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INTERMOLECULAR FORCES (IMF's):
Last semester, we saw the covalent and ionic bonds that hold molecules and ionic compounds together. Now we learn about the forces that work between neighboring atoms and molecules to hold them together in the solid or liquid state. We will look at four of the primary intermolecular forces. They can be broken up into two main categories: those between charged species and those between neutral species.

1. IMF between charged species:

A. Ion - Dipole Forces

Particles: Ions and polar molecules.

Description: The negative end of the dipole is electrostatically attracted to the cation and the positive end of the dipole surrounds and stabalizes the anion.

Example: KCl (aq) in H2O (l)

Drawing:

Comments: Strength increases as the charge on the ion increases or the strength of the dipole increases. These forces are not as strong as the ionic bond between two actual ions.

2. IMFs between neutral atoms or molecules:

A. Dipole- Dipole Forces

Particles: Require polar molecules.

Description: The molecules orient themselves to minimize electrostatic repulsions and maximize attractions between their dipoles.

Example: HCl (l)

Drawing:

Comments: The strength increases with increasing magnitude of the dipole.

B. London Dispersion Forces

Particles: Exist between all particles. It is the only force between non-polar atoms or molecules.

Description: Temporary fluctuation of electron density create an induced dipole which can create and interact with induced dipoles on neighboring atoms or molecules.

Example: He(l), N2(l), CO2(s)

Drawing:

Comments: Dispersion forces tend to increase with increasing molecular weight, since the more electrons a molecule or atom has, the more polarizable it is. Polarizaility is the ease with which a dipole can be induced in a molecule or atom. In molecules with very different molecular weights, the strength of the London forces tends to be the predominate force when determining properties such as boiling point. If the molecular weights are similar, then the other forces are our primary concern. In other words, London forces are usually the weakest of the IMF, though they become significant for larger molecules. Example: N2 and CO have the same molecular weight so the strength of their London forces should be similar. The predominate force then becomes whether the molecules are polar. Since only CO is polar it's molecules are held together by stronger dipole - dipole forces.

C. Hydrogen Bonding

Particles: A special case of dipole - dipole forces requiring an X-H bond, where X = N, O or F.

Description: Bonding a H atom to a small, highly electronegative element (N, O or F) causes the H atom to be effectively striped of its electron density, resulting in a nearly bare proton. This bare proton is then strongly attracted to the lone pair electron density of the N, O or F from a neighboring molecule.

Example: ammonia, NH3 (l)

Drawing:

Comments: It is not enough to have a H and a N, O or F atom in the same molecule. The H atom must be directly bonded to the N, O or F. Hydrogen bonds are generaly stronger than regular dipole - dipole forces though not as strong as ion-dipole forces.

Try this: Determing the type of IMF exhibited for the following:

A) CH3OH
answer: hydrogen bonding (True, it also has dipole-dipole and London forces, but when asked a question like this, we are looking for the strongest of the forces present.)
B) CH3OCH3
answer: Dipole-dipole The H is not directly bonded to the O.
C) CH4
answer: London forces
D) CH3CH2CH2CH3
answer: London forces Also non-polar, but the London forces are stronger than in CH4, which is smaller and less polarizable.

Try this: rank the species HI, HBr, HF and H2 from lowest to highest boiling point based on the strength of their IMFs

answer: H2 < HBr < HI , HF H2 is non-polar and has only London forces. HBr is polar and should have a higher BP. HI is less polar than HBr, but is larger and has stronger London forces. HF exhibits hydrogen bonding,

All atoms and molecules have these IMF acting to attract them to their neighbors. The ultimate question is whether the particles have enough KE at a given temperature to overcome these IMF and enter the gas phase or whether the IMF "glue" is great enough to hold them in one of the condensed phases.