Hybrid orbitals overlap between

Methane forms from orbital overlap between the hydrogen 1 S orbitals and the s hybrid orbitals of the carbon atom. 

The hosts outlined above generally utilize their metal centers to provide interactions with anions. The orbital overlap between unfilled heteroelement orbitals and filled anion orbitals yields bonding interactions. In the case of the mercury based receptors, the role of the metal atoms is partly structural, with their sp hybridization linear geometry specifically yielding large macrocyclic cavities of suitable diameters for anion encapsulation. The metal atoms have often been functionally used as an NMR handle on the anion coordination process. 

In Chapter 3, Section 3.5, molecular orbitals described the bonding in alkanes using the hybridization model. Specifically, sp hybrid orbitals overlap to form a sigma-covalent bond (a o-bond). It is possible to have two covalent bonds between adjacent carbon atoms, a carbon-carbon double bond. One of the two bonds is the usual o-bond, but the other is called a 7i-bond. Hydrocarbons that contain one 7t-bond are called alkenes. In other words, an alkene will have a C=C unit. Each carbon atom of the C=C unit will have four bonds, but only three of the bonds are o-bonds, and the fourth bond is a 7i-bond. 

FIGURE 2.4 A comparison between overlap of a hydrogen Ir orbital with (a) an sp hybrid orbital and (b) an unhybridized 2p orbital. With the sp hybrid orbital, overlap is maximized because the non-overlapping back lobe is small. With the unhybridized 2p orbital, all of the non-overlapping back lobe is wasted. ... 

To describe the multiple bonding in ethylene, we must distinguish between two kinds of bonds. A cylindrical shape about the bond axis. It is formed either when two s orbitals overlap, as in H2 (Figure 10.25A), or when an orbital with directional character, such as a p orbital or a hybrid orbital, overlaps another orbital along their axis (Figure 10.25B). The bonds we discussed in the previous section are a bonds. 

Covalent bonds (the result of the sharing of one or more pairs of electrons in a region of orbital overlap between two or perhaps more atoms) are directional interactions as opposed to ionic and metallic bonds, which are nondirectional. A good example of a covalent network crystal is diamond, shown in Figure 7.3. Note that each carbon atom is best thought of as being r -hybridized and that to maximize the overlap of these hybrid orbitals, a C-C-C bond angle of 109.5° is necessary. 

Although thiols are weak acids, they are far stronger acids than alcohols. The sulfur atom of a thiol is sp -hybridized, and orbital overlap between this orbital and the Is orbital of hydrogen is poor. As a result, the S—H bond is nonpolar, and it is much weaker than the short, polar, O—H bond of alcohols. 

Orbital overlap between two s/j -hybridized atoms can result in a o bond and a n bond. 

Here, the bonding between carbon atoms is briefly reviewed fuller accounts can be found in many standard chemistry textbooks, e.g., [1]. The carbon atom [ground state electronic configuration (ls )(2s 2px2py)] can form sp sp and sp hybrid bonds as a result of promotion and hybridisation. There are four equivalent 2sp hybrid orbitals that are tetrahedrally oriented about the carbon atom and can form four equivalent tetrahedral a bonds by overlap with orbitals of other atoms. An example is the molecule ethane, CjH, where a Csp -Csp (or C-C) a bond is formed between two C atoms by overlap of sp orbitals, and three Csp -Hls a bonds are formed on each C atom. Fig. 1, Al. 

In the third type of hybridisation of the valence electrons of carbon, two linear 2sp orbitals are formed leaving two unhybridised 2p orbitals. Linear a bonds are formed by overlap of the sp hybrid orbitals with orbitals of neighbouring atoms, as in the molecule ethyne (acetylene) C2H2, Fig. 1, A3. The unhybridised p orbitals of the carbon atoms overlap to form two n bonds the bonds formed between two C atoms in this way are represented as Csp Csp, or simply as C C. 

Figure 1.14 The structure of ethylene. Orbital overlap of two sp hybridized carbons forms a carbon-carbon double bond. One part of the double bond results from a (head-on) overlap of sp2 orbitals (green), and the other part results from (sideways) overlap of unhybridized p orbitals (red/blue). The ir bond has regions of electron density on either side of a line drawn between nuclei.

In contrast with amines, amides (RCONH ) are nonbasic. Amides don t undergo substantial protonation by aqueous acids, and they are poor nucleophiles. The main reason for this difference in basicity between amines and amides is that an amide is stabilized by delocalization of the nitrogen lone-pair electrons through orbital overlap with the carbonyl group. In resonance terms, amides are more stable and less reactive than amines because they are hybrids of two resonance forms. This amide resonance stabilization is lost when the nitrogen atom is protonated, so protonation is disfavored. Electrostatic potential maps show clearly the decreased electron density on the amide nitrogen. 

In the interaction of the local 2pv orbitals, two more bonding molecular orbitals are formed against one less bonding. In all previous cases the opposite occurred. This is due to the negative overlap between adjacent 2py orbitals—whether, by convention, all positive lobes point in the clockwise direction, or whether all positive lobes point in the anticlockwise direction. The two bonding 2pv combinations in fact fall below the two antibonding (hybrid 2s, 2px) combinations. The former each have two electrons while the latter are empty. The six electrons of the three C—C bonds are nicely accounted for. The method creates simultaneously the acc and or c molecular orbitals of cyclopropane (note that the latter three lie relatively close in energy). 

This is first illustrated for the two nonbonding -type orbitals n, and n2 of para-benzyne and pyrazine (Fig. 31). These nonbonding orbitals are derived from outer (2s, 2p) sp2 type hybrids which have not been used in any bonding interaction. Although the overlap between n, and n2 is zero each one overlaps with the central CC bond orbitals. All told, there will arise two distinct molecular orbitals in which nj and n2 enter as combinations (symmetric or antisymmetric) and which have different energies, because of selective interactions with the central bonds. 

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