Problem 6 - Chemistry and Chemical Reactivity

Question

Write balanced equations showing how the \(\mathrm{HPO}_{4}^{2-}\) ion of sodium hydrogen phosphate, \(\mathrm{Na}_{2} \mathrm{HPO}_{4},\) can be a Bronsted acid or a Bronsted base.

Step-by-Step Solution

Answer

\(\mathrm{HPO}_{4}^{2-}\) acts as a Bronsted acid: \(\mathrm{HPO}_{4}^{2-} \rightarrow \mathrm{PO}_{4}^{3-} + \mathrm{H}^+\). As a Bronsted base: \(\mathrm{HPO}_{4}^{2-} + \mathrm{H}^+ \rightarrow \mathrm{H}_{2}\mathrm{PO}_{4}^{-}\).

Step 1: Identify Ion Structure

Sodium hydrogen phosphate (\(\mathrm{Na}_{2} \mathrm{HPO}_{4}\)) contains the \((\mathrm{HPO}_{4}^{2-})\) ion. This ion can either donate a hydrogen ion \((\mathrm{H}^+)\) acting as an acid, or accept a hydrogen ion, acting as a base.

Step 2: Write the Bronsted Acid Reaction

When acting as a Bronsted acid, \(\mathrm{HPO}_{4}^{2-}\) can donate a hydrogen ion \((\mathrm{H}^+)\) to become \(\mathrm{PO}_{4}^{3-}\):\[\mathrm{HPO}_{4}^{2-} \rightarrow \mathrm{PO}_{4}^{3-} + \mathrm{H}^+\]This reaction shows that \(\mathrm{HPO}_{4}^{2-}\) is donating a proton to become \(\mathrm{PO}_{4}^{3-}\).

Step 3: Write the Bronsted Base Reaction

When acting as a Bronsted base, \(\mathrm{HPO}_{4}^{2-}\) can accept a hydrogen ion \((\mathrm{H}^+)\) to become \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\):\[\mathrm{HPO}_{4}^{2-} + \mathrm{H}^+ \rightarrow \mathrm{H}_{2}\mathrm{PO}_{4}^{-}\]In this reaction, \(\mathrm{HPO}_{4}^{2-}\) is accepting a proton to form \(\mathrm{H}_{2}\mathrm{PO}_{4}^{-}\).

Key Concepts

Sodium Hydrogen Phosphate

Sodium hydrogen phosphate, chemically represented as \(\mathrm{Na}_2 \mathrm{HPO}_4\), is a salt that combines two sodium ions with one hydrogen phosphate ion \((\mathrm{HPO}_4^{2-})\). This chemical compound is an example of the versatile nature of phosphates in chemical reactions.
Sodium hydrogen phosphate is widely used in many applications, including as food additives and in water treatment. Understanding its dual role in Bronsted acid-base reactions is crucial for grasping its function in these processes.
In a solution, sodium hydrogen phosphate dissociates into its ionic components: two sodium ions \((\mathrm{Na}^+)\) and one hydrogen phosphate ion. The \(\mathrm{HPO}_4^{2-}\) ion is the active participant in Bronsted acid-base reactions, as it can either donate or accept a hydrogen ion \((\mathrm{H}^+)\).

Hydrogen Ion Transfer

Hydrogen ion transfer is a central concept of Bronsted acid-base reactions. It involves the movement of hydrogen ions \((\mathrm{H}^+)\) between different species, determining their ability to act as acids or bases.
When the \(\mathrm{HPO}_4^{2-}\) ion acts as a Bronsted acid, it donates a \((\mathrm{H}^+)\) to form \(\mathrm{PO}_4^{3-}\). In essence, this process reflects the loss of a hydrogen ion from the hydrogen phosphate ion. As a result, \(\mathrm{HPO}_4^{2-}\) is transformed into a phosphate ion \(\mathrm{PO}_4^{3-}\).
Conversely, when the \(\mathrm{HPO}_4^{2-}\) ion acts as a Bronsted base, it accepts a hydrogen ion. This transforms the ion into \(\mathrm{H}_2\mathrm{PO}_4^{-}\), indicating the addition of a \((\mathrm{H}^+)\). By understanding these transformations, it's easier to visualize how substances can switch roles between acids and bases in different chemical environments.

Phosphate Ion

The phosphate ion, noted as \(\mathrm{PO}_4^{3-}\), forms when \(\mathrm{HPO}_4^{2-}\) loses a hydrogen ion. It's an essential participant in many biological and chemical systems.
This ion is known for its significant role in buffering systems and biological energy cycles, such as ATP in cells. In the context of Bronsted acid-base chemistry, recognizing how \(\mathrm{PO}_4^{3-}\) forms enables a deeper understanding of phosphate's behavior in environmental and biological systems.
The ability of \(\mathrm{HPO}_4^{2-}\) to convert into \(\mathrm{PO}_4^{3-}\) by losing a hydrogen ion underscores the dynamic nature of phosphate compounds. These ions not only serve as vital structural components in organisms but also as functional entities in metabolism and ecological nutrient cycles. Understanding these transformations gives insight into the broad implications and applications of phosphate chemistry.