When atoms join to form a molecule, the electrons pair up and occupy the lowest available energy level which in the case of hydrogen is the bonding molecular orbital. The asterisk, *, always indicates an antibonding orbital. We indicate the bonding molecular orbital as σ 1s and the antibonding orbital as σ* 1s. Electrons in bonding molecular orbitals contribute to the bond strength. There is a node between the two nuclei and when an antibonding orbital is occupied, it decreases the bond strength. This is the antibonding molecular orbital, and it is higher in energy (see figure above) than the atomic orbitals. In addition to being summed, the wave functions are also combined in a way which cancels the electron density between the two nuclei– this is destructive combination. The electron density is between the two nuclei. This molecular orbital is lower in energy (see figure below) than the atomic orbitals by themselves. The wave functions for the two s orbitals are added, and a bonding molecular orbital is the result. There are two atomic orbitals therefore there will be two molecular orbitals. The two s orbitals overlap to form a covalent bond. Let’s take a look at the hydrogen molecule, H 2. A molecular orbital is a wave function whose square give the probability of finding an electron within a given region of space in a molecule. Recall, an atomic orbital is a wave function whose square gives the probability of finding an electron within a given region of space in an atom. In molecular orbital theory, molecular orbitals are formed by the combination of atomic orbitals. Molecular orbitals are to molecules as atomic orbitals are to atoms. A molecular oribital, MO, is a mathematical description of the region in a molecule where there is a high probability of finding electrons. An alternative scheme to valence bond theory uses molecular orbitals. In Valence Bond theory, on the the valence orbitals involved in bonding are modified. The valence bond model is not able to explain all features of bonding.
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