Valence bond theory assumes that the orbitals of bonding electron pairs overlap. An important caveat to this theory is that overlapping orbitals, also known as bonding orbitals, are not the same as nonbonding orbitals. According to valence bond theory, a hybrid orbital is an electron orbital that forms when two atomic orbitals combine to form a covalent bond. Hybridization is a basic model that is used to qualitatively describe numbers of VSEPR pairs about various atoms, taking into account where they are bonding and nonbonding as well as their arrangements in space.
The number of hybrid orbitals is equal to the total number of native orbitals (the atomic orbitals in their unbound state) involved in the hybridization. In other words, if four orbitals combine, the result is four hybrid orbitals. The new hybrid orbitals have a potential energy that is between the energy of the native orbitals.
Hybridized orbitals have a different shape than the native orbitals. These hybrid-orbital geometries align with the electron-pair geometries of the VSEPR theory. For example, consider the molecule nitrogen triiodide (NI3). VSEPR theory predicts the electron-pair geometry for the central nitrogen atom in this molecule will be tetrahedral. Nitrogen has five valence electrons: two in the 2s orbital and three in the 2p orbitals. This means the 2s orbital has two electrons and all of the three 2p orbitals have one electron each.A hybrid orbital that forms when one s and three p orbitals combine is an sp3 hybrid orbital. When nitrogen bonds with three iodine atoms, its one s and three p orbitals combine into four sp3 orbitals. Hybrid orbital notation keeps track of the orbitals with superscripts. The superscript 3 indicates that three p orbitals are part of the hybridization. In sp3 hybridization, there are four orbitals with a tetrahedral geometry, and the central nitrogen atom in NI3 has a tetrahedral electron-pair geometry.