Covalent Bonding and Lewis Structures
Draw Lewis structures for molecules and polyatomic ions, including single, double, and triple bonds, lone pairs, formal charge, and resonance at an introductory level.
A lesson of Chemistry, in Chemical Bonding and Molecular Structure, about 35 minutes.
By the end of this lesson, you will be able to draw a Lewis structure for a simple molecule, count bonding and lone pairs, predict the shape from them, and say why that shape decides how the substance behaves.
Atoms Share When Neither Will Give
An ionic bond is a transfer: one atom hands an electron over and both end up charged. That works when one atom wants electrons badly and the other barely holds on to its own.
Two atoms that both want electrons cannot do that. Neither will give one up. So they share.
$$\text{H} \cdot + \cdot \text{H} \longrightarrow \text{H} : \text{H}$$
That shared pair is a covalent bond. Both nuclei are attracted to the same two electrons, and that attraction is what holds the molecule together.
The reason atoms bother is the octet rule: an atom is stable with eight electrons in its outer shell, or two for hydrogen, and sharing is a way to reach that count without anyone losing an electron outright.
Drawing a Lewis Structure
A Lewis structure is a map of where every outer electron is. Four steps, in this order, every time.
- Count the valence electrons of every atom, and add them up
- Put the least electronegative atom in the middle, except hydrogen, which is never central because it can only make one bond
- Join the outer atoms to the center with one pair each
- Spread what is left as lone pairs, outer atoms first, until everyone has an octet
If the center finishes short of eight, pull a lone pair in from an outer atom and make a double or triple bond.
Water, worked through. Oxygen brings 6 valence electrons, each hydrogen brings 1, so 8 in total. Oxygen sits in the middle, two bonds use 4 electrons, and the remaining 4 sit on oxygen as two lone pairs.
H : O : H, with two lone pairs on the oxygen
Two bonding pairs, two lone pairs, and every atom satisfied.
Worked Example 1: Counting Before Drawing
Ammonia, $\text{NH}_3$.
Start with 4 bonded atoms and no lone pairs, then take one bond away and add a lone pair. Watch what happens to the angle.
A Lewis structure is flat on the page, and molecules are not. Bonding pairs and lone pairs all repel each other, so they spread out as far apart as they can get in three dimensions. That is the whole of VSEPR theory.
<!-- SLATE_IF_STRONG: science.chemistry.covalent-bonding --> <!-- END_SLATE_IF_STRONG -->