Bronsted-Lowry of acid and base theory

• A Brønsted-Lowry acid is any species that is capable of donating a proton—$\text{H}^+$H, start superscript, plus, end superscript.
• A Brønsted-Lowry base is any species that is capable of accepting a proton, which requires a lone pair of electrons to bond to the $\text{H}^+$H, start superscript, plus, end superscript.
• Water is amphoteric, which means it can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base.
• Strong acids and bases ionize completely in aqueous solution, while weak acids and bases ionize only partially.
• The conjugate base of a Brønsted-Lowry acid is the species formed after an acid donates a proton. The conjugate acid of a Brønsted-Lowry base is the species formed after a base accepts a proton.
• The two species in a conjugate acid-base pair have the same molecular formula except the acid has an extra $\text H^+$H, start superscript, plus, end superscript compared to the conjugate base.

Brønsted-Lowry theory of acids and bases

The Brønsted-Lowry theory describes acid-base interactions in terms of proton transfer between chemical species. A Brønsted-Lowry acid is any species that can donate a proton, $\text{H}^+$H, start superscript, plus, end superscript, and a base is any species that can accept a proton. In terms of chemical structure, this means that any Brønsted-Lowry acid must contain a hydrogen that can dissociate as $\text H^+$H, start superscript, plus, end superscript. In order to accept a proton, a Brønsted-Lowry base must have at least one lone pair of electrons to form a new bond with a proton.

Using the Brønsted-Lowry definition, an acid-base reaction is any reaction in which a proton is transferred from an acid to a base. We can use the Brønsted-Lowry definitions to discuss acid-base reactions in any solvent, as well as those that occur in the gas phase. For example, consider the reaction of ammonia gas, $\text{NH}_3(g)$N, H, start subscript, 3, end subscript, left parenthesis, g, right parenthesis, with hydrogen chloride gas, $\text{H}\text{Cl}(g)$H, C, l, left parenthesis, g, right parenthesis, to form solid ammonium chloride, $\text{NH}_4 \text{Cl}(s)$N, H, start subscript, 4, end subscript, C, l, left parenthesis, s, right parenthesis:

$\text{NH}_3(g)+\blueD{\text{H}}\text{Cl}(g)\rightarrow\text{N}\blueD{\text{H}}_4\text{Cl}(s)$N, H, start subscript, 3, end subscript, left parenthesis, g, right parenthesis, plus, start color blueD, H, end color blueD, C, l, left parenthesis, g, right parenthesis, right arrow, N, start color blueD, H, end color blueD, start subscript, 4, end subscript, C, l, left parenthesis, s, right parenthesis

This reaction can also be represented using the Lewis structures of the reactants and products, as seen below:
       NH₃ +HCl -----> NH₄⁺   Cl⁻

Lewis structure of ammonia—a nitrogen with a lone pair of electrons that is also bound to 3 hydrogens—plus the Lewis structure of hydrochloric acid forms ammonium chloride.nIn this reaction, 

$\blueD{\text{H}}\text{Cl}$HClstart color blueD, H, end color blueD, C, ldonates its proton—shown in blue—to 
$\text{NH}_3$N, H, start subscript, 3, end subscript. Therefore, $\text{HCl}$H, C, lis acting as a Brønsted-Lowry acid. Since $\text{NH}_3$N, H, start subscript, 3, end subscript has a lone pair which it uses to accept a proton, $\text{NH}_3$N, H, start subscript, 3, end subscript is a Brønsted-Lowry base.

Note that according to the Arrhenius theory, the above reaction would not be an acid-base reaction because neither species is forming $\text{H}^+$H, start superscript, plus, end superscript or $\text{OH}^-$O, H, start superscript, minus, end superscript in water. However, the chemistry involved$-$minusa proton transfer from $\text{HCl}$H, C, l to $\text{NH}_3$N, H, start subscript, 3, end subscript to form $\text{NH}_4 \text{Cl}$N, H, start subscript, 4, end subscript, C, l$-$minusis very similar to what would occur in the aqueous phase.

Identifying Brønsted-Lowry acids and bases

In the reaction between nitric acid and water, nitric acid, $\text{HNO}_3$H, N, O, start subscript, 3, end subscript, donates a proton—shown in blue—to water, thereby acting as a Brønsted-Lowry acid.

$\blueD{\text{H}}\text{NO}_3(aq)+\text{H}_2\text{O}(l)\rightarrow\blueD{\text{H}}_3\text{O}^+(aq)+\text{NO}_3^-(aq)$start color blueD, H, end color blueD, N, O, start subscript, 3, end subscript, left parenthesis, a, q, right parenthesis, plus, H, start subscript, 2, end subscript, O, left parenthesis, l, right parenthesis, right arrow, start color blueD, H, end color blueD, start subscript, 3, end subscript, O, start superscript, plus, end superscript, left parenthesis, a, q, right parenthesis, plus, N, O, start subscript, 3, end subscript, start superscript, minus, end superscript, left parenthesis, a, q, right parenthesis

Since water accepts the proton from nitric acid to form $\blueD{\text{H}}_3\text{O}^+$start color blueD, H, end color blueD, start subscript, 3, end subscript, O, start superscript, plus, end superscript, water acts as a Brønsted-Lowry base. This reaction highly favors the formation of products, so the reaction arrow is drawn only to the right.

Let's now look at a reaction involving ammonia, $\text{NH}_3$N, H, start subscript, 3, end subscript, in water:

$\text{NH}_3(aq)+\blueD{\text{H}}_2\text{O}(l)\rightleftharpoons\text{N}\blueD{\text{H}}_4^+(aq)+\text{OH}^-(aq)$

In this reaction, water is donating one of its protons to ammonia. After losing a proton, water becomes hydroxide, $\text{OH}^-$O, H, start superscript, minus, end superscript. Since water is a proton donor in this reaction, it is acting as a Brønsted-Lowry acid. Ammonia accepts a proton from water to form an ammonium ion, $\text{NH}_4^+$N, H, start subscript, 4, end subscript, start superscript, plus, end superscript. Therefore, ammonia is acting as a Brønsted-Lowry base.

Strong and weak acids: to dissociate, or not to dissociate?

A strong acid is a species that dissociates completely into its constituent ions in aqueous solution. Nitric acid is an example of a strong acid. It dissociates completely in water to form hydronium, $\text{H}_3\text{O}^+$H, start subscript, 3, end subscript, O, start superscript, plus, end superscript, and nitrate, $\text{NO}_3^-$N, O, start subscript, 3, end subscript, start superscript, minus, end superscript, ions. After the reaction occurs, there are no undissociated $\text{HNO}_3$H, N, O, start subscript, 3, end subscript molecules in solution.

By contrast, a weak acid does not dissociate completely into its constituent ions. An example of a weak acid is acetic acid, $\text{CH}_3\text{COOH}$C, H, start subscript, 3, end subscript, C, O, O, H, which is present in vinegar. Acetic acid dissociates partially in water to form hydronium and acetate ions, $\text{CH}_3\text{COO}^-$C, H, start subscript, 3, end subscript, C, O, O, start superscript, minus, end superscript:

$\text{CH}_3\text{COOH}(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_3\text{O}^+(aq)+\text{CH}_3\text{COO}^-(aq)$

Common strong acids

Name	Formula
Hydrochloric acid	$\text{HCl}$H, C, l
Hydrobromic acid	$\text{HBr}$H, B, r
Hydroiodic acid	$\text{HI}$H, I
Sulfuric acid	$\text{H}_2\text{SO}_4$H, start subscript, 2, end subscript, S, O, start subscript, 4, end subscript
Nitric acid	$\text{HNO}_3$H, N, O, start subscript, 3, end subscript
Perchloric acid	$\text{HClO}_4$H, C, l, O, start subscript, 4, end subscript

Strong and weak bases

A strong base is a base that ionizes completely in aqueous solution. An example of a strong base is sodium hydroxide, $\text{NaOH}$N, a, O, H. In water, sodium hydroxide dissociates completely to give sodium ions and hydroxide ions:

$\text{NaOH}(aq)\rightarrow\text{Na}^+(aq)+\text{OH}^-(aq)$N, a, O, H, left parenthesis, a, q, right parenthesis, right arrow, N, a, start superscript, plus, end superscript, left parenthesis, a, q, right parenthesis, plus, O, H, start superscript, minus, end superscript, left parenthesis, a, q, right parenthesis

Thus, if we make a solution of sodium hydroxide in water, only $\text{Na}^+$N, a, start superscript, plus, end superscript and $\text{OH}^-$O, H, start superscript, minus, end superscriptions are present in our final solution. We don't expect any undissociated $\text{NaOH}$N, a, O, H.

Let's now look at ammonia, $\text{NH}_3$N, H, start subscript, 3, end subscript, in water. Ammonia is a weak base, so it will become partially ionized in water:

$\text{NH}_3(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{NH}_4^+(aq)+\text{OH}^-(aq)$

Some of the ammonia molecules accept a proton from water to form ammonium ions and hydroxide ions. A dynamic equilibrium results, in which ammonia molecules are continually exchanging protons with water, and ammonium ions are continually donating the protons back to hydroxide. The major species in solution is non-ionized ammonia, $\text{NH}_3$N, H, start subscript, 3, end subscript, because ammonia will only deprotonate water to a small extent.

Common strong bases include Group 1 and Group 2 hydroxides.

Common weak bases include neutral nitrogen-containing compounds such as ammonia, trimethylamine, and pyridine.

Conjugate acid-base pairs

Now that we have an understanding of Brønsted-Lowry acids and bases, we can discuss the final concept covered in this article: conjugate acid-base pairs. In a Brønsted-Lowry acid-base reaction, a conjugate acid is the species formed after the base accepts a proton. By contrast, a conjugate base is the species formed after an acid donates its proton. The two species in a conjugate acid-base pair have the same molecular formula except the acid has an extra $\text H^+$H, start superscript, plus, end superscriptcompared to the conjugate base.

Summary

• A Brønsted-Lowry acid is any species that is capable of donating a proton—$\text{H}^+$H, start superscript, plus, end superscript.
• A Brønsted-Lowry base is any species that is capable of accepting a proton, which requires a lone pair of electrons to bond to the $\text{H}^+$H, start superscript, plus, end superscript.
• Water is amphoteric, which means it can act as both a Brønsted-Lowry acid and a Brønsted-Lowry base.
• Strong acids and bases ionize completely in aqueous solution, while weak acids and bases ionize only partially in aqueous solution.
• The conjugate base of a Brønsted-Lowry acid is the species formed after an acid donates its proton. The conjugate acid of a Brønsted-Lowry base is the species formed after a base accepts a proton.
• The two species in a conjugate acid-base pair have the same molecular formula except the acid has an extra $\text H^+$H, start superscript, plus, end superscript compared to the conjugate base.

$\text{Ca(OH)}_2$C, a, left parenthesis, O, H, right parenthesis, start subscript, 2, end subscript$\text{OH}^-$O, H, start superscript, minus, end superscriptCommon weak bases include neutral nitrogen-containing compounds such as ammonia, trimethylamine, and pyridine.