Bond formation between atoms occurs primarily to enable each atom to achieve an inert gas electron configuration in the valence level (a valence octet for all elements except hydrogen which requires only two electrons to achieve the electronic configuration of helium). An atom can achieve an inert-gas electronic configuration by giving up electrons, accepting electrons, or sharing electrons with another atom. An ionic bond is formed when one atom gives up one or more electrons to reach an octet electronic configuration (as a positively charged ion) and a second atom accepts one or more electrons to reach an octet electronic configuration (as a negatively charged ion). For example, the reaction of a cesium atom with a fluorine atom occurs by the transfer of an electron from the cesium atom to the chlorine atom. By doing so, both cesium and chlorine have reached a valence octet electron configuration. The cesium atom has been converted to a positively charged cesium ion with the octet electronic configuration of xenon, and the chlorine has been converted to a negatively charged chloride ion with the octet electronic configuration of argon. The “bond” between cesium and chlorine is due to the electrostatic attraction of the cesium and chloride ions. The reaction of potassium metal with tert-butanol gives an ionic bond between the tert-butoxy anion and a potassium cation by transfer of electrons from potassium to the hydroxyl functional group. Hydrogen is evolved as a by-product. By losing an electron, potassium gains the octet electronic configuration of argon, oxygen has an octet structure (three lone pairs and one pair of shared electrons), and hydrogen has the electronic configuration of helium. (Based on functional group behavior, any other alcohol is predicted to react with potassium in the same way—and they do!) Most bonds in organic molecules, however, are covalent bonds in which electronsare shared between two atoms. Sharing electrons is a way to enable each atom of the bonded pair to reach an octet electronic configuration without having to give up or gain an electron. Covalent bonds are formed by the overlap of singly occupied AOs to form new MOs that contain a pair of electrons. Each atom in essence gains an electron by sharing. The reaction of a chlorine atom with a fluorine atom occurs by the overlap of a singly occupied 3p orbital of chlorine with a singly occupied 2p orbital of fluorine to give a bond between the two atoms that contains two electrons. This is shown both by using Lewis structures and by using orbital pictures. The type of bond formed is called a σ bond because the region of greatest electron density falls on the internuclear axis. In addition to electron sharing to reach octet electronic configurations, other considerations such as the number of bonds to an atom, the number of electron pairs that are shared between two bonded atoms, and repulsion energies that are present between electron pairs require some modification of the picture. These factors can be rationalized by the idea that valence shell atomic orbitals (2s and 2p’s) can combine to form hybrid AOs. These hybrid AOs overlap with AOs of other atoms in the usual fashion to form covalent bonds. Hybrid AOs have energies, shapes, and geometries which are intermediate between the atomic orbitals from which they are formed. Hybridization of AOs is an outgrowth of bond formation that enables atoms to derive the greatest amount of bond energy from electron sharing and to allow bonded atoms to achieve octet electronic configurations. If four single bonds and/or electron pairs originate from a single atom, then the s orbital and the three p orbitals of the valence shell combine to form four equivalent sp3 hybrid orbitals that are then used in bond formation to other atoms. Depending on the number of electrons in the valence shell of the atom, these sp3 hybrid orbitals can contain either a single unpaired electron which can be shared with another atom by overlap and bond formation or an unshared pair of electrons which is normally not involved in bond formation. Thus alkanes, which have all single bonds, have carbon atoms which are sp3 hybridized. For example, methane has four single C–H bonds originating at carbon, and these bonds are σ bonds produced by the overlap of four sp3 hybrid orbitals of carbon with four 1s AOs of four hydrogens to give four sp3–1s σ bonds from carbon to hydrogen. The geometry of the four equivalent sp3 hybrid orbitals (and hence the compound produced by overlap with these orbitals) is tetrahedral. Thus methane has four equivalent C–H σ bonds which point toward the corners of a regular tetrahedron and have H–C–H bond angles of 109.5. In a similar fashion each carbon of propane is sp3 hybridized and tetrahedral since each carbon has four single bonds to other atoms originating from it. For example, the central carbon of propane has two equivalent sp3–1s C–H σ bonds and two equivalent sp3 –sp3 C–C σ bonds. (Note that sp3 orbitals from one carbon can overlap with sp3 orbitals from another carbon to produce carbon–carbon bonds.) The geometry is very close to tetrahedral, but the C–C–C bond angle is slightly larger (111◦) to accommodate the bigger CH3 groups.
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