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Factors Affecting Solubility

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The solubility of an ionic compound can vary depending on other solutes present in the solution. The solute may be an ion common to the compound, or an acid or base. Thus, two major factors affecting the solubility are the common ion effect and the pH of a solution.

If lead(II) chloride is added to a solution of sodium chloride, both salts will ionize in the water, producing sodium and lead cations and the common anion, chloride.

Because the chloride ions in solution come from both the complete ionization of sodium chloride and the partial ionization of lead(II) chloride, their concentration would be higher than that of either the sodium ions or the lead ions.

To compensate, the equilibrium between the solid lead(II) chloride and its ions in solution shifts towards the undissociated salt, causing more lead(II) chloride to remain undissolved.

Thus, the presence of a common ion decreases the solubility of a sparingly soluble substance.

For instance, the molar solubility, x, of lead(II) chloride in a 0.100 molar sodium chloride solution can be calculated from the ICE table.

The initial concentration of lead(II) ions in the solution is zero, while chloride is 0.100 molar.

Each molecule of lead(II) chloride will dissociate into one lead and two chloride ions. So, the change in concentration of lead ions will be +x and that of chloride will be +2x.

The equilibrium concentration of lead ions will be x, but for the chloride ions, it will be the sum of 2x and 0.100.

At 25 °C, the value of K_sp for lead(II) chloride is 1.17 × 10⁻⁵, and the equilibrium expression is the product of the equilibrium concentrations of lead and chloride ions, which is x(2x + 0.100)².

Because the K_sp is small, 2x is much smaller than 0.100 molar and their sum can be approximated to 0.100 molar.

Substituting into the expression yields 1.17 × 10⁻⁵= x(0.100)².

Solving for x, the molar solubility of lead(II) chloride in 0.100 molar sodium chloride solution is 1.17 × 10⁻³ molar.

Compared to the molar solubility of lead(II) chloride in water, which is 1.43 × 10⁻² molar, the common ion decreases the solubility of lead(II) chloride 12 fold.

Solubility of a compound can also be affected by the pH of the solution.

Consider the partial dissociation of calcium hydroxide into calcium ions and hydroxide. If the pH is increased, by adding potassium hydroxide for example, the common hydroxide ion will lower the solubility of calcium hydroxide.

Conversely, if the pH is decreased, for example by adding hydrochloric acid, the protons will combine with hydroxide ions, lowering the concentration of hydroxide ions in solution.

The solubility equilibrium of calcium hydroxide would then shift towards the products, leading to an increased solubility in an acidic solution.

Overview

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Compared with pure water, the solubility of an ionic compound is less in aqueous solutions containing a common ion (one also produced by dissolution of the ionic compound). This is an example of a phenomenon known as the common ion effect, which is a consequence of the law of mass action that may be explained using Le Chȃtelier’s principle. Consider the dissolution of silver iodide:

Chemical structure of serotonin, neurotransmitter molecule, diagram for education and research.

This solubility equilibrium may be shifted left by the addition of either silver or iodide ions, resulting in the precipitation of AgI and lowered concentrations of dissolved Ag⁺ and I^–. In solutions that already contain either of these ions, less AgI may be dissolved than in solutions without these ions.

This effect may also be explained in terms of mass action as represented in the solubility product expression:

Chemical structure of benzene; hexagonal ring, alternating double bonds; organic chemistry diagram.

The mathematical product of silver and iodide ion molarities is constant in an equilibrium mixture regardless of the source of the ions, and so an increase in one ion’s concentration must be balanced by a proportional decrease in the other.

The Role of Precipitation in Wastewater Treatment

Solubility equilibria are useful tools in the treatment of wastewater carried out in facilities that may treat the municipal water in a city or town. Specifically, selective precipitation is used to remove contaminants from wastewater before it is released back into natural bodies of water. For example, phosphate ions (PO₄³⁻) are often present in the water discharged from manufacturing facilities. An abundance of phosphate causes excess algae to grow, which impacts the amount of oxygen available for marine life as well as making the water unsuitable for human consumption.

One common way to remove phosphates from water is by the addition of calcium hydroxide, or lime, Ca(OH)₂. As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca₅(PO₄)₃·OH, which then precipitates out of the solution:

Chemical structure diagram of acetaminophen molecule, illustrating aromatic ring and functional groups.

Because the amount of calcium ion added does not result in exceeding the solubility products for other calcium salts, the anions of those salts remain behind in the wastewater. The precipitate is then removed by filtration, and the water is brought back to a neutral pH by the addition of CO₂ in a recarbonation process. Other chemicals can also be used for the removal of phosphates by precipitation, including iron(III) chloride and aluminum sulfate.

This text is adapted from Openstax, Chemistry 2e, Section 15.1: Precipitation and Dissolution.

Transcript

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$$\rightleftharpoonup{xx}$$ $$\longleftharp{xx}$$, $$\longrightharp{xx}$$,

The solubility of an ionic compound can vary depending on other solutes present in the solution. The solute may be an ion common to the compound, or an acid or base. Thus, two major factors affecting the solubility are the common ion effect and the pH of a solution.

If lead(II) chloride is added to a solution of sodium chloride, both salts will ionize in the water, producing sodium and lead cations and the common anion, chloride.

Because the chloride ions in solution come from both the complete ionization of sodium chloride and the partial ionization of lead(II) chloride, their concentration would be higher than that of either the sodium ions or the lead ions.

To compensate, the equilibrium between the solid lead(II) chloride and its ions in solution shifts towards the undissociated salt, causing more lead(II) chloride to remain undissolved.

Thus, the presence of a common ion decreases the solubility of a sparingly soluble substance.

For instance, the molar solubility, x, of lead(II) chloride in a 0.100 molar sodium chloride solution can be calculated from the ICE table.

The initial concentration of lead(II) ions in the solution is zero, while chloride is 0.100 molar.

Each molecule of lead(II) chloride will dissociate into one lead and two chloride ions. So, the change in concentration of lead ions will be +x and that of chloride will be +2x.

The equilibrium concentration of lead ions will be x, but for the chloride ions, it will be the sum of 2x and 0.100.

At 25 °C, the value of K_sp for lead(II) chloride is 1.17 × 10⁻⁵, and the equilibrium expression is the product of the equilibrium concentrations of lead and chloride ions, which is x(2x + 0.100)².

Because the K_sp is small, 2x is much smaller than 0.100 molar and their sum can be approximated to 0.100 molar.

Substituting into the expression yields 1.17 × 10⁻⁵= x(0.100)².

Solving for x, the molar solubility of lead(II) chloride in 0.100 molar sodium chloride solution is 1.17 × 10⁻³ molar.

Compared to the molar solubility of lead(II) chloride in water, which is 1.43 × 10⁻² molar, the common ion decreases the solubility of lead(II) chloride 12 fold.

Solubility of a compound can also be affected by the pH of the solution.

Consider the partial dissociation of calcium hydroxide into calcium ions and hydroxide. If the pH is increased, by adding potassium hydroxide for example, the common hydroxide ion will lower the solubility of calcium hydroxide.

Conversely, if the pH is decreased, for example by adding hydrochloric acid, the protons will combine with hydroxide ions, lowering the concentration of hydroxide ions in solution.

The solubility equilibrium of calcium hydroxide would then shift towards the products, leading to an increased solubility in an acidic solution.

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Common Ion Effect

Solubility Product

Lead II Chloride

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Precipitation Reaction

Wastewater Treatment

Selective Precipitation

Phosphate Removal

Hydroxylapatite Formation