Chemical Bonding and Molecular Structure

1. Introduction to Chemical Bonding

Scientists are constantly discovering new compounds, arranging facts, explaining with existing knowledge, and modifying earlier views or evolving theories for new observations.

Matter is composed of one or more types of elements. Under normal conditions, no element exists as an independent atom in nature, except for noble gases.

A molecule is a group of atoms existing together as one species with characteristic properties.

A chemical bond is the attractive force that holds various constituents (atoms, ions, etc.) together in different chemical species.

The formation of chemical compounds involves atoms combining in various ways, raising questions like: Why do atoms combine? Why are only certain combinations possible? Why do molecules possess definite shapes?

To answer these questions, different theories and concepts have been proposed, including:

• Kössel-Lewis approach.
• Valence Shell Electron Pair Repulsion (VSEPR) Theory.
• Valence Bond (VB) Theory.
• Molecular Orbital (MO) Theory.

The understanding of chemical bonds is closely linked to developments in atomic structure, electronic configuration, and the periodic table.

Bonding is nature’s way of lowering the energy of the system to attain stability.

2. Kössel-Lewis Approach to Chemical Bonding

Developed independently by Kössel and Lewis in 1916, this approach provides a logical explanation of valence based on the inertness of noble gases.

Lewis pictured the atom with a positively charged 'Kernel' (nucleus plus inner electrons) and an outer shell accommodating up to eight electrons. He assumed these eight electrons occupy the corners of a cube surrounding the Kernel. A filled cube (octet of electrons) represents a stable electronic arrangement.

Lewis postulated that atoms achieve a stable octet by linking through chemical bonds, either by electron transfer or by sharing electron pairs.

Lewis Symbols

Notations introduced by G.N. Lewis to represent valence electrons in an atom. The number of dots around the symbol represents the number of valence electrons. This number helps calculate the common or group valence.

Kössel's Contributions

• Noted the separation of highly electronegative halogens and highly electropositive alkali metals by noble gases.
• Observed that forming ions involves the gain and loss of electrons to achieve stable noble gas configurations (an octet).
• Identified that negative and positive ions are stabilized by electrostatic attraction.

The electrovalent bond (ionic bond) is the bond formed as a result of electrostatic attraction between positive and negative ions. Electrovalence is equal to the number of unit charge(s) on the ion.

2.1 Octet Rule

Definition: Atoms combine by transferring valence electrons or by sharing valence electrons to achieve an octet (eight electrons) in their valence shells.

2.2 Covalent Bond

Langmuir (1919) refined Lewis's postulations by abandoning the stationary cubical arrangement and introducing the term covalent bond. In a covalent bond, atoms share a pair of electrons, with each combining atom contributing at least one electron to the shared pair.

Lewis dot structures represent bonding in molecules in terms of shared and lone pairs of electrons and the octet rule.

Types of Covalent Bonds:

• Single covalent bond: Two atoms share one electron pair (e.g., Cl-Cl).
• Double bond: Two atoms share two pairs of electrons (e.g., C=O in CO₂).
• Triple bond: Two atoms share three pairs of electrons (e.g., N≡N in N₂).

2.3 Formal Charge

The difference between the number of valence electrons of an atom in its isolated state and the number of electrons assigned to that atom in a Lewis structure.

Formal charges help in selecting the lowest energy structure, which is generally the one with the smallest formal charges on atoms.

2.4 Limitations of the Octet Rule

The octet rule is useful but not universal. Exceptions include:

• Incomplete octet: Central atom has less than eight electrons (e.g., BeH₂, BCl₃).
• Odd-electron molecules: Molecules with an odd number of electrons (e.g., NO, NO₂).
• Expanded octet: Elements in and beyond the third period can have more than eight valence electrons (e.g., PF₅, SF₆).

Other drawbacks include its inability to explain the formation of compounds from noble gases, the shape of molecules, or the relative stability of molecules.

3. Ionic or Electrovalent Bond

Formation of ionic compounds depends on the ease of formation of positive and negative ions and their arrangement in a crystal lattice.

• Positive ion formation involves ionization enthalpy (endothermic).
• Negative ion formation involves electron gain enthalpy (exothermic or endothermic).

Ionic bonds form more easily between elements with low ionization enthalpies and high negative electron gain enthalpies.

The crystal structure is stabilized by the energy released in the formation of the crystal lattice, known as Lattice Enthalpy. This is the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

4. Bond Parameters

4.5 Resonance Structures

Often, a single Lewis structure is inadequate. When this happens, a number of structures with similar energy, known as canonical structures, are used. The actual structure is a resonance hybrid of these forms.

Resonance stabilizes the molecule and averages the bond characteristics. The canonical forms have no real existence; the molecule has a single, unique structure.

4.6 Polarity of Bonds

Nonpolar Covalent Bond: Formed between similar atoms, the shared electron pair is equally attracted. (e.g., H₂, O₂).

Polar Covalent Bond: Formed between dissimilar atoms, the shared electron pair is displaced towards the more electronegative atom, creating partial charges (δ+ and δ-). (e.g., HF).

The Dipole Moment (µ) is a measure of this polarization. In polyatomic molecules, the net dipole moment is the vector sum of individual bond dipoles. This determines if a molecule is polar or nonpolar overall.

5. Valence Shell Electron Pair Repulsion (VSEPR) Theory

A simple procedure to predict the shapes of covalent molecules. The main idea is that electron pairs in the valence shell repel one another and will arrange themselves to minimize this repulsion.

Order of Repulsive Interactions:
Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp).

This repulsion order explains why molecules with lone pairs have distorted shapes and altered bond angles compared to ideal geometries.

Predicted Geometrical Shapes:

• No Lone Pairs: Linear (AB₂), Trigonal Planar (AB₃), Tetrahedral (AB₄), Trigonal Bipyramidal (AB₅), Octahedral (AB₆).
• With Lone Pairs: Pyramidal (AB₃E), Bent/V-shaped (AB₂E₂), See-saw (AB₄E), T-shape (AB₃E₂).

6. Valence Bond (VB) Theory

Based on quantum mechanics, VB theory explains bond formation in terms of atomic orbitals, overlap, and hybridization.

A covalent bond forms when atomic orbitals undergo partial interpenetration (overlapping), leading to the pairing of electrons with opposite spins. Greater overlap leads to a stronger bond.

Types of Overlapping:

• Sigma (σ) Bond: Formed by end-to-end (axial) overlap. These are stronger bonds.
• Pi (π) Bond: Formed by sidewise overlap. These are weaker bonds.

7. Hybridisation

Introduced by Pauling, it's the process of intermixing atomic orbitals of slightly different energies to form a new set of equivalent orbitals called hybrid orbitals.

Hybrid orbitals are more effective in forming stable bonds and are directed in space to give molecules their characteristic shapes.

Types of Hybridisation:

• sp Hybridisation: Linear geometry (180°). Example: BeCl₂, C₂H₂.
• sp² Hybridisation: Trigonal planar geometry (120°). Example: BCl₃, C₂H₄.
• sp³ Hybridisation: Tetrahedral geometry (109.5°). Example: CH₄, NH₃, H₂O.
• Hybridisation with d Orbitals: sp³d (Trigonal bipyramidal, e.g., PCl₅), sp³d² (Octahedral, e.g., SF₆).

8. Molecular Orbital (MO) Theory

Developed by Hund and Mulliken, this theory states that electrons in a molecule occupy molecular orbitals (MOs) that are spread over the entire molecule.

Atomic orbitals combine to form an equal number of molecular orbitals: bonding MOs (lower energy, stabilizing) and antibonding MOs (higher energy, destabilizing).

MOs are filled according to the Aufbau principle, Pauli’s exclusion principle, and Hund’s rule.

Key Concepts:

• Stability: A molecule is stable if the number of bonding electrons (Nb) is greater than antibonding electrons (Na).
• Bond Order: b.o. = ½ (Nb - Na). A positive bond order indicates a stable molecule.
• Magnetic Nature: Diamagnetic if all MOs are doubly occupied; Paramagnetic if one or more MOs are singly occupied. (e.g., O₂ is paramagnetic).

9. Hydrogen Bonding

An attractive force that binds a hydrogen atom of one molecule with a highly electronegative atom (F, O, or N) of another molecule. It is weaker than a covalent bond but has a strong influence on a compound's properties (like boiling point).

Types of H-Bonds:

• Intermolecular: Between two different molecules (e.g., in water, HF).
• Intramolecular: Within the same molecule (e.g., in o-nitrophenol).