Polar Bonds and Their Consequences

We saw in the last chapter how covalent bonds between atoms are described, and we looked at the hybrid-orbital model used to depict most organic molecules. Before going on to a systematic study of complex organic substances, however, we still need to review a few fundamental topics. In particular, we need to look more closely at how electrons are distributed in covalent bonds and at some of the consequences that arise when the bonding electrons are not shared equally between atoms. 

The continuum in bonding from covalent to ionic as a result of unsymmetrical electron distribution. The symbol S (low/ercase Creek delta) means partial charge, either partial positive (5+) for the electron-poor atom or partial negative (S-) for the electron-rich atom. 

The bond in sodium chloride, for instance, is largely ionic. Sodium has transferred an electron to chlorine to give Na and Cl ions, which are held together in the solid by electrostatic attraction. The C-C bond in ethane, however, is fully covalent. The two bonding electrons are shared equally by the two equivalent carbon atoms, resulting in a symmetrical electron distribution in the bond. Between these two extremes lie the great majority of chemical bonds, in which the electrons are attracted somewhat more strongly by one atom than by the other. We call such bonds, in which the electron distribution is unsymmetrical, polar covalent bonds. 

Carbon, the most important element for our purposes, has an electronegativity value of 2.5. Any element more electronegative than carbon has a value greater than 2.5, and any element less electronegative than carbon has a value less than 2.5. 

A crossed arrow is often used to indicate the direction of bond polarity. By convention, electrons are displaced in the direction of the arrow. The tail of the arrow (which looks like a plus sign) is electron-poor (6-F), and the head of the arrow is electron-rich (5-). 

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