{"id":554,"date":"2012-04-21T10:38:02","date_gmt":"2012-04-21T14:38:02","guid":{"rendered":"http:\/\/www.ttgnet.com\/journal\/?p=554"},"modified":"2012-04-21T10:38:02","modified_gmt":"2012-04-21T14:38:02","slug":"saturday-21-april-2012","status":"publish","type":"post","link":"https:\/\/www.ttgnet.com\/journal\/2012\/04\/21\/saturday-21-april-2012\/","title":{"rendered":"Saturday, 21 April 2012"},"content":{"rendered":"

10:37 –<\/span> Barbara is working on science kits today. I’m doing laundry and working on the forensics book.
\n<\/span><\/p>\n

I should learn to trust myself. One of the solutions included in the chemistry kit is a sodium salicylate solution that’s 200 ppm with respect to salicylate. It’s used as a standard in a lab session about quantitative determination of salicylate in urine by the Trinder reaction. (Ferric ions react with salicylate ions–actually, almost any phenolic ions–to form an intense violet complex; the intensity of the color is proportional to the concentration of salicylate in the specimens.)
\n<\/span><\/p>\n

Unfortunately, that dilute salicylate solution is ideal for mold growth, so we have to add a preservative. As I was making up the solution, referring to the instructions I’d printed months ago, I noticed that they mentioned adding a few crystals of thymol as a preservative. That couldn’t be right. Like salicylate, thymol is a substituted phenol, so thymol should also react with ferric ions. So, I thought, dumping an arbitrary “few crystals” of thymol into the solution would ruin its carefully-measured accuracy.
\n<\/span><\/p>\n

I’d almost decided to substitute a few drops of chloroform as the preservative, but I decided to check just to make sure. So I added one crystal of sodium salicylate to a spot plate, followed by a drop of iron(III) nitrate. As expected, the violet coloration was so intense it looked almost black. I added one crystal of thymol to a second well, followed by a drop of iron(III) nitrate. There was absolutely no visible change.
\n<\/span><\/p>\n

I really should learn to trust myself. I don’t remember doing it those months ago, but there’s no way I would have added thymol to a salicylate standard solution without checking. What occurred to me now would also have occurred to me then.
\n<\/span><\/p>\n

Speaking of organic chemistry, I was talking to Paul Jones yesterday and he mentioned that he’d be grading take-home finals for his organic students. Then he mentioned something about the exam focusing on syntheses. I thought what he described sounded a bit much for first-year organic, so I mentioned that to him. It turns out that Paul isn’t teaching first-year organic this year. The final he was talking about was for his grad students. I told him that was a relief, because I was just thinking about asking him if I could take the final, just to see how much (if any) I remembered. Paul was amused. He said I probably wouldn’t do very well on this final, because the synthesis reactions used hadn’t been invented when I was taking first-year organic 40 years ago. (Take that, old guy!) In fairness, Paul did say that the reactions I learned back when dinosaurs roamed the earth typically actually worked pretty well, while the ones his students were learning were sometimes a bit shaky.
\n<\/span><\/p>\n

\n","protected":false},"excerpt":{"rendered":"

10:37 – Barbara is working on science kits today. I’m doing laundry and working on the forensics book.\n<\/p>\n

I should learn to trust myself. One of the solutions included in the chemistry kit is a sodium salicylate solution that’s 200 ppm with respect to salicylate. It’s used as a standard in a lab session about quantitative determination of salicylate in urine by the Trinder reaction. (Ferric ions react with salicylate ions–actually, almost any phenolic ions–to form an intense violet complex;<\/p>\n

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