| Hybrid Orbitals
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| An s bond is shorter than a p bond. This means that two atoms joined by s bonds are closer to one another than two atoms joined by p bonds. This is a simple geometric pattern. However, many molecules display an intermediate structure. This intermediate structure is explained by first shifting an s electron to an empty p orbital. This excitation of an electron "invests" some energy to increase the number of bonding sites, thereby increasing the stability of the resulting molecule.
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| The study of solid crystal structure displays the symmetry of the resulting molecule. This means that electrons have shifted to a midpoint level, longer than the shorter s bond but shorter than the longer p bond. These hybrid bonds are the result of hybrid orbitals that are an intermediate positioning of the electrons.
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| Hybridization and Hybrid Orbitals
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| Elements in column IIA of the periodic table have one filled s orbital. In an ionic bond, both electrons are transferred to the nonmetal with which the metal is reacting. The atoms are unable to make a covalent bond without having half-filled orbitals. To do this, a single s electron is excited and moves into an empty p orbital. However, the electrons then move into hybrid sp orbitals, enabling the atom to form sp bonds. An example of this is beryllium dichloride (BeCl2). These molecules have a symmetrical linear structure.
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| In the case of IIIA elements, the atom has a filled s orbital and one half-filled p orbital. Again, one of the s electrons is excited into an empty p orbital. The electrons then move into sp2 orbitals, which are able to form bonds of sp2 molecules. An example is aluminum tribromide (AlBr3). These molecules have a trigonal planar structure.
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| Group IVA elements, such as carbon and silicon, have a filled s orbital and two half-filled p orbitals. By exciting one of the s electrons to half-fill the empty p orbital, a new set of sp3 orbitals are formed. These molecules have a tetrahedral structure. Examples are methane (CH4) and silicon tetrachloride (SiCl4).
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| In elements in the third or higher energy levels are able to create other hybrid molecules, since the d orbitals come into play. Phosphorus has a filled s orbital and three half-filled p orbitals. By exciting a single s electron into the d sublevel, there will now be five half-filled sp3d orbitals. This will result in sp3d bonds, such as phosphorus pentachloride (PCl5). The structure of such molecules are triangular bipyramidal.
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| In a similar manner, higher level VIA elements can create sp3d2 bonds by moving both a single s electron and one electron from the filled p orbital into two empty d orbitals. The result is a molecule with six atoms joined to the central atom. An example is sulfur hexafluoride (SF6). This molecule has an octahedral structure.
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| Higher level noble gases, such as xenon, can create another type of hybrid molecule. These elements have a filled s orbital and three filled p orbitals. If enough energy is applied, p electrons can be excited into empty d orbitals. If one p electron is excited into a d orbital, it will create pd bonds. The three filled p orbitals are around the molecule's "equator" and the two pd bonds are at either pole, creating a linear molecule. An example is xenon difluoride (XeF2). If two p electrons are excited into the d orbitals, it will create p2d2 bonds. Two filled orbitals are above and below the central atom and the four hybrid bonds are evenly distributed around the "equator" of the central atom. An example is xenon tetrachloride (XeCl4). This molecule has a square planar structure.
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| Other examples of hybridization exist, but this is a sufficient quantity to offer the concept in a manageable level.
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