Difficulty: Introductory to Intermediate | Prerequisites: Periodic table familiarity, valence electrons by group, the octet rule.
Lewis structures are one of the most fundamental tools in chemistry for representing how atoms bond and where lone pairs sit. This recitation covers the standard five-step method and applies it to fourteen molecules ranging from simple diatomics (HF, N2) to multi-carbon organic fragments (CH3CHO, CH3CCH). You will need to be comfortable counting valence electrons by group number before any of this makes sense. If you are shaky on which group gives how many valence electrons, sort that out first.
Lewis structures show how valence electrons are distributed in a molecule as bonding pairs and lone pairs. Follow a five-step method: count total valence electrons, draw the skeleton, place bonding pairs, complete octets with lone pairs, and form multiple bonds if electrons run short. The molecular formula itself often hints at connectivity.
Lewis structure (Lewis dot structure)
A diagram showing the arrangement of valence electrons in a molecule, with lines for bonding pairs and dots for lone pairs. Think of it as a map of where every valence electron lives.
Valence electrons
The electrons in the outermost shell of an atom, which participate in bonding. For main-group elements, the number of valence electrons equals the group number (using the standard US numbering: Group 1 = 1, Group 14 = 4, Group 17 = 7).
Bonding pair
A pair of electrons shared between two atoms, represented as a line in a Lewis structure. Each single bond uses 2 electrons.
Lone pair
A pair of valence electrons that belong to one atom and are not shared in a bond. Shown as two dots on the atom.
Octet rule
The observation that main-group atoms (except H and B in some cases) tend to form bonds until they are surrounded by 8 electrons (shared and unshared combined). Hydrogen is satisfied with 2.
Multiple bond
A double bond (4 shared electrons, two lines) or triple bond (6 shared electrons, three lines), formed when there are not enough electrons to give every atom an octet using single bonds alone.
Skeleton structure
The arrangement of atoms before electrons are placed. The least electronegative atom usually goes in the centre. Hydrogen is always terminal (it can only form one bond).
Structural isomers
Molecules with the same molecular formula but different connectivity of atoms. C2H6O can be ethanol (C-C-O-H) or dimethyl ether (C-O-C). Lewis structures help you see these.
Cis/trans isomers
For molecules with a double bond (which prevents rotation), substituents can be on the same side (cis) or opposite sides (trans) of the bond. C2H2Cl2 has both cis and trans forms.
Electron-deficient molecule
A molecule where an atom has fewer than 8 electrons in its Lewis structure. BH3 is the classic example: boron has only 6 electrons around it.
Count total valence electrons. Add up the valence electrons for every atom in the molecule. For ions, add electrons for negative charges and subtract for positive charges.
Write the skeleton structure. Place the least electronegative atom in the centre. Hydrogen is always on the outside (it can only form one bond). The molecular formula often hints at connectivity: CH3Cl means C is central with three H and one Cl attached.
Place bonding pairs. Use 2 electrons (one line) for each bond in the skeleton.
Complete octets with lone pairs. Distribute remaining electrons as lone pairs on outer atoms first, then the central atom. Remember: H needs only 2 electrons, not 8.
Form multiple bonds if needed. If the central atom still lacks an octet after step 4, convert lone pairs from adjacent atoms into bonding pairs (double or triple bonds).
a. HF, total valence electrons = 8
H contributes 1, F contributes 7. Single bond between H and F, with three lone pairs on F. Straightforward application of the method.
b. CH3Cl , total valence electrons = 14
C (4) + 3 H (3) + Cl (7) = 14. Carbon is central with three H atoms and one Cl. All single bonds. Cl carries three lone pairs.
c. H2S, total valence electrons = 8
2 H (2) + S (6) = 8. S is central with two H atoms. Two bonding pairs and two lone pairs on S. Similar geometry to water.
d. N2, total valence electrons = 10
2 N (10). A triple bond between the two nitrogen atoms, with one lone pair on each N. This is the classic example of step 5: single and double bonds leave nitrogen short of an octet, so you need a triple bond.
e. BH3, total valence electrons = 6
B (3) + 3 H (3) = 6. Three single bonds to hydrogen, all 6 electrons used. Boron has only 6 electrons around it. This is an electron-deficient molecule; boron is an exception to the octet rule.
f. C2H4 (ethylene), total valence electrons = 12
2 C (8) + 4 H (4) = 12. Each carbon bonds to two H atoms and to the other carbon with a double bond. No lone pairs. This is the simplest alkene.
g. C2H6O, total valence electrons = 20
This formula has two possible structural isomers:
Ethanol: H3C-CH2-O-H. The oxygen sits at the end of the chain with one lone pair visible and one H. Oxygen carries two lone pairs.
Dimethyl ether: H3C-O-CH3. The oxygen bridges two methyl groups. Oxygen carries two lone pairs.
Both use all 20 electrons. This is a key example of structural isomers from the same formula.
h. CH3CH2CH3 (propane), total valence electrons = 20
3 C (12) + 8 H (8) = 20. A three-carbon chain with all single bonds and no lone pairs. The formula itself spells out the connectivity: CH3-CH2-CH3.
i. CH3CHCH2 (propene), total valence electrons = 18
3 C (12) + 6 H (6) = 18. The middle carbon (CH) and the end carbon (CH2) are connected by a double bond. No lone pairs on carbon. The missing 2 electrons compared to propane go into the double bond.
j. PCl3, total valence electrons = 26
P (5) + 3 Cl (21) = 26. Phosphorus is central with three Cl atoms attached by single bonds. Each Cl has three lone pairs. P has one lone pair. Similar shape to NH3.
k. C2H2Cl2, total valence electrons = 24
This molecule has cis and trans isomers because of the C=C double bond:
Trans isomer: the two Cl atoms are on opposite sides of the double bond.
Cis isomer: the two Cl atoms are on the same side of the double bond.
Each Cl has three lone pairs. The double bond prevents rotation, locking the geometry.
l. C2H4Cl2, total valence electrons = 26
No double bond here (all single bonds between the two carbons). Two structural isomers:
1,1-dichloroethane: both Cl atoms on the same carbon.
1,2-dichloroethane: one Cl on each carbon.
Each Cl carries three lone pairs.
m. CH3CHO (acetaldehyde), total valence electrons = 18
The "CHO" end is the aldehyde group: carbon double-bonded to oxygen, with one H on that carbon. Oxygen has two lone pairs. The CH3 group is all single bonds.
n. CH3CCH (propyne), total valence electrons = 16
3 C (12) + 4 H (4) = 16. The two rightmost carbons are connected by a triple bond (C is triple-bonded to CH). No lone pairs on carbon. This is the simplest three-carbon alkyne.
Students often forget to count all valence electrons before starting. If your electron count is wrong, the entire structure will be wrong. Double-check the count before drawing anything.
A common mistake is placing hydrogen in the centre of a molecule. Hydrogen can only form one bond, so it is always terminal.
Students sometimes give every atom lone pairs without checking whether the total electron count still adds up. Every lone pair and bonding pair must be accounted for in the total.
Many students forget step 5: if an atom (usually C, N, or O) does not have an octet after adding lone pairs, you need to form a double or triple bond by converting a lone pair into a bonding pair.
With structural isomers (like C2H6O), students often draw only one structure and assume they are finished. The molecular formula alone does not tell you the connectivity; always check whether more than one valid arrangement exists.
Students confuse cis/trans isomerism with structural isomerism. Cis/trans isomers have the same connectivity but differ in spatial arrangement around a double bond. Structural isomers have different connectivity entirely.
Counting valence electrons is tested directly. Expect a question like "How many valence electrons are in PCl3?" or "Draw the Lewis structure and state the total valence electron count."
Structural isomers are a favourite exam topic. If a molecular formula can produce more than one valid Lewis structure (different connectivity), you may be asked to draw all of them.
Cis/trans isomers around double bonds are testable. Know that a C=C double bond prevents rotation, locking substituents in place.
BH3 is a common exam example of an octet rule exception. Be prepared to explain why boron has only 6 electrons and why that is acceptable.
The five-step method itself may appear as a short-answer question: "Describe the steps for drawing a Lewis structure."
Fill in the blank: A molecule with 5 valence electrons on its central atom and three bonds to hydrogen, with no lone pairs and only 6 electrons around the central atom, is ______ . (BH3, an electron-deficient molecule)
True or false: Hydrogen can be the central atom in a Lewis structure. (False: hydrogen forms only one bond and must be terminal.)
Fill in the blank: C2H6O has ______ total valence electrons and can form ______ structural isomers. (20; two: ethanol and dimethyl ether)
True or false: If you run out of electrons before every atom has an octet, you should form a multiple bond. (True)
Fill in the blank: N2 has a ______ bond and ______ lone pair(s) on each nitrogen. (triple; one)
Q: How many total valence electrons are in CH3CHO, and what type of bond exists between carbon and oxygen?
A: 18 total valence electrons. The carbon-oxygen bond is a double bond (C=O). Oxygen carries two lone pairs.
Q: Draw all valid Lewis structures for C2H6O. How do they differ?
A: Two structures are possible. Ethanol (CH3CH2OH) has the oxygen bonded to a carbon and a hydrogen at the end of the chain. Dimethyl ether (CH3OCH3) has the oxygen bridging two methyl groups. Both have 20 valence electrons, but the atom connectivity differs.
Q: Why does BH3 violate the octet rule, and is that acceptable?
A: Boron has only 3 valence electrons. With three B-H single bonds, all 6 electrons in the molecule are used and boron has only 6 electrons around it. This is acceptable because boron is a known exception to the octet rule; it can be stable with fewer than 8 electrons.
Q: C2H2Cl2 can exist as cis and trans isomers. Explain why.
A: The two carbons are connected by a double bond, which prevents rotation. With one H and one Cl on each carbon, the Cl atoms can be on the same side (cis) or opposite sides (trans) of the double bond, giving two distinct spatial arrangements.
Q: A student draws a Lewis structure for N2 with a single bond and three lone pairs on each nitrogen. What is wrong?
A: Count the electrons: a single bond uses 2, and six lone pairs use 12, totalling 14. N2 has only 10 valence electrons. The correct structure is a triple bond (6 electrons) plus one lone pair on each nitrogen (4 electrons), totalling 10.
Lewis structures are the starting point for VSEPR theory (predicting molecular geometry), which comes next in most general chemistry courses. They also feed directly into discussions of polarity and dipole moments: you need to know where the lone pairs and bonds are before you can assess whether a molecule is polar. In organic chemistry, Lewis structures evolve into line-angle (skeletal) formulas, but the electron-counting logic stays the same.
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