Editing Acid (section)

==Dissociation and equilibrium==
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Reactions of acids are often generalized in the form {{chem2|HA <-> H+ + A-}}, where HA represents the acid and A<sup>−</sup> is the [[conjugate acid|conjugate base]]. This reaction is referred to as '''protolysis'''. The protonated form (HA) of an acid is also sometimes referred to as the '''free acid'''.<ref>{{cite book | editor1-last = Stahl | editor1-first = P. Heinrich | editor2-last = Warmth | editor2-first = Camille G. | last1 = Stahl | first1 = P. Heinrich | last2 = Nakamo | first2 = Masahiro | name-list-style = vanc | title = Handbook of Pharmaceutical Salts: Properties, Selection, and Use | date = 2008 | publisher = Wiley-VCH | location = Weinheim | isbn = 978-3-906390-58-1 | chapter = Pharmaceutical Aspects of the Salt Form | chapter-url = https://books.google.com/books?id=IvSEXUZUON8C&dq=%22free+acid%22+salt&pg=PA92 | pages = 92–94 }}</ref>

Acid–base conjugate pairs differ by one proton, and can be interconverted by the addition or removal of a proton ([[protonation]] and [[deprotonation]], respectively). The acid can be the charged species and the conjugate base can be neutral in which case the generalized reaction scheme could be written as {{chem2|HA+ <-> H+ + A}}. In solution there exists an [[chemical equilibrium|equilibrium]] between the acid and its conjugate base. The [[equilibrium constant]] ''K'' is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that [H<sub>2</sub>O] means ''the concentration of H<sub>2</sub>O''. The [[acid dissociation constant]] ''K''<sub>a</sub> is generally used in the context of acid–base reactions. The numerical value of ''K''<sub>a</sub> is equal to the [[Product (mathematics)|product]] (multiplication) of the concentrations of the products divided by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and H<sup>+</sup>.
:<math chem>K_a = \frac\ce{[H+] [A^{-}]}\ce{[HA]}</math>
The stronger of two acids will have a higher ''K''<sub>a</sub> than the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for ''K''<sub>a</sub> spans many orders of magnitude, a more manageable constant, p''K''<sub>a</sub> is more frequently used, where p''K''<sub>a</sub> = &minus;log<sub>10</sub> ''K''<sub>a</sub>. Stronger acids have a smaller p''K''<sub>a</sub> than weaker acids. Experimentally determined p''K''<sub>a</sub> at 25&nbsp;°C in aqueous solution are often quoted in textbooks and reference material.