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Atomic molecular theory4/11/2024 ![]() ![]() There is an optimal distance.Īt 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. If the two are too close there is repulsion, if the two are too far there is no attraction. Graph of increasing energy against increasting interatomic distance. However, something else is happening at the same time: as the atoms get closer, the repulsive positive-positive interaction between the two nuclei also begins to increase. This lowers the potential energy of the system, as new, attractive positive-negative electrostatic interactions become possible between the nucleus of one atom and the electron of the second. As they move closer and closer together, orbital overlap begins to occur, and a bond begins to form. If they are too far apart, their respective 1 s orbitals cannot overlap, and thus no covalent bond can form - they are still just two separate hydrogen atoms. How far apart are the two nuclei? That is a very important issue to consider. These two electrons are now attracted to the positive charge of both of the hydrogen nuclei, with the result that they serve as a sort of ‘chemical glue’ holding the two nuclei together. This will be the essential principle of valence bond theory. As we will see, the situation is not quite so simple as that, because the electron pair must still obey quantum mechanics - that is, the two electrons must now occupy a shared orbital space. In simple terms, we can say that both electrons now spend more time between the two nuclei and thus hold the atoms together. When we say that the two electrons from each of the hydrogen atoms are shared to form a covalent bond between the two atoms, what we mean in valence bond theory terms is that the two spherical 1 s orbitals overlap, allowing the two electrons to form a pair within the two overlapping orbitals. The simplest case to consider is the hydrogen molecule, H 2. The overlap of bonding orbitals is substantially increased through a process called hybridization, which results in the formation of stronger bonds. ![]() In this section, we present a quantum mechanical description of bonding, in which bonding electrons are viewed as being localized between the nuclei of the bonded atoms. A more sophisticated treatment of bonding is needed for systems such as these. ![]()
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