Editing Common-ion effect

{{short description|Decrease in solubility of an ionic substance in solution when a common ion is added}}

In [[chemistry]], the '''common-ion effect''' refers to the decrease in [[solubility]] of an ionic [[Precipitation (chemistry)|precipitate]] by the addition to the [[Solution (chemistry)|solution]] of a soluble compound with an [[ion]] in common with the precipitate.<ref name=Skoog>{{cite book |last1=Skoog |first1=Douglas A. |last2=West |first2=Donald M. |last3=Holler |first3=F. James |last4=Crouch |first4=Stanley R. |title=Fundamentals of Analytical Chemistry |date=2014 |publisher=Brooks/Cole |isbn=978-0-495-55828-6 |page=209 |edition=9th}}</ref> This behaviour is a consequence of [[Le Chatelier's principle]] for the [[equilibrium reaction]] of the [[Ion association|ionic association]]/[[Dissociation (chemistry)|dissociation]]. The effect is commonly seen as an effect on the solubility of [[salt (chemistry)|salt]]s and other [[weak electrolyte]]s. Adding an additional amount of one of the ions of the salt generally leads to increased precipitation of the salt, which reduces the [[concentration]] of both ions of the salt until the [[solubility equilibrium]] is reached. The effect is based on the fact that both the original salt and the other added chemical have one ion in common with each other.

== Examples of the common-ion effect ==

===Dissociation of hydrogen sulfide in presence of hydrochloric acid===

[[Hydrogen sulfide]] (H<sub>2</sub>S) is a weak [[electrolyte]]. It is partially [[ionize]]d when in [[aqueous solution]], therefore there exists an equilibrium between un-ionized molecules and constituent ions in an aqueous medium as follows:

:H<sub>2</sub>S   {{eqm}}    H<sup>+</sup>  +  HS<sup>−</sup>

By applying the [[law of mass action]], we have

:<math chem>K_\ce{a}=\frac{[\text{H}^+] [\text{HS}^{-}]} {[\text{H}_2\text{S}]}</math>

[[Hydrochloric acid]] (HCl) is a strong electrolyte, which nearly completely ionizes as

:HCl      →  H<sup>+</sup>   +    Cl<sup>−</sup>

If HCl is added to the H<sub>2</sub>S solution, H<sup>+</sup> is a common ion and creates a common ion effect. Due to the increase in concentration of H<sup>+</sup> ions from the added HCl, the equilibrium of the dissociation of H<sub>2</sub>S shifts to the left and keeps the value of  K<sub>a</sub>  constant. Thus the dissociation of H<sub>2</sub>S decreases, the concentration of un-ionized H<sub>2</sub>S increases, and as a result, the concentration of sulfide ions decreases.

===Solubility of barium iodate in presence of barium nitrate===
Barium iodate, Ba(IO<sub>3</sub>)<sub>2</sub>, has a solubility product K<sub>sp</sub> = [Ba<sup>2+</sup>][IO<sub>3</sub><sup>−</sup>]<sup>2</sup> = 1.57 x 10<sup>−9</sup>. Its solubility in pure water is 7.32 x 10<sup>−4</sup> M. However in a solution that is 0.0200 M in barium nitrate, Ba(NO<sub>3</sub>)<sub>2</sub>, the increase in the common ion barium leads to a decrease in iodate ion concentration. The solubility is therefore reduced to 1.40 x 10<sup>−4</sup> M, about five times smaller.<ref name=Skoog/>

==Solubility effects==
{{main|solubility equilibrium}}
A practical example used very widely in areas drawing [[drinking water]] from [[chalk]] or [[limestone]] [[aquifers]] is the addition of sodium carbonate to the raw water to reduce the [[hard water|hardness]] of the water. In the [[water treatment]] process, highly soluble [[sodium carbonate]] salt is added to precipitate out sparingly soluble [[calcium carbonate]]. The very pure and finely divided precipitate of calcium carbonate that is generated is a valuable by-product used in the manufacture of [[toothpaste]].

The [[salting-out]] process used in the manufacture of [[soap]]s benefits from the common-ion effect. Soaps are sodium salts of [[fatty acid]]s. Addition of sodium chloride reduces the solubility of the soap salts. The soaps precipitate due to a combination of common-ion effect and increased [[ionic strength]].

Sea, brackish and other waters that contain appreciable amount of [[sodium]] ions (Na<sup>+</sup>) interfere with the normal behavior of soap because of common-ion effect. In the presence of excess Na<sup>+</sup>, the solubility of soap salts is reduced, making the soap less effective.

== Buffering effect==
{{main|buffer solution}}
A buffer solution contains an acid and its conjugate base or a base and its conjugate acid.<ref>{{VogelQuantitative6th|page=28}}</ref> Addition of the conjugate ion will result in a change of [[pH]] of the buffer solution. For example, if both [[sodium acetate]] and [[acetic acid]] are dissolved in the same solution they both dissociate and ionize to produce [[acetate ion]]s. Sodium acetate is a [[strong electrolyte]], so it dissociates completely in solution. Acetic acid is a [[weak acid]], so it only ionizes slightly. According to [[Le Chatelier's principle]], the addition of acetate ions from sodium acetate will suppress the ionization of acetic acid and shift its equilibrium to the left. Thus the percent dissociation of the acetic acid will decrease, and the pH of the solution will increase. The ionization of an acid or a base is limited by the presence of its conjugate base or acid.

: NaCH<sub>3</sub>CO<sub>2</sub>(s) → Na<sup>+</sup>(aq) + CH<sub>3</sub>CO<sub>2</sub><sup>−</sup>(aq)

: CH<sub>3</sub>CO<sub>2</sub>H(aq) {{eqm}} H<sup>+</sup>(aq) + CH<sub>3</sub>CO<sub>2</sub><sup>−</sup>(aq)

This will decrease the [[hydronium]] concentration, and thus the common-ion solution will be less acidic than a solution containing only acetic acid.

==Exceptions==
Many transition-metal compounds violate this rule due to the formation of complex ions, a scenario not part of the equilibria that are involved in simple precipitation of salts from ionic solution. For example, [[copper(I) chloride]] is insoluble in water, but it dissolves when chloride ions are added, such as when [[hydrochloric acid]] is added. This is due to the formation of soluble CuCl<sub>2</sub><sup>−</sup> [[coordination complex|complex]] ions.

== Uncommon-ion effect ==
Sometimes adding an ion other than the ones that are part of the precipitated salt itself can increase the solubility of the salt. This "[[salting in]]" is called the "uncommon-ion effect" (also "salt effect" or the "diverse-ion effect"). It occurs because as the total ion concentration increases, inter-ion attraction within the solution can become an important factor.<ref name="Boyd2015">{{cite book |author=Claude E. Boyd |title=Water Quality: An Introduction |url=https://books.google.com/books?id=ODwwCgAAQBAJ&pg=PA56 |date=14 July 2015 |publisher=Springer |isbn=978-3-319-17446-4 |pages=56–}}</ref> This alternate equilibrium makes the ions less available for the precipitation reaction. This is also called odd ion effect.

== References ==
{{reflist}}

{{DEFAULTSORT:Common-Ion Effect}}
[[Category:Equilibrium chemistry]]
[[Category:Solutions]]