Question

A buffer is made by dissolving $NaClO$ and $HClO$ in water. How do you write equations to show how this buffer neutralizes added ${H^ + }$ and $O{H^ - }$?

Answer 1

We know that buffer is a solution which could resist change in $pH$ when acidic components (or) basic components are added. It has the ability to neutralize certain amounts of acid (or) base to maintain the pH of the solution which is stable. Usually, buffer solutions contain a working range of pH and capacity that dictate the amount of acid/base could be neutralized before changes in $pH$.

Complete step by step answer:
We know that buffer is a solution that opposes sensational changes in $pH$. Buffers do as by containing a few pairs of solute; either a weak acid along with a salt obtained from that acid that are weak (or) weak base along with salt of that base.
Let us consider the equilibrium of the buffer system. We can write the equation as,
$HClO \rightleftarrows Cl{O^ - } + {H^ + }$
Here, the equilibrium constant ${K_a}$ can be written as,
${K_a} = \dfrac{{\left[ {Cl{O^ - }} \right]\left[ {{H^ + }} \right]}}{{\left[ {HClO} \right]}} \approx 3.0 \times {10^{ - 8}}$
Let us now consider the ionization of water. We can write the equation as,
${H_2}O \rightleftarrows {H^ + } + O{H^ - }$
Here, the equilibrium constant for ionization of water is ${K_w}$. ${K_w}$ is the product of concentration of hydroxide ions and concentration of hydrogen ions. We can write expression for ${K_w}$ as,
${K_w} = \left[ {O{H^ - }} \right]\left[ {{H^ + }} \right] \approx 1.0 \times {10^{ - 14}}$
The preceding equation could be used to understand what happens when hydrogen ions (or) ions of hydroxide are added to the solution of the buffer.
So, when we add hydroxide ${Q_w} > {K_w}$ transiently. The equilibrium constant for ionization is restored by the dissociation of $HClO$. However, when we do so ${Q_w} < {K_w}$, therefore $HClO$ should dissociate to restore its equilibrium. So, $pH$ would decrease slightly.

Note: We have to know that buffer contains some capacity. We have to know that if a strong acid that is a source of ions of hydrogens is added to the solution of the buffer, the ions of hydrogen could react with anions from the salt. Buffers act well for certain quantities of strong acid (or) strong base. Once the solute is completely reacted, the solution does not act as a buffer anymore, and changes in $pH$ could be rapid.