Question

Concentrated nitric acid used in laboratory work is $68\%$ nitric acid by mass in aqueous solution. What should be the molarity of such a sample of the acid if the density of the solution is $1.504\, \text{g mL}^{–1}?$

Answer 1

Concentrated nitric acid used in laboratory work is $68\%$ nitric acid by mass in an aqueous solution.
This means that 68g of nitric acid is dissolved in 100g of the solution.
Molar mass of nitric acid $(\text{HNO}_3) = 1 \times 1 + 1 \times14 + 3 \times 16 = 63 \,\text{g mol}^{ - 1}$
Then, number of moles of $\text{HNO}_3 =\frac{68}{63}\, \text{mol}$
$=1.079\text{mol}$
Given,
Density of solution $= 1.504 \,\text{g mL}^{ - 1 }$
Volume of 100g solution $=\frac{100}{1.504} \text{mL}$
$=66.49\text{mL}$
$=66.49 \times 10-3\text{L}$
Molarity of solution$=\frac{1.079\,\text{mol}}{66.49 \times 10^{-3}\text{L}}$
$=16.3 \text{M}$