A biochemical reaction takes place in a $1\mathrm{.}00mL$ solution of $0\mathrm{.}0250M$

phosphate buffer initially at $pH=7\mathrm{.}35.$

Part A

Are the concentrations of any of the four possible phosphate species negligible?

Check all that apply.

$H_{2}P{O}_4^{-}$

$P{O}_4^{3-}$

$H_{3}P{O}_4$

$HP{O}_4^{2-}$

none of the above

Previous Answers

Correct

At $pH=7\mathrm{.}35[H_{3}P{O}_4]$ and $P{O}_4^{3-}$ will be negligible. The $p{K}_a$ of $H_{3}P{O}_4=2\mathrm{.}14,$ which is $$ five $pH$ units below $7\mathrm{.}35.$ Thus, the Henderson$-$Hasselbalch equation predicts

that $[H_{2}P{O}_4^{-}]>[H_{3}P{O}_4]$ by five orders of magnitude $(100,000$:$1)$ at $pH7\mathrm{.}35.$ Likewise, the $p{K}_a$ of $HP{O}_4^{2-}=12\mathrm{.}4,$ which is $$ five $pH$ units above $7\mathrm{.}35.$ Thus, the

Henderson$-$Hasselbalch equation predicts that $[HP{O}_4^{2-}]>[P{O}_4^{3-}]$ by five orders of magnitude $(100,000$:$1)$ at $pH7\mathrm{.}35.$

Part B

During the reaction, $3\mathrm{.}00mol$ of $HCl$ are produced. Calculate the final $pH$ of the reaction solution. Assume that the $HCl$ is completely neutralized by the buffer.

Express your answer using two decimal places.

$pH=$