Problem 19 - Chemistry An Atoms First Approach

Question

Anions containing hydrogen (for example, \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) ) usually show amphoteric behavior. Write equations illustrating the amphoterism of these two anions.

Step-by-Step Solution

Answer

The amphoteric behavior of \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) can be shown with the following equations: As acids: \[ \mathrm{HCO}_{3}^{-} + \mathrm{OH}^{-} \rightarrow \mathrm{CO}_{3}^{2-} + \mathrm{H}_{2}\mathrm{O} \] \[ \mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{OH}^{-} \rightarrow \mathrm{H}\mathrm{PO}_{4}^{2-} + \mathrm{H}_{2}\mathrm{O} \] As bases: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{2}\mathrm{CO}_{3} \] \[ \mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{3}\mathrm{PO}_{4} \] These equations illustrate how both anions can act as acids and bases, demonstrating their amphoteric behavior.

Step 1: Reaction of HCO3- as an acid

When \(\mathrm{HCO}_{3}^{-}\) acts as an acid, it donates a hydrogen ion (H+) to a base, forming \(\mathrm{CO}_{3}^{2-}\) and releasing a water molecule. The equation for this reaction is: \[ \mathrm{HCO}_{3}^{-} + \mathrm{OH}^{-} \rightarrow \mathrm{CO}_{3}^{2-} + \mathrm{H}_{2}\mathrm{O} \]

Step 2: Reaction of HCO3- as a base

When \(\mathrm{HCO}_{3}^{-}\) acts as a base, it accepts a hydrogen ion (H+) from an acid, forming \(\mathrm{H}_{2}\mathrm{CO}_{3}\), a weak acid. The equation for this reaction is: \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{2}\mathrm{CO}_{3} \]

Step 3: Reaction of H2PO4- as an acid

When \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) acts as an acid, it donates a hydrogen ion (H+) to a base, forming \(\mathrm{H}\mathrm{PO}_{4}^{2-}\) and releasing a water molecule. The equation for this reaction is: \[ \mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{OH}^{-} \rightarrow \mathrm{H}\mathrm{PO}_{4}^{2-} + \mathrm{H}_{2}\mathrm{O} \]

Step 4: Reaction of H2PO4- as a base

When \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) acts as a base, it accepts a hydrogen ion (H+) from an acid, forming \(\mathrm{H}_{3}\mathrm{PO}_{4}\), a weak acid. The equation for this reaction is: \[ \mathrm{H}_{2} \mathrm{PO}_{4}^{-} + \mathrm{H}^{+} \rightarrow \mathrm{H}_{3}\mathrm{PO}_{4} \] These equations show how the two anions, \(\mathrm{HCO}_{3}^{-}\) and \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\), demonstrate amphoteric behavior by acting as both acids and bases.

Key Concepts

acid-base reactions

Acid-base reactions are fundamental chemical processes where an acid and a base interact. An acid is a substance that can donate a proton (hydrogen ion, H+), while a base is a substance that can accept a proton. This is most simply represented by the Arrhenius definition. However, the Brønsted-Lowry definition expands on this by stating that acids and bases can participate in a reaction by donating or accepting a proton, respectively.

Acids and bases that can both donate and accept a proton exhibit amphoteric behavior. This means they can react with both stronger acids and stronger bases, acting as a base in the presence of a stronger acid, and as an acid in the presence of a stronger base. Understanding how these reactions take place is crucial for explaining a wide range of chemical phenomena, from biological processes to industrial applications.

For example, when sodium hydroxide (a strong base) is added to water, it dissociates completely to form hydroxide ions (OH-). If hydrochloric acid (a strong acid) is added, it dissociates to form H+ ions. When these ions meet, they form water in a neutralization reaction. This is represented by the equation: \[ \text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^- \]\[ \text{HCl} \rightarrow \text{H}^+ + \text{Cl}^- \]\[ \text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O} \]These simplifications assist in highlighting the core of acid-base interactions and pave the way to more complex understanding such as acid dissociation constants and pH levels.

hydrogen carbonate ion

The hydrogen carbonate ion (\(\text{HCO}_3^-\)), commonly known as bicarbonate, is a polyatomic anion with considerable importance in both natural and physiological systems. It serves as a crucial component in the body's pH buffering system and is involved in crucial processes such as maintaining acid-base balance within the blood.

As explained in the provided exercise, bicarbonate is amphoteric, which means it can act as both an acid and a base. When bicarbonate acts as an acid, it donates a proton, thereby transforming into its conjugate base, the carbonate ion (\(\text{CO}_3^{2-}\)). Conversely, when it acts as a base, it accepts a proton to form carbonic acid (\(\text{H}_2\text{CO}_3\)), which further highlights its pivotal role in regulating pH levels.

These reactions are vital in environmental science and physiology and have industrial significance, such as in the formation of carbonates in water treatment or the buffering capacity of baking soda in culinary applications. The ability of \(\text{HCO}_3^-\) to readily engage in acid-base reactions makes it an exemplary buffer, preventing significant changes in pH when acids or bases are introduced to a system.

dihydrogen phosphate ion

The dihydrogen phosphate ion (\(\text{H}_2\text{PO}_4^-\)), is another example of an amphoteric species and plays a critical role in biochemistry, particularly in energy transfer through ATP and in buffering blood plasma. It is derived from phosphoric acid (\(\text{H}_3\text{PO}_4\)), a triprotic acid that can donate three protons.

As an acid, dihydrogen phosphate donates a proton to form the hydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), and as a base, it accepts a proton to revert to phosphoric acid. The interconversion between dihydrogen phosphate and its related species is an important part of intracellular and extracellular fluid pH regulation.

The flexibility of \(\text{H}_2\text{PO}_4^-\) in acid-base reactions underpins its essential role in biological systems. For example, the buffering system in blood relies on the phosphate buffer system, which consists of \(\text{H}_2\text{PO}_4^-\) in equilibrium with its deprotonated form, \(\text{HPO}_4^{2-}\). This system works in tandem with other buffering agents to maintain the pH of blood and other bodily fluids within a very narrow, slightly alkaline range.