Solubility is defined as the maximum amount of a substance, the solute, that can completely dissolve into another substance, the solvent, under specific conditions. This measure is typically expressed as the concentration of the solute within a saturated solution, which is the point where the solution is in equilibrium and can dissolve no more material. Temperature represents the most significant external factor influencing this equilibrium. The relationship between temperature and solubility is not uniform, as the effect depends heavily on the physical state of the solute. Understanding this requires separating the behavior of solid solutes from that of gaseous solutes.
Temperature’s Effect on Solid Solutes
For the majority of solid solutes dissolved in a liquid solvent, such as water, solubility tends to increase as the temperature of the system rises. This is rooted in the increased kinetic energy supplied by the heat, which affects both the solute and the solvent molecules. The energetic solvent molecules can more effectively bombard and overcome the attractive forces that hold the solid solute’s crystal lattice structure together. Once the solute particles are freed, the higher temperature helps the solvent molecules surround and stabilize the separated ions or molecules, a process known as solvation. Potassium Nitrate (\(\text{KNO}_3\)) is an example of this trend, being only moderately soluble in cold water but highly soluble as the temperature increases. At \(20^\circ\text{C}\), approximately \(316\) grams of \(\text{KNO}_3\) can dissolve in a liter of water, but this capacity soars to \(2,439\) grams per liter when the water is heated to \(100^\circ\text{C}\). This increase illustrates how temperature changes the dissolving power of a solvent for a solid solute.
Temperature’s Effect on Gaseous Solutes
The solubility of gases dissolved in liquids demonstrates an inverse relationship with temperature, meaning solubility decreases as the temperature rises. When a gas dissolves in a liquid, the gas molecules become trapped by the attractive forces of the solvent, a process that typically releases energy. Increasing the liquid’s temperature introduces more kinetic energy into the system, particularly to the dissolved gas molecules. The increased energy causes the gas molecules to move faster, allowing them to gain enough energy to break free from the solvent’s attractive forces and escape back into the surrounding atmosphere. This phenomenon is easily observed when a carbonated beverage goes flat much faster at room temperature compared to when it is refrigerated. This principle also explains how thermal pollution, where water bodies are warmed by industrial discharge, reduces the amount of dissolved oxygen available for aquatic life.
Understanding the Energy Behind Dissolution
The opposing trends for solids and gases are explained by considering the energy change during dissolution, categorized as either endothermic or exothermic. Endothermic dissolution requires the absorption of heat energy from the surroundings to proceed, meaning the solution process itself feels cold. Since heat is required for the reaction, increasing the temperature drives the dissolution forward, resulting in higher solubility. This behavior, which describes the dissolution of most ionic solids like potassium nitrate, is governed by Le Chatelier’s Principle. This principle states that a system at equilibrium will shift to counteract any imposed stress, meaning that adding heat to an endothermic system causes the equilibrium to shift toward the dissolved state. Conversely, exothermic dissolution releases heat energy into the surroundings, causing the solution to feel warmer. For an exothermic process, heat is a product of the reaction, and increasing the temperature stresses the system. In response, the equilibrium shifts toward the undissolved state, thereby decreasing solubility as the temperature increases. While the dissolution of gases is nearly always exothermic, some solid salts, such as calcium chloride, also exhibit exothermic dissolution, leading to a decrease in their solubility at higher temperatures.
Visualizing Solubility and Supersaturation
The relationship between solubility and temperature is visualized using a solubility curve, which plots the concentration of a solute versus temperature. The line on the curve represents the saturation point, defining the maximum amount of solute the solvent can hold at equilibrium for any given temperature. Any point below this line indicates an unsaturated solution, while a solution exactly on the line is saturated. The concept of supersaturation refers to a temporary, unstable condition where a solution contains more dissolved solute than its normal saturation limit. This state is typically achieved by creating a hot, saturated solution and then carefully cooling it without agitation, trapping the excess solute in the dissolved state. This fragile state is easily disrupted; the introduction of a seed crystal or a slight disturbance will cause the excess solute to rapidly precipitate out of the solution.