Q.Table — Examples: C2H6 (Ethane), C2H4 (Ethene), C2H2 (Ethyne); Structure: -C-C-, C=C, -C≡C-; Type of bond between carbons: single, double, triple; Bond length (nm): 0.154, 0.134, 0.120; Bond enthalpy (kJ mol⁻¹): 348, 612, 837.
What happens to the bond length when unsaturation increases?

Bond angle, bond enthalpy, bond length and bond order are the four standard, experimentally-measurable parameters used to characterise any covalent bond quantitatively. Bond angle is the angle subtended at a central atom by two of its bonds, measured spectroscopically, and reveals how a molecule's orbitals are arranged in space (e.g. 120° in trigonal-planar BF3; 107°18' in pyramidal NH3; 104°35' in bent H2O, progressively compressed by more lone pairs). Bond enthalpy is the energy needed to break one mole of a given bond type in the gaseous state — equal to bond dissociation energy for a diatomic molecule, but an AVERAGE of several slightly different successive dissociation energies for a polyatomic molecule's repeated bond type (e.g. water's two O–H bonds break at 502 and 427 kJ mol⁻¹, averaging to 464.5 kJ mol⁻¹); a larger bond enthalpy means a stronger bond. Bond length is the equilibrium internuclear distance, measured by X-ray/electron diffraction, which INCREASES with the size of the bonded atoms and DECREASES as bond multiplicity increases (single > double > triple bond length, e.g. C–C 154 pm > C=C 133 pm > C≡C 120 pm). Bond order is simply the number of shared electron-pair bonds linking two atoms (1, 2, 3 for single/double/triple); isoelectronic species share identical bond order (N2, CO and NO⁺ all have bond order 3). All four parameters move together in a consistent pattern: as bond order increases, bond enthalpy increases and bond length decreases — so a molecule's relative bond strength and stability can be read directly off any of the three numerical parameters (angle aside, which instead governs shape).