Pi orbital overlap

However, in using VSEPR, we must realize that in a double or triple bond, the sigma and pi orbital overlaps, and the electrons contained... 

Figure 20. Model of heme A with the farnesyl ethyl group in a conformation that allows pi-orbital overlap between adjacent double bonds and between the porphyrin and the C -Cg bond...

The process of electron delocalization is previously observed in many natural and synthetic molecules. For example, in benzene, electrons are delocalized over the carbon ring. In synthetic molecular materials such as polyaniline, electrons become delocalized due to pi orbital overlap between phenyl rings [28]. However, it has long been considered, as proteins are electronic insulators, and metallic-like conductivity or electron delocalization is not possible in proteins [27]. In this chapter, we discuss methods and studies that have revealed metallic-like conductivity in anode biofilms of G. sulfurreducens due to proteinaceous pili nanofilaments present in these biofilms (Fig. 7.2). 

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. 

When p orbitals overlap in a side-by-side configuration, they form a pi bond. 

Two types of bonds may be formed when orbitals overlap. These are named sigma (o) and pi (jr) bonds. 

Fig. 2 (top) Planar regioregular polymer with directly overlapping pi orbitals vs. twisted backbone with 7i orbitals out-of-plane [1]... 

The atomic orbitals overlap laterally or side-on and form a pi (ti) bond. 

Pi bond When atomic orbitals overlap side to side in such away that the resulting molecular orbital is symmetric with the bond axis in only one plane, chemists say that a 7t bond (pi bond) is formed. 

Sigma bonds form when s or p orbitals overlap in a head-on manner. Single bonds cire usually sigma bonds. Pi bonds cire usually double or triple bonds. Figure 5-9 depicts these situations. 

Two parallel p orbitals overlap side-by-side to form a pi (tt) bond. Fig. 2-3(u), or a n bond. Fig. 2-3(6). The bond axis lies in a nodal plane (plane of zero electronic density) perpendicular to the cross-sectional plane of the tt bond. 

Covalent bonds are formed when atomic orbitals overlap. The overlap of atomic orbitals is called hybridization, and the resulting atomic orbitals are called hybrid orbitals. There are two types of orbital overlap, which form sigma (cr) and pi (tt) bonds. Pi bonds never occur alone without the bonded atoms also being joined by a ct bond. Therefore, a double bond consists of a O bond and a tt bond, whereas a triple bond consists of a ct bond and two tt bonds. A sigma overlap occurs when there is one bonding interaction that results from the overlap of two s orbitals or an s orbital overlaps a p orbital or two p orbitals overlap head to head. A tt overlap occurs only when two bonding interactions result from the sideways overlap of two parallel p... 

Structures. The methyl radical is planar and has D symmetry. Probably all other carbon-centerd free radicals with alkyl or heteroatom substituents are best described as shallow pyramids, driven by the necessity to stabilize the SOMO by hybridization or to align the SOMO for more efficient pi-type overlap with adjacent bonds. The planarity of the methyl radical has been attributed to steric repulsion between the H atoms [138]. The C center may be treated as planar for the purpose of constructing orbital interaction diagrams. 

Benzene s relative lack of reactivity is a consequence of its electronic structure. As shown by the orbital picture in Figure 23.3b, each of the six carbons in benzene is sp2-hybridized and has a p orbital perpendicular to the ring. When these p orbitals overlap to form pi bonds, there are two possibilities, shown in Figure 23.3c. 

In situations like this one, the overlap of the carbon p orbital with one of the oxygen p orbitals cannot be ignored as Figure 3.15c attempts to do. Instead, all three p orbitals must be used to form MOs that involve the carbon and both oxygens. The three AOs interact to form three delocalized pi MOs. Two of these three delocalized MOs contain the four electrons the pi electrons and an unshared pair of electrons from the localized picture. Part d of Figure 3.15 shows how the orbitals overlap in the delocalized picture. 

This is an attempt to show an orbital picture for the formate anion. It corresponds to the Lewis structure in part . (The two unshared pairs of electrons on each oxygen that are not involved in resonance have been omitted for clarity.) The two red p orbitals overlap to form the pi bond. One unshared pair of electrons is in the blue p orbital on the other oxygen. This blue p orbital overlaps the red p orbital on the carbon just like the red p orbital on the other oxygen does. 

This is a computer-generated picture of the lowest energy pi bonding MO. The three p orbitals overlap without any nodes to produce this MO. It looks much like a pi bonding MO, with electron density above and below the plane of the atoms, except that it extends over three atoms, rather than two. Pictures like this, showing the lowest-energy pi MO, are provided throughout the book because they help to visualize how the p orbitals overlap to form delocalized MOs. 

In the syn-periplanar conformation, the C—H and C—L sigma bonds are coplanar and on the same side of the C—C bond. As the bonds to the hydrogen and the leaving group start to break, the hybridization at each carbon begins to change to sp2, and the two sp3 orbitals, which have some pi-type overlap initially, change to the p orbitals of the pi bond. 

Second, the ring must be planar so that the p orbitals overlap in pi fashion completely around the cycle. If the ring is not planar, the p orbitals are twisted so that they are not parallel, resulting in a decrease in overlap. This decreases or even eliminates the aromatic or antiaromatic effect of the conjugation. 

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