Acids donate protons (Brønsted) or accept electrons (Lewis); bases do the opposite. pKa tells you how strong an acid is, with lower numbers meaning stronger acids. In any acid-base reaction, equilibrium favours the side with the weaker acid and weaker base, and you can predict which way a reaction goes just by comparing pKa values.
Difficulty: Introductory to Intermediate | Prerequisites: General chemistry (atomic structure, electronegativity, bonding, Lewis dot structures). If you are not comfortable drawing lone pairs and identifying bond polarity, review those topics first.
This topic sits right at the start of Organic Chemistry I and underpins nearly every reaction mechanism you will see for the rest of the course. Organic reactions are, at their core, about electrons moving between atoms, and the language of acids and bases gives you the framework to describe that movement. You need two models here: Brønsted-Lowry (proton transfer) and Lewis (electron-pair transfer). Brønsted tells you who gives up or accepts a proton; Lewis broadens the picture to any electron-pair donor or acceptor, which is how you will later understand nucleophiles and electrophiles. If you can read a pKa table and predict which way an acid-base equilibrium lies, you have the foundation for substitution, elimination, and addition reactions.
Brønsted-Lowry acid
A compound that can donate a proton (H⁺) from an atom–H bond. In simple terms, any molecule with a hydrogen it can let go of.
Brønsted-Lowry base
A compound with a lone pair on an atom that can accept a proton (H⁺). Think of it as: anything with a spare pair of electrons ready to grab a proton.
Conjugate base
The species that remains after an acid has donated its proton. In simple terms, it is what the acid turns into once it loses H⁺.
Conjugate acid
The species formed when a base accepts a proton. Think of it as: the base after it has picked up an H⁺.
Conjugate acid-base pair
Two species that differ by exactly one proton. Every acid-base reaction produces two conjugate pairs.
pKa
The negative log of the acid dissociation constant (Ka). It is a quantitative measure of acid strength: the smaller the pKa, the stronger the acid. Derived from Ka = [H⁺][A⁻] / [HA]. In simple terms, pKa is a number on a scale: low numbers mean the acid gives up its proton easily; high numbers mean it holds on tightly.
Ka (acid dissociation constant)
The equilibrium constant for the dissociation of an acid in water: Ka = [H⁺][A⁻] / [HA]. A larger Ka means a stronger acid.
Lewis acid
An electron-pair acceptor. It has an empty or electron-deficient orbital that can receive a pair of electrons. Think of it as: a molecule that is hungry for electrons.
Lewis base
An electron-pair donor. It has a lone pair (or pi electrons) available to share. Think of it as: a molecule with electrons to spare.
Electron flow (curved-arrow notation)
The convention of drawing curved arrows from an electron source (lone pair or bond) to an electron sink (electrophilic atom or bond). Arrows always flow from electron-rich to electron-poor.
Pi bond as a base
A pi bond (C=C) can act as a Brønsted base because its electron density is available to accept a proton, shown through the resonance form C=C ↔ ⁺C–C⁻.
How a Brønsted acid-base reaction works
The acid (H–A) has a hydrogen bonded to atom A. The base (:B) has a lone pair on atom B.
In the reaction, the base donates its lone pair to grab the proton from the acid: H–A + :B ⇌ A⁻ + H–B⁺
The atom A in the acid gains a negative charge (−1) because it keeps the bonding electrons when H⁺ leaves.
The atom B in the base gains a positive charge (+1) because it has formed a new bond to H⁺.
Conjugate pairs form in every reaction
The acid H–A loses its proton and becomes the conjugate base A⁻.
The base :B accepts the proton and becomes the conjugate acid H–B⁺.
Every Brønsted reaction therefore produces two conjugate pairs: (H–A / A⁻) and (:B / H–B⁺).
Three ways to write the same idea
Acid dissociation alone: H–A ⇌ H⁺ + A⁻
Base picking up a proton: B: + H⁺ ⇌ B⁺–H
Full proton transfer: A–H + :B ⇌ :A⁻ + H–B⁺
Pi bonds can act as bases
A carbon-carbon double bond (C=C) has accessible electron density in its pi system.
That electron density can accept a proton, making the pi bond a Brønsted base.
This is shown through the resonance form: C=C ↔ ⁺C–C⁻
This concept becomes central when you study electrophilic addition reactions later in the course.
pKa as a measure of acid strength
pKa = –log(Ka), where Ka = [H⁺][A⁻] / [HA]
Smaller pKa = stronger acid. The acid gives up its proton more readily.
A pKa below about 3 indicates a strong acid in organic chemistry terms.
A pKa above about 15 indicates a very weak acid (it holds onto its proton tightly).
Using pKa to measure base strength
Base strength is measured indirectly, through the pKa of the base's conjugate acid.
A base whose conjugate acid has a larger pKa is a stronger base. The logic: if the conjugate acid is weak (high pKa, reluctant to donate H⁺), the base itself must be strong (eager to accept H⁺).
A conjugate acid pKa above about 11 indicates a strong base.
The inverse relationship
A weak acid produces a strong conjugate base.
A strong acid produces a weak conjugate base.
This inverse relationship is one of the most tested principles in organic chemistry exams. If you remember one rule from this section, make it this one.
Predicting equilibrium direction
Compare the pKa values of the two acids on each side of the equation.
Equilibrium favours the formation of the weaker acid (the one with the larger pKa).
In other words: the reaction proceeds in the direction that produces the more stable, less reactive conjugate pair.
Worked example
Consider: HA + A₂⁻ ⇌ HA₂ + A⁻
If HA₂ is the weaker acid (larger pKa) than HA, equilibrium lies to the right, because the products are more stable.
The proton transfers from the stronger acid to the stronger base.
Calculating Keq from pKa
The equilibrium constant (Keq) for an acid-base reaction can be estimated from the difference in pKa values of the two acids involved.
Keq = 10^(ΔpKa), where ΔpKa = pKa(product acid) – pKa(reactant acid).
A large positive ΔpKa means equilibrium strongly favours the products.
A negative ΔpKa means equilibrium favours the reactants.
Acid dissociation constant: Ka = [H⁺][A⁻] / [HA]
pKa definition: pKa = –log₁₀(Ka)
Relationship: a smaller pKa corresponds to a larger Ka and a stronger acid.
Equilibrium constant from pKa difference: Keq = 10^(pKa(product acid) – pKa(reactant acid))
General Brønsted acid-base reaction: H–A + :B ⇌ A⁻ + H–B⁺
Electron flow convention: curved arrows always point from electron-rich (lone pair or bond) to electron-poor (the atom accepting electrons).
Blood buffering relies on conjugate acid-base pairs (carbonic acid / bicarbonate) to keep blood pH stable. The same pKa logic you learn here explains why your blood does not swing wildly in pH after a meal.
Drug design depends on pKa: a drug molecule's pKa determines whether it will be ionised or neutral at physiological pH, which controls whether it can cross cell membranes. Pharmacologists use pKa tables the same way you use them on exams.
Household chemistry: vinegar (acetic acid, pKa ~4.75) is a weak acid; hydrochloric acid in your stomach (pKa ~−7) is a strong one. The difference in pKa is why one stings and the other dissolves food.
Students often think a higher pKa means a stronger acid. It is the opposite: lower pKa = stronger acid. Drill this until it is automatic.
Students confuse conjugate base strength with acid strength. A strong acid has a weak conjugate base. If you find yourself saying "strong acid, strong conjugate base," stop and reverse it.
Students sometimes draw curved arrows pointing from the proton to the base. Arrows represent electron flow, and electrons move from the base (electron-rich) toward the proton (electron-poor), not the other way around.
Students assume Lewis acids and Brønsted acids are the same thing. All Brønsted acids can be described as Lewis acids, but Lewis acids include species (such as BF₃ or AlCl₃) that have no proton to donate. The Lewis definition is broader.
⚠️ Predicting equilibrium direction from pKa values is one of the most commonly tested skills in Organic Chemistry I. Expect to see two acids, their pKa values, and a question asking which side of the equilibrium is favoured.
⚠️ Identifying conjugate acid-base pairs in a reaction. You will be given a reaction and asked to label every species. Practise this with at least ten different reactions.
⚠️ Drawing correct curved arrows for proton transfer. Arrows from the lone pair on the base to the H, and from the H–A bond to atom A. Getting the direction wrong is a common mark loss.
⚠️ Distinguishing Brønsted and Lewis acid-base behaviour. Exam questions may give you a reaction and ask whether it is a Brønsted or Lewis acid-base reaction (or both).
⚠️ The inverse relationship between acid strength and conjugate base strength appears on nearly every first exam in this course.
True or False: A pKa of 2 indicates a weaker acid than a pKa of 10. (False. Lower pKa = stronger acid.)
Fill in the blank: When HCl donates a proton to water, Cl⁻ is the ________ of HCl. (conjugate base)
True or False: Equilibrium in an acid-base reaction favours the side with the stronger acid. (False. It favours the side with the weaker acid and weaker base.)
Fill in the blank: A Lewis acid is an electron-pair ________. (acceptor)
True or False: Curved arrows in a proton-transfer mechanism point from the base to the acid. (True. From the lone pair on the base toward the proton on the acid.)
Q: Acetic acid has a pKa of 4.75 and ethanol has a pKa of 16. If acetate ion (CH₃COO⁻) is mixed with ethanol, does a proton transfer occur? Which direction does equilibrium favour?
A: No significant proton transfer occurs. Ethanol (pKa 16) is a much weaker acid than acetic acid (pKa 4.75), so equilibrium strongly favours the reactants. Acetate is too weak a base to deprotonate ethanol.
Q: Identify the conjugate acid-base pairs in the following reaction: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
A: Pair 1: NH₃ (base) and NH₄⁺ (conjugate acid). Pair 2: H₂O (acid) and OH⁻ (conjugate base).
Q: BF₃ reacts with NH₃ to form F₃B–NH₃. Classify each reactant as a Lewis acid or Lewis base, and explain why this is not a Brønsted acid-base reaction.
A: BF₃ is the Lewis acid (electron-pair acceptor; boron has an empty p orbital). NH₃ is the Lewis base (electron-pair donor; nitrogen has a lone pair). No proton is transferred, so this is a Lewis acid-base reaction only, not Brønsted.
Q: Given two acids, HA (pKa = 5) and HB (pKa = 12), estimate Keq for the reaction HA + B⁻ ⇌ A⁻ + HB.
A: ΔpKa = 12 – 5 = 7, so Keq = 10⁷. Equilibrium lies far to the right, strongly favouring products. The stronger acid (HA) donates its proton to the stronger base (B⁻).
Q: Draw the curved arrows for the proton transfer: CH₃O⁻ + H–OH → CH₃OH + OH⁻
A: Arrow 1 goes from the lone pair on the oxygen of CH₃O⁻ to the H of water. Arrow 2 goes from the O–H bond of water to the oxygen of water. Two arrows, both showing electron flow from rich to poor.
This connects directly to nucleophiles and electrophiles: a Lewis base is a nucleophile (electron-pair donor that attacks an electron-poor carbon), and a Lewis acid is an electrophile. The language changes, but the electron logic is identical.
Acid-base chemistry is the foundation for understanding substitution (SN1, SN2) and elimination (E1, E2) reactions. In every one of those mechanisms, you will identify which species is the base or nucleophile and which is the acid or electrophile.
pKa values help you predict leaving-group ability: the conjugate base of a strong acid is a good leaving group. This links acid-base strength directly to reactivity in organic mechanisms.
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