The result of this unequal sharing is what we call a bond dipole, which exists in a polar covalent bond. The same is true for the oxygen-hydrogen bond, as hydrogen is slightly less electronegative than carbon, and much less electronegative than oxygen. In the carbon-oxygen bond of an alcohol, for example, the two electrons in the sigma bond are held more closely to the oxygen than they are to the carbon, because oxygen is significantly more electronegative than carbon. However, quite often in organic chemistry we deal with covalent bonds between two atoms with different electronegativities, and in these cases the sharing of electrons is not equal: the more electronegative nucleus pulls the two electrons closer. Recall from your general chemistry course that electronegativity refers to “ the power of an atom in a molecule to attract electrons to itself” (this is the definition offered by Linus Pauling, the eminent 20 th-century American chemist who was primarily responsible for developing many of the bonding concepts that we have been learning). In these examples, the two atoms have approximately the same electronegativity. Many of the covalent bonds that we have seen – between two carbons, for example, or between a carbon and a hydrogen –involve the approximately equal sharing of electrons between the two atoms in the bond. To understand the nature of noncovalent interactions, we first must return to covalent bonds and delve into the subject of dipoles.
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