CHEMICAL BONDING AND MOLECULAR STRUCTURE

Chemical Bonding and Molecular Structure

Fundamentals of Chemical Bonding

• Atoms combine to complete their octets (8 electrons in outer shell).

  • Methods of bonding:

    • Ionic bond: Transfer of electrons.

    • Covalent bond: Mutual sharing of electrons.

    • Co-ordinate bond: One side sharing of electrons.

• G.N Lewis introduced symbols (dots or crosses) to represent valence electrons.

Covalent Bonding

• Formation of molecules through mutual sharing of electrons results in bonds.

• Covalency: The number of electrons contributed by each atom to the bond.

• Illustrations:

  • Cl2: Formation through shared electron pairs.

  • O2, N2, H2O, NH3: Examples of molecular formation involving different elements.

• Hydrogen: Requires two electrons to complete its duplet (unlike the octet).

Octet Rule & Limitations

• Octet Rule – Central to understanding stability in bonding.

• Lewis Structures – Visual representations of covalent bonds.

• Limitations of the Octet Rule: Does not apply to all compounds, especially electron-deficient or expanded octet compounds.

Formation of Electron Deficient Compounds

• Some compounds, e.g., BeCl2, BF3, AlCl3, have less than 8 electrons in the central atom.

  • Example: Be has 4, B has 6 electrons in bonding scenarios.

Expanded Octet

• Central atoms like P in PCl5 or S in SF6 can have more than 8 electrons (10 for P, 12 for S).

Odd-Electron Molecules

• Molecules like NO, NO2, O2 have an odd number of electrons.

VSEPR Theory

• Proposed by Sidgwick and Powell, further refined by Nyholm and Gillespie.

• Key postulates:

  • Molecular shape depends on the number of Valence Shell Electron Pairs (VSEPs) around the central atom.

  • Repulsion between valence shell electron pairs:

    • Lone pair - lone pair > Lone pair - bond pair > Bond pair - bond pair.

  • Electron pairs adopt geometries to minimize repulsion, defining molecular shapes.

Valence Bond Theory

• Introduced by Heitler, London, and Linus Pauling to explain molecular shapes based on energy considerations.

Covalent Bond Formation

• The formation of an H2 molecule through attractive and repulsive forces between approaching hydrogen atoms.

• Bond Length: Distance at which net attractive and repulsive forces balance (74 pm for H2).

• Bond Enthalpy: Energy release or required to form a bond (435.8 kJ/mol for H2).

Orbital Overlap Concept

• A covalent bond is created when half-filled atomic orbitals overlap, requiring opposite spins.

• Types of Overlapping:

  • Sigma (σ) bond: Formed by end-to-end overlapping.

  • Pi (π) bond: Formed by lateral overlapping.

Bond Characteristics

• Bond Length: Equilibrium distance between nuclei.

• Influencing factors:

  • Size of atoms: Larger atoms result in longer bond lengths (e.g., HI > HBr).

  • Multiplicity of bond: Higher bond order leads to shorter bond lengths (C≡C < C=C < C-C).

• Bond Enthalpy: Energy to break bonds; average bond enthalpies used for polyatomic molecules.

Bond Order

• Definition: Number of bonds between two atoms.

• Related to stability: Higher bond order means stronger bonds and shorter lengths.

Dipole Moments

• Measures the separation of charge in polar molecules.

• Represented by an arrow, with size based on electronegativity differences.

Hybridization

• Hybridization: Mixing of atomic orbitals to form hybrid orbitals for bonding.

  • Characteristics:

    • Number of hybrid orbitals = number of orbitals mixed.

    • Hybrid orbitals shape defines molecular geometry.

• Types:

  • sp3, sp2, sp, sp3d, sp3d2 hybridization with respective molecular shapes and bond angles.

Molecular Orbital Theory

• Explains electron configurations in molecules through Molecular Orbitals (M.Os).

  • Formation through Linear Combination of Atomic Orbitals (LCAO).

  • Bonding MOs (lower energy) vs. Anti-bonding MOs (higher energy).

Conditions for Atomic Orbital Combination

• Same energy and symmetry required for effective overlap.

Magnetic Properties of Molecules

• Diamagnetic: All MOs doubly occupied; repelled by magnetic fields.

• Paramagnetic: Presence of unpaired electrons; attracted to magnetic fields.

Ionic Bonding

• Formation dependent on:

  • Ionization Energy: Lower energy favors cation formation.

  • Electron Gain Enthalpy: More negative values ease anion formation.

• Lattice Energy: Energy released when ionic compounds form from gaseous ions, influenced by:

  • Ionic size and charge.

Fajans' Rules of Polarization

• Rules explaining covalent character in ionic bonds.

  • Smaller cations/larger anions cause greater covalent character.

Coordinate Bonds

• One-sided sharing of electrons, illustrated with examples like ammonium ion formation.

Formal Charge and Resonance Structures

• Formal Charge: Difference between valence electrons in the elemental state and assigned electrons in the Lewis structure.

• Resonance: Combining multiple Lewis structures to represent delocalization in a molecule, such as in carbonate ion (CO3 2-).

Hydrogen Bonding

• Attractions involving H with highly electronegative atoms like F, O, N.

  • Types: Intramolecular and intermolecular hydrogen bonds.