Source: CHEM 2301 Learning Objectives, Ch. 1
Tags: organic chemistry, Lewis structures, bond-line structures, condensed structures, curved arrows, resonance, formal charge, electronegativity, dipole, acid-base, Brønsted-Lowry, Lewis acid, Lewis base, nucleophile, electrophile, pKa, conjugate acid, conjugate base, proton transfer
Difficulty: Introductory to Intermediate | Prerequisites: General Chemistry I and II (atomic structure, electron configuration, periodic trends, basic acid-base concepts).
Chapter 1 is the foundation for everything else in organic chemistry. It covers how to draw and interpret molecular structures, how electrons move (curved arrow notation), and two frameworks for understanding acid-base behaviour: Brønsted-Lowry and Lewis. If you cannot draw a correct Lewis structure, push electrons with curved arrows, or predict which way a proton transfer will go, every later chapter will be harder than it needs to be. This material assumes you are comfortable with electronegativity trends, formal charge, and the basics of chemical bonding from general chemistry.
Chapter 1 teaches you the language of organic chemistry: how to represent molecules (Lewis, condensed, bond-line structures), how to show electron movement (curved arrows), how to evaluate resonance contributors, and how to analyse acid-base reactions using both Brønsted-Lowry and Lewis models. Master these skills first; they recur in every subsequent chapter.
Lewis structure
A diagram showing all valence electrons in a molecule as bonds (shared pairs) and lone pairs (unshared pairs). Think of it as the most detailed electron map of a molecule.
Condensed structure
A shorthand that groups atoms and their attached hydrogens together (e.g. CH₃CH₂OH) without drawing every bond. In simple terms, it is a compact way to write a molecule in a single line of text.
Bond-line structure (skeletal structure)
A minimalist drawing where carbon atoms sit at line endpoints and vertices, hydrogen atoms on carbon are implied, and only heteroatoms and their hydrogens are written out. This is the standard way organic chemists draw molecules.
Curved arrow notation
Arrows that show the movement of electron pairs from an electron-rich site to an electron-poor site. A full curved arrow moves two electrons; a fishhook arrow moves one. Think of it as a map of where electrons go during a reaction or resonance.
Resonance structures
Two or more Lewis structures for the same molecule that differ only in the placement of electrons, not atoms. The real molecule is a weighted average (hybrid) of all contributors. In simple terms, resonance means the electrons are spread out more than any single drawing can show.
Major resonance contributor
The resonance structure that contributes most to the hybrid. It is the one with the most bonds, the fewest formal charges, negative charges on more electronegative atoms, and the most atoms with complete octets.
Formal charge
The charge assigned to an atom in a Lewis structure: formal charge = (valence electrons) − (lone pair electrons) − (½ bonding electrons). It is a bookkeeping tool for tracking electron ownership, not a measure of real electron density.
Electronegativity
The tendency of an atom to attract shared electrons toward itself in a covalent bond. Increases going up and to the right on the periodic table (F is the most electronegative element).
Bond polarity / dipole
When two bonded atoms differ in electronegativity, the shared electrons sit closer to the more electronegative atom, creating a partial negative (δ−) and partial positive (δ+) end. A dipole arrow points from δ+ toward δ−, with a cross-hatch at the positive end.
Brønsted-Lowry acid
A proton (H⁺) donor. In simple terms, any species that can give away a hydrogen ion.
Brønsted-Lowry base
A proton (H⁺) acceptor. Any species with a lone pair or negative charge that can grab a hydrogen ion.
Conjugate acid / conjugate base
When an acid donates a proton it becomes its conjugate base; when a base accepts a proton it becomes its conjugate acid. Every acid-base reaction has two conjugate pairs.
pKa
A logarithmic measure of acid strength. Lower pKa = stronger acid. A difference of 1 pKa unit means a tenfold difference in acidity.
Lewis acid
An electron-pair acceptor (the electrophile in a reaction). Think of it as the electron-hungry species.
Lewis base
An electron-pair donor (the nucleophile in a reaction). Think of it as the electron-rich species.
Electrophile
A Lewis acid: it accepts an electron pair and is attracted to electron-rich sites. Literally "electron-loving."
Nucleophile
A Lewis base: it donates an electron pair and is attracted to electron-poor sites. Literally "nucleus-loving."
Lewis structures show every atom, every bond as a line, and every lone pair as dots. To draw one: count total valence electrons, connect atoms with single bonds, distribute remaining electrons as lone pairs to satisfy octets (duet for H), then convert lone pairs to double or triple bonds if needed.
Condensed structures compress the drawing into text. Atoms are grouped with their hydrogens: CH₃CH₂OH means a methyl group bonded to a CH₂ group bonded to OH. Branching is shown with parentheses: CH₃CH(CH₃)CH₃ for isobutane.
Bond-line (skeletal) structures are the default in organic chemistry. Each line segment is a bond, each vertex or endpoint is a carbon, and hydrogens on carbon are implicit (fill them in mentally to give each carbon four bonds). Heteroatoms (O, N, S, halogens) and their hydrogens are always written out.
Organic chemistry uses several arrows, and confusing them is a common early mistake.
Reaction arrow (→): "goes to" or "yields." Separates reactants from products.
Equilibrium arrows (⇌): the reaction goes in both directions. The longer arrow indicates the favoured direction.
Resonance arrow (↔): a double-headed arrow placed between resonance structures. It does not mean equilibrium; it means the structures are different ways of drawing the same molecule.
Curved arrow: shows the movement of an electron pair from a source (electron-rich) to a sink (electron-poor). The tail starts where the electrons are; the head points to where they go.
Dipole arrow: a small arrow on a bond pointing from δ+ to δ−, with a cross-hatch at the positive end. Shows the direction of bond polarity.
The tail of a curved arrow always starts at an electron source: a lone pair, a bond (sigma or pi), or a negative charge.
The head always points to an electron sink: an atom that can accept electrons (often one bearing a positive charge, or an electronegative atom).
Electrons flow from electron-rich to electron-poor. This is the single most important principle behind curved arrows.
A full curved arrow moves two electrons. A fishhook (half-headed) arrow moves one electron (relevant in radical reactions, covered later).
Move pi electrons or lone pairs using curved arrows. Atoms stay in the same position; only electrons move.
Every valid resonance structure must obey the rules of valence. Second-row elements (C, N, O, F) cannot exceed an octet.
The resonance hybrid (the real molecule) is a weighted blend of all contributors.
Rank resonance structures using these criteria, roughly in order of importance:
Complete octets on every atom (especially C, N, O) beat incomplete octets.
Fewer formal charges are better than more formal charges.
Negative formal charge should sit on the more electronegative atom.
Structures that preserve aromaticity (if applicable) are strongly favoured.
A structure that violates criterion 1 (e.g. a carbon with only six electrons) is always a minor contributor.
When comparing two structures, ask in order:
Same molecular formula? No → they are not isomers at all.
Same connectivity (same atoms bonded to the same atoms)? No → constitutional isomers.
Differ only in electron placement (pi bonds, lone pairs)? Yes → resonance structures (same molecule, different drawings).
Same connectivity but different 3D arrangement? → Stereoisomers (explored fully in Ch. 3 and beyond).
Identical in every respect? → Same structure (you may need to rotate or flip one to see it).
Real-world note: Resonance is why the peptide bond in proteins is planar, why benzene has equal C-C bond lengths, and why carboxylate ions are more stable than expected. Recognising resonance stabilisation is essential for predicting reactivity throughout the course.
Shorter bonds are stronger bonds. A triple bond (e.g. C≡C) is shorter and stronger than a double bond (C=C), which is shorter and stronger than a single bond (C-C).
Bond dissociation energy (BDE) measures the energy required to break one mole of a bond in the gas phase. Higher BDE = stronger bond.
Within the same row of the periodic table, bond strength increases as bond order increases. Down a group, bonds to larger atoms are longer and weaker (e.g. H-F > H-Cl > H-Br > H-I in strength).
A bond is polar when the two atoms differ in electronegativity. The electron density shifts toward the more electronegative atom.
Represent bond polarity with δ+ and δ− symbols on the atoms, or with a dipole arrow: the arrow points from the less electronegative atom (δ+) toward the more electronegative atom (δ−), with a small perpendicular cross at the positive end.
Common polar bonds in organic chemistry: C-O, C-N, C-F, C-Cl, O-H, N-H. The C-H bond is treated as essentially nonpolar for most purposes.
Bond polarity is the basis for understanding molecular polarity, intermolecular forces, and reactivity at electrophilic and nucleophilic sites.
Real-world note: Bond polarity explains why carbonyl groups (C=O) are such reactive sites in organic molecules. The carbon is electron-poor (δ+) and the oxygen is electron-rich (δ−), setting up a target for nucleophilic attack. This pattern drives the chemistry of aldehydes, ketones, carboxylic acids, and their derivatives throughout the course.
Identify the proton donor (acid) and proton acceptor (base).
Transfer one proton from acid to base. The acid loses H⁺ to become its conjugate base; the base gains H⁺ to become its conjugate acid.
Every Brønsted-Lowry acid-base reaction has exactly two conjugate pairs.
Example: CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺. Conjugate pair 1: CH₃COOH / CH₃COO⁻. Conjugate pair 2: H₂O / H₃O⁺.
The equilibrium favours the side with the weaker acid and weaker base (the more stable pair of products). In practice:
Compare pKa values: the proton moves from the stronger acid (lower pKa) to the stronger base, producing the weaker acid (higher pKa).
If pKa(reactant acid) < pKa(product acid), the equilibrium lies to the right.
A Keq > 1 means products are favoured; Keq < 1 means reactants are favoured.
When asked "why is compound A more acidic than compound B?", work through these factors:
Element effects (electronegativity and size): Within a row, acidity increases with electronegativity of the atom bearing the H (e.g. CH₄ < NH₃ < H₂O < HF). Down a group, acidity increases with size of the atom (e.g. HF < HCl < HBr < HI) because the conjugate base is more stable when the negative charge sits on a larger, more polarisable atom.
Resonance stabilisation of the conjugate base: If removing the proton gives a conjugate base whose negative charge is delocalised over multiple atoms, the acid is stronger. This is why carboxylic acids (pKa ~5) are far more acidic than alcohols (pKa ~16).
Inductive effects: Electron-withdrawing groups (e.g. halogens, CF₃, NO₂) near the acidic site stabilise the conjugate base by pulling electron density away from the negative charge, increasing acidity. The effect weakens with distance.
Hybridisation: An sp-hybridised carbon is more electronegative than sp² or sp³ because the s orbital holds electrons closer to the nucleus. Therefore: terminal alkynes (sp, pKa ~25) are more acidic than alkenes (sp², pKa ~44) or alkanes (sp³, pKa ~50).
Nucleophile (Lewis base): has a lone pair or pi electrons to donate. Look for negative charges, lone pairs on N, O, S, or electron-rich pi bonds.
Electrophile (Lewis acid): has an empty or partial-empty orbital, or a δ+ atom that can accept electrons. Look for positive charges, atoms bonded to electronegative groups, or carbons bearing a leaving group.
In any reaction scheme, the curved arrow starts at the nucleophile and points toward the electrophile. If you can identify which species is electron-rich and which is electron-poor, you can predict how they will react.
Draw the curved arrow from the nucleophile's electron pair to the electrophile.
Form the new bond at the arrowhead.
If the electrophile already has a full octet, a bond must break simultaneously (the electrons leave as a curved arrow pointing to the departing atom or group).
Check your product: conserve charge, conserve atoms, obey the octet rule for second-row elements.
This framework underpins every reaction mechanism in the course. Substitution, elimination, addition, and carbonyl reactions all begin with a nucleophile attacking an electrophile.
Real-world note: Drug design relies heavily on nucleophile-electrophile interactions. Many enzyme inhibitors work by presenting a nucleophilic group that attacks an electrophilic site in the enzyme's active site, forming a covalent bond that shuts the enzyme down.
Students often think the resonance arrow (↔) means the molecule flips back and forth between structures. It does not. The molecule exists as one hybrid all the time; the arrow simply connects equivalent drawings.
Students often confuse formal charge with oxidation state. Formal charge is an electron bookkeeping system; it does not tell you the real charge distribution. A nitrogen with a +1 formal charge is not necessarily electron-poor in the way a bare cation would be.
Students often assume that a lower pKa means a more negative number. Be careful: pKa = −10 is a stronger acid than pKa = 5, which is stronger than pKa = 16. The scale runs from very negative (strong acids) through zero and up to ~50+ (essentially non-acidic C-H bonds).
Students often draw curved arrows backward, starting at the electrophile and pointing to the nucleophile. Electrons always flow from rich to poor: tail at the source, head at the sink.
⚠️ Drawing correct Lewis structures with proper formal charges is tested relentlessly. If your formal charges are wrong, every downstream analysis (resonance, acidity, polarity) will be wrong too.
⚠️ Curved arrow problems appear on virtually every exam. You will be asked to draw arrows for resonance, for proton transfers, and for Lewis acid-base reactions.
⚠️ Ranking acidity using the four factors (element effects, resonance, induction, hybridisation) is a classic exam question. Be prepared to explain your reasoning, not just give the answer.
⚠️ "Are these two structures identical, resonance structures, constitutional isomers, or stereoisomers?" is a frequently tested format. Practise systematically: check formula, then connectivity, then electron placement, then 3D arrangement.
⚠️ Predicting the direction of proton transfer using pKa values is straightforward but easy to get backwards under exam pressure. Remember: equilibrium favours the weaker acid (higher pKa) on the product side.
True or false: A resonance arrow (↔) indicates an equilibrium between two species. (False, it connects different electron-placement drawings of the same species.)
Fill in the blank: Curved arrows always point from _______ to _______. (Electron-rich to electron-poor.)
True or false: A bond with a higher bond order is longer and weaker. (False, higher bond order means shorter and stronger.)
Fill in the blank: The pKa of HCl is about −7 and the pKa of acetic acid is about 4.75. The stronger acid is _______. (HCl, because it has the lower pKa.)
True or false: A nucleophile is an electron-pair acceptor. (False, a nucleophile is an electron-pair donor; an electrophile is the acceptor.)
Q: Draw the bond-line structure for CH₃CH(OH)CH₂CH₃. How many carbon atoms are present?
A: The bond-line structure is a four-carbon zigzag chain with an OH group on C2. Four carbon atoms are present.
Q: For the molecule CH₂O (formaldehyde), draw the Lewis structure and assign formal charges to every atom.
A: Carbon forms a double bond to oxygen and two single bonds to hydrogen. Carbon has 0 formal charge (4 − 0 − ½(8) = 0). Oxygen has 0 formal charge (6 − 4 − ½(4) = 0). Each hydrogen has 0 formal charge.
Q: Acetic acid (pKa 4.75) is mixed with sodium ethoxide (conjugate acid: ethanol, pKa 16). Which direction does the equilibrium favour, and why?
A: The equilibrium strongly favours products (ethanol and acetate). The proton transfers from the stronger acid (acetic acid, pKa 4.75) to the stronger base (ethoxide), producing the weaker acid (ethanol, pKa 16) and the weaker base (acetate).
Q: Rank the following in order of increasing acidity: ethane (C₂H₆), ethanol (C₂H₅OH), acetic acid (CH₃COOH). Explain your reasoning.
A: Ethane < ethanol < acetic acid. Ethane has no electronegative atom bearing the proton, so its conjugate base is extremely unstable (pKa ~50). Ethanol is more acidic (pKa ~16) because the proton sits on electronegative oxygen. Acetic acid is the most acidic (pKa ~4.75) because its conjugate base (acetate) is resonance-stabilised, delocalising the negative charge over two oxygen atoms.
Q: In the reaction of hydroxide ion (HO⁻) with bromomethane (CH₃Br), identify the nucleophile and the electrophile. Which way do the curved arrows point?
A: HO⁻ is the nucleophile (electron-rich, has lone pairs and a negative charge). The carbon in CH₃Br is the electrophile (electron-poor, bonded to the electronegative Br). The curved arrow starts at a lone pair on oxygen and points to the carbon of CH₃Br.
The structure-drawing skills from this chapter (especially bond-line structures and Lewis structures) are used in every subsequent chapter without re-explanation. Curved arrow notation returns immediately in Ch. 2 (for showing how orbitals overlap) and becomes the backbone of every reaction mechanism from Ch. 6 onward. Acid-base reasoning (pKa, conjugate stability) is central to understanding leaving group ability (Ch. 7), enolate chemistry, and carbonyl reactivity. The distinction between Brønsted-Lowry and Lewis acid-base frameworks runs through the entire course: Brønsted-Lowry for proton transfers, Lewis for all other electron-pair donation reactions.
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