Difficulty: Intermediate | Prerequisites: Structure, Bonding, and Electron Movement notes (Doc 1 of this series)
This set of notes covers the chemical behaviour that emerges from the structural foundations in the first document. Acid-base chemistry is the single most important reaction type in organic chemistry: nearly every mechanism begins with a proton transfer or a nucleophilic attack. Molecular geometry and intermolecular forces explain why molecules behave differently in solution, in the body, and in the lab. If you are comfortable with Lewis structures, formal charge, and curved arrows, you are ready for this material.
Acid-base reactions in organic chemistry come in two flavours: Brønsted-Lowry (proton transfer) and Lewis (electron-pair donation). Knowing what makes an acid strong or a base good at donating electrons lets you predict which way a reaction goes. Molecular shape, hybridisation, and polarity determine a molecule's physical properties and how it interacts with its neighbours.
Brønsted-Lowry acid
A species that donates a proton (H+) to another species. In simple terms, any molecule that gives away a hydrogen ion.
Brønsted-Lowry base
A species that accepts a proton. Think of it as: any molecule with a lone pair willing to grab a hydrogen ion.
Conjugate acid
The species formed when a base accepts a proton. Every base has a conjugate acid.
Conjugate base
The species formed when an acid loses a proton. Every acid has a conjugate base.
pKa
A logarithmic measure of acid strength. Lower pKa = stronger acid. A difference of one pKa unit means a tenfold difference in acid strength.
Lewis acid
An electron-pair acceptor. In organic chemistry, Lewis acids are electrophiles.
Lewis base
An electron-pair donor. In organic chemistry, Lewis bases are nucleophiles.
Electrophile
An electron-deficient species that accepts an electron pair from a nucleophile. Think of it as: a molecule with a "hole" where electrons are welcome.
Nucleophile
An electron-rich species that donates an electron pair to an electrophile. In simple terms, the electron-giving partner in a reaction.
VSEPR theory (Valence Shell Electron Pair Repulsion)
A model that predicts molecular geometry by assuming electron groups around a central atom repel each other and arrange themselves to be as far apart as possible.
Hybridisation
The mixing of atomic orbitals to form new hybrid orbitals (sp, sp², sp³) that better explain observed molecular geometries and bond angles.
Intermolecular forces (IMFs)
Attractive forces between molecules (not within them). They determine physical properties like boiling point, melting point, and solubility.
Hydrogen bonding
A strong type of dipole-dipole interaction occurring when hydrogen is bonded to oxygen, nitrogen, or fluorine and interacts with a lone pair on a nearby O, N, or F.
Dispersion forces (London dispersion forces / van der Waals forces)
Weak, temporary attractions present in all molecules, arising from momentary fluctuations in electron distribution. Dominant in nonpolar compounds.
Dipole-dipole interactions
Attractive forces between the positive end of one polar molecule and the negative end of another.
Oxidation (in organic chemistry)
An increase in the number of bonds to oxygen (or other electronegative atoms) or a decrease in bonds to hydrogen.
Reduction (in organic chemistry)
A decrease in the number of bonds to oxygen or an increase in bonds to hydrogen. The reverse of oxidation.
A Brønsted-Lowry acid-base reaction is a proton transfer. The acid donates H+ to the base, producing a conjugate base and a conjugate acid.
The equilibrium favours the side with the weaker acid and weaker base. To predict which way a reaction goes, compare pKa values: the proton moves from the stronger acid (lower pKa) to the stronger base.
Four factors that affect acid strength:
Electronegativity: more electronegative atoms stabilise the conjugate base better, making the parent acid stronger. HF is a stronger acid than HCH3 because fluorine holds the negative charge more readily than carbon.
Resonance stabilisation: if the conjugate base can spread its negative charge across multiple atoms through resonance, the acid is stronger. Acetic acid is stronger than ethanol because the acetate ion delocalises charge over two oxygens.
Inductive effects: electron-withdrawing groups near the acidic site stabilise the conjugate base by pulling electron density away. Trichloroacetic acid is stronger than acetic acid because three chlorines withdraw electron density.
Hybridisation: an sp-hybridised carbon holds a negative charge better than sp², which is better than sp³. The greater s-character means electrons are held closer to the nucleus. This is why terminal alkynes (sp C-H) are more acidic than alkenes (sp² C-H) or alkanes (sp³ C-H).
The Lewis definition is broader: a Lewis acid accepts an electron pair; a Lewis base donates one. In organic chemistry, Lewis acids are electrophiles and Lewis bases are nucleophiles.
Curved arrows show electrons flowing from the nucleophile (the Lewis base, electron-rich) to the electrophile (the Lewis acid, electron-poor). This framework lets you predict products: identify the nucleophilic site, identify the electrophilic site, draw the arrow from one to the other, and work out the product.
VSEPR theory predicts geometry by counting electron groups (bonds and lone pairs) around a central atom:
Electron groups | Electron geometry | Molecular geometry (no lone pairs) | Bond angle | Hybridisation |
|---|---|---|---|---|
2 | Linear | Linear | 180° | sp |
3 | Trigonal planar | Trigonal planar | 120° | sp² |
4 | Tetrahedral | Tetrahedral | 109.5° | sp³ |
Lone pairs occupy space but are not "seen" in the molecular geometry. Four electron groups with one lone pair gives a trigonal pyramidal shape (e.g. ammonia). Four electron groups with two lone pairs gives a bent shape (e.g. water).
Polarity depends on two things: polar bonds and molecular symmetry. A molecule can have polar bonds but still be nonpolar overall if the dipoles cancel by symmetry (e.g. CO2). Polar molecules experience dipole-dipole interactions and, if they have O-H, N-H, or F-H bonds, hydrogen bonding.
Alkane names are built from the number of carbon atoms in the longest continuous chain: methane (1), ethane (2), propane (3), butane (4), pentane (5), hexane (6), and so on. Substituents are named as prefixes with their position number.
Cycloalkanes add the prefix "cyclo-" to the parent name: cyclopropane (3-membered ring), cyclobutane (4), cyclopentane (5), cyclohexane (6).
Being able to convert freely between a name and a structure (in both directions) is a basic skill tested early and relied on throughout the course.
In organic chemistry, oxidation and reduction are tracked by changes in bonds to oxygen and hydrogen, not by electron counting as in general chemistry:
Oxidation: more bonds to oxygen (or fewer bonds to hydrogen). Example: alcohol → aldehyde → carboxylic acid.
Reduction: more bonds to hydrogen (or fewer bonds to oxygen). Example: ketone → alcohol.
Recognising whether a transformation is an oxidation or reduction tells you what class of reagent is needed (an oxidising agent or a reducing agent).
Three main types of intermolecular forces, in order of increasing strength:
Dispersion forces: present in every molecule. Strength increases with molecular surface area and molecular weight. Dominant in nonpolar molecules like hexane.
Dipole-dipole interactions: occur between polar molecules. Stronger than dispersion forces alone for molecules of similar size.
Hydrogen bonding: the strongest common IMF. Requires H bonded to O, N, or F interacting with a lone pair on another O, N, or F. This is why water and alcohols have high boiling points relative to their molecular weight.
Stronger intermolecular forces lead to higher boiling points, higher melting points, and greater solubility in polar solvents (like dissolves like).
Acid-base chemistry is central to how your body processes drugs. Most medications are weak acids or weak bases, and their ionisation state at physiological pH (7.4) determines whether they can cross cell membranes. Aspirin, for example, is absorbed in the stomach (low pH) where it remains protonated and nonpolar enough to pass through the stomach lining.
Intermolecular forces explain everyday observations: why oil and water do not mix (nonpolar dispersion forces vs. polar hydrogen bonding), why rubbing alcohol evaporates faster than water (weaker hydrogen bonding), and why geckos can walk on ceilings (dispersion forces between their toe pad surfaces and the wall).
Students often confuse Brønsted-Lowry and Lewis definitions. Every Brønsted-Lowry acid-base reaction is also a Lewis acid-base reaction, but not the other way around. BF3 accepting a lone pair from NH3 is a Lewis acid-base reaction with no proton transfer at all.
A frequent error is assuming that a larger pKa means a stronger acid. The opposite is true: lower pKa = stronger acid.
Students sometimes think hydrogen bonds are a type of covalent bond. They are not. Hydrogen bonds are intermolecular attractions, much weaker than covalent bonds.
Many students forget that molecular polarity depends on both bond polarity and geometry. CO2 has polar C=O bonds, but the molecule is nonpolar because the two dipoles cancel.
⚠️ Predicting the direction of an acid-base equilibrium using pKa values is one of the most commonly tested skills. You will be given two species and asked which way the proton transfers.
⚠️ Ranking acid strength using the four factors (electronegativity, resonance, induction, hybridisation) appears in virtually every exam. Memorise the factors and be able to apply them to unfamiliar molecules.
⚠️ VSEPR geometry questions ask you to identify hybridisation, bond angles, and molecular shape for a given atom in a structure. These are fast marks if you know the table.
⚠️ Intermolecular force questions often ask you to rank boiling points of a set of compounds. The approach: identify which IMFs are present, then rank. Hydrogen bonding > dipole-dipole > dispersion (all else being equal).
⚠️ Oxidation state changes are commonly tested by asking whether a given transformation is an oxidation, a reduction, or neither.
True or False: The equilibrium in an acid-base reaction favours the formation of the stronger acid and base. (Answer: False. It favours the weaker acid and weaker base.)
Fill in the blank: A Lewis acid is an electron-pair ______. (Answer: acceptor.)
True or False: An sp²-hybridised carbon has bond angles of approximately 109.5°. (Answer: False. sp² gives approximately 120°; 109.5° is sp³.)
Fill in the blank: Hydrogen bonding requires H bonded to ______, ______, or ______. (Answer: oxygen, nitrogen, or fluorine.)
True or False: Converting an alcohol to an aldehyde is a reduction. (Answer: False. It is an oxidation.)
Q: Acetic acid (pKa 4.75) is mixed with sodium ethoxide (the conjugate base of ethanol, pKa 16). Which direction does the equilibrium favour, and why?
A: The proton transfers from acetic acid to ethoxide, forming acetate and ethanol. The equilibrium favours the products (right side) because it produces the weaker acid (ethanol, pKa 16) and the weaker base (acetate).
Q: Explain why acetic acid (CH3COOH) is a stronger acid than ethanol (CH3CH2OH).
A: The conjugate base of acetic acid (acetate) is stabilised by resonance: the negative charge is delocalised over two equivalent oxygen atoms. The conjugate base of ethanol (ethoxide) has no such resonance stabilisation, so the negative charge is localised on one oxygen. Better conjugate base stability means stronger acid.
Q: Identify the electrophile and nucleophile in the reaction of BF3 with NH3.
A: BF3 is the electrophile (Lewis acid): boron has an empty p orbital and accepts an electron pair. NH3 is the nucleophile (Lewis base): nitrogen donates its lone pair to boron.
Q: Predict the molecular geometry around the oxygen atom in water (H2O).
A: Oxygen has four electron groups (two bonding pairs, two lone pairs), giving a tetrahedral electron geometry. With two lone pairs hidden, the molecular geometry is bent, with a bond angle slightly less than 109.5° (approximately 104.5°).
Q: Rank the following in order of increasing boiling point: CH4, CH3OH, CH3CH2CH2CH3.
A: CH4 (lowest, small nonpolar, only dispersion forces) < CH3CH2CH2CH3 (nonpolar but larger, stronger dispersion) < CH3OH (highest, hydrogen bonding in addition to dispersion and dipole-dipole).
Acid-base concepts connect directly to substitution and elimination reactions (SN1, SN2, E1, E2), where nucleophilicity and basicity determine which pathway dominates. The stronger the base, the more likely elimination; the better the nucleophile, the more likely substitution.
Intermolecular forces are the bridge between molecular structure and the physical behaviour you see in the lab: why one solvent dissolves a compound and another does not, why distillation separates mixtures at specific temperatures.
Hybridisation and geometry come back in stereochemistry (Chapter 3), where the tetrahedral arrangement of sp³ carbons creates chirality.
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