Lewis DiagramsLewis 结构

One diagram, a lot of information

Lesson 2.4.4 showed how to draw Lewis structures. AP topic 2.5 asks you to draw them quickly and correctly, and also to read information off them: how many electrons surround each atom, whether the central atom has lone pairs, and whether each bond is single or multiple. Resonance and formal charge (U2.6), molecular shape and hybridization (U2.7), and intermolecular forces (Unit 3) all start from a correct Lewis diagram.

Drawing Lewis diagrams: a quick review

1. Count the valence electrons. Add one electron for each negative charge on an anion; subtract one for each positive charge on a cation.
2. Sketch the skeleton. The less electronegative atom usually goes in the middle; H can bond to only one atom, so it is always on the outside.
3. Connect with single bonds. Each single bond uses 2 electrons.
4. Complete the outer atoms. Use the remaining electrons to give each outer atom 8 (H needs only 2), then put any extra electrons on the central atom.
5. If the central atom is short of 8, make multiple bonds. Move a lone pair from an outer atom into the space between the two atoms, turning a single bond into a double bond, and a triple bond if needed.

Lone pairs on the central atom

O and S are both in group 16 with 6 valence electrons. In $\ce{H2O}$ and $\ce{H2S}$, the central atom uses 2 of them to bond to the two H atoms, and the remaining 4 form 2 lone pairs孤电子对 on the central atom. Elements in the same group form molecules with matching Lewis diagrams: the central atoms of $\ce{NH3}$ and $\ce{PH3}$ each have 1 lone pair, and those of $\ce{CH4}$ and $\ce{SiH4}$ have none.

A Lewis diagram is a flat sketch: H–O–H is drawn in a straight line, but that doesn’t mean the water molecule is straight. The real shape comes from the VSEPR model (U2.7).

Exceptions to the octet rule

• Electron-deficient atoms: the B in $\ce{BF3}$ and $\ce{BH3}$ has only 6 electrons around it. Why not move a lone pair from F to make a B=F double bond? That would give F a positive formal charge (U2.6), which makes no sense for the most electronegative element. So B keeps only 3 electron pairs.
• Expanded octets: central atoms from period 3 and beyond (P, S, Cl, Xe, and so on) are larger and can have more than 8 electrons, as in $\ce{PCl5}$ (10) and $\ce{SF6}$ (12). The period 2 elements C, N, O, and F never exceed 8.
• Odd numbers of electrons: $\ce{NO2}$ has 17 valence electrons, so some atom always ends up short of 8. This is a limitation of the Lewis model.

An electron-deficient B meets a lone pair

The B in $\ce{BF3}$ is still one electron pair short, and the N in $\ce{NH3}$ has a lone pair. When the two molecules meet, the lone pair on N is shared with B, forming a new N–B covalent bond:

Both electrons in this bond come from N (a coordinate covalent bond). After the reaction, B also has 8 electrons around it. The +1 and −1 on the diagram are formal charges, covered in detail in U2.6.

Reading a Lewis diagram: hydrocarbons

From these diagrams alone: the more electrons the two C atoms share, the shorter and stronger the C–C bond (U2.2). The three empirical formulas all differ (U1.3), while $\ce{C2H2}$ and benzene, $\ce{C6H6}$, share the empirical formula CH. U2.7 uses the same diagrams to predict shapes: all six atoms of $\ce{C2H4}$ lie in one plane, the four atoms of $\ce{C2H2}$ lie on a straight line, and $\ce{C2H6}$ is not planar.

Isomers

Compounds with the same molecular formula but different connections between atoms are called isomers同分异构体. Ethanol and dimethyl ether are both $\ce{C2H6O}$: in ethanol, O is bonded to one C and one H (an O–H bond), while in dimethyl ether, O is bonded to two C atoms. Different structures give different properties: at room temperature ethanol is a liquid that mixes with water in any proportion, while dimethyl ether is a gas.

Propanal and acetone are both $\ce{C3H6O}$. The difference is where the C=O sits: at the end of the chain in propanal (that C also carries an H), and in the middle in acetone.

To decide whether two diagrams show isomers, first count each kind of atom to confirm the molecular formula is the same, then check whether the atoms are connected differently.

Now draw your own: choose a molecule and click a bond to switch it among single, double, and triple. Lone pairs are filled in automatically by the steps above, and atoms still short of an octet turn orange:

Example 1: The Lewis diagram of PCl₃

Draw the Lewis diagram of $\ce{PCl3}$ and state how many lone pairs are on the central atom.

Example 2: Comparing C₂H₄ and C₂H₂

Using their Lewis diagrams, compare the length and strength of the C–C bonds in $\ce{C2H4}$ and $\ce{C2H2}$, and give the empirical formula of each.

Common mistakes

Quick check

AP-style free response

Summary

• Drawing a Lewis diagram: count valence electrons (adjusting for charge), sketch the skeleton, connect with single bonds, complete the outer atoms and then the central atom, and form multiple bonds if needed.
• Show the lone pairs on the central atom: 2 each in $\ce{H2O}$ and $\ce{H2S}$, 1 in $\ce{NH3}$.
• Exceptions: B can have only 6 electrons; central atoms from period 3 on can exceed 8; molecules with an odd number of electrons always leave some atom short.
• A lone pair can form a new covalent bond with an electron-deficient atom, as in $\ce{H3N-BF3}$.
• A Lewis diagram shows bond types, relative bond lengths and energies, and the empirical formula; compounds with the same molecular formula but different connections are isomers.