A Brønsted-Lowry base is any species that accepts a proton (H⁺) in a reaction, forming its conjugate acid. In AP Chem Unit 8, this definition explains how weak bases like NH₃ react with water to produce OH⁻, and why every acid-base reaction is really a proton transfer.
A Brønsted-Lowry base is a substance that accepts a proton (H⁺) during a chemical reaction. That's the whole definition, and it's deliberately broad. The species doesn't need to contain OH⁻, and it doesn't even need to be in water. It just needs a lone pair available to grab an H⁺. Ammonia (NH₃) is the classic example. It has no hydroxide in it, but it accepts a proton from water to form NH₄⁺ and OH⁻, which makes it a base.
The key move on the AP exam is tracking what happens after the base accepts the proton. Once a Brønsted-Lowry base picks up H⁺, it becomes its conjugate acid (NH₃ becomes NH₄⁺). Meanwhile, whatever donated the proton becomes a conjugate base. Every Brønsted-Lowry reaction is a proton handoff between two conjugate acid-base pairs. Even water gets in on this. In its autoionization, one water molecule acts as the acid (donating H⁺) and another acts as the base (accepting it to form hydronium, H₃O⁺). Water playing both roles is called being amphoteric, and it's the foundation of Kw, pH, and pOH in Topic 8.1.
This term lives in Unit 8: Acids and Bases, specifically Topics 8.1 and 8.3. Topic 8.1 (learning objective 8.1.A) has you calculate pH and pOH from Kw, and that whole framework rests on water acting as a Brønsted-Lowry base when it accepts a proton to form hydronium. Topic 8.3 (learning objective 8.3.A) asks you to relate pH, pOH, and the concentrations of every species in a weak base solution. You can't set up that equilibrium without recognizing the base accepting a proton from water: B + H₂O ⇌ HB⁺ + OH⁻. The Brønsted-Lowry definition is also what makes 'conjugate pairs' a thing, and conjugate pairs are the backbone of buffers, titrations, and basically every problem in the second half of Unit 8. If you can spot the proton acceptor in any equation, you can identify both conjugate pairs and write the correct equilibrium expression, which is half the battle on these questions.
Keep studying AP Chemistry Unit 8
Conjugate Acid (Unit 8)
A Brønsted-Lowry base and its conjugate acid differ by exactly one H⁺. When NH₃ accepts a proton, it becomes NH₄⁺, its conjugate acid. The exam loves asking you to identify these pairs, so memorize the pattern that the base plus H⁺ equals the conjugate acid.
Hydronium Ion and Water Autoionization (Unit 8, Topic 8.1)
Hydronium exists because water can act as a Brønsted-Lowry base. One H₂O accepts a proton from another, giving H₃O⁺ and OH⁻ with Kw = 1.0 × 10⁻¹⁴ at 25°C. Every pH and pOH calculation traces back to this single proton transfer.
Weak Base Equilibria and Ka/Kb (Unit 8, Topic 8.3)
When a weak base accepts a proton from water, only a small fraction of molecules actually react, so you get an equilibrium described by Kb. This mirrors the weak acid story in 8.3.A.2, where Ka describes partial ionization. Same math, just flipped to track OH⁻ instead of H₃O⁺.
Proton Transfer (Unit 8)
The Brønsted-Lowry model reframes every acid-base reaction as one event, a proton moving from donor to acceptor. Seeing reactions this way is what lets you handle non-aqueous bases, amphoteric species, and buffer chemistry that the older hydroxide-based definition can't explain.
No released FRQ has used the phrase 'Brønsted-Lowry base' verbatim, but the concept is baked into nearly every Unit 8 question. Multiple-choice stems will give you a reaction and ask which species acts as the base, or ask you to identify a conjugate acid-base pair. FRQs expect you to use the definition without being told, like writing the equilibrium for a weak base reacting with water (B + H₂O ⇌ HB⁺ + OH⁻), then setting up Kb and solving for pOH and pH. The skill being tested isn't reciting the definition. It's spotting which species gained an H⁺, which lost one, and writing the correct equilibrium expression from that.
An Arrhenius base must produce OH⁻ ions in water, which limits you to compounds like NaOH. A Brønsted-Lowry base just needs to accept a proton, so it covers way more species. NH₃ is the test case. It has no OH⁻ in its formula, so it fails the Arrhenius definition, but it accepts a proton from water (generating OH⁻ in the process), so it's a textbook Brønsted-Lowry base. The AP exam uses the Brønsted-Lowry framework almost exclusively because it explains conjugate pairs and amphoteric behavior.
A Brønsted-Lowry base accepts a proton (H⁺), and after accepting it, the base becomes its conjugate acid.
A base doesn't need OH⁻ in its formula. NH₃ is a Brønsted-Lowry base because it accepts a proton from water to form NH₄⁺ and OH⁻.
Water is amphoteric, meaning it can act as a Brønsted-Lowry base (forming H₃O⁺) or as an acid (forming OH⁻), which is why Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C.
Weak bases only partially accept protons from water, so their solutions are described by an equilibrium constant Kb, parallel to how Ka works for weak acids in Topic 8.3.
To find a conjugate pair, look for two species in an equation that differ by exactly one H⁺. The one with the extra proton is the acid of the pair.
Identifying the proton acceptor in a reaction is the first step to writing the correct equilibrium expression on Unit 8 FRQs.
It's any species that accepts a proton (H⁺) in a reaction. When it accepts the proton, it forms its conjugate acid. For example, NH₃ accepts H⁺ from water to become NH₄⁺, making NH₃ a Brønsted-Lowry base.
No. That's the Arrhenius definition. A Brønsted-Lowry base only needs to accept a proton, so species like NH₃, F⁻, and CO₃²⁻ all count even though none of them contain OH⁻. Many of them generate OH⁻ by pulling a proton off water, but the OH⁻ isn't part of the base itself.
A Brønsted-Lowry base is any proton acceptor in a reaction. A conjugate base is specifically what's left over after an acid donates its proton. For example, when HF donates H⁺, the F⁻ left behind is its conjugate base, and F⁻ can then act as a Brønsted-Lowry base in the reverse reaction.
It can be. Water is amphoteric, so it acts as a base when it accepts a proton (forming H₃O⁺, like when HCl dissolves) and as an acid when it donates one (forming OH⁻, like when NH₃ dissolves). Water's autoionization, with Kw = 1.0 × 10⁻¹⁴ at 25°C, is one water molecule doing each job.
Compare each species to its product and find the one that gained an H⁺. The proton acceptor is the base, and its product is the conjugate acid. In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, NH₃ gained a proton to become NH₄⁺, so NH₃ is the base and H₂O is the acid.
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