= Polar covalent bond
A <polar covalent bond> shares <electrons> unequally, as in <hydrogen chloride>.
The word “polar” qualifies the sharing within a <covalent bond>; it does not require complete transfer of an <electron> or an isolated <cation>–<anion> pair. In $A^{\delta+}-B^{\delta-}$, the signs identify the relatively electron-poor and electron-rich ends, while $\delta$ leaves the amount unspecified. A neutral <Lewis structure> can give both <chemical atoms> zero <formal charge> and still correspond to a <polar covalent bond>: equal division of the <bond pair> is a formal allocation, whereas <bond polarization> concerns the actual <electron density>.
Local and molecular descriptions answer different questions. Calling one <chemical bond> a <polar covalent bond> does not establish that the whole <molecule> is a <polar molecule>. Combine the <bond dipoles> with the <molecular geometry>, then check for <molecular dipole cancellation>. This lets a short <electronegativity> comparison explain a local tendency without confusing it with the observable <electric dipole moment> of the whole <molecule>.
Back to article page
