Question

For preparing a buffer solution of ${{pH = 7}}{{.0}}$, which buffer system will you choose?
A. ${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}}{{,}}{{{H}}_{{2}}}{{PO}}_{{4}}^{{ - }}$
B. ${{{H}}_{{2}}}{{PO}}_{{4}}^{{ - }}{{,HPO}}_{{4}}^{{{2 - }}}$-
C. ${{{H}}_{{2}}}{{PO}}_{{4}}^{{ - }}{{,PO}}_{{4}}^{{{3 - }}}$
D. ${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}}{{,PO}}_{{4}}^{{{3 - }}}$

Answer 1

A buffer is an aqueous solution containing a weak acid and its conjugate base or a weak base and its conjugate acid. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. For example, blood in the human body is a buffer solution.

Complete step by step answer:
As ${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}}$ has three hydrogen attached, so it will be triprotic acid.
When ${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}}$ dissociates, it will give ${{3}}{{{H}}^{{ + }}}$ atoms in three steps:

Step 1:
${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}} \rightleftarrows {{{H}}^{{ + }}}{{ + }}{{{H}}_{{2}}}{{P}}{{{O}}_{{4}}}^{{ - }}$
Step 2:
${{{H}}_{{2}}}{{P}}{{{O}}_{{4}}}^{{ - }} \rightleftarrows {{{H}}^{{ + }}}{{ + HP}}{{{O}}_{{4}}}^{{{2 - }}}$
Step 3:
${{HP}}{{{O}}_{{4}}}^{{{2 - }}}{{ }} \to {{ }}{{{H}}^{{ + }}}{{ + P}}{{{O}}_{{4}}}^{{{3 - }}}$
The ${{pH}}$ formula for buffer solution (Henderson - HasselBalch Equation)
${{pH = p}}{{{k}}_{{a}}}{{ + log}}\dfrac{{\left[ {{s}} \right]}}{{\left[ {{a}} \right]}}$
As it is said in the question that the pH of the buffer is seven. So, there will be a mixture of weak acid and a salt of this weak acid with a strong base.
We assume the concentration of salt is equal to that of acid, then
${{pH = p}}{{{k}}_{{a}}}$
The ${{p}}{{{K}}_{{a}}}$ values for step 1, step 2 and step 3 are
${{p}}{{{k}}_{{a}}}_{{1}}\,{{ = 2}}{{.12,p}}{{{k}}_{{{a2}}}}\,{{ = 7}}{{.2}}\,\,{{and}}\,{{p}}{{{k}}_{{{a3}}}}{{ = 12}}{{.3}}$ respectively.
In Step 3, the ${{pk}}{{{a}}_2}\,{{ = 7}}{{.2}}$ is nearest to the asked ${{pH}}$ value.
In the given options, option (b) has the conjugate acid base system in the step 2.
So the correct answer is option (B).

Additional Information:
${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}}$Is a Bronsted-Lowry Acid.
${{{H}}_{{3}}}{{P}}{{{O}}_{{4}}} \rightleftarrows {{{H}}^{{ + }}} + {{{H}}_2}{{PO}}_4^ -$
${{{H}}_2}{{PO}}_4^ - \rightleftarrows {{{H}}^{{ + }}}{{ + HPO}}_4^{2 - }$
${{HPO}}_4^{2 - } \to {{{H}}^{{ + }}}{{ + PO}}_4^{3 - }$
${{{H}}_{{2}}}{{P}}{{{O}}_{{4}}}^{{ - }}$ commonly called as dihydrogen phosphate which is a monovalent inorganic anion which consists of phosphoric acid where one out of ${{3}}$ ${{OH}}$ groups have been deprotonated ${{{H}}_{{2}}}{{P}}{{{O}}_{{4}}}^{{ - }}$ is a conjugate base of Phosphoric acid having a molecular weight of $96.987{{g/mol}}$ .Most of the Dihydrogen phosphate salts are colourless in nature, water soluble and non-toxic.

Note:
A conjugated pair of an acid and base differs by a proton only.
If Bronsted acid is a strong acid then its conjugate base is a weak base and vice versa. The gases from which acids are formed by mixing them in water are called anhydride of acids.