Difficulty: Introductory to Intermediate | Prerequisites: Basic atomic structure, periodic table groups, electron configuration basics
Lewis structures are dot-and-line diagrams that show how valence electrons are arranged in a molecule or ion. You draw them by counting valence electrons, connecting atoms with bonds, and distributing leftover electrons to satisfy the octet rule. When a single Lewis structure cannot capture the real electron arrangement, you draw resonance structures, and the actual molecule is best described as a blend (resonance hybrid) of all of them.
Lewis structure (electron dot structure)
A diagram showing the bonding between atoms in a molecule and the lone pairs of electrons that may exist. Each line represents a shared pair; each dot represents an unshared electron.
Think of it as a map of where the electrons live in a molecule.
Electronegativity (EN)
A measure of how strongly an atom attracts bonding electrons toward itself. Higher EN means a stronger pull. Differences in EN between two atoms determine whether a bond is ionic, polar covalent, or nonpolar covalent.
In simple terms, this is how greedy an atom is for electrons.
Valence electrons
The electrons in an atom's outermost shell, and the only ones involved in bonding. For main-group elements, the group number on the periodic table tells you how many there are (e.g. Group 6 = 6 valence electrons).
Think of it as the electrons that do the work in chemical reactions.
Octet rule
Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (two for hydrogen). This mirrors the electron configuration of the nearest noble gas.
In simple terms, most atoms want eight electrons around them to be stable.
Resonance structures
Two or more valid Lewis structures for the same molecule or ion that differ only in the placement of electrons (not atoms). No single structure is "correct" on its own.
Think of them as different, equally reasonable snapshots of the same molecule.
Resonance hybrid
The actual, experimentally observed structure of a molecule that exhibits resonance. It is not any one resonance structure, but a weighted blend of all of them, with bond lengths and strengths intermediate between the individual forms.
In simple terms, the real molecule is the average, not any single drawing.
Formal charge
A bookkeeping tool: the charge an atom would have if all bonding electrons were shared equally. Calculated as (valence electrons) minus (lone pair electrons) minus (half of bonding electrons). The best Lewis structure is typically the one where formal charges are minimised and negative formal charges sit on the more electronegative atom.
Think of it as a way to check whether your Lewis structure is the most reasonable one.
Lone pair (non-bonding pair)
A pair of valence electrons that belongs to one atom and is not shared in a bond. Lone pairs take up space around an atom and affect molecular geometry.
In simple terms, electrons that hang around an atom without being part of a bond.
Bonding pair
A pair of electrons shared between two atoms in a covalent bond. A single bond is one bonding pair, a double bond is two, and a triple bond is three.
Think of it as the glue holding two atoms together.
The least electronegative atom goes in the centre. Hydrogen and fluorine are always terminal (outer) atoms, never central. If two atoms have similar electronegativities, the one that can form the most bonds, or the one that appears only once in the formula, is usually central.
Step 1 – Count total valence electrons. Add the valence electrons of every atom. For anions, add electrons equal to the charge; for cations, subtract.
Step 2 – Place the central atom and draw single bonds. Connect each outer atom to the central atom with one single bond (each bond uses two electrons).
Step 3 – Distribute remaining electrons to terminal atoms. Give each outer atom enough lone pairs to complete its octet (or duet for hydrogen).
Step 4 – Place leftover electrons on the central atom. Whatever electrons remain go as lone pairs on the central atom.
Step 5 – Check octets and form multiple bonds if needed. If the central atom has fewer than eight electrons, convert one or more lone pairs from a neighbouring atom into bonding pairs (double or triple bonds) until the octet is satisfied.
Sulfur is less electronegative than bromine, so S is the central atom.
Valence electrons: S has 6 (Group 16), each Br has 7 (Group 17). Total = 6 + 7 + 7 = 20.
Draw S in the centre with single bonds to each Br (uses 4 electrons, leaving 16).
Place three lone pairs on each Br (uses 12 electrons, leaving 4).
The remaining 4 electrons go on S as two lone pairs.
Check: every atom has 8 electrons around it. Done.
If, after distributing electrons, the central atom is short of an octet, move a lone pair from a terminal atom to form a double bond (or two lone pairs for a triple bond). Each double bond replaces one lone pair on the terminal atom with a shared pair.
For polyatomic ions, adjust the electron count before you start: add electrons for negative charges, subtract for positive charges. Enclose the finished structure in square brackets and write the charge outside.
Some molecules and ions cannot be described by a single Lewis structure because the electrons are delocalised (spread out) rather than fixed in one arrangement. Resonance structures are two or more valid Lewis structures for the same species that differ only in where the electrons sit. The atoms stay in the same positions.
The real molecule does not flip between these forms. It exists as a single entity called the resonance hybrid, whose properties are a blend of all the contributing structures.
Resonance is likely whenever you can draw a valid Lewis structure with a double bond in more than one position, or whenever equivalent atoms around a central atom each have a plausible claim to a double bond.
Total valence electrons: C has 4, each O has 6, plus 2 for the charge = 4 + 18 + 2 = 24.
One valid structure places a C=O double bond on one oxygen and C–O single bonds on the other two.
But all three oxygens are equivalent, so you can draw three structures, each with the double bond on a different oxygen.
The resonance hybrid has three identical C–O bonds, each with a bond order of about 1.33 and a bond length of roughly 128 pm, intermediate between a single bond (~143 pm) and a double bond (~120 pm).
Total valence electrons: 3 × 6 = 18.
The central oxygen forms a double bond with one terminal oxygen and a single bond with the other, plus lone pairs to complete all octets.
A second resonance structure swaps which terminal oxygen holds the double bond.
The resonance hybrid has two equivalent O–O bonds, each with a bond order of 1.5 and a length between that of a single O–O bond and a double O=O bond.
Experimental measurements (X-ray crystallography, infrared spectroscopy) confirm that resonance hybrids have bond lengths and bond energies intermediate between the extremes suggested by any one structure. The more equivalent resonance structures a molecule has, the more stabilised it is (resonance stabilisation energy).
Several polyatomic ions that appear throughout general chemistry exhibit resonance. Knowing them saves time on exams and in later courses.
Sulphate (SO₄²⁻) – Four equivalent S–O bonds in the hybrid; formal-charge arguments favour structures with expanded octets on sulphur (double bonds to some oxygens).
Nitrate (NO₃⁻) – Three equivalent resonance structures, each with the double bond on a different oxygen. Bond order per N–O bond is 1.33.
Phosphate (PO₄³⁻) – Similar pattern to sulphate; four P–O bonds in the hybrid.
Acetate (C₂H₃O₂⁻) – Two resonance structures placing the double bond on either of the two oxygens. The two C–O bonds in the carboxylate group are equivalent.
Each of these ions crops up repeatedly in acid-base chemistry, solubility rules, and biochemistry, so recognising their resonance quickly is worth the effort.
Total valence electrons = sum of each atom's group-number electrons ± adjustment for charge (add for anions, subtract for cations).
Formal charge = (valence electrons of the free atom) – (lone pair electrons) – ½(bonding electrons).
Bond order (in resonance) = (total number of bonds across all resonance structures for a given pair of atoms) ÷ (number of resonance structures).
Electrons used by single bonds = 2 × (number of bonds drawn).
Remaining electrons = total valence electrons – electrons used by bonds.
Lewis structures are not just an academic exercise. Pharmaceutical chemists draw them to predict how drug molecules will interact with receptor sites, because the position of lone pairs and the electron density map of a molecule determine its shape and reactivity. Resonance stabilisation explains why the carbonate ion in limestone is so stable, why benzene is far less reactive than you would expect for a molecule with three "double bonds," and why the peptide bond in proteins is planar, which is fundamental to how proteins fold.
Students often think resonance structures are real, separate forms that the molecule flips between. They are not. The molecule exists as a single resonance hybrid at all times; the individual structures are just our limited way of drawing it.
Students frequently place the most electronegative atom in the centre. The opposite is correct: the least electronegative atom (excluding H and F) is the central atom.
A common error is forgetting to adjust the electron count for charged species. For every negative charge, add one electron; for every positive charge, subtract one.
Students sometimes assume that double bonds in resonance structures are "real" double bonds. In the hybrid, the bond order is fractional (e.g. 1.33 for nitrate), and the bond length sits between single and double.
⚠️ Drawing Lewis structures is a near-guaranteed exam question in general chemistry. You need the five-step method to be automatic.
⚠️ Formal charge calculation is commonly tested as a standalone problem and as a way to pick the "best" Lewis structure when several are possible.
⚠️ Resonance is heavily tested in the context of polyatomic ions (carbonate, nitrate, sulphate). Expect to be asked to draw all resonance structures and identify bond order.
⚠️ The distinction between resonance structures and the resonance hybrid is a favourite multiple-choice trap. Remember: the hybrid is the reality, the structures are the sketches.
⚠️ Octet-rule exceptions (expanded octets for Period 3+ elements, incomplete octets for B and Be) appear on most exams. Know which elements break the rule and why.
True or false: The central atom in a Lewis structure is always the most electronegative atom. (False – it is the least electronegative, excluding H and F.)
Fill in the blank: The total number of valence electrons in CO₃²⁻ is ____. (24)
True or false: Resonance structures differ in the positions of the atoms. (False – only the electrons move; atoms stay in the same arrangement.)
Fill in the blank: The bond order of each N–O bond in nitrate (NO₃⁻) is ____. (1.33)
True or false: A resonance hybrid has bond lengths that match one of its contributing resonance structures exactly. (False – the lengths are intermediate.)
Q: Draw the Lewis structure for water (H₂O) and state the total number of valence electrons.
A: Oxygen is the central atom (H is always terminal). Total valence electrons = 6 + 1 + 1 = 8. O forms single bonds to each H (using 4 electrons) and carries two lone pairs (using the remaining 4). Each H has a duet; O has an octet.
Q: How many resonance structures can be drawn for the nitrate ion (NO₃⁻), and what is the bond order of each N–O bond?
A: Three equivalent resonance structures, each placing the double bond on a different oxygen. Bond order = (1 double + 2 singles across three structures) / 3 = 4/3 ≈ 1.33.
Q: Calculate the formal charge on the central carbon in one resonance structure of CO₃²⁻ where carbon has one double bond and two single bonds to oxygen.
A: Carbon has 4 valence electrons. In this structure it has 0 lone pair electrons and 8 bonding electrons. Formal charge = 4 – 0 – ½(8) = 0.
Q: Why is the bond length in the carbonate ion (~128 pm) not equal to a standard C=O double bond (~120 pm) or a C–O single bond (~143 pm)?
A: Because the carbonate ion is a resonance hybrid of three equivalent structures. The real bond is a blend with a bond order of 1.33, so its length falls between the single and double bond values.
Q: Explain why hydrogen is never the central atom in a Lewis structure.
A: Hydrogen can form only one bond (it has one valence electron and needs only a duet). A central atom must bond to at least two other atoms, which hydrogen cannot do.
Lewis structures feed directly into VSEPR theory (Valence Shell Electron Pair Repulsion), which uses the number of bonding and lone pairs around the central atom to predict molecular geometry. If you can draw the Lewis structure, you can predict the shape.
Resonance and formal charge also set the stage for molecular orbital (MO) theory later in the course, which provides a more complete picture of electron delocalisation than resonance structures alone.
Understanding electronegativity and bond polarity here connects to intermolecular forces (dipole-dipole, hydrogen bonding, London dispersion), which in turn explain boiling points, solubility, and many physical properties covered in later chapters.
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