NCERT Solutions for Class 12 Chemistry Chapter 5 Coordination Compounds
NCERT Solutions for Class 12 Chemistry Chapter 5 Coordination Compounds provides step-by-step explanations covering Werner's theory, ligands, IUPAC nomenclature, isomerism, Valence Bond Theory (VBT), Crystal Field Theory (CFT), and practical applications of coordination complex compounds.
Overview of Topics & Subtopics (Click to Jump)
1. Werner's Theory, Terms & IUPAC Nomenclature
Short Note:
• Werner's Theory: Central metal atoms exhibit two types of valencies: Primary Valency (ionizable, corresponds to oxidation state) and Secondary Valency (non-ionizable, corresponds to coordination number, fixed directional geometry).
• Important Terms:
- Central Atom/Ion: Lewis acid that accepts lone pairs of electrons from donor ligands.
- Ligands: Ion/molecule donating electron pairs (Lewis base). Classified into monodentate ($\text{Cl}^-, \text{H}_2\text{O}$), didentate ($\text{ox}^{2-}, \text{en}$), polydentate ($\text{EDTA}^{4-}$), and ambidentate ($\text{NO}_2^-/\text{ONO}^-, \text{SCN}^-/\text{NCS}^-$).
- Chelate Effect: Ring formation by di-/polydentate ligands significantly enhances complex stability.
• IUPAC Nomenclature Rules: Cation is named first. Ligands are named in alphabetical order followed by central metal name and its oxidation state in Roman numerals in parentheses. (If complex ion is anionic, metal name ends in -ate).
2. Isomerism in Coordination Compounds
Short Note:
• Structural Isomerism:
- Ionisation Isomerism: Exchange of counter ions with ligating ions inside coordination sphere (e.g., $[\text{Co(NH}_3)_5\text{SO}_4]\text{Br}$ vs $[\text{Co(NH}_3)_5\text{Br}]\text{SO}_4$).
- Solvate/Hydrate Isomerism: Water acts as ligand vs crystal water (e.g., $[\text{Cr(H}_2\text{O})_6]\text{Cl}_3$ vs $[\text{Cr(H}_2\text{O})_5\text{Cl}]\text{Cl}_2\cdot\text{H}_2\text{O}$).
- Linkage Isomerism: Involves ambidentate ligands bonding through different donor atoms (e.g., $-\text{NO}_2$ vs $-\text{ONO}$).
- Coordination Isomerism: Interchange of ligands between cationic and anionic entities of different metal ions.
• Stereoisomerism:
- Geometrical Isomerism: Cis/Trans positioning of ligands in square planar ($\text{MA}_2\text{B}_2$) and octahedral ($\text{MA}_4\text{B}_2$) complexes, or Facial ($fac$) / Meridional ($mer$) in $\text{MA}_3\text{B}_3$ complexes.
- Optical Isomerism: Non-superimposable mirror images ($d$- and $l$- forms), typically seen in chiral octahedral complexes containing chelate rings like $[\text{Co(en)}_3]^{3+}$.
3. Bonding Theories (VBT & CFT) & Applications
Short Note:
• Valence Bond Theory (VBT): Explains geometry and magnetic behaviour via orbital hybridization:
- $sp^3$ (Tetrahedral), $dsp^2$ (Square planar), $d^2sp^3$ (Inner orbital Octahedral), $sp^3d^2$ (Outer orbital Octahedral).
- Strong field ligands ($\text{CN}^-, \text{CO}, \text{en}$) force electron pairing (diamagnetic / low-spin); weak field ligands ($\text{F}^-, \text{Cl}^-, \text{H}_2\text{O}$) do not (paramagnetic / high-spin).
• Crystal Field Theory (CFT): Treats ligands as point charges causing electrostatic repulsion that splits degenerate $d$-orbitals:
- Octahedral Splitting ($\Delta_o$): $d$-orbitals split into lower energy $t_{2g}$ ($d_{xy}, d_{yz}, d_{zx}$) and higher energy $e_g$ ($d_{x^2-y^2}, d_{z^2}$) sets.
- Spectrochemical Series: Orders ligands by field strength: $\text{I}^- < \text{Br}^- < \text{Cl}^- < \text{F}^- < \text{OH}^- < \text{H}_2\text{O} < \text{NH}_3 < \text{en} < \text{CN}^- < \text{CO}$.
• Applications: Metallurgy (Extraction of Gold/Silver via cyanide complexation), Biological systems (Chlorophyll with $\text{Mg}^{2+}$, Hemoglobin with $\text{Fe}^{2+}$), Medicine (Cisplatin $[\text{Pt(NH}_3)_2\text{Cl}_2]$ as anti-cancer drug).
Intext & Conceptual Questions
1. $[\text{Co(NH}_3)_6]\text{Cl}_3$: Hexaamminecobalt(III) chloride
2. $[\text{Pt(NH}_3)_2\text{Cl(NO}_2)]$: Diamminechloridonitrito-N-platinum(II)
3. $\text{K}_3[\text{Fe(CN)}_6]$: Potassium hexacyanoferrate(III)
| Feature | Inner Orbital Complex | Outer Orbital Complex |
|---|---|---|
| d-Orbital Used | Uses inner $(n-1)d$ orbitals for hybridization. | Uses outer $nd$ orbitals for hybridization. |
| Hybridization Type | $d^2sp^3$ hybridization. | $sp^3d^2$ hybridization. |
| Ligand Nature | Formed in the presence of strong field ligands ($\text{CN}^-, \text{CO}$). | Formed in the presence of weak field ligands ($\text{F}^-, \text{H}_2\text{O}$). |
| Spin & Magnetism | Low-spin complex; usually diamagnetic or less paramagnetic. | High-spin complex; strongly paramagnetic. |
Textbook Chapter End Exercise Questions
Given complex: $[\text{Fe(H}_2\text{O})_6]\text{Cl}_3$
• Oxidation Number of Fe: Let oxidation state of $\text{Fe}$ be $x$. $\text{H}_2\text{O}$ is neutral ligand (charge = 0), $\text{Cl}^-$ has charge $-1$.
$x + 6(0) + 3(-1) = 0 \implies x = +3$. Thus, oxidation state of $\text{Fe}$ is +3.
• Coordination Number: Central $\text{Fe}^{3+}$ is bonded to 6 monodentate $\text{H}_2\text{O}$ molecules. Hence, Coordination Number = 6.
• d-Electron Configuration: Atomic number of $\text{Fe} = 26$. Configuration of $\text{Fe} = [\text{Ar}] 3d^6 4s^2$.
For $\text{Fe}^{3+}$ ion: $[\text{Ar}] 3d^5 4s^0$. Since $\text{H}_2\text{O}$ is a weak field ligand ($\Delta_o < P$), electrons do not pair up.
CFT configuration in octahedral field: $t_{2g}^3 e_g^2$. Total $d$-electrons = $3d^5$ ($t_{2g}^3 e_g^2$).
In both complex ions, Iron is present in the $+3$ oxidation state ($\text{Fe}^{3+}$ with $3d^5$ configuration).
• In $[\text{Fe(H}_2\text{O})_6]^{3+}$:
Water ($\text{H}_2\text{O}$) is a weak field ligand. The crystal field splitting energy ($\Delta_o$) is less than the electron pairing energy ($\Delta_o < P$). Thus, electrons enter $e_g$ orbitals rather than pairing up in $t_{2g}$.
- Electronic Distribution: $t_{2g}^3 e_g^2$
- Number of unpaired electrons ($n$) = 5
- Spin-only magnetic moment: $\mu = \sqrt{5(5+2)} = \sqrt{35} \approx \mathbf{5.92\text{ BM}}$ (Strongly Paramagnetic).
• In $[\text{Fe(CN)}_6]^{3-}$:
Cyanide ($\text{CN}^-$) is a strong field ligand. The crystal field splitting energy is greater than pairing energy ($\Delta_o > P$). Thus, all electrons pair up in the lower energy $t_{2g}$ set.
- Electronic Distribution: $t_{2g}^5 e_g^0$
- Number of unpaired electrons ($n$) = 1
- Spin-only magnetic moment: $\mu = \sqrt{1(1+2)} = \sqrt{3} \approx \mathbf{1.73\text{ BM}}$ (Weakly Paramagnetic).
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