Strong and weak acids
At 0.100 M, hydrochloric acid is fully ionized and has pH 1.00, while acetic acid has given up only 1.3% of its protons and has pH 2.87.
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Hydrochloric acid and vinegar are both acids, but nobody cooks with hydrochloric acid. The difference is not only how much acid is in the bottle. Even at the same concentration, some acids give up their protons almost completely and others hold on to most of them. Chemists call the first kind strong acids and the second kind weak acids. Strength describes how far the acid ionizes in water; concentration describes how much acid is dissolved. The two are separate ideas, and a dilute strong acid can be less acidic than a concentrated weak one.
Strong acids ionize completely
OpenStax lists six common strong acids: perchloric acid HClO₄, hydrochloric acid HCl, hydrobromic acid HBr, hydroiodic acid HI, nitric acid HNO₃ and sulfuric acid H₂SO₄. For monoprotic strong acids such as HCl, ionization is essentially complete in dilute solution, so [H₃O⁺] is approximately the acid concentration. Sulfuric acid loses its first proton essentially completely, but its second ionization is an equilibrium that can add more H₃O⁺. A 0.100 M solution of HCl has [H₃O⁺] = 0.100 M and pH 1.00; a 0.0010 M solution has pH 3.00. The six common strong bases are the hydroxides of lithium, sodium, potassium, calcium, strontium and barium, which release their OH⁻ completely.
Weak acids and Ka
Most acids are weak. Acetic acid, the acid in vinegar, sets up an equilibrium in water: CH₃COOH + H₂O ⇌ H₃O⁺ + CH₃COO⁻. Its strength is measured by the acid ionization constant, Ka = [H₃O⁺][CH₃COO⁻] / [CH₃COOH]. For acetic acid Ka = 1.8 × 10⁻⁵, which means pKa = 4.74; PubChem lists a measured pKa of 4.76 at 25 °C. Formic acid is about ten times stronger with Ka = 1.8 × 10⁻⁴. The larger the Ka, the stronger the acid.
To find the pH of a weak acid, solve the equilibrium. For 0.100 M acetic acid, let x be the amount that ionizes: x² / (0.100 − x) = 1.8 × 10⁻⁵. Since x is small, x² ≈ 1.8 × 10⁻⁶ and x ≈ 1.3 × 10⁻³ M. The pH is 2.87, and the percent ionization is (1.3 × 10⁻³ / 0.100) × 100 = 1.3%. Only about one molecule in seventy-five has given up its proton, which is what the lab shows as one split molecule out of a hundred. Solving the equation without the small-x shortcut gives pH 2.88, and the lab uses that exact solution.
Dilution changes the percentage
Slide the concentration of acetic acid down in the lab. At 0.010 M about 4.2% is ionized, and at 0.0010 M about 12.5%. Ka stays the same, but the percent ionization grows as the acid gets more dilute, because the equilibrium shifts towards the side with more particles. A strong acid stays at 100% at every concentration. This is why percent ionization is not a fixed property of an acid, while Ka is.
Weak bases work the same way with a base ionization constant, Kb. Ammonia takes a proton from water to form NH₄⁺ and OH⁻ with Kb = 1.8 × 10⁻⁵, and PubChem gives the pH of 0.1 N ammonia as 11.1, far below the 13 of a strong base at the same concentration.
- Which is a strong acid?
- HNO₃
- CH₃COOH
- HCOOH
- H₂CO₃
- What is the pH of 0.0010 M HCl?
- What percentage of 0.100 M acetic acid is ionized?
- What is the pH of 0.100 M acetic acid?
- Which acid is stronger?
- Formic acid, Ka = 1.8 × 10⁻⁴
- Acetic acid, Ka = 1.8 × 10⁻⁵
- They are equal
- It depends only on concentration
- Acetic acid is diluted from 0.100 M to 0.0010 M. Its percent ionization…
- Increases
- Decreases
- Stays the same
- Becomes zero
- What is pKa for Ka = 1.8 × 10⁻⁵, to two decimal places?
- "Strong" in strong acid refers to…
- How completely it ionizes
- How concentrated it is
- How fast it pours
- How dangerous it smells
- Which is a strong base?
- NaOH
- NH₃
- CH₃COO⁻
- H₂O
- How many common strong acids does OpenStax list?
- HClO₄, HCl, HBr, HI and HNO₃ ionize essentially completely in dilute solution. H₂SO₄ loses its first proton essentially completely; its second ionization and weak-acid ionization are equilibria.
- Ka measures strength: acetic 1.8 × 10⁻⁵, formic 1.8 × 10⁻⁴.
- Percent ionization of a weak acid rises on dilution: 1.3% at 0.100 M acetic acid.
Sources for this course (10)
- OpenStax Chemistry 2e (openstax.org/books/chemistry-2e): 4.2 neutralization and antacids (stomach roughly 0.03 M HCl, milk of magnesia Mg(OH)2, Al(OH)3)
- 14.1 Brønsted-Lowry acids and bases (Boyle 1680, Davy 1815, Arrhenius 1884, Brønsted and Lowry 1923, autoionization about two in a billion molecules, Kw = 1.0 × 10⁻¹⁴ at 25 °C, about 5.6 × 10⁻¹³ at 100 °C)
- 14.2 pH and pOH (pH = −log[H3O+], stomach acid 1.2 × 10⁻³ M gives 2.92, vinegar about 4 × 10⁻³ M gives 2.4, blood 7.3, KOH 0.0125 M gives 12.10, neutral pH 6.31 at 80 °C, rainwater 5 to 6, acid rain below 5, pH meter resolution 0.001, Figure 14.2: 1 M HCl 0, gastric juice about 1.5, lime juice about 2, wine about 3.5, coffee about 5, milk of magnesia about 10.5)
- 14.3 relative strengths (six strong acids and six strong bases, Ka, percent ionization, 0.100 M acetic acid 1.3%)
- 14.6 buffers (normal blood pH about 7.4)
- 14.7 titrations (25.00 mL of 0.100 M HCl or acetic acid with 0.100 M NaOH, Table 14.2, equivalence pH 7.00 and 8.72, half-equivalence pH = pKa, Ka acetic acid 1.8 × 10⁻⁵, methyl orange Ka 4.0 × 10⁻⁴ red HIn and yellow In⁻, interval about pKa ± 1, phenolphthalein colourless below pH 8.3)
- LibreTexts Acid-Base Indicators (litmus about 5 to 8, methyl orange 3.1 to 4.4, phenolphthalein 8.3 to 10.0; chem.libretexts.org) and bromothymol blue 6.0 to 7.6 yellow to blue (Prairie South indicator table via LibreTexts Chemistry 30)
- PubChem acetic acid pKa 4.76 at 25 °C, sodium bicarbonate 0.1 M pH 8.3, ammonia 0.1 N pH 11.1, sodium hydroxide (pubchem.ncbi.nlm.nih.gov)
- NOAA Ocean acidification (noaa.gov: surface ocean pH has fallen by 0.1 since the industrial revolution, about a 30 percent increase in acidity)
- IUPAC CIAAW abridged standard atomic weights.
Checked: 2026-10-17