Today’s lecture discussed how to name these classes of compounds: RX, ROH, RSH, ROR, RSR. An important point: the complexity of naming/drawing problems will be limited to compounds containing a single principal group.
Slides for today’s lecture are located here.
I just came across this in the Editor’s Choice section of Science magazine (Oct 2, 2009, p. 19):
Getting a Grip on Nitrogen
Chirality is associated more with carbon than with nitrogen centers, as the latter atoms tend to invert their configurations fairly rapidly on account of their unbonded electron pair. Recently, relatively slow nitrogen inversion was observed in cyclic oligomers of four or six … amino acids … on account of hydrogen bonding among the substituents. Mocquet et al. now show that swapping in a single analogous [amino acid] residue that is chiral at carbon disrupts the collective inversion mechanism and thereby dramatically stabilizes the chrial nitrogen conformations throughout the ring.
Notice how the “editor” misuses the word “conformation” when referring to nitrogen “configuration”? Oh well. Probably a physicist.
This was forwarded to me today from the EHS office. Acrylonitrile’s formula H2C=CH-CN. Nerd humor based on a CN triple bond.
“Acrylic” is a common textile material. It is favored for its wool-like
feel and commonly used in socks. “Acrylic’ is really polyacrylonitrile.
Polyacrylonitrile is a nice way of saying polyvinyl cyanide. Do you
think anyone would buy socks made from polyvinyl cyanide? If you knock
some socks off the shelf in your local discount store, do not run up to
the clerk and announce, “There is a polyvinyl cyanide spill in aisle
10′.
Enjoy Fall Break!
A great question came my way right after lecture: can resolving agents by achiral?
The short answer is no.
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Slides for today’s lecture are located here.
I’ll be updating this page with information on the rotation barrier in HO-OH later today. Let me know if you would like links to information about conumption/manufacture of polymers, plastics, and petroleum.
I thought I was setting the record straight on energy yesterday, but upon reflection, I made an egregious mistake. Well, I probably made several, but there’s only one that I’m currently aware of.
In drawing a distinction between free energy (G) and enthalpy (H), I unintentionally conflated enthalpy with potential energy. Fortunately, the situation is easily corrected.
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On p. 146 Loudon asks “why an sp2-sp3 C-C bond is stronger than an sp3-sp3 C-C bond”. It’s a good question because it might help explain why more substituted alkenes are more stable than their less substituted isomers, a topic that was covered in today’s lecture. Unfortunately, the explanation provided by Loudon strikes me as incorrect.
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Slides for today’s lecture are located here. The table of common acidic functional groups and related properties are located here (note: you need to memorize the information in this table for next Friday’s exam + future exams). The table will also be included with the study guide for chapter 3 whenever that gets posted (hoping to have it ready in the next 24 hours).
Loudon mentions two van der Waals forces in Chapter 2, an attractive force and a repulsive force. This is confusing because, other than the fact that both forces have something to do with electrons, they are not related.
The attractive van der Waals force is described nicely in Figure 2.8. It appears when the distance between two atoms (or molecules) is short, but not too short (if the atoms’ space-filling models are just touching or are separated by a small gap, then that’s perfect). The force is created by temporary imbalances in each atom’s electron distribution caused by electron movement around the nucleus. The imbalances produce small temporary electrostatic fields and two atoms will naturally correlate their imbalances (fields) so that they attract one another. The attractive force is very weak.
The repulsive force is another beast entirely. It suddenly replaces the attractive force when the distance between the atoms is less than the sum of the atoms’ van der Waals radii. The repulsive force is very strong and we blame it on Pauli repulsion, the tendency of same-spin electrons to avoid each other (or “die” trying).