Covalent bonding hybridization

Crystalline carbon exists in natural deposits in three crystalline modifications a and 8-graphite and diamond. Synthetic graphites contain only the a-form, from which the (3 can be made by mechanical working. This introduces the possibility of the transformation a p during a hardness test. Localized transformation from diamond to a metallic carbon 8-graphite could also be considered, but in this case a radical rearrangement of covalent bond hybridization would be required from sp to sp + p such that time would be a problem. 

When elements in Period 2 form covalent bonds, the 2s and 2p orbitals can be mixed or hybridised to form new, hybrid orbitals each of which has. effectively, a single-pear shape, well suited for overlap with the orbital of another atom. Taking carbon as an example the four orbitals 2s.2p.2p.2p can all be mixed to form four new hybrid orbitals (called sp because they are formed from one s and three p) these new orbitals appear as in Figure 2.9. i.e. they... 

Organic molecules are generally composed of covalent bonded atoms with several well-defined hybridization states tending to have well-understood preferred geometries. This makes them an ideal case for molecular mechanics parameterization. Likewise, organic molecules are the ideal case for semiempirical parameterization. 

The compounds of carbon and silicon with hydrogen would be expected to be completely covalent according to these models, but the dhectionality of the bonds, which is towards the apices of a regular tetrahedron, is not explained by these considerations. Another of Pauling s suggestions which accounts for this type of directed covalent bonding involves so-called hybrid bonds. 

An extreme example of hybidization is the structure proposed for sulphur hexafluoride, SFe. The six S-F bonds are dhected to the apices of a regular octahedron. An aiTangement which would satisfy this number of covalent bonds is sp d hybridization. The ground state of the sulphur atom is s p° and... 

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

Examine the eleetrostatie potential map for ketene. Which (non-hydrogen) atom is most eleetron poor, and which regions around this atom are most electron poor After oxygen, which atom is most electron rich, and which regions are most electron rich Account for these data with a diagram that shows the orbitals on each atom, their orientation and electron occupancy, and whether or not they participate in covalent bonds (assume that oxygen is sp hybridized). 

The valence-bond concept of orbital hybridization described in the previous four sections is not limited to carbon compounds. Covalent bonds formed by-other elements can also be described using hybrid orbitals. Look, for instance, at the nitrogen atom in methylamine, CH3NH2, an organic derivative of ammonia (NH3) and the substance responsible for the odor of rotting fish. 

Phosphorus and sulfur are the third-row analogs of nitrogen and oxygen, and the bonding in both can be described using hybrid orbitals. Because of their positions in the third row, however, both phosphorus and sulfur can expand their outer-shell octets and form more than the typical number of covalent bonds. Phosphorus, for instance, often forms five covalent bonds, and sulfur occasionally forms four. 

We said in Section 1.5 that chemists use two models for describing covalent bonds valence bond theory and molecular orbital theory. Having now seen the valence bond approach, which uses hybrid atomic orbitals to account for geometry and assumes the overlap of atomic orbitals to account for electron sharing, let s look briefly at the molecular orbital approach to bonding. We ll return to the topic in Chapters 14 and 15 for a more in-depth discussion. 

Molecular orbital (MO) theory describes covalent bond formation as arising from a mathematical combination of atomic orbitals (wave functions) on different atoms to form molecular orbitals, so called because they belong to the entire molecule rather than to an individual atom. Just as an atomic orbital, whether un hybridized or hybridized, describes a region of space around an atom where an electron is likely to be found, so a molecular orbital describes a region of space in a molecule where electrons are most likely to be found. 

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the valence bond model, which uses hybrid orbitals to account for the observed shapes of organic molecules. Before going on to a systematic study of organic chemistry, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the electrons in a bond are not shared equally between atoms. 

In diamond, each carbon atom forms single bonds with four other carbon atoms arranged tetrahedrally around it The hybridization in diamond is sp3. The three-dimensional covalent bonding contributes to diamond s unusual hardness. Diamond is one of the hardest substances known it is used in cutting tools and quality grindstones (Figure 9.12). 

In Chapter 7, we used valence bond theory to explain bonding in molecules. It accounts, at least qualitatively, for the stability of the covalent bond in terms of the overlap of atomic orbitals. By invoking hybridization, valence bond theory can account for the molecular geometries predicted by electron-pair repulsion. Where Lewis structures are inadequate, as in S02, the concept of resonance allows us to explain the observed properties. 

FIGURE 5.21 The structure of diamond, Each sphere represents the location of the center of a carbon atom. Each atom is at the center of a tetrahedron formed hy the sp1 hybrid covalent bonds to each of its four neighbors. 

It is customary to use the line between two atomic symbols, A—B, to represent a normal covalent bond, with the usual amount of ionic character. In the discussion that follows, A B is used to represent a pure covalent single bond, and A+B to represent the ionic structure that is hybridized with A B to give A—B. A pure covalent double bond is represented by A=B. Thus in the molecule NF3 we might describe each NF bond as involving N+ F , N F , and N F +, that is, as having some covalent double-bond character. 

---