Ionic Bonding and Fajan’s Rules
An ionic bond is formed by the complete transfer of one or more valence electrons from an electropositive atom (metal) to an electronegative atom (non-metal). This results in electrostatic attraction between oppositely charged ions.
This process is favored by low ionization energy of the metal, high electron affinity of the non-metal, and a high negative lattice enthalpy. Compounds like $\text{NaCl}$ and $\text{MgO}$ are classic examples of ionic lattices.
However, no ionic bond is 100% pure; every ionic bond possesses a certain degree of covalent character due to polarization—the distortion of an anion’s electron cloud by a neighboring cation. This phenomenon is quantitatively analyzed using Fajan’s Rules, which predict covalent character based on specific ionic parameters. Aspirants must thoroughly memorize these rules as they frequently appear in conceptual JEE and NEET questions regarding melting points, solubility, and thermal stability.
- Cation Size: Smaller cations have higher charge density, leading to greater polarizing power and increased covalent character.
- Anion Size: Larger anions are more polarizable because their outermost electrons are loosely bound to the nucleus.
- Charge on Ions: Higher charges on either the cation or the anion increase electrostatic attraction and polarization, enhancing covalent character.
- Electronic Configuration: Cations with a pseudo-inert gas configuration ($n\text{s}^2n\text{p}^6n\text{d}^{10}$) possess higher polarizing power than those with an inert gas configuration ($n\text{s}^2n\text{p}^6$) due to poor shielding by d-electrons.
“Polarization leads to partial sharing of electrons, bridging the gap between extreme ionic bonding and true covalent bonding.”
Covalent Bonding, VSEPR Theory, and Hybridisation
A covalent bond involves the mutual sharing of one or more electron pairs between two atoms. To predict the precise three-dimensional geometry of these covalent molecules, the Valence Shell Electron Pair Repulsion (VSEPR) Theory is employed.
According to VSEPR theory, electron pairs around a central atom repel one another and arrange themselves spatially to minimize this repulsion. Lone pair-lone pair repulsion remains greater than lone pair-bond pair repulsion, which in turn exceeds bond pair-bond pair repulsion.
Building upon molecular geometry, the concept of hybridisation explains the mixing of atomic orbitals of comparable energies to form new equivalent orbitals known as hybrid orbitals. Common hybridisation states include $\text{sp}$ (linear, $180^\circ$), $\text{sp}^2$ (trigonal planar, $120^\circ$), $\text{sp}^3$ (tetrahedral, $109.5^\circ$), $\text{sp}^3\text{d}$ (trigonal bipyramidal), and $\text{sp}^3\text{d}^2$ (octahedral). Understanding lone pair contributions to these geometries is critical for solving structural problems in competitive examinations.
Resonance structures and formal charges are equally vital tools for evaluating stability. When a single Lewis structure cannot adequately account for all molecular properties, a hybrid of multiple resonance structures is used.
The formal charge ($\text{FC}$) of an atom in a molecule is calculated using the formula: $\text{FC} = V – N – \frac{B}{2}$, where $V$ represents valence electrons, $N$ denotes non-bonding valence electrons, and $B$ indicates total bonding electrons.
Molecular Orbital Theory (MOT) and Bond Order
While Valence Bond Theory and Lewis structures provide qualitative descriptions, Molecular Orbital Theory (MOT), developed by Hund and Mulliken, treats electrons in a molecule as belonging to molecular orbitals spanning the entire nuclear framework. Atomic orbitals linearly combine through Linear Combination of Atomic Orbitals (LCAO) to form Bonding Molecular Orbitals (BMOs) and Antibonding Molecular Orbitals (ABMOs).
The stability and existence of a molecule are determined by its bond order, which is calculated as half the difference between the number of electrons in bonding and antibonding orbitals: $\text{Bond Order} = \frac{N_b – N_a}{2}$. A positive bond order indicates a stable species, while a bond order of zero indicates non-existence.
For diatomic molecules like $\text{O}_2$ and $\text{N}_2$, MOT successfully explains magnetic properties. It confirms that $\text{O}_2$ is paramagnetic due to unpaired electrons in its $\pi^*$ antibonding orbitals.
- Bond Order Relationship: Higher bond order corresponds to higher bond dissociation energy and shorter bond length.
- Sp-Mixing Effect: For diatomic molecules up to $\text{N}_2$ ($\le 14$ electrons), s-p mixing alters the filling order of molecular orbitals, placing $\pi 2\text{p}_x = \pi 2\text{p}_y$ below $\sigma 2\text{p}_z$.
- Paramagnetism vs Diamagnetism: Presence of unpaired electrons confers paramagnetism; all paired electrons denote diamagnetism.
Coordinate Bonds, Dipole Moment, and Hydrogen Bonding
A coordinate bond (or dative bond) is a special type of covalent bond where both shared electrons originate from a single atom (the donor), while the other atom acts as the acceptor. Represented by an arrow pointing from donor to acceptor, these bonds are fundamentally identical to covalent bonds once formed.
Coordinate bonds are observed extensively in coordination compounds and complex ions like $\text{NH}_4^+$ and $[\text{PtCl}_4]^{2-}$.
Polarity in molecules is measured via the dipole moment ($\mu$), defined as the product of magnitude of charge ($q$) and the distance between centers ($d$). Measured in Debyes ($\text{D}$), dipole moment is a vector quantity useful in predicting molecular symmetry and differentiating cis-trans isomers. Symmetrical molecules such as $\text{CO}_2$ and $\text{CCl}_4$ exhibit a net dipole moment of zero despite having polar bonds.
Intermolecular forces also significantly alter physical properties. Hydrogen bonding is a strong electrostatic dipole-dipole attraction occurring when hydrogen is covalently bonded to highly electronegative atoms such as Fluorine, Oxygen, or Nitrogen ($\text{F}, \text{O}, \text{N}$).
Intermolecular hydrogen bonding increases boiling points and solubility in water (e.g., $\text{H}_2\text{O}$). Conversely, intramolecular hydrogen bonding (seen in ortho-nitrophenol) suppresses these effects by forming stable ring structures.
Important Facts / Formulas
| Property / Concept | Mathematical Expression / Rule | Significance |
|---|---|---|
| Formal Charge | $\text{FC} = V – N – \frac{B}{2}$ | Determines most stable resonance structure |
| Bond Order (MOT) | $\text{B.O.} = \frac{N_b – N_a}{2}$ | Measures bond strength and stability |
| Dipole Moment | $\mu = q \times d$ | Determines molecular polarity and geometry |
| Fajan’s Rule (Covalent) | $\text{Small cation} + \text{Large anion} + \text{High charge}$ | Predicts covalent character in ionic compounds |
Key Points to Remember
- Fajan’s rules dictate that covalent character increases with smaller cation size and larger anion size.
- VSEPR geometry accounts for lone pairs; lone pair-lone pair repulsions are maximum.
- Hybridisation states ($\text{sp}$, $\text{sp}^2$, $\text{sp}^3$) determine standard bond angles in polyatomic species.
- Molecules with odd numbers of electrons or unpaired electrons in MOT are paramagnetic.
- Intramolecular hydrogen bonding decreases boiling point, whereas intermolecular hydrogen bonding increases it.
- Coordinate bonds act identically to covalent bonds after formation but involve unidirectional electron donation.
- Bond order is directly proportional to bond dissociation enthalpy and inversely proportional to bond length.
- Symmetrical molecules with polar bonds often have a net dipole moment of zero due to vector cancellation.
Previous Year Question Hints
- Hint 1: When asked to compare thermal stability or melting points of alkaline earth metal carbonates, apply Fajan’s rules to identify which compound possesses higher covalent character.
- Hint 2: For questions involving the magnetic behavior of diatomic species (e.g., $\text{O}_2^-, \text{O}_2^{2-}$), construct the Molecular Orbital electron configuration to count unpaired electrons accurately.
- Hint 3: Identify hybridisation and geometry changes in reactions involving Lewis acid-base adduct formation, such as $\text{BF}_3 + \text{F}^- \rightarrow \text{BF}_4^-$.
Quick Revision Summary
- Ionic bonds involve complete electron transfer; covalent bonds involve electron sharing.
- Fajan’s rules correlate high polarization with increased covalent character.
- VSEPR theory minimizes spatial repulsion between valence electron pairs.
- Hybridisation explains directional bonding and specific molecular geometries.
- Molecular Orbital Theory combines atomic orbitals to evaluate bond order and paramagnetism.
- Coordinate bonds feature a single atom donating both electrons in the shared pair.
- Dipole moment measures molecular polarity and verifies molecular symmetry.
- Hydrogen bonding involving $\text{F}, \text{O}$, and $\text{N}$ drastically impacts boiling points and solubility.
