Worked example 1
Identify the acid, the base and the two conjugate pairs in HNO₂ + H₂O ⇌ NO₂⁻ + H₃O⁺.
Try it first: Find the species that lost an H going left to right.
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Acids, Bases & Aqueous EquilibriaAcids and BasesContent level: Core 16 min
Acid-Base Definitions and Conjugate Pairs
What you'll be able to do: Identify Bronsted-Lowry acids and bases in a reaction and write the conjugate partner of any species.
Introduction
Before you can calculate a pH you have to be able to say which species donated the proton and which one accepted it. Every later topic in this unit rests on that one skill.
These are recommended, not required. You can start this lesson at any time.
Learning objectives
- Distinguish the Arrhenius and Bronsted-Lowry definitions
- Identify the acid, the base and both conjugate partners in a reaction
- Write the conjugate acid or conjugate base of a given species
- Explain why a strong acid has a very weak conjugate base
Lesson
Two definitions
Arrhenius defined an acid as a substance that increases H⁺ in water and a base as one that increases OH⁻. That works for HCl and NaOH but fails for ammonia, which contains no OH⁻ yet clearly makes a basic solution. The Bronsted-Lowry definition fixes this: an acid donates a proton and a base accepts one.
Proton transfer and conjugate pairs
Because a proton has to go somewhere, an acid and a base always react together. HA + B gives A⁻ + HB⁺. The pair HA/A⁻ differs by exactly one proton, and so does B/HB⁺. Members of a conjugate pair never differ by anything else: not by an oxygen, not by two protons.
HA + B ⇌ A⁻ + HB⁺
To write a conjugate base, delete one H and lower the charge by one. To write a conjugate acid, add one H and raise the charge by one.
Water plays both roles
Water is amphiprotic. With HCl it accepts a proton to become H₃O⁺, and with NH₃ it donates a proton to become OH⁻. The hydrated proton is properly written H₃O⁺, though H⁺ is used as shorthand; both mean the same thing in a calculation.
Strength and the conjugate relationship
Strong acids such as HCl, HBr, HI, HNO₃, HClO₄ and H₂SO₄ transfer their proton essentially completely in water, so their conjugate bases (Cl⁻, NO₃⁻ and so on) have almost no tendency to take it back and are spectator ions. A weak acid such as CH₃COOH holds its proton well, so its conjugate base CH₃COO⁻ is a meaningful base.
The stronger the acid, the weaker its conjugate base. This inverse relationship is what makes salt solutions acidic or basic later in this unit.
Lewis acids and bases
A third definition is broader still: a Lewis acid accepts an electron pair and a Lewis base donates one. Every Bronsted acid is a Lewis acid, but species with an empty orbital, such as BF₃ or Al³⁺, are Lewis acids without donating any proton at all.
Key ideas
Definition
Bronsted-Lowry acid
A proton donor.
Definition
Bronsted-Lowry base
A proton acceptor.
Rule
Conjugate base
The species left after an acid donates one proton: one fewer H, charge lower by one.
Rule
Amphiprotic
A species such as water or HCO₃⁻ that can either donate or accept a proton.
Rule
Inverse strength
A strong acid has a negligibly weak conjugate base.
Equation
Proton transfer
HA + H₂O ⇌ H₃O⁺ + A⁻
- HA = acid, the proton donor
- A⁻ = its conjugate base
Equation
Base with water
B + H₂O ⇌ HB⁺ + OH⁻
- B = base, the proton acceptor
- HB⁺ = its conjugate acid
Worked examples
Worked example 2
Write the conjugate base of H₂PO₄⁻ and the conjugate acid of H₂PO₄⁻.
Try it first: Apply the add-one-H and remove-one-H rules separately.
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Worked example 3
Explain why a solution of NaCl is neutral while a solution of NaF is basic.
Try it first: Ask which ion, if any, is the conjugate base of a weak acid.
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Common mistakes
Calling any hydrogen-containing species an acid.
Why it's wrong: CH₄ and NH₃ contain hydrogen but do not donate protons in water.
Check instead: Ask whether the proton is actually transferred.
Writing the conjugate base of H₂SO₄ as SO₄²⁻.
Why it's wrong: That removes two protons, not one.
Check instead: Change exactly one H and one unit of charge.
Assuming strong means concentrated.
Why it's wrong: Strength describes the extent of ionization; concentration describes how much solute is present.
Check instead: A dilute HCl solution is still a strong acid.
Saying the conjugate base of a strong acid is a strong base.
Why it's wrong: The relationship is inverse: strong acid, negligible conjugate base.
Check instead: Cl⁻ has no measurable basicity in water.
Treating H⁺ and H₃O⁺ as different species in a calculation.
Why it's wrong: H₃O⁺ is just the hydrated form of H⁺.
Check instead: Use either symbol consistently.
Practice this skill
No practice questions are available for this topic yet. You can still practice the whole unit.
What you should now know
An Arrhenius acid produces H⁺ in water and an Arrhenius base produces OH⁻. The Bronsted-Lowry definition is broader and far more useful: an acid is a proton donor and a base is a proton acceptor. Removing one proton from an acid gives its conjugate base, and adding one proton to a base gives its conjugate acid, so every proton transfer involves two conjugate pairs. Water is amphiprotic, acting as an acid with bases and as a base with acids. Strong acids ionize essentially completely and have very weak conjugate bases, while weak acids ionize only partially and have conjugate bases strong enough to matter.
- A Bronsted acid donates a proton; a Bronsted base accepts one
- Conjugate partners differ by exactly one H and one unit of charge
- Water is amphiprotic and appears on both sides of this unit
- The stronger the acid, the weaker its conjugate base
- Strength is about extent of ionization, not concentration
Sources and further reading
This lesson is original Chem Help content. No external sources were adapted.