|Text: Chapter 16 sections 1,2||
|Examples of London forces and dipole forces effects|
|Competition between induced dipole effect and permanent dipole effect must be considered for this set.|
|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.
|Participants on either side of a "hydrogen bond" must almost invariably be F, O, or N.|
|The strength of hydrogen bonding relative to the other interaction strengths we've been looking at.|
|Hydrogen bonding in water (ice).|
|Boiling point effects owed to hydrogen bonding|
|Multiple hydrogen bonds|
|A comparison of melting points of two similar looking compounds, both are benzene rings with "hydroxy" and "nitro" substituents in different positions. For comparison, below on the left, the "hydroxy" group has been replaced by a "methyl" group.|
|Melting points are also affected by hydrogen bonding.|
|The melting points of these two structural isomers of nearly identical electronic volumes differ substantially. Why?|
|An illustration of an "intramolecular" hydrogen bond within a molecule (rather than between molecules).|
|Less interesting, chemically are salt solubilities. For ionic compounds, as shown here, the attractive forces are electrostatic and increase with the amount of charge, like 2+ vs 1+, and decrease as opposite charges get further apart, as when larger ions are used. The last salt shown is actually liquid at room temperature and is an example of an ionic liquid, a newly discovered class of salts that have tremendous applications anticipated.|