Amphoteric Substances — Acting as Both Acid and Base
Some substances refuse to pick a side. Zinc oxide dissolves in hydrochloric acid like a base and dissolves in sodium hydroxide like an acid. Water donates a proton to ammonia and accepts one from HCl. The hydrogencarbonate ion in your blood does both at once. These are the amphoteric substances, and they show up in the p-block, in ionic equilibrium, in biomolecules and in every qualitative analysis scheme. This page separates amphoteric from amphiprotic, gives the equations both ways, and works through the pH calculation students most often get wrong.
Two words that are not quite synonyms
Amphiprotic — can both donate a proton and accept a proton.
Every amphiprotic species is amphoteric, but the reverse is not true. Al₂O₃ has no ionisable hydrogen to donate, so it cannot be amphiprotic — yet it reacts happily with both HCl and NaOH, so it is amphoteric in the wider Lewis sense. Water and HCO₃⁻ are amphiprotic and amphoteric. Writing the right word earns the mark in a definition question.
Water — the standard example
Water is amphiprotic, and it proves it against itself:
- Acting as an acid: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (water donates a proton)
- Acting as a base: HCl + H₂O → H₃O⁺ + Cl⁻ (water accepts a proton)
In pure water at 298 K, [H₃O⁺] = [OH⁻] = √(1.0 × 10⁻¹⁴) = 1.0 × 10⁻⁷ mol L⁻¹, giving pH 7. Worth knowing for a stretch question: Kw increases with temperature, so the neutral pH of hot water falls below 7 — around 6.1 at 100 °C in commonly tabulated data. The water is still neutral, because [H₃O⁺] still equals [OH⁻]. "Neutral" means equal, not "pH 7".
Amphoteric oxides and hydroxides — write both equations
Examiners almost always want two balanced equations, one with an acid and one with a base. One equation alone does not demonstrate amphoteric character.
| Substance | With acid (behaves as a base) | With alkali (behaves as an acid) |
|---|---|---|
| ZnO | ZnO + 2HCl → ZnCl₂ + H₂O | ZnO + 2NaOH + H₂O → Na₂[Zn(OH)₄] |
| Al₂O₃ | Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O | Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄] |
| Al(OH)₃ | Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O | Al(OH)₃ + NaOH → Na[Al(OH)₄] |
| Zn(OH)₂ | Zn(OH)₂ + 2HCl → ZnCl₂ + 2H₂O | Zn(OH)₂ + 2NaOH → Na₂[Zn(OH)₄] |
| PbO | PbO + 2HNO₃ → Pb(NO₃)₂ + H₂O | PbO + 2NaOH + H₂O → Na₂[Pb(OH)₄] |
| BeO | BeO + 2HCl → BeCl₂ + H₂O | BeO + 2NaOH → Na₂BeO₂ + H₂O |
An honest note on the alkali products. Older textbooks and many NCERT-style answers write the products as meta-salts — NaAlO₂ (sodium meta-aluminate) and Na₂ZnO₂ (sodium zincate) — while modern treatments write the hydroxo-complexes Na[Al(OH)₄] and Na₂[Zn(OH)₄]. Both appear in current syllabi. The complex forms are the better description of what exists in aqueous solution; the meta forms are dehydrated versions of the same chemistry. Follow the form your own textbook uses, and do not be alarmed when a different book writes the other one.
Worked example 1 — a stoichiometry check on Al(OH)₃
Question. What mass of aluminium hydroxide will react with exactly 0.100 mol of NaOH?
From the equation Al(OH)₃ + NaOH → Na[Al(OH)₄], the ratio is 1 : 1, so 0.100 mol of Al(OH)₃ is needed.
Molar mass of Al(OH)₃:
Al: 1 × 26.982 = 26.982
O: 3 × 15.999 = 47.997
H: 3 × 1.008 = 3.024
M = 26.982 + 47.997 + 3.024 = 78.003 ≈ 78.00 g/mol
Cross-check by a second route: Al + 3 × (OH), where OH = 15.999 + 1.008 = 17.007, so 26.982 + 3 × 17.007 = 26.982 + 51.021 = 78.003 ✓
Mass = n × M = 0.100 × 78.003 = 7.80 g
Now the same 0.100 mol against acid: Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O needs 0.300 mol of HCl — three times as much, because here the hydroxide is behaving as a base with three OH groups to neutralise. Same compound, two completely different stoichiometries. That contrast is the exam point.
Worked example 2 — the pH of an amphiprotic salt
Sodium hydrogencarbonate is the classic case: HCO₃⁻ can donate a proton to become CO₃²⁻ or accept one to become H₂CO₃. For a salt like this the pH is set by the two acid constants of the parent acid, not by its concentration.
For carbonic acid, commonly tabulated values are pKa1 = 6.35 and pKa2 = 10.33.
pH = ½ (6.35 + 10.33) = ½ × 16.68 = 8.34
Converting back to a hydrogen ion concentration:
[H⁺] = 10^(−8.34) = 10^(0.66) × 10⁻⁹ = 4.6 × 10⁻⁹ mol L⁻¹
Two honest limits on this formula: it assumes the salt is not extremely dilute, and it ignores activity corrections. A real 0.1 M solution of baking soda measures close to this value, which is why it is mildly alkaline and safe to swallow, while sodium carbonate — where only the higher pKa applies — is far more alkaline and is not.
Worked example 3 — amino acids and the isoelectric point
An amino acid carries a –COOH group (acidic) and an –NH₂ group (basic) in one molecule. In solution it exists mainly as a zwitterion, ⁺H₃N–CHR–COO⁻, which has a positive and a negative charge but zero net charge.
In acid: the –COO⁻ is protonated, the ion becomes net positive.
In base: the –NH₃⁺ loses its proton, the ion becomes net negative.
The pH at which the net charge is zero is the isoelectric point:
pI = ½ (pKa1 + pKa2)
For glycine, with typical tabulated values pKa1 = 2.34 (carboxyl) and pKa2 = 9.60 (ammonium):
pI = ½ (2.34 + 9.60) = ½ × 11.94 = 5.97
At pH 5.97 glycine does not migrate in an electric field — the basis of electrophoresis, and also the reason amino acids are high-melting, water-soluble solids rather than the oily liquids their molar mass would suggest.
The periodic pattern behind amphoteric character
Amphoteric behaviour is not random. It sits at the boundary between metal and non-metal.
| Period 3 oxide | Na₂O | MgO | Al₂O₃ | SiO₂ | P₄O₁₀ | SO₃ | Cl₂O₇ |
|---|---|---|---|---|---|---|---|
| Character | Basic | Basic | Amphoteric | Acidic (weakly) | Acidic | Acidic | Strongly acidic |
Two rules follow, and both are examinable:
- Across a period, oxides go basic → amphoteric → acidic as the element becomes less metallic. Aluminium sits exactly on the changeover.
- For one element in several oxidation states, the higher the oxidation state, the more acidic the oxide. Chromium is the standard illustration: CrO is basic, Cr₂O₃ is amphoteric and CrO₃ is acidic.
Beryllium is amphoteric while the rest of Group 2 is basic — an instance of the diagonal relationship between beryllium and aluminium, and a favourite one-mark question.
Mistakes that cost marks
- Saying amphoteric means neutral. It means reactive towards both, not unreactive towards either. Water is amphoteric and neutral; ZnO is amphoteric and not.
- Giving only one equation. Amphoteric character is proved by two reactions. One equation answers half the question and earns half the marks.
- Confusing amphoteric with amphiprotic. Al₂O₃ is amphoteric but not amphiprotic — it has no proton to donate.
- Assuming every metal oxide is basic. ZnO, Al₂O₃, PbO, SnO, BeO and Cr₂O₃ are all metal oxides and all amphoteric.
- Using pH = ½(pKa1 + pKa2) for an ordinary salt. It applies to the intermediate amphiprotic species — HCO₃⁻, H₂PO₄⁻, HPO₄²⁻ — not to CO₃²⁻ or to a simple weak acid.
- Writing the zwitterion as uncharged. It carries both a + and a − charge; the net charge is zero. Drawing a plain neutral structure loses the mark.
- Forgetting the water on the left in alkali reactions. ZnO + 2NaOH + H₂O → Na₂[Zn(OH)₄] will not balance without it.
Where this is asked in exams
| Exam | Typical question |
|---|---|
| ICSE / CBSE Class 10 | Define amphoteric; give the two reactions of ZnO or Al₂O₃ |
| CBSE Class 11 | Periodic trends in oxide character; the diagonal relationship of Be with Al |
| CBSE Class 12 | p-block oxides, amino acids and zwitterions in biomolecules |
| JEE / NEET | Identify the amphoteric oxide from a list; predict products with NaOH |
| IIT-JAM / CUET-PG | pH of amphiprotic salts, isoelectric points, buffer regions of polyprotic acids |
| GATE / CSIR-NET | Acid–base behaviour of oxides and hydroxides, Lewis acidity, hydrolysis of metal ions |
Move between pH, pOH and concentration without slips. Every calculation on this page ends in a conversion — pH 8.34 back to 4.6 × 10⁻⁹ mol L⁻¹, or [OH⁻] back to a pH. The pH / pOH tool does all four conversions, so you can check the step where sign errors usually creep in.
Open the pH / pOH Calculator →Acid–base chemistry runs from Class 10 right through to Class 12 and beyond, and gaps early on show up later. ABC Chemistry teaches Class 11–12 chemistry at the Gurugram centre and online across India — details at abcchemistry.in. Students in Delhi, Noida or Gurgaon who want one-to-one teaching at home can arrange it through delhihometutor.com.