Solubility of a solute can be determined by comparing the polarities of the solute and the solvent molecules. Polar solvents will dissolve polar solutes and nonpolar solvents will dissolve non-polar solutes due to the presence of similar intermolecular forces.Continue Reading
Covalent molecules have three types of intermolecular forces: van der Waals, dipole-dipole interactions and hydrogen bonding. Van der Waals forces are the weakest of all intermolecular forces and are found in nonpolar molecules. Dipole-dipole interactions are found in polar molecules containing a partial positive charge and a partial negative charge due to the difference in electronegativities of the atoms in the molecule. The partial positive charge of one polar molecule attracts the partial negative charge of a different polar molecule. Hydrogen bonds are very strong dipole interactions that occur in molecules where hydrogen is bonded to very highly electronegative atoms such as fluorine, nitrogen, oxygen or chlorine. Hydrogen bonds are the strongest of all three intermolecular forces.
Polar solvents have molecules that can engage in dipole-dipole interactions or hydrogen bonding. When a polar solute is added to the solvent, the solute is able to disrupt the interactions and disperse itself among the solvent molecules. In these cases, the partial positive charge of a solvent molecule attracts the partial negative charge of a solute molecule. However, when a non-polar solute is added to a polar solvent, the van der Waals forces present in the solute are too weak to disrupt the interactions between the solvent molecules, thus leaving the solute particles bonded together and insoluble.
A non-polar solvent can dissolve nonpolar solutes because both the solvent and the solute have weak van der Waals forces. However, when a polar solute is added to a nonpolar solvent, the dipole interactions between the solute molecules is stronger than the van der Waals forces of the solvent and therefore will not separate in solution, making the solute insoluble.Learn more about Solutions & Mixtures