Source: Lecture notes, The Ohio State University
Tags: Lewis structures, electron dot model, formal charge, octet rule, octet exceptions, valence shell expansion, Kekulé structures, bonding patterns, lone pairs, radicals, Lewis acids
Difficulty: Introductory to intermediate Prerequisites: Sections 1.1 and 1.3 (functional groups, ionic and covalent bonds, electronegativity, valence electrons).
Drawing Lewis structures is one of the most fundamental skills in organic chemistry. This is where you move from "atoms bond" to "here is exactly how the electrons are arranged." The rules here, counting valence electrons, distributing them as bonds and lone pairs, checking formal charges, and knowing when the octet rule breaks down, will be used in virtually every topic for the rest of the course. If you can draw a correct Lewis structure quickly and assign formal charges without hesitation, reaction mechanisms and molecular geometry become far more approachable.
Lewis structures show how valence electrons are distributed in a molecule as bonds and lone pairs. You draw them by counting valence electrons, sketching the molecular skeleton, filling octets, and checking formal charges. The octet rule works most of the time, but there are three important exceptions: radicals (odd electron count), electron-deficient compounds (B, Be), and valence shell expansion (elements in Period 3 and beyond).
Lewis structure (electron-dot structure)
A diagram showing how valence electrons are distributed among atoms in a molecule, using dots for lone pairs and lines (or dot pairs) for bonds.
Kekulé structure
A way of representing covalent bonds as straight lines: a single line for a single bond, two lines for a double bond. Think of it as the "shorthand" version of a Lewis structure.
Lone pair
A pair of valence electrons that belongs to one atom and is not shared in a bond. Also called a non-bonding pair.
Formal charge (FC)
A bookkeeping tool that assigns a hypothetical charge to each atom in a Lewis structure. Calculated as: FC = (number of valence electrons) − (number of lone pair electrons) − (number of bonds). A Lewis structure is said to be "charge-separated" when atoms carry non-zero formal charges.
Octet rule
The principle that atoms tend to bond in ways that give them eight electrons in their outer shell (a full octet), except for hydrogen, which requires only two (a duet).
Radical
A species with an odd number of valence electrons, meaning at least one electron is unpaired. Examples include nitrogen oxide (NO) and the methyl radical (·CH₃).
Lewis acid
A species that can accept an electron pair. Electron-deficient compounds like BH₃ and BeH₂ are Lewis acids because they have fewer than eight electrons around the central atom.
Valence shell expansion
The ability of elements in Period 3 and beyond (e.g. P, S) to accommodate more than eight electrons around them, because they have accessible d orbitals. In simple terms, these larger atoms have room for extra electrons.
Draw the molecular skeleton. Place hydrogen and halogen atoms on the outside; carbon and other non-hydrogen atoms go in the interior (carbon is bonded, not terminal).
Count the total number of available valence electrons. Add up the valence electrons for every atom in the molecule. For ions, add electrons for negative charges or subtract for positive charges.
Example, CH₄: C has 4, each H has 1 → 4 + 4(1) = 8 valence electrons.
Example, HBr: H has 1, Br has 7 → 1 + 7 = 8 valence electrons.
Example, NH₄⁺: N has 5, each H has 1, subtract 1 for the positive charge → 5 + 4(1) − 1 = 8 valence electrons.
Distribute electrons as bonds and lone pairs, giving as many atoms as possible a surrounding electron octet. Hydrogen is the exception: it only requires a duet (2 electrons).
Non-bonding (lone) pairs fill remaining positions.
If an atom lacks an octet after placing single bonds and lone pairs, convert a lone pair on an adjacent atom into a shared pair (forming a double or triple bond).
Assign formal charges to each atom using the formula above, and check the result against common bonding patterns.
Atom | H | C | N | O | F |
|---|---|---|---|---|---|
Valence electrons | 1 | 4 | 5 | 6 | 7 |
Number of bonds | 1 | 4 | 3 | 2 | 1 |
Lone pairs | 0 | 0 | 1 | 2 | 3 |
These patterns follow the octet rule: each atom forms enough bonds to reach 8 electrons (or 2 for hydrogen).
Formal charge = (valence electrons) − (lone pair electrons) − (number of bonds)
Each lone pair electron counts individually (i.e. a lone pair = 2 lone pair electrons).
Example, CH₃NO₂:
FC(N) = 5 − 0 − 4 = +1
FC(O, double-bonded) = 6 − 4 − 2 = 0
FC(O, single-bonded with 3 lone pairs) = 6 − 6 − 1 = −1
Example, CO₃²⁻:
Total valence electrons: 4 + 3(6) + 2 = 24
The charge-separated Lewis structure distributes these with formal charges on certain oxygens.
A structure is "charge-separated" when it contains atoms with non-zero formal charges, even if the molecule as a whole is neutral.
Count valence electrons: 2C = 2(4) = 8; 6H = 6(1) = 6; 1O = 1(6) = 6. Total = 20.
Draw the skeleton with C atoms bonded centrally and O bridging or terminal.
Distribute electrons to satisfy octets.
Result: H₃C–C–O–H with lone pairs on oxygen. Check all formal charges are zero.
Species with an odd total number of valence electrons cannot have all electrons in pairs.
At least one electron is unpaired, making the species a radical.
Examples:
Nitrogen oxide (:N=O, odd electron)
Methyl radical (·CH₃, 7 electrons around the system)
For CH₅: FC(C) = 4 − 1 − 3 = 0. The carbon has only 7 electrons, not a full octet.
Some compounds of boron and beryllium have fewer than 8 electrons around the central atom.
BeH₂: beryllium is surrounded by only 4 electrons (two bonds, no lone pairs).
BH₃: boron is surrounded by only 6 electrons (three bonds, no lone pairs).
These are classified as Lewis acids because they can accept an electron pair to complete an octet.
Not having a full octet (Exceptions 1 and 2) makes these species unusually reactive.
Elements from Period 3 onward (P, S, Cl, etc.) can be surrounded by more than 8 electrons.
This is called valence shell expansion, sometimes referred to as an "expanded octet."
Examples:
H₃PO₄ (phosphoric acid): phosphorus has 5 bonds and is surrounded by 10 electrons.
H₂SO₄ (sulfuric acid): sulfur has 6 bonds and is surrounded by 12 electrons.
This happens because these elements have accessible d orbitals that can accommodate the extra electrons.
Single bonds are drawn as a single line; double bonds as two lines.
Examples: methane (H–C–H with 4 single bonds), ethene (C=C with a double bond), diatomic nitrogen (N≡N with a triple bond).
Formal charge formula:
FC = (number of valence electrons) − (number of lone pair electrons) − (number of bonds)
Common bonding patterns (neutral atoms):
H: 1 bond, 0 lone pairs
C: 4 bonds, 0 lone pairs
N: 3 bonds, 1 lone pair
O: 2 bonds, 2 lone pairs
F (and other halogens): 1 bond, 3 lone pairs
Lewis structures are the standard language chemists use to communicate molecular structure, from undergraduate problem sets to pharmaceutical patents. Formal charge assignments are essential for understanding reaction mechanisms: electrons flow from regions of negative formal charge toward regions of positive formal charge, which is the basis for predicting how reactions proceed. Valence shell expansion explains the bonding in everyday substances like phosphoric acid (used in soft drinks and fertilisers) and sulfuric acid (one of the most produced industrial chemicals worldwide).
Students often forget to count the charge when calculating valence electrons for ions. A cation means you subtract electrons; an anion means you add them.
A frequent error is placing hydrogen in the interior of a molecular skeleton. Hydrogen always goes on the outside, bonded to one other atom only.
Students sometimes assume that any atom with fewer than 8 electrons in a Lewis structure is "wrong." Boron and beryllium compounds are legitimate exceptions, not drawing errors.
Formal charge is not the same as oxidation state. Formal charge is a bookkeeping tool for Lewis structures; oxidation state is a different accounting system used in redox chemistry.
⚠️ Be able to draw a Lewis structure for any small molecule in under two minutes. This is a foundational skill tested throughout the course.
⚠️ Know the formal charge formula cold. Expect to be asked to calculate formal charges on specific atoms.
⚠️ Memorise the common bonding patterns for H, C, N, O, and F. If an atom deviates from its usual pattern, it likely carries a formal charge.
⚠️ Know all three octet rule exceptions and be able to give an example of each (radical, electron-deficient compound, valence shell expansion).
⚠️ Understand that charge-separated structures are valid Lewis structures, not errors.
Fill in the blank: The formal charge formula is FC = valence electrons − ______ − number of bonds.
True or false: Hydrogen can be surrounded by up to 8 electrons.
Fill in the blank: BH₃ is an example of a ______ acid because boron has fewer than 8 electrons.
True or false: Sulfur can have more than 8 electrons around it because it is in Period 3.
Fill in the blank: A species with an odd number of valence electrons is called a ______.
Answers: 1. Lone pair electrons. 2. False (hydrogen only needs 2, a duet). 3. Lewis. 4. True. 5. Radical.
Q: How do you calculate the total number of valence electrons for NH₄⁺?
A: Nitrogen contributes 5 valence electrons, each hydrogen contributes 1 (for a total of 4), and you subtract 1 for the positive charge. Total: 5 + 4 − 1 = 8 valence electrons.
Q: What is the formal charge on nitrogen in the ammonium ion (NH₄⁺)?
A: FC = 5 (valence electrons) − 0 (lone pairs) − 4 (bonds) = +1.
Q: Why is BH₃ classified as a Lewis acid?
A: Boron in BH₃ has only 6 electrons around it (three bonds, no lone pairs), making it electron-deficient. It can accept an electron pair from a donor to complete its octet, which is the defining behaviour of a Lewis acid.
Q: Give an example of valence shell expansion and explain why it occurs.
A: In H₂SO₄, sulfur forms 6 bonds and is surrounded by 12 electrons, exceeding the octet. This is possible because sulfur is in Period 3 and has accessible d orbitals that can hold additional electrons.
Q: What is the difference between a Lewis structure and a Kekulé structure?
A: A Lewis structure shows all valence electrons (bonds and lone pairs) as dots or lines. A Kekulé structure is a simplified version that represents bonds as lines but typically omits lone pairs.
Q: What are the three exceptions to the octet rule?
A: (1) Radicals, which have an odd number of electrons and at least one unpaired electron. (2) Electron-deficient compounds of early 2nd-row elements like B and Be, which have fewer than 8 electrons. (3) Valence shell expansion in elements from Period 3 onward (e.g. P, S), which can exceed 8 electrons.
Lewis structures connect directly to molecular geometry (VSEPR theory) and hybridisation, which are typically covered next in Chapter 1. Formal charge is revisited whenever you study resonance structures and reaction mechanisms. The concept of Lewis acids and bases introduced here through BH₃ and BeH₂ reappears prominently in acid-base chemistry and in reactions like hydroboration. Valence shell expansion is relevant whenever you encounter phosphorus- or sulfur-containing reagents in later organic reactions.
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