Editing Acid dissociation constant (section)

== Monoprotic acids ==
{{See also|Acid#Monoprotic acids}}
[[File:Weak acid speciation.svg|thumb|200px|right|alt=This figure plots the relative fractions of the protonated form A H of an acid to its deprotonated form, A minus, as the solution p H is varied about the value of the acid's p K A. When the p H equals the p K a, the amounts of the protonated and deprotonated forms are equal. When the p H is one unit higher than the p K A, the ratio of concentrations of protonated to deprotonated forms is 10 to 1. When the p H is two units higher that ratio is 100 to 1. Conversely, when the p H is one or two unit lower than the p K A, the ratio is 1 to ten or 1 to 100. The exact percentage of each form may be determined from the Henderson–Hasselbalch equation.|Variation of the % formation of a monoprotic acid, AH, and its conjugate base, A<sup>−</sup>, with the difference between the pH and the p''K''<sub>a</sub> of the acid.]]

After rearranging the expression defining ''K''<sub>a</sub>, and putting {{nowrap|pH {{=}} −log<sub>10</sub>[H<sup>+</sup>]}}, one obtains<ref>{{cite web | last1= Mehta| first1= Akul| title= Henderson–Hasselbalch Equation: Derivation of p''K''<sub>a</sub> and p''K''<sub>b</sub>| url= http://pharmaxchange.info/press/2012/10/henderson%E2%80%93hasselbalch-equation-derivation-of-pka-and-pkb/| website = PharmaXChange| date= 22 October 2012| access-date = 16 November 2014}}</ref>

:<math alt="p H equals p K A plus the logarithm (base ten) of a ratio of chemical concentrations, namely the concentration of the protonated form A H divided by that of the deprotonated form A minus." >
 \mathrm{pH} = \mathrm{p}K_\text{a} + \log\mathrm{\frac{[A^-]}{[HA]}}
</math>

This is the [[Henderson–Hasselbalch equation]], from which the following conclusions can be drawn.
* At half-neutralization the ratio {{nowrap|{{sfrac|[A<sup>−</sup>]|[HA]}} {{=}} 1}}; since {{nowrap|log(1) {{=}} 0}}, the pH at half-neutralization is numerically equal to p''K''<sub>a</sub>. Conversely, when {{nowrap|pH {{=}} p''K''<sub>a</sub>}}, the concentration of HA is equal to the concentration of A<sup>−</sup>.
* The [[buffer solution|buffer region]] extends over the approximate range p''K''<sub>a</sub>&nbsp;±&nbsp;2. Buffering is weak outside the range p''K''<sub>a</sub>&nbsp;±&nbsp;1. At pH ≤ p''K''<sub>a</sub> − 2 the substance is said to be fully protonated and at pH ≥ p''K''<sub>a</sub> + 2 it is fully dissociated (deprotonated).
* If the pH is known, the ratio may be calculated. This ratio is independent of the analytical concentration of the acid.

In water, measurable p''K''<sub>a</sub> values range from about −2 for a strong acid to about 12 for a very weak acid (or strong base).

A [[buffer solution]] of a desired pH can be prepared as a mixture of a weak acid and its conjugate base. In practice, the mixture can be created by dissolving the acid in water, and adding the requisite amount of strong acid or base. When the p''K''<sub>a</sub> and analytical concentration of the acid are known, the extent of dissociation and pH of a solution of a monoprotic acid can be easily calculated using an [[ICE table]].