In this SparkNote we introduce several tools to describe covalent bonds,
which are formed when atoms share electrons in a mutual effort to attain full
valence shells. The most important representation that organic chemists use to
depict covalently bonded molecules is the Lewis structure. Lewis structures
depict the valence electrons of all atoms in the molecule, as either bonded
electron pairs or lone pairs. Atoms can be held together not just by
single bonds but by double and triple bonds; this bond order affects the bond's
strength and length.
Because the atoms a bond holds together can have different
electronegativities, some covalent bonds are shared unequally. Such
polar
covalent bonds have a partial positive and a partial negative end,
giving the
bond a dipole. We can estimate the charge of an entire atom by adding the
separate charges of the dipoles.
We provide a systematic method for writing Lewis structures. An
interesting result is that there is often more than one valid way to place
electrons on a given atomic framework. The different Lewis structures are
resonance structures. They represent an actual molecule which is the
resonance hybrid of those structures. The resonance hybrid is a weighted
average of its contributors, with more stable contributors giving greater
weight.
