The bigger drawback brought by this energy/force-based mathematical definitions of the chemical bond is the fact that if you stretch any pair of atoms in any molecule you will observe a force that will attract them back to the optimal geometry. This is a necessary conclusion of defining the molecule as a minimum in the potential energy surface, which makes any geometrical distortion raise the molecular energy. It will happen even if these two atoms are far apart and removed by many 'proper' bonds. In other words, according to this definition, there is a pairwise attraction between all the atoms of a molecule, and therefore we will erroneously infer that all its atoms would be directly bonded, even the ones on opposite sides of the molecule. This is exemplified in Fig. , which shows that if you elongate the distance between two hydrogens in ethene, the energy goes up and a force will be created that will push back the hydrogens to their lowest energy. However, it is inconceivable to think under our typical conceptualization that there is a proper chemical bond between these two hydrogens separated by two carbons. How would you call the 'force' between the hydrogens? Certainly not a covalent or ionic bond, not even a van der Waals interaction (although it might be considered a 'through-bond interaction' or in certain cases a 'mechanical bond' ). Without the carbons holding the hydrogens in their positions (as can be done in a model H2 dimer) they would repel each other, clearly indicating that there is no chemical bond between hydrogens in ethene, no matter what the stretching force tells us.