Organic Chemistry · Polar Covalent Bonds; Acids and Bases

Acids and Bases: The Brønsted–Lowry Definition

7 min read

Science note: K_w = 1.0 × 10^-14 at 25 °C and K_a = 1.8 × 10^-5 for acetic acid are standard literature values; no experimental data were fabricated. Lab instructions appear only as general principles.

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In 30 seconds

The Brønsted–Lowry definition recasts acid–base chemistry in terms of transfer. An acid is any species that donates a proton (a hydrogen ion, H⁺), and a base is any species that accepts one. Every acid–base reaction therefore produces a (the acid minus its proton) and a (the base plus a proton). This definition works in any solvent and covers far more than the Arrhenius acids and bases of general chemistry — including the carbon acids, ammonium salts, and protonated heteroatoms that dominate organic mechanisms. The strength of an acid is quantified by its acid-dissociation constant K_a and its pK_a, which are introduced here and developed in Topic 8.

Why this matters

Proton transfer is the most frequent elementary step in organic chemistry. Nearly every mechanism — acid-catalyzed ester hydrolysis, enamine formation, aldol condensation, enzyme catalysis — begins or ends with a proton being handed from one molecule to another. In medicine, the fraction of a drug that is neutral rather than ionized at physiological pH (about 7.4) determines whether it crosses membranes and reaches its target; that fraction is controlled by the drug's pK_a through the Henderson–Hasselbalch relationship. The same equilibrium chemistry maintains blood pH through bicarbonate buffers. Understanding Brønsted–Lowry acids and bases is therefore not just a chapter exercise — it is the operating system for organic reactions and the key to drug absorption and physiological buffering.

The college version

Core Concepts

The definition

A is a proton donor; a is a proton acceptor. The proton is a bare hydrogen nucleus, H⁺, with no electrons, so it can never exist alone in solution: it must be transferred from one bond to another. In water, an acid transfers its proton to a water molecule, forming the hydronium ion, H₃O⁺.

Conjugate pairs

When an acid HA loses its proton it becomes A^-, the conjugate base of HA. When a base B gains a proton it becomes BH⁺, the conjugate acid of B. Acid and conjugate base always differ by one proton; base and conjugate acid always differ by one proton. Every acid–base reaction has two conjugate pairs: HA/A^- and BH⁺/B. A strong acid has a weak conjugate base, and a strong base has a weak conjugate acid.

Water: amphoteric behavior

Water can act as an acid (donating a proton to ammonia to give OH^- and NH₄⁺) or as a base (accepting a proton from HCl to give H₃O⁺ and Cl^-). Its self-ionization,

$2 H 2 O ⇌ H 3 O + + OH -$

has the equilibrium constant K_w = [H₃O⁺][OH^-] = 1.0 × 10^-14 at 25 °C. This single constant anchors all pH and pK_a calculations.

Acid strength and Ka

For an acid HA in water, HA + H₂O ⇌ H₃O⁺ + A^-, the acid-dissociation constant is:

$K a = [H 3 O +][A -] [HA]$

A larger K_a means more dissociation, hence a stronger acid. Because K_a values span many orders of magnitude, chemists use pK_a = -log₁₀ K_a: the smaller the pK_a, the stronger the acid. Water itself has pK_a ≈ 15.7 (using [H₂O] = 55.5 M), acetic acid has pK_a ≈ 4.74, and HCl has pK_a ≈ -7.

Common Confusions

Do Not Confuse	With	Difference
Acid strength	Acid concentration	Strength is an equilibrium property (K_a); concentration is how much acid is present.
Conjugate base	A base in general	The conjugate base is specifically the acid minus a proton; any species can act as a base.
Large K_a	Large pK_a	Larger K_a = stronger acid; larger pK_a = weaker acid. They move in opposite directions.
H⁺ free in solution	H₃O⁺ in water	Protons cannot exist alone; in water they are always attached to water as hydronium.
Brønsted–Lowry acid	Lewis acid	Brønsted acids donate protons; Lewis acids accept electron pairs (Topic 11).
"Strong" acid	"Dangerous" acid	Strong means fully ionized; hazards depend on concentration, quantity, and exposure.

Eli explains

The same idea, in plain words

Explain it like I’m 10

A Brønsted–Lowry acid is a molecule that hands a proton (a tiny hydrogen core) to a partner, and a base is the partner that catches it. It works like a relay race: the acid passes the baton and becomes the empty-handed conjugate base, while the base catches it and becomes the loaded conjugate acid. Every great acid makes a calm, weak conjugate base — the better the pass, the more comfortable the catcher.

Worked examples

Write the reaction of acetic acid with ammonia and label all four participants:

$CH 3 COOH + NH 3 ⇌ CH 3 COO - + NH 4 +$

Acetic acid donates a proton to ammonia, so acetic acid is the acid and ammonia is the base. The acetate ion is acetic acid's conjugate base; the ammonium ion is ammonia's conjugate acid. The two conjugate pairs are CH₃COOH/CH₃COO^- and NH₄⁺/NH₃. Because acetic acid is a much stronger acid than ammonium ion (pK_a 4.74 vs. 9.3), the equilibrium lies to the right — acetic acid protonates ammonia almost completely.

Hydrochloric acid is a strong acid: in water it ionizes completely, HCl + H₂O → H₃O⁺ + Cl^-. For a 0.010 M HCl solution, the hydronium concentration equals the acid concentration:

$[H 3 O +] = 0.010 M$

The pH is defined by pH = -log₁₀[H₃O⁺]. Substituting:

$pH = - log 10 (0.010) = 2.00$

Unit check: [H₃O⁺] has units mol/L, and the log of a dimensionless ratio yields a dimensionless pH, so the result carries no units. The strong-acid case is the fastest pH calculation in chemistry and the benchmark against which weak-acid problems are compared.

Acetic acid in water does not dissociate fully. Its K_a = 1.8 × 10^-5 at 25 °C. For a 1.0 M solution, let x = [H₃O⁺]. Then the equilibrium expression is:

$K a = x 2 1.0 - x \approx x 2 1.0$

The approximation is valid because x is small compared with 1.0 M. Solving:

$x = \sqrt K a \times 1.0 M = \sqrt 1.8 \times 10 -5 \times 1.0 M 2 = \sqrt 1.8 \times 10 -5 M 2 = 4.2 \times 10 -3 M$

Dimensional analysis confirms the algebra: K_a has units M, multiplied by 1.0 M gives M², and the square root restores M — a concentration. Only about 0.42% of the acetic acid is ionized:

$4.2 \times 10 -3 M 1.0 M \times 100 % = 0.42 %$

Weak acids are mostly undissociated, which is why vinegar tastes sour rather than burning: the concentration of free H₃O⁺ stays low.

Key takeaways

Brønsted–Lowry acid = proton donor; base = proton acceptor; water is both (amphoteric).
Acid and conjugate base differ by one proton; base and conjugate acid differ by one proton.
Every acid–base reaction contains two conjugate pairs and proceeds by proton transfer.
K_a = [H₃O⁺][A^-]/[HA]; pK_a = -log₁₀ K_a; smaller pK_a = stronger acid.
K_w = [H₃O⁺][OH^-] = 1.0 × 10^-14 at 25 °C; pure water has [H₃O⁺] = 1.0 × 10^-7 M.
Strong acid ⇄ weak conjugate base; strong base ⇄ weak conjugate acid.
A proton has no electrons; it cannot exist free in solution and must transfer directly from one bond to another.
Carboxylic acids, ammonium salts, and protonated alcohols are the acid forms most common in organic mechanisms.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

State the Brønsted–Lowry definitions of acid and base.
Show answer
An acid is a proton donor; a base is a proton acceptor.
In HF + H₂O ⇌ H₃O⁺ + F^-, identify the two conjugate pairs.
Show answer
HF/F^- and H₃O⁺/H₂O.
Why can a proton never exist alone in solution?
Show answer
A proton has no electrons, so it cannot exist as a free particle in solution; it must transfer directly from one atom to another.
K_w for water is 1.0 × 10^-14 at 25 °C. What is [H₃O⁺] in pure water?
Show answer
In pure water [H₃O⁺] = [OH^-] = 1.0 × 10^-14 = 1.0 × 10^-7 M.
What is the pH of a 0.0010 M HCl solution?
Show answer
pH = -log₁₀(0.0010) = 3.00.
True or false: a strong acid always has a strong conjugate base.
Show answer
False. A strong acid has a weak conjugate base (e.g., Cl^- from HCl).

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Brønsted–Lowry acid: Any species that donates a proton, H⁺.
Brønsted–Lowry base: Any species that accepts a proton.
conjugate base: What remains after an acid loses a proton (A^- from HA).
conjugate acid: What forms when a base gains a proton (BH⁺ from B).
proton: A hydrogen nucleus, H⁺, with no electrons.
amphoteric: Able to act as both acid and base (e.g., water).
Kₐ / pKₐ: Acid-dissociation constant and its negative base-10 log.
Kw: Ion-product constant of water, 1.0 × 10^-14 at 25 °C.

Sources & references

openstax.org — Organic Chemistry

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

The definition

Conjugate pairs

Water: amphoteric behavior

Acid strength and Ka

Common Confusions

The same idea, in plain words

Worked examples

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references