Q.Define the following

Four measurable quantities characterise every covalent bond.

Bond length is the distance between the nuclei of two bonded atoms, equal to the sum of their bonded radii ($r_A + r_B$), measurable by spectroscopy, X-ray diffraction, or electron diffraction. It GROWS with the size of the bonded atoms (C-C, 1.54 Å, is longer than C-N, 1.43 Å) and SHRINKS as bond multiplicity increases (C-C single 1.54 Å > C=C double 1.33 Å > C≡C triple 1.20 Å).

Bond order, in simple Lewis-theory terms, is the number of electron pairs shared between two bonded atoms: 1 for a single bond (H-H), 2 for a double bond (O=O), 3 for a triple bond (N≡N). Molecular orbital theory later generalises this to a formula, $BO = (N_b - N_a)/2$, that can also yield FRACTIONAL values (e.g. NO, bond order 2.5) — see Molecular Orbital Theory.

Bond angle is the fixed angle between two bonds meeting at a common atom, a consequence of covalent bonds being directional in space, measurable by spectroscopy. CH₄ (no lone pairs on C) has bond angle 109°28'; NH₃ (one lone pair on N) has 107°18'; H₂O (two lone pairs on O) has 104°35' — the progressive shrinkage as lone-pair count rises is explained by VSEPR theory's repulsion ordering.

Bond enthalpy is the minimum energy needed to break one mole of a given bond in the gaseous state (kJ mol⁻¹) — larger enthalpy means a stronger bond, and it falls as the bonded atoms grow larger (matching the bond-length trend). For molecules with more than one bond of the same type that are not truly equivalent (e.g. water's two O-H bonds, which individually require 502 and 427 kJ mol⁻¹ to break), an AVERAGE bond enthalpy — here, 464.5 kJ mol⁻¹ — is quoted instead.

These four parameters are deeply interlinked: greater bond order generally means shorter bond length AND greater bond enthalpy (a stronger, tighter bond), while bond angle is governed less by the bond itself and more by the repulsion among ALL the electron pairs (bonding and lone) around the central atom.