bidentate ligands arranged in a cis configuration or in complexes with multiple chiral centers.

• Octahedral complexes:
  • M(AA)₃ type: Example: [Co(en)₃]³⁺. This complex exists as a pair of enantiomers.
  • MA₂(BB)₂ type: Example: [Co(NH₃)₂(en)₂]³⁺. The cis isomer is optically active, while the trans isomer is optically inactive due to a plane of symmetry.
  • M(AA)₂B₂ type: Example: [Co(en)₂Cl₂]⁺. The cis isomer is optically active, while the trans isomer is optically inactive.
• Square planar complexes: Generally, square planar complexes of the type MAB (where A, B, C, D are different monodentate ligands) or M(AA)BC can exhibit optical isomerism if they are chiral. However, complexes of the type MA₂B₂ or MA₂BC do not show optical isomerism because they possess a plane of symmetry.

Bonding Approaches in Coordination Compounds

Several theories have been proposed to explain the nature of bonding, structure, and properties of coordination compounds. The most important ones are Valence Bond Theory (VBT), Crystal Field Theory (CFT), and Ligand Field Theory (LFT).

1. Valence Bond Theory (VBT)

Proposed by Linus Pauling, VBT explains the formation of coordinate bonds as the overlap between filled ligand orbitals and empty hybrid orbitals of the central metal atom. The key aspects are:

• Hybridization: The central metal atom undergoes hybridization to form empty hybrid orbitals of appropriate geometry and energy. These hybrid orbitals are directed towards the ligands.
• Coordinate Bond Formation: The filled ligand orbitals overlap with these empty hybrid orbitals to form coordinate covalent bonds.
• Inner and Outer Orbital Complexes: Based on the type of d-orbitals involved in hybridization, complexes are classified as inner orbital (or low spin) complexes (using (n-1)d orbitals) and outer orbital (or high spin) complexes (using nd orbitals).

Examples:

• [Co(NH₃)₆]³⁺: Co(III) is d⁶. It uses (n-1)d orbitals (3d) for hybridization, forming inner orbital complex: 3d⁶ → 4s¹4p³4d² (d²sp³ hybridization). This is a low spin complex.
• [Fe(H₂O)₆]³⁺: Fe(III) is d⁵. It uses nd orbitals (4d) for hybridization, forming outer orbital complex: 4s¹4p³4d¹ (sp³d² hybridization). This is a high spin complex.

Limitations of VBT: VBT fails to explain the magnetic properties (especially quantitative aspects), the colors of coordination compounds, and the stability of complexes.

2. Crystal Field Theory (CFT)

Developed by Hans Bethe and John Hasbrouck van Vleck, CFT treats the metal-ligand bond as purely ionic, arising from electrostatic attraction between the metal ion and the charged ligands (or the partial charges on polar ligands). It focuses on the splitting of d-orbitals in the presence of ligands.

Assumptions of CFT:

• The metal-ligand bond is purely electrostatic.
• Ligands are point charges (anionic) or point dipoles.
• The d-orbitals of the central metal ion split into different energy levels due to the electrostatic field created by the ligands.
• The splitting pattern depends on the geometry of the complex and the nature of the ligands.

Splitting of d-orbitals:

In an isolated gaseous metal ion, the five d-orbitals (d_xy, d_yz, d_zx, d_x²-y², d_z²) are degenerate (have the same energy). When ligands approach the metal ion, they repel the electrons in the d-orbitals. The d-orbitals whose electron clouds are directed towards the ligands experience greater repulsion and are raised in energy, while those directed between the ligands experience less repulsion and are lowered in energy.

1. Splitting in Octahedral Complexes:

In an octahedral field, the ligands approach along the x, y, and z axes. The d_x²-y² and d_z² orbitals (called e_g set) point directly towards the ligands and are raised in energy. The d_xy, d_yz, and d_zx orbitals (called t_2g set) point between the ligands and are lowered in energy. The energy difference between the e_g and t_2g levels is called the crystal field splitting energy, denoted as Δ_o (or 10 Dq).

The e_g orbitals are raised by 0.6 Δ_o above the mean energy, and the t_2g orbitals are lowered by 0.4 Δ_o below the mean energy. The mean energy is called the barycenter.

2. Splitting in Tetrahedral Complexes:

In a tetrahedral field, the ligands are positioned between the axes. The splitting pattern is inverted compared to octahedral complexes, and the magnitude of splitting (Δ_t) is smaller (approximately 4/9 of Δ_o). The d_xy, d_yz, d_zx orbitals are raised in energy (t₂ set), and the d_x²-y², d_z² orbitals are lowered (e set). Since Δ_t is small, pairing of electrons rarely occurs, and tetrahedral complexes are generally high spin.

3. Splitting in Square Planar Complexes:

Square planar complexes can be considered as a distorted octahedral complex where the ligands along the z-axis are moved infinitely far away. This results in a more complex splitting pattern of the d-orbitals, typically involving five levels: d_z², d_xz, d_yz, d_xy, and d_x²-y². The d_x²-y² orbital is raised to the highest energy level due to strong repulsion from ligands in the xy plane.

Crystal Field Stabilization Energy (CFSE):

CFSE is the stabilization energy gained by the complex due to the splitting of d-orbitals. It affects the stability, magnetic properties, and spectral properties of coordination compounds.

Spectrochemical Series:

CFT helps explain the colors of coordination compounds by relating them to the energy of the electronic transitions between split d-orbitals. The spectrochemical series ranks ligands in order of their ability to cause d-orbital splitting (Δ_o):

I^- < Br^- < S^2- < SCN^- < Cl^- < NO₃^- < F^- < OH^- < C₂O₄^2- < H₂O < NCS^- < CH₃CN < py < NH₃ < en < bipy < NO₂^- < PPh₃ < CN^- < CO

Ligands on the left are weak-field ligands (small Δ_o), while those on the right are strong-field ligands (large Δ_o).

3. Ligand Field Theory (LFT)

LFT is a more advanced theory that combines aspects of VBT and CFT. It considers the metal-ligand bond to be covalent but acknowledges the electrostatic interactions. LFT treats ligands as having molecular orbitals that interact with the metal d-orbitals, leading to the formation of bonding, non-bonding, and antibonding molecular orbitals. This theory provides a more accurate explanation for the magnetic and spectral properties of coordination compounds, especially those with covalent character in metal-ligand bonds.

Colour and Magnetic Properties

The colour and magnetic properties of coordination compounds are well explained by CFT.

Colour

Most coordination compounds are coloured. According to CFT, the colour arises from the absorption of certain wavelengths of visible light, which causes electronic transitions between the split d-orbitals (d-d transitions). The absorbed light promotes an electron from a lower energy d-orbital (t_2g in octahedral) to a higher energy d-orbital (e_g in octahedral). The colour we observe is the complementary colour of the light absorbed.

Factors Affecting Colour:

• Nature of the metal ion: Oxidation state and identity of the metal.
• Nature of the ligands: Strong-field ligands cause larger splitting (Δ_o), requiring higher energy (shorter wavelength) light absorption, leading to different observed colours.
• Geometry of the complex: Octahedral, tetrahedral, and square planar complexes have different splitting patterns and thus different colours.

Example: [Cu(H₂O)₆]²⁺ is blue because it absorbs light in the orange-red region of the spectrum. [Ni(H₂O)₆]²⁺ is green because it absorbs light in the violet-red region.

Absence of Colour:

• Complexes with d⁰ or d¹⁰ configurations do not show d-d transitions and are usually colourless (e.g., [Sc(H₂O)₆]³⁺, [Zn(H₂O)₆]²⁺).
• Complexes where the ligand absorbs light in the visible region (ligand-to-metal charge transfer, LMCT) or the metal absorbs light (metal-to-ligand charge transfer, MLCT) can also be coloured.

Magnetic Properties

The magnetic properties of coordination compounds depend on the presence and number of unpaired electrons. CFT helps predict whether a complex will be paramagnetic (unpaired electrons) or diamagnetic (all electrons paired).

• Paramagnetic: Compounds with one or more unpaired electrons are attracted to a magnetic field.
• Diamagnetic: Compounds with all electrons paired are weakly repelled by a magnetic field.

In CFT, the splitting energy Δ_o and the pairing energy (P) determine the electron configuration and thus the magnetic property.

• Low spin complexes (strong-field ligands, large Δ_o): Electrons preferentially pair up in the lower energy t_2g orbitals before occupying the higher energy e_g orbitals. This often leads to fewer unpaired electrons and diamagnetism or lower paramagnetism.
• High spin complexes (weak-field ligands, small Δ_o): Electrons occupy the higher energy e_g orbitals before pairing up in the t_2g orbitals, leading to the maximum number of unpaired electrons and higher paramagnetism.

Example:

• [Co(NH₃)₆]³⁺ (d⁶, strong field ligand NH₃): Low spin complex. Configuration is (t_2g)⁶. Zero unpaired electrons. Diamagnetic.
• [Fe(H₂O)₆]³⁺ (d⁵, weak field ligand H₂O): High spin complex. Configuration is (t_2g)³(e_g)². Five unpaired electrons. Strongly paramagnetic.

Importance of Coordination Compounds

Coordination compounds play vital roles in various fields, including analytical chemistry, industrial processes, and biological systems.

1. Importance in Qualitative Analysis

Coordination compounds are extensively used in the qualitative analysis of metal ions.

• Detection of ions: Formation of characteristic colours upon reaction with specific ligands is used to identify metal ions. For example, the formation of a blood-red colour with potassium thiocyanate (KSCN) indicates the presence of Fe³⁺ ions due to the formation of [Fe(SCN)(H₂O)₅]²⁺. The formation of a deep blue precipitate with ammonium thiocyanate indicates Co²⁺.
• Separation of ions: Differences in the stability and solubility of coordination complexes are exploited for separating metal ions. For example, in the separation of Group II cations in qualitative analysis, the formation of stable thiocayanato complexes of Cu²⁺ and Cd²⁺ keeps them in solution, preventing their precipitation as sulfides.
• Complexometric Titrations: EDTA (ethylenediaminetetraacetic acid) is a powerful chelating agent that forms stable, colourless complexes with most metal ions. It is used in complexometric titrations to determine the concentration of metal ions. The endpoint is often detected using a metallochromic indicator that forms a coloured complex with the metal ion.

2. Importance in Biological Systems

Coordination compounds are fundamental to life processes. Many essential biological molecules are coordination complexes.

• Haemoglobin: This protein in red blood cells contains iron as a central atom in a porphyrin ring (heme). Haemoglobin is responsible for transporting oxygen from the lungs to the tissues and carbon dioxide from the tissues to the lungs. The iron atom in haemoglobin can reversibly bind to oxygen.
• Chlorophyll: This pigment in plants is responsible for photosynthesis. It contains magnesium as the central metal atom in a porphyrin ring. Chlorophyll absorbs light energy, which is crucial for converting carbon dioxide and water into glucose.
• Vitamin B₁₂ (Cyanocobalamin): This essential vitamin contains cobalt as the central metal atom in a corrin ring. Vitamin B₁₂ plays a vital role in the formation of red blood cells, DNA synthesis, and the functioning of the nervous system.
• Enzymes: Many enzymes contain metal ions as cofactors, which are essential for their catalytic activity. These metal ions are coordinated to amino acid residues in the protein structure. Examples include carbonic anhydrase (Zn²⁺), superoxide dismutase (Cu²⁺/Zn²⁺ or Mn²⁺), and cytochrome oxidase (Fe²⁺/Cu⁺).

3. Industrial and Other Applications

Coordination compounds have numerous applications in industry and technology.

• Electroplating: Metal plating, such as gold or chromium plating, is often carried out using solutions containing complex ions (e.g., K₄[Fe(CN)₆] for iron plating, or complexes of gold and silver). This allows for controlled deposition of metals.
• Catalysis: Many coordination compounds are used as catalysts in industrial processes. Examples include Ziegler-Natta catalysts (based on titanium and aluminum complexes) for polymerization of alkenes, Wilkinson's catalyst ([RhCl(PPh₃)₃]) for hydrogenation, and complexes of palladium in organic synthesis (e.g., Heck reaction, Suzuki coupling).
• Pigments: Certain coordination compounds are used as pigments due to their vibrant colours. For example, Prussian blue (iron hexacyanoferrate) is a deep blue pigment. Cadmium sulfide (CdS) is used as a yellow pigment.
• Extraction of Metals: Coordination chemistry is used in the extraction and purification of metals. For instance, in the purification of nickel, the volatile complex Ni(CO)₄ is formed and then decomposed to obtain pure nickel.
• Medicine: Besides cisplatin as an anticancer drug, other coordination compounds are used in medical imaging (e.g., gadolinium complexes as MRI contrast agents) and as therapeutic agents.