Category: science kits

07:55 – We’re all caught up on shipping kits. The only orders outstanding are the ones that came in overnight and this morning, which we’ll ship this afternoon. Meanwhile, I need to make up a bunch of solutions today and get started on bottling them.

One of those solutions is 4 liters of 6M sodium hydroxide, which has gotten me thinking about chemical storage. I’m down to my last three 500 g bottles of sodium hydroxide. When I finish those, I have to open a new container of sodium hydroxide, which in this case is a 10 kilo bucket rather than a 500 g bottle. Right now, that 10 kilo bucket is sitting on the floor because it won’t fit my storage shelves.

When we got started building science kits a few years ago, I put up shelves for chemical storage. Most of them are 4″ (10 cm) wide with vertical separation of 6″ (15 cm). Those worked fine when I was buying chemicals in 25 g, 100 g, and 500 g bottles. They’re not wide enough now that I’m buying a lot of chemicals in 1-kilo, 2- or 2.5-kilo, 5-kilo, and 10-kilo containers. That’s why there are still a couple of cartons of chemicals from Fisher Scientific sitting on the floor where UPS delivered them. I thought about repackaging them into 500 g and one kilo bottles, but that’s just too much work. Instead, I think I’ll remove some of the smaller shelves and replace them with wider shelves with more vertical separation. But that’ll have to wait for things to calm down a bit around here.

11:02 – I’d forgotten how obnoxious lead acetate is. We provide a 0.1 M solution of lead acetate in many of our kits, and I was just making up four liters of the stuff. I weighed out the appropriate mass of reagent-grade lead acetate and added it to distilled water. One might expect a nice, clear water-like solution to result. Instead, one gets a solution that looks like milk, literally.

The problem is that most common lead salts, with the exceptions of the acetate and the nitrate, are extremely insoluble in water. And water exposed to air just loves to suck up carbon dioxide. At room temperature, a liter of water dissolves about 1.6 grams of carbon dioxide. That doesn’t sound like much, but with the molar mass of carbon dioxide about 44 g/mol, that means that plain water exposed to air is actually about 0.036 molar with respect to carbon dioxide. That carbon dioxide reacts with water in a reversible reaction to form carbonic acid, the acid whose salts are carbonates. And lead carbonate is extremely insoluble in water, which is why my solution looks like milk. That 0.036 molar carbonic acid reacts 1:1 with my 0.1 molar lead acetate solution precipitating out nearly a third of the lead ions as insoluble carbonate. What’s worse is that that reaction removes the carbon dioxide from the solution, so it promptly sucks more carbon dioxide out of the air, until all the lead is precipitated and the solution reaches equilibrium with about 1.6 g/L of dissolved carbon dioxide. Basically, my dilute solution of lead acetate eventually turns into a dilute solution of acetic acid with most of the lead precipitated out as lead carbonate.

Fortunately, one can use Le Chatelier’s principle to shift the equilibrium by dissolving the lead acetate in a dilute solution of acetic acid rather than plain water. Although it’s a weak acid in absolute terms, acetic acid is a much stronger acid than carbonic acid. That forces the equilibrium of the reversible carbon dioxide <-> carbonic acid reaction to the left, keeping the dissolved carbon dioxide in the form of the molecular gas rather than the carbonate ion. And the lead acetate remains in solution as lead acetate.

10:41 – Barbara is ironing and cleaning house while I do kit stuff. Right now, I’m dehydrating some magnesium sulfate heptahydrate to anhydrous form. The hydrated form, AKA Epsom salts, is cheap, a couple bucks a kilo in USP/FCC form, which is as pure as reagent grade. Buying the anhydrate from a chemical supplier runs $80 to $100 per kilo, which is outrageous.

So I just spread about a kilo of Epsom salts in a large casserole dish and stuck it in the oven at 500F (260C) for an hour. That removes most of the water of hydration and forms a thin glassy layer of magnesium sulfate. I break that into chunks and toss it into a blender that I reserve for such work. I then blend it on high to break up the chunks into mostly powder, run it through a flour sifter, repeat as necessary until all the chunks are broken up, and then put the powder back in the oven for another hour at 500F to finish drying it out. I do this while I’m doing other stuff, so the whole process requires maybe ten minutes of my time. Add the cost of my time to maybe two or three bucks in materials cost and electricity, and I end up with half a kilo of magnesium sulfate anhydrate for less than $20, even billing myself at $100/hour.

The price of many chemicals has gotten ridiculously high. For example, we use copper(II) acetate to make up Barfoed’s reagent, which is essentially a 0.5 molar solution of copper(II) acetate with 10 mL of glacial acetic acid added per liter. I was about to order some copper(II) acetate, but found my regular supplier wanted $120/kilo. Geez.

So, the next time I need to make up Barfoed’s reagent I’ll do it from scratch on the fly. I generally make up four liters at a time, so I’ll start with two clean 2-liter Coke bottles. I’ll transfer two moles of copper(II) sulfate to one bottle. That copper(II) sulfate is from Home Depot, which sells a 2-pound (907 g) bottle of the stuff for about $10 under the name of Root Kill. The assay on the bottle says it’s 99% copper(II) sulfate, which I’ve verified gravimetrically. The remaining <1% by mass is mostly insoluble copper oxide. The molar mass of copper(II) sulfate is 249.68 g/mole. Dividing that by 0.99 gives 250.22 g/mole, so I'll transfer 500.44 g of the Root Kill to the two liter bottle and dissolve it in hot water. (It dissolves quickly in hot water; in room temperature water it can take literally a week to dissolve.) I'll then filter the resulting two liters of pretty blue solution into the second bottle, rinse out the first bottle, and divide the solution with one liter in each of the two bottles. So far, I'll have used up maybe five minutes of actual working time and about $5 worth of the Root Kill.

I’ll then add either sodium carbonate or sodium bicarbonate, both of which are cheap, to precipitate the copper ions as insoluble copper(II) carbonate. Once the precipitate settles, I’ll decant off the supernatant liquid, which contains mainly sodium sulfate with a small amount of the excess sodium carbonate or bicarbonate in solution. If I decant 90% of the supernatant liquid and refill the bottle with tap water, I’ve diluted the original level of soluble contaminants to 10% of what they were. Repeating that process a few times, ending with a distilled water wash, reduces the soluble contaminants to 1%, 0.1%, 0.01%, and finally 0.001%, which is better than good enough.

I don’t even need to filter out the copper(II) carbonate and dry it. I can simply wet it with a liter or so of distilled water and add glacial acetic acid stoichimetrically to convert the copper(II) carbonate to copper(II) acetate in situ, add an extra 20 mL of the glacial acetic acid, and then bottle the resulting Barfoed’s reagent.

And don’t get me started on ammonium metavanadate. The last time I bought it, maybe three or four years ago, I paid something like $15 for a 25 gram bottle. I thought $0.60/gram was pretty high then, but that’s now tripled to nearly $2/gram, and that doesn’t even include the required poison-pack container and hazardous shipping surcharge. Geez. I can synthesize the stuff from scratch here for something like $0.05/gram, and it’s no more difficult than the copper(II) acetate synthesis.