See a table of bond lengths and bond energies For the H 2 molecule, the distance is 74 pm (picometers, 10 -12 meters). Likewise, the difference in potential energy between the lowest energy state (at the optimal internuclear distance) and the state where the two atoms are completely separated is called the bond dissociation energy, or, more simply, bond strength. For the hydrogen molecule, the H-H bond strength is equal to about 435 kJ/mol.Įvery covalent bond in a given molecule has a characteristic length and strength. In general, the length of a typical carbon-carbon single bond in an organic molecule is about 150 pm, while carbon-carbon double bonds are about 130 pm, carbon-oxygen double bonds are about 120 pm, and carbon-hydrogen bonds are in the range of 100 to 110 pm. The strength of covalent bonds in organic molecules ranges from about 234 kJ/mol for a carbon-iodine bond (in thyroid hormone, for example), about 410 kJ/mole for a typical carbon-hydrogen bond, and up to over 800 kJ/mole for a carbon-carbon triple bond. There is a defined optimal distance between the nuclei in which the potential energy is at a minimum, meaning that the combined attractive and repulsive forces add up to the greatest overall attractive force. This optimal internuclear distance is the bond length. When the two nuclei are ‘too close’, we have an unstable, high-energy situation. Electrons fill the lower-energy bonding orbital before the higher-energy antibonding orbital, just as they fill lower-energy atomic orbitals before they fill higher-energy atomic orbitals.At first this repulsion is more than offset by the attraction between nuclei and electrons, but at a certain point, as the nuclei get even closer, the repulsive forces begin to overcome the attractive forces, and the potential energy of the system rises quickly. Hence, these orbitals are called antibonding orbitals.

The attractive force between the nuclei and these electrons pulls the two nuclei apart. Electrons in the σ ∗ s orbitals are located well away from the region between the two nuclei. Adding electrons to these orbitals creates a force that holds the two nuclei together, so we call these orbitals bonding orbitals. Electrons in a σ s orbital are attracted by both nuclei at the same time and are more stable (of lower energy) than they would be in the isolated atoms. The asterisk signifies that the orbital is an antibonding orbital.

The out-of-phase addition (which can also be thought of as subtracting the wave functions) produces a higher energy σ* s molecular orbital (read as “sigma-s-star”) molecular orbital in which there is a node between the nuclei. The two types are illustrated/ The in-phase combination produces a lower energy σ s molecular orbital (read as “sigma-s”) in which most of the electron density is directly between the nuclei. There are two types of molecular orbitals that can form from the overlap of two atomic s orbitals on adjacent atoms.

In phase combination of atomic orbitals