Aqueous Solutions & Solubility | Chemistry for EmSAT Achieve PDF Download

When certain substances dissolve in water, they undergo a change, either physical or chemical, resulting in the production of ions in the solution. These substances belong to a crucial category known as electrolytes. Conversely, substances that do not produce ions upon dissolution are termed nonelectrolytes. If the process generating ions is highly efficient, resulting in the complete conversion of the dissolved compound into ions, it is classified as a strong electrolyte. On the other hand, if only a small portion of the dissolved substance undergoes ion production, it is termed a weak electrolyte.
The classification of substances as strong, weak, or nonelectrolytes can be determined by assessing the electrical conductivity of an aqueous solution containing the substance. For a substance to conduct electricity, it must contain charged species that are freely mobile. The most familiar example is the conduction of electricity through metallic wires, where electrons serve as the mobile, charged entities. Similarly, solutions can conduct electricity if they contain dissolved ions, and the conductivity increases with higher ion concentration. By applying a voltage to electrodes immersed in a solution, one can evaluate the relative concentration of dissolved ions either quantitatively, by measuring electrical current flow, or qualitatively, by observing the brightness of a light bulb included in the circuit.

Ionic Electrolytes

Water and other polar molecules exhibit an attraction towards ions, as illustrated in Figure. This electrostatic attraction between an ion and a molecule possessing a dipole is referred to as an ion-dipole attraction. Such interactions hold significance in the process of dissolving ionic compounds in water.

As potassium chloride (KCl) dissolves in water, the ions are hydrated. The polar water molecules are attracted by the charges on the K⁺ and Cl⁻ ions. Water molecules in front of and behind the ions are not shown. The diagram shows eight purple spheres labeled K superscript plus and eight green spheres labeled C l superscript minus mixed and touching near the center of the diagram. Outside of this cluster of spheres are seventeen clusters of three spheres, which include one red and two white spheres. A red sphere in one of these clusters is labeled O. A white sphere is labeled H. Two of the green C l superscript minus spheres are surrounded by three of the red and white clusters, with the red spheres closer to the green spheres than the white spheres. One of the K superscript plus purple spheres is surrounded by four of the red and white clusters. The white spheres of these clusters are closest to the purple spheres. 

• When ionic compounds dissolve in water, the solid's ions separate and disperse evenly throughout the solution because water molecules surround and solvate them, thereby diminishing the intense electrostatic forces between them. This process, known as dissociation, represents a physical alteration. Typically, ionic compounds dissociate almost entirely upon dissolution, thus classifying them as strong electrolytes.
• Consider what occurs on a microscopic level when solid KCl is added to water. Ion-dipole forces draw the positive (hydrogen) ends of polar water molecules towards the negative chloride ions on the solid's surface, while they attract the negative (oxygen) ends towards the positive potassium ions. Consequently, water molecules infiltrate between individual K⁺ and Cl⁻ ions, enveloping them and mitigating the robust interionic forces that hold the ions together, allowing them to disperse into solution as solvated ions, as depicted in Figure. 
• This reduction in electrostatic attraction facilitates the independent movement of each hydrated ion in a diluted solution, leading to an increase in the system's disorder. This transition occurs as ions shift from their fixed and ordered positions in the crystal to mobile and considerably more disordered states in solution. Such escalated disorder is accountable for the dissolution of numerous ionic compounds, including KCl, which dissolve with the absorption of heat.
• In contrast, in some cases, the electrostatic attractions between ions in a crystal are so formidable, or the ion-dipole attractive forces between the ions and water molecules are so feeble, that the rise in disorder cannot offset the energy required to separate the ions, rendering the crystal insoluble. This circumstance applies to compounds such as calcium carbonate (limestone), calcium phosphate (the inorganic component of bone), and iron oxide (rust).

Solubility Rules

Certain combinations of aqueous reactants lead to the formation of a solid precipitate as a product, while others do not yield such a product. For instance, when solutions of sodium nitrate and ammonium chloride are mixed, no reaction takes place. Although one could theoretically depict a molecular equation indicating a double-replacement reaction, both resulting products, sodium chloride and ammonium nitrate, are soluble and would remain in the solution as ions. Consequently, every ion acts as a spectator ion, rendering the absence of a net ionic equation. It proves beneficial to predict when a precipitate will form in a reaction. This prediction can be facilitated through a set of guidelines known as solubility rules, as outlined in Tables .

Solubility Rules for Soluble Substances

Solubility Rules for Sparingly Soluble Substances

As an example on how to use the solubility rules, predict if a precipitate will form when solutions of cesium bromide and lead (II) nitrate are mixed.

The potential precipitates from a double-replacement reaction are cesium nitrate and lead (II) bromide. According to the solubility rules table, cesium nitrate is soluble because all compounds containing the nitrate ion, as well as all compounds containing the alkali metal ions, are soluble. Most compounds containing the bromide ion are soluble, but lead (II) is an exception. Therefore, the cesium and nitrate ions are spectator ions and the lead (II) bromide is a precipitate. The balanced net ionic reaction is:
Pb²⁺(aq) + 2Br⁻ (aq) → PbBr₂(s)