Question

When carbon dioxide dissolves in water, it is in equilibrium with carbonic acid \(\mathrm{H}_{2} \mathrm{CO}_{3},\) which is a weak electrolyte. What solutes are present in aqueous solution of this compound? Write the chemical equation for the ionization of \(\mathrm{H}_{2} \mathrm{CO}_{3}\)

Short answer

The solutes present in the aqueous solution of carbonic acid (H₂CO₃) are H⁺ ions and HCO₃⁻ ions. The chemical equation for the ionization of carbonic acid is: \[H_2CO_3 \rightleftharpoons H^+ + HCO_3^-\]

Step-by-step solution

Identify the solutes present in the aqueous solution of carbonic acid.

Carbonic acid is a weak acid, and when it dissolves in water, it partially ionizes, meaning it splits into its constituent ions. In the case of H₂CO₃, the solutes present in an aqueous solution will be H⁺ ions and HCO₃⁻ ions.

Write the chemical equation for the ionization of carbonic acid.

To write the chemical equation, we will begin with the chemical formula of carbonic acid (H₂CO₃) and show it ionizing into its constituent ions: \[H_2CO_3 \rightleftharpoons H^+ + HCO_3^-\] This equation represents the ionization of carbonic acid when it dissolves in water, forming H⁺ and HCO₃⁻ ions. The double arrow (\(\rightleftharpoons\)) indicates that the reaction is in equilibrium, meaning it's both forward (ionization) and reverse (recombination) reactions occur simultaneously.

Key concepts

Weak Acid Equilibrium

Carbonic acid (\(\mathrm{H}_2\mathrm{CO}_3\)) is a great example of a weak acid in chemistry. Weak acids differ from strong acids because they do not fully ionize in solution.
This means only a small fraction of carbonic acid molecules in water will release hydrogen ions (\(\mathrm{H}^+\)). The rest remain intact as \(\mathrm{H}_2\mathrm{CO}_3\).
Because the dissociation (ionization) is incomplete, the solution reaches a dynamic equilibrium. This is represented by the double arrow symbol \((\rightleftharpoons)\) in the chemical equation that shows both forward and reverse reactions occurring simultaneously.

• In the forward reaction, some \(\mathrm{H}_2\mathrm{CO}_3\) molecules split into \(\mathrm{H}^+\) and \(\mathrm{HCO}_3^-\) ions.
• In the reverse reaction, these ions recombine to form \(\mathrm{H}_2\mathrm{CO}_3\) again.

Understanding weak acid equilibrium is crucial because it controls the pH of solutions where weak acids are present, including natural processes like carbonic acid's behavior in the ocean and bloodstream.

H2CO3 Chemical Equation

The ionization of carbonic acid involves writing its chemical equation. Specifically, the equation illustrates how \(\mathrm{H}_2\mathrm{CO}_3\) ionizes to form \(\mathrm{H}^+\) and \(\mathrm{HCO}_3^-\) ions.
Here's the chemical equation:\[\mathrm{H}_2\mathrm{CO}_3 \rightleftharpoons \mathrm{H}^+ + \mathrm{HCO}_3^-\] This equation highlights the weak acid nature of \(\mathrm{H}_2\mathrm{CO}_3\) by showing how it splits into ions in water.
The equation uses the double arrow \((\rightleftharpoons)\) to indicate equilibrium, meaning it's a reversible process. While forward reactions increase \(\mathrm{H}^+\) and \(\mathrm{HCO}_3^-\) concentration, the reverse reaction lets them recombine into \(\mathrm{H}_2\mathrm{CO}_3\), maintaining balance.
This balance means that at equilibrium, the concentration of \(\mathrm{H}_2\mathrm{CO}_3 \) and its ions stays relatively constant. Understanding this balance helps chemists predict how the solution's pH will behave when carbonic acid is dissolved in water.

Solutes in Aqueous Solution

In aqueous solutions, solutes are the substances dissolved in a solvent like water. When carbon dioxide dissolves, it forms carbonic acid \(\mathrm{H}_2\mathrm{CO}_3\) in water, which partially ionizes providing specific solutes.
These solutes are important to understand:

• Hydrogen ions \((\mathrm{H}^+)\): These ions make the solution acidic, lowering its pH.
• Bicarbonate ions \((\mathrm{HCO}_3^-)\): They are a common component in buffering systems, stabilizing pH by absorbing excess hydrogen ions.

In equilibrium, both these ions interact dynamically with undissociated carbonic acid molecules, harmonizing the solution.
This dynamic equilibrium is essential for many biological and environmental processes. For instance, in the human body, bicarbonate ions help maintain the pH of blood, while in oceans, they contribute to the carbon balance influencing marine life. Understanding the nature of solutes in aqueous solutions helps in predicting reactions and outcomes in different chemical environments.