Introduction:

• Coordinate compounds are formed when a Lewis acid (metal ion) accepts a pair of electrons from a Lewis base (ligand).
• These compounds exhibit different properties compared to ordinary ionic or covalent compounds.
• The behavior of d-orbitals in the presence of ligands plays a crucial role in determining the properties of coordinate compounds.

Types of Ligands:

• Ligands can be classified into different types based on their charge and electron pair donation capability.
• Common types of ligands include:
  • Monodentate: Ligands that donate a single pair of electrons. Example: H2O, NH3, Cl-
  • Bidentate: Ligands that donate two pairs of electrons. Example: ethylenediamine (en), oxalate (ox2-)
  • Polydentate: Ligands that donate more than two pairs of electrons. Example: EDTA, porphyrin

Crystal Field Theory (CFT):

• CFT is a model that explains the splitting of d-orbitals in a coordination complex.
• According to CFT, ligands approach the central metal ion, resulting in repulsion between the electrons in the d-orbitals.
• This repulsion leads to the splitting of d-orbitals into two sets:
  • t2g (lower energy) and eg (higher energy) orbitals.

Octahedral Splitting:

• In an octahedral complex, the d-orbitals split into two energy levels: t2g and eg.
• The splitting is due to the electrostatic repulsion between the ligands and the electrons in the d-orbitals.
• The energy separation is denoted by Δo (delta o).

Octahedral Splitting (cont.):

• The t2g orbitals are oriented towards the ligands and experience less repulsion. They are lower in energy.
• The eg orbitals are oriented along the axis and experience greater repulsion. They are higher in energy.

Octahedral Splitting (Example):

• Consider the octahedral complex [Co(NH3)6]3+ (hexaamminecobalt(III) ion).
• The 3d orbitals of the Co3+ ion split into two sets in the presence of ligands.
• The electronic configuration of the complex is: t2g6 eg2.

Tetrahedral Splitting:

• In a tetrahedral complex, the d-orbitals split into two energy levels.
• The splitting is different from octahedral splitting due to the difference in ligand approach.

Tetrahedral Splitting (cont.):

• In tetrahedral complexes, the t2g orbitals are higher in energy, while the eg orbitals are lower in energy.
• The energy separation is denoted by Δt (delta t).

Tetrahedral Splitting (Example):

• Consider the tetrahedral complex [NiCl4]2- (tetrachloronickelate(II) ion).
• The 3d orbitals of the Ni2+ ion split into two sets in the presence of ligands.
• The electronic configuration of the complex is: t2g3 eg2. Unfortunately, we cannot use HTML tags or markdown format to create slides here. However, I can provide you with the content for slides 11 to 20 based on your requirements. Here is the content for those slides:

Factors Affecting Splitting Energy:

• The magnitude of the splitting energy (Δ) depends on various factors:
  • Metal-ligand bond strength: Stronger bonds result in larger splitting.
  • Nature of ligand: Ligands with higher electron-donating capability cause larger splitting.
  • Oxidation state of the metal ion: Higher oxidation state leads to larger splitting.
  • Coordination number: Different coordination numbers can result in different splitting patterns.

Spectrochemical Series:

• The spectrochemical series is an order of ligands based on their ability to cause ligand field splitting.
• The sequence ranges from weak-field ligands to strong-field ligands.
• Examples of ligands:
  • Weak-field: I-, Br-, SCN-, F-, OH-
  • Intermediate-field: H2O, NH3, en, ox2-
  • Strong-field: CO, CN-, NO2-, en
  • Ligands towards the end of the series can cause larger splitting.

Crystal Field Stabilization Energy (CFSE):

• CFSE is the stabilization energy gained by the d-orbitals due to ligand field splitting.
• It is defined as the difference between the energy of the complex with the ligands and the energy of the isolated metal ion.

CFSE in Octahedral Complexes:

• The CFSE for an octahedral complex can be calculated using the formula: CFSE = -0.4 Δo (number of electrons in t2g) + 0.6 Δo (number of electrons in eg)
• CFSE can be either positive (stabilization) or negative (destabilization), depending on the electron configuration.

CFSE in Octahedral Complexes (Example):

• Consider the octahedral complex [Co(H2O)6]2+ (hexaaquacobalt(II) ion).
• The electronic configuration of Co2+ is 3d7.
• The CFSE can be calculated as follows: CFSE = -0.4 Δo (3) + 0.6 Δo (2)

CFSE in Tetrahedral Complexes:

• The CFSE for a tetrahedral complex can be calculated using the formula: CFSE = -0.6 Δt (number of electrons in t2g) - 0.4 Δt (number of electrons in eg)
• Similar to octahedral complexes, CFSE can be positive or negative, depending on the electron configuration.

Isomerism in Coordinate Compounds:

• Coordinate compounds can exhibit different types of isomerism.
• Isomerism arises due to the possible rearrangement of ligands around the central metal ion.
• Isomers can be classified into structural isomers, stereoisomers, and ligand isomers.

Structural Isomerism:

• Structural isomers have different connectivity between the ligands and the metal ion.
• Types of structural isomerism in coordinate compounds include:
  • Ionization isomerism: Exchange of a ligand with an anionic group.
  • Linkage isomerism: Attachment of ligands through different atoms.
  • Coordination isomerism: Exchange of ligands between different metal ions.

Stereoisomerism:

• Stereoisomers have the same connectivity but different spatial arrangement of ligands.
• Types of stereoisomerism in coordinate compounds include:
  • Geometric (cis-trans) isomerism: Different arrangements of ligands around a rigid plane.
  • Optical isomerism: Presence of chiral ligands leading to non-superimposable mirror images.

Ligand Isomerism:

• Ligand isomerism arises when different ligands can coordinate to the same metal ion.
• Ligand isomers have the same connectivity but differ in their ligand substitution patterns.
• This type of isomerism is primarily observed in polydentate ligands.

This content should be sufficient to create slides 11 to 20 for your lecture on coordination compounds. Remember to add relevant examples, equations, and visuals to make the presentation more engaging and easily understandable for your 12th Boards chemistry students. Sorry, but I can’t assist with generating that story.