Ch 5Chemical Bonding — Class 11 Chemistry, concept-first.

Atoms in a compound are held together by chemical bonds — the forces responsible for that binding. A chemical bond can form in one of two ways: by the loss and gain of electrons between two atoms, or…

W. Kossel and G. N. Lewis, working independently, gave the first satisfactory electron-based account of chemical valence in 1916, building their Electronic Theory of Valence on the observed chemical i…

An ionic (or electrovalent) bond forms when one atom transfers one or more valence electrons completely to another atom, and the resulting oppositely-charged ions are then held together purely by elec…

Ionic solids are crystalline arrays of cations and anions held together purely by coulombic (electrostatic) attraction, with every cation surrounded by a fixed number of anions and vice versa in a reg…

Lewis (1919) proposed that some atoms instead reach a stable noble-gas configuration by SHARING one or more electron pairs with another atom, rather than by transferring electrons outright — each atom…

A Lewis dot structure is a diagram that pictures a molecule's or ion's bonding purely in terms of shared electron pairs (bonds) and lone pairs, built around the octet rule.

Writing a Lewis dot structure follows a fixed eight-step procedure. First, add up the TOTAL number of valence electrons contributed by every atom in the molecule (adjusting for any overall ionic charg…

When a polyatomic species can be drawn as more than one valid Lewis structure — differing only in exactly where its double or triple bond(s) sit, never in which atoms are bonded to which — formal char…

Despite its usefulness, the octet rule fails to hold — or fails to explain a molecule's real behaviour — in three distinct ways.

Because a Lewis structure alone says nothing about a molecule's three-dimensional shape, the Valence Shell Electron Pair Repulsion (VSEPR) theory — proposed by Sidgwick and Powell — was developed spec…

Valence bond (VB) theory, developed by Heitler and London on the basis of wave mechanics and subsequently extended by Pauling and Slater, is a more advanced, quantum-mechanically grounded theory of co…

Valence bond theory rests on eight postulates. (i) A covalent bond forms when a half-filled valence orbital of one atom overlaps with a half-filled valence orbital of another atom.

Bond formation lowers a system's total energy because of a changing balance between newly-arising attractive and repulsive forces as two atoms approach each other.

The strength of a covalent bond is governed by how much its two contributing atomic orbitals overlap — a GREATER extent of overlap gives a STRONGER bond — and how much overlap is geometrically possibl…

Plain (unhybridized) valence-bond overlap correctly explains simple diatomics like H2, F2 and HF, but it cannot explain several well-known observations: how beryllium forms TWO covalent bonds, how bor…

Depending on how many 's' and 'p' orbitals combine, hybridization on one atom produces one of three types of hybrid orbital, each with its own characteristic geometry.

Valence bond theory, despite the specific limitations catalogued in the next section, introduced five genuinely new ideas into the theory of chemical bonding that are still used today: (i) the DELOCAL…

Valence bond theory, for all its successes, has three specific, well-documented limitations. First, it explains only the ORDINARY covalent bond, where each of the two bonded atoms contributes exactly…

Molecular Orbital (MO) theory is introduced as a further theoretical advance beyond valence bond theory.

Molecular orbitals are built from atomic orbitals through the Linear Combination of Atomic Orbitals (LCAO) method.

Not every pair of atomic orbitals can combine into molecular orbitals — three conditions must ALL be satisfied first.

In a diatomic molecule, every molecular orbital formed from combining atomic orbitals falls into one of two symmetry classes: a sigma (σ) molecular orbital, which is symmetrical all the way around the…

Combining two 1s atomic orbitals gives the pair σ1s (bonding) and σ1s (antibonding); combining two 2s orbitals similarly gives σ2s and σ2s; and the three 2p orbitals on one atom combine with the three…

Seven key ideas summarise molecular orbital theory as a whole. First, molecular orbitals (MOs) are to a MOLECULE what atomic orbitals (AOs) are to a single ATOM — each is a region of space, derived fr…

Applying the molecular-orbital energy-level ordering (section 5.5.4) and the bond-order formula to five real second-row homonuclear diatomic molecules gives a consistent, testable picture of each one'…

Having established two competing theories of covalent bonding (valence bond theory and molecular orbital theory), this section turns to four measurable PARAMETERS that together characterise any covale…

Bond angle is defined as the angle between the orbitals that hold a central atom's bonding electrons — equivalently, the angle subtended at the central atom by any two of its bonds.

Bond enthalpy is the amount of energy required to break one mole of a given TYPE of bond between two atoms, with both atoms in the gaseous state.

Bond length is the equilibrium distance between the nuclei of two covalently bonded atoms in a molecule, measured experimentally by X-ray diffraction and electron diffraction techniques.

In the simple Lewis-structure picture, bond order is just the NUMBER of shared electron-pair bonds linking two atoms in a molecule — 1 for a single bond (H2), 2 for a double bond (O2), 3 for a triple…

A covalent bond's shared electron pair is not always shared perfectly equally between its two atoms — whether it is depends on whether the two bonded atoms have the same or different electronegativity…

Building on the qualitative picture of polar bonds already developed in section 5.6.5, dipole moment is formally DEFINED as the product of the magnitude of the separated charge (Q) and the distance of…

NOTE ON SOURCE NUMBERING: the source PDF prints this section's heading number as '5.8.7', which does not fit the chapter's own numbering scheme (it directly follows 5.7 Dipole moment, and the source P…

NOTE ON SOURCE NUMBERING: the source PDF heads this section '5.8', immediately after the section renumbered 5.8 above (see that section's note) — kept here as its own distinct section, renumbered 5.9,…