How is a reaction such as LiOH(aq)+HBr(aq)→H2 O(l)+LiBr(aq) considered an acid-base reactions by this new definition. LiOH does not accept a proton and become LiH2O.

I see what you are saying, but it is the OH- of the LiOH that is accepting the proton to give H2O, so the OH- is acting as a Brønsted-Lowry base. LiOH will completely dissociate in water to give Li+ and OH-. So LiOH is the source of the OH- ions that accept the protons. LiOH is regarded as being a base despite the Li+ ions not being involved in proton transfer.

Can someone explain why LiOH + HBr --> H20 + LiBr is an acid base reaction (practice problem 1)?

The HBr is an acid, because it donates a proton to the OH⁻ of the LiOH. The OH⁻ of the LiOH is a base because it accepts the proton from the HBr.

I have a very basic question. Why did chemists select the concentration of H+ ions to calculate acidity?

Acidity is just the name we give to the presence of extra H+ ions.

At what point would you balance the equation you write to show the reaction of acids or bases in water? Also, when do you use equilibrium arrows, for an acid-base reaction or for an acid or a base reacting with water?

You must balance an equation before you do any calculations that will need the coefficients in the balanced equation (i.e. stoichiometry or equilibrium calculations). You use a reaction arrow (→) for any acid-base neutralization in which you have at last one strong acid or strong base. NaOH + HCl → NaCl + H₂O NaOH + CH₃COOH → CH₃COONa + H₂O HCl + NH₃ → NH₄Cl You use equilibrium arrows (⇌) if none of the reactants is a strong electrolyte: CH₃COOH + NH₃ ⇌ CH₃ COONH₄ CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺ NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

How do we know which reactions are one way and which are reversible? Is there some fundamental concept that I should study?

Where the _activation energy_ is low, the reaction is reversible. This typically occurs in reaction where the equilibrium (equilibrium constant) is affected by factors so that it can change from >1 (forward direction) to <1 (backward direction). This is mostly the case where concentration of solutions or pressure in gases has an effect.

Does ammonia act both as an acid and a base in aqueous solution? because it is given that it repeatedly accepts and donates protons

Ammonia is a weak base in water. It behaves as a Bronsted-Lowry base because its lobe pair accepts a H+ ion. Its conjugate acid NH4+ behaves as an acid a Bronsted-Lowry acid.....

Why aren't strong acid and base reactions also reversible?

Because they fully dissociate, or fully react. Their Equilibrium constant is so high, and the equilibrium is shifted so far right that all of a strong acid dissociates and does't want to be put back together.

In the "Example 3: Ionization of a weak base", why NH3 acts like a base and not like an acid?

I was wondering as well, but possibly that ammonia likes to form ammonium ion by accepting a proton, and that's why it acts like a base.

The word "species" in this article means "compound", right?

Chemical species are atoms, molecules, molecular fragments, ions, etc., being subjected to a chemical process or to a measurement. Generally, a chemical species can be defined as an ensemble of chemically identical molecular entities that can explore the same set of molecular energy levels on a defined time scale.

Main content

AP®︎/College Chemistry

Course: AP®︎/College Chemistry > Unit 12

Lesson 1: Acids, bases, and pH

Brønsted-Lowry acids and bases

Q: Why aren't strong acid and base reactions also reversible?

Because they fully dissociate, or fully react. Their Equilibrium constant is so high, and the equilibrium is shifted so far right that all of a strong acid dissociates and does't want to be put back together.

Q: In the "Example 3: Ionization of a weak base", why NH3 acts like a base and not like an acid?

I was wondering as well, but possibly that ammonia likes to form ammonium ion by accepting a proton, and that's why it acts like a base.

Q: The word "species" in this article means "compound", right?

Chemical species are atoms, molecules, molecular fragments, ions, etc., being subjected to a chemical process or to a measurement. Generally, a chemical species can be defined as an ensemble of chemically identical molecular entities that can explore the same set of molecular energy levels on a defined time scale.

Google Classroom

Key points

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

Introduction

In a previous article on Arrhenius acids and bases, we learned that an Arrhenius acid is any species that can increase the concentration of

H^{+}

in aqueous solution and an Arrhenius base is any species that can increase the concentration of

{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,

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

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,

{NH}_{3} (g)

, with hydrogen chloride gas,

H Cl (g)

, to form solid ammonium chloride,

{NH}_{4} Cl (s)

${NH}_{3} (g) + H Cl (g) \to N H_{4} Cl (s)$ ‍

This reaction can also be represented using the Lewis structures of the reactants and products, as seen below:

In this reaction,

H Cl

donates its proton—shown in blue—to

{NH}_{3}

. Therefore,

HCl

is acting as a Brønsted-Lowry acid. Since

{NH}_{3}

has a lone pair which it uses to accept a proton,

{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

H^{+}

{OH}^{-}

in water. However, the chemistry involved

-

a proton transfer from

HCl

{NH}_{3}

to form

{NH}_{4} 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,

{HNO}_{3}

, donates a proton—shown in blue—to water, thereby acting as a Brønsted-Lowry acid.

$H {NO}_{3} (a q) + H_{2} O (l) \to H_{3} O^{+} (a q) + {NO}_{3}^{-} (a q)$ ‍

Since water accepts the proton from nitric acid to form

H_{3} 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,

{NH}_{3}

, in water:

${NH}_{3} (a q) + H_{2} O (l) ⇌ N H_{4}^{+} (a q) + {OH}^{-} (a q)$ ‍

In this reaction, water is donating one of its protons to ammonia. After losing a proton, water becomes hydroxide,

{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,

{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.

[What does the prefix amphi mean?]

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,

H_{3} O^{+}

, and nitrate,

{NO}_{3}^{-}

, ions. After the reaction occurs, there are no undissociated

{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,

{CH}_{3} COOH

, which is present in vinegar. Acetic acid dissociates partially in water to form hydronium and acetate ions,

{CH}_{3} {COO}^{-}

${CH}_{3} COOH (a q) + H_{2} O (l) ⇌ H_{3} O^{+} (a q) + {CH}_{3} {COO}^{-} (a q)$ ‍

Notice that in this reaction, we have arrows pointing in both directions:

⇋

. 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

{CH}_{3} COOH

molecules as well as in the form of the dissociated ions,

H^{+}

and

{CH}_{3} {COO}^{-}

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

Name	Formula
Hydrochloric acid	$HCl$ ‍
Hydrobromic acid	$HBr$ ‍
Hydroiodic acid	$HI$ ‍
Sulfuric acid	$H_{2} {SO}_{4}$ ‍
Nitric acid	${HNO}_{3}$ ‍
Perchloric acid	${HClO}_{4}$ ‍

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,

NaOH

. In water, sodium hydroxide dissociates completely to give sodium ions and hydroxide ions:

$NaOH (a q) \to {Na}^{+} (a q) + {OH}^{-} (a q)$ ‍

Thus, if we make a solution of sodium hydroxide in water, only

{Na}^{+}

and

{OH}^{-}

ions are present in our final solution. We don't expect any undissociated

NaOH

Let's now look at ammonia,

{NH}_{3}

, in water. Ammonia is a weak base, so it will become partially ionized in water:

${NH}_{3} (a q) + H_{2} O (l) ⇌ {NH}_{4}^{+} (a q) + {OH}^{-} (a q)$ ‍

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,

{NH}_{3}

, because ammonia will only deprotonate water to a small extent.

Common strong bases include Group 1 and Group 2 hydroxides.

[How can an insoluble compound be 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,

{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

H_{2} {PO}_{4}^{-}

${HPO}_{4}^{2 -} (a q) + H^{+} (a q) \to H_{2} {PO}_{4}^{-} (a q)$ ‍

Since hydrogen phosphate is acting as a weak base in this particular example, we will need to use equilibrium arrows,

⇌

, 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:

${HPO}_{4}^{2 -} (a q) + H_{2} O (l) ⇌ H_{2} {PO}_{4}^{-} (a q) + {OH}^{-} (a q)$ ‍

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?

[Show the answer.]

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

H^{+}

compared to the conjugate base.

Example 2: Dissociation of a strong acid

Let's reconsider the strong acid

HCl

reacting with water:

$HCl (a q) + H_{2} O (l) \to H_{3} O^{+} (a q) + {Cl}^{-} (a q)$ ‍

acid base acid base

In this reaction,

HCl

donates a proton to water; therefore,

HCl

is acting as a Brønsted-Lowry acid. After

HCl

donates its proton, the

{Cl}^{-}

ion is formed; thus,

{Cl}^{-}

is the conjugate base of

HCl

$Conjugate pair 1 = HCl and {Cl}^{-}$ ‍

Because water accepts a proton from

HCl

, water is acting as a Brønsted-Lowry base. When water accepts a proton,

H_{3} O^{+}

is formed. Therefore,

H_{3} O^{+}

is the conjugate acid of

H_{2} O

$Conjugate pair 2 = H_{2} O and H_{3} O^{+}$ ‍

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:

${NH}_{3} (a q) + H_{2} O (l) ⇌ {NH}_{4}^{+} (a q) + {OH}^{-} (a q)$ ‍

base acid acid 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

{NH}_{4}^{+}

. Therefore,

{NH}_{4}^{+}

is the conjugate acid of ammonia.

$Conjugate pair 1 = {NH}_{3} and {NH}_{4}^{+}$ ‍

Water, by donating a proton to ammonia, acts as a Brønsted-Lowry acid. After water donates its proton to ammonia,

{OH}^{-}

is formed. Therefore,

{OH}^{-}

is the conjugate base of water.

$Conjugate pair 2 = H_{2} O and {OH}^{-}$ ‍

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— $H^{+}$ ‍.
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 $H^{+}$ ‍.
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 $H^{+}$ ‍ compared to the conjugate base.

[Attributions and references]

Practice 1: Identifying acid-base reactions

Based on Brønsted-Lowry theory, which of the following are acid-base reactions?

[Hint 1]

[Hint 2]

[Hint 3]

Practice 2: Identifying conjugate acid-base pairs

Hydrofluoric acid,

HF

, is a weak acid that dissociates in water according to the following equation:

$HF (a q) + H_{2} O (l) ⇌ H_{3} O^{+} (a q) + F^{-} (a q)$ ‍

What is the conjugate base of $HF$ ‍ in this reaction?

[Hint 1]

[Hint 2]

Want to join the conversation?

Sort by:

Mujtaba Hussain
Posted 7 years ago. Direct link to Mujtaba Hussain's post “How is a reaction such as...”
How is a reaction such as LiOH(aq)+HBr(aq)→H2
O(l)+LiBr(aq) considered an acid-base reactions by this new definition. LiOH does not accept a proton and become LiH2O.
Button navigates to signup pageButton navigates to signup page
(15 votes)
Answer
- RogerP
  Posted 7 years ago. Direct link to RogerP's post “I see what you are saying...”
  I see what you are saying, but it is the OH- of the LiOH that is accepting the proton to give H2O, so the OH- is acting as a Brønsted-Lowry base. LiOH will completely dissociate in water to give Li+ and OH-. So LiOH is the source of the OH- ions that accept the protons. LiOH is regarded as being a base despite the Li+ ions not being involved in proton transfer.
  Button navigates to signup page
  (23 votes)
Cindy
Posted 8 years ago. Direct link to Cindy's post “Can someone explain why L...”
Can someone explain why LiOH + HBr --> H20 + LiBr is an acid base reaction (practice problem 1)?
Button navigates to signup pageButton navigates to signup page
(12 votes)
Answer
- Ernest Zinck
  Posted 8 years ago. Direct link to Ernest Zinck's post “The HBr is an acid, becau...”
  The HBr is an acid, because it donates a proton to the OH⁻ of the LiOH.
  The OH⁻ of the LiOH is a base because it accepts the proton from the HBr.
  Comment on Ernest Zinck's post “The HBr is an acid, becau...”
  (8 votes)
Amit Mukherjee
Posted 8 years ago. Direct link to Amit Mukherjee's post “I have a very basic quest...”
I have a very basic question. Why did chemists select the concentration of H+ ions to calculate acidity?
Button navigates to signup pageButton navigates to signup page
(8 votes)
Answer
- Andrew M
  Posted 8 years ago. Direct link to Andrew M's post “Acidity is just the name ...”
  Acidity is just the name we give to the presence of extra H+ ions.
  Button navigates to signup page
  (9 votes)
twentyonellamas
Posted 8 years ago. Direct link to twentyonellamas's post “At what point would you b...”
At what point would you balance the equation you write to show the reaction of acids or bases in water?
Also, when do you use equilibrium arrows, for an acid-base reaction or for an acid or a base reacting with water?
Button navigates to signup pageButton navigates to signup page
(4 votes)
Answer
- Ernest Zinck
  Posted 8 years ago. Direct link to Ernest Zinck's post “You must balance an equat...”
  You must balance an equation before you do any calculations that will need the coefficients in the balanced equation (i.e. stoichiometry or equilibrium calculations).
  You use a reaction arrow (→) for any acid-base neutralization in which you have at last one strong acid or strong base.
  NaOH + HCl → NaCl + H₂O
  NaOH + CH₃COOH → CH₃COONa + H₂O
  HCl + NH₃ → NH₄Cl
  You use equilibrium arrows (⇌) if none of the reactants is a strong electrolyte:
  CH₃COOH + NH₃ ⇌ CH₃ COONH₄
  CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺
  NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
  Comment on Ernest Zinck's post “You must balance an equat...”
  (7 votes)
Amit Mukherjee
Posted 8 years ago. Direct link to Amit Mukherjee's post “How do we know which reac...”
How do we know which reactions are one way and which are reversible? Is there some fundamental concept that I should study?
Button navigates to signup pageButton navigates to signup page
(4 votes)
Answer
- Matt B
  Posted 8 years ago. Direct link to Matt B's post “Where the _activation ene...”
  Where the activation energy is low, the reaction is reversible. This typically occurs in reaction where the equilibrium (equilibrium constant) is affected by factors so that it can change from >1 (forward direction) to <1 (backward direction). This is mostly the case where concentration of solutions or pressure in gases has an effect.
  Comment on Matt B's post “Where the _activation ene...”
  (4 votes)
sanzoo123456
Posted 8 years ago. Direct link to sanzoo123456's post “Does ammonia act both as ...”
Does ammonia act both as an acid and a base in aqueous solution? because it is given that it repeatedly accepts and donates protons
Button navigates to signup pageButton navigates to signup page
(3 votes)
Answer
- Davide Ghazal
  Posted 8 years ago. Direct link to Davide Ghazal's post “Ammonia is a weak base in...”
  Ammonia is a weak base in water. It behaves as a Bronsted-Lowry base because its lobe pair accepts a H+ ion. Its conjugate acid NH4+ behaves as an acid a Bronsted-Lowry acid.....
  Button navigates to signup page
  (3 votes)
Kyler
Posted 7 years ago. Direct link to Kyler's post “Why aren't strong acid an...”
Why aren't strong acid and base reactions also reversible?
Button navigates to signup pageButton navigates to signup page
(3 votes)
Answer
- Zenu Destroyer of Worlds (AK)
  Posted 7 years ago. Direct link to Zenu Destroyer of Worlds (AK)'s post “Because they fully dissoc...”
  Because they fully dissociate, or fully react. Their Equilibrium constant is so high, and the equilibrium is shifted so far right that all of a strong acid dissociates and does't want to be put back together.
  Comment on Zenu Destroyer of Worlds (AK)'s post “Because they fully dissoc...”
  (4 votes)
Jorge Daniel Paez Vivas
Posted 8 years ago. Direct link to Jorge Daniel Paez Vivas's post “In the "Example 3: Ioniza...”
In the "Example 3: Ionization of a weak base", why NH3 acts like a base and not like an acid?
Button navigates to signup pageComment on Jorge Daniel Paez Vivas's post “In the "Example 3: Ioniza...”
(3 votes)
Answer
- Gavin Peng
  Posted 8 years ago. Direct link to Gavin Peng's post “I was wondering as well, ...”
  I was wondering as well, but possibly that ammonia likes to form ammonium ion by accepting a proton, and that's why it acts like a base.
  Comment on Gavin Peng's post “I was wondering as well, ...”
  (2 votes)
Jonathan Ziesmer
Posted 8 years ago. Direct link to Jonathan Ziesmer's post “The word "species" in thi...”
The word "species" in this article means "compound", right?
Button navigates to signup pageButton navigates to signup page
(2 votes)
Answer
- saadia peerzada
  Posted 8 years ago. Direct link to saadia peerzada's post “Chemical species are atom...”
  Chemical species are atoms, molecules, molecular fragments, ions, etc., being subjected to a chemical process or to a measurement. Generally, a chemical species can be defined as an ensemble of chemically identical molecular entities that can explore the same set of molecular energy levels on a defined time scale.
  Button navigates to signup page
  (2 votes)
Shruthi K
Posted 8 years ago. Direct link to Shruthi K's post “Why do there have to be 2...”
Why do there have to be 2 electrons for the hydrogen ions to bond?
Button navigates to signup pageButton navigates to signup page
(3 votes)
Answer
- Matt B
  Posted 8 years ago. Direct link to Matt B's post “Each covalent bond needs ...”
  Each covalent bond needs to have two electrons. (It is one of these rules from quatum chemistry.) Hydrogens will therefore each donate one electron to form this single covalent bond.
  Button navigates to signup page
  (1 vote)