Hi Dr. Bob, I need some clarification about the concepts of Freezing Point Depression and Boiling Point Elevation. I understand when it's a dissociation of solutes, the ions fully dissociate. In this instance, is it true that we should take into account of the numbers of moles or ions? And which one should we do so (ions vs mole)? And should we distinguish the instances when we take the number of mols into account by whether or not the substance is a solid or not? so if it's a liquid or aqueous substance (or even substances that don't dissociate fully), we can disregard the number of mols and just use deltaTf=Kfm (or kb)? Thank you so much!

Also, another question in regards to the intermolecular forces. I just want to check some of my understanding. Am i correct in saying stronger forces = greater boiling point = more soluble in water? And in the instances of having to compare the boiling points or the solubilities of two substances with the same type of intermolecular force (for example two H-bonds), how could we compare them? Is it based on molar mass of the substance? An example i could find is question 2 on our Exam 1, between 1.0 m KNO2 and 1.0 m K2SO4. Since they are both in aqueous solutions, aren't they both ion-dipole? How could we know then that K2SO4 has a higher bp than KNO2? Thanks again.

Mariana.....Both of your questions are related to the effect of solute dissociation on colligative properties. In general, colligative properties of solutions depend on the number of moles of solute particles (molecules, ions, etc.) in solution. Strong electrolytes like KNO2, K2SO4, etc. are completely ionized so, e.g., a 1 m solution of K2SO4 contains 3 moles of ions so a colligative property like BP elevation will be about 3 times greater than a 1 m solution of a non-electrolyte like CH3OH. A 1 m solution of KNO2 or KNO3 or HNO3 (all strong electrolytes) would be about 2 times greater.