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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.