mol – Similarly the bond enthalpy for molecules containing multiple bonds, for example O and N will be as under : O (O = O) (g) → O(g) + O(g); ∆ a H = kJ mol – N (N ≡ N) (g) → N(g) + N(g); ∆ a H = . kJ mol – It is important that larger the bond dissociation enthalpy, stronger will be the bond in the molecule. For a heteronuclear diatomic molecules like HCl, we have HCl (g) → H(g) + Cl (g); ∆ a H = . kJ mol – In case of polyatomic molecules, the measurement of bond strength is more complicated.
For example in case of H O molecule, the enthalpy needed to break the two O – H bonds is not the same. Table . Average Bond Lengths for Some Single, Double and Triple Bonds Bond Type Covalent Bond Length (pm) O–H C–H N–O C–O C–N C–C C=O N=O C=C C=N C ≡ N C ≡ C Table . Bond Lengths in Some Common Molecules Molecule Bond Length (pm) H (H – H) F (F – F) Cl (Cl – Cl) Br (Br – Br) I (I – I) N (N ≡ N) O (O = O) HF (H – F) HCl (H – Cl) HBr (H – Br) HI (H – I) Table .
Covalent Radii, * r cov /(pm) * The values cited are for single bonds, except where otherwise indicated in parenthesis. (See also Unit for periodic trends). H O(g) → H(g) + OH(g); ∆ a H = kJ mol – OH(g) → H(g) + O(g); ∆ a H = kJ mol – The difference in the ∆ a H value shows that the second O – H bond undergoes some change because of changed chemical environment. This is the