Difficulty: Introductory to Intermediate | Prerequisites: General chemistry (atomic structure, electron configuration, periodic trends)
This material forms the bedrock of organic chemistry. Every reaction mechanism, spectroscopy problem, and synthesis question you will encounter later depends on being fluent with structural representations, electron-flow arrows, resonance, and bond properties. If you skipped or forgot general chemistry's coverage of Lewis dot structures and electronegativity, revisit those first. The concepts here are not difficult in isolation, but they interlock: a weak grasp of formal charge, for example, will hobble your ability to evaluate resonance contributors or predict acid-base behaviour.
Organic chemistry uses three main structural drawing styles (Lewis, condensed, and bond-line) to represent molecules, and four types of arrows to show what electrons are doing. Resonance structures let you map out electron delocalisation within a single molecule, and formal charge tells you which resonance form matters most. Knowing how to classify isomers and relate bond strength, length, and polarity rounds out the toolkit you need before tackling reactions.
Lewis structure (electron dot structure)
A diagram showing every valence electron in a molecule as dots or lines, including lone pairs and bonding pairs. Used to work out formal charges and spot bonding interactions.
In simple terms, it is the most detailed way to draw a molecule, where every electron is accounted for.
Condensed structure
A shorthand that drops the bonds and lone pairs, writing atoms and their attached hydrogens in a linear string (e.g. CH3CH2OH). Useful for quick notation of larger molecules.
Think of it as: the texting abbreviation of a Lewis structure.
Bond-line structure (skeletal structure)
A drawing where lines represent bonds and vertices or line endpoints represent carbon atoms. Hydrogens on carbon are implied. This is the standard way organic chemists draw molecules.
Think of it as: the default shorthand of organic chemistry, where you only draw the skeleton.
Dipole arrow
An arrow showing bond polarity, pointing from the less electronegative atom toward the more electronegative one. A small plus sign sits at the tail end, and the arrowhead carries the partial negative charge.
Equilibrium arrows
A pair of opposing arrows (⇌) indicating a reversible reaction. The relative lengths can indicate which direction is favoured.
Resonance arrow (double-headed arrow)
A double-headed arrow (↔) placed between resonance structures of the same molecule. It means "these are alternative electron arrangements," not "these interconvert."
Curved arrow (curly arrow)
An arrow showing the movement of an electron pair during a reaction or resonance interconversion. The tail starts at the electron source (a lone pair or a bond) and the head points to the electron sink (an electron-poor centre).
Think of it as: a tracking device for where electrons go.
Resonance structures
Alternative Lewis structures of the same molecule that differ only in the placement of electrons, not atoms. The true molecule is a weighted average (a hybrid) of all valid resonance forms.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: valence electrons minus lone-pair electrons minus half the bonding electrons. It helps identify the most stable resonance contributor.
Identical structures
The same molecule drawn in different orientations or formats. Rotating or flipping the drawing does not create a new compound.
Constitutional isomers (structural isomers)
Molecules with the same molecular formula but different connectivity of atoms. For example, butane and isobutane are both C4H10 but have different carbon skeletons.
Stereoisomers
Molecules with the same connectivity but different spatial arrangement of atoms. Covered in depth in Chapter 3.
Bond strength (bond dissociation energy)
The energy required to break one mole of a particular bond in the gas phase. Stronger bonds release more energy when formed and require more energy to break.
Bond length
The distance between the nuclei of two bonded atoms. Inversely related to bond strength: shorter bonds are stronger.
Bond polarity
The uneven distribution of electron density in a bond caused by a difference in electronegativity between the two atoms. Larger electronegativity differences produce more polar bonds.
Electronegativity
A measure of how strongly an atom attracts bonding electrons toward itself. Increases across a period (left to right) and up a group on the periodic table.
Three drawing formats appear throughout organic chemistry. The choice depends on what you need to communicate.
Lewis structures show every valence electron. Draw them when you need to see lone pairs, assign formal charges, or count electrons around an atom. They are essential for molecules where electron distribution drives the chemistry (e.g. working out why one oxygen is more nucleophilic than another).
Condensed structures compress the information into a single line of text. Bonds and lone pairs are omitted. Each carbon is written with its attached hydrogens: CH3CH2OH for ethanol. Useful in written answers and textbook shorthand, but they hide geometry.
Bond-line (skeletal) structures are the workhorse notation. Lines are bonds, vertices and endpoints are carbons, and hydrogens on carbon are implied. Heteroatoms (O, N, S, halogens) and their hydrogens are always written explicitly. Master this format early: it is how organic chemists think on paper.
Four arrow types, each with a distinct meaning:
Dipole arrows show bond polarity. The arrow points from the less electronegative atom (partial +) toward the more electronegative atom (partial –). A crosshatch at the tail marks the positive end.
Equilibrium arrows (⇌) indicate a reversible reaction. Both forward and reverse processes are occurring.
Resonance arrows (↔) sit between resonance structures. They do not mean a reaction is happening. They mean the molecule is a hybrid of the forms shown.
Curved (curly) arrows track the movement of electron pairs in mechanisms and resonance interconversions. The tail always starts at the electron source (a lone pair or a bond). The head always points at the electron sink (an electron-poor atom or a bond being formed). A full curved arrow moves two electrons; a half-headed (fishhook) arrow moves one electron (radical chemistry).
The cardinal rule: electrons flow from electron-rich to electron-poor. Getting the direction wrong is one of the most common mistakes in mechanism problems.
Resonance structures are different Lewis drawings of the same molecule. Only the electrons move; the atoms stay put. The real molecule is a resonance hybrid, a weighted blend of all contributors.
How to evaluate resonance contributors:
The structure with the fewest formal charges is generally the major contributor.
Negative formal charges are more stable on more electronegative atoms (oxygen rather than carbon, for instance).
Positive formal charges are better tolerated on less electronegative atoms.
All atoms should have complete octets where possible (second-row elements especially).
Formal charge is calculated per atom:
Formal charge = (valence electrons) – (lone pair electrons) – (1/2 × bonding electrons)
Resonance stabilisation matters because it lowers a molecule's energy. This directly affects reactivity, acid-base strength, and spectroscopic behaviour.
When two structures have the same molecular formula, you need to classify the relationship:
Identical structures: the same molecule, just drawn differently. Rotate or flip the drawing mentally.
Resonance structures: same molecule, same atom connectivity, different electron placement. Linked by a double-headed arrow.
Constitutional isomers: same formula, different connectivity. The atoms are bonded in a different order.
Stereoisomers: same formula, same connectivity, different spatial arrangement. Covered fully in the stereochemistry unit.
A systematic approach: first check if the molecular formulae match. If they do, check whether the atom connectivity is the same. If connectivity matches, the difference must be spatial (stereoisomers) or electronic (resonance forms).
Three bond properties that recur throughout the course:
Bond order and strength: triple bonds > double bonds > single bonds in strength. A C≡C triple bond (~837 kJ/mol) is stronger than a C=C double bond (~614 kJ/mol), which is stronger than a C–C single bond (~347 kJ/mol).
Bond order and length: higher bond order means shorter bond length. The triple bond is the shortest of the three.
Polarity: determined by the electronegativity difference between the bonded atoms. A C–O bond is polar (oxygen pulls electron density toward itself). A C–C bond is nonpolar. Polar bonds are shown with dipole arrows.
These properties feed directly into predicting intermolecular forces, physical properties (boiling point, solubility), and reactivity patterns.
\text{Formal Charge} = V - L - \frac{B}{2}Where V = valence electrons of the free atom, L = lone pair electrons on the atom, B = bonding electrons around the atom.
Bond order relationship (for carbon-carbon bonds):
Property | Single bond (C–C) | Double bond (C=C) | Triple bond (C≡C) |
|---|---|---|---|
Bond length | ~154 pm | ~134 pm | ~120 pm |
Bond energy | ~347 kJ/mol | ~614 kJ/mol | ~837 kJ/mol |
Hybridisation | sp³ | sp² | sp |
The pattern: as bond order increases, length decreases and strength increases.
Curved-arrow notation is the language pharmaceutical chemists use to design drug molecules: every new reaction in a synthesis plan is mapped with electron-flow arrows before anyone steps into a lab. Bond polarity determines whether a molecule dissolves in water or in fat, which is why medicinal chemists care about it when predicting whether a drug can cross a cell membrane. Resonance stabilisation explains why aspirin's carboxylate group stays deprotonated at physiological pH, which controls how the drug is absorbed.
Students often draw curved arrows pointing from the electrophile to the nucleophile. This is backwards. Electrons always flow from the electron-rich species (nucleophile) to the electron-poor species (electrophile).
Resonance structures are sometimes mistaken for equilibria or for distinct molecules that interconvert. They are not. The molecule does not flicker between forms; it exists as a single hybrid.
A common error is moving atoms when drawing resonance structures. Only electrons move. If you have relocated an atom, you have drawn a constitutional isomer, not a resonance structure.
Students frequently assume that the resonance contributor with the most bonds is always the major one. This is often true, but a structure that places a negative charge on a very electropositive atom (like carbon instead of oxygen) may not be the dominant contributor, even if it has more bonds.
⚠️ Formal charge calculation is a near-guaranteed exam question. Practise until you can do it in your head for any atom in a Lewis structure.
⚠️ Expect at least one question asking you to classify pairs of structures as identical, constitutional isomers, stereoisomers, or resonance structures.
⚠️ Curved arrow direction (electron-rich to electron-poor) is tested both directly and indirectly. If your arrows point the wrong way, the entire mechanism answer is marked wrong.
⚠️ Questions often ask you to rank resonance contributors by stability. Remember: fewest formal charges wins, and negative charge on electronegative atoms wins.
⚠️ Bond-strength and bond-length trends (single vs. double vs. triple) appear in multiple-choice sections regularly.
True or False: In a bond-line structure, every carbon atom and its attached hydrogens are written explicitly. (Answer: False. Carbon atoms are implied at vertices and endpoints; their hydrogens are implied too.)
Fill in the blank: Curved arrows always point from ______ to ______. (Answer: electron-rich to electron-poor.)
True or False: Resonance structures differ in the positions of both atoms and electrons. (Answer: False. Only electrons move; atoms stay in the same positions.)
Fill in the blank: The formal charge formula is: valence electrons minus ______ minus half of ______. (Answer: lone pair electrons; bonding electrons.)
True or False: A triple bond is longer than a double bond. (Answer: False. A triple bond is shorter and stronger.)
Q: Calculate the formal charge on nitrogen in NH4+. Nitrogen has five valence electrons, zero lone pairs, and eight bonding electrons.
A: Formal charge = 5 – 0 – (8/2) = 5 – 0 – 4 = +1. The nitrogen carries a formal charge of +1.
Q: Two structures have the molecular formula C4H10. Structure A is a straight chain of four carbons. Structure B has a three-carbon chain with a methyl branch. What is the relationship between A and B?
A: They are constitutional isomers. They share the same molecular formula but differ in atom connectivity.
Q: Draw the curved arrow for the reaction of hydroxide ion (HO–) with a proton on HCl. Where does the tail start, and where does the head point?
A: The tail starts on a lone pair of the oxygen in HO– (the electron-rich nucleophile). The head points at the hydrogen of HCl (the electron-poor site). This shows the electrons flowing from nucleophile to electrophile.
Q: You are given three resonance structures for the same molecule. Structure I has no formal charges. Structure II has a negative charge on oxygen and a positive charge on carbon. Structure III has a negative charge on carbon and a positive charge on oxygen. Rank them from most to least stable.
A: Structure I (no formal charges) > Structure II (negative on electronegative oxygen) > Structure III (negative on less electronegative carbon). Fewer formal charges and negative charge on more electronegative atoms stabilise a resonance contributor.
Q: Rank the following bonds from shortest to longest: C–C, C=C, C≡C.
A: C≡C (shortest) < C=C < C–C (longest). Higher bond order means shorter bond length.
This connects to acid-base chemistry because formal charge and electronegativity directly determine acid strength and conjugate base stability. The resonance concepts here reappear whenever you evaluate whether a conjugate base is stabilised by delocalisation.
Bond polarity and intermolecular forces (covered in the next set of notes) build on the electronegativity and dipole concepts introduced here. Understanding polarity is the gateway to predicting boiling points, solubility, and reaction environments.
Curved arrow notation is the foundation for every reaction mechanism you will study from substitution and elimination reactions through to carbonyl chemistry.
Lewis dot structure, electron dot diagram, skeletal structure, line-angle structure, bond-line notation, condensed formula, structural formula, curly arrow, electron pushing, electron flow, arrow pushing mechanism, resonance hybrid, delocalisation, formal charge calculation, constitutional isomers, structural isomers, stereoisomers, identical molecules, bond dissociation energy, BDE, bond order, bond length, bond polarity, electronegativity difference, dipole moment, organic chemistry I, ochem 1 study notes, CHEM 2301