Editing Solubility (section)

===Temperature===
The solubility of a given solute in a given solvent is function of temperature. Depending on the change in [[enthalpy]] (Δ''H'') of the dissolution reaction, ''i.e.'', on the [[Endothermic process|endothermic]] (Δ''H''&nbsp;>&nbsp;0) or [[Exothermic process|exothermic]] (Δ''H''&nbsp;<&nbsp;0) character of the dissolution reaction, the solubility of a given compound may increase or decrease with temperature.  The [[van 't Hoff equation]] relates the change of solubility [[equilibrium constant]] (''K''<sub>sp</sub>) to temperature change and to reaction [[enthalpy]] change. 

*For most [[solid]]s and liquids, their solubility increases with temperature because their dissolution reaction is endothermic (Δ''H''&nbsp;>&nbsp;0).<ref name = hill>John W. Hill, Ralph H. Petrucci, ''General Chemistry'', 2nd edition, Prentice Hall, 1999.</ref> In liquid water at high temperatures, (e.g. that approaching the [[critical temperature]]), the solubility of ionic solutes tends to decrease due to the change of properties and structure of liquid water; the lower [[dielectric constant]] results in a less [[polar solvent]] and in a change of hydration energy affecting the Δ''G'' of the dissolution reaction.

*[[Gas]]eous solutes exhibit more complex behavior with temperature. As the temperature is raised, gases usually become less soluble in water (exothermic dissolution reaction related to their hydration) (to a minimum, which is below 120&nbsp;°C for most permanent gases<ref>{{cite book|editor=P. Cohen|title=The ASME Handbook on Water Technology for Thermal Power Systems|publisher=The American Society of Mechanical Engineers|year=1989| page =442}}</ref>), but more soluble in organic solvents (endothermic dissolution reaction related to their solvation).<ref name=hill/>

The chart shows solubility curves for some typical solid inorganic [[salt (chemistry)|salts]] in liquid water (temperature is in degrees [[Celsius]], i.e. [[kelvin]]s minus 273.15).<ref>{{cite book|title=Handbook of Chemistry and Physics| edition= 27th|location= Cleveland, Ohio|year=1943 |publisher= Chemical Rubber Publishing Co.}}</ref> Many salts behave like [[barium nitrate]] and [[disodium hydrogen arsenate]], and show a large increase in solubility with temperature (Δ''H''&nbsp;>&nbsp;0). Some solutes (e.g. [[sodium chloride]] in water) exhibit solubility that is fairly independent of temperature (Δ''H''&nbsp;≈&nbsp;0). A few, such as [[calcium sulfate]] ([[gypsum]]) and [[cerium(III) sulfate]], become less soluble in water as temperature increases (Δ''H''&nbsp;<&nbsp;0).<ref name="Scientific American">{{cite web|title=What substances, such as cerium sulfate, have a lower solubility when they are heated?|website=[[Scientific American]] |url=http://www.scientificamerican.com/article/what-substances-such-as-c/|access-date=28 May 2014}}</ref> This is also the case for [[calcium hydroxide]] ([[portlandite]]), whose solubility at 70&nbsp;°C is about half of its value at 25&nbsp;°C. The dissolution of calcium hydroxide in water is also an exothermic process (Δ''H''&nbsp;<&nbsp;0). As dictated by the [[van 't Hoff equation]] and [[Le Chatelier's principle]], low temperatures favor dissolution of Ca(OH)<sub>2</sub>. Portlandite solubility increases at low temperature. This temperature dependence is sometimes referred to as "retrograde" or "inverse" solubility.{{cn|date=June 2024}} Occasionally, a more complex pattern is observed, as with [[sodium sulfate]], where the less soluble deca[[hydrate]] crystal ([[mirabilite]]) loses [[water of crystallization]] at 32&nbsp;°C to form a more soluble [[anhydrous]] phase ([[thenardite]]) with a smaller change in [[Gibbs free energy]] (Δ''G'') in the dissolution reaction.{{Citation needed|date=July 2008}}

[[File:Temperature dependence solublity of solid in liquid water high temperature.svg|right|400px]]

The solubility of [[organic compounds]] nearly always increases with temperature. The technique of [[Recrystallization (chemistry)|recrystallization]], used for purification of solids, depends on a solute's different solubilities in hot and cold solvent. A few exceptions exist, such as certain [[cyclodextrin]]s.<ref>{{cite journal|title=A highly water-soluble 2+1 b-cyclodextrin–fullerene conjugate|author=Salvatore Filippone, Frank Heimanna and André Rassat|journal=[[Chem. Commun.]]|volume=2002|pages=1508–1509|doi=10.1039/b202410a|year=2002|issue=14|pmid=12189867}}</ref>