# Question #6f539

Jul 1, 2015

The effect of strong base on water is to dramatically increase the concentration of $O {H}^{-}$ ions and decrease the concentration of ${H}_{3} {O}^{+}$ ions.

#### Explanation:

Water always contains at least small concentrations of both $O {H}^{-}$ (hydroxide) and ${H}_{3} {O}^{+}$ (hydronium) ions. This is because water can react with itself in a self-ionization reaction:

$2 {H}_{2} O \leftrightarrow {H}_{3} {O}^{+} + O {H}^{-}$

At equilibrium, which is attained quickly for this reaction at room temperature, the product of $\left[{H}_{3} {O}^{+}\right]$ and $\left[O {H}^{-}\right]$ concentrations is approximately $1 \times {10}^{- 14}$.

Pure water contains equal concentrations of the two ions, so at neutral pH conditions (pH=7),
$\left[{H}_{3} {O}^{+}\right] = \left[O {H}^{-}\right] \approx 1 \times {10}^{- 7} M$

Addition of strong base (e.g., NaOH) releases hydroxide ions directly into solution because NaOH dissolves completely as $N {a}^{+}$ and $O {H}^{-}$ ions. A 0.01 M solution of NaOH has
$\left[O {H}^{-}\right] = 0.01 M$ and $\left[{H}_{3} {O}^{+}\right] = 1 \times {10}^{- 12} M$
The pH of this solution is 12.

A buffer solution consists of moderately high concentrations of an acid and its conjugate base. An example is a solution containing 0.1M carbonic acid $\left({H}_{2} C {O}_{3}\right)$ and 0.1M bicarbonate ion $\left(H C {O}_{3}^{-}\right)$, for example from dissolving potassium bicarbonate.

${H}_{2} C {O}_{3} + {H}_{2} O \leftrightarrow {H}_{3} {O}^{+} + H C {O}_{3}^{-}$

Addition of strong base to this buffer has the effect of reacting with some of the buffer acid $\left({H}_{2} C {O}_{3}\right)$ and converting it to base $\left(H C {O}_{3}^{-}\right)$. As long as the concentration of added strong base is much smaller than the concentrations of buffer acid and base, the pH will not change very much. We say that the pH is buffered by the presence of the acid/base combination, which can absorb small concentrations of added strong base without much effect.

In the same way, addition of small amounts of strong acid to a buffer solution will react with the buffer base and convert it to acid, leaving the pH nearly unchanged.

Blood contains many different buffer acid/base combinations, but the most important one is carbonic acid / bicarbonate ion, which keeps the pH of blood nearly constant at about 7.35.