Topic 2.5 Lewis Diagrams: draw Lewis diagrams for molecules and polyatomic ions, applying the octet rule and accounting for valence electrons, multiple bonds, and common exceptions.

What this topic is asking

The College Board (Topic 2.5) wants you to draw Lewis diagrams for molecules and polyatomic ions: structures that show every bonding pair and lone pair of valence electrons. You apply the octet rule (most atoms want eight valence electrons around them, hydrogen wants two), use multiple bonds when needed, and recognize common exceptions. A correct Lewis diagram is the gateway to resonance, formal charge (Topic 2.6) and molecular shape (Topic 2.7).

What a Lewis diagram shows

The goal is a structure in which each atom (except hydrogen) is surrounded by eight valence electrons, the octet rule, reflecting the stability of a full outer shell. Hydrogen is the exception that wants only two (a full first shell).

The drawing procedure

For a polyatomic ion, draw the final structure in brackets with the overall charge written outside. Counting electrons correctly is the step students most often get wrong, so add for negative charge and subtract for positive before you start placing electrons.

Multiple bonds and exceptions

When there are not enough electrons to give the central atom an octet using single bonds, convert outer-atom lone pairs into double or triple bonds. Carbon dioxide ($\text{O}=\text{C}=\text{O}$) and nitrogen ($\text{N}\equiv\text{N}$) are examples.

A few species break the octet rule, and the College Board expects you to recognize them:

• Incomplete octets: boron and beryllium often have fewer than eight electrons (for example $\text{BF}_3$ has six around boron).
• Expanded octets: atoms in period 3 and beyond (such as sulfur or phosphorus) can hold more than eight electrons (for example $\text{SF}_6$).
• Odd-electron species (radicals): molecules with an odd total, such as $\text{NO}$, cannot give every atom a full octet.

A good Lewis diagram is the foundation for the rest of the unit. Once you have the bonding and lone pairs right, formal charge (Topic 2.6) tells you which of several possible structures is best, and the count of bonding regions and lone pairs on the central atom feeds straight into VSEPR (Topic 2.7) to predict the shape. Because so much downstream reasoning depends on it, it is worth checking your electron total twice, especially the charge adjustment for ions.

Try this

Q1. State the total number of valence electrons to use when drawing the Lewis diagram of the ammonium ion, $\text{NH}_4^{+}$. [1 point]

• Cue. N (5) $+$ 4 H ($4 \times 1$) $-$ 1 (for the $+$ charge) $= 8$ electrons.

Q2. Identify the octet-rule exception shown by boron trifluoride, $\text{BF}_3$. [1 point]

• Cue. An incomplete octet; boron has only six valence electrons around it.