Q.Define the four bond parameters — bond length, bond angle, bond enthalpy and bond order — and state how bond order is related to bond length and bond enthalpy.

[!TLDR] Bond length, bond angle, bond enthalpy and bond order are the four standard descriptors of a covalent bond, and a higher bond order predicts both a shorter length and a larger enthalpy. [!ANSWER] Bond length = internuclear distance; bond angle = angle between two bonds at a common atom; bond enthalpy = energy to break one mole of the bond; bond order = number of bonds between the two atoms; higher bond order means shorter, stronger (higher-enthalpy) bonds.

Bond length is the equilibrium distance between the nuclei of two covalently bonded atoms, the separation at which attractive and repulsive forces balance to give the lowest potential energy. Bond angle is the angle subtended between two bonds that share a common central atom, set mainly by the number and type (bonding versus lone) of electron pairs on that atom. Bond enthalpy (bond dissociation energy) is the energy that must be supplied to break one mole of a specific bond in the gas phase, producing separated atoms or fragments; it is always positive. Bond order is simply the count of shared electron-pair bonds between two atoms — one for a single bond, two for a double bond, three for a triple bond. These are correlated: increasing the bond order packs more shared electron density between the same two nuclei, pulling them closer together (shorter bond length) and requiring more energy to separate them (higher bond enthalpy). [!ANSWER] Bond order and bond length are inversely related, while bond order and bond enthalpy are directly related — a higher bond order gives a shorter, stronger bond.

Define each of the four parameters in turn, then state and briefly justify the length/enthalpy correlation with bond order.

Confusing bond enthalpy (energy to break the bond, always positive) with the negative enthalpy of bond formation; assuming bond angle depends only on the number of bonded atoms, ignoring lone pairs.