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Title: Amphoterism  
Author: World Heritage Encyclopedia
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Subject: Chromium, Infobox hydrogen/testcases, Infobox plutonium/testcases, Infobox titanium/testcases, Infobox germanium/testcases
Collection: Acid–base Chemistry, Chemical Properties
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In chemistry, an amphoteric compound is a molecule or ion that can react as an acid as well as a base.[1] The word is derived from the Greek word amphoteroi (ἀμφότεροι) meaning "both". Many metals (such as copper, zinc, tin, lead, aluminium, and beryllium) form amphoteric oxides or hydroxides. Amphoterism depends on the oxidation state of the oxide.

One type of amphoteric species are amphiprotic molecules, which can either donate or accept a proton (H+). Examples include amino acids and proteins, which have amine and carboxylic acid groups, and self-ionizable compounds such as water and ammonia.

Ampholytes are amphoteric molecules that contain both acidic and basic groups and will exist mostly as zwitterions in a certain range of pH. The pH at which the average charge is zero is known as the molecule's isoelectric point. Ampholytes are used to establish a stable pH gradient for use in isoelectric focusing.


  • Amphoteric oxides and hydroxides 1
  • Amphiprotic molecules 2
    • Examples 2.1
    • Not all amphoteric substances are amphiprotic 2.2
  • See also 3
  • References 4

Amphoteric oxides and hydroxides[2]

Zinc oxide (ZnO) reacts with both acids and with bases:

  • In acid: ZnO + 2H+ → Zn2+ + H2O
  • In base: ZnO + H2O + 2 OH → [Zn(OH)4]2−

This reactivity can be used to separate different cations, such as zinc(II), which dissolves in base, from manganese(II), which does not dissolve in base.

Aluminium hydroxide is also amphoteric:

  • As a base (neutralizing an acid): Al(OH)3 + 3 HCl → AlCl3 + 3 H2O
  • As an acid (neutralizing a base): Al(OH)3 + NaOH → Na[Al(OH)4]

Some other amphoteric compounds include:

  • Beryllium hydroxide
    • with acid: Be(OH)2 + 2 HCl → BeCl2 + 2 H2O
    • with base: Be(OH)2 + 2 NaOH → Na2[Be(OH)4]
  • Aluminium oxide
    • with acid: Al2O3 + 3 H2O + 6 H3O+(aq) → 2 [Al(H2O)6]3+(aq)
    • with base: Al2O3 + 3 H2O + 2 OH(aq) → 2 [Al(OH)4](aq)
  • Lead(II) oxide
    • with acid: PbO + 2 HCl → PbCl2 + H2O
    • with base: PbO + 2 NaOH + H2O → Na2[Pb(OH)4]

Some other elements which form amphoteric oxides are gallium, indium, scandium, titanium, zirconium, vanadium, chromium, iron, cobalt, copper, silver, gold, germanium, tin, antimony and bismuth, tellurium.[3]

Amphiprotic molecules

According to the Brønsted-Lowry theory of acids and bases: acids are proton donors and bases are proton acceptors.[4] An amphiprotic molecule (or ion) can either donate or accept a proton, thus acting either as an acid or a base. Water, amino acids, hydrogen carbonate ions and hydrogen sulfate ions are common examples of amphiprotic species. Since they can donate a proton, all amphiprotic substances contain a hydrogen atom. Also, since they can act like an acid or a base, they are amphoteric.


A common example of an amphiprotic substance is the hydrogen carbonate ion, which can act as a base:

HCO3 + H3O+ → H2CO3 + H2O

or as an acid:

HCO3 + OH → CO32− + H2O

Thus, it can effectively accept or donate a proton.

Water is the most common example, acting as a base when reacting with an acid such as hydrogen chloride:

H2O + HCl → H3O+ + Cl,

and acting as an acid when reacting with a base such as ammonia:

H2O + NH3 → NH4+ + OH

Not all amphoteric substances are amphiprotic

Although an amphiprotic species must be amphoteric, the converse is not true. For example, the metal oxide ZnO contains no hydrogen and cannot donate a proton. Instead it is a Lewis acid whose Zn atom accepts an electron pair from the base OH. The other metal oxides and hydroxides mentioned above also function as Lewis acids rather than Brønsted acids.

See also


  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "amphoteric".
  2. ^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. pp. 173–4.  
  3. ^ CHEMIX School & Lab - Software for Chemistry Learning, by Arne Standnes (program download required)
  4. ^ R.H. Petrucci, W.S. Harwood, and F.G. Herring, "General Chemistry" (8th edn, Prentice-Hall 2002), p.669
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