Brønsted–Lowry acid–base theory

The Brønsted–Lowry theory (also called proton theory of acids and bases^[1]) is an acid–base reaction theory which was proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923.^[2]^[3] The fundamental 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 is a generalization of the Arrhenius theory.

Definitions of acids and bases[]

Johannes Nicolaus Brønsted and Thomas Martin Lowry, independently, formulated the idea that acids are proton (H⁺) donors while bases are proton acceptors.

In the Arrhenius theory, acids are defined as substances that dissociate in aqueous solution to give H⁺ (hydrogen ions), while bases are defined as substances that dissociate in aqueous solution 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 that carries their names.^[5]^[6]^[7] In the Brønsted–Lowry theory acids and bases are defined by the way they react with each other, which allows for greater generality. The definition is expressed in terms of an equilibrium expression

acid + base ⇌ conjugate base + conjugate acid.

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

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

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

Aqueous solutions[]

Acetic acid, CH3COOH, is composed of a methyl group, CH3, bound chemically to a carboxylate group, COOH. The carboxylate group can lose a proton and donate it to a water molecule, H2O, leaving behind an acetate anion CH3COO− and creating a hydronium cation H3O+. This is an equilibrium reaction, so the reverse process can also take place.

Acetic acid, a weak acid, donates a proton (hydrogen ion, highlighted in green) to water in an equilibrium reaction to give the acetate ion and the hydronium ion. Red: oxygen, black: carbon, white: hydrogen.

Consider the following acid–base reaction:

{\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, acetate is the base of the reverse reaction and hydronium ion is the acid.

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

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

Amphoteric substances[]

The amphoteric nature of water

The essence of Brønsted–Lowry theory is that an acid only exists as 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 furnished by substances like aluminium hydroxide, Al(OH)₃.

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

, acting as an acid

{\ce {3H+ + Al (OH)3<=> 3H2O + Al^{3+}(aq)}}

, acting as a base

Non-aqueous solutions[]

The hydrogen ion, or hydronium ion, is a Brønsted–Lowry acid in aqueous solutions, and the hydroxide ion is a base, by virtue of the self-dissociation reaction

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

An analogous reaction occurs in liquid ammonia

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

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

Some non-aqueous solvents can behave as bases, that is, proton acceptors, in relation to Brønsted–Lowry acids.

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

where S stands for a solvent molecule. The most important such solvents are dimethylsulfoxide, DMSO, and acetonitrile, CH₃CN, as these solvents have been widely used to measure the acid dissociation constants of organic molecules. Because DMSO is a stronger proton acceptor than H₂O the acid becomes a stronger acid in this solvent than in water.^[11] Indeed, many molecules behave as acids in non-aqueous solution that do not do so in aqueous solution. An extreme case occurs with carbon acids, where a proton is extracted from a C-H bond.

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

{\ce {HCl + CH3COOH <=> Cl- + CH3C(OH)2+}}

Comparison with Lewis acid–base theory[]

In the same year that Brønsted and Lowry published their theory, G. N. Lewis proposed an alternative theory of acid–base reactions. The Lewis theory is based on electronic structure. A Lewis base is defined as a compound that can donate an electron pair to a Lewis acid, a compound that can accept an electron pair.^[13]^[14] Lewis's proposal gives an explanation to the Brønsted–Lowry classification in terms of electronic structure.

{\ce {HA + B <=> A- + BH+}}

In this representation both the base, B, and the conjugate base, A⁻, are shown carrying a lone pair of electrons and the proton, which is a Lewis acid, is transferred between them.

Adduct of ammonia and boron trifluoride

Lewis later wrote in "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen."^[14] In Lewis theory an acid, A, and a base, B:, form an adduct, AB, in which the electron pair is used to form a dative covalent bond between A and B. This is illustrated with the formation of the adduct H₃N−BF₃ from ammonia and boron trifluoride, a reaction that cannot occur in aqueous solution because boron trifluoride reacts violently with water in a hydrolysis reaction.

{\ce {BF3 + 3H2O -> B (OH)3 + 3HF}}

{\ce {HF <=> H+ + F-}}

These reactions illustrate that BF₃ is an acid in both Lewis and Brønsted–Lowry classifications and emphasizes the consistency between both theories.^{[citation needed]}

Boric acid is recognized as a Lewis acid by virtue of the reaction

{\ce {B(OH)3 + H2O <=> B(OH)4- + H+}}

In this case the acid does not dissociate, it is the base, H₂O that dissociates. A solution of B(OH)₃ is acidic because hydrogen ions are liberated in this reaction.

There is strong evidence that dilute aqueous solutions of ammonia contain negligible amounts of the ammonium ion

{\ce {H2O + NH3 -> OH- + NH+4}}

and that, when dissolved in water, ammonia functions as a Lewis base.^[15]

Comparison with the Lux–Flood theory[]

The reactions between oxides in the solid or liquid state are not included in Brønsted–Lowry theory. For example, the reaction

{\ce {2MgO + SiO2 -> Mg2 SiO4}}

does not fall within the scope of the Brønsted–Lowry definition of acids and bases. On the other hand, magnesium oxide acts as a base when it reacts with an aqueous solution of an acid.

{\ce {2H+ + MgO(s) -> Mg^{2+}(aq) + 2H2O}}

Dissolved SiO₂ has been predicted to be a weak acid in the Brønsted–Lowry sense.^[16]

{\ce {SiO2(s) + 2H2O <=> Si(OH)4 (solution)}}

{\ce {Si(OH)4 <=> Si(OH)3O- + H+}}

According to the Lux–Flood theory compounds such as MgO and SiO₂ in the solid state may be classified as acids or bases. For example, the mineral olivine may be regarded as a compound of a basic oxide, MgO, with an acidic oxide, silica, SiO₂. This classification is important in geochemistry.

References[]

^ "Brønsted–Lowry theory | chemistry". Encyclopedia Britannica. Retrieved 2021-03-07.
^ Brönsted, J. N. (1923). "Einige Bemerkungen über den Begriff der Säuren und Basen" [Some observations about the concept of acids and bases]. Recueil des Travaux Chimiques des Pays-Bas. 42 (8): 718–728. doi:10.1002/recl.19230420815.
^ Lowry, T. M. (1923). "The uniqueness of hydrogen". Journal of the Society of Chemical Industry. 42 (3): 43–47. doi:10.1002/jctb.5000420302.
^ Myers, Richard (2003). The Basics of Chemistry. Greenwood Publishing Group. pp. 157–161. ISBN 978-0-313-31664-7.
^ Masterton, William; Hurley, Cecile; Neth, Edward (2011). Chemistry: Principles and Reactions. Cengage Learning. p. 433. ISBN 978-1-133-38694-0.
^ Ebbing, Darrell; Gammon, Steven D. (2010). General Chemistry, Enhanced Edition. Cengage Learning. pp. 644–645. ISBN 978-0-538-49752-7.
^ Whitten, Kenneth; Davis, Raymond; Peck, Larry; Stanley, George (2013). Chemistry. Cengage Learning. p. 350. ISBN 978-1-133-61066-3.
^ Lew, Kristi (2009). Acids and Bases. Infobase Publishing. ISBN 9780791097830.
^ Patrick, Graham (2012). Instant Notes in Organic Chemistry. Taylor & Francis. p. 76. ISBN 978-1-135-32125-3.
^ Holliday, A.K.; Massy, A.G. (1965). Inorganic Chemistry in Non-Aqueous Solvents. Pergamon Press.
^ Reich, Hans J. "Bordwell pKa Table (Acidity in DMSO)". Department of Chemistry, University of Wisconsin, U.S. Archived from the original on 9 October 2008. Retrieved 2008-11-02.
^ Waddington, T.C. (1965). Non-Aqueous Solvent Systems. New York: Academic Press.
^ Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172
^ ^a ^b Hall, Norris F. (March 1940). "Systems of Acids and Bases". Journal of Chemical Education. 17 (3): 124–128. Bibcode:1940JChEd..17..124H. doi:10.1021/ed017p124.
^ Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 187. ISBN 978-0-13-039913-7.
^ Pauling, Linus (1960). The Nature of the Chemical Bond (3rd ed.). Ithaka: Cornell University Press. p. 557.

[1] "Brønsted–Lowry theory | chemistry". Encyclopedia Britannica. Retrieved 2021-03-07.

[2] Brönsted, J. N. (1923). "Einige Bemerkungen über den Begriff der Säuren und Basen" [Some observations about the concept of acids and bases]. Recueil des Travaux Chimiques des Pays-Bas. 42 (8): 718–728. doi:10.1002/recl.19230420815.

[3] Lowry, T. M. (1923). "The uniqueness of hydrogen". Journal of the Society of Chemical Industry. 42 (3): 43–47. doi:10.1002/jctb.5000420302.

[4] Myers, Richard (2003). The Basics of Chemistry. Greenwood Publishing Group. pp. 157–161. ISBN 978-0-313-31664-7.

[5] Masterton, William; Hurley, Cecile; Neth, Edward (2011). Chemistry: Principles and Reactions. Cengage Learning. p. 433. ISBN 978-1-133-38694-0.

[6] Ebbing, Darrell; Gammon, Steven D. (2010). General Chemistry, Enhanced Edition. Cengage Learning. pp. 644–645. ISBN 978-0-538-49752-7.

[7] Whitten, Kenneth; Davis, Raymond; Peck, Larry; Stanley, George (2013). Chemistry. Cengage Learning. p. 350. ISBN 978-1-133-61066-3.

[8] Lew, Kristi (2009). Acids and Bases. Infobase Publishing. ISBN 9780791097830.

[9] Patrick, Graham (2012). Instant Notes in Organic Chemistry. Taylor & Francis. p. 76. ISBN 978-1-135-32125-3.

[10] Holliday, A.K.; Massy, A.G. (1965). Inorganic Chemistry in Non-Aqueous Solvents. Pergamon Press.

[11] Reich, Hans J. "Bordwell pKa Table (Acidity in DMSO)". Department of Chemistry, University of Wisconsin, U.S. Archived from the original on 9 October 2008. Retrieved 2008-11-02.

[12] Waddington, T.C. (1965). Non-Aqueous Solvent Systems. New York: Academic Press.

[13] Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172

[review1940-14] Hall, Norris F. (March 1940). "Systems of Acids and Bases". Journal of Chemical Education. 17 (3): 124–128. Bibcode:1940JChEd..17..124H. doi:10.1021/ed017p124.

[15] Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 187. ISBN 978-0-13-039913-7.

[16] Pauling, Linus (1960). The Nature of the Chemical Bond (3rd ed.). Ithaka: Cornell University Press. p. 557.

[1]

[2]

[3]

[4]

[5]

[6]

[7]

[8]

[9]

[10]

[11]

[12]

[13]

[14]

[15]

[16]