Editing Acid–base reaction (section)

==Acid–base equilibrium==
{{Main|Acid dissociation constant}}
The reaction of a strong acid with a strong base is essentially a quantitative reaction. For example,
<math chem display=block>\ce{HCl_{(aq)} {} + Na(OH)_{(aq)} -> H2O + NaCl_{(aq)} }</math>
 
In this reaction both the sodium and chloride ions are spectators as the neutralization reaction,
<math chem display=block>\ce{H + OH- -> H2O}</math>
does not involve them. With weak bases addition of acid is not quantitative because a solution of a weak base is a [[buffer solution]]. A solution of a weak acid is also a buffer solution. When a weak acid reacts with a weak base an equilibrium mixture is produced. For example, [[adenine]], written as AH, can react with a hydrogen [[phosphate]] ion, {{chem2|HPO4(2-)}}.
<math chem display=block>\ce{AH + HPO4^2- <=> A- + H2PO4-}</math>

The equilibrium constant for this reaction can be derived from the acid dissociation constants of adenine and of the dihydrogen phosphate ion.
<math chem display=block>\begin{align}
\left[\ce{A-}\right] \! \left[\ce{H+}\right] &= K_{a1}\bigl[\ce{AH}\bigr] \\[4pt]
\left[\ce{HPO4^2-}\right] \! \left[\ce{H+}\right] &= K_{a2}\left[\ce{H2PO4-}\right]
\end{align}</math>

The notation [X] signifies "concentration of X". When these two equations are combined by eliminating the hydrogen ion concentration, an expression for the equilibrium constant, {{mvar|K}} is obtained.
<math chem display=block>\left[\ce{A-}\right] \! \left[\ce{H2PO4-}\right] = K \bigl[\ce{AH}\bigr] \! \left[\ce{HPO4^2-}\right]; \quad K = \frac{K_{a1}}{K_{a2}}</math>