Silicon-hybridized orbitals

Why does siiicon not form itt bonds as readiiy as carbon, its Group 14 neighbor The answer can be found by examining the sizes of the vaience orbitais of siiicon and carbon. Figure 10-23 compares side-by-side p-orbital overiap for carbon atoms and siiicon atoms. Notice that for the larger silicon atom, the 3 orbitals are too far apart for strong side-by-side overiap. The resuit is a very weak ttbond. Consequently, silicon makes more effective use of its vaience orbitais to form four a bonds, which can be described using s p hybrid orbitals. 

Both sand and silt surfaces are dominated by oxygen and its lone pairs of electrons in p orbitals. In some instances, broken surfaces may also have silicon-hybridized sp3 orbitals4 available for bonding. Comparison of sand, silt, and clay reveals the surface area of sand and silt to be low and the interaction between surface bonding orbitals and components in the surrounding medium relatively weak. 

It can be imagined that the bonds can be broken at any location, that is, with an oxygen, hydroxy, silicon, or aluminum exposed. In this case, it could further be imagined that s-, p-, and sp3-hybridized orbitals would be on the surface. This... 

Clays contain aluminum oxides in addition to silicon as Si02 and its polymeric forms. Again, there are the p orbitals of oxygen and sp-hybridized orbitals from aluminum, which may result in end-on-end or side-by-side bonding with the same restrictions encountered with silicon. 

Fig. 2-12. Electron energy band formation of silicon crystals from atomic frontier orbitals number of silicon atoms in crystal r = distance between atoms rg = stable atom-atom distance in crystals, sp B8 = bonding band (valence band) of sp hybrid orbitals sp ABB = antibonding band (conduction band) of sp hybrid orbitals.

The formation of a Si crystal is shown in Fig. 1.10. Aside from the core, each Si atom has four valence electrons two 3s electrons and two 3p electrons. To form a Si crystal, one of the 3s electrons is excited to the 3p orbital. The four valence electrons form four sp hybrid orbitals, each points to a vertex of a tetrahedron, as shown in Fig. 1.10. Thpse four sp orbitals are unpaired, that is, each orbital is occupied by one electron. Since the electron has spin, each orbital can be occupied by two electrons with opposite spins. To satisfy this, each of the directional sp orbitals is bonded with an sp orbital of a neighboring Si atom to form electron pairs, or a valence bond. Such a valence bonding of all Si atoms in a crystal form a structure shown in (b) of Fig. 1.10, the so-called diamond structure. As seen, it is a cubic crystal. Because all those tetrahedral orbitals are fully occupied, there is no free electron. Thus, similar to diamond, silicon is not a metal. 

The basic structural element in both vitreous and crystalline silica is the Si04-4 tetrahedron, which arises from the sp hybrid orbitals of the silicon. Each silicon atom sits in the center of the tetrahedron surrounded by four oxygen atoms that hold the comer positions. Tetrahedrons bond together by comer sharing. In a properly developed structure, each oxygen is shared by only two tetrahedrons. This bonding scheme can produce a large variety of three-dimensional structures and is the reason that silica has a number of crystalline phases. 

A similar diagram (II) may be drawn for species that make use of carbon or silicon as the bridging atom. In these systems the orbital used for bridge formation may be described in terms of a hybrid orbital from the bridging atom which becomes five- (or six)-coordinate. 

With atoms such as carbon and silicon, the valence-state electronic configuration to form four covalent bonds has to be (s)1(px)1(py)1(ps)1. Repulsion between the electron pairs and between the attached nuclei will be minimized by formation of a tetrahedral arrangement of the bonds. The same geometry is predicted from hybridization of one s and three p orbitals, which gives four sp3-hybrid orbitals directed at angles of 109.5° to each other. The predicted relative overlapping power of s -hybrid orbitals is 2.00 (Figure 6-10). 

Unlike carbon, the valence shell of the silicon atom has available d orbitals. In many silicon compounds, the d orbitals of Si contribute to the hybrid orbitals and Si forms more than four 2c-2e covalent bonds. For example, Sib s- uses sp3d hybrid orbitals to form five Si-F bonds, and SiF62 uses sp3d2 hydrid orbitals to form six Si-F bonds. 

D A barrier height of 0.8 kcal/mole is obtained in the absence of silicon d orbitals, i) Such a picture is reminiscent of the Walsh model of aziridine -where the nitrogen atom is sp2 hybridized as in methylene imine 171.172),... 

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