Calculate the solubility of aluminum hydroxide, Al(OH), in a solution buffered at pH 11.00 .

Name: Problem 20, Chapter 15: Chemistry
Uploaded: 2019-09-20T12:29:57-08:00
Duration: 1 min 35 s
Channel: Eric Ferrara
Description: Calculate the solubility of aluminum hydroxide, Al(OH), in a solution buffered at pH 11.00 .

step 1 So here we need to start by finding the initial concentration of hydroxide, and we can get that

Calculate the solubility of aluminum hydroxide, Al(OH), in a...

Key Concepts

Common Ion Effect

The common ion effect refers to the decrease in the solubility of a sparingly soluble salt when one of its constituent ions is added to the solution from another source. This effect shifts the dissolution equilibrium according to Le Chatelier's principle, reducing the amount of the solid that can dissolve. It is especially important in problems where an external source of an ion, like a buffer providing hydroxide ions, is present.

Relationship Between pH and Hydroxide Ion Concentration

The pH of a solution is directly related to its pOH by the equation pH + pOH = 14. In systems where the pH is controlled, such as buffered solutions, the hydroxide ion concentration can be readily determined from the pOH. This concentration is a key factor in the solubility equilibria of hydroxide salts and thereby influences their calculated solubility.

Chemical Equilibrium

Chemical equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant concentrations of reactants and products. In solubility problems, establishing the equilibrium expression for the dissolution reaction of a compound allows one to link the concentrations of dissolved ions to the solubility product constant.

Solubility Product Constant (Ksp)

The solubility product constant is a specific type of equilibrium constant used for sparingly soluble ionic compounds. It represents the product of the molar concentrations of the ions, each raised to the power of their stoichiometric coefficients, at equilibrium. This value is critical in predicting whether a precipitate will form under given conditions and in calculating the solubility of the compound.

Buffer Solutions and pH

Buffers are solutions that resist changes in pH when small amounts of acids or bases are added. They maintain a nearly constant pH, which is crucial in solubility calculations because the pH can influence the speciation and concentration of ions, such as hydroxide ions. A buffered pH ensures that the hydroxide ion concentration remains fixed, thereby affecting the solubility equilibrium of hydroxide compounds.

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