Question

Which equilibrium can be described as an acid base reaction using the Lewis acid - base definition but not using the Bronsted- Lowry definition? (1) \(2 \mathrm{NH}_{3}+\mathrm{H}_{2} \mathrm{SO}_{4} \rightleftharpoons 2 \mathrm{NH}_{4}^{+}+\mathrm{SO}_{4}^{2-}\) (2) \(\mathrm{NH}_{3}+\mathrm{CH}_{3} \mathrm{COOH} ?\mathrm{NH}_{4}{ }^{*}+\mathrm{CH}_{3} \mathrm{COO}^{-}\) (3) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{CH}_{3} \mathrm{COOH} \quad \mathrm{H}_{3} \mathrm{O}++\mathrm{CH}_{3} \mathrm{COO}^{-}\) (4) \(\left[\mathrm{Cu}\left(\mathrm{H}_{2} \mathrm{O}\right)_{4}\right]^{2+}+4 \mathrm{NH}_{3}=\left[\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) \(+4 \mathrm{H}_{2} \mathrm{O}\)

Short answer

Option 4: \( \left[ \mathrm{Cu}(\mathrm{H}_2\, \mathrm{O})_4 \right]^{2+} + 4 \mathrm{NH}_3 \rightleftharpoons \left[ \mathrm{Cu}(\mathrm{NH}_3)_4 \right]^{2+} + 4 \mathrm{H}_2\, \mathrm{O} \)

Step-by-step solution

Understand the Definitions

Recall that a Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor. In contrast, a Bronsted-Lowry acid is a proton donor and a Bronsted-Lowry base is a proton acceptor.

Analyze Option 1

\[ 2 \, \mathrm{NH}_3 + \mathrm{H}_2\, \mathrm{SO}_4 \rightleftharpoons 2 \, \mathrm{NH}_4^+ + \mathrm{SO}_4^{2-} \]In this reaction, \( \mathrm{NH}_3 \) is a Bronsted-Lowry base, accepting a proton from \( \mathrm{H}_2\, \mathrm{SO}_4 \), a Bronsted-Lowry acid. Since proton transfer occurs, it can be described by both definitions.

Analyze Option 2

\[ \mathrm{NH}_3 + \mathrm{CH}_3\, \mathrm{COOH} \rightleftharpoons \mathrm{NH}_4^+ + \mathrm{CH}_3\, \mathrm{COO}^- \]In this reaction, \( \mathrm{NH}_3 \) accepts a proton from \( \mathrm{CH}_3\, \mathrm{COOH} \). This involves proton transfer, fitting both Bronsted-Lowry and Lewis definitions.

Analyze Option 3

\[ \mathrm{H}_2\, \mathrm{O} + \mathrm{CH}_3\, \mathrm{COOH} \rightleftharpoons \mathrm{H}_3\, \mathrm{O}^+ + \mathrm{CH}_3\, \mathrm{COO}^- \]Here, \( \mathrm{CH}_3\, \mathrm{COOH} \) donates a proton to \( \mathrm{H}_2\, \mathrm{O} \). This proton transfer can be described by both the Bronsted-Lowry and Lewis definitions.

Analyze Option 4

\[ \left[ \mathrm{Cu}(\mathrm{H}_2\, \mathrm{O})_4 \right]^{2+} + 4 \mathrm{NH}_3 \rightleftharpoons \left[ \mathrm{Cu}(\mathrm{NH}_3)_4 \right]^{2+} + 4 \mathrm{H}_2\, \mathrm{O} \] In this reaction, \( \mathrm{NH}_3 \) donates electron pairs to \( \left[ \mathrm{Cu}(\mathrm{H}_2\, \mathrm{O})_4 \right]^{2+} \). No proton transfer occurs, so it only fits the Lewis definition, not the Bronsted-Lowry definition.

Determine the Correct Answer

Based on the analysis, the only reaction that can be described as an acid-base reaction using the Lewis definition but not using the Bronsted-Lowry definition is Option 4: \[ \left[ \mathrm{Cu}(\mathrm{H}_2\, \mathrm{O})_4 \right]^{2+} + 4 \mathrm{NH}_3 \rightleftharpoons \left[ \mathrm{Cu}(\mathrm{NH}_3)_4 \right]^{2+} + 4 \mathrm{H}_2\, \mathrm{O} \]

Key concepts

Lewis acid

A Lewis acid is a substance that can accept an electron pair. This definition expands on the traditional concept of acids by focusing on electron pair interactions instead of protons.
For example, in the reaction \[\text{[Cu(H}_2\text{O})}_4\text{]}^{2+} + 4 \text{NH}_3 \rightleftharpoons \text{[Cu(NH}_3\text{)}_4\text{]}^{2+} + 4 \text{H}_2 \text{O}\], the \[ \text{[Cu(H}_2\text{O})}_4\text{]}^{2+} \] acts as a Lewis acid because it accepts electron pairs from the \[ \text{NH}_3 \] molecules.
This definition does not involve the transfer of protons, which is why it’s distinct from the Bronsted-Lowry concept.
Keep in mind:

• Lewis acids accept electron pairs
• They can be molecules, ions, or atoms
• They do not necessarily involve proton transactions

Bronsted-Lowry acid

In the Bronsted-Lowry framework, an acid is defined as a proton donor. This theory is especially useful for understanding a wide range of acid-base reactions that involve the transfer of protons.
For instance, in the reaction \[ \text{H}_2 \text{O} + \text{CH}_3 \text{COOH} \rightleftharpoons \text{H}_3 \text{O}^+ + \text{CH}_3 \text{COO}^- \], \[ \text{CH}_3 \text{COOH} \] acts as a Bronsted-Lowry acid because it donates a proton to \[ \text{H}_2\text{O} \], forming \[ \text{H}_3\text{O}^+ \].
Principal points to remember:

• Bronsted-Lowry acids are proton donors
• They often lead to the formation of ions
• This involves transfer of \[ \text{H}^+ \] ions, unlike Lewis acids

Electron pair donor

An electron pair donor, often referred to as a Lewis base, is a substance that donates an electron pair during a reaction.
For example, in \[ \text{NH}_3 + \text{CH}_3 \text{COOH} \rightleftharpoons \text{NH}_4^+ + \text{CH}_3 \text{COO}^- \], \[ \text{NH}_3 \] acts as an electron pair donor, which makes it a Lewis base and a Bronsted-Lowry base simultaneously.
Core aspects to note:

• Electron pair donors are Lewis bases
• They can share their electron pairs with Lewis acids
• This sharing results in the formation of coordinate covalent bonds

Proton transfer

Proton transfer is the fundamental mechanism of Bronsted-Lowry acid-base reactions. It involves the movement of a \[ \text{H}^+ \] ion from one molecule to another.
For example, in the reaction \[ 2 \text{NH}_3 + \text{H}_2 \text{SO}_4 \rightleftharpoons 2 \text{NH}_4^+ + \text{SO}_4^{2-} \], protons are transferred from \[ \text{H}_2 \text{SO}_4 \] to \[ \text{NH}_3 \].
Key processes to understand:

• Proton donors are acids, and proton acceptors are bases in Bronsted-Lowry theory
• Proton transfer results in conjugate acid-base pairs
• This mechanism is central to many chemical reactions