Chemical bonding is a fundamental idea in chemistry that helps to explain other concepts like molecules and reactions. Scientists would be unable to explain why atoms are attracted to one another or how products are generated following a chemical reaction if it were not for it. To comprehend the notion of bonding, one must first comprehend the fundamentals of atomic structure.
The Rule of the Octet
Chemical Combination Modes
Bonds that are ionic or electrovalent
I.E = 496 kJ/mole in the production of the Na+ ion, for example.
(i) Ion size: The smaller the ions, the higher the lattice energy.
(ii) Ion charge: The larger the ionic charge, the stronger the interionic attraction and the higher the lattice energy.
Characteristics of Ionic Compounds
(i) Physical'State: They are crystalline solids with a crystal lattice structure. Ionic compounds, unlike other gaseous molecules such as H2, N2, 02, and Cl2, do not exist as single molecules.
(ii) Melting and boiling points: Ionic compounds have high melting and boiling points due to the high interionic force between them.
(iii) Solubility: They dissolve in polar solvents like water but not in organic solvents like benzene, CCl4, and so forth.
Lewis-Langmuir Concept of Covalent Bond
The following steps can be used to write the Lewis dot Structure:
(i) Determine the combined atoms' total number of valence electrons.
(ii) Each anion adds one electron to the equation, whereas each cation subtracts one electron. This provides you the total quantity of electrons you'll need to distribute.
(iii) By understanding the merging atoms' chemical symbols.
(iv) After inserting shared pairs of electrons for a single bond, the leftover electrons can be used to account for multiple bonds or lone pairs. It should be emphasised that each atom's octet should be finished.
Dipole Moment's Applications
(i) To figure out how polar the molecules are.
(ii) Identifying the molecular shapes.
Molecules having zero dipole moment, for example, will be linear or symmetrical. Molecules with asymmetrical shapes will be bent or angular in form.
(For example, NH3 has a D of 1.47).
(iii) When computing the fraction of polar bonds that are ionic.
Sidgwick and Powell introduced the Valence Shell Electron Pair Repulsion (VSEPR) Theory in 1940, based on the repulsive behaviour of electron pairs in the valence shell of atoms. Nyholm and Gillespie worked on it further (1957).
The following are the main postulates:
I The amount of electron pairs (bound or non-bonded) surrounding the core atoms determines the precise structure of the molecule.
(ii) Because the electron pairs exist surrounding the centre atom and the electron clouds are negatively charged, they have a propensity to reject one other.
(iii) Electron couples strive to align themselves in such a way that the rupulsion between them is minimised.
(iv) The valence shell is modelled as a sphere with electron pairs arranged at their greatest separation.
(v) A multiple bond is considered as a single electron pair, and the electron pairs that make up the bond are treated as single pairs as well.
Bond Theory of Valence
(i) Each atom's nucleus is attracted to its own electron as well as the electron of the other atom, and vice versa.
(ii) Between the electrons of two atoms and the nuclei of two atoms, repulsive forces emerge. Attractive forces pull the two atoms closer together, while repulsive forces pull them apart.
Orbital Overlap Types
The covalent bonds are classified as sigma () or pi () bonds, depending on the type of overlapping.
(i) Bond (Sigma): The end-to-end (head-on) overlap of bonding orbitals along the internuclear axis forms the sigma bond.
There are three forms of axial overlap between these orbitals:
• s-s overlapping: Two half-filled s-orbitals overlap along the internuclear axis in this example, as seen below:
• s-p overlapping: This sort of overlapping happens when one atom's half-filled s-orbitals overlap with the half-filled p-orbitals of another atom.
• p-p overlapping: This sort of overlapping occurs between half-filled p-orbitals of two approaching atoms.
(ii) bond (pi): When atomic orbitals overlap in such a way that their axes stay parallel to each other and perpendicular to the internuclear axis, a bond is established. The orbital is generated as a result of lateral or sidewise overlapping.
Sigma and pf Bond Strength
Hybridisation is the process of combining orbitals with somewhat varying energies in order to redistribute their energies and generate a new set of orbitals with comparable energy and shapes.
Hybridization's Key Characteristics:
The hybridisation involves orbitals of almost equal energy.
The number of hybrid orbitals created is equal to the number of atomic orbitals combined; (iii) The kind of hybridisation can reveal the geometry of a covalent molecule.
Hybrid orbitals are more effective than pure atomic orbitals in creating stable bonds. Conditions required for hybridisation: I Valence shell orbitals participate in hybridisation.
The energy of the orbitals engaged in hybridisation should be almost equal.
Prior to hybridisation, electron promotion is not a need.
In some circumstances, filled valence shell orbitals are also involved in hybridization.
Molecular Orbital Formation: Atomic Orbitals in a Linear Combination (LCAO)
(1) The combining atomic orbitals must have about identical energy.
(2) Around the molecule axis, the combining atomic orbitals must have the same symmetry.
(3) As much as possible, the combining atomic orbitals must overlap.