Bond Parameters

Once atoms are joined by a covalent bond, that bond can be described quantitatively using four standard bond parameters.

Bond length is the equilibrium distance between the nuclei of two covalently bonded atoms, usually measured in picometres (pm) or Ångströms. It represents the internuclear separation at which the attractive forces (nucleus–electron) and repulsive forces (nucleus–nucleus, electron–electron) are perfectly balanced, giving the lowest possible potential energy for the bonded pair.

Bond angle is the angle between two bonds sharing a common atom, expressed in degrees. It is determined chiefly by the number and arrangement of electron pairs (bonding and lone) around the central atom — the subject of the VSEPR theory section later in this chapter.

Bond enthalpy (or bond dissociation energy) is the energy required to break one mole of a particular covalent bond in a gaseous molecule, producing gaseous atoms or fragments. It is always a positive quantity (bond breaking requires energy input) and is usually quoted in $\text{kJ mol}^{-1}$. A larger bond enthalpy signals a stronger, harder-to-break bond.

Bond order is the number of covalent bonds (counting a single bond as one, a double bond as two, and a triple bond as three) linking two given atoms. For simple Lewis structures it is easy to read off directly: a $\text{C}-\text{C}$ single bond has bond order $1$, a $\text{C}=\text{C}$ double bond has bond order $2$, and a $\text{C}\equiv\text{C}$ triple bond has bond order $3$.

These four parameters are not independent of one another — they are strongly correlated:

• Higher bond order means shorter bond length. As more electron pairs are shared between the same two nuclei, the electron density between them increases and pulls the nuclei closer together. A $\text{C}\equiv\text{C}$ triple bond ($\approx 120\ \text{pm}$) is shorter than a $\text{C}=\text{C}$ double bond ($\approx 134\ \text{pm}$), which in turn is shorter than a $\text{C}-\text{C}$ single bond ($\approx 154\ \text{pm}$). …