# Brønsted-Lowry acid base theory

Definition of Brønsted-Lowry acids and bases, strong and weak acids and bases, and how to identify conjugate acid-base pairs.

## Key points

• A Brønsted-Lowry acid is any species that is capable of donating a proton—$\text{H}^+$.
• 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 $\text{H}^+$.
• 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 $\text H^+$ compared to the conjugate base.

## Introduction

A fish market where a variety of fresh and packaged fish are displayed on ice.
Seafood contains compounds that can break down to form amines, which are weak bases with a characteristic "fishy" odor. Image credit: from pixabay, CC0 public domain
In a previous article on Arrhenius acids and bases, we learned that an Arrhenius acid is any species that can increase the concentration of $\text{H}^+$ in aqueous solution and an Arrhenius base is any species that can increase the concentration of $\text{OH}^-$ in aqueous solution. A major limitation of Arrhenius theory is that we can only describe acid-base behavior in water. In this article, we'll move on to look at the more general Brønsted-Lowry theory, which applies to a broader range of chemical reactions.

## 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, $\text{H}^+$, 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 $\text H^+$. 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, $\text{NH}_3(g)$, with hydrogen chloride gas, $\text{H}\text{Cl}(g)$, to form solid ammonium chloride, $\text{NH}_4 \text{Cl}(s)$:
$\text{NH}_3(g)+\blueD{\text{H}}\text{Cl}(g)\rightarrow\text{N}\blueD{\text{H}}_4\text{Cl}(s)$
This reaction can also be represented using the Lewis structures of the reactants and products, as seen below:
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.
In this reaction, $\blueD{\text{H}}\text{Cl}$ donates its proton—shown in blue—to $\text{NH}_3$. Therefore, $\text{HCl}$ is acting as a Brønsted-Lowry acid. Since $\text{NH}_3$ has a lone pair which it uses to accept a proton, $\text{NH}_3$ 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 $\text{H}^+$ or $\text{OH}^-$ in water. However, the chemistry involved$-$a proton transfer from $\text{HCl}$ to $\text{NH}_3$ to form $\text{NH}_4 \text{Cl}$$-$is very similar to what would occur in the aqueous phase.
To get more familiar with these definitions, let's examine some more examples.

## Identifying Brønsted-Lowry acids and bases

In the reaction between nitric acid and water, nitric acid, $\text{HNO}_3$, donates a proton—shown in blue—to water, thereby acting as a Brønsted-Lowry acid.
$\blueD{\text{H}}\text{NO}_3(aq)+\text{H}_2\text{O}(l)\rightarrow\blueD{\text{H}}_3\text{O}^+(aq)+\text{NO}_3^-(aq)$
Since water accepts the proton from nitric acid to form $\blueD{\text{H}}_3\text{O}^+$, 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, $\text{NH}_3$, 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, $\text{OH}^-$. 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, $\text{NH}_4^+$. Therefore, ammonia is acting as a Brønsted-Lowry base.
In the two previous reactions, we see water behaving both as a Brønsted-Lowry base—in the reaction with nitric acid—and as a Brønsted-Lowry acid—in the reaction with ammonia. Because of its ability to both accept and donate protons, water is known as an amphoteric or amphiprotic substance, meaning that it can act as either a Brønsted-Lowry acid or a Brønsted-Lowry base.
The word amphoteric is derived from a Greek word meaning both. It might sound like a really weird, random word that is hard to say, but you actually see the prefix in many other words! Examples of related words include amphibian, which means "both lives", and amphitheatre, which means "a place for viewing on both sides".

## 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, $\text{H}_3\text{O}^+$, and nitrate, $\text{NO}_3^-$, ions. After the reaction occurs, there are no undissociated $\text{HNO}_3$ 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, $\text{CH}_3\text{COOH}$, which is present in vinegar. Acetic acid dissociates partially in water to form hydronium and acetate ions, $\text{CH}_3\text{COO}^-$:
$\text{CH}_3\text{COOH}(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_3\text{O}^+(aq)+\text{CH}_3\text{COO}^-(aq)$
Notice that in this reaction, we have arrows pointing in both directions: $\leftrightharpoons$. This indicates that dissociation of acetic acid is a dynamic equilibrium where there will be a significant concentration of acetic acid molecules that are present as neutral $\text{CH}_3\text{COOH}$ molecules as well as in the form of the dissociated ions, $\text H^+$ and $\text{CH}_3\text{COO}^-$.
On left: zoomed-in representation of hydrochloric acid solution, where the acid is fully dissociated as protons and chloride ions. On right: zoomed-in representation of hydrofluoric acid solution showing most of the hydrofluoric acid is still in the neutral molecule form, HF, while a few are dissociated as protons and fluoride ions.
Aqueous solutions of a strong acid, left, and a weak acid, right. (a) Hydrochloric acid is a strong acid that fully dissociates in water. (b) Hydrofluoric acid is a weak acid that partially dissociates into protons and fluoride ions.
A common question is, “When do you know when something is a strong or a weak acid?” That is an excellent question! The short answer is that there are only a handful of strong acids, and everything else is considered a weak acid. Once we are familiar with the common strong acids, we can easily identify both weak and strong acids in chemistry problems.
The following table lists some examples of common strong acids.

### Common strong acids

NameFormula
Hydrochloric acid$\text{HCl}$
Hydrobromic acid$\text{HBr}$
Hydroiodic acid$\text{HI}$
Sulfuric acid$\text{H}_2\text{SO}_4$
Nitric acid$\text{HNO}_3$
Perchloric acid$\text{HClO}_4$

## 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, $\text{NaOH}$. In water, sodium hydroxide dissociates completely to give sodium ions and hydroxide ions:
$\text{NaOH}(aq)\rightarrow\text{Na}^+(aq)+\text{OH}^-(aq)$
Thus, if we make a solution of sodium hydroxide in water, only $\text{Na}^+$ and $\text{OH}^-$ ions are present in our final solution. We don't expect any undissociated $\text{NaOH}$.
Let's now look at ammonia, $\text{NH}_3$, 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, $\text{NH}_3$, because ammonia will only deprotonate water to a small extent.
Common strong bases include Group 1 and Group 2 hydroxides.
Even insoluble compounds such as $\text{Ca(OH)}_2$ are often classified as strong bases because they have some tiny fraction that can dissolve in water, and the fraction that is in solution fully dissociates to form $\text{OH}^-$ ions.
Another way to think about this is that complete solubility is not a requirement for something to be a strong base! As long as some of it gets into solution and fully dissociates, we can classify it as a strong base.
Common weak bases include neutral nitrogen-containing compounds such as ammonia, trimethylamine, and pyridine.

## Example 1: Writing an acid-base reaction with hydrogen phosphate

Hydrogen phosphate, $\text{HPO}_4^{2-}$, can act as a weak base or as a weak acid in aqueous solution.
What is the balanced equation for the reaction of hydrogen phosphate acting as a weak base in water?
Since hydrogen phosphate is acting as a Brønsted-Lowry base, water must be acting as a Brønsted-Lowry acid. This means that water will donate a proton to generate hydroxide. The addition of a proton to hydrogen phosphate results in the formation of $\text{H}_2 \text {PO}_4^{-}$:
$\text{HPO}_4^{2-}(aq)+\text H^+(aq) \rightarrow \text{H}_2\text {PO}_4^{-}(aq)$
Since hydrogen phosphate is acting as a weak base in this particular example, we will need to use equilibrium arrows, $\rightleftharpoons$, in our overall reaction to show that the reaction is reversible. That gives the following balanced equation for the reaction of hydrogen phosphate acting as a weak base in water:
$\text{HPO}_4^{2-}(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_2 \text{PO}_4^{-}(aq)+\text{OH}^-(aq)$
How do we know when something like hydrogen phosphate will act like an acid or a base? The short answer is that when different reactions are possible, the different equilibrium reactions have different equilibrium constants as well. Which equilibrium will be favored depends on factors such as the pH of the solution and what other species are in solution. This question will be addressed in more detail when we learn about buffers and titrations!
Concept check: What would our balanced equation look like if hydrogen phosphate acted as a weak acid in aqueous solution?
If hydrogen phosphate is acting as a Brønsted-Lowry acid, then water must be acting as a Brønsted-Lowry base. This means that water will accept a proton to generate hydronium. The loss of a proton from hydrogen phosphate results in the formation of phosphate, $\text {PO}_4^{3-}$:
$\text{HPO}_4^{2-}(aq) \rightarrow \text {PO}_4^{3-}(aq)+\text H^+(aq)$
Since hydrogen phosphate is acting as a weak acid in this particular example, we will need to use equilibrium arrows, $\rightleftharpoons$, in our overall reaction to show that the reaction is reversible. That gives the following balanced equation:
$\text{HPO}_4^{2-}(aq)+\text{H}_2\text{O}(l)\rightleftharpoons \text{PO}_4^{3-}(aq)+\text{H}_3 \text O^+(aq)$

## 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 $\text H^+$ compared to the conjugate base.

### Example 2: Dissociation of a strong acid

Let's reconsider the strong acid $\text{HCl}$ reacting with water:
$\text{HCl}(aq)+\text{H}_2\text{O}(l)\rightarrow \text{H}_3\text{O}^+(aq)+\text{Cl}^-(aq)$
$~~~~~~~~~~\greenD{\text{acid}}~~~~~~~~~~~~\purpleC{\text{base}}~~~~~~~~~~~~~~\purpleC{\text{acid}}~~~~~~~~~~~\greenD{\text{base}}$
In this reaction, $\text{HCl}$ donates a proton to water; therefore, $\text{HCl}$ is acting as a Brønsted-Lowry acid. After $\text{HCl}$ donates its proton, the $\text{Cl}^-$ ion is formed; thus, $\text{Cl}^-$ is the conjugate base of $\text{HCl}$.
$\greenD{\text{Conjugate pair 1}}=\text{HCl}\text{ and }\text{Cl}^-$
Because water accepts a proton from $\text{HCl}$, water is acting as a Brønsted-Lowry base. When water accepts a proton, $\text{H}_3\text{O}^+$ is formed. Therefore, $\text{H}_3\text{O}^+$ is the conjugate acid of $\text{H}_2\text{O}$.
$\purpleC{\text{Conjugate pair 2}}=\text{H}_2 \text O\text{ and }\text{H}_3\text{O}^+$
Each conjugate acid-base pair in our reaction contains one Brønsted-Lowry acid and one Brønsted-Lowry base; the acid and base differ by a single proton. It will generally be true that a reaction between a Brønsted-Lowry acid and base will contain two conjugate acid-base pairs.

### Example 3: Ionization of a weak base

Let's consider the reaction of the weak base ammonia in water:
$\text{NH}_3(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{NH}_4^+(aq)+\text{OH}^-(aq)$
$~~~~~~~~~~\greenD{\text{base}}~~~~~~~~~~~~\purpleC{\text{acid}}~~~~~~~~~~~~\greenD{\text{acid}}~~~~~~~~~~~~~\purpleC{\text{base}}$
Ammonia accepts a proton from water in this reaction, and thereby acts as a Brønsted-Lowry base. Upon accepting a proton from water, ammonia forms $\text{NH}_4^+$. Therefore, $\text{NH}_4^+$ is the conjugate acid of ammonia.
$\greenD{\text{Conjugate pair 1}}=\text{NH}_3\text{ and }\text{NH}_4^+$
Water, by donating a proton to ammonia, acts as a Brønsted-Lowry acid. After water donates its proton to ammonia, $\text{OH}^-$ is formed. Therefore, $\text{OH}^-$ is the conjugate base of water.
$\purpleC{\text{Conjugate pair 2}}=\text{H}_2 \text O\text{ and }\text{OH}^-$
Since ammonia is a weak base, the ammonium ion can donate a proton back to hydroxide to reform ammonia and water. Thus, a dynamic equilibrium exists. This will always be true for reactions involving weak acids and bases.

## Summary

• A Brønsted-Lowry acid is any species that is capable of donating a proton—$\text{H}^+$.
• 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 $\text{H}^+$.
• 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 $\text H^+$ compared to the conjugate base.

1. "Acid/Base Basics" from UC Davis ChemWiki, CC BY-NC-SA 3.0
2. Conjugate Acid-Base Pairs” from UC Davis ChemWiki, CC BY-NC-SA 3.0
3. "The Arrhenius Definition" from Boundless Learning, CC BY-SA 4.0

Zumdahl, S.S., and S. A. Zumdahl S.A. "Atomic Structure and Periodicity." In Chemistry, 290-294. 6th ed. Boston, MA: Houghton Mifflin Company, 2003.
Kotz, J. C., P. m. Treichel, J. R. Townsend, and D. A. Treichel. "The Brønsted-Lowry Concept of Acids and Bases." In Chemistry and Chemical Reactivity, Instructor's Edition, 586-588. 9th ed. Stamford, CT: Cengage Learning, 2015.

## Practice 1: Identifying acid-base reactions

Based on Brønsted-Lowry theory, which of the following are acid-base reactions?
An acid-base reaction is a reaction between a proton donor—a Brønsted-Lowry acid—and a proton acceptor—a Brønsted-Lowry base.
In the reaction $\text{LiOH}(aq)+\text{HBr}(aq)\rightarrow\text{H}_2\text{O}(l)+\text{LiBr}(aq)$, $\text{HBr}$ is donating a proton to $\text{LiOH}$, thus acting as a Brønsted-Lowry acid. Since $\text{LiOH}$ accepts the proton, it is acting as a Brønsted-Lowry base.
In the reaction $2\text{ NH}_3\rightleftharpoons\text{NH}_4^++\text{NH}_2^-$, one molecule of $\text{NH}_3$ donates a proton to another molecule of $\text{NH}_3$. Therefore, one molecule of $\text{NH}_3$ is acting as a Brønsted-Lowry acid—proton donor—and another molecule of $\text{NH}_3$ is acting as a Brønsted-Lowry base—proton acceptor. This is an example of an autoionization reaction.
The acid-base reactions are below:
$\text{LiOH}(aq)+\text{HBr}(aq)\rightarrow\text{H}_2\text{O}(l)+\text{LiBr}(aq)$
$2\text{ NH}_3\rightleftharpoons\text{NH}_4^++\text{NH}_2^-$

## Practice 2: Identifying conjugate acid-base pairs

Hydrofluoric acid, $\text{HF}$, is a weak acid that dissociates in water according to the following equation:
$\text{HF}(aq)+\text{H}_2\text{O}(l)\rightleftharpoons\text{H}_3\text{O}^+(aq)+\text{F}^-(aq)$
What is the conjugate base of $\text{HF}$ in this reaction?
In this reaction, $\text{HF}$ acts as a Brønsted-Lowry acid by donating its proton to $\text{H}_2\text{O}$. After donating a proton, the anion $\text{F}^-$ is formed.
The fluoride anion, $\text{F}^-$, is the conjugate base of $\text{HF}$.