Acids and bases
Acceptor number Acid Acid–base reaction Acid–base homeostasis Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Donor number Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self-ionization of water Titration Lewis acid catalysis Frustrated Lewis pair Chiral Lewis acid
Acid types
Brønsted–Lowry Lewis Mineral Organic Oxide Strong Superacids Weak Solid
Base types
Brønsted–Lowry Lewis Organic Oxide Strong Superbases Non-nucleophilic Weak
v t e

The Brønsted–Lowry theory (also called proton theory of acids and bases^[1]) is an acid–base reaction theory which was first developed by Johannes Nicolaus Brønsted and Thomas Martin Lowry independently in 1923.^[2][3] The basic concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H⁺). This theory generalises the Arrhenius theory.

Definitions of acids and bases

Johannes Nicolaus Brønsted and Thomas Martin Lowry, independently, formulated the idea that acids donate protons (H⁺) while bases accept protons.

In the Arrhenius theory, acids are defined as substances that dissociate in aqueous solutions to give H⁺ (hydrogen ions or protons), while bases are defined as substances that dissociate in aqueous solutions to give OH⁻ (hydroxide ions).^[4]

In 1923 physical chemists Johannes Nicolaus Brønsted in Denmark and Thomas Martin Lowry in England both independently proposed the theory named after them.^[5][6][7] In the Brønsted–Lowry theory acids and bases are defined by the way they react with each other, generalising them. This is best illustrated by an equilibrium equation.

acid + base ⇌ conjugate base + conjugate acid.

With an acid, HA, the equation can be written symbolically as:

${\displaystyle {\ce {HA + B <=> A- + HB+}}}$

The equilibrium sign, ⇌, is used because the reaction can occur in both forward and backward directions (is reversible). The acid, HA, is a proton donor which can lose a proton to become its conjugate base, A⁻. The base, B, is a proton acceptor which can become its conjugate acid, HB⁺. Most acid–base reactions are fast, so the substances in the reaction are usually in dynamic equilibrium with each other.^[8]

Aqueous solutions

Consider the following acid–base reaction:

${\displaystyle {\ce {CH3 COOH + H2O <=> CH3 COO- + H3O+}}}$

Acetic acid, CH₃COOH, is an acid because it donates a proton to water (H₂O) and becomes its conjugate base, the acetate ion (CH₃COO⁻). H₂O is a base because it accepts a proton from CH₃COOH and becomes its conjugate acid, the hydronium ion, (H₃O⁺).^[9]

The reverse of an acid–base reaction is also an acid–base reaction, between the conjugate acid of the base in the first reaction and the conjugate base of the acid. In the above example, ethanoate is the base of the reverse reaction and hydronium ion is the acid.

${\displaystyle {\ce {H3O+ + CH3 COO- <=> CH3COOH + H2O}}}$

One feature of the Brønsted–Lowry theory in contrast to Arrhenius theory is that it does not require an acid to dissociate.

Amphoteric substances

The essence of Brønsted–Lowry theory is that an acid is only such in relation to a base, and vice versa. Water is amphoteric as it can act as an acid or as a base. In the image shown at the right one molecule of H₂O acts as a base and gains H⁺ to become H₃O⁺ while the other acts as an acid and loses H⁺ to become OH⁻.

Another example is illustrated by substances like aluminium hydroxide, Al(OH)₃.

${\displaystyle {\ce {{\overset {(acid)}{Al(OH)3}}{}+ OH- <=> Al(OH)4^-}}}$

${\displaystyle {\ce {3H+{}+ {\overset {(base)}{Al(OH)3}}<=> 3H2O{}+ Al_{(aq)}^3+}}}$

Non-aqueous solutions

The hydrogen ion, or hydronium ion, is a Brønsted–Lowry acid when dissolved in H₂O and the hydroxide ion is a base because of the autoionization of water reaction

${\displaystyle {\ce {H2O + H2O <=> H3O+ + OH-}}}$

An analogous reaction occurs in liquid ammonia

${\displaystyle {\ce {NH3 + NH3 <=> NH4+ + NH2-}}}$

Thus, the ammonium ion, NH+4, in liquid ammonia corresponds to the hydronium ion in water and the amide ion, NH−2 in ammonia, to the hydroxide ion in water. Ammonium salts behave as acids, and amides behave as bases.^[10]

Some non-aqueous solvents can behave as bases, i.e. accept protons, in relation to Brønsted–Lowry acids.

${\displaystyle {\ce {HA + S <=> A- + SH+}}}$

where S stands for a solvent molecule. The most important of such solvents are dimethylsulfoxide, DMSO, and acetonitrile, CH₃CN, as these solvents have been widely used to measure the acid dissociation constants of carbon-containing molecules. Because DMSO accepts protons more strongly than H₂O the acid becomes stronger in this solvent than in water.^[11] Indeed, many molecules behave as acids in non-aqueous solutions but not in aqueous solutions. An extreme case occurs with carbon acids, where a proton is extracted from a C−H bond.^{[page needed]}

Some non-aqueous solvents can behave as acids. An acidic solvent will make dissolved substances more basic. For example, the compound CH₃COOH is known as acetic acid since it behaves as an acid in water. However it behaves as a base in liquid hydrogen chloride, a much more acidic solvent.^[12]

Zdroj:https://en.wikipedia.org?pojem=Brønsted–Lowry_acid–base_theory
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