Editing Solubility (section)

==Solubility of ionic compounds in water==
{{main|Solubility chart|Solubility table}}
Some ionic compounds ([[salts]]) dissolve in water, which arises because of the attraction between positive and negative charges (see: [[solvation]]). For example, the salt's positive ions (e.g. Ag<sup>+</sup>) attract the partially negative oxygen atom in {{chem2|H2O}}. Likewise, the salt's negative ions (e.g. Cl<sup>−</sup>) attract the partially positive hydrogens in {{chem2|H2O}}. Note: the oxygen atom is partially negative because it is more [[electronegativity|electronegative]] than hydrogen, and vice versa (see: [[chemical polarity]]).

:{{chem2|AgCl_{(s)} <–> Ag+_{(aq)} + Cl–_{(aq)}|}}

However, there is a limit to how much salt can be dissolved in a given volume of water. This concentration is the solubility and related to the [[solubility product]], ''K''<sub>sp</sub>. This equilibrium constant depends on the type of salt ({{chem2|AgCl}} vs. {{chem2|NaCl}}, for example), temperature, and the common ion effect.

One can calculate the amount of {{chem2|AgCl}} that will dissolve in 1 liter of pure water as follows:

:''K''<sub>sp</sub> = [Ag<sup>+</sup>] × [Cl<sup>−</sup>] / M<sup>2</sup> (definition of solubility product; M = mol/L)
:''K''<sub>sp</sub> = 1.8 × 10<sup>−10</sup> (from a table of solubility products)
[Ag<sup>+</sup>] = [Cl<sup>−</sup>], in the absence of other silver or chloride salts, so
:[Ag<sup>+</sup>]<sup>2</sup> = 1.8 × 10<sup>−10</sup> M<sup>2</sup>
:[Ag<sup>+</sup>] = 1.34 × 10<sup>−5</sup> mol/L

The result: 1 liter of water can dissolve 1.34 × 10<sup>−5</sup> [[mole (unit)|moles]] of {{chem2|AgCl|}} at room temperature. Compared with other salts, {{chem2|AgCl}} is poorly soluble in water. For instance, table salt ({{chem2|NaCl}}) has a much higher ''K''<sub>sp</sub> = 36 and is, therefore, more soluble. The following table gives an overview of solubility rules for various ionic compounds.

{| class="wikitable" style="margin:0.5em auto"
|-
! style="width:50%" | Soluble
! style="width:50%" | Insoluble<ref>{{cite book|editor=C. Houk |editor2=R. Post|title=Chemistry, Concept and Problems|publisher=John Wiley & Sons|year=1997|page=[https://archive.org/details/chemistryconcept00houk/page/121 121]|isbn=978-0-471-12120-6|url=https://archive.org/details/chemistryconcept00houk|url-access=registration}}</ref>
|-
| [[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds (except [[lithium phosphate]])||[[Carbonate]]s (except [[Alkali metal|Group I]], [[Ammonium|NH<sub>4</sub><sup>+</sup>]] and [[uranyl]] compounds)
|-
| [[Nitrate]]s||[[Sulfite]]s (except [[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds)
|-
| [[Acetate]]s (ethanoates) (except [[Silver|Ag<sup>+</sup>]] compounds)||[[Phosphate]]s (except [[Alkali metal|Group I]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds (excluding [[Lithium|Li]]<sup>+</sup>))
|-
| [[Chloride]]s (chlorates and perchlorates), [[bromide]]s and [[iodide]]s (except [[Silver|Ag<sup>+</sup>]], [[Lead|Pb<sup>2+</sup>]], [[Copper|Cu<sup>+</sup>]] and [[Mercury (element)|Hg<sub>2</sub><sup>2+</sup>]])||[[Hydroxide]]s and [[oxide]]s (except [[Alkali metal|Group I]], [[Ammonium|NH<sub>4</sub><sup>+</sup>]], [[Barium|Ba<sup>2+</sup>]], [[Strontium|Sr<sup>2+</sup>]] and [[Thallium|Tl<sup>+</sup>]])
|-
| [[Sulfate]]s (except [[Silver|Ag<sup>+</sup>]], [[Lead|Pb<sup>2+</sup>]], [[Barium|Ba<sup>2+</sup>]], [[Strontium|Sr<sup>2+</sup>]] and [[Calcium|Ca<sup>2+</sup>]])||[[Sulfide]]s (except [[Alkali metal|Group I]], [[Alkaline earth metal|Group II]] and [[Ammonium|NH<sub>4</sub><sup>+</sup>]] compounds)
|}