Bronsted Lowry Acid and Base Theory

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Bronsted Lowry Acid and Base
A Bronsted Lowry acid is a proton or hydrogen donor, while a Bronsted Lowry base is a proton or hydrogen acceptor.

The Bronsted Lowry acid and base theory states that an acid donates a proton (hydrogen ion, H+), while a base accepts a proton. The reaction forms the conjugate base of the acid and the conjugate acid of the base. Other names for the theory are the Brønsted–Lowry theory or proton theory of acids and bases. Johannes Nicolaus Brønsted and Thomas Martin Lowry independently outlined the theory in 1923 as a generalization of the Arrhenius theory of acids and bases.

  • The Brønsted–Lowry theory defines acids as proton donors and bases as proton acceptors.
  • A proton is essentially a H+ ion, so all Bronsted Lowry acids contain hydrogen.
  • Acids and bases exist as conjugate pairs. When the acid donates a proton, it forms its conjugate base. When a base accepts a proton, it forms its conjugate acid.
  • Some compounds act as either an acid or a base, depending on the reaction. Compounds which are both acids and bases are amphoteric.

Defining Bronsted Lowry Acids and Bases

According to the Bronsted Lowry theory, an acid is a proton donor. Since a proton is essentially the H+ ion, all Bronsted-Lowry acids contain hydrogen. A base is a proton acceptor. When the acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it forms its conjugate acid. An amphoteric compound is species that can either donate or accept a proton.

For example, consider the reaction between hydrochloric acid (HCl) and ammonia (NH3) that forms the ammonium ion (NH4+) and chloride ion (Cl–).

HCl(aq) + NH3(aq) → NH4+(aq) + Cl–(aq)

In this reaction, HCl donates a hydrogen to NH3. HCl is the Bronsted Lowry acid and NH3 is the Bronsted Lowry base. When HCl donates its proton, it forms its conjugate base, Cl–. When NH3 accepts a proton, it forms its conjugate acid, NH4+. So, the reaction contains two conjugate pairs:

  • HCl (acid) and Cl– (conjugate base)
  • NH3​ (base) and NH4+ (conjugate acid)

Strong and Weak Bronsted Lowry Acids and Bases

An acid or base is either strong or weak.

A strong acid or base fully dissociates into its ion in its solvent, which is usually water. All of a strong acid converts into its conjugate base, while all of a strong base converts into its conjugate acid. The conjugate base of a strong acid is a very weak base. The conjugate acid of a strong base is a very weak acid. Examples of strong Bronsted Lowry acids include hydrochloric acid (HCl), nitric acid (HNO3), sulfuric acid (H2SO4), and hydrobromic acid (HBr). Examples of strong bases include sodium hydroxide (NaOH), potassium hydroxide (KOH), lithium hydroxide (LiOH), and calcium hydroxide (Ca(OH2)).

A weak acid or base incompletely dissociates, reaching an equilibrium condition where both the weak acid and its conjugate base or weak base and its conjugate acid both remain in solution. Examples of weak Bronsted Lowry acids include phosphoric acid (H3PO4), nitrous acid (HNO2), and acetic acid (CH3COOH). Examples of weak bases include ammonia (NH3), copper hydroxide (Cu(OH)2), and methylamine (CH₃NH₂).

Remember that water is amphoteric and acts as an acid in some reactions and as a base in other reactions. When you dissolve a strong acid in water, the water acts as a base. When you dissolve a strong base in water, the water acts as an acid.

For example:

HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)

The conjugate pairs are as follows:

  • HCl (acid) and Cl- (conjugate base)
  • H2O (base) and H3O+ (conjugate acid)

NaOH(s) + H2O(l) → Na+(aq) + OH–(aq)

The conjugate pairs are as follows:

  • NaOH (base) and Na+ (conjugate acid)
  • H2O (acid) and OH– (conjugate base)
To calculate pH, take the log of the hydrogen ion concentration and change the sign of the answer.

How to Find pH

Acids have a pH under 7, while bases have a pH above 7. Here’s how you find the value.

Comparison to Arrhenius Acids and Bases

The Bronsted Lowry theory is less restrictive than the Arrhenius theory of acids and bases. For one thing, it allows for solvents other than water. Another difference relates to the defining properties of acids and bases. According to the Arrhenius theory, acids increase hydrogen ion (H+) concentration in water, while bases increase hydroxide ion (OH–) concentration in water. Bronsted Lowry theory allows for bases that do not contain OH or at least form its ion in water. For example, ammonia (NH3) is an Arrhenius base because even though it does not contain OH, it increases the concentration of hydroxide ions in water. Ammonia is also a Bronsted Lowry base. However, methylamine (CH₃NH₂) is a Bronsted Lowry base, but not an Arrhenius base. It neither contains hydroxide nor raises its ion concentration in water.

Mostly, the list of Arrhenius and Bronsted Lowry acids is the same, but there are exceptions. For example, dimethylamine [(CH3)2NH] is never an Arrhenius acid because its pKa value is lower than water. It does not increase H+ or H3O+ concentration in water. It’s usually a Bronsted Lowry base, but it can be a Bronsted Lowry acid. Dimethylamine can donate a proton when it reacts with a sufficiently strong base, such as butyllithium (C4H9Li)

Comparison to Lewis Acids and Bases

Gilbert Lewis proposed the Lewis theory of acids and bases the same year that Bronsted and Lowry published their theories. The big difference between the two theories is that the Bronsted Lowry theory deals with protons, while the Lewis theory focuses on electrons. According to the Lewis theory, an acid is an electron pair receptor, while a base is an electron pair donor. Both theories include conjugate acids and bases.

All Bronsted Lowry acids are Lewis acids, but not all Lewis acids are Bronsted Lowry acids. The Lewis theory allows for acids that do not contain hydrogen atoms. For example, BF3 and AlCl3 are Lewis acids, but not Bronsted Lowry acids.

References

  • 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
  • Hall, Norris F. (March 1940). “Systems of Acids and Bases”. Journal of Chemical Education. 17 (3): 124–128. doi:10.1021/ed017p124
  • Lowry, T. M. (1923). “The uniqueness of hydrogen”. Journal of the Society of Chemical Industry. 42 (3): 43–47. doi:10.1002/jctb.5000420302
  • Masterton, William; Hurley, Cecile; Neth, Edward (2011). Chemistry: Principles and Reactions. Cengage Learning. ISBN 978-1-133-38694-0.
  • Myers, Richard (2003). The Basics of Chemistry. Greenwood Publishing Group. ISBN 978-0-313-31664-7.