Chapter 16, Problem 16.3QP

(a)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction:

Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (a) is both Bronsted acid and Bronsted base.

Explanation of Solution

To classify: ${\text{H}}_{\text{2}}\text{O}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

  ${\text{H}}_{\text{2}}\text{O}\to {\text{H}}^{\text{+}}\text{+}{\text{OH}}^{\text{-}}$

Water molecule loses a proton to form a conjugate base as shown above.  Therefore, water can act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{H}}_{\text{2}}\text{O}+{\text{H}}^{\text{+}}\to {\text{H}}_{\text{3}}{\text{O}}^{\text{+}}$

  Water molecule accepts a proton to form hydronium ion.  Therefore, water can act as Bronsted base.

From this we can conclude that water can act as both Bronsted acid and Bronsted base.

(b)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (b) is Bronsted base.

Explanation of Solution

To classify: ${\text{OH}}^{\text{-}}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

Hydroxide ion cannot lose a proton to form a conjugate base.  Therefore, hydroxide ion cannot act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{OH}}^{\text{-}}+{\text{H}}^{\text{+}}\to {\text{H}}_{\text{2}}\text{O}$

Hydroxide ion accepts a proton to form water molecule.  Therefore, hydroxide ion can act as Bronsted base.

From this we can conclude that hydroxide ion can only act as Bronsted base.

(c)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (c) is Bronsted acid.

Explanation of Solution

To classify: ${\text{H}}_{\text{3}}{\text{O}}^{\text{+}}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

  ${\text{H}}_{\text{3}}{\text{O}}^{\text{+}}\to {\text{H}}_{\text{2}}\text{O}\text{+}{\text{H}}^{\text{+}}$

The hydronium ion can lose a proton to form a conjugate base as shown above.  Therefore, hydronium ion can act as Bronsted acid.

To identify the species as Bronsted base.

Hydronium ion cannot accept proton to form a conjugate acid.

From this we can conclude that hydronium ion can act only as Bronsted acid.

(d)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (d) is Bronsted base.

Explanation of Solution

To classify: ${\text{NH}}_{\text{3}}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

Ammonia cannot lose a proton to form a conjugate base.  Therefore, ammonia cannot act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{NH}}_{\text{3}}+{\text{H}}^{\text{+}}\to {\text{NH}}_{\text{4}}{}^{+}$

Ammonia accepts a proton to form ammonium ion.  Therefore, ammonia ion can act as Bronsted base.

From this we can conclude that ammonia can act only as Bronsted base.

(e)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (e) is Bronsted acid.

Explanation of Solution

To classify: ${\text{NH}}_{\text{4}}{}^{+}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

  ${\text{NH}}_{\text{4}}{}^{\text{+}}\to {\text{NH}}_{\text{3}}\text{+}{\text{H}}^{\text{+}}$

The ammonium ion can lose a proton to form a conjugate base as shown above.  Therefore ammonium ion can act as Bronsted acid.

To identify the species as Bronsted base.

Ammonium ion cannot accept proton to form a conjugate acid.

From this we can conclude that ammonium ion can act only as Bronsted acid.

(f)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (f) is Bronsted base.

Explanation of Solution

To classify: ${\text{NH}}_{\text{2}}{}^{-}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

${\text{NH}}_{\text{2}}{}^{-}$ cannot lose a proton to form a conjugate base.  Therefore, ${\text{NH}}_{\text{2}}{}^{-}$ cannot act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{NH}}_{\text{2}}{}^{-}+{\text{H}}^{\text{+}}\to {\text{NH}}_{\text{3}}$

${\text{NH}}_{\text{2}}{}^{-}$ accepts a proton to form ammonia.  Therefore, ${\text{NH}}_{\text{2}}{}^{-}$ ion can act as Bronsted base.

From this we can conclude that ${\text{NH}}_{\text{2}}{}^{-}$ can act only as Bronsted base.

(g)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (g) is Bronsted base.

Explanation of Solution

To classify: ${\text{NO}}_{\text{3}}{}^{-}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

${\text{NO}}_{\text{3}}{}^{-}$ does not have a proton at all.  Therefore, this cannot lose a proton to form a conjugate base.  Hence, ${\text{NO}}_{\text{3}}{}^{-}$ cannot act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{NO}}_{\text{3}}{}^{-}+{\text{H}}^{\text{+}}\to {\text{HNO}}_{\text{3}}$

${\text{NO}}_{\text{3}}{}^{-}$ accepts a proton to form nitric acid.  Therefore, ${\text{NO}}_{\text{3}}{}^{-}$ ion can act as Bronsted base.

From this we can conclude that ${\text{NO}}_{\text{3}}{}^{-}$ can act only as Bronsted base.

(h)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (h) is Bronsted base.

Explanation of Solution

To classify: ${\text{CO}}_{\text{3}}{}^{2-}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

Explanation:  ${\text{CO}}_{\text{3}}{}^{2-}$ does not have a proton at all.  Therefore, this cannot lose a proton to form a conjugate base.  Hence, ${\text{CO}}_{\text{3}}{}^{2-}$ cannot act as Bronsted acid.

To identify the species as Bronsted base.

  ${\text{CO}}_{\text{3}}{}^{2-}+{\text{H}}^{\text{+}}\to {\text{HCO}}_{\text{3}}{}^{-}$

${\text{CO}}_{\text{3}}{}^{2-}$ accepts a proton to form ${\text{HCO}}_{\text{3}}{}^{-}$ ion.  Therefore, ${\text{CO}}_{\text{3}}{}^{2-}$ ion can act as Bronsted base.

From this we can conclude that ${\text{CO}}_{\text{3}}{}^{2-}$ can act only as Bronsted base.

(i)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (i) is Bronsted acid.

Explanation of Solution

To classify: $\text{HBr}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

  $\text{HBr}\to {\text{Br}}^{\text{-}}\text{+}{\text{H}}^{\text{+}}$

The $\text{HBr}$ can lose a proton to form a conjugate base as shown above.  Therefore, $\text{HBr}$ can act as Bronsted acid.

To identify the species as Bronsted base.

$\text{HBr}$ cannot accept proton to form a conjugate acid.

From this we can conclude that $\text{HBr}$ can act only as Bronsted acid.

(j)

Interpretation Introduction

Interpretation:

Given set of species has to be classified as Bronsted acid or base, or both.

Concept Introduction: Bronsted's definition is based on the chemical reaction that occurs when both acids and bases are added with each other.  In Bronsted's theory acid donates proton, while base accepts proton from acid resulting in the formation of water.

Example: Consider the following reaction.

  $\text{HCl}\text{+}{\text{NH}}_{\text{3}}\to {\text{NH}}_{\text{4}}{}^{\text{+}}\text{+}{\text{Cl}}^{\text{-}}$

Hydrogen chloride donates a proton, and hence it is a Bronsted acid.  Ammonia accepts a proton, and hence it is a Bronsted base.

Bronsted base accepts a proton to give a protonated species known as conjugate acid and Bronsted acid loses a proton deprotonated species is known as conjugate base.  When a proton is removed the resulting species will have a negative charge and when a proton is added the resulting species will have a positive charge.

Answer to Problem 16.3QP

The species (j) is Bronsted acid.

Explanation of Solution

To classify: $\text{HCN}$ as Bronsted acid or base, or both.

To identify the species as Bronsted acid.

  $\text{HCN}\to {\text{CN}}^{\text{-}}\text{+}{\text{H}}^{\text{+}}$

The $\text{HCN}$ can lose a proton to form a conjugate base as shown above.  Therefore, $\text{HCN}$ can act as Bronsted acid.

To identify the species as Bronsted base.

$\text{HCN}$ cannot accept proton to form a conjugate acid.

From this we can conclude that $\text{HCN}$ can act only as Bronsted acid.