View source for Dissociation (chemistry)

{{Short description|Separation of molecules or ionic compounds into smaller constituent entities}}
{{Redirect2|Dissociate|Dissociating|other uses|Dissociation (disambiguation){{!}}Dissociation}}
{{More citations needed|date=June 2014}}{{Multiple image
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| alt1 = Dissociation of [[Gomberg's dimer]]
| caption1 = Dissociation of [[Gomberg's dimer]]
| image2 = Phosphoric acid dissociation.png
| caption2 = Dissociation diagram of [[phosphoric acid]]
| alt2 = Dissociation diagram of [[phosphoric acid]]
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'''Dissociation''' in [[chemistry]] is a general process in which molecules (or ionic compounds such as [[salt (chemistry)|salts]], or [[coordination complex|complexes]]) separate or split into other things such as atoms, ions, or [[radical (chemistry)|radicals]], usually in a reversible manner. For instance, when an [[acid]] dissolves in water, a [[covalent bond]] between an [[electronegativity|electronegative]] atom and a hydrogen atom is broken by [[heterolytic fission]], which gives a proton (H<sup>+</sup>) and a negative [[ion]]. Dissociation is the opposite of '''association''' or '''recombination'''.

==Dissociation constant==
{{Main|Dissociation constant}}
For reversible dissociations in a [[chemical equilibrium]]
:<chem>AB <=> A + B</chem>
the [[dissociation constant]] ''K''<sub>d</sub> is the ratio of dissociated to undissociated compound
:<math>K_d = \mathrm{\frac{[A] [B]}{[AB]}}</math>
where the brackets denote the equilibrium concentrations of the species.<ref name=Atkins/>

==Dissociation degree==
The dissociation degree <math>\alpha</math> is the fraction of original solute molecules that have dissociated. It is usually indicated by the Greek symbol α. More accurately, degree of dissociation refers to the amount of solute dissociated into ions or radicals per mole. In case of very strong acids and bases, degree of dissociation will be close to 1. Less powerful acids and bases will have lesser degree of dissociation. There is a simple relationship between this parameter and the [[van 't Hoff factor]] '''<math>i</math>'''. If the solute substance dissociates into <math>n</math> ions, then
:<math>i = 1 + \alpha (n - 1)</math>

For instance, for the following dissociation
:<chem>KCl <=> K+ + Cl-</chem>

As <math>n = 2</math>, we would have that <math>i = 1 + \alpha</math>.

==Salts==
{{See also|Solubility equilibrium}}
[[File:Dissolution of sodium chloride crystal under a microscope.ogv|thumb|A video of [[sodium chloride]] crystals dissolving and dissociating in water]]
The dissociation of salts by [[solvation]] in a [[Solution (chemistry)|solution]], such as [[water (molecule)|water]], means the separation of the [[anion]]s and [[cation]]s. The salt can be recovered by [[evaporation]] of the solvent.

An ''[[electrolyte]]'' refers to a substance that contains free ions and can be used as an [[electrically conductive]] medium. Most of the solute does not dissociate in a weak electrolyte, whereas in a strong electrolyte a higher ratio of solute dissociates to form free ions.

A weak electrolyte is a substance whose solute exists in solution mostly in the form of molecules (which are said to be "undissociated"), with only a small fraction in the form of ions. Simply because a substance does not readily dissolve does not make it a weak electrolyte. [[Acetic acid]] ({{chem2|CH3COOH}}) and [[ammonium]] ({{chem2|NH4+}}) are good examples. Acetic acid is extremely soluble in water, but most of the compound dissolves into molecules, rendering it a weak electrolyte. Weak bases and weak acids are generally weak electrolytes. In an aqueous solution there will be some {{chem2|CH3COOH}} and some {{chem2|CH3COO-}} and {{chem2|H+}}.

A strong electrolyte is a solute that exists in solution completely or nearly completely as ions. Again, the strength of an electrolyte is defined as the percentage of solute that is ions, rather than molecules. The higher the percentage, the stronger the electrolyte. Thus, even if a substance is not very soluble, but does dissociate completely into ions, the substance is defined as a strong electrolyte. Similar logic applies to a weak electrolyte. Strong acids and bases are good examples, such as HCl and {{chem2|H2SO4}}. These will all exist as ions in an aqueous medium.

==Gases==
The degree of dissociation in [[gas]]es is denoted by the symbol {{mvar|α}}, where {{mvar|α}} refers to the percentage of gas molecules which dissociate. Various relationships between {{mvar|K{{sub|p}}}} and {{mvar|α}} exist depending on the [[stoichiometry]] of the equation. The example of [[dinitrogen tetroxide]] ({{chem2|N2O4}}) dissociating to [[nitrogen dioxide]] ({{chem2|NO2}}) will be taken.
<math chem display=block>\ce{N2O4 <=> 2NO2}</math>

If the initial concentration of dinitrogen tetroxide is 1&nbsp;[[Mole (unit)|mole]] per [[litre]], this will decrease by {{mvar|α}} at equilibrium giving, by stoichiometry, {{mvar|α}} moles of {{chem2|NO2}}. The [[equilibrium constant]] (in terms of pressure) is given by the equation
<math chem display=block>K_p = \frac{p\bigl(\ce{NO2}\bigr)^2}{p \, \ce{N2O4}}</math>

where {{mvar|p}} represents the [[partial pressure]]. Hence, through the definition of partial pressure and using {{mvar|p{{sub|T}}}} to represent the total pressure and {{mvar|x}} to represent the [[mole fraction]];
<math chem display=block>K_p = \frac{p_T^2 \bigl(x \, \ce{NO2} \bigr)^2}{p_T\cdot x \, \ce{N2O4}} = \frac{p_T \bigl(x \, \ce{NO2} \bigr)^2}{x \, \ce{N2O4}}</math>

The total number of moles at equilibrium is {{math|(1 – ''α'') + 2''α''}}, which is equivalent to {{math|1 + ''α''}}. Thus, substituting the mole fractions with actual values in term of {{mvar|α}} and simplifying;
<math display=block>K_p = \frac{p_T(4\alpha^2)}{(1+\alpha)(1-\alpha)} = \frac{p_T(4\alpha^2)}{1-\alpha^2}</math>

This equation is in accordance with [[Le Chatelier's principle]]. {{mvar|K{{sub|p}}}} will remain constant with temperature. The addition of pressure to the system will increase the value of {{mvar|p{{sub|T}}}}, so {{mvar|α}} must decrease to keep {{mvar|K{{sub|p}}}} constant. In fact, increasing the pressure of the equilibrium favours a shift to the left favouring the formation of dinitrogen tetroxide (as on this side of the equilibrium there is less pressure since pressure is proportional to number of moles) hence decreasing the extent of dissociation {{mvar|α}}.

==Acids in aqueous solution==
{{Main|Acid dissociation constant}}
The reaction of an acid in water solvent is often described as a dissociation
:<chem>HA <=> H+ + A-</chem>
where HA is a proton acid such as acetic acid, CH<sub>3</sub>COOH. The double arrow means that this is an equilibrium process, with dissociation and recombination occurring at the same time. This implies that the [[acid dissociation constant]]
:<math chem>K_{\ce a} = \ce{\frac{[H^+] [A^-]}{[HA]}}</math>

However a more explicit description is provided by the [[Brønsted–Lowry acid–base theory]], which specifies that the ''proton'' H+ does not exist as such in solution but is instead ''accepted'' by (bonded to) a water molecule to form the [[hydronium]] ion H<sub>3</sub>O<sup>+</sup>.

The reaction can therefore be written as
:<chem>HA + H2O <=> H3O+ + A-</chem>
and better described as an ''ionization'' or formation of ions (for the case when HA has no net charge). The equilibrium constant is then
:<math chem>K_{\ce a} = \ce{\frac{[H_3O^+] [A^-]}{[HA]}}</math>
where <chem>[H_2O]</chem> is not included because in dilute solution the solvent is essentially a pure liquid with a [[thermodynamic activity]] of one.<ref name=":0">{{cite book|title=General chemistry: principles and modern applications|url=https://archive.org/details/generalchemistry00hill|url-access=registration|last1=Petrucci|first1=Ralph H.|last2=Harwood|first2=William S.|last3=Herring|first3=F. Geoffrey|date=2002|publisher=Prentice Hall|isbn=978-0-13-014329-7|edition=8th|location=Upper Saddle River, N.J|lccn=2001032331|oclc=46872308}}</ref>{{Rp|668}}

K<sub>a</sub> is variously named a ''dissociation constant'',<ref>[[Keith Laidler|Laidler K.J.]] ''Physical Chemistry with Biological Applications'' (Benjamin/Cummings) 1978, p.307 {{ISBN|978-0-8053-5680-9}}</ref> an ''acid ionization constant'',<ref name=":0" />{{Rp|668}} an ''acidity constant''<ref name="Atkins">Atkins P. and de Paula J. ''Physical Chemistry'' (8th ed. W.H.Freeman 2006) p.763 {{ISBN|978-0-7167-8759-4}}</ref> or an ''ionization constant''.<ref name=":0" />{{Rp|708}} It serves as an indicator of the acid strength: stronger acids have a higher ''K''<sub>a</sub> value (and a lower p''K''<sub>a</sub> value).

==Fragmentation==
[[Fragmentation (chemistry)|Fragmentation]] of a molecule can take place by a process of [[heterolysis (chemistry)|heterolysis]] or [[homolysis (chemistry)|homolysis]].

==Receptors==
[[Receptor (biochemistry)|Receptors]] are [[protein]]s that bind small [[ligand]]s. The dissociation constant ''K''<sub>d</sub> is used as indicator of the [[affinity (pharmacology)|affinity]] of the ligand to the receptor. The higher the affinity of the ligand for the receptor the lower the ''K''<sub>d</sub> value (and the higher the p''K''<sub>d</sub> value).

==See also==
* [[Bond-dissociation energy]]
* [[Photodissociation]], dissociation of molecules by photons (light, gamma rays, x-rays)
* [[Radiolysis]], dissociation of molecules by ionizing radiation
* [[Thermal decomposition]]

==References==
{{Reflist}}

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[[Category:Chemical processes]]
[[Category:Equilibrium chemistry]]