Editing Acid (section)

===Polyprotic acids===
{{See also|Acid dissociation constant#Polyprotic acids}}
Polyprotic acids, also known as polybasic acids, are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic (or dibasic) acid (two potential protons to donate), and triprotic (or tribasic) acid (three potential protons to donate). Some macromolecules such as proteins and nucleic acids can have a very large number of acidic protons.<ref>{{cite book |title=Biophysical Chemistry - Volume 1 |first1=Jeffries|last1= Wyman|first2= John |last2=Tileston Edsall |chapter=Chapter 9: Polybasic Acids, Bases, and Ampholytes, Including Proteins | page=477 }}</ref>

A diprotic acid (here symbolized by H<sub>2</sub>A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, K<sub>a1</sub> and K<sub>a2</sub>.
:{{chem2|H2A (aq) + H2O (l) <-> H3O+ (aq) + HA- (aq)}} &nbsp;&nbsp;&nbsp; ''K''<sub>a1</sub>
:{{chem2|HA- (aq) + H2O (l) <-> H3O+ (aq) + A(2−) (aq)}} &nbsp;&nbsp;&nbsp; &nbsp; ''K''<sub>a2</sub>

The first dissociation constant is typically greater than the second (i.e., ''K''<sub>a1</sub> > ''K''<sub>a2</sub>). For example, [[sulfuric acid]] (H<sub>2</sub>SO<sub>4</sub>) can donate one proton to form the [[bisulfate]] anion (HSO{{su|b=4|p=−}}), for which ''K''<sub>a1</sub> is very large; then it can donate a second proton to form the [[sulfate]] anion (SO{{su|b=4|p=2−}}), wherein the ''K''<sub>a2</sub> is intermediate strength. The large ''K''<sub>a1</sub> for the first dissociation makes sulfuric a strong acid. In a similar manner, the weak unstable [[carbonic acid]] {{nowrap|(H<sub>2</sub>CO<sub>3</sub>)}} can lose one proton to form [[bicarbonate]] anion {{nowrap|(HCO{{su|b=3|p=−}})}} and lose a second to form [[carbonate]] anion (CO{{su|b=3|p=2−}}). Both ''K''<sub>a</sub> values are small, but ''K''<sub>a1</sub> > ''K''<sub>a2</sub> .

A triprotic acid (H<sub>3</sub>A) can undergo one, two, or three dissociations and has three dissociation constants, where ''K''<sub>a1</sub> > ''K''<sub>a2</sub> > ''K''<sub>a3</sub>.
:{{chem2|H3A (aq) + H2O (l) <-> H3O+ (aq) + H2A− (aq)}} &nbsp;&nbsp;&nbsp;&nbsp; ''K''<sub>a1</sub>
:{{chem2|H2A− (aq) + H2O (l) <-> H3O+ (aq) + HA(2−) (aq)}} &nbsp; &nbsp; &nbsp; ''K''<sub>a2</sub>
:{{chem2|HA(2−) (aq) + H2O (l) <-> H3O+ (aq) + A(3−) (aq)}} &nbsp; &nbsp;&nbsp; ''K''<sub>a3</sub>

An [[inorganic]] example of a triprotic acid is orthophosphoric acid (H<sub>3</sub>PO<sub>4</sub>), usually just called [[phosphoric acid]]. All three protons can be successively lost to yield H<sub>2</sub>PO{{su|b=4|p=−}}, then HPO{{su|b=4|p=2−}}, and finally PO{{su|b=4|p=3−}}, the orthophosphate ion, usually just called [[phosphate]]. Even though the positions of the three protons on the original phosphoric acid molecule are equivalent, the successive ''K''<sub>a</sub> values differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged. An [[organic compound|organic]] example of a triprotic acid is [[citric acid]], which can successively lose three protons to finally form the [[citrate]] ion.

Although the subsequent loss of each hydrogen ion is less favorable, all of the conjugate bases are present in solution. The fractional concentration, ''α'' (alpha), for each species can be calculated. For example, a generic diprotic acid will generate 3 species in solution: H<sub>2</sub>A, HA<sup>−</sup>, and A<sup>2−</sup>. The fractional concentrations can be calculated as below when given either the pH (which can be converted to the [H<sup>+</sup>]) or the concentrations of the acid with all its conjugate bases:
:<math chem>\begin{align}
\alpha_\ce{H2A}   &= \frac{\ce{[H+]^2}}{\ce{[H+]^2}  + [\ce{H+}]K_1 + K_1 K_2}  = \frac{\ce{[H2A]}}{\ce{{[H2A]}} + [HA^-] + [A^{2-}]}\\ 
\alpha_\ce{HA^-}  &= \frac{[\ce{H+}]K_1}{\ce{[H+]^2} + [\ce{H+}]K_1 + K_1 K_2}  = \frac{\ce{[HA^-]}}{\ce{[H2A]}+{[HA^-]}+{[A^{2-}]}}\\
\alpha_\ce{A^{2-}}&= \frac{K_1 K_2}{\ce{[H+]^2} + [\ce{H+}]K_1 + K_1 K_2} = \frac{\ce{[A^{2-}]}}{\ce{{[H2A]}}+{[HA^-]}+{[A^{2-}]}}
\end{align}</math>
A plot of these fractional concentrations against pH, for given ''K''<sub>1</sub> and ''K''<sub>2</sub>, is known as a [[Bjerrum plot]]. A pattern is observed in the above equations and can be expanded to the general ''n'' -protic acid that has been deprotonated ''i'' -times:
:<math chem>
\alpha_{\ce H_{n-i} A^{i-} }= { {[\ce{H+}]^{n-i} \displaystyle \prod_{j=0}^{i}K_j} \over { \displaystyle \sum_{i=0}^n \Big[ [\ce{H+}]^{n-i} \displaystyle \prod_{j=0}^{i}K_j} \Big] }
</math>

where ''K''<sub>0</sub> = 1 and the other K-terms are the dissociation constants for the acid.