Overlap sideways

Sometimes each of two atoms has a partly filled p-orbital, and instead of overlapping end-to-end to form a sigma bond they overlap sideways in a less effective manner, to form what is called a m (tt) bond (Figure 2.2.1.3). 

The two sides of alkenes are due to the nature of the four-electron double bond. In the simplest picotre, we assign two of its electrons to a basic, garden-variety CT bond between the atoms. The other two electrons are then placed in two parallel p orbitals, overlapping sideways to fomi the ir bond. This tr-type overlap prevents the carbons at each end of the double bond from rotating with respect to one another. Ethene, therefore, is a perfectly flat molecule, and, in general, the carbons of the alkene functional group and all the atoms attached to them will lie in a plane, with the n electrons above and below. 

The p orbitals overlap sideways with their neighbours to form tt clouds or doughnuts of electron density, above and below the... 

Unhybridised 2p arbitals on each carbon overlap sideways to produce a k bond. 

The other 2p orbitals overlap sideways and thus concentrate electron density above and below the internuclear axis. MOs of this type are called pi ( -) molecular oibitals by... 

The remaining two valence electrons reside in the unhybridized 2p orbitals, one electron on each carbon. These two orbitals can overlap sideways with each other, as shown in Figure 9.23. The resultant electron density is concentrated above and below the C—C bond axis It is a tt bond (Figure 9.21). Thus, the C=C double bond in ethylene consists of one a bond and one tt bond. 

A pi (tt) bond, part of double and triple carbon-carbon bonds, is one in which the electron densities of two adjacent parallelp orbitals overlap sideways to form a bonding pi molecular orbital. 

Ethene has one double bond (H2C=CH2) and each carbon atom is coimected via three sp orbitals two spin pair with Ij hydrogen electrons and the third is shared with the other carbon atom, giving a total of five ct bonds in one plane. Each carbon atom has, at right angles to this plane, two further p orbitals, each with one electron and these orbitals take no part in the a hybridization. The p orbitals can overlap sideways and form a n orbital in the form of two dogbone or sausage shaped sections, one above the ct orbital plane and the other below (Figure 1.7) [5], where the plane represents a node for the n orbital. Maximum overlap of the p orbitals prevents rotation about the double bond. 

How, then, do we account for the bonding of the C atoms As Figure 4.14(a) shows, each carbon atom uses the three sp hybrid orbitals to form two sigma (o-) bonds with the two hydrogen Is orbitals and one tr bond with the sp hybrid orbital of the adjacent C atom. In addition, the two unhybridized 2p orbitals of the C atoms form a pi (it) bond by overlapping sideways [Figure 4.14(b)]. 

PROBLEM 2.8 Why not allow a 2p orbital to overlap sideways with an s orbital Why would this not solve the problem of waste of the rear lobe (see below and Fig. 2.10) ... 

WThe formation of a cr bond by the overlap of two s orbitals. A cr bond can also be formed by the overlap of two p orbitals along their axes, fc) When twop orbitals overlap sideways, a tt bond is formed. 

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