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 g mL⁻¹?

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Solution

Concentrated nitric acid used in laboratory work is 68% nitric acid by mass in an aqueous solution. This means that 68 g of nitric acid is dissolved in 100 g of the solution.

Molar mass of nitric acid (HNO₃) = 1 × 1 + 1 × 14 + 3 × 16 = 63 g mol⁻¹

Then, number of moles of HNO₃

Given,

Density of solution = 1.504 g mL⁻¹

Volume of 100 g solution =

Molarity of solution

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QUESTION 2.4

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

Nitric acid is 68% by weight and density of solution is 1.504 g/ml

Calculate Molarity

Q. (a) Nitric acid cannot be concentrated beyond 68% by the distillation of a dilute solution of HNO₃. State the reason.
(b) How can 68% nitric acid be concentrated to 98% nitric acid?

Q. How many grams of concentrated nitric acid solution should be used to prepare

250 mL

2.0 M H N O_{3}

. The concentrated nitric acid is

70 % H N O_{3}

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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 g mL−1?

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 g mL⁻¹?