COVALENT BONDING AND ORBITAL OVERLAP

Because the octahedral molecular geometry is symmetrical, the bond dipoles cancel, and the molecule is nonpolar, meaning that fi = 0. 

In Lewis theory, covalent bonding occurs when atoms share electrons because the sharing concentrates electron density between the nuclei. In valence-bond theory, we visualize the buildup of electron density between two nuclei as occurring when a valence atomic orbital of one atom shares space, or overlaps, with a valence atomic orbital of another atom. The overlap of orbitals allows two electrons of opposite spin to share the space between the nuclei, forming a covalent bond. 

The coming together of two H atoms to form H2 is depicted in FIGURE 9.13. Each atom has a single electron in a Is orbital. As the orbitals overlap, electron density is concentrated between the nuclei. Because the electrons in the overlap region are simultaneously attracted to both nuclei, they hold the atoms together, forming a covalent bond. 

A FIGURE 9.13 Covalent bonds in H2, HCl, and CI2 result from overlap of atomic orbitals. 

On the left part of the curve the potential energy rises above zero. What causes this to happen  

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COVALENT BONDiNG AND ORBiTAL OVERLAP We recognize that eiectrons are shared between atoms in a covaient bond, in valence-bond theory, the bonding eiectrons are visuaiized as originating in atomic orbitais on two atoms. A covaient bond is formed when these orbitais overiap. 

Modes of Orbital Overlap and the Types of Covalent Bonds 335 Orbital Overlap in Single and Multiple Bonds 335... 

Figure 2.8. The iwo orbitals overlap giving a covalent bond and ihe tvvv electrons are (>ir in a molecular orbital. (If the t o nuclei could be pushed together completely, the...

The concepts of directed valence and orbital hybridization were developed by Linus Pauling soon after the description of the hydrogen molecule by the valence bond theory. These concepts were applied to an issue of specific concern to organic chemistry, the tetrahedral orientation of the bonds to tetracoordinate carbon. Pauling reasoned that because covalent bonds require mutual overlap of orbitals, stronger bonds would result from better overlap. Orbitals that possess directional properties, such as p orbitals, should therefore be more effective than spherically symmetric 5 orbitals. 

However, in sulphides and related minerals, the effects of covalent bonding predominate and orbital overlap must be taken into account. Thus, concepts of molecular orbital theory are described in chapter 11 and applied to aspects of the sulfide mineralogy of transition elements. Examples of computed energy diagrams for molecular clusters are also presented in chapter 11. There, it is noted that the fundamental 3d orbital energy splitting parameter of crystal field theory, A, receives a similar interpretation in the molecular orbital theory. 

The coverage of covalent bonding and overlap of orbitals is a study in its own right. To go further with this topic, you might like to read a more advanced text and consider the topic of hybridisation . 

In the previous paper (13), we have discussed the chemical bonding nature of uranyl nitrate dihydrate and found that the bonding interaction is mainly due to the U 5f, 6d - O 2p components. In the present work, we carry out orbital overlap population analysis to understand contribution of each atomic orbital to the chemical bonding. The orbital overlap populations indicate strength of covalent bonds (19,20). 

Oxygen has an electronic structure ls22s22p22py2p. Now a seemingly satisfactory picture of the oxygen molecule could be obtained by considering the formation of two covalent bonds, by the overlap of the partly filled 2py and 2pz orbitals respectively, from two oxygen atoms. 

Orbitals overlap directly in sigma bonds. Parallel orbitals overlap in pi bonds. A single covalent bond is a sigma bond but multiple covalent bonds are made of both sigma and pi bonds. 

In the initial coordination, the 7r-electron cloud of the olefin overlaps with the outer bond orbital of the metal cation. This causes stretching and eventual rupture of the R-M (metal) bond. An intramolecular rearrangement follows with the migration of the caibanion (R to the most electron-deficient carbon atom of the double bond. A new covalent bond and a new caibanion are formed simultaneously ... 

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