The species: H2O, HCO3 –, HSO4 – and NH3 can act b – Class 11 Chemistry

Exercises · 6.39

Q. The species: H2O, HCO3 -, HSO4 - and NH3 can act both as Brönsted acids and bases. For each case give the corresponding conjugate acid and base.

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Concept understanding — Brønsted Lowry Conjugate Pairs

From Intuition to Precision: Brønsted–Lowry Conjugate Pairs

Imagine you're at a party where people are passing around a single, very special coin. The coin represents a proton ( $\text{H}^+$ ). The game is simple: someone can give the coin to someone else, but only if the other person is willing to take it. You can't just throw the coin; you need a willing receiver.

In chemistry, acids and bases play exactly this game. An acid is the person who gives away the proton. A base is the person who accepts it. But here's the twist: the moment the acid gives away its proton, it transforms into something new — something that can now accept a proton back. That transformed form is called the conjugate base. Similarly, the base, after accepting the proton, becomes something that can donate it back — the conjugate acid.

This is the core of the Brønsted–Lowry theory: every acid-base reaction involves a pair of substances that are linked by the gain or loss of a single proton.

The Precise Statement

\text{Acid} \rightleftharpoons \text{Conjugate Base} + \text{H}^+

A conjugate acid-base pair consists of two species that differ by exactly one proton ( $\text{H}^+$ ). The acid has the proton; the conjugate base does not.

When an acid donates a proton, it becomes its conjugate base.

When a base accepts a proton, it becomes its conjugate acid.

Important

In any Brønsted–Lowry reaction, there are two conjugate pairs: one on the reactant side and one on the product side. They always appear together.

Seeing It in Action

Take the classic reaction between hydrochloric acid and water:

\text{HCl} + \text{H}_2\text{O} \rightarrow \text{Cl}^- + \text{H}_3\text{O}^+

Let's identify the pairs:

Pair 1: $\text{HCl}$ (acid) and $\text{Cl}^-$ (conjugate base). They differ by one $\text{H}^+$ .
Pair 2: $\text{H}_2\text{O}$ (base) and $\text{H}_3\text{O}^+$ (conjugate acid). They also differ by one $\text{H}^+$ .

Notice: water acted as a base here — it accepted the proton from HCl. But water can also act as an acid in other reactions. That's the beauty of the Brønsted–Lowry theory: a substance's role depends on the reaction, not on a fixed label.

Tip

To find the conjugate base of any acid, simply remove one $\text{H}^+$ and reduce the charge by +1.

To find the conjugate acid of any base, add one $\text{H}^+$ and increase the charge by +1.

A Quick Reference Table

Acid	Conjugate Base	Base	Conjugate Acid
$\text{HCl}$	$\text{Cl}^-$	$\text{NH}_3$	$\text{NH}_4^+$
$\text{H}_2\text{SO}_4$	$\text{HSO}_4^-$	$\text{H}_2\text{O}$	$\text{H}_3\text{O}^+$
$\text{NH}_4^+$	$\text{NH}_3$	$\text{OH}^-$	$\text{H}_2\text{O}$
$\text{H}_2\text{O}$	$\text{OH}^-$	$\text{CO}_3^{2-}$	$\text{HCO}_3^-$

Notice how water appears in both columns — it's amphoteric, meaning it can act as either an acid or a base depending on its partner.

Why This Matters …

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Concept: Brønsted–Lowry theory defines an acid as a proton donor and a base as a proton acceptor. Every species that can donate or accept a proton has a conjugate partner differing by exactly one $\mathrm{H}^+$ .

Reasoning:

Conjugate acid = species + $\mathrm{H}^+$ (what forms when it accepts a proton)
Conjugate base = species − $\mathrm{H}^+$ (what remains after it donates a proton)

Apply to each amphoteric species:

Species	Acts as base → Conjugate acid	Acts as acid → Conjugate base
$\mathrm{H_2O}$	$\mathrm{H_3O}^+$	$\mathrm{OH}^-$
$\mathrm{HCO_3}^-$	$\mathrm{H_2CO_3}$	$\mathrm{CO_3}^{2-}$

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Amphiprotic species can donate or accept a proton. For each species, the conjugate acid is formed by adding $\text{H}^+$ and the conjugate base by removing $\text{H}^+$ : $\text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ / \text{OH}^-$ , $\text{HCO}_3^- \rightleftharpoons \text{H}_2\text{CO}_3 / \text{CO}_3^{2-}$ , $\text{HSO}_4^- \rightleftharpoons \text{H}_2\text{SO}_4 / \text{SO}_4^{2-}$ , $\text{NH}_3 \rightleftharpoons \text{NH}_4^+ / \text{NH}_2^-$ .

The Brønsted–Lowry theory defines acids as proton donors and bases as proton acceptors. A conjugate pair differs by exactly one proton: when an acid donates $\text{H}^+$ , what remains is its conjugate base; when a base accepts $\text{H}^+$ , it becomes its conjugate acid.

Species that can both donate and accept protons are called amphiprotic (or amphoteric in the Brønsted sense). They sit in the middle of a conjugate acid-base trio. To find the conjugate acid of any species, imagine it accepting a proton—add $\text{H}^+$ . To find the conjugate base, imagine it donating a proton—remove $\text{H}^+$ .

Let's work through each species systematically.

1. Water, $\text{H}_2\text{O}$

Water is the textbook amphiprotic molecule.

As a Brønsted base (accepting $\text{H}^+$ ):

\text{H}_2\text{O} + \text{H}^+ \longrightarrow \text{H}_3\text{O}^+

The conjugate acid is the hydronium ion, $\text{H}_3\text{O}^+$ .

As a Brønsted acid (donating $\text{H}^+$ ):

\text{H}_2\text{O} \longrightarrow \text{H}^+ + \text{OH}^-

The conjugate base is the hydroxide ion, $\text{OH}^-$ .

2. Bicarbonate ion, $\text{HCO}_3^-$

This ion plays a central role in blood buffering and carbonate equilibria.

As a Brønsted base (accepting $\text{H}^+$ ):

\text{HCO}_3^- + \text{H}^+ \longrightarrow \text{H}_2\text{CO}_3

The conjugate acid is carbonic acid, $\text{H}_2\text{CO}_3$ .

As a Brønsted acid (donating $\text{H}^+$ ):

\text{HCO}_3^- \longrightarrow \text{H}^+ + \text{CO}_3^{2-}

The conjugate base is the carbonate ion, $\text{CO}_3^{2-}$ .

3. Hydrogen sulfate ion, $\text{HSO}_4^-$

Sulfuric acid is diprotic, and $\text{HSO}_4^-$ is its intermediate form.

As a Brønsted base (accepting $\text{H}^+$ ):

\text{HSO}_4^- + \text{H}^+ \longrightarrow \text{H}_2\text{SO}_4

The conjugate acid is sulfuric acid, $\text{H}_2\text{SO}_4$ .

As a Brønsted acid (donating $\text{H}^+$ ):

\text{HSO}_4^- \longrightarrow \text{H}^+ + \text{SO}_4^{2-}

The conjugate base is the sulfate ion, $\text{SO}_4^{2-}$ .

Note

In aqueous solution, $\text{HSO}_4^-$ is a moderately strong acid ( $K_a \approx 10^{-2}$ ), so it more readily acts as an acid than a base under typical conditions.

4. Ammonia, $\text{NH}_3$

Ammonia is usually thought of as a base, but it can donate a proton under strongly basic conditions.

As a Brønsted base (accepting $\text{H}^+$ ): …

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Q.The species: H2O, HCO3 –, HSO4 – and NH3 can act both as Brönsted acids and bases. For each case give the corresponding conjugate acid and base.

1. Water, H2O\text{H}_2\text{O}H2​O

2. Bicarbonate ion, HCO3−\text{HCO}_3^-HCO3−​

3. Hydrogen sulfate ion, HSO4−\text{HSO}_4^-HSO4−​

4. Ammonia, NH3\text{NH}_3NH3​

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1. Water, $\text{H}_2\text{O}$

2. Bicarbonate ion, $\text{HCO}_3^-$

3. Hydrogen sulfate ion, $\text{HSO}_4^-$

4. Ammonia, $\text{NH}_3$

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