Ampholyte

Acid-base ampholytes (composed of Greek αμφίς ( amphis ) = on both sides and λύσις ( lysis ) = resolution) or acid-base Amphoteric or amphiprotic compounds are chemical compounds having both as a Brønsted - acid as a Brønsted base can react. This behavior is also known as acid-base amphoteric . Amphoterics can both accept and release protons.

properties

The water solubility of the ampholytes depends strongly on the pH value. Some ampholytes react with themselves, the best known example of this is water . It reacts with acids to form H ₃ O ⁺ or with bases to form OH ^- , this behavior is also shown in pure water as autoprotolysis :

${\displaystyle \mathrm {2\ H_{2}O\ \rightleftharpoons \ H_{3}O^{+}+OH^{-}} }$

Examples of ampholytes

Compounds prone to autoprotolysis

Examples (autoprotolysis constants pK _au nach):

• Water H ₂ O (pK _au = 14)
• Ammonia NH ₃ (pK _au = 29 at −50 ° C)
• Sulfuric acid H ₂ SO ₄ (pK _au = 3.85)
• Acetic acid CH ₃ COOH (pK _au = 14.45)
• Formic acid HCOOH (pK _au = 6.2)
• Methanol CH ₃ OH (pK _au = 16.9)
• Ethanol CH ₃ CH ₂ OH (pK _au = 19.5)
• Hydrogen fluoride HF (pK _au = 10.7 at 0 ° C)

The specified autoprotolysis constants correspond to the negative decadic logarithm (see also pH value ) of the ionic product of the substances. The extent of autoprotolysis usually increases with increasing temperature.

Reaction example : water

Reacts with acid as a base:

${\displaystyle \mathrm {HCl+H_{2}O\longrightarrow H_{3}O^{+}+Cl^{-}} }$

Reacts with base as an acid:

${\displaystyle \mathrm {NH_{3}+H_{2}O\longrightarrow NH_{4}^{+}+OH^{-}} }$

Partially deprotonated multi-protonic acids

Examples:

• Monohydrogen phosphate HPO ₄ ²⁻
• Dihydrogen phosphate H ₂ PO ₄ ^-
• Hydrogen sulfate HSO ₄ ^-
• Hydrogen carbonate HCO ₃ ^-

Reaction example: dihydrogen phosphate

Reacts with acid as a base:

${\displaystyle \mathrm {HCl+H_{2}PO_{4}^{-}\longrightarrow H_{3}PO_{4}+Cl^{-}} }$

Reacts with base as an acid:

${\displaystyle \mathrm {NH_{3}+H_{2}PO_{4}^{-}\longrightarrow NH_{4}^{+}+HPO_{4}^{2-}} }$

Partially protonated polyvalent bases

Examples:

• basic magnesium chloride Mg (OH) Cl or Mg (OH) ⁺ Cl ^-
• Hydrazine monohydrochloride H ₂ N-NH ₂ · HCl or H ₂ N-NH ₃ ⁺ Cl ^-

Reaction example : basic magnesium chloride

Reacts with acid as a base:

${\displaystyle \mathrm {HCl+Mg(OH)Cl\longrightarrow H_{2}O+MgCl_{2}} }$

Reacts with base as an acid:

${\displaystyle \mathrm {Mg(OH)Cl+NaOH\longrightarrow NaCl+Mg(OH)_{2}} }$

Compounds with acidic and basic functional groups

Compounds with at least one acidic and one basic functional group are also amphoteric substances, for example:

• Amino acids with their acidic carboxy groups and basic amino groups (and thus also peptides and most proteins )
• Zwitterions

Reaction example : glycine (simplest amino acid)

Reacts with acid as a base:

${\displaystyle \mathrm {HCl+H_{2}N{-}CH_{2}{-}COOH\longrightarrow H_{3}N^{+}{-}CH_{2}{-}COOH+Cl^{-}} }$

Reacts with base as an acid:

${\displaystyle \mathrm {NaOH+H_{2}N{-}CH_{2}{-}COOH\longrightarrow H_{2}O+H_{2}N{-}CH_{2}{-}COO^{-}+Na^{+}} }$

Calculating the intrinsic pH of ampholytes

If ampholytes (with two functional groups) are dissolved in water, an average pH value is established , which can be calculated using the following approximation formula, also known as the "ampholyte equation" (for not too strong dilutions).

${\displaystyle pH={\frac {1}{2}}\ (pK_{S1}+pK_{S2})={\frac {1}{2}}\ (pK_{S1}+14-pK_{B2})}$

Here, pK are _S1 and pK _S2 , the acidity constants (pK _S values) of the respective Dissoziationsmöglichkeiten of the ampholyte.

Electrically neutral ampholytes, e.g. B. Amino acids also have the lowest solubility at this pH value; if the pH falls or rises, the solubility increases again, since the solvation shell is stabilized with the charge .

See also

• Dissociation (chemistry)
• Protolysis

Web links

Wiktionary: Ampholyt  - explanations of meanings, word origins, synonyms, translations

Individual evidence

1. ↑ Lothar Kolditz : Inorganic Chemistry. Volume 1. 2nd edition. VEB Deutscher Verlag der Wissenschaften, Berlin 1983, p. 188.
2. ^ AF Holleman , E. Wiberg , N. Wiberg : Textbook of Inorganic Chemistry . 101st edition. Walter de Gruyter, Berlin 1995, ISBN 3-11-012641-9 , p. 457.