Question

Write two chemical equations, one in which \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) is a Bronsted acid (in reaction with the carbonate ion, \(\left.\mathrm{CO}_{3}^{2-}\right),\) and a second in which \(\mathrm{HPO}_{4}^{2-}\) is a Bronsted base (in reaction with acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) ).

Short answer

H2PO4- reacts with CO3^2- to form HPO4^2- and HCO3-. HPO4^2- reacts with CH3CO2H to form H2PO4- and CH3CO2-.

Step-by-step solution

Identify the reaction for H2PO4- as a Bronsted acid

Bronsted acids donate protons (H+ ions). The dihydrogen phosphate ion, \( \mathrm{H}_2\mathrm{PO}_4^{-}\), can donate a hydrogen ion to the carbonate ion, \( \mathrm{CO}_3^{2-} \). Thus, the Bronsted acid reaction is:\[ \mathrm{H}_2\mathrm{PO}_4^{-} + \mathrm{CO}_3^{2-} \rightarrow \mathrm{HPO}_4^{2-} + \mathrm{HCO}_3^{-} \]

Identify the reaction for HPO4^2- as a Bronsted base

Bronsted bases accept protons. The hydrogen phosphate ion, \( \mathrm{HPO}_4^{2-} \), can accept a hydrogen ion from acetic acid, \( \mathrm{CH}_3\mathrm{CO}_2\mathrm{H} \). Thus, the Bronsted base reaction is:\[ \mathrm{HPO}_4^{2-} + \mathrm{CH}_3\mathrm{CO}_2\mathrm{H} \rightarrow \mathrm{H}_2\mathrm{PO}_4^{-} + \mathrm{CH}_3\mathrm{CO}_2^{-} \]

Key concepts

Acid-Base Reactions

Understanding acid-base reactions is crucial in chemistry, especially when dealing with different ions. According to the Bronsted-Lowry theory, acids are substances that can donate protons (i.e., hydrogen ions, H extsuperscript{+}), while bases are those that can accept these protons. This concept is essential for identifying the roles of various substances in chemical reactions.
In the given problem, dihydrogen phosphate ion (\(\mathrm{H}_2\mathrm{PO}_4^{-}\)) acts as a Bronsted acid because it donates a proton to the carbonate ion (\(\mathrm{CO}_3^{2-}\)), forming hydrogen carbonate (\(\mathrm{HCO}_3^{-}\)). The balanced reaction is:

• \(\mathrm{H}_2\mathrm{PO}_4^{-} + \mathrm{CO}_3^{2-} \rightarrow \mathrm{HPO}_4^{2-} + \mathrm{HCO}_3^{-}\)

This showcases a typical acid-base reaction where the acid loses a proton, forming a conjugate base.
On the flip side, hydrogen phosphate ion (\(\mathrm{HPO}_4^{2-}\)) acts as a Bronsted base when it accepts a proton from acetic acid (\(\mathrm{CH}_3\mathrm{CO}_2\mathrm{H}\)), forming dihydrogen phosphate (\(\mathrm{H}_2\mathrm{PO}_4^{-}\)) and acetate ion (\(\mathrm{CH}_3\mathrm{CO}_2^{-}\)).

• \(\mathrm{HPO}_4^{2-} + \mathrm{CH}_3\mathrm{CO}_2\mathrm{H} \rightarrow \mathrm{H}_2\mathrm{PO}_4^{-} + \mathrm{CH}_3\mathrm{CO}_2^{-}\)

This reaction illustrates the classic action of a base gaining a proton to form a conjugate acid.

Dihydrogen Phosphate

Dihydrogen phosphate, or \(\mathrm{H}_2\mathrm{PO}_4^{-}\), is an anion with diverse roles in biological and chemical systems. It is a polyatomic ion, formed from a combination of hydrogen, phosphorus, and oxygen.
This ion plays a noteworthy role as a Bronsted acid. In its role as an acid, \(\mathrm{H}_2\mathrm{PO}_4^{-}\) can donate a hydrogen ion to a base, leading to the formation of hydrogen phosphate (\(\mathrm{HPO}_4^{2-}\)). This specific conversion exemplifies its ability to act within buffer solutions.

• Helps maintain pH levels in solutions by acting as an acid or base when necessary.
• Used in biological systems as part of energy transfers and enzymatic reactions.

Its dualistic nature, being able to act as an acid or base, makes \(\mathrm{H}_2\mathrm{PO}_4^{-}\) a vital component in various chemical and biological settings.

Carbonate Ion

The carbonate ion, or \(\mathrm{CO}_3^{2-}\), is an important polyatomic ion in both the geological and biological spheres. It consists of one carbon atom centrally bonded to three oxygen atoms, creating a trigonal planar shape.
In the context of Bronsted-Lowry acid-base reactions, \(\mathrm{CO}_3^{2-}\) behaves predominantly as a base because it has the tendency to accept protons. When reacting with a Bronsted acid like dihydrogen phosphate (\(\mathrm{H}_2\mathrm{PO}_4^{-}\)), it forms the hydrogen carbonate ion (\(\mathrm{HCO}_3^{-}\)).

• In nature, carbonate ions are found in mineral deposits and are a part of the Earth's carbon cycle.
• They play a key role in buffering systems, helping to maintain consistent pH levels in environmental and physiological processes.

The flexibility of the carbonate ion to act in acid-base reactions enhances its significance, particularly in buffering environments and regulating acidity.