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Chapter 17: Problem 53

Explain how the dissolving of an ionic solute in water represents an equilibrium process.

Short Answer

Expert verified
The dissolving of an ionic solute in water represents an equilibrium process because it is a reversible process where the forward (dissolution) and reverse (precipitation) reactions occur at the same rate, resulting in no net change in the concentrations of the reactants and products. The solute dissolves due to electrostatic interactions between its ions and polar water molecules, and it can also re-form through precipitation. When the dissolution and precipitation rates become equal, the system reaches a dynamic equilibrium where the concentrations of dissolved ions and undissolved solute remain constant over time.

Step by step solution

01

Define equilibrium process

An equilibrium process is a reversible process in which the forward and reverse reactions occur at the same rate, resulting in no net change in the concentrations of the reactants and products. In other words, at equilibrium, the system is dynamically balanced, and the concentrations of the substances involved remain constant over time.
02

Describe the process of dissolving an ionic solute

When an ionic solute is added to water, it dissolves due to the electrostatic interactions between the positive and negative ions of the solute and the polar water molecules. The solute's positive ions interact with the oxygen atoms of water (with partial negative charge), while the solute's negative ions interact with the hydrogen atoms of water (with partial positive charge). This process is known as solvation or hydration of ions.
03

Explain the reversibility of the process

The dissolving process is reversible because the hydrated ions in solution can collide with one another and re-form the ionic solute. This re-formation is called precipitation. The ions continuously dissolve and precipitate in the solution, depending on their solubility.
04

Explain the dynamic equilibrium in the dissolving process

As the solute dissolves in the water, its solubility and precipitation rates will eventually become equal. When this happens, the concentrations of the dissolved ions and the undissolved solute remain constant, even though the dissolution and precipitation processes are still occurring. This balance between the dissolution and precipitation rates is called dynamic equilibrium.
05

Relate the equilibrium process of dissolving an ionic solute to the overall concept

In summary, the dissolving of an ionic solute in water represents an equilibrium process because it is reversible and reaches a dynamic equilibrium. At equilibrium, the dissolution and precipitation rates become equal, and the concentrations of the dissolved ions and the undissolved solute remain constant over time. The concept of equilibrium is an essential aspect of many chemical processes, including solubility and precipitation reactions.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Ionic Solute Dissolution
When an ionic solute, such as table salt, is added to water, it undergoes a fascinating process called dissolution. This happens because of the interactions between the solute ions and the water molecules. Water is a polar molecule, meaning it has a partial positive charge on one side and a partial negative charge on the other. These polar regions are what make the water such an excellent solvent for ionic compounds.
  • The positive ions of the solute, like sodium (\( Na^+ \)), are attracted to the oxygen part (which has a negative charge) of water molecules.
  • The negative ions, like chloride (\( Cl^- \)), are drawn to the hydrogen part (which has a positive charge) of water molecules.
This interaction, where water molecules surround the ions, is called solvation or hydration. Solvation helps the ions separate and disperse, essentially pulling the ionic compound apart into its individual ions.
Reversible Processes
In chemistry, many processes are not just one-way streets; they can go both forward and backward. This is what we call reversible processes. The dissolving of ionic solutes in water is a perfect example. While ions move from the solid into solution as they dissolve, they can also come together to re-form the solid, which is known as precipitation.
  • The forward process involves dissolving the solute where the solid forms ions in the solution.
  • The reverse process involves precipitation where the ions return to their solid form.
This reversible nature ensures that the process can achieve equilibrium, where neither dissolution nor precipitation can dominate unless conditions such as temperature change.
Dynamic Equilibrium
Dynamic equilibrium represents a state where two opposing processes occur at equal rates, leading to no net change in system concentration. For an ionic solute like salt in water, dynamic equilibrium is achieved when the rate at which the salt dissolves equals the rate at which it precipitates. This means, on a macroscopic level, nothing changes.
  • The number of ions dissolving from the solid into the liquid state equals the number of ions recombining into the solid.
  • The concentration of ions in the solution and the amount of undissolved solute remain constant, even though ions continue to move between the dissolved and precipitated states.
Understanding dynamic equilibrium helps in predicting the behavior of solutions under different conditions.
Solubility and Precipitation
One of the fundamental properties in chemistry is solubility, which refers to how much of a substance can dissolve in a solvent at a given temperature and pressure. For ionic compounds, solubility dictates the extent of their dissolution in water. Precipitation, on the other hand, is the process of forming a solid from a solution.
  • If a solution contains more dissolved ions than it can normally hold at equilibrium, it is supersaturated, and additional solid will precipitate out.
  • Conversely, if the conditions change to allow more solute to dissolve, such as an increase in temperature, more of the solid will dissolve.
The balance between solubility and precipitation is crucial in many natural and industrial processes, such as forming stalactites and stalagmites in caves or purifying chemicals through recrystallization.

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Most popular questions from this chapter

\(K_{\text {sp }}\) for magnesium carbonate, \(\mathrm{MgCO}_{3}\), has a value \(3.5 \times 10^{-8}\) at 25 ' \(\mathrm{C}\). Calculate the solubility of magnesium carbonate in \(\mathrm{mol} / \mathrm{L}\) and \(\mathrm{g} / \mathrm{L}\) at \(25 \mathrm{C}\).

\(K_{\text {sp }}\) for copper(II) hydroxide, \(\mathrm{Cu}(\mathrm{OH})_{2},\) has a value \(2.2 \times 10^{-20}\) at 25 " \(\mathrm{C}\). Calculate the solubility of copper(II) hydroxide in \(\mathrm{mol} / \mathrm{L}\) and \(\mathrm{g} / \mathrm{L}\) at \(25 \mathrm{C}\)

Magnesium fluoride dissolves in water to the extent of \(8.0 \times 10^{-2} \mathrm{~g} / \mathrm{L}\) at \(25 \mathrm{C}\). Calculate the solubility of \(\mathrm{MgF}_{2}(s)\) in moles per liter, and calculate \(K_{\mathrm{sp}}\) for \(\mathrm{MgF}_{2}\) at 25 .

Under what circumstances can we compare the solubilities of two salts by directly comparing the values of their solubility products?

For a given reaction at a given temperature, the special ratio of products to reactants defined by the equilibrium constant is always equal to the same number. Explain why this is true, no matter what initial concentrations of reactants (or products) may have been taken in setting up an experiment.
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