Chemical Bonding and Molecular Structure – Comprehensive Notes

This chapter is fundamental to understanding the properties and reactivity of all chemical compounds. It explains why atoms combine and the resulting shapes of molecules.

1. Introduction to Chemical Bonding

• Why do atoms combine? Atoms combine to achieve stability, primarily by acquiring a stable electron configuration (usually a duplet or octet) similar to that of noble gases. This process leads to a lower energy state for the combined system.
• Octet Rule: Atoms tend to gain, lose, or share electrons so as to achieve eight electrons in their outermost shell.
  • Limitations of Octet Rule:
    • Incomplete Octet of the Central Atom: e.g., LiCl, BeH2, BCl3
    • Odd Electron Molecules: e.g., NO, NO2
    • Expanded Octet: e.g., PCl5, SF6, H2SO4 (central atom has more than 8 valence electrons).
    • Noble gas compounds like XeF2, XeF4, XeOF4.
• Types of Bonds:
  • Ionic Bond (Electrovalent Bond): Formed by the complete transfer of one or more electrons from one atom (typically a metal) to another (typically a non-metal).
  • Covalent Bond: Formed by the mutual sharing of electrons between two atoms (typically non-metals).
  • Coordinate (Dative) Bond: A type of covalent bond where both shared electrons are contributed by only one of the combining atoms (donor atom). The acceptor atom has an empty orbital.
    • Represented by an arrow (→). Example: NH3 -> BF3 (in NH3.BF3).

2. Ionic Bonding

• Formation: Occurs between atoms with large electronegativity difference. Metal loses electrons to form cations, non-metal gains electrons to form anions. These oppositely charged ions are held together by strong electrostatic forces of attraction.
  • Example: Formation of NaCl
    • Na -> Na+ + e- (Ionization Enthalpy)
    • Cl + e- -> Cl- (Electron Gain Enthalpy)
    • Na+ (g) + Cl- (g) -> NaCl (s) (Lattice Energy)
• Lattice Energy: The energy released when one mole of an ionic compound is formed from its constituent gaseous ions. It is a measure of the strength of the ionic bond.
  • Factors affecting Lattice Energy:
    • Charge on Ions: Higher charge, greater lattice energy (MgO > NaCl).
    • Size of Ions: Smaller ions, greater lattice energy.
• Characteristics of Ionic Compounds:
  • Typically solids with high melting and boiling points.
  • Hard and brittle.
  • Soluble in polar solvents (like water).
  • Good conductors of electricity in molten state or in aqueous solution (due to free ions), but non-conductors in solid state.

3. Covalent Bonding

3.1. Lewis Theory of Covalent Bonding (Lewis Dot Structures)

• Atoms share electrons to complete their octet (or duplet for H).
• Shared pair(s) constitute covalent bond(s).
• Formal Charge: A hypothetical charge assigned to an atom in a molecule assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.Formal Charge (FC)=Valence Electrons (V)−Non-bonding Electrons (N)−21​×Bonding Electrons (B)
  • Helps in selecting the most stable Lewis structure (structure with lowest formal charges, or negative formal charge on more electronegative atom).

3.2. VSEPR Theory (Valence Shell Electron Pair Repulsion Theory)

• Principle: Electron pairs (both bond pairs and lone pairs) in the valence shell of the central atom repel each other and arrange themselves in space to minimize repulsion, thus determining the molecular geometry.
• Repulsion Order: Lone Pair-Lone Pair (LP-LP) > Lone Pair-Bond Pair (LP-BP) > Bond Pair-Bond Pair (BP-BP).
• Steps to Predict Geometry:
  1. Count total valence electrons.
  2. Draw Lewis structure.
  3. Count steric number (SN = No. of Bond Pairs + No. of Lone Pairs around central atom).
  4. Predict electron geometry and molecular geometry based on SN and LP.
• Examples:

• Effect of Lone Pairs on Bond Angles: Lone pairs occupy more space than bond pairs, causing compression of bond angles. (e.g., CH4 (109.5°), NH3 (107°), H2O (104.5°)).

3.3. Valence Bond Theory (VBT)

• Principle: Covalent bond is formed by the overlapping of atomic orbitals of two atoms. The greater the overlap, the stronger the bond.
• Types of Orbital Overlap:
  • Sigma (σ) Bond: Formed by head-on (axial) overlap of orbitals. Stronger, free rotation.
    • Examples: s-s, s-p, p-p (head-on overlap).
  • Pi (π) Bond: Formed by sideways (lateral) overlap of unhybridized p-orbitals. Weaker, no free rotation (leads to geometric isomerism).
    • Always present along with a σ bond in multiple bonds (e.g., double bond: one σ, one π; triple bond: one σ, two π).
• Hybridization: The concept of mixing atomic orbitals of slightly different energies to form new, equivalent hybrid orbitals of equal energy and shape, suitable for effective overlap and bond formation.
  • Need for Hybridization: To explain the observed geometries and bond angles which cannot be explained by simple atomic orbital overlap (e.g., CH4 should have 90° angles with pure p-orbitals, but has 109.5°).
  • Characteristics of Hybrid Orbitals:
    • Number of hybrid orbitals equals number of atomic orbitals mixed.
    • Hybrid orbitals are identical in energy and shape.
    • More effective in forming stable bonds than pure atomic orbitals.
    • Oriented in specific directions to minimize repulsion.
  • Types of Hybridization & Geometry:

• Hybridization in Molecules with Multiple Bonds:
  • In C2H4 (Ethene): Each carbon is sp2 hybridized, forming three σ bonds and one π bond (from unhybridized p-orbitals). (Diagram: Ethene structure showing σ bonds from sp2 overlap and π bond from parallel p-orbital overlap)
  • In C2H2 (Ethyne): Each carbon is sp hybridized, forming two σ bonds and two π bonds. (Diagram: Ethyne structure showing σ bonds from sp overlap and two π bonds from two sets of parallel p-orbital overlap)
• Bond Parameters:
  • Bond Length: Equilibrium distance between the nuclei of two bonded atoms.
    • Factors: Decreases with multiple bonds (C#C < C=C < C-C), decreases with increasing s-character (sp-sp < sp2-sp2 < sp3-sp3), decreases with smaller atomic size.
  • Bond Angle: Angle between the orbitals containing bonding electron pairs around the central atom.
    • Influenced by hybridization and lone pairs (VSEPR).
  • Bond Enthalpy (Bond Energy): Energy required to break one mole of a particular type of bond in gaseous state.
    • Higher bond order means higher bond enthalpy.
    • Factors: Bond multiplicity, bond length.
  • Bond Order: Number of bonds between two atoms (e.g., single bond = 1, double bond = 2, triple bond = 3). For resonance structures, it’s average.

3.4. Molecular Orbital Theory (MOT)

• Principle: Atomic orbitals (AOs) of combining atoms combine to form new orbitals called Molecular Orbitals (MOs). Electrons occupy MOs, similar to how they occupy AOs.
• LCAO (Linear Combination of Atomic Orbitals): Atomic orbitals combine either constructively (addition) to form Bonding Molecular Orbitals (BMOs) (lower energy, stabilize) or destructively (subtraction) to form Anti-bonding Molecular Orbitals (ABMOs) (higher energy, destabilize).
  • Represented as σ,π for BMOs and σ∗,π∗ for ABMOs.
• Rules for Filling MOs:
  • Aufbau Principle: Fill lower energy MOs first.
  • Pauli’s Exclusion Principle: Max two electrons per MO with opposite spins.
  • Hund’s Rule: Degenerate MOs are singly occupied before pairing.
• Energy Level Diagrams for Homonuclear Diatomic Molecules:
  • For molecules up to N2 (total electrons ≤ 14): σ1s<σ∗1s<σ2s<σ∗2s<π2px​=π2py​<σ2pz​<π∗2px​=π∗2py​<σ∗2pz​
  • For molecules O2 and F2 (total electrons > 14): σ1s<σ∗1s<σ2s<σ∗2s<σ2pz​<π2px​=π2py​<π∗2px​=π∗2py​<σ∗2pz​
• Bond Order (BO):Bond Order=21​(Number of electrons in BMOs (Nb)−Number of electrons in ABMOs (Na))
  • BO > 0 ⟹ molecule is stable.
  • Higher BO ⟹ greater stability, shorter bond length, higher bond enthalpy.
• Magnetic Properties:
  • Paramagnetic: Contains unpaired electrons in MOs (attracted by magnetic field).
  • Diamagnetic: All electrons are paired in MOs (repelled by magnetic field).
• Examples:
  • H2 (2 electrons): σ1s2. BO = 21​(2−0)=1. Diamagnetic.
  • He2 (4 electrons): σ1s2σ∗1s2. BO = 21​(2−2)=0. Does not exist.
  • O2 (16 electrons): σ1s2σ∗1s2σ2s2σ∗2s2σ2pz2​π2px2​=π2py2​π∗2px1​=π∗2py1​.
    • BO = 21​(10−6)=2. Paramagnetic (due to two unpaired electrons in π∗ MOs).

4. Hydrogen Bonding

• Definition: An attractive force between a hydrogen atom covalently bonded to a highly electronegative atom (F, O, or N) and another highly electronegative atom in the same or different molecule.
  • Represented by a dotted line (…).
• Conditions:
  1. Presence of a highly electronegative atom (F, O, N) directly bonded to hydrogen.
  2. Presence of another highly electronegative atom (F, O, N) with a lone pair to accept the hydrogen.
• Types:
  • Intermolecular H-bonding: Occurs between different molecules of the same or different compounds.
    • Increases association, leading to higher melting point, boiling point, and solubility in water.
    • Example: Water (H2O), Alcohols, Ammonia.(Diagram: Intermolecular H-bonding in water)
  • Intramolecular H-bonding: Occurs within the same molecule.
    • Leads to cyclization, decreasing intermolecular forces, volatility increases, solubility decreases.
    • Example: o-Nitrophenol(Diagram: Intramolecular H-bonding in o-Nitrophenol)

5. Dipole Moment (μ)

• Definition: A measure of the polarity of a covalent bond or a molecule. It is a vector quantity.μ=q×dwhere q = magnitude of charge, d = distance between charges. Units: Debye (D). 1D=3.33564×10−30Cm.
• Polar vs. Non-polar Bonds:
  • Polar Bond: Formed between atoms with different electronegativities (e.g., H-Cl).
  • Non-polar Bond: Formed between atoms with same electronegativity (e.g., Cl-Cl).
• Net Dipole Moment in Polyatomic Molecules: Vector sum of individual bond dipole moments.
  • Non-polar Molecules (Net μ=0): Symmetrical structures where individual bond dipoles cancel out.
    • Example: CO2 (Linear), CCl4 (Tetrahedral), BF3 (Trigonal Planar).(Diagram: Linear CO2 with canceling dipole arrows)
  • Polar Molecules (Net μ=0): Asymmetrical structures or where bond dipoles do not cancel out.
    • Example: H2O (Bent), NH3 (Pyramidal), CHCl3 (Tetrahedral, asymmetric).(Diagram: Bent H2O with net dipole arrow)
• Applications:
  • Distinguishing between cis and trans isomers (cis-isomers often have non-zero dipole moment while trans-isomers often have zero, e.g., cis-1,2-dichloroethene vs trans-1,2-dichloroethene).
  • Predicting polarity of molecules.
  • Determining molecular geometry (e.g., CO2 vs H2O).