Ethylene sigma bond orbitals

Bonding orbitals in ethylene (CH2=CH2) and acetylene (CH=CH). The sigma bond backbones are shown in blue. The pi bonds (one in ethylene and two in acetylene) are shown in red. Note that a pi bonding orbital consists of two lobes. 

When two p orbitals overlap in a side-by-side configuration, they form a pi bond, shown in Figure 7.7. This bond is named after the Greek letter 7t. The electron clouds in pi bonds overlap less than those in sigma bonds, and they are correspondingly weaker. Pi bonds are often found in molecules with double or triple bonds. One example is ethene, commonly known as ethylene, a simple double-bonded molecule (Figure 7.8). The two vertical p orbitals form a pi bond. The two horizontal orbitals form a sigma bond. 

In ethylene, there are two types of bonds. Sigma (tr) bonds have the overlap of the orbitals on a line between the two atoms involved in the covalent bond. In ethylene, the C-H bonds and one of the C-C bonds are sigma bonds. Pi (ir) bonds have the overlap of orbitals above and below a line through the two nuclei of the atoms involved in the bond. A double bond is always composed of one sigma and one pi bond. A carbon-to-carbon triple bond results from the... 

Rule 3 If two or three pairs of electrons form a multiple bond between two atoms, the first bond is a sigma bond formed by a hybrid orbital. The second bond is a pi bond, consisting of two lobes above and below the sigma bond, formed by two unhybridized p orbitals (see the structure of ethylene in Figure 2-17). The third bond of a triple bond is another pi bond, perpendicular to the first pi bond (shown in Figure 2-18). 

Each carbon atom has an octet, and there is a double bond between the carbon atoms. Each carbon is bonded to three other atoms (three sigma bonds), and there are no lone pairs. The carbon atoms are sp2 hybridized, and the bond angles are trigonal about 120°. The double bond is composed of a sigma bond formed by overlap of two sp2 hybridized orbitals, plus a pi bond formed by overlap of the unhybridized p orbitals remaining on the carbon atoms. Because the pi bond requires parallel alignment of its two p orbitals, the ethylene molecule must be planar (Figure 2-17). 

The double bond between carbon and oxygen looks just like the double bond in ethylene. There is a sigma bond formed by overlap of sp2 hybrid orbitals and a pi bond formed by overlap of the unhybridizedp orbitals on carbon and oxygen (Figure 2-19). 

Each of the carbon-hydrogen sigma bonds is formed by overlap of an sp2 hybrid orbital on carbon with the Is orbital of a hydrogen atom. The C—H bond length in ethylene (1.08 A) is slightly shorter than the C—H bond in ethane (1.09 A) because the sp2 orbital in ethylene has more s character (one-third, v) than an sp3 orbital (one-fourth, v). The s orbital is closer to the nucleus than the p orbital, contributing to shorter bonds. 

These three views of the ethylene molecule emphasize different aspects of the disposition of shared electron pairs in the various bonding orbitals of ethene (ethylene), (a) The backbone structure consisting of sigma (a) bonds formed from the three sp2-hybridized orbitals on each carbon, (b) The % (pi) bonding system formed by overlap of the unhybridized pz orbital on each carbon. The pi orbital has two regions of electron density extending above and below the plane of the molecule, (c) A cutaway view of the combined sigma and pi system. 

Optical purity, by NMR, 13, 14 Orbital correlation diagrams, 196-203 cycloaddition reactions, 197-196 Diels-Alder, 198 ethylene -E ethylene, 198 electrocyclic reactions, 198-200 butadienes, 199 hexatrienes, 199 limitations, 203 photochemical, 201 Woodward-Hoffinann, 197 Orbital energies, see also Energies, orbital degeneracy, 27, 90 Orbital interaction theory, 34-71 diagram, 40, 42, 47 limitations, 69-71 sigma bonds, 72-86 Orbitals... 

Figure 2.3 The bonding in ethylene is explained by combining two sp hybridized carbon atoms. The C—H sigma bonds all lie in the same plane. The unhybridized p orbitals of the carbon atoms overlap above and below the molecular plane to form the pi bond.

Not all bonds allow free rotation ethylene, for example is quite rigid. In ethylene, the double bond between the two CH2 groups consists of a sigma bond and a pi bond. When we twist one of the two CH2 groups, the sigma bond is unaffected but the pi bond loses its overlap. The two p orbitals cannot overlap when the two ends of the molecule are at right angles, and the pi bond is effectively broken. 

In a similar manner, VBT can be used to explain why ethylene is a planar molecule. As in O2, the two C atoms are sp hybridized and lie in the xy-plane—only in H2C=CH2, each sp hybrid contains a single electron. One of these hybrids is used to form the C-C sigma bond, while the other two hybrids are used to form the C-H sigma bonds. The C-C pi bond forms as a result of the sideways overlap between the two p orbitals. In order for the two p orbitals to overlap with each other, there can be no free rotation around the C=C double bond. Therefore, both CH2 halves of the ethylene molecule must lie in the same plane, as shown in Figure 10.10. 

According to the orbital overlap model, a carbon-carbon double bond consists of one sigma bond formed by the overlap of sp hybrid orbitals and one pi bond formed by the overlap of parallel 2p atomic orbitals. It takes approximately 264 kJ/mol (63 kcal/mol) to break the pi bond in ethylene. 

Figure 4.14 Bonding in ethylene, C2H4. (a) Top view of the sigma bonds between carbon atoms and between carbon and hydrogen atoms. All of the atoms lie in the same plane, making C2H4 a linear molecule, (b) Side view showing how the two 2p orbitals on the two carbon atoms overlap, leading to the formation of a pi bond. Blue and red colors denote + and -regions of the orbitals, (c) The interactions in (a) and (b) lead to the formation of the sigma bonds and pi bond in ethylene. Note that the pi bond lies above and below the plane of the molecule.

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