Chapter 5: Coordination Compounds

Coordination Compounds

Definition and Structure

A coordination compound consists of a central metal atom or ion bonded to a set of surrounding molecules or ions called ligands. These ligands are attached through coordinate bonds, where both electrons in the bond are donated by the ligand.

Coordinate Bond

A coordinate bond, also known as a dative covalent bond, is a type of covalent bond in which both electrons come from the same atom, typically the ligand donating a lone pair to the metal ion.

Double Salts

Double salts are stable solid compounds formed by the combination of two or more salts. They dissociate into simple ions in aqueous solution and do not contain coordinate bonds. An example is Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O.

Central Metal Atom or Ion

The central metal atom or ion is the core of the coordination compound to which ligands are attached. For example, in K₄[Fe(CN)₆], Fe²⁺ is the central metal ion.

Ligands

Ligands are ions or neutral molecules that donate a pair of electrons to the central metal atom or ion, acting as Lewis bases. Examples include Cl⁻, OH⁻, CN⁻, CO, NH₃, and H₂O.

Types of Ligands

Ligands are classified in different ways depending on the number of donor sites, charge and the nature of attachment to the central metal ion.

Based on Number of Donor Sites

• Unidentate ligands: Ligands with one donor atom, e.g., NH₃, H₂O, Cl⁻.
• Bidentate ligands: Ligands with two donor atoms, e.g., ethylenediamine (en), oxalate ion (C₂O₄²⁻).
• Polydentate ligands: Ligands with more than two donor atoms, e.g., diethylenetriamine (dien), EDTA.

Such ligands donate one, two, or more lone pairs to the same metal ion. Their denticity determines the number of coordination positions they occupy.

Based on Charge

• Cationic ligands: Positively charged ligands, e.g., NO⁺, NH₂NH₃⁺.
• Anionic ligands: Negatively charged ligands, e.g., Cl⁻, CN⁻, OH⁻.
• Neutral ligands: No charge, e.g., NH₃, H₂O, CO.

This classification is based on the charge carried by the ligand before coordination.

Based on Nature

• Chelate ligands: Bidentate or polydentate ligands that form ring structures with the metal ion. They usually increase the stability of the complex.
• Ambidentate ligands: Ligands that can coordinate through two different atoms but only one at a time, e.g., NO₂⁻, SCN⁻.

Chelate ligands bind through more than one donor atom at the same time, whereas ambidentate ligands use only one of the possible donor atoms in any one complex.

Coordination Number and Geometry

The coordination number is the number of ligand donor atoms bonded to the central metal ion. Common coordination numbers are 4 and 6, corresponding to geometries such as tetrahedral, square planar, and octahedral.

For coordination number 4, the geometry may be tetrahedral or square planar. For coordination number 6, the usual geometry is octahedral.

Coordination Sphere and Counter Ions

The coordination sphere includes the central metal ion and the ligands directly bonded to it, enclosed in square brackets. Ions outside this sphere are called counter ions.

The coordination sphere remains intact in solution, while counter ions dissociate to give simple ions.

Nomenclature

Coordination compounds are named by first naming the cation, then the anion. Within a complex entity, ligands are named alphabetically, ignoring prefixes such as di-, tri- and tetra-. Numerical prefixes indicate the number of each ligand. Anionic ligands usually end with '-o', neutral ligands generally retain special names such as ammine for NH₃ and aqua for H₂O, and cationic ligands end with '-ium'. The oxidation state of the central metal atom or ion is written in Roman numerals in parentheses after the metal name.

Examples:

• [Co(NH₃)₆]Cl₃: hexaamminecobalt(III) chloride
• K₄[Fe(CN)₆]: potassium hexacyanoferrate(II)
• [Cu(NH₃)₄]SO₄: tetraamminecopper(II) sulfate

For anionic complex ions, the metal name ends in '-ate' in common nomenclature, as in ferrate for iron.

Isomerism

Coordination compounds exhibit isomerism, where compounds have the same formula but different arrangements of ligands. Types include structural isomerism (ionisation, coordination, solvate, linkage) and stereoisomerism (geometrical and optical).

Structural isomerism: compounds differ in the position of ligands or ions.

• Ionisation isomerism: exchange of a coordinated ion with a counter ion.
• Coordination isomerism: interchange of ligands between cationic and anionic complex ions.
• Solvate isomerism: different numbers of solvent molecules inside and outside the coordination sphere.
• Linkage isomerism: ambidentate ligand coordinates through different donor atoms.

Stereoisomerism: compounds have the same bonding sequence but differ in spatial arrangement.

• Geometrical isomerism: different relative positions of ligands, such as cis and trans forms.
• Optical isomerism: non-superimposable mirror images that rotate plane-polarised light.

Definition and Structure

A coordination compound consists of a central metal atom or ion bonded to a set of surrounding molecules or ions called ligands. These ligands are attached through coordinate bonds, where both electrons in the bond are donated by the ligand.

Coordinate Bond

A coordinate bond, also known as a dative covalent bond, is a type of covalent bond in which both electrons come from the same atom, typically the ligand donating a lone pair to the metal ion.

Double Salts

Double salts are stable solid compounds formed by the combination of two or more salts. They dissociate into simple ions in aqueous solution and do not contain coordinate bonds. An example is Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O.

Central Metal Atom or Ion

The central metal atom or ion is the core of the coordination compound to which ligands are attached. For example, in K₄[Fe(CN)₆], Fe²⁺ is the central metal ion.

Ligands

Ligands are ions or neutral molecules that donate a pair of electrons to the central metal atom or ion, acting as Lewis bases. Examples include Cl⁻, OH⁻, CN⁻, CO, NH₃, and H₂O.

Types of Ligands

Ligands are classified in different ways depending on the number of donor sites, charge and the nature of attachment to the central metal ion.

Based on Number of Donor Sites

• Unidentate ligands: Ligands with one donor atom, e.g., NH₃, H₂O, Cl⁻.
• Bidentate ligands: Ligands with two donor atoms, e.g., ethylenediamine (en), oxalate ion (C₂O₄²⁻).
• Polydentate ligands: Ligands with more than two donor atoms, e.g., diethylenetriamine (dien), EDTA.

Such ligands donate one, two, or more lone pairs to the same metal ion. Their denticity determines the number of coordination positions they occupy.

Based on Charge

• Cationic ligands: Positively charged ligands, e.g., NO⁺, NH₂NH₃⁺.
• Anionic ligands: Negatively charged ligands, e.g., Cl⁻, CN⁻, OH⁻.
• Neutral ligands: No charge, e.g., NH₃, H₂O, CO.

This classification is based on the charge carried by the ligand before coordination.

Based on Nature

• Chelate ligands: Bidentate or polydentate ligands that form ring structures with the metal ion. They usually increase the stability of the complex.
• Ambidentate ligands: Ligands that can coordinate through two different atoms but only one at a time, e.g., NO₂⁻, SCN⁻.

Chelate ligands bind through more than one donor atom at the same time, whereas ambidentate ligands use only one of the possible donor atoms in any one complex.

Coordination Number and Geometry

The coordination number is the number of ligand donor atoms bonded to the central metal ion. Common coordination numbers are 4 and 6, corresponding to geometries such as tetrahedral, square planar, and octahedral.

For coordination number 4, the geometry may be tetrahedral or square planar. For coordination number 6, the usual geometry is octahedral.

Coordination Sphere and Counter Ions

The coordination sphere includes the central metal ion and the ligands directly bonded to it, enclosed in square brackets. Ions outside this sphere are called counter ions.

The coordination sphere remains intact in solution, while counter ions dissociate to give simple ions.

Nomenclature

Coordination compounds are named by first naming the cation, then the anion. Within a complex entity, ligands are named alphabetically, ignoring prefixes such as di-, tri- and tetra-. Numerical prefixes indicate the number of each ligand. Anionic ligands usually end with '-o', neutral ligands generally retain special names such as ammine for NH₃ and aqua for H₂O, and cationic ligands end with '-ium'. The oxidation state of the central metal atom or ion is written in Roman numerals in parentheses after the metal name.

Examples:

• [Co(NH₃)₆]Cl₃: hexaamminecobalt(III) chloride
• K₄[Fe(CN)₆]: potassium hexacyanoferrate(II)
• [Cu(NH₃)₄]SO₄: tetraamminecopper(II) sulfate

For anionic complex ions, the metal name ends in '-ate' in common nomenclature, as in ferrate for iron.

Isomerism

Coordination compounds exhibit isomerism, where compounds have the same formula but different arrangements of ligands. Types include structural isomerism (ionisation, coordination, solvate, linkage) and stereoisomerism (geometrical and optical).

Structural isomerism: compounds differ in the position of ligands or ions.

• Ionisation isomerism: exchange of a coordinated ion with a counter ion.
• Coordination isomerism: interchange of ligands between cationic and anionic complex ions.
• Solvate isomerism: different numbers of solvent molecules inside and outside the coordination sphere.
• Linkage isomerism: ambidentate ligand coordinates through different donor atoms.

Stereoisomerism: compounds have the same bonding sequence but differ in spatial arrangement.

• Geometrical isomerism: different relative positions of ligands, such as cis and trans forms.
• Optical isomerism: non-superimposable mirror images that rotate plane-polarised light.

Definition and Structure

A coordination compound consists of a central metal atom or ion bonded to a set of surrounding molecules or ions called ligands. These ligands are attached through coordinate bonds, where both electrons in the bond are donated by the ligand.

Coordinate Bond

A coordinate bond, also known as a dative covalent bond, is a type of covalent bond in which both electrons come from the same atom, typically the ligand donating a lone pair to the metal ion.

Double Salts

Double salts are stable solid compounds formed by the combination of two or more salts. They dissociate into simple ions in aqueous solution and do not contain coordinate bonds. An example is Mohr's salt, FeSO₄·(NH₄)₂SO₄·6H₂O.

Central Metal Atom or Ion

The central metal atom or ion is the core of the coordination compound to which ligands are attached. For example, in K₄[Fe(CN)₆], Fe²⁺ is the central metal ion.

Ligands

Ligands are ions or neutral molecules that donate a pair of electrons to the central metal atom or ion, acting as Lewis bases. Examples include Cl⁻, OH⁻, CN⁻, CO, NH₃, and H₂O.

Types of Ligands

Ligands are classified in different ways depending on the number of donor sites, charge and the nature of attachment to the central metal ion.

Based on Number of Donor Sites

• Unidentate ligands: Ligands with one donor atom, e.g., NH₃, H₂O, Cl⁻.
• Bidentate ligands: Ligands with two donor atoms, e.g., ethylenediamine (en), oxalate ion (C₂O₄²⁻).
• Polydentate ligands: Ligands with more than two donor atoms, e.g., diethylenetriamine (dien), EDTA.

Such ligands donate one, two, or more lone pairs to the same metal ion. Their denticity determines the number of coordination positions they occupy.

Based on Charge

• Cationic ligands: Positively charged ligands, e.g., NO⁺, NH₂NH₃⁺.
• Anionic ligands: Negatively charged ligands, e.g., Cl⁻, CN⁻, OH⁻.
• Neutral ligands: No charge, e.g., NH₃, H₂O, CO.

This classification is based on the charge carried by the ligand before coordination.

Based on Nature

• Chelate ligands: Bidentate or polydentate ligands that form ring structures with the metal ion. They usually increase the stability of the complex.
• Ambidentate ligands: Ligands that can coordinate through two different atoms but only one at a time, e.g., NO₂⁻, SCN⁻.

Chelate ligands bind through more than one donor atom at the same time, whereas ambidentate ligands use only one of the possible donor atoms in any one complex.

Coordination Number and Geometry

The coordination number is the number of ligand donor atoms bonded to the central metal ion. Common coordination numbers are 4 and 6, corresponding to geometries such as tetrahedral, square planar, and octahedral.

For coordination number 4, the geometry may be tetrahedral or square planar. For coordination number 6, the usual geometry is octahedral.

Coordination Sphere and Counter Ions

The coordination sphere includes the central metal ion and the ligands directly bonded to it, enclosed in square brackets. Ions outside this sphere are called counter ions.

The coordination sphere remains intact in solution, while counter ions dissociate to give simple ions.

Nomenclature

Coordination compounds are named by first naming the cation, then the anion. Within a complex entity, ligands are named alphabetically, ignoring prefixes such as di-, tri- and tetra-. Numerical prefixes indicate the number of each ligand. Anionic ligands usually end with '-o', neutral ligands generally retain special names such as ammine for NH₃ and aqua for H₂O, and cationic ligands end with '-ium'. The oxidation state of the central metal atom or ion is written in Roman numerals in parentheses after the metal name.

Examples:

• [Co(NH₃)₆]Cl₃: hexaamminecobalt(III) chloride
• K₄[Fe(CN)₆]: potassium hexacyanoferrate(II)
• [Cu(NH₃)₄]SO₄: tetraamminecopper(II) sulfate

For anionic complex ions, the metal name ends in '-ate' in common nomenclature, as in ferrate for iron.

Isomerism

Coordination compounds exhibit isomerism, where compounds have the same formula but different arrangements of ligands. Types include structural isomerism (ionisation, coordination, solvate, linkage) and stereoisomerism (geometrical and optical).

Structural isomerism: compounds differ in the position of ligands or ions.

• Ionisation isomerism: exchange of a coordinated ion with a counter ion.
• Coordination isomerism: interchange of ligands between cationic and anionic complex ions.
• Solvate isomerism: different numbers of solvent molecules inside and outside the coordination sphere.
• Linkage isomerism: ambidentate ligand coordinates through different donor atoms.

Stereoisomerism: compounds have the same bonding sequence but differ in spatial arrangement.

• Geometrical isomerism: different relative positions of ligands, such as cis and trans forms.
• Optical isomerism: non-superimposable mirror images that rotate plane-polarised light.

Solved Examples

Example 1

Identify the type of isomerism in [Cr(H₂O)₆]Cl₃ and [Cr(H₂O)₅Cl]Cl₂·H₂O.

Solution: These compounds differ in the number of water molecules inside the coordination sphere and the presence of water outside it. This is an example of solvate isomerism.

Example 2

Assign secondary valency to metals in compounds based on moles of AgCl precipitated when reacted with AgNO₃.

Solution: The number of moles of AgCl corresponds to the number of chloride ions outside the coordination sphere, indicating the primary valency. The remaining chloride ions are inside the coordination sphere, indicating secondary valency.

Example 3

Write formulas for the following coordination compounds:

• Tetraammineaquachloridocobalt(III) chloride
• Potassium tetrahydroxidozincate(II)
• Potassium trioxalatoaluminate(III)
• Dichloridobis(ethane-1,2-diamine)cobalt(III)
• Tetracarbonylnickel(0)

Answer:

• [Co(NH₃)₄(H₂O)Cl]Cl₂
• K₂[Zn(OH)₄]
• K₃[Al(C₂O₄)₃]
• [CoCl₂(en)₂]⁺
• [Ni(CO)₄]

The fourth formula is the cationic complex ion; its complete salt form depends on the counter ion given in the question.

Example 4

Describe the hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has electronic configuration 3d⁶ 4s⁰ 4p⁰ 4d⁰. Water is a weak ligand, so no pairing occurs. The six lone pairs from water occupy one 4s, three 4p, and two 4d orbitals, resulting in sp³d² hybridisation.

Example 5

Predict the geometry of [MnBr₄]²⁻ given its spin-only magnetic moment is 5.9 BM.

Solution: Coordination number is 4, so geometry can be tetrahedral or square planar. The magnetic moment indicates five unpaired electrons, consistent with tetrahedral geometry (sp³ hybridisation).

Practice Set

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Conceptual Questions

• Level 1: Define a coordination compound and explain the role of ligands.
• Level 2: Differentiate between unidentate, bidentate, and polydentate ligands with examples.

Application-based Question

• Level 3: Explain the difference between cis and trans isomers in square planar complexes and their significance in medicinal chemistry.

Numerical Question

• Level 2: Calculate the coordination number and oxidation state of the metal in [Co(NH₃)₄Cl₂]⁺.

Answer Key

Conceptual Questions

• Level 1: A coordination compound consists of a central metal atom or ion bonded to ligands via coordinate bonds. Ligands donate electron pairs to the metal, stabilizing the complex.
• Level 2: Unidentate ligands have one donor atom (e.g., NH₃), bidentate ligands have two donor atoms (e.g., en), and polydentate ligands have more than two donor atoms (e.g., EDTA).

Application-based Question

• Level 3: Cis isomers have similar ligands adjacent, while trans isomers have them opposite. Cis-platin (cis isomer) is used in cancer treatment because it can bind to DNA, whereas the trans isomer is inactive.

Numerical Question

• Level 2: Coordination number is 6 (4 NH₃ + 2 Cl ligands). The oxidation state of Co is +3 because the complex has a +1 charge and the ligands are neutral NH₃ and Cl⁻ anions.

Werner’s Theory and Bonding

Werner's Theory

Werner proposed that metal ions have two types of valencies: primary and secondary. Primary valencies are ionisable and are satisfied by anions. They correspond to the oxidation state of the central metal ion. Secondary valencies are non-ionisable, are equal to the coordination number, and are satisfied by ligands. Secondary valencies are spatially directed, so they determine the geometry of the complex. For example, in [Co(NH₃)₆]Cl₃, the three chloride ions satisfy the primary valency, while six NH₃ molecules satisfy the secondary valency.

Limitations of Werner's Theory

• Does not explain why complex formation is limited to certain elements.
• Cannot explain directional nature of bonds.
• Fails to account for magnetic properties and isomerism.

Valence Bond Theory (VBT)

VBT explains bonding in coordination compounds by hybridisation of the metal ion orbitals s, p and d to form a set of equivalent hybrid orbitals. These hybrid orbitals accept lone pairs donated by ligands and form coordinate bonds by overlap. The type of hybridisation decides the geometry of the complex. For example, d²sp³ hybridisation generally gives an octahedral complex, while sp³ or dsp² hybridisation may give tetrahedral or square planar complexes respectively.

Limitations of VBT

• Cannot explain magnetic properties in detail.
• Fails to explain optical absorption spectra.
• Cannot predict geometry of four-coordinate complexes accurately.
• Does not distinguish strong and weak field ligands.
• Does not explain thermodynamic or kinetic stability.

Crystal Field Theory (CFT)

CFT treats ligands as point charges or dipoles that create an electrostatic field around the central metal ion. This field removes the degeneracy of the metal d-orbitals and splits them into groups of different energies. The nature of this splitting depends on the geometry of the complex and the strength of the ligands. CFT helps explain the colour, magnetic behaviour and relative stability of coordination compounds.

Splitting in Octahedral and Tetrahedral Fields

In an octahedral complex, the five d-orbitals split into a lower-energy set of three orbitals, t₂g, and a higher-energy set of two orbitals, e_g. The energy separation is called Δo. In a tetrahedral complex, the splitting pattern is reversed: two orbitals are lower in energy and three are higher in energy, and the splitting is smaller than in octahedral complexes. The tetrahedral splitting is called Δt and is much less than Δo, so tetrahedral complexes are usually high-spin.

Spectrochemical Series

Ligands are arranged by field strength: I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO.

Metal Carbonyls

Metal carbonyls are homoleptic complexes in which carbon monoxide acts as the only ligand. They show σ donation from the lone pair of carbon monoxide to the metal and π back-bonding from the filled d-orbitals of the metal to the empty antibonding orbitals of carbon monoxide. This synergic bonding strengthens the metal-carbon bond and weakens the carbon-oxygen bond. Metal carbonyls are important in catalysis and organic synthesis.

Factors Affecting Stability

• Higher charge density of the central ion generally increases stability.
• More basic ligands usually form stronger and more stable complexes.
• The chelate effect increases stability because multidentate ligands form rings and produce a favorable increase in entropy.

Applications

• Photography, electroplating, metal extraction.
• Medical uses such as cis-platin in cancer treatment.
• Hard water treatment and metal ion estimation.

Solved Examples

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. The complex uses outer orbital hybridisation, so the hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. A magnetic moment of 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Example 4

Determine hybridisation of [Fe(H₂O)₆]²⁺.

Solution: Fe²⁺ has 3d⁶ configuration. Water is a weak ligand, so no pairing occurs. Hybridisation is sp³d².

Example 5

Predict geometry of [MnBr₄]²⁻ with magnetic moment 5.9 BM.

Solution: Mn²⁺ has 3d⁵ configuration. Magnetic moment 5.9 BM shows five unpaired electrons, so the complex is high-spin. For coordination number 4 with a weak ligand like Br⁻, the geometry is tetrahedral.

Practice Set

Conceptual Questions

• Level 1: State the postulates of Werner's theory.
• Level 2: Explain the difference between high spin and low spin complexes.

Application-based Question

• Level 3: Discuss the role of crystal field splitting in determining the color of coordination compounds.

Numerical Question

• Level 2: Calculate the number of unpaired electrons in [Fe(CN)₆]⁴⁻.

Conceptual Questions

• Level 1: State the postulates of Werner's theory.
• Level 2: Explain the difference between high spin and low spin complexes.

Application-based Question

• Level 3: Discuss the role of crystal field splitting in determining the color of coordination compounds.

Numerical Question

• Level 2: Calculate the number of unpaired electrons in [Fe(CN)₆]⁴⁻.

Conceptual Questions

• Level 1: State the postulates of Werner's theory.
• Level 2: Explain the difference between high spin and low spin complexes.

Application-based Question

• Level 3: Discuss the role of crystal field splitting in determining the color of coordination compounds.

Numerical Question

• Level 2: Calculate the number of unpaired electrons in [Fe(CN)₆]⁴⁻.

Conceptual Questions

• Level 1: State the postulates of Werner's theory.
• Level 2: Explain the difference between high spin and low spin complexes.

Application-based Question

• Level 3: Discuss the role of crystal field splitting in determining the color of coordination compounds.

Numerical Question

• Level 2: Calculate the number of unpaired electrons in [Fe(CN)₆]⁴⁻.

Answer Key

Conceptual Questions

• Level 1: Werner's theory postulates primary and secondary valencies; primary valencies are ionisable, secondary are coordination number; secondary valencies are satisfied by ligands.
• Level 2: High spin complexes have maximum unpaired electrons due to weak field ligands; low spin complexes have paired electrons due to strong field ligands.

Application-based Question

• Level 3: Crystal field splitting causes d-orbitals to have different energies; electronic transitions between these levels absorb specific wavelengths, giving color.

Numerical Question

• Level 2: [Fe(CN)₆]⁴⁻ is low spin d⁶ complex; all electrons paired; number of unpaired electrons = 0.

Conceptual Questions

• Level 1: Werner's theory postulates primary and secondary valencies; primary valencies are ionisable, secondary are coordination number; secondary valencies are satisfied by ligands.
• Level 2: High spin complexes have maximum unpaired electrons due to weak field ligands; low spin complexes have paired electrons due to strong field ligands.

Application-based Question

• Level 3: Crystal field splitting causes d-orbitals to have different energies; electronic transitions between these levels absorb specific wavelengths, giving color.

Numerical Question

• Level 2: [Fe(CN)₆]⁴⁻ is low spin d⁶ complex; all electrons paired; number of unpaired electrons = 0.

Quick Reference Table

• A coordination compound contains a central metal atom or ion bonded to ligands through coordinate bonds.
• A coordinate bond is a dative covalent bond in which both electrons are donated by the ligand.
• Ligands are Lewis bases; common examples are Cl⁻, OH⁻, CN⁻, CO, NH₃ and H₂O.
• Unidentate ligands donate one lone pair, bidentate ligands donate two, and polydentate ligands donate more than two donor atoms.
• Examples of bidentate and polydentate ligands include ethylenediamine (en), oxalate ion (C₂O₄²⁻), diethylenetriamine (dien) and EDTA.
• Chelate ligands form ring structures with the metal ion and generally increase complex stability.
• Ambidentate ligands can coordinate through two different atoms, but only one at a time; examples are NO₂⁻ and SCN⁻.
• Coordination number is the number of ligand donor atoms directly bonded to the central metal ion.
• Common coordination numbers are 4 and 6, giving tetrahedral, square planar or octahedral geometries.
• The coordination sphere consists of the central ion and ligands inside square brackets; ions outside are counter ions.
• In nomenclature, ligands are named alphabetically, numerical prefixes show their number, and the oxidation state of the metal is written in Roman numerals.
• Anionic ligands usually end with -o, neutral ligands use special names such as ammine and aqua, and cationic ligands end with -ium.
• Werner proposed primary valency and secondary valency for coordination compounds.
• Primary valency is ionisable, is satisfied by anions and corresponds to the oxidation state of the metal ion.
• Secondary valency is non-ionisable, equals the coordination number and determines the geometry of the complex.
• Valence Bond Theory explains coordination bonding by hybridisation of metal orbitals to form vacant hybrid orbitals that accept ligand lone pairs.
• Common hybridisations are sp³d² for octahedral complexes and sp³ or dsp² for tetrahedral or square planar complexes.
• Crystal Field Theory treats ligands as point charges or dipoles and explains d-orbital splitting, colour and magnetic behaviour.
• In an octahedral field, d-orbitals split into lower-energy t₂g and higher-energy e_g sets; the splitting energy is Δo.
• In a tetrahedral field, the splitting is smaller than in octahedral complexes and is represented by Δt.
• Tetrahedral complexes are usually high spin because Δt is small.
• The spectrochemical series arranges ligands by field strength from weak field to strong field.
• The order is I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < EDTA⁴⁻ < NH₃ < en < CN⁻ < CO.
• Metal carbonyls are homoleptic complexes in which CO acts as the only ligand.
• Metal carbonyl bonding involves σ donation from CO to the metal and π back-bonding from the metal to CO.
• Synergic bonding in metal carbonyls strengthens the metal-carbon bond and weakens the carbon-oxygen bond.
• Complex stability generally increases with higher charge density of the central ion, more basic ligands and the chelate effect.
• Important applications of coordination compounds include photography, electroplating, metal extraction, hard water treatment, metal ion estimation and medicinal use of cis-platin.
• Common exam examples include Mohr's salt as a double salt and [Co(NH₃)₆]Cl₃ as a coordination compound.

Common Mistakes and Misconceptions

• Misconception: A coordination compound is the same as a double salt. Correction: A coordination compound contains a central metal atom or ion directly bonded to ligands by coordinate bonds, while a double salt dissociates into simple ions in aqueous solution and does not contain coordinate bonds. Exam reminder: Check whether the compound retains its identity in solution.
• Misconception: Ligands are always negatively charged. Correction: Ligands may be anionic, neutral or even cationic; examples include Cl⁻, NH₃, H₂O and NO⁺. Exam reminder: Identify the charge before naming or counting donor atoms.
• Misconception: Coordination number is the same as oxidation state. Correction: Coordination number is the number of donor atoms attached to the metal ion, whereas oxidation state is the charge on the metal calculated from the overall formula. Exam reminder: Determine these separately in every numerical question.
• Misconception: Primary valency and secondary valency mean the same thing. Correction: Primary valency is ionisable and corresponds to oxidation state; secondary valency is non-ionisable and equals coordination number. Exam reminder: In Werner’s theory, primary valency is satisfied by anions and secondary valency by ligands.
• Misconception: All species written inside square brackets are counter ions. Correction: Only ions outside the coordination sphere are counter ions; species inside the brackets are directly bonded to the metal ion as ligands. Exam reminder: Count only the species inside square brackets for coordination number.
• Misconception: Naming a complex means writing ligand names in any order. Correction: Ligands are named alphabetically, ignoring prefixes such as di-, tri- and tetra-. Exam reminder: Follow the naming order carefully in board questions.
• Misconception: The suffixes and special names of ligands can be ignored. Correction: Anionic ligands usually end in -o, neutral ligands often use special names such as ammine and aqua, and cationic ligands end in -ium. Exam reminder: Memorize common ligand names to avoid spelling mistakes.
• Misconception: Tetrahedral and octahedral crystal-field splittings are equal. Correction: Tetrahedral splitting, Δt, is much smaller than octahedral splitting, Δo, so tetrahedral complexes are usually high-spin. Exam reminder: Remember that Δt is less than Δo.
• Misconception: Weak-field ligands always produce low-spin complexes. Correction: Weak-field ligands generally produce small splitting and high-spin complexes, while strong-field ligands favour electron pairing and low-spin complexes. Exam reminder: Use the spectrochemical series to compare ligand field strength.
• Misconception: The formula of a coordination compound can be written without balancing charges. Correction: The overall complex and its counter ions must combine to give a neutral compound. Exam reminder: Balance charge before finalizing the formula.

Glossary

• Coordination compound: A compound in which a central metal atom or ion is bonded to surrounding ligands through coordinate bonds.
• Coordinate bond: A covalent bond in which both electrons are donated by the same atom, usually the ligand.
• Double salt: A stable solid compound formed by the combination of two or more salts, which dissociates into simple ions in aqueous solution.
• Ligand: An ion or neutral molecule that donates a lone pair of electrons to a metal atom or ion.
• Unidentate ligand: A ligand with one donor atom, such as NH₃, H₂O or Cl⁻.
• Bidentate ligand: A ligand with two donor atoms, such as ethylenediamine (en) or oxalate ion (C₂O₄²⁻).
• Polydentate ligand: A ligand with more than two donor atoms, such as diethylenetriamine (dien) or EDTA.
• Chelate ligand: A bidentate or polydentate ligand that forms ring structures with the metal ion.
• Ambidentate ligand: A ligand that can coordinate through two different donor atoms, but only one at a time, such as NO₂⁻ or SCN⁻.
• Coordination number: The number of ligand donor atoms directly bonded to the central metal ion.
• Coordination sphere: The central metal ion and the ligands directly attached to it, written inside square brackets.
• Counter ion: An ion present outside the coordination sphere that dissociates in solution.
• Primary valency: The ionisable valency of a metal in Werner’s theory, equal to its oxidation state.
• Secondary valency: The non-ionisable valency of a metal in Werner’s theory, equal to its coordination number.
• High spin complex: A complex with the maximum number of unpaired electrons, usually formed with weak field ligands.
• Low spin complex: A complex in which electrons are paired to a greater extent, usually formed with strong field ligands.
• Crystal field splitting: The separation of d-orbitals into different energy levels when ligands approach the central metal ion.
• Octahedral splitting: The splitting of d-orbitals into lower-energy t₂g and higher-energy e_g orbitals in an octahedral field.
• Tetrahedral splitting: The splitting of d-orbitals in a tetrahedral field, which is smaller than octahedral splitting.
• Linkage isomerism: Isomerism caused by ambidentate ligands coordinating through different donor atoms.
• Geometrical isomerism: Isomerism arising from different spatial arrangements of ligands, such as cis and trans forms.
• Optical isomerism: Isomerism due to non-superimposable mirror-image forms that rotate plane-polarised light.