Difficulty: Introductory to Intermediate | Prerequisites: Bonding fundamentals (Part 1), valence electrons, electronegativity
Once you know why atoms bond (Part 1), the next step is learning how to represent those bonds on paper. Lewis structures are the standard notation for showing bonding and nonbonding electrons in a molecule. Formal charge tells you whether your Lewis structure is reasonable, and resonance explains why some molecules cannot be captured by a single drawing. This material is essential groundwork for VSEPR geometry and molecular polarity, which come next.
Lewis structures show how valence electrons are distributed in a molecule. Formal charge checks whether each atom has a sensible number of electrons assigned to it. When more than one valid Lewis structure exists for the same molecule, the real electron distribution is a blend (resonance hybrid) of all of them.
Lewis structure (Lewis dot structure)
A diagram showing all valence electrons in a molecule as dots (lone pairs) and lines (bonding pairs). In simple terms, it is the standard way to draw a molecule on paper, with lines for bonds and dots for unshared electrons.
Bonding electrons
Electrons that are shared between two atoms in a covalent bond, represented as lines in a Lewis structure. Think of them as the electrons doing the work of holding two atoms together.
Nonbonding electrons (lone pairs)
Valence electrons that remain on a single atom and are not involved in bonding, shown as paired dots. In simple terms, these are "leftover" electrons that sit on one atom after all the bonds have been made.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons minus nonbonding electrons minus half of the bonding electrons. Think of it as a bookkeeping tool that checks whether each atom "owns" the right number of electrons in your drawing.
Oxidation number
The charge an atom would have if all bonding electrons were assigned to the more electronegative atom. In simple terms, formal charge splits bonding electrons 50/50; oxidation number gives them all to whichever atom pulls harder. Formal charge underestimates electronegativity effects, oxidation number overestimates them.
Resonance
The concept that the true electron distribution in some molecules is a weighted average of two or more valid Lewis structures, none of which alone represents reality. Think of it as a molecule that cannot be captured by a single drawing, so you draw several and understand the real structure as a blend.
Resonance hybrid
The actual structure of a molecule that exhibits resonance, representing the average of all contributing resonance structures. In simple terms, the hybrid is what the molecule looks like in reality, not any one of the individual drawings.
Lewis acid
A species that accepts a lone pair of electrons in a coordinate covalent bond. Think of it as the atom or molecule with an empty orbital, ready to receive electrons.
Lewis base
A species that donates a lone pair of electrons in a coordinate covalent bond. In simple terms, it is the atom or molecule with a spare lone pair to give away.
Draw the backbone. Place the least electronegative atom in the centre. Hydrogens and halogens are almost always terminal (outer) atoms.
Example: for HCN, carbon is the central atom, giving H–C–N
Count total valence electrons. Add up the valence electrons for every atom in the molecule.
Example: HCN has C = 4, H = 1, N = 5, total = 10
Convert bonds to electron pairs. Each bond line represents two shared electrons (two dots).
Complete octets on the outer atoms first. Use remaining valence electrons as lone pairs on the atoms surrounding the central atom.
Example: after the two bonds in H–C–N use four electrons, six remain to distribute
Check the central atom. If it has fewer than eight electrons, convert lone pairs on adjacent atoms into double or triple bonds until the central atom has a full octet.
Example: HCN requires a triple bond between C and N to give carbon its octet, producing H–C≡N
Formal charge is a way to verify that a Lewis structure is reasonable
Formula: FC = V − N(nonbonding) − ½ N(bonding)
V = number of valence electrons the atom normally has (from its group number)
N(nonbonding) = number of nonbonding (lone pair) electrons on the atom
N(bonding) = number of bonding electrons around the atom
The best Lewis structure is the one where formal charges are minimised (ideally zero on every atom)
Formal charge underestimates the effect of electronegativity differences, while oxidation numbers overestimate them
Formal charge assumes bonding electrons are shared equally
Oxidation numbers assume the more electronegative atom has 100% of the bonding pair
Example: in CO₂, the oxidation number on C is +4 and on each O is −2, but the formal charges are all zero
Some molecules have more than one valid Lewis structure. The real electron distribution is a hybrid (weighted average) of all of them.
The resonance hybrid is the actual structure; individual resonance structures are just contributors.
Stability guidelines for resonance structures:
A structure with no formal charges is more stable than one with formal charges
Smaller separation between opposite formal charges is more stable than large separation
Negative formal charge on the more electronegative atom is more stable than negative charge on a less electronegative atom
The more stable a resonance structure, the more it contributes to the hybrid (major contributor). Less stable structures with formal charges are minor contributors.
Formal charge:
FC = V − N(nonbonding) − ½ N(bonding)
Where V is the number of valence electrons the isolated atom normally has, N(nonbonding) is the count of lone-pair electrons on that atom, and N(bonding) is the count of electrons in bonds to that atom.
Lewis structures are the starting point for predicting molecular geometry, which in turn determines how drugs bind to receptors and how catalysts work at the molecular level.
Resonance explains why the carbonate ion (CO₃²⁻) has three identical C–O bonds rather than one short double bond and two long single bonds. This matters in materials science and mineralogy, where the symmetry of the carbonate ion determines how crystals like calcite form.
Students often think resonance structures are different molecules that the compound flips between. They are not. The molecule does not switch; it exists as the hybrid at all times.
A frequent error is placing the most electronegative atom in the centre of a Lewis structure. The least electronegative atom goes in the centre (with hydrogen and halogens always terminal).
Students sometimes believe that a formal charge of zero means the atom has no charge at all. Formal charge is a bookkeeping tool for a specific Lewis structure; the real charge distribution depends on electronegativity.
Students often forget to check formal charges after drawing a Lewis structure. If your structure gives carbon a formal charge of −1 while oxygen has +1, that is a sign to redraw with the negative charge on the more electronegative atom.
⚠️ Drawing Lewis structures is one of the most frequently tested skills in general chemistry. Practise the five-step method until it is automatic.
⚠️ Formal charge calculations appear regularly on exams. Be ready to calculate FC for each atom and identify which Lewis structure is most stable.
⚠️ Resonance questions often ask you to draw all valid structures and identify the major contributor. Remember: fewest formal charges wins, and negative charge belongs on the more electronegative atom.
⚠️ The distinction between formal charge and oxidation number is a common exam trap. Know that formal charge splits electrons 50/50, while oxidation number assigns them all to the more electronegative atom.
True or false: the most electronegative atom should be placed at the centre of a Lewis structure. (False, the least electronegative atom goes in the centre)
Fill in the blank: formal charge = V − ______ − ½ N(bonding). (N(nonbonding), i.e. lone pair electrons)
True or false: resonance structures represent different molecules that rapidly interconvert. (False, the molecule exists as a single hybrid at all times)
Fill in the blank: a resonance structure with no formal charges is ______ stable than one with formal charges. (more)
True or false: in CO₂, the formal charge on carbon is +4. (False, the formal charge on carbon in CO₂ is zero; +4 is the oxidation number)
Q: Walk through drawing the Lewis structure for HCN, step by step.
A: Carbon is the central atom (least electronegative of C, H, N). Total valence electrons: H = 1, C = 4, N = 5, total = 10. Draw H–C–N (two bonds = four electrons, six remaining). Place remaining electrons as lone pairs on nitrogen first. Carbon still lacks a full octet, so convert two lone pairs from nitrogen into bonding pairs, forming a triple bond: H–C≡N. Nitrogen keeps one lone pair. All atoms now satisfy the octet rule (hydrogen has two).
Q: Calculate the formal charge on each atom in carbon dioxide (O=C=O).
A: Carbon: V = 4, nonbonding = 0, bonding = 8. FC = 4 − 0 − 4 = 0. Each oxygen: V = 6, nonbonding = 4, bonding = 4. FC = 6 − 4 − 2 = 0. All formal charges are zero, confirming this is a stable Lewis structure.
Q: Two resonance structures for a molecule are proposed. Structure A has zero formal charges on all atoms. Structure B has a +1 on carbon and a −1 on oxygen. Which is the major contributor?
A: Structure A is the major contributor because it has no formal charges. Structure B, while valid, is less stable and contributes less to the resonance hybrid.
Q: What is the difference between formal charge and oxidation number?
A: Formal charge assumes bonding electrons are shared equally (split 50/50), which underestimates the effect of electronegativity differences. Oxidation number assumes the more electronegative atom takes 100% of the bonding electrons, which overestimates the effect. Neither is the "real" charge, but both are useful bookkeeping tools.
Lewis structures feed directly into VSEPR theory (Part 3 of these notes): the number of bonding pairs and lone pairs around a central atom determines molecular geometry. Formal charge and resonance also connect to acid-base chemistry, because Lewis acid-base reactions are defined by lone-pair donation (coordinate covalent bonds).
Resonance is particularly important in organic chemistry, where delocalised electrons explain the stability of benzene rings and the behaviour of reactive intermediates like carbocations.
Lewis structure, Lewis dot structure, electron dot diagram, formal charge, oxidation number, resonance, resonance hybrid, resonance structures, contributing structures, major contributor, minor contributor, bonding electrons, nonbonding electrons, lone pairs, Lewis acid, Lewis base, coordinate covalent bond, dative bond, central atom, terminal atom, valence electrons, octet completion, double bond, triple bond, CHEM 3, Michigan State University, general chemistry