NEET Chemistry — Coordination Compounds Revision Notes

These comprehensive revision notes cover the essential aspects of Coordination Compounds, a crucial chapter for NEET Chemistry. This resource is designed to provide a structured and in-depth understanding, ensuring you master the concepts required to tackle related questions effectively in the exam.

Coordination Compounds: NEET Revision Notes

Introduction to Coordination Compounds

Coordination compounds, also known as complex compounds, are a special class of compounds where a central metal atom or ion is bonded to a cluster of ions or molecules called ligands. These compounds play vital roles in biological systems (e.g., chlorophyll, hemoglobin, Vitamin B12) and industrial applications (e.g., catalysts, pigments, analytical reagents). Understanding their structure, bonding, and properties is fundamental for NEET.

Exam Weightage and Importance

Coordination Compounds is a high-yield chapter for NEET. Typically, 2-4 questions are asked from this chapter, contributing 8-16 marks to your total score. The questions can range from basic definitions and nomenclature to more complex topics like crystal field theory and isomerism. A thorough understanding can significantly boost your Chemistry score.

Key Concepts and Definitions

1. Werner's Theory of Coordination Compounds

Alfred Werner, the father of coordination chemistry, proposed the following postulates:

• Primary Valency (Ionizable Valency): This corresponds to the oxidation state of the central metal ion and is satisfied by anions. It is ionizable and can be represented by dotted lines.
• Secondary Valency (Non-ionizable Valency): This corresponds to the coordination number of the central metal ion and is satisfied by ligands. It is non-ionizable and can be represented by solid lines. Each metal ion has a fixed secondary valency.
• The secondary valencies are directed towards fixed positions in space, giving a definite geometry to the complex.
• A metal can exhibit both primary and secondary valencies.

Example: In $[Co(NH_3)_6]Cl_3$, Co has a primary valency of +3 (satisfied by three $Cl^-$ ions) and a secondary valency of 6 (satisfied by six $NH_3$ molecules).

2. Terminology in Coordination Compounds

• Central Metal Atom/Ion: The atom or ion (usually a transition metal) to which a fixed number of ions or molecules are directly attached. It acts as a Lewis acid (electron pair acceptor).
• Ligands: Ions or molecules that are capable of donating a pair of electrons to the central metal atom/ion. They act as Lewis bases (electron pair donors).
  • Classification based on number of donor atoms:
    • Monodentate (Unidentate): Ligands possessing one donor atom. E.g., $Cl^-$, $H_2O$, $NH_3$, $CN^-$.
    • Bidentate (Didentate): Ligands possessing two donor atoms. E.g., Ethylenediamine (en, $NH_2CH_2CH_2NH_2$), Oxalate ($C_2O_4^{2-}$).
    • Polydentate: Ligands possessing more than two donor atoms. E.g., EDTA (Ethylenediaminetetraacetate, a hexadentate ligand).
  • Ambidentate Ligands: Ligands that can coordinate to the central metal atom through two different donor atoms. E.g., $NO_2^-$ (can link via N or O), $SCN^-$ (can link via S or N).
• Coordination Number (CN): The total number of ligand donor atoms directly bonded to the central metal atom/ion in a complex.
  • For monodentate ligands, CN = number of ligands.
  • For polydentate ligands, CN = number of ligands $\times$ denticity.
  • Common CNs are 4 and 6.
    • CN=4: Tetrahedral or Square planar geometry.
    • CN=6: Octahedral geometry.
• Coordination Entity: The central metal atom/ion and the ligands attached to it, enclosed in a square bracket. E.g., $[Co(NH_3)_6]^{3+}$, $[Fe(CN)_6]^{4-}$.
• Coordination Sphere (Inner Sphere): The central metal atom and the ligands directly attached to it, which do not dissociate in solution.
• Counter Ions (Outer Sphere): The ions present outside the coordination sphere, which balance the charge of the coordination entity and dissociate in solution. E.g., $Cl^-$ in $[Co(NH_3)_6]Cl_3$.
• Homoleptic Complex: A complex where all the ligands attached to the central metal atom are identical. E.g., $[Co(NH_3)_6]^{3+}$.
• Heteroleptic Complex: A complex where more than one type of ligand is attached to the central metal atom. E.g., $[Co(NH_3)_4Cl_2]^+$.

3. Oxidation State of Central Metal Atom/Ion

The charge an atom would have if all bonds were ionic. It is calculated by summing the charges of the ligands and the overall charge of the complex.

Formula: $$ \text{Oxidation State of Metal} + \sum (\text{Charge of Ligand} \times \text{Number of Ligands}) = \text{Overall Charge of Complex} $$

Example: Calculate the oxidation state of Fe in $[Fe(CN)_6]^{4-}$. Let oxidation state of Fe be $x$. $CN^-$ is a monodentate ligand with charge -1. $x + (6 \times -1) = -4$ $x - 6 = -4$ $x = +2$

4. Nomenclature of Coordination Compounds (IUPAC Rules)

1. Cation before Anion: In both simple and complex compounds, the cation is named first, followed by the anion.
2. Ligands before Metal: Within the coordination sphere, ligands are named first in alphabetical order, followed by the central metal atom.
3. Naming Ligands:
   • Anionic Ligands: End with '-o'. E.g., chloride ($Cl^-$) becomes chloro, cyanide ($CN^-$) becomes cyano, hydroxide ($OH^-$) becomes hydroxo, sulfate ($SO_4^{2-}$) becomes sulfato, oxalato ($C_2O_4^{2-}$).
   • Neutral Ligands: Retain their common names, but some have special names:
     • $H_2O$: aqua
     • $NH_3$: ammine
     • $CO$: carbonyl
     • $NO$: nitrosyl
     • $CH_2CH_2(NH_2)_2$: ethylenediamine (en)
     • $C_5H_5N$: pyridine (py)
   • Cationic Ligands: End with '-ium'. E.g., $NO_2^+$ (nitronium). (Rarely encountered in NEET).
4. Prefixes for Number of Ligands:
   • If the ligand name is simple (e.g., chloro, ammine): di-, tri-, tetra-, penta-, hexa-
   • If the ligand name already contains a numerical prefix (e.g., ethylenediamine, triphenylphosphine) or is complex: bis-, tris-, tetrakis-
5. Oxidation State of Metal: Written in Roman numerals in parentheses immediately after the name of the central metal atom.
6. Naming the Central Metal Atom:
   • If the complex is a Cationic or Neutral Coordination Entity: The metal name is used as is. E.g., Cobalt, Platinum, Iron.
   • If the complex is an Anionic Coordination Entity: The suffix '-ate' is added to the name of the metal. For some metals, their Latin names are used.
     • Fe: Ferrate
     • Cu: Cuprate
     • Ag: Argentate
     • Au: Aurate
     • Pb: Plumbate
     • Sn: Stannate
     • Co: Cobaltate
     • Cr: Chromate
     • Pt: Platinate
7. Ambidentate Ligands: The donor atom is specified by an italicized letter after the name of the ligand. E.g., $NO_2^-$ can be nitro (via N) or nitrito-O (via O); $SCN^-$ can be thiocyanato (via S) or isothiocyanato (via N).
8. No Spaces: The entire name of the coordination entity is written as a single word.

Example 1: $[Co(NH_3)_5Cl]Cl_2$

• Coordination entity: $[Co(NH_3)_5Cl]^2+$ (cationic)
• Ligands: Pentaammine, chloro (alphabetical: ammine before chloro)
• Metal: Cobalt
• Oxidation state of Co: $x + (5 \times 0) + (1 \times -1) = +2 \Rightarrow x = +3$.
• Counter ion: Chloride
• Name: Pentaamminechlorocobalt(III) chloride

Example 2: $K_4[Fe(CN)_6]$

• Cation: Potassium
• Coordination entity: $[Fe(CN)_6]^{4-}$ (anionic)
• Ligand: Hexacyano
• Metal: Ferrate (since it's an anionic complex)
• Oxidation state of Fe: $4 + x + (6 \times -1) = 0 \Rightarrow x = +2$.
• Name: Potassium hexacyanoferrate(II)

5. Isomerism in Coordination Compounds

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

A. Structural Isomerism (Constitutional Isomerism)

Isomers have different connections between atoms.

1. Ionization Isomerism: Occurs when the counter ion in a complex salt is itself a potential ligand and can swap positions with a ligand inside the coordination sphere.
   • E.g., $[Co(NH_3)_5Br]SO_4$ (pentaamminebromocobalt(III) sulfate) and $[Co(NH_3)_5SO_4]Br$ (pentaamminesulfatocobalt(III) bromide).
     • The first gives $SO_4^{2-}$ ions in solution, the second gives $Br^-$ ions.
2. Hydrate Isomerism (Solvate Isomerism): Occurs when water molecules are involved as either a ligand or a molecule of crystallization.
   • E.g., $[Cr(H_2O)_6]Cl_3$ (violet), $[Cr(H_2O)_5Cl]Cl_2 \cdot H_2O$ (green), $[Cr(H_2O)_4Cl_2]Cl \cdot 2H_2O$ (dark green).
3. Linkage Isomerism: Occurs in complexes containing ambidentate ligands (ligands that can bind through two different donor atoms).
   • E.g., $[Co(NH_3)_5(NO_2)]^{2+}$ (pentaamminenitrocobalt(II), N-bonded) and $[Co(NH_3)_5(ONO)]^{2+}$ (pentaamminenitrito-O-cobalt(II), O-bonded).
4. Coordination Isomerism: Occurs in compounds where both cation and anion are complex entities, and there is an interchange of ligands between the cationic and anionic entities.
   • E.g., $[Co(NH_3)_6][Cr(CN)_6]$ and $[Cr(NH_3)_6][Co(CN)_6]$.
5. Ligand Isomerism: Occurs when the ligands themselves are isomers.
   • E.g., Complexes with propylenediamine (pn) and trimethylenediamine (tn) ligands.

B. Stereoisomerism (Space Isomerism)

Isomers have the same connections between atoms but different spatial arrangements.

1. Geometrical Isomerism (cis-trans isomerism): Arises due to the different spatial arrangements of ligands around the central metal atom.
   • Square planar complexes (CN=4): Exhibit cis-trans isomerism for $MA_2B_2$ type. A and B are ligands.
     • cis-isomer: Identical ligands are adjacent to each other ($90^\circ$).
     • trans-isomer: Identical ligands are opposite to each other ($180^\circ$).
     • E.g., $[Pt(NH_3)_2Cl_2]$ (cis-platin and trans-platin).
   • Octahedral complexes (CN=6): Exhibit cis-trans isomerism for $MA_4B_2$ type and $MA_3B_3$ type.
     • $MA_4B_2$ type: E.g., $[Co(NH_3)_4Cl_2]^+$
       • Two B ligands at $90^\circ$ (cis) or $180^\circ$ (trans).
     • $MA_3B_3$ type (fac-mer isomerism): E.g., $[Co(NH_3)_3Cl_3]$
       • fac (facial) isomer: Three identical ligands occupy adjacent positions at the corners of an octahedral face.
       • mer (meridional) isomer: Three identical ligands lie in a plane passing through the central metal atom.
   • Tetrahedral complexes: Do not exhibit geometrical isomerism because all positions are adjacent to each other.
2. Optical Isomerism (Enantiomerism): Occurs when a coordination compound is non-superimposable on its mirror image. Such isomers are called enantiomers and are optically active.
   • Chiral complexes: Optical isomers are possible when the complex and its mirror image are non-superimposable.
   • No plane of symmetry: A molecule with an axis of symmetry but no plane of symmetry can be chiral.
   • Common in octahedral complexes, especially those with bidentate ligands like 'en':
     • $[M(AA)_3]$ type: E.g., $[Co(en)_3]^{3+}$
     • $[M(AA)_2B_2]$ type: E.g., $[Co(en)_2Cl_2]^+$ (cis-isomer is optically active, trans-isomer is not due to a plane of symmetry).
   • Square planar complexes are rarely optical isomers. Tetrahedral complexes with four different monodentate ligands ($MABCD$) can exhibit optical isomerism (like organic compounds).

6. Valence Bond Theory (VBT)

Proposed by Linus Pauling, VBT explains the bonding in coordination compounds based on the overlap of atomic orbitals.

Postulates:

• The central metal atom provides vacant s, p, or d orbitals for hybridization.
• Ligands donate electron pairs to these vacant hybrid orbitals to form metal-ligand coordinate covalent bonds.
• The number of such hybrid orbitals equals the coordination number.
• The geometry of the complex is determined by the type of hybridization.
• Inner orbital complexes (low spin): Formed when the $d$ orbitals involved in hybridization are from the $(n-1)d$ subshell. E.g., $d^2sp^3$. Strong field ligands cause inner orbital complexes by forcing electron pairing.
• Outer orbital complexes (high spin): Formed when the $d$ orbitals involved are from the $nd$ subshell. E.g., $sp^3d^2$. Weak field ligands form outer orbital complexes.
• The magnetic properties (paramagnetic or diamagnetic) are determined by the number of unpaired electrons.
  • Unpaired electrons = paramagnetic.
  • All electrons paired = diamagnetic.
• Magnetic Moment ($\mu$): Calculated using the spin-only formula: $$ \mu = \sqrt{n(n+2)} \text{ BM (Bohr Magneton)} $$ where $n$ is the number of unpaired electrons.

Hybridization and Geometry Table:

Limitations of VBT:

• It does not explain color and spectra of complexes.
• It does not adequately explain the preference for specific geometries.
• It does not provide a quantitative explanation for the strength of ligands (strong vs. weak field).
• It does not predict distortions from ideal geometries.

7. Crystal Field Theory (CFT)

Proposed by Bethe and Van Vleck, CFT is an electrostatic model that explains the bonding, geometry, magnetic properties, and optical properties of coordination compounds.

Postulates:

• The bond between the metal and ligand is purely ionic, arising from electrostatic interaction between the metal ion and the ligands.
• Ligands are treated as point charges (anions) or dipoles (neutral molecules).
• Central metal d-orbitals are degenerate in an isolated gaseous metal atom.
• In a complex, the degeneracy of these d-orbitals is lifted due to the electrostatic field of the ligands, resulting in splitting of d-orbitals. This is called Crystal Field Splitting.
• The extent of splitting depends on the nature of the metal ion, its oxidation state, and the nature of the ligand.

A. Crystal Field Splitting in Octahedral Complexes

• In an octahedral environment, the six ligands approach the central metal ion along the x, y, and z axes.
• The $d_x^2-y^2$ and $d_z^2$ orbitals (axial orbitals, collectively called the $e_g$ set) have lobes directed along the axes, so they experience greater repulsion from the ligands.
• The $d_{xy}$, $d_{yz}$, and $d_{zx}$ orbitals (non-axial orbitals, collectively called the $t_{2g}$ set) have lobes directed between the axes, so they experience less repulsion.
• Result: The $e_g$ orbitals are raised in energy, and the $t_{2g}$ orbitals are lowered in energy.
• The energy difference between the $e_g$ and $t_{2g}$ sets is called the Crystal Field Splitting Energy ($\Delta_o$ for octahedral).
• The barycenter rule applies: $2(E_{e_g}) + 3(E_{t_{2g}}) = 0$.
  • $E_{e_g} = +0.6\Delta_o$
  • $E_{t_{2g}} = -0.4\Delta_o$

B. Crystal Field Splitting in Tetrahedral Complexes

• In a tetrahedral environment, the four ligands approach the central metal ion from alternate corners of a cube.
• The $t_2$ set ($d_{xy}, d_{yz}, d_{zx}$) are oriented closer to the direction of ligand approach than the $e$ set ($d_x^2-y^2, d_z^2$).
• Result: The $t_2$ orbitals are raised in energy, and the $e$ orbitals are lowered in energy.
• The energy difference is $\Delta_t$.
• $\Delta_t = (4/9)\Delta_o$.
• Since $\Delta_t$ is much smaller than $\Delta_o$, tetrahedral complexes are almost always high-spin (weak field ligands), as pairing energy (P) is usually greater than $\Delta_t$.

C. Crystal Field Splitting in Square Planar Complexes

• Square planar geometry can be considered as an octahedral complex where the two ligands along the z-axis are removed.
• The splitting is more complex, but generally, the $d_x^2-y^2$ orbital is highest in energy, followed by $d_{xy}$, then $d_{z^2}$, and finally $d_{xz}, d_{yz}$ (lowest).
• Square planar complexes are typically formed by $d^8$ metals (e.g., $Ni^{2+}, Pt^{2+}, Pd^{2+}$).
• They are almost always low-spin, due to large splitting energy.

D. Spectrochemical Series

An experimentally determined series that ranks ligands based on their ability to cause crystal field splitting. Strong field ligands cause large splitting ($\Delta > P$, favour pairing), while weak field ligands cause small splitting ($\Delta < P$, favour unpaired electrons).

$$ I^- < Br^- < S^{2-} < SCN^- < Cl^- < NO_3^- < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < EDTA^{4-} < NH_3 < en < CN^- < CO $$

• Weak field ligands: $I^-, Br^-, Cl^-, F^-, H_2O, OH^-, C_2O_4^{2-}$ (generally halides and oxygen donors)
  • High spin complexes (outer orbital, $sp^3d^2$ for CN=6)
  • $\Delta_o < P$ (pairing energy), electrons prefer to remain unpaired.
• Strong field ligands: $CN^-, CO, en, NH_3$ (generally carbon and nitrogen donors)
  • Low spin complexes (inner orbital, $d^2sp^3$ for CN=6)
  • $\Delta_o > P$, electrons pair up in the lower energy $t_{2g}$ orbitals.

E. Crystal Field Stabilization Energy (CFSE)

CFSE is the stabilization energy gained by the central metal ion when its d-orbitals are split in the ligand field.

Formula for Octahedral Complexes: $$ CFSE = [-0.4x + 0.6y]\Delta_o + nP $$ where $x$ = number of electrons in $t_{2g}$ orbitals, $y$ = number of electrons in $e_g$ orbitals, $n$ = number of extra electron pairs formed due to ligand field (pairing energy, P).

Formula for Tetrahedral Complexes: $$ CFSE = [-0.6x + 0.4y]\Delta_t + nP $$ where $x$ = number of electrons in $e$ orbitals, $y$ = number of electrons in $t_2$ orbitals.

Color of Coordination Compounds:

• The color of transition metal complexes arises from d-d transitions.
• When white light falls on a complex, it absorbs light of certain wavelengths from the visible region.
• The complementary color is transmitted or reflected, which gives the observed color.
• The energy of the absorbed photon corresponds to $\Delta_o$ or $\Delta_t$. $$ \Delta E = h\nu = hc/\lambda $$
• Stronger ligands lead to larger $\Delta$, absorbing higher energy (shorter wavelength) light, and thus transmitting lower energy (longer wavelength) light.

Example:

• $[Ti(H_2O)_6]^{3+}$ is purple. $Ti^{3+}$ is $d^1$. It absorbs green-yellow light, reflecting purple.
• $[Ni(H_2O)_6]^{2+}$ is green. $Ni^{2+}$ is $d^8$.
• $[Cu(NH_3)_4]^{2+}$ is deep blue.

Limitations of CFT:

• Treats ligands as point charges, not considering their covalent character.
• Doesn't explain bonding in complexes having $\pi$-bonding ligands ($CO, CN^-$).
• Doesn't explain why neutral ligands like CO and $NH_3$ are strong field ligands, while $OH^-$ and $F^-$ (anions) are weak field ligands. (This is addressed by Ligand Field Theory, which is beyond NEET scope).

8. Bonding in Metal Carbonyls

Metal carbonyls are coordination compounds where carbon monoxide (CO) acts as a ligand. They are typically zero oxidation state metals.

• Synergic Bonding: The metal-carbon bond in metal carbonyls possesses both $\sigma$ and $\pi$ character.
  • $\sigma$-bond: Formed by the donation of electrons from the lone pair of the carbonyl carbon into a vacant orbital of the metal. (Ligand to metal).
  • $\pi$-bond: Formed by the donation of a pair of electrons from a filled d-orbital of the metal into the vacant antibonding ($\pi^*$) orbital of the carbonyl ligand. (Metal to ligand, back bonding).
• This synergic effect (M-C $\sigma$ bond and M-C $\pi$ back bond) strengthens the metal-ligand bond, resulting in stable complexes.
• Back bonding reduces the bond order of CO and increases the bond order of M-C. This is reflected in the stretching frequencies of CO.

Examples:

• $Ni(CO)_4$ (tetrahedral)
• $Fe(CO)_5$ (trigonal bipyramidal)
• $Cr(CO)_6$ (octahedral)

9. Stability of Coordination Compounds

• Thermodynamic Stability (Stability Constant): Refers to the equilibrium constant for the formation of a complex. A higher stability constant ($K_{stab}$ or $\beta$) indicates a more stable complex. $$ M^{n+} + L \rightleftharpoons ML^{n+} \quad K_1 = \frac{[ML^{n+}]}{[M^{n+}][L]} $$ $$ ML^{n+} + L \rightleftharpoons ML_2^{n+} \quad K_2 = \frac{[ML_2^{n+}]}{[ML^{n+}][L]} $$ And so on for stepwise formation constants. Overall stability constant: $\beta_n = K_1 K_2 ... K_n$.

• Factors Affecting Stability:

  1. Nature of the Central Metal Ion:
     • Higher charge on the metal ion increases stability. ($Fe^{3+}$ complexes are more stable than $Fe^{2+}$).
     • Smaller size of the metal ion increases stability (higher charge density).
  2. Nature of the Ligand:
     • Basicity of ligand: More basic ligands (stronger Lewis bases) form more stable complexes.
     • Chelate Effect: Polydentate ligands (chelating ligands) form more stable complexes than monodentate ligands with similar donor atoms. This is primarily due to a favorable entropy change ($\Delta S > 0$), as one polydentate ligand replaces multiple monodentate ones, increasing the number of molecules/ions in solution.
       • E.g., $[Ni(en)_3]^{2+}$ is more stable than $[Ni(NH_3)_6]^{2+}$.
  3. Nature of the Solvent: Solvent molecules can compete with ligands.

• Kinetic Stability (Lability and Inertness): Refers to the rate at which a complex undergoes ligand substitution reactions.

  • Labile complexes: Undergo rapid ligand substitution.
  • Inert complexes: Undergo slow ligand substitution.
  • Note: Thermodynamic stability and kinetic stability are not directly related. A thermodynamically stable complex can be labile (e.g., $[Mn(CN)_6]^{3-}$), and a thermodynamically unstable complex can be inert (e.g., $[Cr(CO)_6]$).

Important Formulas and Concepts Summary

• Oxidation State Calculation: Metal charge + $\sum$(ligand charge * no. of ligands) = complex charge
• Magnetic Moment (spin-only): $\mu = \sqrt{n(n+2)}$ BM
• CFT Splitting Energies:
  • Octahedral: $E_{e_g} = +0.6\Delta_o$, $E_{t_{2g}} = -0.4\Delta_o$
  • Tetrahedral: $E_{t_2} = +0.4\Delta_t$, $E_e = -0.6\Delta_t$ (Note: signs reversed compared to octahedral, and $e$ is lower)
  • Relation: $\Delta_t = (4/9)\Delta_o$
• CFSE: Refer to formulas above based on electron configuration.
• Color Correlation: $\Delta E = hc/\lambda$. Larger $\Delta$, shorter $\lambda_{absorbed}$, longer $\lambda_{transmitted}$ (complementary color).
• Isomerism Types: Structural (Ionization, Hydrate, Linkage, Coordination, Ligand), Stereoisomerism (Geometrical, Optical).
• Spectrochemical Series: $I^- < Br^- < S^{2-} < SCN^- < Cl^- < NO_3^- < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < EDTA^{4-} < NH_3 < en < CN^- < CO$

Worked Examples

Example 1: Oxidation State and Coordination Number

Determine the oxidation state of the central metal ion and the coordination number in the following complexes: a) $[Cr(NH_3)_4Cl_2]Br$ b) $K_3[Fe(C_2O_4)_3]$

Solution: a) $[Cr(NH_3)_4Cl_2]Br$ * Let the oxidation state of Cr be $x$. * $NH_3$ is a neutral ligand (0 charge). * $Cl^-$ is an anionic ligand (-1 charge). * $Br^-$ is a counter ion (-1 charge). * The overall charge of the coordination sphere $[Cr(NH_3)_4Cl_2]$ must be +1 to balance the $Br^-$. * $x + (4 \times 0) + (2 \times -1) = +1$ * $x - 2 = +1 \Rightarrow x = +3$. * Oxidation state of Cr = +3. * Coordination Number (CN): Number of donor atoms. $NH_3$ is monodentate, $Cl^-$ is monodentate. * CN = $4 (\text{from } NH_3) + 2 (\text{from } Cl^-) = 6$. * Coordination Number = 6.

b) $K_3[Fe(C_2O_4)_3]$ * Let the oxidation state of Fe be $x$. * $K^+$ is a counter ion (+1 charge). Three $K^+$ make total charge +3. * The overall charge of the coordination sphere $[Fe(C_2O_4)_3]$ must be -3 to balance the three $K^+$. * $C_2O_4^{2-}$ (oxalate) is a bidentate anionic ligand (-2 charge). * $x + (3 \times -2) = -3$ * $x - 6 = -3 \Rightarrow x = +3$. * Oxidation state of Fe = +3. * Coordination Number (CN): $C_2O_4^{2-}$ is a bidentate ligand. * CN = $3 (\text{ligands}) \times 2 (\text{denticity}) = 6$. * Coordination Number = 6.

Example 2: IUPAC Nomenclature

Write the IUPAC names for the following complexes: a) $[Co(en)_2Cl_2]Cl$ b) $Na_3[CoF_6]$

Solution: a) $[Co(en)_2Cl_2]Cl$ * Cationic complex: $[Co(en)_2Cl_2]^+$ * Ligands: ethylenediamine (en, neutral), chloro ($Cl^-$, anionic). Alphabetical order: chloro, ethylenediamine. * Prefixes: 'bis' for 'en' (because 'ethylenediamine' already has a numerical prefix), 'di' for 'chloro'. * Metal: Cobalt (since it's a cationic complex). * Oxidation state of Co: $x + (2 \times 0) + (2 \times -1) = +1 \Rightarrow x - 2 = +1 \Rightarrow x = +3$. * Counter ion: Chloride. * Name: Dichlorobis(ethylenediamine)cobalt(III) chloride.

b) $Na_3[CoF_6]$ * Cation: Sodium. * Anionic complex: $[CoF_6]^{3-}$ * Ligand: fluoro ($F^-$, anionic). Prefix: 'hexa'. * Metal: Cobaltate (since it's an anionic complex). * Oxidation state of Co: $3 \times (+1) + x + (6 \times -1) = 0 \Rightarrow 3 + x - 6 = 0 \Rightarrow x = +3$. * Name: Sodium hexafluorocobaltate(III).

Example 3: VBT and Magnetic Properties

Predict the hybridization, magnetic nature, and geometry of $[Ni(CN)_4]^{2-}$.

Solution:

• Central metal ion: $Ni^{2+}$
• Electronic configuration of Ni: $[Ar] 3d^8 4s^2$
• Electronic configuration of $Ni^{2+}$: $[Ar] 3d^8$

  3d  [ ↑↓  ↑↓  ↑↓  ↑   ↑ ]
  4s  [   ]
  4p  [   ] [   ] [   ]
  

• Ligand: $CN^-$. From the spectrochemical series, $CN^-$ is a strong field ligand. It causes pairing of electrons.
• After pairing:

  3d  [ ↑↓  ↑↓  ↑↓  ↑↓  _ ]  (one empty d orbital for hybridization)
  4s  [   ]
  4p  [   ] [   ] [   ]
  

• Since CN=4, and $CN^-$ is a strong field ligand, $dsp^2$ hybridization is expected. This happens when the empty 3d orbital, 4s, and two 4p orbitals participate.
• Hybridization: $dsp^2$
• Geometry: Square planar
• Magnetic nature: All electrons are paired.
• Magnetic moment: $n=0$ (number of unpaired electrons). $\mu = \sqrt{0(0+2)} = 0$ BM. So, it is diamagnetic.

Example 4: CFT and CFSE

Calculate the CFSE for $d^6$ octahedral complex in both high spin and low spin states.

Solution: For an octahedral complex, the $d$-orbitals split into $t_{2g}$ (lower energy, -0.4$\Delta_o$) and $e_g$ (higher energy, +0.6$\Delta_o$).

a) High Spin ($d^6$)

• High spin occurs with weak field ligands, where electrons prefer to occupy separate orbitals before pairing up.
• Electron configuration: $t_{2g}^4 e_g^2$

  e_g         [ ↑   ↑ ]
           /       \
          /         \
  t_2g   [ ↑↓  ↑   ↑ ]
  

• Number of electrons in $t_{2g}$ ($x$) = 4
• Number of electrons in $e_g$ ($y$) = 2
• Number of extra pairs ($n$) = 0 (as pairing occurs after filling all 5 orbitals in high spin, no extra pairing is induced by ligands)
• CFSE = $[-0.4x + 0.6y]\Delta_o + nP$ CFSE = $[-0.4(4) + 0.6(2)]\Delta_o + 0P$ CFSE = $[-1.6 + 1.2]\Delta_o = -0.4\Delta_o$

b) Low Spin ($d^6$)

• Low spin occurs with strong field ligands, where electrons pair up in the lower energy $t_{2g}$ orbitals before occupying $e_g$ orbitals.
• Electron configuration: $t_{2g}^6 e_g^0$

  e_g         [ _   _ ]
           /       \
          /         \
  t_2g   [ ↑↓  ↑↓  ↑↓ ]
  

• Number of electrons in $t_{2g}$ ($x$) = 6
• Number of electrons in $e_g$ ($y$) = 0
• Number of extra pairs ($n$) = 2 (in $d^6$, three electrons are paired by Hund's rule, a fourth pairing is formed in the $t_{2g}$ subshell. Here, the $t_{2g}$ subshell receives 6 electrons, so there are 3 pairs. In the free ion, $d^6$ has 4 unpaired electrons (2 pairs), but in low spin, it has 0 unpaired electrons (3 pairs). So 1 extra pair is formed. Wait, let me re-evaluate, it's easier to think by comparison to high spin. In high spin $d^6$, there are 4 unpaired electrons and 1 pair. In low spin $d^6$, there are 0 unpaired electrons and 3 pairs. So there are 2 additional pairs formed due to the ligand field that would not be present in the high spin configuration.)
• CFSE = $[-0.4(6) + 0.6(0)]\Delta_o + 2P$ (Here, P refers to pairing energy, 2P because two extra pairs are formed compared to the free ion (or high-spin case)). CFSE = $[-2.4 + 0]\Delta_o + 2P = -2.4\Delta_o + 2P$

Example 5: Isomerism

How many geometrical isomers are possible for $[Co(en)_2Cl_2]^+$? Draw their structures and indicate if any are optically active.

Solution:

• The complex is of type $M(AA)_2B_2$, where 'AA' is a bidentate ligand (en) and 'B' is a monodentate ligand (Cl). This is an octahedral complex.

• Geometrical Isomers:

  1. cis-isomer: The two chloride ligands ($Cl$) are adjacent to each other ($90^\circ$). The two 'en' ligands are also adjacent.

     • This isomer lacks a plane of symmetry and is chiral. Therefore, it will exist as a pair of enantiomers (d- and l-forms) and will be optically active.

     •   Cl
        / \
       Co --en
          /     \
       

     en ----- Cl

     (Imagine Co at the center, Cl's on top and front right, en on front left and back left/right. Hard to draw without proper 3D rendering. The two Cls are at 90 degrees)
     

  2. trans-isomer: The two chloride ligands ($Cl$) are opposite to each other ($180^\circ$).

     • This isomer possesses a plane of symmetry (passing through Co and the two 'en' ligands, bisecting the $Cl-Co-Cl$ axis). Therefore, it is achiral and optically inactive.

     •   en
        / \
       Cl -- Co -- Cl
          /     \
       

     en

     (Cl's on top and bottom, en's in the plane).
     

• Total number of geometrical isomers = 2 (cis and trans).

• Total number of stereoisomers (including optical) = 3 (cis-d, cis-l, trans).

Example 6: Color and Spectrochemical Series

The aqueous solution of $[Co(H_2O)_6]^{3+}$ is pink, while that of $[Co(CN)_6]^{3-}$ is blue. Explain why.

Solution:

1. Central Metal Ion: In both complexes, the central metal ion is $Co^{3+}$.

   • Electronic configuration of Co: $[Ar] 3d^7 4s^2$
   • Electronic configuration of $Co^{3+}$: $[Ar] 3d^6$

2. Ligands and Crystal Field Splitting:

   • For $[Co(H_2O)_6]^{3+}$: Water ($H_2O$) is a weak field ligand. For $d^6$ in an octahedral field, a weak field ligand causes high spin pairing of electrons.
     • Electron configuration: $t_{2g}^4 e_g^2$.
     • The splitting energy ($\Delta_{o1}$) is relatively small.
   • For $[Co(CN)_6]^{3-}$: Cyanide ($CN^-$) is a strong field ligand. For $d^6$ in an octahedral field, a strong field ligand causes low spin pairing of electrons.
     • Electron configuration: $t_{2g}^6 e_g^0$.
     • The splitting energy ($\Delta_{o2}$) is relatively large.

3. Color Explanation (d-d transitions):

   • The color of coordination compounds arises from the absorption of light in the visible region, causing d-d electron transitions. The absorbed energy corresponds to the crystal field splitting energy ($\Delta_o$).
   • We know $\Delta E = hc/\lambda$. A larger $\Delta_o$ means absorption of higher energy (shorter wavelength) light.
   • From the spectrochemical series, $CN^-$ is a stronger field ligand than $H_2O$.
   • Therefore, $\Delta_{o2} (for CN^-) > \Delta_{o1} (for H_2O)$.
   • When white light passes through $[Co(H_2O)_6]^{3+}$, it absorbs lower energy light (longer wavelength, e.g., green-yellow part of the spectrum). The transmitted light is a mix of the remaining colors, appearing pink.
   • When white light passes through $[Co(CN)_6]^{3-}$, it absorbs higher energy light (shorter wavelength, e.g., orange-red part of the spectrum). The transmitted light is a mix of the remaining colors, appearing blue.

Thus, the difference in ligand strength leads to different crystal field splitting energies and, consequently, absorption of different wavelengths of light, resulting in different observed colors.

Memorization Tricks and Mnemonics

1. Spectrochemical Series (Weak to Strong):

   • "I Brought Some Salty Clam Noodles For Our Ox With No Ethylene Narration, Eating Crabs Carelessly Only (rarely found in NEET: $NO^+, PPh_3$, etc.)"
   • Or, simply remember that halides and oxygen donors are generally weak field, while carbon and nitrogen donors are strong field.
   • $I^- < Br^- < S^{2-} < SCN^- < Cl^- < NO_3^- < F^- < OH^- < C_2O_4^{2-} < H_2O < NCS^- < EDTA^{4-} < NH_3 < en < CN^- < CO$.

2. Naming Ambidentate Ligands:

   • N-bonded = Nitro. $\rightarrow NO_2^-$
   • O-bonded = Onitrito. $\rightarrow ONO^-$
   • S-bonded = Thiocyanato. $\rightarrow SCN^-$
   • N-bonded = Isothiocyanato. $\rightarrow NCS^-$

3. VBT Hybridization & Geometry (CN=4 & CN=6):

   • CN=4:
     • If strong field ligand $\Rightarrow dsp^2$ (Square planar).
     • If weak field ligand $\Rightarrow sp^3$ (Tetrahedral).
   • CN=6:
     • If strong field ligand $\Rightarrow d^2sp^3$ (Inner orbital, Octahedral).
     • If weak field ligand $\Rightarrow sp^3d^2$ (Outer orbital, Octahedral).
   • This general rule works for most $3d$ series elements. For $4d$ and $5d$ elements, strong field ligand behavior is almost universal.

4. $MA_3B_3$ Isomers (Octahedral):

   • Fac (facial): Like a face on a cube. Three identical ligands are on one face of the octahedron. They are mutually at $90^\circ$.
   • Mer (meridional): All three identical ligands are in the same plane, i.e., around the 'meridional' line (like the equator if you consider the central metal). Two are $180^\circ$ apart, one is $90^\circ$ from both.

Common Mistakes to Avoid

1. Oxidation State Calculation Errors: Double-check charges of neutral and anionic ligands. Remember that a complex's overall charge must be balanced by counter ions.
2. Incorrect Coordination Number: For polydentate ligands, remember to multiply the number of ligands by their denticity. (e.g., three oxalate ligands ($C_2O_4^{2-}$), denticity 2, so CN=6, not 3).
3. Nomenclature Errors:
   • Not using prefixes 'bis', 'tris' for complex ligand names (e.g., ethylenediamine).
   • Incorrectly assigning '-ate' suffix to metal (only for anionic complexes).
   • Incorrect Roman numeral for oxidation state.
   • Not naming ligands alphabetically.
4. Confusing Cis-Trans with Fac-Mer: Cis-trans applies to $MA_2B_4$ or $MA_4B_2$ types (where 'A' and 'B' are single atoms/groups). Fac-mer applies to $MA_3B_3$ type.
5. Optical Isomerism:
   • Tetrahedral complexes only show optical isomerism if they have four different monodentate ligands (i.e., NO plane of symmetry similar to chiral organic molecules).
   • Square planar complexes rarely show optical isomerism.
   • Octahedral complexes with bidentate ligands, especially $[M(AA)_3]$ types and cis- $[M(AA)_2B_2]$ types, are often optically active. The trans- isomer of $[M(AA)_2B_2]$ is usually optically inactive due to a plane of symmetry.
6. VBT vs. CFT: Don't mix postulates. Remember VBT focuses on orbital overlap and hybridization; CFT on electrostatic splitting of d-orbitals. VBT explains magnetic properties and geometry; CFT explains color, CFSE, and provides a quantitative aspect to ligand field strength.
7. High Spin vs. Low Spin:
   • Occurs for $d^4, d^5, d^6, d^7$ configurations in Octahedral complexes.
   • Crucially depends on the ligand field strength ($\Delta_o$ vs. pairing energy P).
   • Tetrahedral complexes are almost always high spin because $\Delta_t$ is very small.
8. Crystal Field Splitting Diagrams: For octahedral, $e_g$ is higher; for tetrahedral, $t_2$ is higher. Relationship between $\Delta_o$ and $\Delta_t$ is $\Delta_t = (4/9)\Delta_o$.

Previous Year Question Patterns

NEET questions from Coordination Compounds typically cover:

• Nomenclature: Writing names from formulas and vice-versa. (Very common)
• Isomerism: Identifying isomer types, drawing isomers, or determining the number of possible isomers (especially geometrical and optical). (Common)
• VBT-based questions: Predicting hybridization, geometry, and magnetic properties (paramagnetic/diamagnetic, magnetic moment calculation) for given complexes. (Very common)
• CFT-based questions: Explaining the color of complexes, calculating CFSE, or comparing weak vs. strong field ligands based on electron configuration and splitting. (Common)
• Werner's Theory: Questions mostly involve primary/secondary valency and their implications. (Less common, but fundamental)
• Stability of Complexes: Factors affecting stability, chelate effect. (Occasional)

Example PYQ types (conceptual):

• "Which of the following complexes exhibits linkage isomerism?" (Requires knowledge of ambidentate ligands)
• "Which of the following is an outer orbital complex and diamagnetic?" (Requires VBT application with strong/weak field ligands)
• "The magnetic moment of $[MnBr_4]^{2-}$ is 5.92 BM. Predict its geometry and hybridization." (Work backward from magnetic moment to unpaired electrons, then use VBT rules)
• "Explain why $[NiCl_4]^{2-}$ is tetrahedral, while $[Ni(CN)_4]^{2-}$ is square planar." (Ligand strength, VBT)
• "Arrange the following ligands in increasing order of crystal field splitting energy: CO, Cl-, H2O." (Spectrochemical series)

Revision Checklist

1. Werner's Theory: Primary vs. Secondary Valency.
2. Terminology: Metal, ligand, CN, coordination sphere, homoleptic/heteroleptic.
3. Oxidation State of Metal.
4. IUPAC Nomenclature: All rules, practice with examples.
5. Isomerism:
   • Structural: Ionization, Hydrate, Linkage, Coordination. (Definitions & examples)
   • Stereo: Geometrical (cis/trans, fac/mer), Optical (enantiomers, conditions for chirality). (Drawings & conditions)
6. Valence Bond Theory (VBT):
   • Postulates.
   • Hybridization ($sp^3, dsp^2, d^2sp^3, sp^3d^2$) and geometries.
   • Magnetic properties ($\mu = \sqrt{n(n+2)}$).
   • Weak vs. strong field ligands (pairing/no pairing).
   • Limitations.
7. Crystal Field Theory (CFT):
   • Postulates.
   • Orbital splitting diagrams (Octahedral, Tetrahedral).
   • Spectrochemical series (memorize critical ligands).
   • High spin vs. low spin conditions ($P$ vs. $\Delta$).
   • CFSE calculation.
   • Color of complexes (d-d transitions, complementary colors, $\Delta E = hc/\lambda$).
   • Limitations.
8. Bonding in Metal Carbonyls: Synergic bonding (sigma donation, pi back-donation).
9. Stability of Complexes: Thermodynamic vs. Kinetic, Chelate Effect.

Last-Day Strategy

• Quick Scan through Notes: Focus on bolded terms, formulas, and tables.
• Review all Mnemonics: Especially for spectrochemical series and naming rules.
• Solve 2-3 previous NEET questions: Pick questions that cover different aspects (e.g., one on VBT, one on nomenclature, one on isomerism).
• Focus on Diagrams: Quickly redraw the d-orbital splitting diagrams for octahedral and tetrahedral complexes to reinforce understanding.
• Check common mistakes: Remind yourself of the pitfalls to avoid.
• Don't learn new concepts: Stick to what you've studied.
• Visualize structures: Especially for geometrical and optical isomerism, mentally rehearse what the structures look like.