CHEM1011 · Chemistry 1a
Chemical Bonding: Lewis Structures and Formal Charge
Week 3 opens bonding with the electrostatic picture and the vocabulary — bond energy and length, lone versus bonding pairs, sigma and pi bonds, bond order — then drills Lewis structures for inorganic ions and organic functional groups. Assigning formal charges to choose the most realistic structure, and recognising octet exceptions, are reliable in-term-test and exam items where a careful electron count is the whole battle.
What this chapter covers
- 01The electrostatic model: attraction vs repulsion and the potential-energy curve (bond length at the minimum, bond energy = well depth)
- 02Vocabulary: bond energy, bond length, lone pair, bonding pair, valence shell, core electrons, sigma (σ) and pi (π) bonds, bond order
- 03Lewis-structure procedure: count valence electrons (adjust for charge), least electronegative atom central, complete octets, add multiple bonds
- 04Formal charge FC = (valence e⁻) − (lone-pair e⁻) − ½(bonding e⁻); the best structure minimises FC and puts negative FC on the more electronegative atom
- 05Octet exceptions: electron-deficient (BF₃), expanded octet (PCl₅, SF₆), odd-electron species
- 06Bonding in organic functional groups: alkanes, alkenes, alkynes, alcohols, carboxylic acids, esters, ketones, aldehydes, amines, amides
- 07Resonance and delocalisation (e.g. CO₃²⁻, NO₃⁻); real bond orders are intermediate (Mastery)
- 08Electron domains: a single, double or triple bond each count as one domain (feeds VSEPR)
Choosing the best Lewis structure by formal charge
- +1State the rule: the best structure minimises the magnitude of formal charges and places any negative formal charge on the more electronegative atom (here oxygen, χ_O > χ_N > χ_C).
- +1Structure A, N≡C–O (N has 1 lone pair, O has 3 lone pairs). N: 5 − 2 − ½(6) = 0. C: 4 − 0 − ½(8) = 0. O: 6 − 6 − ½(2) = −1. Charges sum to −1 ✓.
- +1Structure C, O=C=N (O has 2 lone pairs, N has 2 lone pairs). O: 6 − 4 − ½(4) = 0. C: 4 − 0 − ½(8) = 0. N: 5 − 4 − ½(4) = −1. Charges sum to −1 ✓.
- +1Both keep formal charges small (a single −1), so the tie-breaker is electronegativity: structure A puts the −1 on oxygen (more electronegative), structure C puts it on nitrogen. Structure A is the better Lewis structure.
Key terms
- Lewis structure
- A dot-and-line diagram showing bonding and lone-pair electrons. Built by counting valence electrons (adjusting for charge), connecting atoms with the least electronegative one central, completing octets, then adding multiple bonds if needed.
- Formal charge (FC)
- FC = (valence electrons) − (lone-pair electrons) − ½(bonding electrons). The preferred structure minimises formal-charge magnitudes and puts negative FC on the more electronegative atom; charges must sum to the species' overall charge.
- Sigma (σ) and pi (π) bonds
- A σ bond is head-on overlap with density along the internuclear axis (every single bond, and one bond of any multiple bond); a π bond is side-on p-orbital overlap above and below the axis (the extra bonds in double/triple bonds).
- Bond order
- The number of bonding electron pairs between two atoms; higher bond order gives a shorter, stronger bond. Resonance can give fractional bond orders (CO₃²⁻ is 1⅓).
- Octet exception
- A stable structure that breaks the octet rule: electron-deficient species (BF₃, fewer than eight), expanded octets (PCl₅, SF₆, more than eight), and odd-electron radicals.
- Resonance / delocalisation
- When several equivalent Lewis structures exist, the true structure is a resonance hybrid with electrons delocalised over the molecule; real bond lengths and orders are intermediate between the drawn forms (Mastery).
Chemical Bonding: Lewis Structures and Formal Charge FAQ
What is the difference between formal charge and oxidation number?
They split the bonding electrons differently. Formal charge divides each bond evenly (half the electrons to each atom) and is used to compare and rank candidate Lewis structures. Oxidation number assigns both bonding electrons to the more electronegative atom and is used for redox bookkeeping (that comes in the electrochemistry topic). For choosing a Lewis structure, formal charge is the tool you want.
How do I decide which Lewis structure is 'best'?
Apply two rules in order. First, minimise the magnitude of formal charges across the molecule. If structures tie, put any negative formal charge on the more electronegative atom and positive formal charge on the less electronegative one. Also check that the formal charges sum to the overall charge of the species. This is exactly how you'd pick between the cyanate structures, and it is a common exam step.
When is it acceptable to break the octet rule?
For three recognised cases: electron-deficient molecules like BF₃ (boron is happy with six electrons), expanded octets like PCl₅ and SF₆ (central atoms in period 3 or below can hold more than eight), and odd-electron species (radicals). Second-row atoms (C, N, O, F) essentially never exceed an octet, so an expanded octet on carbon or nitrogen is a red flag that your structure is wrong.
Can Sia help me draw Lewis structures and assign formal charges?
Yes. Sia can count valence electrons, build the skeleton, complete octets, add multiple bonds, and then compute formal charges atom by atom to rank the candidate structures. It can also flag octet exceptions and walk resonance cases. It explains the method and checks your working; it does not do graded assessment, and UNSW academic-integrity rules apply.
Exam move
Turn Lewis structures into a fixed algorithm you never improvise: total the valence electrons (adjust for the ion charge), put the least electronegative atom in the centre, connect with single bonds, complete the outer octets, then move lone pairs into multiple bonds if the central atom is short. Then always run the formal-charge check — FC = valence − lone-pair − ½ bonding for each atom — and use it to choose between candidate structures (smallest |FC|, negative FC on the more electronegative atom, charges summing correctly). Build fluency across the common organic functional groups, because 'draw and describe the bonding' questions reuse the same handful. Keep the three octet exceptions (BF₃, PCl₅/SF₆, radicals) memorised as recognisable cases. For the Mastery layer, practise drawing all resonance forms of ions like CO₃²⁻ and stating the delocalised, fractional-bond-order picture. When an electron count won't balance, ask Sia to recount it step by step.
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