A conjugate base is what’s left after an acid gives away a proton (H⁺), and a conjugate acid is what forms after a base picks one up. The entire concept comes down to tracking a single hydrogen ion as it moves from one species to another. Once you see that pattern, identifying conjugate pairs in any reaction becomes straightforward.
The One-Proton Rule
Every conjugate pair differs by exactly one hydrogen ion. That’s the only rule you need to memorize. An acid donates H⁺ and becomes its conjugate base. A base accepts H⁺ and becomes its conjugate acid. The word “conjugate” just means “partner”: it’s the species on the other side of that proton transfer.
Take hydrochloric acid (HCl) dissolving in water:
HCl + H₂O → H₃O⁺ + Cl⁻
HCl donates a proton, so it’s the acid. After losing that proton, it becomes Cl⁻, its conjugate base. Water accepts the proton, so it’s the base. After gaining the proton, it becomes H₃O⁺, its conjugate acid. Two conjugate pairs exist in every acid-base reaction: one on each side.
How to Find the Conjugate Base of Any Acid
Start with the acid’s formula and remove one H⁺. That’s the conjugate base. The charge on the conjugate base will always be one unit more negative than the original acid, because a positively charged proton left without taking any electrons with it.
- HCl → Cl⁻ (neutral acid becomes negatively charged base)
- CH₃COOH → CH₃COO⁻ (acetic acid loses a proton to form the acetate ion)
- NH₄⁺ → NH₃ (positive ion becomes neutral)
- H₂O → OH⁻ (neutral molecule becomes negatively charged)
- HBr → Br⁻ (hydrobromic acid becomes bromide)
Notice the pattern: in every case, the conjugate base is more negative (or less positive) than the acid by exactly one charge.
How to Find the Conjugate Acid of Any Base
This works in reverse. Take the base’s formula and add one H⁺. The result is the conjugate acid, and its charge will be one unit more positive than the original base.
- NH₃ + H⁺ → NH₄⁺ (ammonia’s conjugate acid is ammonium)
- OH⁻ + H⁺ → H₂O (hydroxide’s conjugate acid is water)
- F⁻ + H⁺ → HF (fluoride’s conjugate acid is hydrofluoric acid)
- CN⁻ + H⁺ → HCN (cyanide’s conjugate acid is hydrogen cyanide)
- CH₃COO⁻ + H⁺ → CH₃COOH (acetate’s conjugate acid is acetic acid)
Identifying Both Pairs in a Reaction
In a full equation, there are always two conjugate pairs. Here’s how to find them both, step by step.
Look at the reaction of ammonia with water:
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
First, find which reactant gained a proton and which lost one. NH₃ starts with three hydrogens and ends up as NH₄⁺ with four, so it gained a proton. That makes NH₃ the base and NH₄⁺ its conjugate acid (pair one). H₂O starts with two hydrogens and ends up as OH⁻ with one, so it lost a proton. That makes H₂O the acid and OH⁻ its conjugate base (pair two).
You can apply this same logic to any equation. Compare each reactant to the product it transforms into, count the hydrogens, and check the charge. The species that lost an H⁺ is the acid, and what it became is the conjugate base. The species that gained an H⁺ is the base, and what it became is the conjugate acid.
When a Substance Can Be Either
Some molecules can donate or accept a proton depending on what they’re reacting with. Water is the most common example. When HCl dissolves in water, water accepts a proton and acts as a base. When water reacts with ammonia, it donates a proton and acts as an acid. Water even reacts with itself in a process called autoionization, where one water molecule donates a proton to another:
H₂O + H₂O ⇌ H₃O⁺ + OH⁻
Species that behave this way are called amphiprotic. The bicarbonate ion (HCO₃⁻) is another common one. It can donate a proton to become CO₃²⁻ (its conjugate base) or accept a proton to become H₂CO₃ (its conjugate acid). When you encounter an amphiprotic species on a test, the reaction context tells you which role it’s playing.
Acids With More Than One Proton
Some acids, called polyprotic acids, have multiple hydrogen ions they can donate. They lose protons one at a time, creating a chain of conjugate pairs. Sulfuric acid (H₂SO₄) is a classic example:
Step 1: H₂SO₄ → H⁺ + HSO₄⁻
Step 2: HSO₄⁻ → H⁺ + SO₄²⁻
In the first step, H₂SO₄ is the acid and HSO₄⁻ is its conjugate base. In the second step, that same HSO₄⁻ now acts as an acid, and SO₄²⁻ is its conjugate base. Each step strips away one proton and shifts the charge by one unit. Hydrogen sulfide (H₂S) works the same way: H₂S loses a proton to become HS⁻, and HS⁻ can lose another to become S²⁻. Each proton is harder to remove than the last, which is why these acids dissociate in stages rather than all at once.
The Strength Relationship
There’s an inverse relationship between how strong an acid is and how strong its conjugate base is. A strong acid like HCl donates its proton very easily, which means Cl⁻ has almost no tendency to grab that proton back. Cl⁻ is therefore a very weak base. On the flip side, a very weak acid like water barely donates protons at all, which means its conjugate base, OH⁻, holds onto protons aggressively and is a strong base.
This pattern is useful as a quick check. If you know you’re dealing with a strong acid, you can immediately predict its conjugate base will be weak, and vice versa. Acetic acid is a weak acid, so its conjugate base, the acetate ion, is a relatively strong base compared to something like Cl⁻. The weaker the acid, the stronger the conjugate base.
Quick-Reference Summary
- Conjugate base: remove one H⁺ from the acid, charge decreases by one
- Conjugate acid: add one H⁺ to the base, charge increases by one
- In a reaction: the acid and its conjugate base are one pair; the base and its conjugate acid are the other
- Amphiprotic species: can go either direction depending on the reaction partner
- Strength rule: strong acid = weak conjugate base, weak acid = strong conjugate base