Question

Write stepwise chemical equations for protonation or deprotonation of each of these polyprotic acids and bases in water. (a) \(\mathrm{CO}_{3}^{2-}\) (b) \(\mathrm{H}_{3} \mathrm{AsO}_{4}\) (c) \(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{COO}^{-}\) (glycinate ion, a diprotic base)

Short answer

Write the equations for stepwise protonation/deprotonation to understand each acid or base process.

Step-by-step solution

Understanding Polyprotic Species

First, recognize that polyprotic acids and bases can donate or accept protons in multiple steps. For example, carbonate ion \( \mathrm{CO}_{3}^{2-} \) can be protonated to bicarbonate \( \mathrm{HCO}_{3}^{-} \) and then to carbonic acid \( \mathrm{H}_{2}\mathrm{CO}_{3} \). Similarly, \( \mathrm{H}_{3}\mathrm{AsO}_{4} \) can lose three protons stepwise.

Protonation of Carbonate Ion

Write the first protonation reaction for carbonate:\[ \mathrm{CO}_{3}^{2-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{HCO}_{3}^{-} + \mathrm{OH}^{-} \] Then, the bicarbonate ion can be further protonated:\[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{H}_{2}\mathrm{CO}_{3} + \mathrm{OH}^{-} \]

Deprotonation of Arsenic Acid

Write the stepwise deprotonation for \( \mathrm{H}_{3}\mathrm{AsO}_{4} \):\[ \mathrm{H}_{3}\mathrm{AsO}_{4} \rightarrow \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} + \mathrm{H}^{+} \] Next, deprotonate \( \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} \):\[ \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} \rightarrow \mathrm{HAsO}_{4}^{2-} + \mathrm{H}^{+} \] Finally, deprotonate \( \mathrm{HAsO}_{4}^{2-} \):\[ \mathrm{HAsO}_{4}^{2-} \rightarrow \mathrm{AsO}_{4}^{3-} + \mathrm{H}^{+} \]

Deprotonation of Glycinate Ion

Glycinate ion \( \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COO}^{-} \) can gain protons since it's a diprotic base. First, protonate the carboxylate group:\[ \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COO}^{-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{OH}^{-} \] Then, protonate the amino group:\[ \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{NH}_{3}^{+}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{OH}^{-} \]

Key concepts

Protonation

Protonation is a chemical process where an atom, molecule, or ion gains a proton (hydrogen ion, \( \mathrm{H}^{+} \)). Polyprotic acids and bases can undergo multiple protonation steps. Let's explore this with some examples.

For the carbonate ion \( \mathrm{CO}_{3}^{2-} \), protonation occurs in two steps:

• First, the carbonate ion reacts with water to form bicarbonate \( \mathrm{HCO}_{3}^{-} \):
  \[ \mathrm{CO}_{3}^{2-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{HCO}_{3}^{-} + \mathrm{OH}^{-} \]
• Then, the bicarbonate can further react with water, forming carbonic acid \( \mathrm{H}_{2}\mathrm{CO}_{3} \):
  \[ \mathrm{HCO}_{3}^{-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{H}_{2}\mathrm{CO}_{3} + \mathrm{OH}^{-} \]

Each protonation step reduces the negative charge, making the ion neutral eventually or simply less negative.

For the glycinate ion \( \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COO}^{-} \), as a diprotic base, it can undergo protonation at two sites:

• First, the carboxylate group accepts a proton to form \( \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COOH} \):
  \[ \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COO}^{-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{OH}^{-} \]
• Second, the amino group is protonated, resulting in \( \mathrm{NH}_{3}^{+}\mathrm{CH}_{2}\mathrm{COOH} \):
  \[ \mathrm{NH}_{2}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{NH}_{3}^{+}\mathrm{CH}_{2}\mathrm{COOH} + \mathrm{OH}^{-} \]

This series of steps is crucial for understanding how polyprotic ions behave in solutions.

Deprotonation

Deprotonation is the reverse process of protonation, where a proton \( \mathrm{H}^{+} \) is removed from a molecule. This process is common in polyprotic acids, which can donate multiple protons sequentially.

Consider the deprotonation of arsenic acid \( \mathrm{H}_{3}\mathrm{AsO}_{4} \). It occurs in the following steps:

• First, arsenic acid loses a proton to form dihydrogen arsenate \( \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} \):
  \[ \mathrm{H}_{3}\mathrm{AsO}_{4} \rightarrow \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} + \mathrm{H}^{+} \]
• Next, dihydrogen arsenate loses a proton to become monohydrogen arsenate \( \mathrm{HAsO}_{4}^{2-} \):
  \[ \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} \rightarrow \mathrm{HAsO}_{4}^{2-} + \mathrm{H}^{+} \]
• Finally, monohydrogen arsenate releases its proton to form arsenate ion \( \mathrm{AsO}_{4}^{3-} \):
  \[ \mathrm{HAsO}_{4}^{2-} \rightarrow \mathrm{AsO}_{4}^{3-} + \mathrm{H}^{+} \]

These reactions highlight the stepwise reduction in proton numbers, illustrating how acids can donate protons in stages.

Understanding these steps helps in predicting the behavior of acids in different chemical environments.

Chemical Equations

Chemical equations are essential for representing the protonation and deprotonation processes of polyprotic acids and bases. These equations use symbols and formulas to depict what occurs chemically during a reaction.

For example, the chemical equation for the first protonation of the carbonate ion \( \mathrm{CO}_{3}^{2-} \) is:
\[ \mathrm{CO}_{3}^{2-} + \mathrm{H}_{2}\mathrm{O} \rightarrow \mathrm{HCO}_{3}^{-} + \mathrm{OH}^{-} \]
This equation shows that one proton from water is added to the carbonate ion.

Another example is the deprotonation of arsenic acid, where:
\[ \mathrm{H}_{3}\mathrm{AsO}_{4} \rightarrow \mathrm{H}_{2}\mathrm{AsO}_{4}^{-} + \mathrm{H}^{+} \]
This equation indicates the removal of a proton from arsenic acid.

We write chemical equations in balance to ensure the same number of each type of atom exists on both the reactant and product sides, preserving mass. Furthermore, including states of matter can provide additional context about the reactants and products involved, e.g., whether they are in aqueous form or gaseous.

Creating balanced chemical equations is crucial for understanding and predicting chemical reactions in labs and industries.