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\(=\frac{1.079\,\text{mol}}{66.49 \times 10^{-3}\text{L}}\)
\(=16.3 \text{M}\)