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Week of 14 January 2008

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Monday, 14 January 2008
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08:10 - This morning I'll make the final few edits that need to be completed before the entire chem lab book goes to production. At that point, I'm finished with the book, other than any final edits I make when I get the PDF galley proofs. The book is on track to appear in the bookstores in April.

Before then, I have a lot to do to bring up the web site and messageboard for the book, as well as a lot of other administrative stuff. I'll be interleaving all that stuff with work on the next book, which is already well in progress although we haven't yet signed a contract for it. And I'm already doing some planning and outlining for the book that will follow that one.


Tuesday, 15 January 2008
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08:59 - The book is now complete as far as I'm concerned, and off to O'Reilly's production folks. In a few weeks, I'll get the PDF galley proofs to review and make minor corrections to, but other than that I'm finished with Illustrated Guide to Home Chemistry Experiments.

I also got the home chemistry lab messageboard/forums stubbed out yesterday and sent invites to my editors and technical reviewers to join it so they'd have coveted single-digit membership numbers. I'll open it to my readers soon, and eventually to anyone who wants to join. I also did some work on the homechemlab.com web site, although at this point I'm just creating pages and getting material incorporated rather than trying to make them look pretty.

UPS showed up yesterday with a 20-pound box of (mostly) chemicals from Elemental Scientific. I haven't checked off the items against the packing list yet, but eyeballing it it looks as though everything is there. The one casualty was a one-ounce bottle of ammonium oxalate. The lid came off during shipping and the bottle was empty. That's not a major problem. I have oxalic acid and aqueous ammonia, so it's easy enough to make the stuff up myself.

As usual, I found myself wondering why I ordered a couple of the chemicals. For example, there's a bottle of bismuth nitrate. I remember ordering it. In fact, I added it to my order after I'd originally submitted it, so I must have had good reason. I seem to remember that it's one component of Dragendorff Reagent for alkaloids. Oh, well. I'll figure it out.


Wednesday, 16 January 2008
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08:35 - It looks like we're in for our first real winter weather of the season. The "wintry mix" is supposed to start falling around midnight and continue through tomorrow afternoon, changing over before daybreak to freezing rain and then to rain in the afternoon. The forecast high temperatures for tomorrow are right at freezing, so anything that falls is likely to be frozen or very nearly so.

I haven't checked, but I'm sure that school is already canceled for tomorrow. Unfortunately, Barbara has a follow-up doctor visit tomorrow morning to check progress after her knee surgery last month.

Because we're in the lee of the Appalachian Mountains, reality is often far different from the forecasts. Sometimes, the forecast is for a lot of frozen precipitation, but the mountains stop it before it arrives here. Other times, the forecast is for a little frozen precipitation, but we end up getting a lot. The most extreme example of that I remember was the time all the weather weasels were agreed that we'd get at most 2" (5 cm) of snow and we ended up getting close to 20" (50 cm). We don't have much in the way of snow plows and other snow removal equipment around here, so even a few inches pretty much brings things to a stop. At least the dogs will have a good time.

Freezing rain always raises the specter of power outages. Even a quarter inch of freezing rain will cause some outages, and if we get an inch the outages will likely be widespread. Fortunately, we have high-output unvented natural gas logs in our den fireplace, so we're able to stay warm even if the power goes out for an extended period.

My guess, though, is that nothing much will happen. The ground is still quite warm, so any frozen precipitation that falls isn't likely to stay around for long.


Thursday, 17 January 2008
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08:50 - We have about half an inch (12.7 mm) of snow on the ground this morning, with a thin coat of ice on top of it. The frozen precipitation is supposed to continue through this afternoon, changing over to freezing rain. The NOAA winter storm warning text notes that this is the first time in several years that most North Carolina drivers have experienced such conditions, so there are likely to be more than a few accidents during the morning commute.

Barbara is off to her doctor's appointment this morning to follow up on her knee surgery. She's recovering pretty well, but has some pain both in the leg that had surgery and in the other leg. I suspect that in accommodating her injured leg she's putting extra stress on the other leg.

To make things easier for readers of Illustrated Guide to Home Chemistry Experiments, I've configured several kits that contain the glassware, lab equipment, and chemicals necessary to do the experiments. A few days ago, I got a call from the supplier I'm working with. One of the chemicals I'd specified in one of the kits was 5 grams of iodine crystals. A few days ago, he emailed me to let me know that the DEA's recent action in moving elemental iodine from List II to List I meant that it was impractical to include iodine in the kit. The paperwork would have been horrendous, and every kit buyer would have ended up on a DEA list of suspects.

My first inclination was to tell him to substitute either tincture of iodine or Lugol's solution for the iodine crystals. (The DEA exempts single transactions of one fluid ounce or less of iodine solutions that contain 2.2% or less iodine.) Although either iodine tincture or Lugol's would have worked for the experiments in the book, neither is ideal, particularly for doing other experiments that aren't in the book.

For example, tincture of iodine is iodine crystals dissolved in ethanol. For most purposes, that'd be fine. But it wouldn't work for the iodoform test, because one of the things that yields a positive for the iodoform test is ethanol. Duh. There are similar objections to Lugol's, which is an aqueous solution of elemental iodine in potassium iodide. For some purposes, the presence of iodide makes Lugol's useless.

So, I decided to take the DEA lemons and make lemonade, which is a polite way of saying screw the DEA. I'm writing a supplemental lab session that will be distributed freely as a PDF. This lab session involves isolating elemental iodine from potassium iodide (which is not listed by the DEA), using the common chemicals hydrochloric acid (muriatic acid from the hardware store), hydrogen peroxide (from the drugstore), and diethyl ether (automobile starting fluid from the auto parts store).

When I told my friend and technical advisor, Dr. Paul Jones, what I was doing, he started to laugh. I asked him what was so funny, and he said that I may end up being single-handedly responsible for the DEA moving potassium iodide and starting fluid to List I. If that sounds ridiculous, remember that at the drugstore you now have to present photo ID and sign to buy OTC drugs that include pseudoephedrine. Some day soon, at the auto parts store you may have to present photo ID and sign to buy ether-based starting fluid.


Friday, 18 January 2008
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09:14 - We got only a half inch or so of snow yesterday, and the main roads were in good shape. Barbara arrived at the doctor's office in time for her 8:40 appointment, only to find that they were running a skeleton crew and that her doctor wouldn't be in until 9:30. She was there, so she waited. She got a generally good report, although the doctor was a bit concerned about her persistent pain. He sent her to the hospital for a scan to rule out a blood clot, which it turns out wasn't the problem.

Today is to be near freezing all day, which with the roads still wet means there'll probably be quite a few minor accidents. Barbara called an hour or so ago to let me know she'd arrived at work safely. She went in early to make up some of the time from yesterday.

I should have listened to Paul Jones, who suggested that I try isolating the iodine by filtration before I used solvent extraction. I didn't do that because I made an assumption that turned to be wrong. I was lead astray by the relative volumes of the reactant and product. The 2.00 g of potassium iodide I started with filled the test tube to a depth of only a centimeter or so when tapped down. The iodine produced by the reaction settled slowly to a depth of perhaps 4 centimeters.

Iodine is considerably denser than potassium chloride, so I assumed that the iodine produced by the reaction was very finely divided, perhaps almost colloidal. That didn't bode well for using filtration, so I did the solvent extraction instead. But, as a belt and suspenders kind of guy, I decided that I'd better repeat the experiment and try using filtration to isolate the iodine, even though I "knew" the iodine was so fine that it'd pass right through the filter paper.

As it turns out, it didn't. The filter paper captured essentially all of the iodine, other than the minor amount that dissolves in the aqueous reaction solution and the rinse water. That's just enough to stain the filtrate a very pale brown color. (It takes about 3 liters of water to dissolve one gram of iodine, so the ~10 mL aliquots I was using dissolved only milligram quantities of the product.)

So, I reshot some images and just finished re-writing the lab to use filtration rather than solvent extraction. Frankly, I'm glad I was wrong. Diethyl ether is nasty stuff. Most people know it's extremely flammable, but they have no idea just how flammable it is. Its autoignition temperature is about 170 C, which means no flame is necessary to ignite the vapors. A hotplate or stove burner can ignite ether vapor, or even a lit incandescent light bulb, which is probably why illegal meth labs blow up and burn down so frequently.


Saturday, 19 January 2008
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09:33 - The weather weasels say we're to be struck by the Great Blizzard of 2008 today. Two days ago, they were forecasting 2 to 4 inches (5 to 10 cm) of snow. Yesterday, the forecast had changed to 90% and 1 to 3 inches. Last night, it was 80% and 1 to 2 inches. This morning, it's 70% and 1 inch. Still, if that much turns up, it'll be a blizzard for us. We haven't had any measurable snow for several years.

I like to tell stories on myself occasionally, because it helps me avoid taking myself too seriously. As anyone who knows me well will tell you, I'm a pompous, arrogant git.

So, yesterday I was working in the lab. I'm writing up some supplemental lab sessions that I'll publish as a freely-distributable PDF file to help promote the book. I'd finished work on the session about isolating elemental iodine, and decided to do one on qualitative analysis of some sugars.

I decided to use three reagents for that lab session: Barfoed's Reagent, Benedict's Solution, and Seliwannoff Reagent. I mentioned in the introduction to the lab session that all of these reagents are available commercially but can also be made up at home. As much as possible, I'm trying to limit the equipment and chemicals I use in these supplemental lab sessions to ones that were used in the book itself or can be obtained easily and cheaply locally, and therein lay a problem.

Benedict's Reagent is a solution of copper(II) sulfate and sodium carbonate, both of which readers will have or can obtain easily. But Benedict's also requires sodium citrate, which isn't readily available other than from specialty chemical suppliers. But it's easy enough to make sodium citrate from sodium hydroxide, which readers will already have, and citric acid, which is sold in drugstores and grocery stores as "sour salt" (despite the fact that it's not a salt at all...).

So, in the "Substitutions and Modifications" section, I included instructions for making up Benedict's with sodium hydroxide and citric acid. For a moment, I felt the satisfaction of making things easier (and cheaper) for readers, but then I started thinking about the Seliwanoff Reagent and Barfoed's Reagent.

There's not a lot I can do about Seliwanoff's. It's a solution of resorcinol in relatively concentrated hydrochloric acid. The hydrochloric acid is no problem. Readers will already have it, or they can use muriatic acid from the hardware store (which is really 10.2 M or 31.45% hydrochloric acid). But resorcinol isn't something you can buy at a hardware store or drugstore, at least not nowadays. So, if readers wanted to do the part of the lab that requires Seliwanoff Reagent, they'd simply have to bite the bullet and order a small quantity of resorcinol. Oh, well.

But then I started thinking about Barfoed's Reagent, which is a solution of copper(II) acetate with a bit of acetic acid added. Readers will have acetic acid, but copper(II) acetate is a specialty chemical. Then I had an ah-hah moment. In Laboratory Session 10.1 in the book, readers will produce copper(II) carbonate, which they later combine with charcoal to smelt metallic copper. But only a small amount of copper(II) carbonate is needed for the smelting session, so most readers would have excess copper(II) carbonate on hand.

Copper(II) carbonate reacts with acetic acid to form copper(II) acetate, which is just what we need to make Barfoed's Reagent. So I rushed down to the lab, thinking I had some unused copper(II) carbonate, still in the filter paper from when I did Lab 10.1. I did! A whole lot of it, in fact. So I scooped quite a bit of the light green copper(II) carbonate powder into a 150 mL beaker, added 100 mL or so of distilled water, swirled the beaker to suspend the insoluble copper(II) carbonate, and added a few mL of glacial acetic acid.

Everything worked just as I expected. The solution began fizzing as carbon dioxide gas was evolved by the reaction. The solution turned from a light green chalky appearance to the strong blue color of copper(II) acetate, and the light green solid disappeared. But as the light green solid disappeared, it was replaced by a dark blue solid. The solution became noticeably viscous, with chunks of dark blue junk suspended in it. Crap.

My first thought was that perhaps I hadn't washed the copper(II) carbonate thoroughly enough, so it might be contaminated with the sodium sulfate that was the other product when I precipitated the copper(II) sulfate solution with sodium carbonate to produce the copper(II) carbonate. I couldn't think of any obvious reason why a little sodium sulfate would have the effect of turning the copper(II) acetate solution that I expected to be clear blue into a viscous mess with embedded blue chunks, but neither could I think of anything else that might cause the problem.

So, I transferred the remaining copper(II) carbonate from the filter paper to another 150 mL beaker, added about 100 mL of distilled water, swirled it, and allowed the insoluble copper(II) carbonate to settle out. I decanted off the supernatant liquid, which presumably contained any excess sodium sulfate. Just to be safe, I repeated the wash twice more, ending up with a soupy mass of (presumably) reasonably clean copper(II) carbonate.

I then repeated my earlier steps, adding about 100 mL of distilled water to the copper(II) carbonate, swirling the beaker to suspend the solid matter, and adding a few mL of glacial acetic acid. I was completely surprised when, you guessed it, I ended up with a viscous dark blue solution with chunks of dark blue solid suspended in it. Arrrrghhh.

Obviously, there must have been something wrong with the copper(II) carbonate I was using, or perhaps the acetic acid had somehow rotted or something. At any rate, this soupy blue crap with blue chunks was obviously useless, so I dumped the contents of the first beaker down the drain with plenty of water, rinsed that beaker, and picked up the second beaker full of soupy blue crap with blue chunks. I dumped the contents of that beaker down the drain and started rinsing it. Just as the last of the bluish water was disappearing down the drain, I had a shit-Shit-SHIT moment.

What is the solubility of copper(II) acetate? I ran upstairs and pulled out the CRC handbook. As it turns out, the solubility of copper(II) acetate in water at room temperature is only about 7.2 g per 100 mL. SHIT! There was considerably more than 7.2 g of copper(II) carbonate in each of those beakers, so OF COURSE it precipitated out. Duh.

In a telephone conversation with Paul Jones later that day, I told him what I'd done. He just laughed, but it wasn't at my expense. He laughed, I'm sure, because he's had his share of shit-Shit-SHIT moments just like mine. So has every chemist.

So, I went back to the drawing board. As it turns out, Barfoed's Reagent is really just a nearly saturated solution of copper(II) acetate with 1 mL of glacial acetic acid added per 100 mL. That made it easy to write an alternative procedure for making up Barfoed's, which I did.


Sunday, 20 January 2008
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Copyright 1998, 1999, 2000, 2001, 2002, 2003, 2004, 2005, 2006, 2007, 2008 by Robert Bruce Thompson. All Rights Reserved.