Problem 40 - Chemistry A Molecular Approach

Question

Both \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{HS}^{-}\) are amphoteric. Write an equation to show how each substance can act as an acid and another equation to show how each can act as a base.

Step-by-Step Solution

Answer

As acids, \(\mathrm{HCO}_{3}^{-}\) forms \(\mathrm{CO}_{3}^{2-} + \mathrm{H}^{+}\) and \(\mathrm{HS}^{-}\) forms \(\mathrm{S}^{2-} + \mathrm{H}^{+}\). As bases, \(\mathrm{HCO}_{3}^{-}\) forms \(\mathrm{H}_{2}\mathrm{CO}_{3}\) and \(\mathrm{HS}^{-}\) forms \(\mathrm{H}_{2}\mathrm{S}\) when reacting with \(\mathrm{H}^{+}\).

Step 1: Acting as an Acid for \(\mathrm{HCO}_{3}^{-}\)

To act as an acid, the \(\mathrm{HCO}_{3}^{-}\) must donate a proton \(\mathrm{(H^{+})}\). Write down the ionization reaction where \(\mathrm{HCO}_{3}^{-}\) donates a proton to form \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{H}^{+}\).\[\mathrm{HCO}_{3}^{-} \rightarrow \mathrm{CO}_{3}^{2-} + \mathrm{H}^{+}\]

Step 2: Acting as a Base for \(\mathrm{HCO}_{3}^{-}\)

To act as a base, the \(\mathrm{HCO}_{3}^{-}\) must accept a proton. Write down the ionization reaction where \(\mathrm{HCO}_{3}^{-}\) accepts a proton to form \(\mathrm{H}_{2}\mathrm{CO}_{3}\).\[\mathrm{HCO}_{3}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{2}\mathrm{CO}_{3}\]

Step 3: Acting as an Acid for \(\mathrm{HS}^{-}\)

To act as an acid, the \(\mathrm{HS}^{-}\) must donate a proton. Write down the ionization reaction where \(\mathrm{HS}^{-}\) donates a proton to form \(\mathrm{S}^{2-}\) and \(\mathrm{H}^{+}\).\[\mathrm{HS}^{-} \rightarrow \mathrm{S}^{2-} + \mathrm{H}^{+}\]

Step 4: Acting as a Base for \(\mathrm{HS}^{-}\)

To act as a base, the \(\mathrm{HS}^{-}\) must accept a proton. Write down the ionization reaction where \(\mathrm{HS}^{-}\) accepts a proton to form \(\mathrm{H}_{2}\mathrm{S}\).\[\mathrm{HS}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{2}\mathrm{S}\]

Key Concepts

Acid-Base Reactions

Acid-base reactions are essential processes in chemistry involving the exchange of protons between reactants. These reactions are fundamental to understanding the behavior of many substances, including amphoteric compounds, which can act as both acids and bases. In acid-base chemistry, the substance that donates a proton (a hydrogen ion, \( \mathrm{H}^{+} \)) is called an acid, and the substance that accepts a proton is referred to as a base. This proton exchange dramatically affects the pH levels of solutions and plays a critical role in chemical reactions and biological processes alike.

For example, \( \mathrm{HCO}_{3}^{-} \) acts as an acid when it donates a proton to form \( \mathrm{CO}_{3}^{2-} \) and a \( \mathrm{H}^{+} \) ion. Conversely, it behaves as a base when it accepts a proton to form \( \mathrm{H}_{2}\mathrm{CO}_{3} \). Understanding these reactions helps decipher the intricate balance of various biological and environmental systems.

Proton Transfer Reactions

Proton transfer reactions, or protolysis, represent a specific type of acid-base reaction where a proton is transferred from one molecule to another. It's a critical concept that explains how acids and bases interact in solution. The simplicity of these reactions belies their importance; the transfer of a single, charged particle can lead to significant changes in the properties of substances involved.

It's worth noting that the solvent plays a key role in facilitating proton transfers. For example, water can act as a mediator for proton exchange, holding onto the proton momentarily to ease the transfer between reactants. These reactions are reversible and reach a point of dynamic chemical equilibrium, where the rate of the forward reaction equals the rate of the reverse reaction, stabilizing the concentration of products and reactants.

Bicarbonate Ion

The bicarbonate ion (\( \mathrm{HCO}_{3}^{-} \)) is a pivotal species in physiological processes and environmental systems. It acts as a buffer, helping to maintain pH stability in blood and other biological fluids. Bicarbonate's amphoteric nature allows it to neutralize excess acids or bases, which is vital for keeping biological systems functioning correctly.

When \( \mathrm{HCO}_{3}^{-} \) acts as an acid, it loses a proton to become carbonate (\( \mathrm{CO}_{3}^{2-} \)). On the other hand, when acting as a base, bicarbonate gains a proton to form carbonic acid (\( \mathrm{H}_{2}\mathrm{CO}_{3} \)). These reactions are essential for carbon dioxide transport and excretion in respiration and are integral to the Earth's carbon cycle.

Hydrosulfide Ion

The hydrosulfide ion (\( \mathrm{HS}^{-} \)) is another example of an amphoteric substance. It occurs in various sulfur-containing compounds and is involved in some biochemical reactions. Its ability to donate or accept protons enables it to partake in different chemical reactions, depending on the pH of the environment.

As an acid, \( \mathrm{HS}^{-} \) donates a proton to form sulfide (\( \mathrm{S}^{2-} \)) while producing a \( \mathrm{H}^{+} \) ion. Conversely, when acting as a base, it accepts a proton to become hydrogen sulfide (\( \mathrm{H}_{2}\mathrm{S} \)), an important compound in the natural sulfur cycle but also known for its characteristic rotten egg smell.

Chemical Equilibrium

Chemical equilibrium is a state in which the rates of the forward and reverse reactions are equal, leading to no overall change in the amounts of reactants and products over time. It is a dynamic balance, not a static situation, meaning that the reactions continue to occur, but the concentrations remain constant. In the context of amphoteric substances like \( \mathrm{HCO}_{3}^{-} \) and \( \mathrm{HS}^{-} \), the concept of chemical equilibrium helps us understand why these substances can maintain a relatively stable pH in solutions.

Both the reactions showing how \( \mathrm{HCO}_{3}^{-} \) and \( \mathrm{HS}^{-} \) can act as acids or bases involve equilibrium processes. The equilibrium lies at the heart of their buffering capacities, ensuring that even when small amounts of acids or bases are added to a system, the pH doesn't change drastically. This principle is fundamental to many natural and industrial processes, from blood pH regulation to wastewater treatment.