Calculate the solubility of Mg(OH)2 in a solution buffered to a pH of 7.00. How does this compar

Calculate the solubility of Mg(OH)2 in a solution buffered...

Key Concepts

Dissolution Equilibrium

Dissolution equilibrium describes the reversible process by which a solid salt dissociates into its constituent ions in a solvent. Setting up and solving the equilibrium expression for this process, accounting for factors like the solubility product, pH, and the presence of common ions, is fundamental in predicting the salt’s solubility under various conditions, such as in a buffered versus an unbuffered solution.

Buffer Systems

A buffer solution maintains a relatively constant pH despite the addition of acids or bases. By stabilizing the pH, a buffer can control the concentration of ions involved in acid–base equilibria. In solubility problems, buffers can limit the change in hydroxide or hydrogen ion concentrations, leading to differences in solubility compared to unbuffered systems where such concentrations can vary freely.

Solubility Product (Ksp)

The solubility product is an equilibrium constant (Ksp) that quantifies the extent to which a sparingly soluble salt dissolves in water. It is defined as the product of the molar concentrations of the ions each raised to the power of their stoichiometric coefficients in the dissolution reaction. Understanding Ksp allows us to calculate the concentration of ions in solution at equilibrium and predict whether precipitation will occur under given conditions.

Common Ion Effect

The common ion effect refers to the shift in the equilibrium of a sparingly soluble salt when an ion common to its dissolution equilibrium is added to the solution. When a solution is buffered at a certain pH, the concentration of hydroxide ions may be fixed or influenced by the buffer, reducing the salt’s solubility due to Le Chatelier’s principle. This phenomenon also explains why the solubility in a buffered solution can differ from that in unbuffered water.

pH and Equilibrium

pH plays a crucial role in solubility equilibria, especially for salts that produce or consume hydroxide or hydrogen ions upon dissolution. Changes in pH can alter the ionization state of dissolved species and shift the equilibrium. In the case of salts that produce hydroxide ions, fixing the pH with a buffer can limit the solubility by reducing the concentration of the dissociation product, thereby affecting the overall equilibrium.

Transcript

00:01 Hello, today we're going to be looking at magnesium hydroxide and looking at its molar solubility in different situations.

00:10 So first, let's talk about the solubility of magnesium hydroxide in just water.

00:20 So just normal circumstances, we're just taking magnesium hydroxide, we're dissolving it in water.

00:26 So let's draw the equilibrium that's taking place.

00:34 Magnesium hydroxide is dissociating partially into magnesium hydroxide and two hydroxide ions.

00:47 So initially we have zero concentrations of both.

00:53 Then we're going to dissolve a certain amount of moles of magnesium hydroxide.

01:02 Which will create one, there will be on a one to one ratio.

01:07 So if we dissolve one mole of magnesium hydroxide, then we will get one mole of magnesium ions per volume.

01:17 So if we are dissolving the maximum that we can, that would just be the molar solubility.

01:24 But then notice for every one mole of magnesium hydroxide that's dissolved, we get two hydroxide ions.

01:31 So that means we're going to have to do it'll increase by twice the molar cellubility.

01:40 And so now let's draw our ksp expression.

01:46 So the ksp is equal to the concentration of magnesium times the concentration of hydroxide squared, because it has a two coefficient in front.

02:04 So our ksp for magnesium hydroxide is 1 .2 times 10 to the negative 11.

02:18 Now we plug in our molar cellulity for the concentrations.

02:26 We'll get 4 s cubed will be equal to 1 .2 times 10 to the negative 11...

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