Question

A solution contains three anions with the following concentrations: \(0.20 M \mathrm{CrO}_{4}^{2-}, 0.10 M \mathrm{CO}_{3}^{2-}\) , a n d 0.010\(M \mathrm{Cl}^{-} .\) If a dilute AgNO \(_{3}\) solution is slowly added to the solution, what is the first compound to precipitate: \(\mathrm{Ag}_{2} \mathrm{CrO}_{4}\left(K_{s p}=1.2 \times 10^{-12}\right), \mathrm{Ag}_{2} \mathrm{CO}_{3}\left(K_{s p}=8.1 \times 10^{-12}\right)\) or \(\mathrm{AgCl}\left(K_{s p}=1.8 \times 10^{-10}\right) ?\)

Short answer

When a dilute AgNO₃ solution is slowly added to the given solution, the first compound to precipitate is AgCl, as it requires the smallest amount of Ag⁺ (1.8 x 10⁻⁸ M) to surpass its solubility product constant (Kₛₚ).

Step-by-step solution

Write balanced net ionic equations for each compound

For the formation of Ag₂CrO₄, Ag₂CO₃, and AgCl, the balanced net ionic equations are: 1. Ag₂CrO₄(s) ↔ 2Ag⁺(aq) + CrO₄²⁻(aq) 2. Ag₂CO₃(s) ↔ 2Ag⁺(aq) + CO₃²⁻(aq) 3. AgCl(s) ↔ Ag⁺(aq) + Cl⁻(aq)

Calculate the Reaction Quotient (Q) for each compound

Let 'x' be the concentration of Ag⁺ in the solution. The Q expression and Q for each compound are as follows: 1. For Ag₂CrO₄ \(Q = [Ag⁺]^2[CrO₄²⁻]\) Substituting the concentration values, we get: \(Q= (2x)^2(0.20)\) 2. For Ag₂CO₃ \(Q = [Ag⁺]^2[CO₃²⁻]\) Substituting the concentration values, we get: \(Q= (2x)^2(0.10)\) 3. For AgCl \(Q = [Ag⁺][Cl⁻]\) Substituting the concentration values, we get: \(Q= x(0.010)\)

Compare Q values with K_sp values

We want to find the smallest concentration of Ag⁺(x) necessary to cause precipitation in each case. For this, we compare the Q expressions with K_sp values: 1. For Ag₂CrO₄, \[(2x)^2(0.20) = 1.2 \times 10^{-12}\] 2. For Ag₂CO₃, \[(2x)^2(0.10) = 8.1 \times 10^{-12}\] 3. For AgCl, \[x(0.010) = 1.8 \times 10^{-10}\] Now, let's find the minimum concentration of Ag⁺(x) necessary for precipitation: 1. For Ag₂CrO₄, \[x^2 = \frac{1.2 \times 10^{-12}}{0.40}\] => \(x = 1.7 \times 10^{-6}\) 2. For Ag₂CO₃, \[x^2 = \frac{8.1 \times 10^{-12}}{0.20}\] => \(x = 4.0 \times 10^{-6}\) 3. For AgCl, \[x = \frac{1.8 \times 10^{-10}}{0.010}\] => \(x = 1.8 \times 10^{-8}\) Comparing the Ag⁺ concentrations, the smallest amount of Ag⁺ is 1.8 x 10⁻⁸ M which can cause AgCl to precipitate first.

Conclusion

When a dilute AgNO₃ solution is slowly added to the given solution, the first compound to precipitate is AgCl.

Key concepts

Solubility Product Constant

Understanding the Solubility Product Constant (K_sp) is pivotal when studying precipitation reactions in chemistry. It is a special type of equilibrium constant that applies to the dissolution of sparingly soluble salts. These salts, when dissolved in water, reach a point where the ions in solution are in a dynamic equilibrium with the solid undissolved salt.

The K_sp value gives us an idea of the extent to which a compound will dissolve in water. A lower K_sp indicates that less of the compound dissolves to form ions, making it less soluble. Using the example from the exercise, where different compounds of silver were considered, the K_sp values provided rank the solubility of these salts. A compound with a lower K_sp is less soluble and is likely to precipitate first when compared with other salts in the solution, assuming equal ionic concentrations.

Reaction Quotient

The Reaction Quotient (Q) helps determine the direction in which a reaction will proceed to reach equilibrium. It compares the current concentrations of the reactants and products in a reaction to the equilibrium concentrations as expressed by the equilibrium constant (K).

Mathematically, Q is expressed in the same way as the K_sp but with current concentrations instead of equilibrium concentrations. When Q is less than K_sp, the system will favor the formation of products (dissolving of solids, in the case of solubility equilibria). When Q equals K_sp, the system is at equilibrium, and when Q is greater than K_sp, the system will favor the formation of reactants (precipitation). In the exercise, by calculating Q for each potential precipitate and comparing it to the respective K_sp, we predict which salt will precipitate first.

Net Ionic Equations

Net ionic equations are a chemist’s tool to succinctly represent only the entities involved in a chemical change during a reaction, omitting the spectator ions. These equations show the actual chemical species that are participating in the reaction to form the precipitate.

Writing net ionic equations involves identifying the ionic species that change physical states or combine to form a precipitate in a solution. In the provided exercise, the net ionic equation for each silver compound clearly shows the ions that lead to the formation of each respective solid precipitate. This selective representation is crucial for understanding the stoichiometry and the formation of compounds during precipitation reactions.

Equilibrium Expressions

Equilibrium expressions are mathematical representations of the state of equilibrium for reactions. For precipitation reactions, the equilibrium expression is defined by the K_sp, which reflects the concentrations of the ionic species once the reaction has reached a stable state, where the rate of dissolution equals the rate of precipitation.

In the context of our exercise, the K_sp expressions for Ag₂CrO₄, Ag₂CO₃, and AgCl were derived from their respective balanced net ionic equations. By setting up these expressions and manipulating them, chemists can calculate the specific concentration of ions needed for the initiation of precipitation. With these capacities, scientists can predict and control aspects of reactions, which is invaluable in areas such as material science, pharmaceuticals, and environmental engineering.