Class 12 Chemistry – Coordination Compounds
Quick Revision Notes
1. Basic Terms
Coordination compound: A compound in which a
central metal atom/ion is surrounded by ions or molecules
called ligands, which donate electron pairs to the metal.
Example:
\[
[Co(NH_3)_6]Cl_3
\]
Central metal ion → Co³⁺
Ligand → NH₃
Coordination number → 6
Coordination sphere → [Co(NH₃)₆]³⁺
Counter ions → 3Cl⁻
2. Important Definitions
Term Meaning
Ligand Ion/molecule that donates an
electron pair to metal
Coordination number Number of donor atoms
directly attached to central metal
Coordination sphere Species written inside
square brackets
Oxidation state Charge on metal after considering
ligand charges
Denticity Number of donor atoms of a ligand attached
to metal
Chelate Ring formed when a multidentate ligand
attaches to metal
3. Types of Ligands
According to denticity:
Monodentate: One donor atom
Examples: NH₃, H₂O, Cl⁻, CN⁻
Bidentate: Two donor atoms
Examples: en (ethane-1,2-diamine), \(C_2O_4^{2-}\)
Polydentate: More than two donor atoms
Example: EDTA⁴⁻
Ambidentate ligands: Can coordinate through two different atoms.
Examples:
\(NO_2^-\) → nitro / nitrito
\(SCN^-\) → thiocyanato / isothiocyanato
4. Werner's Theory
Werner proposed two types of valencies:
Primary valency
Corresponds to oxidation state.
Ionisable.
Satisfied by negative ions.
Secondary valency
Corresponds to coordination number.
Non-ionisable.
Satisfied by ligands.
Example:
[Co(NH_3)_6]Cl_3
Primary valency = 3
Secondary valency = 6
5. Coordination Number
Coordination number = Number of donor atoms directly
bonded to the central metal ion.
Examples:
[Co(NH_3)_6]^{3+}
CN = 6
[PtCl_4]^{2-}
CN = 4
For:
[Co(en)_3]^{3+}
Each en is bidentate.
CN = 3times2=6
6. Oxidation State
Use:
Oxidation state of metal}+charges of ligands
=charge on complex
Example:
\[
[Fe(CN)_6]^{4-}
\]
Let oxidation state of Fe = x.
\[
x+6(-1)=-4
\]
\[
x=+2
\]
Therefore, Fe = +2.
---
7. Nomenclature
Basic order:
Ligands + metal + oxidation state
Important ligand names:
NH₃ → ammine
H₂O → aqua
CO → carbonyl
NO → nitrosyl
Cl⁻ → chlorido
Br⁻ → bromido
OH⁻ → hydroxido
CN⁻ → cyanido
\(C_2O_4^{2-}\) → oxalato
Prefixes:
2 → di
3 → tri
4 → tetra
5 → penta
6 → hexa
Example:
\[
[Co(NH_3)_6]Cl_3
\]
Hexaamminecobalt(III) chloride
For an anionic complex, metal name ends in -ate.
\[
K_4[Fe(CN)_6]
\]
Potassium hexacyanidoferrate(II)
---
8. Isomerism
Coordination compounds show:
A. Structural isomerism
1. Ionisation isomerism
2. Hydrate/Solvate isomerism
3. Linkage isomerism
4. Coordination isomerism
B. Stereoisomerism
1. Geometrical isomerism
cis
trans
2. Optical isomerism
Non-superimposable mirror images
Called enantiomers
Example:
\[
[Pt(NH_3)_2Cl_2]
\]
shows cis-trans isomerism.
---
9. Valence Bond Theory (VBT)
VBT explains:
Hybridisation
Geometry
Magnetic nature
Common hybridisations:
Hybridisation Geometry
\(sp^3\) Tetrahedral
\(dsp^2\) Square planar
\(sp^3d^2\) Octahedral – outer orbital
\(d^2sp^3\) Octahedral – inner orbital
---
10. Strong and Weak Ligands
Weak-field ligands generally do not cause pairing.
Examples:
\[
F^-, Cl^-, Br^-, I^-, H_2O
\]
Strong-field ligands cause pairing of electrons.
Examples:
\[
CN^-, CO, NH_3
\]
A useful simplified spectrochemical series:
\[
I^-<Br^-<Cl^-<F^-<H_2O<NH_3<CN^-<CO
\]
---
11. Inner and Outer Orbital Complexes
Inner orbital complex:
Uses \((n-1)d\) orbitals.
Example:
\[
d^2sp^3
\]
Outer orbital complex:
Uses \(nd\) orbitals.
Example:
\[
sp^3d^2
\]
---
12. Crystal Field Theory (CFT)
According to CFT, ligands approach the metal ion and cause splitting of d-orbitals.
Octahedral complex
Five d-orbitals split into:
Lower energy → \(t_{2g}\)
Higher energy → \(e_g\)
Energy gap = \(\Delta_o\)
Tetrahedral complex
Lower energy → \(e\)
Higher energy → \(t_2\)
Energy gap = \(\Delta_t\)
\[
\Delta_t \approx \frac{4}{9}\Delta_o
\]
---
13. Magnetic Properties
Paramagnetic: Has one or more unpaired electrons.
Diamagnetic: All electrons are paired.
Magnetic moment:
\[
\boxed{\mu=\sqrt{n(n+2)}\ BM}
\]
where n = number of unpaired electrons.
Examples:
n = 0 → 0 BM
n = 1 → 1.73 BM
n = 2 → 2.83 BM
n = 3 → 3.87 BM
n = 4 → 4.90 BM
n = 5 → 5.92 BM
---
14. Colour of Coordination Compounds
Colour is generally due to d–d transitions.
When an electron absorbs energy, it moves from a lower-energy d-orbital to a higher-energy d-orbital.
The absorbed wavelength determines the colour observed.
Important: \(d^0\) and \(d^{10}\) complexes generally do not show d–d transitions.
---
15. Stability of Coordination Compounds
Stability is related to the formation/stability constant.
For:
\[
M+4L\rightleftharpoons ML_4
\]
\[
K_f=\frac{[ML_4]}{[M][L]^4}
\]
Higher \(K_f\) generally means greater stability.
---
16. Chelate Effect
Complexes containing multidentate ligands are generally more stable than comparable complexes containing monodentate ligands.
Example:
\[
[Ni(en)_3]^{2+}
\]
is a chelate complex.
Reason: Formation of stable rings and favourable entropy change.
---
17. Applications
Coordination compounds are important in:
Biological systems: Haemoglobin, chlorophyll, vitamin B₁₂
Medicine: Cisplatin
Metallurgy: Extraction of metals
Qualitative analysis: Detection/separation of metal ions
Photography: Silver complexes
Electroplating: Metal complexes
---
⭐ Must-Remember for Exams
1. Ox
idation state ≠ coordination number.
2. Denticity counts donor atoms, not ligands.
3. Strong-field ligands → pairing → often low spin.
4. Weak-field ligands → less pairing → often high spin.
5. \(\boxed{\mu=\sqrt{n(n+2)}\ BM}\)
6. Octahedral: \(\boxed{t_{2g}<e_g}\)
7. Tetrahedral: \(\boxed{e<t_2}\)
8. \(\boxed{\Delta_t=\frac49\Delta_o}\)
9. Anionic complex → metal name generally ends in -ate.
10. Chelating ligands form rings with the central metal ion.
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