To use knowledge of polarity to explain why “like dissolves like”.
To apply knowledge of polarity and intermolecular forces to predict properties of chemical compounds.
We have looked at intermolecular forces already; understanding how polarity works is important to understanding intermolecular forces and the dissolving process in chemistry.
A molecule is polar when it contains atoms with a difference in electronegativity, arranged in a way that creates an unequal charge distribution. It can be predicted from the bonds in the molecule:
A polar bond is a bond between two atoms of different electronegativity – this creates an unbalanced charge distribution because more electrons (negative charge) are closer to the more electronegative atom than the less. As we saw in our lesson on intermolecular forces, this is called a dipole because it creates slightly positive and negative charges (δ+ / δ-) like a north/south pole.
The polar bond creates a permanent imbalance where electrons are closer to the electronegative atom – the charge is not equally distributed across the molecule. This causes polarity in a molecule. See the example of CH3Cl below.
In some cases, compounds with multiple polar bonds can still be non-polar because the polarizing effect of equal bonds can cancel each other out. An example is CCl4; C-Cl is a polar bond where Cl has a δ- charge. However, because all of these bonds are equally polar and equally spaced around the central carbon atom, their effects cancel out in all directions and the molecule is actually non-polar. This is what happens to CO2 and to any molecule with all its electron pairs bonding to equivalent atoms equally spaced around the central atom – symmetrical molecules. See the diagram below:
We saw in earlier lessons on solubility that “like dissolves like”, meaning polar solvents dissolve polar solutes, and non-polar solvents dissolve non-polar solutes. This is because of intermolecular forces (IMFs):
Solvents are used in the liquid phase. In the liquid phase, molecules are in close contact with the ability to move, and the millions of identical solvent molecules mean any IMFs occur across all the individual solvent molecules in the container. The solvent molecules will pack together to maximize these attractive forces for stability, whether it is stability from London forces, dipole-dipole interactions or strong hydrogen bonds.
These attractive IMFs have to be overcome by the solute molecules if they are going to spread throughout the solvent molecules and dissolve. To do this, the solute molecules need to have IMFs of a similar strength! If they don’t, the stronger IMFs will not be overcome by the weaker IMFs.
When solute and solvent molecules do have similar strength IMFs, the overcoming and breaking of solvent IMFs is replaced by forming IMFs between solute and solvent molecules. This allows molecules of both to move and interact with one another freely. This will form one homogenous phase of the solute dispersed throughout the solvent - that is a solution!
This means that solvation – the process of forming a solution – can equally happen with two non-polar species or two polar species, but not a polar and non-polar substance.
As mentioned in our “Intermolecular forces” lesson, the relative strengths of intermolecular forces are (strongest to weakest): hydrogen bonding > dipole-dipole interaction > London dispersion forces.
Remember that any molecule which contains electrons experiences London dispersion forces. Generally, the more electrons a molecule has, the stronger its London dispersion forces.
Because they are considerably weaker than dipoles and hydrogen bonding, London forces are generally ignored in molecules that contain stronger forces, because these stronger forces are a lot stronger; they determine properties.
Anything that contains polar bonds that is not symmetrical will contain permanent dipoles and have dipole-dipole interactions.
Hydrogen bonding occurs whenever molecules contain H-F, H-O, or H-N bonds. These are the strongest intermolecular forces – in molecules that use hydrogen bonding, this is the interaction most relevant to finding properties because it is the strongest.
In the same way that ‘polar’ interactions are stronger than nonpolar London forces and can’t be overcome by them, compounds that make these polar interactions need more heat energy to be overcome than molecules with just London forces. This is observed by experiment, in some enormous shifts in melting and boiling point from the change in one atom. Below is a list of related compounds with their intermolecular forces listed and the number of electrons in the compound, to account for London forces:
Compound formula (name)
Strongest IMFs present (# of electrons)
Boiling point (oC)
F2 (Fluorine gas)
London forces (18 electrons)
HF (Hydrogen fluoride)
Hydrogen bonding (10 electrons)
N2 (Nitrogen gas)
London forces (14 electrons)
Hydrogen bonding (10 electrons)
London forces (10 electrons)
Dipole-dipole interactions (42 electrons)
Along with many other experimental results, chemists use evidence like this to determine the strongest intermolecular forces and the extent that they affect properties of chemical compounds.
For example, the difference between the boiling points of ammonia and hydrogen fluoride (despite them having the same number of electrons) is that the hydrogen bonding in HF is stronger and therefore more energy is needed to overcome these stronger attractive forces than is needed to overcome the weaker hydrogen bonds in ammonia. London forces are weak and therefore very little energy is needed to overcome them in F2, N2 or CH4.
Polarity in chemical substances
Apply knowledge of intermolecular forces to predict properties of chemical compounds.
Compare the chemicals in each question and predict which has a higher boiling point. Explain your answer.
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