Problem 7 - Chemistry: Principles and Reactions

Question

Write a balanced equation showing how the \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ion can be either a Bronsted-Lowry acid or a Bronsted-Lowry base.

Step-by-Step Solution

Answer

Question: Write two balanced equations showing \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) acting as a Bronsted-Lowry acid and a Bronsted-Lowry base. Answer: As an acid: \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-} \longleftrightarrow \mathrm{H}^+ + \mathrm{H}\mathrm{PO}_{4}^{2-}\); As a base: \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-} + \mathrm{H}^+ \longleftrightarrow \mathrm{H}_{3}\mathrm{PO}_{4}\).

Step 1: Acting as an Acid

When \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) acts as an acid, it donates a proton, releasing \(\mathrm{H}^+\) and forming \(\mathrm{H}\mathrm{PO}_{4}^{2-}\). The reaction can be represented as: \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-} \longleftrightarrow \mathrm{H}^+ + \mathrm{H}\mathrm{PO}_{4}^{2-}\)

Step 2: Acting as a Base

When \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) acts as a base, it accepts a proton, forming \(\mathrm{H}_{3}\mathrm{PO}_{4}\). The reaction can be represented as: \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-} + \mathrm{H}^+ \longleftrightarrow \mathrm{H}_{3}\mathrm{PO}_{4}\)

Key Concepts

Understanding Acid-Base Reactions

Acid-base reactions are central concepts in chemistry that explain how substances interact through the donation and acceptance of protons. In the Bronsted-Lowry theory, an acid is defined as a species that can donate a proton, while a base is one that can accept a proton.
In our example, the \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) species can act as both an acid and a base. This dual ability is characteristic of molecules that can undergo diprotic behavior.

• When acting as an acid, it gives away a proton and creates \(\mathrm{H}\mathrm{PO}_{4}^{2-}\).
• When behaving as a base, it accepts a proton to form \(\mathrm{H}_{3}\mathrm{PO}_{4}\).

This kind of reaction can easily be represented through balanced chemical equations, showing both the proton donor and acceptor in action.

Balanced Chemical Equations in Acid-Base Reactions

In any chemical reaction, balancing equations is a crucial step. This ensures that the same number of each type of atom appears on both sides of the equation, conserving mass and complying with the law of conservation of matter.
When \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) acts as an acid, the balanced equation is:\[\mathrm{H}_{2}\mathrm{PO}_{4}^{-} \rightarrow \mathrm{H}^+ + \mathrm{HPO}_{4}^{2-}\].

• This shows one proton is released and balances with the formation of the \(\mathrm{HPO}_{4}^{2-}\) ion.
• In the reverse, when \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) behaves as a base, it accepts a proton:\[\mathrm{H}_{2}\mathrm{PO}_{4}^{-} + \mathrm{H}^+ \rightarrow \mathrm{H}_{3}\mathrm{PO}_{4}\].

Both equations illustrate how the proton transfer maintains balance, showing how these concepts work in harmony.

Proton Transfer in Acid-Base Chemistry

Proton transfer is a fundamental aspect of acid-base reactions, defined very clearly in the Bronsted-Lowry theory. It is essentially the movement of a proton (\(\mathrm{H}^+\)) between two chemical species. This transfer is what distinguishes a Bronsted-Lowry acid from a base.
For \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\) , proton transfer shows its dual nature:

• When it loses a proton, it turns into \(\mathrm{HPO}_{4}^{2-}\), exemplifying its role as a Bronsted-Lowry acid.
• Conversely, accepting a proton transforms it into \(\mathrm{H}_{3}\mathrm{PO}_{4}\), aligning with its function as a Bronsted-Lowry base.

This simplified movement embodies the elegance of chemical reactions, where even a single proton transfer has profound implications on substance properties. Understanding proton transfer simplifies the grasp of intricate acid-base reactions in chemistry.