Friday, 13 July 2012

10:03 – Another lab day today. As usual, I’ve saved the most obnoxious solutions for last. Stuff like concentrated acetic acid, ammonia, hydrochloric acid, and so on. Very concentrated solutions of some chemicals are called “fuming”, as in “fuming nitric acid” or “fuming sulfuric acid”. No one talks about “fuming acetic acid” or “fuming ammonia” or “fuming hydrochloric acid” because the usual concentrations already emit noxious fumes, and those fumes are sufficient to knock your socks off. That’s one major reason I decided to include 6 M solutions of those chemicals in the kits. It would actually have been easier to provide standard concentrated versions–17 M acetic acid, 15 M ammonia, and 12 M hydrochloric acid–but the 6 M solutions are usable for our purposes and the fumes are a lot less obnoxious.

I also need to start getting solutions prepared for the forensics kits. Several of those are hazardous–three or four are basically concentrated sulfuric acid with minor additions–but at least none of them are particularly obnoxious in terms of fumes.

14:05 – I just made something I didn’t know existed: iodine syrup. I was making up three liters of IKI (iodine/potassium iodide) solution that’s 0.1 molar with respect to both iodine and iodide. That meant I needed 38.07 g of iodine and 49.80 g of potassium iodide.

Now, the thing is, iodine is almost insoluble in water. But in the presence of equimolar or more iodide ions, each iodine molecule bonds with an iodide ion to form a triiodide ion, which is extremely soluble in water. But the speed of dissolution depends on the concentration of the iodide ion. If I’d simply dissolved that 49.80 g of potassium iodide in about three liters of water, added the 38.07 g of iodine, and made up the solution to three liters, the iodine would have dissolved. Eventually. It might have taken literally a month to dissolve, but it would have dissolved.

But iodine dissolves very quickly in a concentrated iodide solution, the more concentrated, the better. Potassium iodide is extremely soluble. At room temperature, that 49.80 g of potassium iodide will dissolve in about 36 mL of water. So I weighed out 49.80 g of potassium iodide in a glass 250 mL beaker, added just enough DI water to dissolve the salt, and then added 38.07 g of iodine. As far as I could tell, the iodine crystals dissolved instantly. I say as far as I could tell, because the liquid in the beaker instantly turned an opaque black, so opaque that I couldn’t see any light through the liquid even holding the beaker up against an overhead fluorescent tube and looking through the bottom of the beaker and a couple centimeters of liquid.

If there’d still been iodine crystals in the bottom of the beaker, I wouldn’t have been able to tell. I tried tilting the beaker back and forth to see if any crystals were visible on the bottom of the beaker, but there weren’t. Still, on general principles, I kept swirling the beaker for a few seconds every minute for ten minutes or so. That’s how I discovered that there is such a thing as iodine syrup. The stuff was viscous, kind of like vegetable oil. Not surprising, I guess, with almost 90 g of solids dissolved in maybe 50 mL of water.

Finally, I decided to give it go, so I carefully poured the liquid into a 1 L volumetric flask, which was the largest I have. I was kind of expecting iodine crystals to reveal themselves in the bottom of the beaker, but there weren’t any. So I made up the solution to 1.0 L and transferred it to the storage container, adding two more 1 L flasks’ worth of water. My volumetric flasks are calibrated to-contain rather than to-deliver, but I know from previous tests that the flasks actually deliver about 999+ mL. So I added just enough water to the storage container to make it up to exactly 3.0 L, give or take a mL.