A conjugate base is the species that remains after a Brønsted-Lowry acid donates its proton (H+). Every acid HA has a conjugate base A-, and the pair differs by exactly one proton. In AP Chem, conjugate bases anchor weak acid equilibria (Ka), buffer chemistry, and the Henderson-Hasselbalch equation.
A conjugate base is what an acid becomes after it gives away its proton. When acetic acid (CH3COOH) donates H+ to water, the leftover acetate ion (CH3COO-) is its conjugate base. The two species form a conjugate acid-base pair, and the only difference between them is one H+ (which also means a charge difference of one). That's the whole test for spotting them: HA and A-, H2PO4- and HPO4^2-, NH4+ and NH3.
The deeper idea, and the one AP Chem actually cares about, is that strength is inversely related across the pair. Strong acids like HCl have conjugate bases (Cl-) so weak they basically never grab a proton back. Weak acids have conjugate bases strong enough to matter, which is exactly why a weak acid in water sits at equilibrium between HA and A- (described by Ka) instead of ionizing completely. Per the CED, the conjugate bases of strong acids are stabilized by electronegativity, inductive effects, or resonance, which is why they're so unreactive (8.6.A.1).
Conjugate bases first appear in Topic 4.8, where LO 4.8.A asks you to identify Brønsted-Lowry acids, bases, and conjugate pairs in a proton-transfer reaction. But the term does its heaviest lifting in Unit 8. LO 8.3.A frames a weak acid solution as an equilibrium between the un-ionized acid and its conjugate base, governed by Ka. LO 8.7.A uses the pH vs. pKa comparison to predict whether the acid form (HA) or the conjugate base form (A-) dominates in solution. And the entire buffer arc (LOs 8.8.A, 8.9.A, 8.10.A) is built on conjugate pairs. A buffer is just large amounts of both members of a conjugate pair, where the conjugate base neutralizes added acid and the conjugate acid neutralizes added base (8.8.A.1). If you can't identify the conjugate base in a system, you can't write buffer reactions, use Henderson-Hasselbalch, or explain buffer capacity. It's one of the most load-bearing terms in the course.
Keep studying AP Chemistry Unit 4
Conjugate Acid (Units 4 and 8)
Conjugate acid and conjugate base are two snapshots of the same pair. Add a proton to the conjugate base and you get the conjugate acid; remove it and you're back. On the exam, you'll often label both in one equation, so check which species gained H+ (that one's the conjugate acid) and which lost it.
Properties of Buffers (Unit 8)
A buffer works because the conjugate base is on standby to eat added H3O+. That's the literal mechanism in EK 8.8.A.1, and it's why a buffer with more conjugate base than acid has greater capacity against added acid (8.10.A.2). Buffer FRQs are really conjugate-base FRQs in disguise.
Molecular Structures of Acids and Bases (Unit 8)
Why is HCl strong but HF weak? Look at the conjugate base. Strong acids have conjugate bases stabilized by electronegativity, inductive effects, or resonance, so they won't take the proton back (8.6.A.1). Acid strength questions are conjugate base stability questions flipped around.
Introduction to Titration (Unit 4)
Titrating a weak acid with strong base converts HA into A- one proton at a time. At the half-equivalence point the conjugate pair is 50/50 and pH = pKa, and past equivalence the solution's pH is set by the conjugate base alone. The whole titration curve is a map of the HA-to-A- ratio.
Multiple-choice questions test conjugate bases in two main ways. First, straight identification: given a proton-transfer equation, pick out the conjugate pairs (LO 4.8.A). Second, buffer reasoning: questions hand you a buffer like 0.10 M CH3COOH with 0.20 M CH3COONa, or an H2PO4-/HPO4^2- mixture, and ask which reaction neutralizes added acid or base. The answer is always the conjugate base reacting with added H3O+ or the conjugate acid reacting with added OH-, and you need to write that specific net ionic equation. On FRQs, conjugate bases show up in buffer prep problems (the 2023 short FRQ had a student build a buffer from CH3NH2 and CH3NH3Cl, an equimolar conjugate pair) and in Henderson-Hasselbalch calculations where the [A-]/[HA] ratio sets the pH. You should be able to write the conjugate base of any given acid, justify acid strength using conjugate base stability, and explain in words why the conjugate base stabilizes pH against added acid.
Both come from the same conjugate pair, and students mix up which is which constantly. The fix is to track the proton's direction. The acid donates H+ and becomes the conjugate base (CH3COOH → CH3COO-). The base accepts H+ and becomes the conjugate acid (NH3 → NH4+). So in NH3 + H2O ⇌ NH4+ + OH-, NH4+ is the conjugate acid of NH3, and OH- is the conjugate base of H2O. One reaction, two pairs. Label the proton transfer first and the names fall out automatically.
A conjugate base is the species left over after a Brønsted-Lowry acid donates its proton, so HA and A- differ by exactly one H+.
Acid strength and conjugate base strength are inverses: strong acids like HCl and HNO3 have very weak conjugate bases, stabilized by electronegativity, inductive effects, or resonance.
A weak acid solution is an equilibrium between the un-ionized acid and its conjugate base, and Ka (or pKa) describes that equilibrium.
In a buffer, the conjugate base reacts with added acid and the conjugate acid reacts with added base, which is the entire reason buffers resist pH change.
When solution pH is above the acid's pKa, the conjugate base form (A-) dominates; when pH is below pKa, the acid form (HA) dominates.
A buffer with more conjugate base than conjugate acid has greater capacity to absorb added acid than added base.
It's the species that remains after an acid donates its proton. For example, when acetic acid (CH3COOH) gives up H+, the acetate ion (CH3COO-) is its conjugate base. The acid and its conjugate base differ by exactly one proton.
No, it's the opposite. Strong acids like HCl, HNO3, and H2SO4 have very weak conjugate bases (Cl-, NO3-, HSO4-) because those ions are stabilized by electronegativity, inductive effects, or resonance and won't take the proton back. That inverse relationship is exactly what Topic 8.6 tests.
Follow the proton. An acid loses H+ and becomes a conjugate base; a base gains H+ and becomes a conjugate acid. In NH3 + H2O ⇌ NH4+ + OH-, NH4+ is the conjugate acid of NH3 and OH- is the conjugate base of H2O.
The conjugate base is the buffer's defense against added acid. When you add H3O+ to an acetate buffer, CH3COO- reacts with it to form CH3COOH, so the pH barely moves. Without large amounts of both members of the conjugate pair, there's no buffering (EK 8.8.A.1).
Remove one H+ from the formula and lower the charge by one. H2PO4- becomes HPO4^2-, NH4+ becomes NH3, HF becomes F-. If two species in an equation differ by exactly one proton, they're a conjugate pair.