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This is a handout which I use in an advanced general chemistry course, but which could be used in an inorganic course as well. It is a mini-periodic table with common cations and their charge to size ratios expressed as Q/r2, where Q is in integer charges (+1, +2), and r is in Angstroms. Conveniently, Na+ is an easy to remember 1.0, and Al3+ and Be2+ are easy to remember values of 10. This corresponds to the polarizing power of these ions, and is a crude proxy for how covalent their interactions with a given anion tend to be. Values above 10 are excluded as being "essentially covalent" (the proton has a value of 4.4 billion!). I also tie this to HSAB and solubility rules. The larger the value of the cation, the less soluble it will be with a given ion unless the ion is unusually soft.
| Attachment | Size |
|---|---|
| Polarizing Power of Common Cations.doc | 41 KB |
A student should be able to articulate, using this chart, why some of the upper main groups are so covalent (H, Be, B, Al), why covalency increases as you move to higher charges, why the solubility rules are what they are (both the hard-hard insolubilities and the soft-soft insolubilities), and why the very soft metals to the right of the table are so unusual.
This requires a knowledge of Lewis acids and bases, as well as a good command of Coulomb's Law. Students like that it explains the solubility rules.
Evaluation
This fails to capture why the alkali metals are not very soft, and they seem to want to explain all of these trends in terms of the hard-hard ionic & covalent interactions. Students have a real hard time with the idea that as ionic forces increase, the bonding becomes more covalent.