# Molarity of the solution?

A solution of KCl has a density of 1.011 g/mL and has a concentration of 0.273 M. What is the molality of the solution? The molar mass of KCl is 74.5513 g.

So first I got g of KCl solute:

.273 (m) * 74.5513 (g/m) = 20.35 g

Total mass of solution = 1020.35g (since 1000g solvent per .273m solute).

Now density = mass/volume
Thus, volume = mass/density:

V = 1020.35 (g) / 1.011 (g/ml) = 1009 mL or 1.009 L

Now, molarity = moles/volume

So, M = .273/1.009 = .271

However, the accepted solution says .276 and does the problem differently, so I'm having trouble understanding if this is the correct or incorrect way of doing things.

Thank you!

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chemisttree
Homework Helper
Gold Member
By inspection, I see that a liter of solution must weigh 1011g yet you have 1020.35g using logic that escapes me. You also show that 1000g of solution must have 0.273 moles of KCl ("since 1000g solvent per .273m solute") yet the definition of MOLARITY is M = #moles/L. You are also redefining the concentration by forcing the volume to be 1.009 L yet assuming that this volume must also contain 0.273 moles of KCl.

You seem to be confusing molality with molarity.

Why must a liter of solution weigh 1011 g?

Borek
Mentor
What is the density definition?

OK, so 1.011 (g/mL) = mass (g)/1000 (mL)

Correct?

Ah, chemisttree, you were right. Molality and molarity -__-

Thank you!

Can someone explain why the solution weighs 1011 g?

If I have 1000 g of solvent and 20.35 g of solute, why isn't the weight 1020.35?

Borek
Mentor
Can someone explain why the solution weighs 1011 g?

If I have 1000 g of solvent and 20.35 g of solute, why isn't the weight 1020.35?
You have 20.35 g of solute, but you don't have 1000 g of solvent.

However, you have 1L of the solution of density 1.011 g/mL. That's all you need to know to calculate the solution mass.

And to move a step further - mass of solvent is not 1000 g, but mass of the solution minus mass of the solute.

chemisttree