Valence bond theory covalent configuration

A vexing puzzle in the early days of valence bond theory concerned the bonding in methane (CH4). Since covalent bonding requires the overlap of half-filled orbitals of the connected atoms, carbon with an electron configuration of ls 2s 2p 2py has only two half-filled orbitals (Figure 1.20a), so how can it have bonds to four hydrogens ... 

The basic tenet of covalent bonding between atoms in molecules is called the valence bond theory. The basic idea is that covalent bonds are formed by the overlap of atomic orbitals of different atoms. The two electrons are involved in a paired spin that is shared by an atomic orbital of each of the two involved atoms. The more overlap you see, the stronger will be the covalent bond between the atoms. The atomic orbitals can be the same as the original orbitals of each atom however, the bond can t happen if the original orbitals are maintained. The orbitals must come together into a new configuration, called reconfigured orbitals or hybridized orbitals . 

The VSEPR model correctly predicts that Bep2 is linear with two identical Be—F bonds. How can we use valence-bond theory to describe the bonding The electron configuration of F (ls 2s 2p ) indicates an unpaired electron in a 2p orbital. This electron can be paired with an unpaired Be electron to form a polar covalent bond. Which orbitals on the Be atom, however, overlap with those on the F atoms to form the Be — F bonds ... 

The problem of directed valence is treated from a group theory point of view. A method is developed by which the possibility of formation of covalent bonds in any spatial arrangement from a given electron configuration can be tested. The same method also determines the possibilities of double and triple bond formation. Previous results in the field of directed valence are extended to cover all possible configurations from two to eight s, p, or d electrons, and the possibilities of double bond formation in each case. A number of examples are discussed. 

An important property of covalent molecules is that these covalent bonds have directional properties, and the molecules have three-dimensional shape. What determines this shape is the number of electron pair bonds in the valence shell configuration of the central atom (Figure 6.1), about which the shape of the molecule is described. The VSEPR theory... 

The molecule sulfur hexafluoride (SFA has recently challenged both molecular spectroscopy with its unexpected rotational spectra 29) and electronic structure theories with novel correlation effects (30,31,5). The electronic structure must explain the molecule s high stability, octahedral symmetry, and, most importantly, provide a simple picture of the bonding. At first glance, the traditional chemical models do not appear to be appropriate because sulfur seemingly forms six bonds to fluorines, yet the sulfur s2pA valence configuration allows for at most two covalent bonds. 

The electrons responsible for bonding are those in the outer shell, or valence shell, of an atom. Valence shell electrons are those that were not present in the preceding noble gas orbitals. We will focus attention on these electrons in our discussion of covalent bonding. The electrons in the lower energy nohle gas configuration are not directly involved in covalent bonding and are often referred to as core electrons. Lewis formulas show the number of valence shell electrons in a polyatomic molecule or ion (see Sections 7-5 through 7-9). We will write Lewis formulas for each molecule or polyatomic ion we discuss. The theories introduced in this chapter apply equally well to polyatomic molecules and to ions. 

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