Amphoterism

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In chemistry, an amphoteric compound is a molecule or ion that can react both as an acid and as a base.[1] What exactly this can mean depends on which definitions of acids and bases are being used. The prefix of the word 'amphoteric' is derived from a Greek prefix amphi which means "both".

One type of amphoteric species are amphiprotic molecules, which can either donate or accept a proton (H+). This is what "amphoteric" means in Brønsted–Lowry acid–base theory. 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. For example, an amino acid H2N–RCH–CO2H has both a basic group NH2 and an acidic group COOH, and exists as several structures in chemical equilibrium:

H2N–RCH–CO2H + H2O ⇌ H2N–RCH–COO + H3O+ ⇌ H3N+–RCH–COOH + OH ⇌ H3N+–RCH–COO + H2O.

In approximately neutral aqueous solution (pH ≅ 7), the basic amino group is mostly protonated and the carboxylic acid is mostly deprotonated, so that the predominant species is the zwitterion H3N+–RCH–COO. 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.

Metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides. Many metals (such as zinc, tin, lead, aluminium, and beryllium) form amphoteric oxides or hydroxides. Al2O3 is an example of an amphoteric oxide. Amphoterism depends on the oxidation states of the oxide. Amphoteric oxides include lead (II) oxide and zinc (II) oxide, among many others.[2]

Etymology[edit]

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.[3]

Amphiprotic molecules[edit]

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 ion (or bicarbonate ion) HCO3, dihydrogen phosphate ion H2PO4, and hydrogen sulfate ion (or bisulfate ion) HSO4 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.

Examples[edit]

The water molecule is amphoteric in aqueous solution. It can either gain a proton to form a hydronium ion H3O+, or else lose a proton to form a hydroxide ion OH.[5]

Another possibility is the molecular autoionization reaction between two water molecules, in which one water molecule acts as an acid and another as a base.

H2O + H2O ⇌ H3O+ + OH

The bicarbonate ion, HCO3, is amphoteric as it can act as either an acid or a base:

As an acid, losing a proton: HCO3 + OH ⇌ CO32− + H2O
As a base, accepting a proton: HCO3 + H+ ⇌ H2CO3

Note: in dilute aqueous solution the formation of the hydronium ion, H3O+(aq), is effectively complete, so that hydration of the proton can be ignored in relation to the equilibria.

Other examples of inorganic polyprotic acids include anions of sulphuric acid, phosphoric acid, EDTA and hydrogen sulphide that have lost one or more protons. In organic chemistry and biochemistry, important examples include amino acids and derivatives of citric acid.

Although an amphiprotic species must be amphoteric, the converse is not true. For example, a metal oxide such as zinc oxide, ZnO, contains no hydrogen and so cannot donate a proton. Nevertheless, it can act as an acid by reacting with the hydroxide ion, a base:

ZnO(s) + 2OH + H2O → Zn(OH)42- (aq)

This reaction is not covered by Brønsted–Lowry acid–base theory. As zinc oxide can also act as a base

ZnO(s) + 2H+ → Zn2+(aq) + H2O

it is classified as amphoteric rather than amphiprotic.

Oxides[edit]

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, for instance 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 → Na2[Pb(OH)4]

Aluminium oxide (Al2O3):

  • In acid: Al2O3 + 6 HCl → 2 AlCl3 + 3 H2O
  • In base: Al2O3 + 2 NaOH + 3 H2O → 2 Na[Al(OH)4] (hydrated sodium aluminate)

Stannous oxide (SnO):

  • In acid : SnO + 2 HCl ⇌ SnCl2 + H2O
  • In base : SnO + 4 NaOH + H2O ⇌ Na4[Sn(OH)6]

Vanadium dioxide (VO2):

  • In acid: VO2 + 2 HCl → VOCl2 + H2O
  • In base: 4 VO2 + 2 NaOH → Na2V4O9 + H2O

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

Hydroxides[edit]

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]

Beryllium hydroxide:

  • with acid: Be(OH)2 + 2 HCl → BeCl2 + 2 H2O
  • with base: Be(OH)2 + 2 NaOH → Na2[Be(OH)4].[6]

Chromium hydroxide:

  • with acid:
  • with base:

See also[edit]

References[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "amphoteric". doi:10.1351/goldbook.A00306
  2. ^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. pp. 173–4. ISBN 978-0-13-039913-7.
  3. ^ Penguin Science Dictionary 1994, Penguin Books
  4. ^ Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General chemistry: principles and modern applications (8th ed.). Upper Saddle River, N.J: Prentice Hall. p. 669. ISBN 978-0-13-014329-7. LCCN 2001032331. OCLC 46872308.
  5. ^ Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2014). Fundamentals of analytical chemistry (Ninth ed.). Belmont, CA. p. 200. ISBN 978-0-495-55828-6. OCLC 824171785.
  6. ^ CHEMIX School & Lab - Software for Chemistry Learning, by Arne Standnes (program download required)