Carbon sp2 hybrid orbitals

The vinyl H2C=CH radical can be produced by cleavage of a C-H bond in ethene, and has been studied in the gas phase. The unpaired electron clearly occupies a carbon sp2 hybrid orbital, to lapse into the language of descriptive organic chemistry, but there are regions of space where the /1-spin electrons have... 

Overlap of carbon sp2 hybrid orbitals with hydrogen Is orbitals to produce the ethylene a bonded framework. (The axes have been changed for the sake of clarity—i.e. the z axis would be out of the plane of the paper.)... 

Carbon-hydrogen framework in benzene produced by overlap of carbon sp2 hybrid orbitals with hydrogen Is orbitals. All bonds are or bonds. 

Figure 1.13 An sp hybridized carbon. The three equivalent sp2 hybrid orbitals (green) lie in a plane at angles of 120° to one another, and a single unhybridized p orbital (red/blue) is perpendicular to the sp2 plane.

When we discussed sp3 hybrid orbitals in Section 1.6, we said that the four valence-shell atomic orbitals of carbon combine to form four equivalent sp3 hybrids. Imagine instead that the 2s orbital combines with only two of the three available 2p orbitals. Three sp2 hybrid orbitals result, and one 2p orbital remains unchanged- The three sp2 orbitals lie in a plane at angles of 120° to one another, with the remaining p orbital perpendicular to the sp2 plane, as shown in Figure 1.13. 

We use different hybridization schemes to describe other arrangements of electron pairs (Fig. 3.16). For example, to explain a trigonal planar electron arrangement, like that in BF, and each carbon atom in ethene, we mix one s-orbital with two /7-orbitals and so produce three sp2 hybrid orbitals ... 

An oxygen atom can also form a double bond to carbon thus in propanone (acetone), Me2C=Q , the oxygen atom could use three sp2 hybrid orbitals one to form a a bond by overlap with an sp2 orbital of the carbon atom, and the other two to accommodate the two lone pairs of electrons. This leaves an unhybridised p orbital on both oxygen and carbon, and these can overlap with each other laterally (cf. C=C, p. 9) to form a n bond ... 

Arynes present structural features of some interest. They clearly cannot be acetylenic in the usual sense as this would require enormous deformation of the benzene ring in order to accommodate the 180° bond angle required by the sp1 hybridised carbons in an alkyne (p. 9). It seems more likely that the delocalised 7i orbitals of the aromatic system are left largely untouched (aromatic stability thereby being conserved), and that the two available electrons are accommodated in the original sp2 hybrid orbitals (101) ... 

The planar framework has a bonds as just shown, which involve sp2 hybrid orbitals on the boron atoms. This leaves one unhybridized p orbital that is perpendicular to the plane. The B2H6 molecule can be considered being made by adding two H+ ions to a hypothetical B21142 ion that is isoelectronic with C2H4 because each carbon atom has one more electron than does a boron atom. In the B2I l42 ion, the two additional electrons reside in a tt bond that lies above and below the plane of the structure just shown. When two H+ ions are added, they become attached to the lobes of the n bond to produce a structure, the details of which can be shown as... 

When a carbon atom undergoes sp2 hybridization, sp2 hybrid orbitals form n bonds, but the unhybridized p orbital forms a pi bond. 

Both carbon atoms in ethylene molecule undergo sp2 hybridization and form three identical sp2 hybrid orbitals. One p orbital remains unhybridized. Two sp2 hybrid orbitals from each carbon atom overlap end to end with the Is orbital of a hydrogen atom and four C — Ho bonds are formed in total. Also, between the two carbon atoms, a C — Co bond is formed as a result of the overlap between two sp2 hybrid orbitals. So, in the C2H4 molecule in total there are five o bonds. Meanwhile, the unhybridized p orbitals of the two carbon atoms overlap side by side and form a rt bond. So between the two carbon atoms in the C2H4 molecule there is one o bond, formed by the overlapping of sp2 hybrid orbitals and one n bond, formed by the side by side overlapping of the unhybridized p orbitals. In total, two bonds are formed, hence a double bond exists between the two carbon atoms. 

Figure 10-4 shows the hybridization that occurs in ethylene, H2C=CH2. Each carbon has sp2 hybridization. On each carbon, two of the hybrid orbitals overlap with an s-orbital on a hydrogen atom to form a carbon-to-hydrogen covalent bond. The third sp2 hybrid orbital overlaps with the sp2 hybrid on the other carbon to form a carbon-to-carbon covalent bond. Note that each carbon has a remaining p-orbital that has not undergone hybridization. These are also overlapping above and below a line joining the carbons. 

The element before carbon in Period 2, boron, has one electron less than carbon, and forms many covalent compounds of type BX3 where X is a monovalent atom or group. In these, the boron uses three sp2 hybrid orbitals to form three trigonal planar bonds, like carbon in ethene, but the unhybridised 2p orbital is vacant, i.e. it contains no electrons. In the nitrogen atom (one more electron than carbon) one orbital must contain two electrons—the lone pair hence spJ hybridisation will give four tetrahedral orbitals, one containing this lone pair. Oxygen similarly hybridised will have two orbitals occupied by lone pairs, and fluorine, three. Hence the hydrides of the elements from carbon to fluorine have the structures... 

A wide variety of carbon materials, natural as well as synthetic, exist. Carbon is a polymer consisting of a hexagonal network of carbon atoms bonded to each other by sp2 hybrid orbitals the tr-bonds are parallel to the carbon network with a rc-bond perpendicular to it. The network structure of carbon is illustrated in Figure l.14... 

It will thus be apparent why the use of hybrid orbitals, e.g. sp2 hybrid orbitals in the combination of one carbon and four hydrogen atoms to form methane, results in the formation of stronger bonds. 

In simple conjugated hydrocarbons, carbon utilizes sp2 hybrid orbitals to form a-bonds and the pure px orbital to give the it-MOs. Since the c-skeleton of the hydrocarbon is perpendicular to the wave functions of it-MO, only px AOs need be considered for the formation of it-MOs of interest for photochemists. Let us consider the case of butadiene with px AO contributed by 4 carbon atoms. The possible combinations are given in Figure 2.18. The energy increases with the number of nodes so that Et < < E3 < Et. 

Each carbon of ethylene uses two of its sp2 hybrid orbitals to form hydrogen atoms, as illustrated in the first part of Figure 2.17. The remaining sp2 orbitals, one on each carbon, overlap along the intemuclear axis to give a bond connecting the two carbons. 

FIGURE 2.17 The carbon-carbon double bond in ethylene has a a component and a -it component. The ct component arises from overlap of sp2-hybridized orbitals along the internudear axis. The tt component results from a side-by-side overlap of 2p orbitals. 

The structure of ethylene and the orbital hybridization model for its double bond were presented in Section 2.20 and are briefly reviewed in Figure 5.1. Ethylene is planar, each carbon is rp2-hybridized, and the double bond is considered to have a a component and a it component. The a component arises from overlap of sp2 hybrid orbitals along a line connecting the two carbons, the tt component via a side-by-side overlap of two p orbitals. Regions of high electron density, attributed to the tt electrons, appear above and below the plane of the molecule and are clearly evident in the electrostatic potential map. Most of the reactions of ethylene and other alkenes involve these electrons. 

Section 5.2 Bonding in alkenes is described according to an sp2 orbital hybridization model. The double bond unites two sp2-hybridized carbon atoms and is made of a cr component and a tt component. The cr bond arises by overlap of an sp2 hybrid orbital on each carbon. The tt bond is weaker than the cr bond and results from a side-by-side overlap of p orbitals. 

In addition to its three sp2 hybrid orbitals, each carbon has a half-filled 2p orbital that can participate in tt bonding. Figure 11.3b shows the continuous tt system that encompasses all of the carbons that result from overlap of these 2p orbitals. The six tt electrons of benzene are delocalized over all six carbons. 

FIGURE 14.3 (a) The unshared electron pair occupies an sp2-hybridized orbital in dichlorocarbene. There are no electrons in the unhybridized p orbital. ( >) An electrostatic potential map of dichlorocarbene shows negative charge is concentrated in the region of the unshared pair, and positive charge above and below the carbon. 

The C—C single bond in vinylacetylene is a a bond generated by overlap of an sp2-hybridized orbital on one carbon with an sp-hybridized orbital on the other. Vinylacetylene has three cr bonds and three it bonds. 

One electron in each of the three hybrid orbitals on the C atom is available for bonding the fourth valence electron of each C atom occupies the unhybridized 2p-orbital, which lies perpendicular to the plane formed by the hybrids. The two carbon atoms form a cr-bond by overlap of an sp2 hybrid orbital on each atom. The H atoms form cr-bonds with the remaining lobes of the sp2 hybrids. This arrangement of orbitals leaves the electrons in the two unhybridized 2p-orbitals free to pair and form a TT-bond by side-by-side overlap. Notice that the electron density in the Tr-bond is found above and below the axis of the C—C cr-bond (Fig. 3.23). 

Doubly bonded carbons are sp2-hybridized. Carbon has three sp2 hybrid orbitals, which lie in a plane and point toward the comers of an equilateral triangle, and one unhybridized p orbital, which is oriented at a 90° angle to the plane of the sp2 hybrids. When two s/r-hybridized carbon atoms approach each other with sp2 orbitals aligned head-on for sigma bonding, the unhybridized p orbitals on each carbon overlap to form a pi bond, resulting in a net carbon-carbon double bond. 

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