1. Basic Atomic And Molecular Concepts (Bonding Basics)
Atoms combine to form molecules to achieve greater stability, often by attaining a stable electron configuration like that of noble gases (octet rule). Chemical bonding refers to the attractive forces that hold atoms together in molecules and compounds. These forces arise from the interactions between valence electrons, the electrons in the outermost shell of an atom, which are involved in chemical reactions and bonding.
2. Introduction to Bonding
Chemical bonds are broadly classified into primary bonds (ionic, covalent, metallic) and secondary bonds (like hydrogen bonding and van der Waals forces). The type of bond formed depends on the electronegativity difference between the atoms involved. Ionic bonds typically form between metals and non-metals with a large electronegativity difference, covalent bonds between non-metals with similar electronegativity, and metallic bonds occur in metals where valence electrons are delocalized.
3. Lewis Approach To Chemical Bonding and Covalent Bonds
The Lewis approach uses dots to represent valence electrons of atoms. Covalent bonds are formed by the sharing of electron pairs between atoms to achieve a stable electron configuration. These shared pairs form a covalent bond, holding the atoms together in a molecule. Single, double, and triple covalent bonds represent the sharing of one, two, or three pairs of electrons, respectively. Lewis structures are a simple way to depict the bonding in molecules.
4. Ionic Bond and Lattice Enthalpy
An ionic bond is formed by the electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions), typically resulting from the transfer of electrons from a metal to a non-metal. The strength of this ionic bond in a crystalline solid is quantified by lattice enthalpy, which is the energy released when gaseous ions combine to form one mole of an ionic compound. Higher lattice enthalpy generally indicates stronger ionic bonding and greater stability.
5. Bond Parameters and Resonance
Bond parameters describe the characteristics of a chemical bond, including bond length (average distance between nuclei of bonded atoms) and bond enthalpy (energy required to break one mole of bonds). Resonance occurs when a molecule or ion cannot be adequately represented by a single Lewis structure; instead, its properties are described as an average of multiple contributing Lewis structures, called resonance structures. The actual structure is a resonance hybrid, more stable than any single contributing structure.
6. Valence Shell Electron Pair Repulsion (VSEPR) Theory
The Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts the molecular geometry of molecules based on the repulsion between electron pairs in the valence shell of the central atom. Electron pairs (both bonding and lone pairs) arrange themselves as far apart as possible to minimize repulsion, dictating the molecule's shape. For example, molecules with four electron groups around the central atom tend to have a tetrahedral electron geometry.
7. Valence Bond Theory and Hybridisation
Valence Bond Theory (VBT) describes covalent bonds as the overlap of atomic orbitals. When atomic orbitals overlap, they form molecular orbitals. Hybridisation is the concept of mixing atomic orbitals to form new hybrid orbitals with different shapes and energies, which are better suited for bonding. For example, the carbon atom in methane (CH$_4$) undergoes sp$^3$ hybridization, forming four equivalent sp$^3$ hybrid orbitals that overlap with hydrogen 1s orbitals to create a tetrahedral molecular geometry.
8. Molecular Orbital Theory
Molecular Orbital Theory (MOT) views molecules as formed by the combination of atomic orbitals to create new molecular orbitals that encompass the entire molecule. These molecular orbitals can be bonding (lower energy, stabilizing) or antibonding (higher energy, destabilizing). The filling of these orbitals with electrons determines the molecule's properties, including bond order, magnetic behavior, and stability. MOT provides a more complete explanation for bonding than VBT, especially for diatomic molecules.
9. Bonding In Diatomic Molecules
The bonding in diatomic molecules, such as O$_2$, N$_2$, and H$_2$, is effectively explained by Molecular Orbital Theory. By constructing molecular orbital diagrams and filling them with electrons according to the Pauli exclusion principle and Hund's rule, we can determine the bond order (half the difference between the number of electrons in bonding and antibonding orbitals), magnetic properties (paramagnetic or diamagnetic), and stability. For example, MOT correctly predicts that O$_2$ is paramagnetic, a feat not explained by simpler theories.
10. Hydrogen Bonding
Hydrogen bonding is a special type of intermolecular attraction that occurs when a hydrogen atom covalently bonded to a highly electronegative atom (like O, N, or F) is attracted to a lone pair of electrons on another electronegative atom in a nearby molecule. This strong intermolecular force significantly affects the physical properties of substances like water, raising its boiling point and melting point, and influencing its unique solvent properties, which are vital for life and many chemical processes in India and globally.