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Chemistry 2e · Acid-Base Equilibria

Brønsted-Lowry Acids and Bases

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On this page 8 sections
  1. In 30 seconds
  2. Why this matters
  3. The college version
  4. Eli explains
  5. Key takeaway
  6. Check yourself
  7. Study tools
  8. Sources & references

In 30 seconds

The Brønsted-Lowry theory defines acids and bases by what they do with a proton (a hydrogen ion, H+). An acid is a proton donor; a base is a proton acceptor. This definition expands on the older Arrhenius idea (limited to water and OH− production) so acid-base behavior can be recognized in any solvent — even the gas phase.

The key structural idea is the conjugate pair. When an acid donates a proton, what remains is its conjugate base; when a base accepts one, the result is its conjugate acid. Every acid-base reaction therefore involves two pairs:

HClacid + H2Obase ⇌ H3O+conj. acid + Cl−conj. base

Water is special: it can act as either an acid or a base depending on its partner. Species that can do both are . This chapter uses the Brønsted-Lowry framework to build the pH scale, compare acid strengths, and understand buffers and titrations.

Why this matters

Acid-base chemistry is everywhere in the body and the environment, and the Brønsted-Lowry view is its language:

  • Human physiology: Exhaled carbon dioxide exists in blood mainly as carbonic acid and bicarbonate, H2CO3 / HCO3−, a conjugate pair that keeps blood pH near 7.4. Kidney and lung function are, in essence, acid-base management.
  • Drug behavior: Many medications are weak acids or bases. Their ability to cross membranes depends on which protonated form dominates at body pH — a direct consequence of these equilibria (developed further in later topics).
  • Environmental chemistry: Acid rain, ocean acidification from dissolved CO₂, and soil chemistry are all proton-transfer problems. Understanding who donates and who accepts a proton lets you reason about these systems without memorizing reactions.

The definitions also clear up a common confusion: a substance need not contain OH− to be a base. Ammonia, NH3, is a classic base because it accepts a proton.

The college version

Core Concepts

Acids donate protons; bases accept them

An acid is anything that can give up H+; a base is anything that can take it. The proton is a bare hydrogen nucleus — tiny and strongly attracted to electron pairs — so in water it never exists alone; it is always attached to a water molecule as the hydronium ion, H3O+. Writing H+(aq) is shorthand for H3O+(aq).

Consider ammonia in water:

NH3(aq) + H2O(l)  ⇌ NH4+(aq) + OH-(aq)

Here NH3 accepts a proton (it is the base) and H2O donates one (it is the acid). Note that NH3 contains no oxygen and no OH− — yet it is unmistakably a base under this definition.

Conjugate acid-base pairs

A conjugate pair differs by exactly one proton. Remove a proton from an acid and you get its conjugate base; add a proton to a base and you get its conjugate acid.

  • HCl / Cl−: chloride is the conjugate base of hydrochloric acid.
  • NH4+ / NH3: ammonium is the conjugate acid of ammonia.
  • H2O / OH−: hydroxide is the conjugate base of water.

In any acid-base reaction, the acid and base on the left are converted into their conjugates on the right. Identifying the two pairs is the fundamental skill: find the species that differ by one H+, and they form a pair.

Amphiprotic species

Water, bicarbonate (HCO3−), dihydrogen phosphate (H2PO4−), and hydrogen sulfate (HSO4−) can donate or accept a proton:

  • With a strong base like OH−, bicarbonate acts as an acid: HCO3− + OH−  ⇌ CO32− + H2O.
  • With a strong acid, bicarbonate acts as a base: HCO3− + H3O+  ⇌ H2CO3 + H2O.

This dual behavior makes bicarbonate the central buffer species in blood.

Strong versus weak: complete versus partial proton transfer

Strong acids and bases transfer protons essentially completely in water; weak ones transfer only partially, establishing an equilibrium. Common strong acids include HCl, HBr, HI, HNO3, H2SO4 (first proton), and HClO4. Common strong bases are the hydroxides of group 1 metals (like NaOH, KOH) and of the heavier group 2 metals (like Ba(OH)2). Everything else commonly encountered — acetic acid, carbonic acid, ammonia — is weak and described by an equilibrium constant (next topics). Strong means "goes to completion"; weak means "use Ka or Kb."

Water's self-ionization: the connection to pH

Water autoionizes:

2H2O(l) ⇌ H3O+(aq) + OH-(aq)

At 25 °C the ion product is

Kw = [H3O+][OH−] = 1.0 × 10-14

This single equation links every acid-base calculation: knowing one of [H3O+] or [OH−] gives the other, and it is the foundation of the pH and pOH scale in the next topic.

How It Works / Step-by-Step Process

Worked example 1: labeling the species in an acid-base reaction

Problem. For HF(aq) + H2O(l) ⇌ H3O+(aq) + F-(aq), label the acid, base, conjugate acid, and conjugate base.

Solution.

  1. Find species that differ by one proton: HF/F− form a conjugate pair; H2O/H3O+ form the other.
  2. Track the proton: HF gives one up, so it is the acid and F− its conjugate base; H2O gains it, so it is the base and H3O+ its conjugate acid.
  3. Summary: acid HF, base H2O, conjugate acid H3O+, conjugate base F−.

Worked example 2: predicting products from a conjugate-pair analysis

Problem. Predict the products when ammonia accepts a proton from water, and identify the pairs.

Solution.

  1. Ammonia accepts a proton: NH3 is the base; it becomes NH4+ (its conjugate acid).
  2. Water donates the proton: H2O is the acid; it becomes OH− (its conjugate base).
  3. Balanced reaction:

NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)

Pairs: NH4+/NH3 and H2O/OH−. The product side contains hydroxide even though no starting base "contained" OH− — proton transfer created it.

Worked example 3: using Kw to find a missing concentration

Problem. In a solution at 25 °C, [H3O+] = 2.5 × 10-4 M. Find [OH−].

Solution.

  1. Write the ion-product expression: Kw = [H3O+][OH−] = 1.0 × 10-14.
  2. Solve for the unknown by dividing both sides:

[OH−] = Kw[H3O+] = 1.0 × 10-142.5 × 10-4 = 4.0 × 10-11 M

Dimensional analysis: (M2)/(M) = M, so the answer is in molarity. The solution is acidic (more H3O+ than OH−) — quantified by the pH scale next.

Common Confusions

Do Not ConfuseWithDifference
Base containing OH−Base accepting a protonNH3 and CO32− are bases with no hydroxide; basicity is about accepting H+, not containing OH−.
Conjugate pair membersAny acid and base in the equationPair members differ by exactly one proton; HCl and H2O in the intro reaction are NOT a pair.
H+ alone in waterH3O+A bare proton cannot exist in water; writing H+(aq) is shorthand for hydronium.
Strong acidConcentrated acidStrength is about how completely a substance ionizes, not how much is present; a dilute strong acid is still "strong."
AmphiproticAmphoteric (neutral in all senses)Amphiprotic specifically means able to donate or accept a proton; the term applies to species, not to reactions.
Eli, the EliExplains learning guide

Eli explains

The same idea, in plain words

Explain it like I’m 10

Think of a proton as a ball that molecules pass to each other. An acid is a molecule that passes the ball away; a base is one that catches it. Whenever the ball moves, the passer turns into its "empty-handed" partner and the catcher turns into its "holding-the-ball" partner — those partners are the conjugate pair. Some molecules, like water, can do both jobs depending on who they meet.

Key takeaways

  • Brønsted-Lowry: acid = proton donor; base = proton acceptor. No OH− required for basicity.
  • H+(aq) is shorthand for hydronium, H3O+(aq).
  • A conjugate pair differs by exactly one proton: acid ⇌ conjugate base + H+.
  • Every acid-base reaction contains two conjugate pairs; label all four species.
  • Amphiprotic species (e.g., H2O, HCO3−, H2PO4−, HSO4−) can donate or accept a proton.
  • Strong acids/bases ionize completely; weak ones establish equilibria (quantified by Ka/Kb in later topics).
  • Water autoionizes with Kw = [H3O+][OH−] = 1.0 × 10-14 at 25 °C.
  • The stronger an acid, the weaker its conjugate base (and vice versa) — a preview of the relative-strengths topic.

Check yourself

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

  1. Define acid and base in the Brønsted-Lowry sense.

    Show answer

    An acid donates a proton; a base accepts a proton.

  2. In HNO2(aq) + H2O(l) ⇌ H3O+(aq) + NO2-(aq), identify all four species by role.

    Show answer

    Acid HNO2, base H2O, conjugate acid H3O+, conjugate base NO2−.

  3. Why can water act as both acid and base? What is the term?

    Show answer

    Water can lose a proton (acid → OH−) or gain one (base → H3O+); dual behavior is called amphiprotic.

  4. Write the two conjugate pairs in the reaction of HCO3− with OH−.

    Show answer

    HCO3− + OH− ⇌ CO32− + H2O: pairs are HCO3-/CO32− and H2O/OH−.

  5. At 25 °C, if [OH−] = 1.0 × 10-3 M, what is [H3O+]? Acidic or basic?

    Show answer

    [H3O+] = Kw/[OH−] = 1.0 × 10-14/1.0 × 10-3 = 1.0 × 10-11 M. Since [OH−] > [H3O+], it is basic.

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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 acid-base pair
Two species differing by one proton
hydronium ion (ceH3O+)
The form a proton takes in water
amphiprotic
Able to act as either an acid or a base
Kw
Ion product of water, 1.0 × 10-14 at 25 °C
conjugate acid–base pair
Two species differing by one proton (e.g., NH4+/NH3).

Sources & references

  1. openstax.org — Chemistry 2e

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

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