Difficulty: Intermediate | Prerequisites: Lewis structures, formal charges, resonance, electronegativity.
Acid-base chemistry is the single most important framework in organic chemistry. Nearly every reaction mechanism you will study involves a proton transfer, a Lewis acid-base interaction, or both. This material teaches you to predict which way a proton transfer will go, compare the strength of acids and bases using structural reasoning, and recognise Lewis acid-base reactions where a lone pair donor meets an electron-pair acceptor. If you can look at two molecules and say which is the stronger acid (and explain why), you have the foundation for reaction prediction.
Stronger acids have more stable conjugate bases. Conjugate base stability depends on electronegativity, resonance delocalisation, inductive effects, and atom size. In Lewis acid-base reactions, the base donates an electron pair to the acid. Equilibrium favours the side with the weaker acid and weaker base.
Brønsted-Lowry acid
A proton (H⁺) donor. In simple terms, the molecule that gives up a hydrogen ion.
Brønsted-Lowry base
A proton (H⁺) acceptor. The molecule that takes a hydrogen ion using a lone pair.
Conjugate base
What remains after an acid loses a proton. The stability of the conjugate base determines how strong the acid is.
Conjugate acid
What forms after a base gains a proton.
Lewis acid
An electron-pair acceptor. It has an empty or electron-deficient orbital that can accept a lone pair. Think of it as the molecule with the "hole" that needs electrons.
Lewis base
An electron-pair donor. It has a lone pair available to share. Think of it as the molecule with electrons to give.
Lewis acid-base adduct
The product formed when a Lewis base donates its electron pair to a Lewis acid, creating a new covalent bond.
Inductive effect
The pull or push of electron density through sigma bonds caused by electronegative or electropositive atoms. Electronegative atoms nearby stabilise a negative charge by pulling electron density toward themselves through the chain.
Resonance stabilisation (of a conjugate base)
When the negative charge on a conjugate base can be delocalised across multiple atoms through resonance, the base is more stable and the parent acid is stronger.
The central principle: a stronger acid has a more stable conjugate base. To compare two acids, remove a proton from each and ask which conjugate base is more stable.
Factors that stabilise a conjugate base (in rough order of importance for this course):
Resonance delocalisation. If the negative charge can spread over multiple atoms through resonance, the conjugate base is more stable. Example from the exam: a carboxylic acid (with C=O adjacent to OH) is stronger than an alcohol (OH without adjacent C=O), because the carboxylate conjugate base delocalises the charge over two oxygens.
Electronegativity of the atom bearing the charge. A negative charge on oxygen is more stable than on nitrogen, which is more stable than on carbon.
Inductive effects. Electronegative atoms near the acidic site pull electron density away, stabilising the negative charge. The closer the electronegative atom, the stronger the effect. Example from the exam: between two diols, the one with the OH groups closer together is a stronger acid because the nearby oxygen stabilises the conjugate base through induction.
Charge on the acid. A positively charged acid (like H₃O⁺) is stronger than its neutral counterpart (H₂O), because losing a proton relieves a positive charge.
A stronger base has a less stable conjugate acid (or equivalently, it holds onto its lone pair less tightly).
Example from the exam: comparing RO⁻ (alkoxide) and RS⁻ (thiolate), the alkoxide (oxygen-based) is the stronger base. Sulfur is larger and more polarisable, so it stabilises a negative charge better, making RS⁻ a weaker base.
Equilibrium favours the formation of the weaker acid and weaker base (the more stable side).
Identify the acid and base on each side. The stronger acid reacts with the stronger base to form the weaker acid and weaker base.
Example from the exam: an alkoxide (strong base) reacts with a protonated amine (acid) to form an alcohol (weaker acid) and a neutral amine (weaker base). The reaction proceeds to the right.
The Lewis base donates a lone pair to the Lewis acid. A new covalent bond forms.
The Lewis acid is the species with an empty or electron-poor orbital (often a metal or a carbocation). The Lewis base is the species with a lone pair to give.
Example from the exam: AlCl₃ (Lewis acid, electron-deficient aluminium) reacts with formaldehyde H₂C=O (Lewis base, oxygen donates a lone pair). The oxygen lone pair forms a bond to aluminium, giving aluminium a formal negative charge and oxygen a formal positive charge in the adduct.
Another example: a carbocation CH₃⁺ (Lewis acid) reacts with water H₂O (Lewis base). The oxygen lone pair bonds to the carbon, forming a protonated ether with a positive charge on oxygen.
Students often confuse Brønsted-Lowry and Lewis definitions. Every Brønsted-Lowry base is also a Lewis base (both donate electrons), but a Lewis acid need not involve protons at all.
A common mistake is reasoning about acid strength by looking at the acid itself rather than its conjugate base. Always remove the proton first, then assess stability.
Students sometimes assume that inductive effects are stronger than resonance. In most organic acid-base comparisons, resonance is the more powerful stabilising factor. Induction matters when resonance is equal or absent.
When drawing Lewis acid-base reactions, students sometimes draw the arrow going from the acid to the base. The arrow always starts from the lone pair (the base) and points to the acid.
⚠️ Acid-strength comparison is the highest-point topic on the exam (19 points in this exam). Expect multiple comparison questions.
⚠️ Be able to explain your reasoning, not just pick the answer. "Conjugate base of B is stabilised by resonance" is the kind of sentence the grader wants.
⚠️ Lewis acid-base reactions require you to draw the adduct with correct formal charges and electron flow arrows. Practise with AlCl₃, BF₃, and carbocations.
⚠️ Know the equilibrium rule: the reaction always favours the weaker acid/base pair. Draw both sides and compare.
True or false: a positively charged species (like H₃O⁺) is a stronger acid than its neutral form (H₂O). (True.)
Fill in the blank: equilibrium in an acid-base reaction favours the ____ acid and ____ base. (weaker; weaker)
True or false: RO⁻ is a stronger base than RS⁻. (True. Oxygen is smaller and less able to stabilise the charge than sulfur.)
Fill in the blank: in a Lewis acid-base reaction, the electron-flow arrow starts from the ____ and points to the ____. (Lewis base / lone pair; Lewis acid)
Q: Which is a stronger acid: a neutral alcohol (R-OH) or a protonated alcohol (R-OH₂⁺)?
A: The protonated alcohol. Losing a proton from a positive species is easier because it relieves a positive charge, producing a neutral molecule.
Q: Which is the stronger base: an alkoxide (RO⁻) or a thiolate (RS⁻)?
A: The alkoxide. Sulfur is larger and more polarisable, which stabilises the negative charge, making the thiolate a weaker (more stable) base.
Q: Compound A is a carboxylic acid (C=O next to OH). Compound B is an alcohol (plain OH, no adjacent C=O). Which is the stronger acid, and why?
A: Compound B (the carboxylic acid) is the stronger acid. Its conjugate base is stabilised by resonance: the negative charge is delocalised over two oxygen atoms. The alcohol's conjugate base has no such resonance stabilisation.
Q: AlCl₃ reacts with H₂C=O. Identify the Lewis acid and Lewis base, and describe the adduct.
A: AlCl₃ is the Lewis acid (aluminium is electron-deficient). Formaldehyde (H₂C=O) is the Lewis base (oxygen donates a lone pair). In the adduct, oxygen is bonded to aluminium with a formal positive charge on oxygen and a formal negative charge on aluminium.
Acid-base reasoning underpins every reaction mechanism in the rest of the course. Nucleophilicity (how reactive a base is in attacking carbon) builds directly on base strength. The stability arguments you learn here (resonance, induction, electronegativity) apply identically when predicting carbocation stability, leaving-group ability, and enolate chemistry.
Brønsted-Lowry acid, Brønsted-Lowry base, Lewis acid, Lewis base, conjugate acid, conjugate base, proton transfer, acid strength, base strength, pKa, resonance stabilisation, inductive effect, electronegativity, atom size, charge stabilisation, alkoxide, thiolate, carboxylic acid, alcohol acidity, Lewis adduct, electron flow, curved arrows, equilibrium, AlCl3, carbocation, protonated alcohol