Lecture #28
Text: Chapter 16 sections 1,2
These "slides" represent highlights from lecture and are neither complete nor meant to replace lecture. It is advised not to use these as a reliable means to replace missed lecture material. Do so at risk to healthy academic performance in 09-105.
Lecture Outline Intermolecular Interactions (condensed phases)

Induced dipoles

Polarizability

Van der Waals interactions

Permanent dipoles

Temporary (induced) dipoles

Boiling points, melting points, densities

Effects of size

Effects of shape
Previously, we looked at Van der Waals' equation that folded intermolecular interactions into the ideal gas picture. The "a" parameter was a measure of the strength of the attractive forces.
The instantaneous dipole in one molecule induces a temporary dipole in the nearby molecule.
The ease with which a valence electron cloud shape is distorted is determined by the atom or molecule's "polarizability".
Comparison of various interaction strengths
Comparing chemical bond formation with van der Waals interaction
Melting points and boiling points should correlate with strengths of intermolecular attractions.
Boiling points of the inert gases.
Boiling points of (linear) hydrocarbons
A graph of the smooth dependence of (linear) hydrocarbon boiling point on size.
London forces are always present.
A comparison of linear to branched C5H12
A ball and stick picture of the linear isomer.
A space filling model of the linear isomer.
Van der Waals interaction between two linear isomers.
Space filling model of the branched hydrocarbon.
Neighboring branched hydrocarbons interacting through the van der Waals interaction.
What's expected for cyclopropane?
Examples of London forces effects.
Examples of London forces and dipole forces effects
Does this comparison suggest a contradiction?
Boiling point trends and predictions for water and for HF.
Obviously, you know water's boiling point is way off following the prediction shown.
The explanation is a extraordinary phenomenon called hydrogen bonding.