A. Orbital hybridization and chemical bonding
Stable, covalent bonds between nonmetal atoms are produced when orbitals of the two atoms form molecular orbitals that are occupied by one electron from each of the atoms. Thus, the four bonding electrons of the carbon atom occupy 2s and 2p atomic orbitals (1a ).The 2s orbital is spherical in shape, while the three 2p orbitals ares hapedl iked umbbells arranged along the x, y, and z axes. It might therefore be assumed that carbon atoms should form at least two different types of molecular orbital. However, this is not normally the case. The reason is an effect known as orbital hybridization .Combination of the s orbital and the three p orbitals of carbon gives rise to four equivalent, tetrahedrally arranged sp 3 atomic orbitals (sp 3 hybridization ).When these overlap with the 1s orbitals of H atoms, four equivalent σ -molecular orbitals (1b )a reformed. For this reason, carbon is capable of forming four bonds —i.e., it has a valency of four. Single bonds between nonmetal atoms arise in the same way as the four σ or single bonds in methane (CH4).For example, the hydrogen phosphate ion (HPO4 2 –)and the ammonium ion (NH4 + )are also tetrahedral in structure (1c ).
A second common type of orbital hybridization involves the 2s orbital and only two of the three 2p orbitals (2a).This process is therefore referred to as sp 2 hybridization .
The result is three equivalent sp 2 hybrid orbitals lying in one plane at an angle of 120 °to one another. The remaining 2px orbital is oriented perpendicular to this plane. In contrast to their sp 3 counterparts, sp 2 –hybridized atoms form two different types of bond when they combine into molecular orbitals (2b ).The three sp 2 orbitals enter into σ bonds, as described above. In addition, the electrons in the two 2px orbitals, known as .electrons , combine to give an additional, elongated π molecular orbital, which is located above and below the plane of the σ bonds. Bonds of this type are called double bonds .They consist of a σ bond and a π bond, and arise only when both of the atoms involved are capable of sp 2 hybridization. In contrast to single bonds, double bonds are not freely rotatable, since rotation would distort the π molecular orbital. This is why all of the atoms lie in one plane (2c );in addition, cis –trans isomerism arises in such cases .
Double bonds that are common in biomolecules are C=C and C=O.C=N double bonds are found in aldimines (Schiff bases) .
Many molecules that have several double bonds are much less reactive than might be expected. The reason for this is that the double bonds in these structures cannot be localized unequivocally. Their π orbitals are not confined to the space between the double-bonded atoms, but form a shared, extended .-molecular orbital .Structures with this property are referred to as resonance hybrids ,because it is impossible to describe their actual bonding structure using standard formulas. One can either use what are known as resonance structures —i.e., idealized configurations in which π electrons are assigned to specific atoms (cf.pp.32 and 66,for example)—or one can use dashed lines as in Fig.B to suggest the extent of the delocalized orbitals.
Resonance-stabilized systems include carboxylate groups, as in formate ;a liphatic hydrocarbons with conjugated double bonds,
such as 1,3-butadiene ;and the systems known as aromatic ring systems .The best-known aromatic compound is benzene, which has six delocalized π electrons in its ring. Extended resonance systems with 10 or more π electrons absorb light within the visible spectrum and are therefore colored. This group includes the aliphatic carotenoids ,for example, as well as the heme group, in which 18 π electrons occupy an extended molecular orbital .