Hydrogen orbital overlap

The bonding hydrogen orbital overlaps simultaneously with the / -orbitals on both the terminal carbon atoms and a cyclic transition structure is formed in which the C-l becomes a / -orbital and... 

In structure (a) the hydrogen orbital overlaps suprafacially with the terminal p orbitals of the n system while in structure (b) the overlap is antarafacially. Therefore the geometry of the two transition systems becomes different. While the suprafacial overlap has a plane of symmetry, the antarafacial migration has two fold axis. 

The characteristic feature of valence bond theory is that it pictures a covalent bond between two atoms in terms of an m phase overlap of a half filled orbital of one atom with a half filled orbital of the other illustrated for the case of H2 m Figure 2 3 Two hydrogen atoms each containing an electron m a Is orbital combine so that their orbitals overlap to give a new orbital associated with both of them In phase orbital overlap (con structive interference) increases the probability of finding an electron m the region between the two nuclei where it feels the attractive force of both of them... 

FIGURE 2 4 Valence bond picture of bonding in H2 as illustrated by electro static potential maps The Is orbitals of two hydrogen atoms overlap to give an or bital that contains both elec trons of an H2 molecule... 

FIGURE 2 9 Each half filled sp orbital overlaps with a half filled hydrogen Is or bital along a line between them giving a tetrahedral arrangement of four ct bonds Only the major lobe of each sp orbital is shown Each orbital contains a smaller back lobe which has been omitted for clarity... 

Section 2 6 Bonding m methane is most often described by an orbital hybridization model which is a modified form of valence bond theory Four equiva lent sp hybrid orbitals of carbon are generated by mixing the 2s 2p 2py and 2p orbitals Overlap of each half filled sp hybrid orbital with a half filled hydrogen Is orbital gives a ct bond... 

Dienes would be expected to adopt conformations in which the double bonds are coplanar, so as to permit effective orbital overlap and electron delocalization. The two alternative planar eonformations for 1,3-butadiene are referred to as s-trans and s-cis. In addition to the two planar conformations, there is a third conformation, referred to as the skew conformation, which is cisoid but not planar. Various types of studies have shown that the s-trans conformation is the most stable one for 1,3-butadiene. A small amount of one of the skew conformations is also present in equilibrium with the major conformer. The planar s-cis conformation incorporates a van der Waals repulsion between the hydrogens on C—1 and C—4. This is relieved in the skew conformation. 

Each carbon in propane is bonded to four atoms and is -hybridized. The C—C bonds are cr bonds involving overlap of a half-filled sp hybrid orbital of one carbon with a half-filled sp hybrid orbital of the other. The C—H bonds are cr bonds involving overlap of a half-filled sp hybrid oribital of carbon with a half-filled hydrogen orbital. 

The Mulliken scheme places the negative charge more or less evenly on the three carbons, and splits the positive charge among the hydrogens. Mulliken population analysis computes charges by dividing orbital overlap evenly between the two atoms involved. 

When two sp-hybridized carbon atoms approach each other, sp hybrid orbitals on each carbon overlap head-on to form a strong sp-sp a bond. In addition, the pz orbitals from each carbon form a pz-pz it bond by sideways overlap and the py orbitals overlap similarly to form a py-py tt bond. The net effect is the sharing of six electrons and formation of a carbon-carbon triple bond. The two remaining sp hybrid orbitals each form a bond with hydrogen to complete the acetylene molecule (Figure 1.16). 

FIGURE 3.8 When electrons with opposite spins (depicted as t and 1) in two hydrogen 1s-orbitals pair and thes-orbitals overlap, they lorm a boundary surface of the electron cloud. The cloud has cylindrical symmetry around the internuclear axis and spreads over both nuclei. In the illustrations in this book, cr-bonds are usually colored blue... 

Finally, note that no combination of ligand a orbitals interacts with members of the metal t2g set. The vanishing overlap between any ligand a orbital and, say, the dxy orbital is illustrated in Fig. 6-8. Overall, therefore, the metal t2g orbitals are nonbonding in this scheme. Recall how the 2p orbital of oxygen is similarly nonbonding to the hydrogen orbitals in water. 

This follows from the principle that bonds are formed only by overlap of orbitals of the same sign. Since this is a concerted reaction, the hydrogen orbital in the transition state... 

Because electrons have wave-like properties, orbital interactions involve similar addition or subtraction of wave functions. When two orbitals are superimposed, one result is a new orbital that is a composite of the originals, as shown for molecular hydrogen in Figure 10-2. This interaction is called orbital overlap, and it is the foundation of the bonding models described in this chapter. 

Consider the orbital interactions of a hydrogen atom and a fluorine atom as they combine to form a molecule of hydrogen fluoride. The electron in the hydrogen atom occupies the 1 S orbital. According to the orbital overlap... 

Hydrogen sulfide is a toxic gas with the foul odor of rotten eggs. The Lewis structure of H2 S shows two bonds and two lone pairs on the S atom. Experiments show that hydrogen sulfide has a bond angle of 92.1°. We can describe the bonding of H2 S by applying the orbital overlap model. 

A complete orbital overlap view of methane appears in Figure 10-10. Hybridization gives each carbon orbital a strongly favored direction for overlap with an atomic 1. S orbital from an approaching hydrogen atom. Four such interactions generate four localized bonds that use all the valence electrons of the five atoms involved. 

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

The opening pages of Chapter 10 introduce the principles of orbital overlap by describing molecular hydrogen as a composite of overlapping spherical 1 S orbitals. This Interaction Is the starting point for our discussion of molecular orbital theory. 

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