Difficulty: Introductory to Intermediate | Prerequisites: Lewis structures, electronegativity, resonance basics
Acid-base chemistry runs through nearly every topic in Organic Chemistry I. It determines which protons are removed, which nucleophiles attack, and how reaction equilibria sit. The course moves through three progressively broader definitions of acids and bases (Arrhenius, Brønsted-Lowry, Lewis), and you are expected to apply the right one in context. If you are comfortable with Lewis structures and electronegativity, you have the tools to follow this material. If resonance is still shaky, revisit those notes first, because resonance stabilisation of conjugate bases is central to understanding acid strength.
Acids donate protons (or accept electron pairs); bases accept protons (or donate electron pairs). Three theories define what counts as an acid or base, each broader than the last. Acid and base strength is measured by how readily the transfer happens, quantified by pH and pKa.
Acid (Brønsted-Lowry)
A substance that donates a proton (H⁺) to another substance. In simple terms, it is the molecule that gives up a hydrogen ion in a reaction.
Base (Brønsted-Lowry)
A substance that accepts a proton (H⁺) from another substance. Think of it as the molecule with a lone pair ready to grab a hydrogen ion.
Lewis acid
An electron-pair acceptor. This is the broadest definition of an acid and covers species with no protons to donate, such as BF₃ or AlCl₃.
Lewis base
An electron-pair donor. Any species with an available lone pair (e.g. NH₃, OH⁻) qualifies.
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 donates a proton. Every acid has a conjugate base.
pH
A logarithmic scale (0 to 14) measuring the concentration of H⁺ ions in aqueous solution. Below 7 is acidic, 7 is neutral, above 7 is basic.
pKa
The negative log of the acid dissociation constant (Ka). A lower pKa means a stronger acid. This is the single most useful number for comparing acid strengths in organic chemistry.
Neutralisation reaction
A reaction between an acid and a base that produces a salt and (usually) water.
Arrhenius (narrowest): An acid increases H⁺ concentration in water; a base increases OH⁻ concentration. Limited to aqueous solutions.
Brønsted-Lowry (broader): An acid is a proton donor; a base is a proton acceptor. Works in any solvent and is the default framework for most organic chemistry problems.
Lewis (broadest): An acid is an electron-pair acceptor; a base is an electron-pair donor. Captures reactions with no proton transfer at all (e.g. BF₃ accepting a lone pair from NH₃).
Every Arrhenius acid is a Brønsted-Lowry acid, and every Brønsted-Lowry acid is a Lewis acid. The reverse is not true.
A strong acid donates its proton readily and essentially completely (e.g. HCl, H₂SO₄). A weak acid donates reluctantly and reaches equilibrium with significant undissociated acid remaining (e.g. acetic acid).
Strength is quantified by Ka (the acid dissociation constant) and, more usefully, pKa = −log(Ka). A lower pKa means a stronger acid.
Four factors govern acid strength in organic chemistry:
Element effects (bond strength and electronegativity down a group or across a period)
Inductive effects (electron-withdrawing groups stabilise the conjugate base)
Resonance effects (delocalisation of the negative charge in the conjugate base)
Hybridisation effects (more s-character in the orbital holding the lone pair stabilises the conjugate base)
pH = −log[H⁺]. The scale runs from 0 to 14 in aqueous solution.
pH < 7 is acidic; pH = 7 is neutral; pH > 7 is basic.
Each whole-number step represents a tenfold change in H⁺ concentration.
Neutralisation: Acid + base → salt + water. The classic example is HCl + NaOH → NaCl + H₂O.
Proton-transfer (Brønsted-Lowry): A proton moves from the acid to the base, forming a conjugate acid and a conjugate base. Equilibrium favours the side with the weaker acid and weaker base.
Lewis acid-base reactions: An electron pair is donated from the Lewis base to the Lewis acid, forming a new covalent bond. No proton need be involved.
Acid-base chemistry governs the pH balance in your blood (buffered near 7.4 by the carbonic acid/bicarbonate system) and the digestion of food in your stomach (hydrochloric acid at roughly pH 1.5 to 2). In industry, acids and bases are essential for manufacturing fertilisers (sulfuric acid in the Haber process chain), pharmaceuticals (controlling drug protonation states to improve absorption), and dyes.
Students often assume that "strong acid" means "concentrated acid." Strength (how completely it dissociates) and concentration (how much is dissolved) are independent properties. Dilute HCl is still a strong acid.
Students confuse the three acid-base theories and apply the wrong one. In organic chemistry, Brønsted-Lowry is the default unless the question explicitly involves a non-proton-transfer reaction (Lewis).
Students sometimes think a low pH means the substance is dangerous. pH measures H⁺ concentration, not toxicity. Orange juice (pH ~3.5) is acidic but perfectly safe to drink.
Students forget that equilibrium in a proton-transfer reaction favours the side with the weaker acid and weaker base. Comparing pKa values on each side of the equation tells you the direction.
⚠️ Expect to be given two compounds and asked to predict which is the stronger acid, using pKa reasoning (element effects, inductive effects, resonance, hybridisation).
⚠️ You will be asked to identify conjugate acid-base pairs in a reaction and to predict which direction the equilibrium favours.
⚠️ Lewis acid-base interactions appear in mechanism questions throughout the course. If a nucleophile attacks an electrophile, that is a Lewis base donating to a Lewis acid.
⚠️ Know the approximate pKa ranges for common functional groups (e.g. carboxylic acids ~4 to 5, alcohols ~16, terminal alkynes ~25). Exam questions often require you to compare these from memory.
True or false: Every Brønsted-Lowry acid is also a Lewis acid. (True.)
Fill in the blank: A substance with a pKa of 4 is a ______ acid than one with a pKa of 10. (Stronger.)
True or false: pH 3 is ten times more acidic than pH 4. (True – each unit is a tenfold change in H⁺ concentration.)
Fill in the blank: When a base accepts a proton, it becomes the ______ acid of that base. (Conjugate.)
True or false: The Arrhenius definition of acids and bases applies in all solvents. (False – it is limited to aqueous solutions.)
Q: Acetic acid (CH₃COOH, pKa ~4.75) is mixed with sodium ethoxide (NaOCH₂CH₃), the conjugate base of ethanol (pKa ~16). Predict the direction of equilibrium and explain why.
A: Equilibrium favours the products (acetate ion + ethanol). The reaction moves the proton from the stronger acid (acetic acid, pKa 4.75) to form the weaker acid (ethanol, pKa 16). Equilibrium always favours the weaker acid side.
Q: Explain the difference between a Brønsted-Lowry acid and a Lewis acid, and give an example of a Lewis acid that is not a Brønsted-Lowry acid.
A: A Brønsted-Lowry acid donates a proton; a Lewis acid accepts an electron pair. BF₃ is a Lewis acid (it accepts a lone pair into its empty p-orbital) but has no proton to donate, so it is not a Brønsted-Lowry acid.
Q: Why is the carboxylate ion (RCOO⁻) a weaker base than an alkoxide ion (RO⁻)?
A: The negative charge in the carboxylate ion is delocalised by resonance across two oxygen atoms, stabilising it. The alkoxide ion has the charge localised on a single oxygen. Greater stabilisation of the conjugate base means the parent acid (carboxylic acid) is stronger, and its conjugate base is correspondingly weaker.
Q: Classify the reaction of BF₃ with NH₃ under each acid-base theory that applies.
A: It is a Lewis acid-base reaction: BF₃ (Lewis acid, electron-pair acceptor) forms a coordinate bond with NH₃ (Lewis base, electron-pair donor). It is not a Brønsted-Lowry reaction because no proton is transferred, and it is not an Arrhenius reaction because it need not occur in aqueous solution.
Acid-base chemistry ties directly to resonance: the ability of a conjugate base to delocalise its charge through resonance is one of the strongest predictors of acid strength. This is why carboxylic acids are so much more acidic than alcohols.
It also connects to nucleophilicity and electrophilicity. In organic mechanisms, a nucleophile is a Lewis base and an electrophile is a Lewis acid. Every substitution and addition mechanism you learn later is, at its core, a Lewis acid-base reaction.
Understanding pKa values will be essential when you reach elimination vs. substitution decisions, because the choice of base (strong vs. weak, bulky vs. compact) drives the reaction pathway.
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