Question

Use Le Châtelier's principle to predict the effect of the following changes on the extent of hydrolysis of sodium nitrite \(\left(\mathrm{NaNO}_{2}\right)\) solution: (a) \(\mathrm{HCl}\) is added, (b) \(\mathrm{NaOH}\) is added, (c) \(\mathrm{NaCl}\) is added, (d) the solution is diluted.

Short answer

(a) Decreases hydrolysis, (b) Increases hydrolysis, (c) No significant effect, (d) Increases hydrolysis.

Step-by-step solution

Understand the Reaction

Sodium nitrite (NaNO₂) can be represented by the dissociation reaction: \[ \text{NaNO}_2 \rightleftharpoons \text{Na}^+ + \text{NO}_2^- \]. The NO₂⁻ anion can further react with water in a hydrolysis reaction: \[ \text{NO}_2^- + \text{H}_2\text{O} \rightleftharpoons \text{HNO}_2 + \text{OH}^- \]. Le Châtelier's principle states that if a system at equilibrium experiences a change in concentration, temperature, or pressure, the system will adjust to counteract that change and re-establish equilibrium.

Effect of Adding HCl

Adding HCl increases the concentration of H⁺ ions in the solution. According to Le Châtelier's principle, the equilibrium will shift to the left to produce more NO₂⁻ and consume OH⁻ ions, decreasing the extent of hydrolysis, while also forming water by reacting with H⁺ ions.

Effect of Adding NaOH

Adding NaOH increases the concentration of OH⁻ ions in the solution. The equilibrium will shift to the right to produce more HNO₂ and consume the added OH⁻ ions to decrease the amount of excess OH⁻, increasing the extent of hydrolysis.

Effect of Adding NaCl

Adding NaCl, a common ion, increases the concentration of Na⁺ ions. According to Le Châtelier's principle, adding a common ion generally decreases the extent of a solubility reaction. However, since Na⁺ does not participate directly in the hydrolysis reaction of NO₂⁻, the extent of hydrolysis is largely unaffected by adding NaCl.

Effect of Dilution

Diluting the solution decreases the concentration of all species, shifting the equilibrium towards the side with a greater number of particles to re-establish equilibrium. This means the hydrolysis reaction of NO₂⁻ will shift to the right, increasing the extent of hydrolysis.

Key concepts

Chemistry Equilibrium

Chemistry equilibrium refers to a state in a chemical reaction where the rate of the forward reaction equals the rate of the backward reaction, leading to a stable concentration of reactants and products. In the context of sodium nitrite hydrolysis, we observe an equilibrium reaction, where ²⁻ and water react to form HNO₂ and OH⁻ ions. Adjustments in this balance, due to external changes, follow Le Châtelier's principle. Essentially, when a change is introduced to the system, such as adding HCl or NaOH, the system naturally adjusts to minimize the impact of that change. This could result in the reaction shifting to favor the creation of products or reactants to bring the system back to equilibrium. When carefully understood, equilibrium helps predict the outcome of reactions when subjected to different conditions.

Acid-Base Reactions

Acid-base reactions are fundamental to understanding many chemical processes, including the behavior of ²⁻ during hydrolysis. Acidic or basic substances dissolve in water to alter the pH of the solution. When ON₂ d is added to an aqueous solution, it interacts with water, causing hydrolysis, an acid-base reaction producing HNO₂ (a weak acid) and OH⁻ ions. This is further influenced by adding acids like HCl, which increases H⁺ concentration, thereby driving the equilibrium to lower OH⁻ by consuming it. Furthermore, the addition of a base such as NaOH will add more OH⁻ ions, promoting the formation of more HNO₂ in an attempt to counterbalance the increased OH⁻, enhancing hydrolysis. Recognizing these interactions aids in comprehending how solutions respond to variations in acidity and basicity.

Solution Concentration

Solution concentration plays a critical role in determining the direction of equilibrium shifts. By altering concentrations, like when NaCl is added sourcing a common ion (Na⁺), we might expect changes due to the common ion effect. However, in our specific reaction, Na⁺ is not directly involved in the hydrolysis process. Hence, its presence, while increasing primary ion concentration, does not impact the ongoing equilibrium between O₂⁻ and water. For dilution, reducing overall concentration nudges the system toward the formation of more ions (in this case HNO₂ and OH⁻), because the system attempts to reach a new equilibrium by producing more output to compensate for the decrease in concentration density. Therefore, dynamic solutions facilitate gaining insight into how systems adjust upon concentration shifts.

Ionic Compounds

Ionic compounds, like sodium nitrite (NaNO₂), dissociate in water to form ions that participate in various chemical reactions. These compounds consist of ions held together by electrostatic forces, disbanding in solution to forge new pathways, such as hydrolysis in this scenario. When NaNO₂ dissociates, it forms Na⁺ and NO₂⁻ ions, the latter engaging in equilibrium reactions with water. The behavior of such ions under varying conditions—whether in presence of acids, bases, or changes in concentration—is remarkably predictable using principles like Le Châtelier's. They highlight how equilibrium can pivotingly hinge on bond strength and ion interaction property. Understanding ionic behavior is integral to mastering the nuances of chemical reactions.