Brønsted–Lowry acid–base theory

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

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In the Arrhenius theory, acids are defined as substances that dissociate in aqueous solutions to give H⁺ (hydrogen cations or protons), while bases are defined as substances that dissociate in aqueous solutions to give OH⁻ (hydroxide ions).{{cite book|first=Richard |last=Myers|title=The Basics of Chemistry|url=https://archive.org/details/basicsofchemistr0000myer|url-access=registration |date=2003|publisher=Greenwood Publishing Group|pages=[https://archive.org/details/basicsofchemistr0000myer/page/157 157]–161|isbn=978-0-313-31664-7}}

In 1923, physical chemists Johannes Nicolaus Brønsted in Denmark and Thomas Martin Lowry in England both independently proposed the theory named after them.{{cite book|first1=William |last1=Masterton|first2=Cecile |last2=Hurley|first3=Edward |last3=Neth|title=Chemistry: Principles and Reactions|url=https://books.google.com/books?id=udUIAAAAQBAJ&pg=PT433|date=2011|publisher=Cengage Learning|page=433|isbn=978-1-133-38694-0}}{{cite book|first1=Darrell |last1=Ebbing|first2=Steven D. |last2=Gammon|title=General Chemistry, Enhanced Edition|url=https://books.google.com/books?id=8_7iGj4QLR4C&pg=PA644|date=2010|publisher=Cengage Learning|isbn=978-0-538-49752-7|pages=644–645}}{{cite book|first1=Kenneth |last1=Whitten|first2=Raymond |last2=Davis|first3=Larry |last3=Peck|first4=George |last4=Stanley|title=Chemistry|url=https://books.google.com/books?id=XQiu_E72llkC&pg=PA350|date=2013|publisher=Cengage Learning|page=350|isbn=978-1-133-61066-3}} 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:

:{{chem2|1=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.{{cite book|last1=Lew|first1=Kristi|title=Acids and Bases|date=2009|publisher=Infobase Publishing|isbn=9780791097830|url=https://books.google.com/books?id=QHrzG6CaVIwC}}

File:Acetic-acid-dissociation-3D-balls.png, {{chem2|CH3}}, chemically bound to a carboxyl group, {{chem2|\sCOOH}}. The carboxyl group can lose a proton and donate it to a water molecule, {{chem2|H2O}}, leaving behind an ethanoate anion {{chem2|CH3COO-}} and creating a hydronium cation {{chem2|H3O+}}. This is an equilibrium reaction, so the reverse process can also take place.|Acetic acid, a weak acid, donates a proton (hydron) to water in an equilibrium reaction to give the acetate ion and the hydronium ion.

{{legend|red|Oxygen}}

{{legend|black|Carbon}}

{{legend|white|Hydrogen}}

{{legend|lime|Proton (hydron)}}]]

Consider the following acid–base reaction:

:CH3 COOH + H2O <=> CH3 COO- + H3O+

Acetic acid, {{chem2|CH3COOH}}, is an acid because it donates a proton to water ({{chem2|H2O}}) and becomes its conjugate base, the acetate ion ({{chem2|CH3COO-}}). {{chem2|H2O}} is a base because it accepts a proton from {{chem2|CH3COOH}} and becomes its conjugate acid, the hydronium ion, ({{chem2|H3O+}}).{{Sfn|Patrick|2004|p=76}}

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.

: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.

File:Autoprotolyse eau.svg

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 {{chem2|H2O}} acts as a base and gains {{chem2|H+}} to become {{chem2|H3O+}} while the other acts as an acid and loses {{chem2|H+}} to become {{chem2|OH-}}.

Another example is illustrated by substances like aluminium hydroxide, {{chem2|Al(OH)3}}.

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

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

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

:H2O + H2O <=> H3O+ + OH-

An analogous reaction occurs in liquid ammonia

:NH3 + NH3 <=> NH4+ + NH2-

Thus, the ammonium ion, {{chem2|NH4+}}, in liquid ammonia corresponds to the hydronium ion in water and the amide ion, {{chem2|NH2-}} in ammonia, to the hydroxide ion in water. Ammonium salts behave as acids, and metal amides behave as bases.{{cite book|last1=Holliday|first1=A.K.|last2=Massy|first2=A.G.|title=Inorganic Chemistry in Non-Aqueous Solvents|date=1965|publisher=Pergamon Press}}

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

:HA + S <=> A- + SH+

where S stands for a solvent molecule. The most important of such solvents are dimethylsulfoxide, DMSO, and acetonitrile, {{chem2|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 {{chem2|H2O}} the acid becomes stronger in this solvent than in water.{{cite web |url=http://www.chem.wisc.edu/areas/reich/pkatable/ |title=Bordwell pKa Table (Acidity in DMSO) |access-date=2008-11-02 | archive-url= https://web.archive.org/web/20081009060809/http://www.chem.wisc.edu/areas/reich/pkatable/| archive-date= 9 October 2008|author=Reich, Hans J. |publisher=Department of Chemistry, University of Wisconsin, U.S.}} 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 {{chem2|C\sH}} bond.{{Cite book |last1=Smith |first1=Michael |title=March's advanced organic chemistry: reactions, mechanisms, and structure |last2=March |first2=Jerry |date=2001 |publisher=Wiley |isbn=978-0-471-58589-3 |edition=5th |location=New York |pages=328}}

Some non-aqueous solvents can behave as acids. An acidic solvent will make dissolved substances more basic. For example, the compound {{chem2|CH3COOH}} is known as acetic acid since it behaves as an acid in water. However, it behaves as a base in liquid hydrogen fluoride, a much more acidic solvent.{{cite book |last1=Waddington |first1=T.C. |title=Non-Aqueous Solvent Systems |date=1965 |publisher=Academic Press |location=New York |pages=69}}

:CH3COOH + 2HF <=> CH3C(OH)2+ + HF2-

{{See also|Lewis acids and bases}}

In the same year that Brønsted and Lowry published their theory, G. N. Lewis created an alternative theory of acid–base reactions. The Lewis theory is based on electronic structure. A Lewis base is a compound that can give an electron pair to a Lewis acid, a compound that can accept an electron pair.Miessler, G. L., Tarr, D. A., (1991) "Inorganic Chemistry" 2nd ed. Pearson Prentice-Hall pp. 170–172{{Cite journal|title=Systems of Acids and Bases|last=Hall|first=Norris F.|journal=Journal of Chemical Education|date=March 1940|volume=17|issue=3|doi=10.1021/ed017p124|pages=124–128|bibcode = 1940JChEd..17..124H }} Lewis's proposal explains the Brønsted–Lowry classification using electronic structure.

: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.

File:NH3-BF3-adduct-bond-lengthening-2D.png

Lewis later wrote "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." In Lewis theory an acid, A, and a base, B, form an adduct, AB, where the electron pair forms a dative covalent bond between A and B. This is shown when the adduct H₃N−BF₃ forms from ammonia and boron trifluoride, a reaction that cannot occur in water because boron trifluoride hydrolizes in water.

: 4BF3 + 3H2O -> B(OH)3 + 3HBF4 {{Cite journal |last=Wamser |first=Christian A. |date=1951-01-01 |title=Equilibria in the System Boron Trifluoride—Water at 25° |url=https://pubs.acs.org/doi/abs/10.1021/ja01145a134 |journal=Journal of the American Chemical Society |language=en |volume=73 |issue=1 |pages=409–416 |doi=10.1021/ja01145a134 |bibcode=1951JAChS..73..409W |issn=0002-7863|url-access=subscription }}

The reaction above illustrates that BF₃ is an acid in both Lewis{{Sfn|Patrick|2004|p=91}} and Brønsted–Lowry classifications and shows that the theories agree with each other.{{citation needed|date=August 2014}}

Boric acid is recognised as a Lewis acid because of the reaction

:B(OH)3 + H2O <=> B(OH)4- + H+

In this case the acid does not split up but the base, H₂O, does. A solution of B(OH)₃ is acidic because hydrogen ions are given off in this reaction.

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

:H2O + NH3 -> OH- + NH+4

and that, when dissolved in water, ammonia functions as a Lewis base.{{Housecroft2nd|page=187}}

The reactions between oxides in the solid or liquid states are excluded in the Brønsted–Lowry theory. For example, the reaction

:2MgO + SiO2 -> Mg2 SiO4

is not covered in 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 (instead, the reaction can be considered a lewis acid-base reaction).

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

Dissolved silicon dioxide, SiO₂, has been predicted to be a weak acid in the Brønsted–Lowry sense.{{cite book |author-link=Linus Pauling| last1=Pauling |first1=Linus |title=The Nature of the Chemical Bond |url=https://archive.org/details/natureofchemical00paul |url-access=registration |date=1960 |publisher=Cornell University Press |location=Ithaca |page=[https://archive.org/details/natureofchemical00paul/page/557 557] |edition=3rd}}

:SiO2(s) + 2H2O <=> Si(OH)4 (solution)

:Si(OH)4 <=> Si(OH)3O- + H+

According to the Lux–Flood theory, oxides like MgO and SiO₂ in the solid state may be called acids or bases. For example, the mineral olivine may be known as a compound of a basic oxide, MgO, and silicon dioxide, SiO₂, as an acidic oxide. This is important in geochemistry.{{cite journal |last1=Ottonello |first1=G |last2=Moretti |first2=R |title=Lux-Flood basicity of binary silicate melts |journal= Journal of Physics and Chemistry of Solids|date=2004 |volume=65 |issue=8–9 |pages=1609–1614 |doi=10.1016/j.jpcs.2004.01.012 |bibcode=2004JPCS...65.1609O |url=https://www.sciencedirect.com/science/article/pii/S0022369704000551 |access-date=2 January 2025|url-access=subscription }}

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{{DEFAULTSORT:Bronsted-Lowry acid-base theory}}

Category:Acid–base chemistry

Category:Equilibrium chemistry