|
Daynotes
Journal
Week of 14 January 2008
Latest
Update: Saturday, 19 January 2008
09:33 -0500 |
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.
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
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
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.
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
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.
00:00
-
Copyright
© 1998,
1999, 2000, 2001, 2002, 2003, 2004, 2005, 2006, 2007, 2008 by Robert
Bruce
Thompson. All
Rights Reserved.