Difficulty: Introductory to Intermediate. Prerequisites: Lewis dot structures, basic atomic structure, periodic table familiarity.
This lecture covers two foundational ideas in organic chemistry: resonance and electronegativity. Resonance explains why certain molecules cannot be represented by a single Lewis structure, while electronegativity describes how strongly atoms attract electrons. Both concepts are essential for understanding molecular behaviour, polarity, and reactivity in later chapters. If you are unfamiliar with Lewis dot structures or the layout of the periodic table, review those first.
Some molecules have electrons that are spread across multiple bonds rather than locked in one place, and no single Lewis structure captures that. This is resonance. Electronegativity tells you how strongly a given atom pulls on shared electrons, which in turn determines how charge is distributed across a bond.
Resonance
The phenomenon where no single Lewis dot structure fully describes the electron distribution in a molecule or ion. Multiple valid structures (resonance structures) are drawn, connected by double-headed arrows.
Think of it as: the molecule is a blend of all the structures at once, not flipping between them.
Resonance hybrid
The actual structure of a molecule that exhibits resonance. It is the weighted average of all contributing resonance structures, not any one of them individually.
In simple terms, this means the real molecule sits somewhere between the individual drawings you put on paper.
Resonance structures (contributing structures)
The individual Lewis dot structures drawn to represent different possible electron arrangements for the same molecule. Only electrons move between structures; atoms stay fixed.
Think of it as: different "snapshots" of where the electrons could be, all layered on top of each other.
Lewis dot structure
A diagram showing all valence electrons in a molecule, with bonding pairs as lines and lone pairs as dots.
Electronegativity
A measure of the strength of attraction between an atom's nucleus and its valence (outer-shell) electrons. More electronegative atoms pull bonding electrons closer to themselves.
In simple terms, this means electronegativity is how greedy an atom is for shared electrons.
Coulomb's Law (as applied to atomic interactions)
The energy of interaction between two charged particles is proportional to the product of their charges and inversely proportional to the distance between them, scaled by the dielectric constant of the medium.
Think of it as: closer charges and bigger charges mean a stronger pull, but the surrounding material can weaken it.
Dielectric constant (insulating capacity)
A property of the medium (such as water or vacuum) that describes how much it reduces the electrostatic interaction between charged particles. Higher dielectric constant means more insulation between charges.
When a molecule or ion cannot be accurately represented by a single Lewis dot structure, we draw multiple resonance structures.
These structures are connected by double-headed arrows (not equilibrium arrows).
Only electrons move between resonance structures. The atoms stay in exactly the same positions.
The molecule does not switch back and forth between structures. It exists as all of them simultaneously.
Three equivalent resonance structures can be drawn for the nitrate ion.
In each structure, nitrogen carries a formal positive charge (+1), one oxygen carries a formal negative charge, and the double bond sits on a different oxygen.
The resonance hybrid (the real structure) has three equivalent N–O bonds, each with a bond order of 1⅓.
Each oxygen carries a partial charge of ⅔ of a negative charge (since in two out of three structures, any given oxygen carries the full negative charge).
The resonance hybrid is the true picture of the molecule: a weighted average of all resonance structures.
Partial bonds (drawn as dashed lines) represent the averaged bond order.
Partial charges (δ+ and δ−) represent the averaged charge distribution.
Electronegativity measures how strongly an atom's nucleus attracts its valence electrons.
Periodic table trend (increases going):
Left to right across a period (more protons in the nucleus, stronger pull)
Bottom to top within a group (valence electrons are closer to the nucleus)
Fluorine is the most electronegative element.
Describes the energy of interaction between charged particles.
The interaction energy depends on:
The magnitude of both charges (larger charges, stronger interaction)
The distance between them (shorter distance, stronger interaction)
The dielectric constant of the medium (higher dielectric constant, weaker interaction)
This law underpins why electronegativity exists: the nucleus-electron attraction follows the same charge-distance relationship.
E = \frac{q_1 \times q_2}{\varepsilon \times r}Where:
E = energy of interaction between the two charged particles
q₁ and q₂ = charges on the two particles
ε (epsilon) = dielectric constant of the medium (its insulating capacity)
r = distance between the two charges
Increases from left to right across a period.
Increases from bottom to top within a group.
The top-right corner of the periodic table (fluorine) has the highest electronegativity.
The bottom-left corner (francium, caesium) has the lowest.
Three resonance structures, each with one N=O double bond and two N–O single bonds.
Average bond order per N–O bond = (1 + 1 + 2) / 3 = 1⅓.
Students often think the molecule "flips" between resonance structures. It does not. The molecule exists as one fixed hybrid at all times; the individual structures are just our limited way of drawing it.
Students sometimes move atoms when drawing different resonance structures. Only electrons move. If you have moved an atom, you have drawn a different molecule, not a resonance structure.
Students often assume that one resonance structure is "more real" than the others. All equivalent resonance structures contribute equally to the hybrid. (Non-equivalent structures do contribute unequally, but that is a later topic.)
Students frequently confuse electronegativity with electron affinity. Electronegativity is a tendency within a bond (how an atom pulls on shared electrons). Electron affinity is the energy change when a free atom gains an electron.
⚠️ You will almost certainly be asked to draw resonance structures for a given ion or molecule. Practise the nitrate ion, carbonate ion, and carboxylate group until you can do them from memory.
⚠️ Expect a question asking you to calculate or describe partial bond order from a set of resonance structures (e.g. "What is the bond order of each N–O bond in nitrate?").
⚠️ The electronegativity periodic table trend (up and to the right) is tested constantly, often as a ranking question: "Arrange these elements in order of increasing electronegativity."
⚠️ Coulomb's Law may appear as a conceptual question: "What happens to the interaction energy if the distance doubles?" (It halves.)
True or False: Resonance structures show the molecule rapidly switching between different forms.
Answer: False. The molecule exists as one hybrid at all times.
Fill in the blank: In resonance, only ______ move; atoms remain fixed.
Answer: Electrons.
True or False: Electronegativity increases as you move down a group on the periodic table.
Answer: False. It increases as you move up a group.
Fill in the blank: In Coulomb's Law, the dielectric constant (ε) appears in the ______ of the equation.
Answer: Denominator.
True or False: Each N–O bond in the nitrate ion has a bond order of exactly 1.
Answer: False. The bond order is 1⅓.
Q: What is a resonance hybrid, and how does it differ from an individual resonance structure?
A: A resonance hybrid is the actual, real structure of the molecule, representing the weighted average of all contributing resonance structures. An individual resonance structure is just one of several possible Lewis drawings and does not fully describe the molecule on its own.
Q: Draw the three resonance structures of the nitrate ion (NO₃⁻). What is the bond order of each N–O bond in the hybrid?
A: Each structure has nitrogen at the centre bonded to three oxygens, with one N=O double bond and two N–O single bonds. The double bond rotates to a different oxygen in each structure. The average bond order is (2 + 1 + 1) / 3 = 1⅓.
Q: Using the periodic table trend, rank the following in order of increasing electronegativity: Na, Cl, F, O.
A: Na < Cl < O < F. Electronegativity increases going up and to the right.
Q: In Coulomb's Law, what happens to the interaction energy if you double the distance between two charges while keeping everything else constant?
A: The interaction energy is halved, because E is inversely proportional to distance (r).
Q: Why does resonance involve movement of electrons but not atoms?
A: Resonance structures describe different ways of distributing electrons within the same molecular framework. If you moved an atom, you would have a structural isomer, which is a different molecule entirely.
Resonance connects directly to aromaticity (benzene and its derivatives), which relies on delocalised electrons across a ring. Understanding resonance is also essential for predicting acid/base strength: a conjugate base stabilised by resonance is more stable, making the parent acid stronger.
Electronegativity feeds into bond polarity and dipole moments (covered in the second part of this lecture), and later into understanding reaction mechanisms where electron-rich sites attack electron-poor sites.
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