Question

Describe the effects of the salts \(\mathrm{KI}\) and \(\mathrm{AgNO}_{3}\) on the solubility of AgI in water.

Short answer

The presence of both KI and AgNO3 in solution decreases the solubility of AgI due to the common ion effect. Higher concentration of I- ions from KI and Ag+ ions from AgNO3 shifts the AgI dissolution equilibrium to the left, reducing the solubility.

Step-by-step solution

Understand the reaction of AgI in water

Firstly, it's important to comprehend the chemical reaction of Silver iodide (AgI) in water. AgI partially dissolves in water producing silver ions (Ag+) and iodide ions (I-). The equation is represented as: \[ AgI_{(s)} \rightleftharpoons Ag^+_{(aq)} + I^-_{(aq)} \] This reaction is an equilibrium process- meaning it can go both forward and backward.

Understand the effect of KI on the solubility of AgI

Adding Potassium iodide (KI) to the solution increases the concentration of I- ions. According to Le Chatelier's principle, the equilibrium will shift to the left to reduce the effect of increased I- ions, thus reducing the solubility of AgI. Hence, presence of KI decreases the solubility of AgI in water.

Understand the effect of AgNO3 on the solubility of AgI

Adding Silver nitrate (AgNO3) to the solution increases the concentration of Ag+ ions. According to Le Chatelier's principle, the equilibrium will shift to the left to reduce the effect of increased Ag+ ions, thus reducing the solubility of AgI. Therefore, presence of AgNO3 decreases the solubility of AgI in water.

Key concepts

Le Chatelier's Principle

When a chemical system at equilibrium experiences a change in concentration, temperature, or pressure, Le Chatelier's principle provides a prediction on how the system will respond to restore a new equilibrium. Essentially, when you disturb a system in equilibrium, it will shift in that direction which tends to minimize the disturbance. This principle not only applies to gases but also helps explain the behavior of dissolving ionic compounds in water.

For example, if we increase the concentration of a reactant, the system will adjust by creating more products. On the other hand, if we increase the product concentration, the system will try to produce more reactants. This shifting is what maintains the delicate balance in chemical reactions and is crucial in understanding solubility equilibria, such as the dissolution of silver iodide (AgI) in water.

Silver Iodide Reaction

The silver iodide reaction in water is a classic example of a solubility equilibrium. Silver iodide is poorly soluble in water, which means only a small amount can dissolve at a given temperature. In chemical terms, the reaction can be summarized as:\[\begin{equation}AgI_{(s)} \rightleftharpoons Ag^+_{(aq)} + I^-_{(aq)}\end{equation}\]This reversible process indicates that silver iodide (AgI_{(s)}) can dissolve by separating into silver (Ag^+) and iodide (I^-) ions, and these ions can also come back together to form solid silver iodide. The forward process defines dissolving, and the reverse denotes precipitation. Understanding this reversible reaction is fundamental in grasping how adding different salts affects the solubility of silver iodide.

Ionic Solubility

Ionic solubility is the extent to which an ionic compound can dissolve in a solvent, typically water. It is greatly influenced by changes in the system’s conditions, such as the presence of other ions. In our example, we examine how adding potassium iodide (KI) or silver nitrate (AgNO_3) affects the solubility of silver iodide (AgI). These additives contribute extra ions that interact with the ions coming from AgI, guiding the equilibrium back and forth according to Le Chatelier's principle.

The process is governed by the solubility product constant (K_{sp}), a unique value for each compound at a certain temperature, which determines just how much of it can dissolve in water. This concept is crucial in predicting the behavior of ions in solution and their interactions with different substances.

Chemical Equilibrium

Chemical equilibrium pertains to the state in a reaction where the rates of the forward and backward reactions are equal, hence there is no net change in the concentrations of reactants and products over time. It’s a dynamic balance where the chemical species are continuously reacting, yet remain at a constant ratio. When a reaction reaches equilibrium, it doesn’t imply that the reactants and products are equal in concentration but that their ratios remain stable.

In the context of our silver iodide example, this equilibrium can be thrown out of balance by adding salts that provide ions already present in the equilibrium system. Such changes will cause the reaction to shift its position to re-establish equilibrium, reducing the solubility of the silver iodide in water. The ability to manipulate this balance by altering conditions is a powerful tool in controlling reactions and is essential for practical applications, such as water purification and pharmaceutical synthesis.