Question

Concentrated nitric acid used in the 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\[1.504\text{g m}{{\text{L}}^{-1}}\]?

Answer 1

Hint:This question can be solved from the concept of Molarity of a solution which is the number of moles of a substance present per litre of the solution.
Formulae Used:
 $\text{mass=volume }\!\!\times\!\!\text{ density}$
Molarity of a solution = $\dfrac{\text{no}\text{. of moles}}{\text{volume}}$expressed in moles/litre.

Complete step by step answer:
The percentage of nitric acid used for laboratory purposes is 68%. Hence, in 100 g of dilute nitric acid, 68 g is nitric acid and the rest is water.
Molecular weight of nitric acid, $\text{HN}{{\text{O}}_{3}}$=$\left[ 1+14+\left( 16\times 3 \right) \right]=63\text{g/mol}$
So, according to the formula, no. of moles of nitric acid in $68$g = $\dfrac{68}{63}=1.08$moles
The density of the solution as given in the question = \[1.504\text{g m}{{\text{L}}^{-1}}\]
From the relation of mass, density, and volume, we get,
$\text{mass=volume }\!\!\times\!\!\text{ density}$
Therefore volume of 100 g dilute nitric acid having density \[1.504\text{g m}{{\text{L}}^{-1}}\]
=$\dfrac{100}{1.504}=66.5$mL = 0.0665 L of nitric acid.
So, the molarity of the nitric acid solution:
$\dfrac{\text{no}\text{. of moles}}{\text{volume}}$= $\dfrac{1.08}{0.0665}=16.22$$\left( \text{M} \right)$or $16.22$moles/litre.

Note:
Nitric acid is also known by the name “Aqua Fortis” and is a highly corrosive acid. It reacts with the protein of the skin turning the skin yellow.
The acid is itself colourless, but when exposed to sunlight some of the molecules may decompose to form nitrogen dioxide gas, which is brownish gas and hence the concentrated acid appears yellowish when kept for long.The amount of the nitric acid used in the laboratory is 68% by weight.