|Acids and bases|
In chemistry, an amphoteric compound is a molecule or ion that can react both as an acid as well as a base. Many metals (such as copper, zinc, tin, lead, aluminium, and beryllium) form amphoteric oxides or hydroxides. Amphoterism depends on the oxidation states 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.
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 is derived from the Greek word amphoteroi (ἀμφότεροι) meaning "both". Related words in acid-base chemistry are amphichromatic and amphichroic, both describing substances such as acid-base indicators which give one colour on reaction with an acid and another colour on reaction with a base.
According to the Brønsted-Lowry theory of acids and bases: acids are proton donors and bases are proton acceptors. 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.
Zinc oxide (ZnO) reacts with both acids and with bases:
- In acid: ZnO + H2SO4 → ZnSO4 + H2O
- In base: ZnO + 2 NaOH + H2O → Na2[Zn(OH)4]
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.
Lead oxide (PbO):
- In acid: PbO + 2 HCl → PbCl2 + H2O
- In base: PbO + 2 NaOH + H2O → Na2Pb[(OH)3]
Aluminium oxide (Al2O3)
- In acid: Al2O3 + 6 HCl→ 2 AlCl3 + H2O
- In base: Al2O3 +2 NaOH + H2O → 2 Na[Al(OH)4]
Stannous Oxide (SnO)
In acid : SnO +2HCl = SnCl2 + H2O
In base : SnO +4NaOH + H2O = Na4[Sn(OH)6]
Some other elements which form amphoteric oxides are gallium, indium, scandium, titanium, zirconium, vanadium, chromium, iron, cobalt, copper, silver, gold, germanium, tin, antimony, bismuth, and tellurium
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:
- with acid: Be(OH)2 + 2 HCl → BeCl2 + 2 H2O
- with base: Be(OH)2 + 2 NaOH → Na2[Be(OH)4].
|Wikimedia Commons has media related to Amphoteric oxides.|
- IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "amphoteric".
- Penguin Science Dictionary 1994, Penguin Books
- R.H. Petrucci, W.S. Harwood, and F.G. Herring, "General Chemistry" (8th edn, Prentice-Hall 2002), p.669
- Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. pp. 173–4. ISBN 978-0130399137.
- CHEMIX School & Lab - Software for Chemistry Learning, by Arne Standnes[permanent dead link] (program download required)