Sparingly Soluble Salts
Sparingly Soluble Salts - Related Videos
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Factors Affecting Solubility
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2020
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: This solubility equilibrium may be shifted left by the addition of either silver or iodide ions, resulting in...
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Solubility - Student Protocol
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2020
Source: Smaa Koraym at Johns Hopkins University, MD, USA Creating a Saturated SolutionExpand In this experiment, you will create a saturated solution of sodium tetraborate decahydrate, also called borax. In water, borax dissociates into two sodium cations and one tetraborate anion. When an aqueous borax solution is saturated, it means that it contains the maximum amount of dissolved solute, which is borax, for that volume of solvent, which is water. Any additional solid won't appear to...
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Solubility - Concepts
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2020
Solubility Solubility describes how much of a solute can dissolve in a given volume of a specific solvent. Solubility is usually reported in terms of solute mass per solvent volume or solute mass per solvent mass. For example, the solubility of sodium chloride in water at room temperature is reported as 36 g per 100 mL of water. If solubility is reported in solute mass per solvent mass, the solvent mass will need to be converted to volume for further calculations. Solubility changes with...
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Solubility of Ionic Compounds
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2020
Solubility is the measure of the maximum amount of solute that can be dissolved in a given quantity of solvent at a given temperature and pressure. Solubility is usually measured in molarity (M) or moles per liter (mol/L). A compound is termed soluble if it dissolves in water. When soluble salts dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution; this process represents a physical change known as dissociation. Potassium chloride (KCl) is an example...
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Solubility Equilibria
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2020
Solubility equilibria are established when the dissolution and precipitation of a solute species occur at equal rates. These equilibria underlie many natural and technological processes, ranging from tooth decay to water purification. An understanding of the factors affecting compound solubility is, therefore, essential to the effective management of these processes. This section applies previously introduced equilibrium concepts and tools to systems involving dissolution and precipitation. The...
FAQs
Ksp serves as the quantitative link between a sparingly soluble salt’s solid phase and its dissolved ions. At equilibrium, the small amount that enters solution establishes ion concentrations consistent with this constant. Comparing conditions through Ksp helps chemists determine whether a solution can remain in equilibrium with the solid or whether precipitation is likely, making the constant central to prediction and calculation.
A common ion changes the dissolution balance because it adds one of the ions already associated with the salt to solution. The equilibrium then shifts toward conditions that limit further dissolution, so the salt’s dissolved amount can decrease. This effect is useful when controlling precipitation or improving the selectivity of a chemical separation involving more than one ionic compound.
pH and complex formation influence solubility by changing the chemical environment of the dissolved ions. A pH change can alter the dissolution equilibrium, while complex formation changes how an ion exists in solution. Consequently, the dissolved-ion concentrations and the tendency of the solid to remain or form can differ from conditions without these effects.
To analyze a system, identify the sparingly soluble solid, consider the ions produced when a small amount dissolves, and use the relevant Ksp relationship to evaluate their equilibrium concentrations. The calculation should also account for common ions, pH, or complex formation when those conditions are present. This approach connects specified solution conditions with predicted ion levels.
In a precipitation test, chemists assess whether solution conditions favor formation of a solid. Ion concentrations are evaluated in relation to the salt’s Ksp, while common ions and other equilibrium shifts may alter the result. Because different salts respond differently, controlled precipitation can help separate substances in qualitative and analytical chemistry, where selective solid formation provides useful chemical information.
These equilibria provide a chemical basis for studying mineral formation, water treatment, and environmental contamination. In natural systems, they help describe how poorly soluble compounds behave as conditions change; in laboratory systems, they support prediction of precipitation and dissolved-ion concentrations. The same framework therefore connects aqueous chemistry with practical separation and analysis of contaminated or treated water.