Problem 5 - Chemistry and Chemical Reactivity

Question

Write balanced equations showing how the hydrogen oxalate ion, \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-},\) can be both a Bronsted acid and a Bronsted base.

Step-by-Step Solution

Answer

\( \text{HC}_2\text{O}_4^- \) can donate a proton to form \( \text{C}_2\text{O}_4^{2-} \) or accept a proton to form \(\text{H}_2\text{C}_2\text{O}_4 \).

Step 1: Understanding Bronsted Acid and Base

In the Bronsted-Lowry theory, an acid is any species that can donate a proton (H⁺), and a base is any species that can accept a proton. The hydrogen oxalate ion (\( ext{HC}_2 ext{O}_4^-\)) must be shown acting as both an acid and a base.

Step 2: As a Bronsted Acid

To show \( ext{HC}_2 ext{O}_4^-\) as an acid, it must donate a proton. When \( ext{HC}_2 ext{O}_4^-\) donates a proton, it becomes \( ext{C}_2 ext{O}_4^{2-}\) (oxalate ion). Hence, a potential equation is \[ ext{HC}_2 ext{O}_4^- ightarrow ext{C}_2 ext{O}_4^{2-} + ext{H}^+\].

Step 3: As a Bronsted Base

To show \( ext{HC}_2 ext{O}_4^-\) as a base, it must accept a proton. When \( ext{HC}_2 ext{O}_4^-\) accepts a proton, it forms \( ext{H}_2 ext{C}_2 ext{O}_4\) (oxalic acid). Thus, a potential equation is \[ ext{HC}_2 ext{O}_4^- + ext{H}^+ ightarrow ext{H}_2 ext{C}_2 ext{O}_4\].

Step 4: Combining the Reactions

The reactions can be combined to demonstrate that \( ext{HC}_2 ext{O}_4^-\) can act both ways: \[ ext{HC}_2 ext{O}_4^- ightleftharpoons ext{C}_2 ext{O}_4^{2-} + ext{H}^+\] and \[ ext{H}^+ + ext{HC}_2 ext{O}_4^- ightleftharpoons ext{H}_2 ext{C}_2 ext{O}_4\]. This demonstrates both losing and gaining a proton.

Key Concepts

Hydrogen Oxalate Ion

The hydrogen oxalate ion, represented as \( \text{HC}_2\text{O}_4^- \), is a fascinating chemical species that plays a versatile role in acid-base chemistry. As an anion derived from oxalic acid, this ion contains hydrogen, carbon, and oxygen, which allow it to participate in proton transfer reactions.
It has the capability to behave both as an acid and a base, making it amphiprotic. This means that, under different conditions, the hydrogen oxalate ion can either donate or accept an \( \text{H}^+ \) ion (proton), depending on the chemical environment.
In essence, the hydrogen oxalate ion serves as an intermediate form between oxalic acid \( \text{H}_2\text{C}_2\text{O}_4 \) and the oxalate ion \( \text{C}_2\text{O}_4^{2-} \). This dual function is central when studying acid-base reactions in chemistry.

Bronsted Acid

A Bronsted acid is defined as any chemical species capable of donating a proton (\( \text{H}^+ \)). The hydrogen oxalate ion, \( \text{HC}_2\text{O}_4^- \), showcases this property in certain reactions.
When acting as a Bronsted acid, the \( \text{HC}_2\text{O}_4^- \) donates an \( \text{H}^+ \) ion to another species. This donation transforms the hydrogen oxalate ion into an oxalate ion \( \text{C}_2\text{O}_4^{2-} \).
This process is represented by the equation:
\[ \text{HC}_2\text{O}_4^- \rightarrow \text{C}_2\text{O}_4^{2-} + \text{H}^+ \]
This reaction underscores the principle of Bronsted acid behavior, where the ability to donate a proton is a defining characteristic.

Bronsted Base

Conversely, a Bronsted base is defined as a species capable of accepting a proton. In certain reactions, the hydrogen oxalate ion, \( \text{HC}_2\text{O}_4^- \), operates as a Bronsted base.
When \( \text{HC}_2\text{O}_4^- \) accepts a proton, it becomes the molecule oxalic acid \( \text{H}_2\text{C}_2\text{O}_4 \). This transformation is indicative of its ability to act as a base by gaining a proton.
The reaction representing this behavior is:
\[ \text{HC}_2\text{O}_4^- + \text{H}^+ \rightarrow \text{H}_2\text{C}_2\text{O}_4 \]
This illustrates the principle of Bronsted base behavior, highlighting its role in accepting protons during chemical reactions.

Chemical Equilibrium

Chemical equilibrium refers to the state of a reaction where the rates of the forward and reverse reactions are equal, leading to a stable mixture of reactants and products.
For the hydrogen oxalate ion, its role as both a Bronsted acid and base equilibrates through reversible reactions:
\[ \text{HC}_2\text{O}_4^- \rightleftharpoons \text{C}_2\text{O}_4^{2-} + \text{H}^+ \]
and
\[ \text{H}^+ + \text{HC}_2\text{O}_4^- \rightleftharpoons \text{H}_2\text{C}_2\text{O}_4 \]
These equations show that both forward and reverse reactions occur at the same rate, achieving equilibrium in the system.
Recognizing chemical equilibrium is vital for understanding how concentrations of reactants and products persist over time, and it allows for predictions about the shifting nature of reactions under different conditions.