Question

A 36.3 mL aliquot (sample) of 0.529 M H2SO4 (aq) is to be titrated with 0.0411 M NaOH. What volume of base will it take to reach the equivalence point?

Answer 1

Explanation:

Sulfuric acid ($H_{2}SO_{4}$) is a diprotic acid. So, it means that we need twice as much NaOH.

As we known that an equivalence point is reached when moles of an acid equals the moles of a base.

As moles of $H_{2}SO_{4}$ is calculated as follows.

     $36.3 mL \times \frac{1 L}{1000 mL} \times  0.0529 M H_{2}SO_{4}$ = 0.00192 moles of $H_{2}SO_{4}$

Therefore, moles of NaOH needed will be as follows.

            $2 \times 0.00192 moles$

          = 0.00384 mol

Hence, volume of NaOH is calculated as follows.

     $\frac{1000 mL}{1 L} \times \frac{0.00384 moles}{0.0411 M NaOH}$                  

                      = 93.43 mL

Thus, we can conclude that the volume of base required to reach the equivalence point is 93.43 mL.

Answer 2

Answer: 934.4 ml

Explanation:

According to the neutralization law,

$n_1M_1V_1=n_2M_2V_2$

where,

$M_1$ = molarity of $H_2SO_4$ solution = 0.529 M

$V_1$ = volume of $H_2SO_4$ solution = 36.3 ml

$M_2$ = molarity of $NaOH$ solution = 0.0411 M

$V_2$ = volume of $NaOH$ solution = ?

$n_1$ = valency of $H_2SO_4$ = 2

$n_2$ = valency of $NaOH$ = 1

$2\times 0.529M\times 36.3=1\times 0.0411\times V_2$

$V_2=934.4ml$

Therefore, the volume of the base required to reach the equivalence point is 934.4 ml