Question

How would you expect the presence of each of the following solutes to affect the molar solubility of \(\mathrm{CaCO}_{3}\) in water: (a) \(\mathrm{Na}_{2} \mathrm{CO}_{3} ;\) (b) \(\mathrm{HCl} ;\) (c) \(\mathrm{NaHSO}_{4}\) ? Explain.

Short answer

(a) The presence of Na2CO3 will decrease the molar solubility of CaCO3. (b) The presence of HCl will increase the molar solubility of CaCO3. (c) The presence of NaHSO4 will increase the molar solubility of CaCO3.

Step-by-step solution

Analyzing Na2CO3 Effect on CaCO3 Solubility

Na2CO3 will dissociate into Na+ and CO32- ions in solution. Adding it to a solution of CaCO3 introduces a 'common ion' into the mix. This increases the concentration of CO32-, one of the ions in equilibrium with solid CaCO3. According to Le Chatelier's principle, the reaction will shift to the left, decreasing the solubility of CaCO3.

Analyzing HCl Effect on CaCO3 Solubility

HCl is a strong acid and will completely dissociate into H+ and Cl- ions in solution. The added H+ ions will react with CO32- ions from CaCO3 to form HCO3-. This reduces CO32- ion concentration and, according to Le Chatelier's principle, the equilibrium will shift to the right to increase the solubility of CaCO3.

Analyzing NaHSO4 Effect on CaCO3 Solubility

NaHSO4 will dissociate into Na+ and HSO4- ions in solution. The HSO4- can donate a proton (H+) to form SO42- and that H+ can react with CO32- to form HCO3-. This decreases CO32- concentration, and consistent with Le Chatelier's principle, the equilibrium will shift to the right, increasing the solubility of CaCO3.

Key concepts

Le Chatelier's principle

In chemical systems at equilibrium, Le Chatelier's principle is a fundamental concept to understand. This principle states that if a stress is applied to a system at equilibrium, the system will adjust to counteract the stress and restore a new equilibrium. In simpler terms, the system will try to "undo" the effect of the change.
When additional reactants or products are added to a chemical reaction in equilibrium, the reaction shifts to restore balance. For example, when the concentration of carbonate ions (\(\mathrm{CO}_{3}^{2-}\)) increases, due to the addition of a solute like \(\mathrm{Na}_2\mathrm{CO}_3\), the equilibrium shifts to the left, favoring the formation of solid \(\mathrm{CaCO}_3\), thus reducing its solubility.
Conversely, if carbonate ion concentration decreases, as when acids like \(\mathrm{HCl}\) or \(\mathrm{NaHSO}_4\) are added, the equilibrium shifts to the right, increasing \(\mathrm{CaCO}_3\) solubility. These shifts are the system's way of maintaining a dynamic balance, adjusting according to the changes imposed.

common ion effect

The common ion effect plays a crucial role in understanding solubility dynamics in solutions. It occurs when the addition of a solute introduces an ion already present in the solution, leading to a shift in the chemical equilibrium. This common ion modifies the solubility of the solute.
For instance, consider the dissolution of \(\mathrm{CaCO}_3\) in water, where \(\mathrm{CO}_{3}^{2-}\) is part of the equilibrium system. If \(\mathrm{Na}_2\mathrm{CO}_3\) is added, it introduces additional \(\mathrm{CO}_{3}^{2-}\) ions, hence the term "common ion."
This causes the equilibrium \(\mathrm{CaCO}_3\, (s) ightleftharpoons \mathrm{Ca}^{2+}\, (aq) + \mathrm{CO}_{3}^{2-}\, (aq)\) to shift left, decreasing ionization and reducing the solubility of \(\mathrm{CaCO}_3\).
The effect highlights that adding substances with a common ion typically reduces the solubility of a solute, a concept that has far-reaching implications in various chemical processes.

chemical equilibrium

Chemical equilibrium is a condition in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant, although not necessarily equal. Achieving equilibrium does not mean the reactions stop; instead, they continue at a steady pace in both directions.
In the context of solubility, when a solid like \(\mathrm{CaCO}_3\) dissolves in water, equilibrium is established between the solid and its ions in solution, described by the reaction:\[\mathrm{CaCO}_3\, (s) \rightleftharpoons \mathrm{Ca}^{2+}\, (aq) + \mathrm{CO}_{3}^{2-}\, (aq)\]The position of this equilibrium can be influenced by various factors, including the addition of new chemical species into the system as seen with solutes like \(\mathrm{HCl}\) or \(\mathrm{NaHSO}_4\) which interact with the carbonate ions.

• By reducing carbonate ion concentration, the solubility of the calcium carbonate increases as the equilibrium shifts to the right, highlighting the dynamic nature of equilibrium.

Understanding chemical equilibrium allows scientists and students alike to predict and manipulate the behavior of a chemical system under different conditions.