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

Hint: 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.