For dilute aqueous solutions in which the density of the solution is roughly equal to that of the pure solvent, the molarity of the solution is equal to its molality. Show that this statement is correct for a 0.010 M aqueous urea [(NH2)2CO] solution.

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Respuesta :

Mergus Mergus · Answer 1 · 2019-08-07T04:41:04+02:00

Explanation:

Given: Molarity = 0.01 M

Let the volume of the solution = 1 L

Hence, moles = 0.01 moles (Molarity*Volume)

Molar mass of urea = 60 g/mol

So, mass of urea = moles x molar mass = 0.01 moles x 60 g/mol = 0.6 g = 0.0006 kg ,

Given: Density of solvent = Density of solution = 1 Kg/liter (Water).

So,

The mass of solution = vol x density = 1 L x 1 kg/L= 1 kg

Also,

Mass of solution = Mass of solute + Mass of solvent

1 kg= 0.0006 kg + mass of solvent

Mass of solvent = 1-0.0006 = 0.9994 kg

Molality is the moles of solute present in 1 kg of the solvent. So,

Molaity = ( 0.01/0.9994) ≈ 0.01 m

Molarity = Molality

Hence proved.

pstnonsonjoku pstnonsonjoku · Answer 2 · 2022-03-25T14:20:30+01:00

From the calculation, the molarity of the solution is equal to its molality.

Molality refers to the umber of moles divide by the mass of the solvent.

We know that the molar mass of urea is 60 g/mol. Let us arbitrarily set the volume of solution at 1L.

Since molarity of solution = 0.010 M, number of moles of urea = 0.010 moles

Mass of urea = 0.01 moles x 60 g/mol = 0.6 g = 0.0006 kg

If the solvent is water, mass of solvent = 1 Kg/L * 1 L = 1 Kg

Molality = number of moles/mas of solvent in Kg = 0.010 moles/1 Kg

= 0.010 m

Hence, the molarity of the solution is equal to its molality.

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