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Chemistry Master Class Notes: Chemical Bonding — High-Yield Exam Guide

Chemical Bonding & Molecular Geometry: Complete Master Notes

Part 1: Why Atoms Form Chemical Bonds

Atoms bond to lower their potential energy and attain a stable octet (8 valence electrons) or duplet (2 valence electrons for H and He) electronic configuration, mimicking noble gases.

Part 2: Primary Types of Chemical Bonds

1. Ionic (Electrovalent) Bonding

Formed by complete transfer of one or more valence electrons from an electropositive metal atom to an electronegative non-metal atom, resulting in strong electrostatic attraction between cations and anions.

\[ \text{Na} (2,8,1) + \text{Cl} (2,8,7) \rightarrow \text{Na}^+ (2,8) + \text{Cl}^- (2,8,8) \implies \text{NaCl} \]

• Properties of Ionic Compounds: High melting and boiling points, high lattice energy, soluble in polar solvents (water), non-conductors in solid state but excellent conductors in molten or aqueous state.

2. Covalent Bonding

Formed by mutual sharing of electron pairs between non-metal atoms of comparable electronegativity.

• Single Bond (1 shared pair): \( \text{H}-\text{H} \), \( \text{Cl}-\text{Cl} \), \( \text{CH}_4 \)
• Double Bond (2 shared pairs): \( \text{O}=\text{O} \), \( \text{CO}_2 \)
• Triple Bond (3 shared pairs): \( \text{N}\equiv\text{N} \), \( \text{HC}\equiv\text{CH} \)

3. Coordinate (Dative) Covalent Bonding

A special type of covalent bond where the shared electron pair is donated entirely by one atom (Donor / Lewis Base) to an electron-deficient atom (Acceptor / Lewis Acid).

\[ \text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+ \quad | \quad \text{H}_2\text{O} + \text{H}^+ \rightarrow \text{H}_3\text{O}^+ \]

4. Metallic Bonding

Described by the Electron Sea Model (Lorentz & Drude): Positive metal ions (kernels) are immersed in a sea of mobile, delocalized valence electrons. Explains high thermal & electrical conductivity, metallic lustre, malleability, and ductility.

5. Hydrogen Bonding

An electrostatic attractive force between a hydrogen atom covalently bonded to a highly electronegative atom (\( \text{F}, \text{O}, \text{N} \)) and another electronegative atom with a lone pair.

• Intermolecular H-Bonding: Occurs between separate molecules (e.g., \( \text{H}_2\text{O} \), \( \text{HF} \), \( \text{NH}_3 \)). Causes elevated boiling points and anomalous expansion of water (maximum density at \( 4^\circ\text{C} \)).
• Intramolecular H-Bonding: Occurs within the same molecule (e.g., o-nitrophenol). Lowers boiling point compared to p-nitrophenol.

Part 3: VSEPR Theory & Molecular Geometries

Valence Shell Electron Pair Repulsion (VSEPR) Theory (Gillespie & Nyholm) predicts 3D molecular shapes based on minimizing electron pair repulsions:

Repulsion Order: \( \text{Lone Pair - Lone Pair (lp-lp)} > \text{Lone Pair - Bond Pair (lp-bp)} > \text{Bond Pair - Bond Pair (bp-bp)} \)

Part 4: Molecular Orbital Theory (MOT) & Magnetism

Formulated by Robert Mulliken and Friedrich Hund (1932), Molecular Orbital Theory treats electrons as delocalized over the entire molecule in Molecular Orbitals formed by Linear Combination of Atomic Orbitals (LCAO).

1. Bonding vs Anti-Bonding Molecular Orbitals

• Bonding MO (\( \sigma, \pi \)): Formed by constructive in-phase overlap of atomic orbital wavefunctions (\( \psi_A + \psi_B \)). Lower energy, higher electron density between nuclei, stabilizes the molecule.
• Anti-Bonding MO (\( \sigma^*, \pi^* \)): Formed by destructive out-of-phase overlap (\( \psi_A - \psi_B \)). Higher energy, nodal plane between nuclei, destabilizes the molecule.

2. Bond Order & Magnetic Character Formulas

\[ \text{Bond Order (B.O.)} = \frac{N_b - N_a}{2} \]

Where \( N_b \) is the number of electrons in bonding orbitals, and \( N_a \) is the number of electrons in anti-bonding orbitals.

• If \( \text{B.O.} > 0 \), the molecule is stable; if \( \text{B.O.} = 0 \), the molecule cannot exist (e.g., \( \text{He}_2 \)).
• Higher Bond Order indicates shorter bond length and higher Bond Dissociation Energy.
• Paramagnetic: Contains unpaired electrons in MOs (attracted by magnetic fields, e.g., \( \text{O}_2, \text{B}_2, \text{NO} \)).
• Diamagnetic: All electrons are paired in MOs (repelled by magnetic fields, e.g., \( \text{N}_2, \text{F}_2, \text{C}_2 \)).

Part 5: Fajan's Rules for Covalent Character in Ionic Bonds

No chemical bond is 100% ionic or 100% covalent. Kasimir Fajans established rules predicting covalent character in ionic compounds based on polarization:

1. Small Cation Size: Smaller cations have higher charge density and greater polarizing power (e.g., \( \text{LiCl} \) is more covalent than \( \text{NaCl} \)).
2. Large Anion Size: Larger anions are more easily distorted (higher polarizability) by cations (e.g., \( \text{AgI} \) is more covalent than \( \text{AgF} \)).
3. High Ionic Charge: Higher positive charge on cation or negative charge on anion increases polarization (e.g., \( \text{AlCl}_3 \) is more covalent than \( \text{MgCl}_2 \) and \( \text{NaCl} \)).
4. Pseudo-Noble Gas Configuration: Cations with \( 18 \)-electron outer shell (\( d^{10} \) configuration, e.g., \( \text{Cu}^+, \text{Ag}^+ \)) have greater polarizing power than cations with 8-electron noble gas configuration (e.g., \( \text{Na}^+, \text{K}^+ \)).

Part 6: High-Yield Bonding & Molecular Geometry Question Bank

Part 7: Dipole Moments, Resonance & Formal Charge Masterclass

1. Molecular Dipole Moment (\( \vec{\mu} \))

Dipole moment measures the overall electrical polarity of a molecule as a vector quantity pointing from electropositive to electronegative atom:

\[ \vec{\mu} = q \times d \quad (\text{SI Unit: C}\cdot\text{m}, \text{ Common Unit: Debye } 1 \text{ D} = 3.33564 \times 10^{-30} \text{ C}\cdot\text{m}) \]

• Non-Polar Symmetrical Molecules (\( \mu = 0 \)): Carbon Dioxide (\( \text{CO}_2 \), linear, opposing dipoles cancel), Boron Trifluoride (\( \text{BF}_3 \), trigonal planar), Carbon Tetrachloride (\( \text{CCl}_4 \), tetrahedral).
• Polar Asymmetrical Molecules (\( \mu > 0 \)): Water (\( \text{H}_2\text{O}, \mu = 1.85 \text{ D} \)), Ammonia (\( \text{NH}_3, \mu = 1.47 \text{ D} \)), Chloroform (\( \text{CHCl}_3, \mu = 1.04 \text{ D} \)).

2. Resonance & Formal Charge Calculation

When a single Lewis structure cannot explain all experimental properties of a molecule (e.g., equal C-O bond lengths in Carbonate ion \( \text{CO}_3^{2-} \)), the actual molecule is represented as a Resonance Hybrid of multiple canonical structures.

Formula for Formal Charge (F.C.) on an atom in a Lewis structure:

\[ \text{F.C.} = V - L - \frac{1}{2} B \]

Where \( V \) is the number of valence electrons in free atom, \( L \) is the number of non-bonding lone pair electrons, and \( B \) is the number of bonding shared electrons.

Part 8: Comprehensive Chemical Bonding Master Comparison & Exam Review

High-Yield Practice Questions & Concept Review

Part 9: Advanced VSEPR Geometry Masterclass & Hybridization Calculation

Method for Determining Hybridization State & Geometry

To determine the hybridization of any central atom in a molecule or polyatomic ion, calculate the Steric Number (\( Z \)):

\[ Z = \frac{1}{2} [ V + M - C + A ] \]

• \( V \) = Number of valence electrons on central atom
• \( M \) = Number of monovalent atoms bonded to central atom (H, F, Cl, Br, I)
• \( C \) = Positive charge on cation
• \( A \) = Negative charge on anion

Frequently Asked Questions (FAQ) & High-Yield Exam Tips