CR-Scientific Minerals & Experimental Science newsletter

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August / September 2003
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In this issue:

I. Action of Enzymes on Ethylene Glycol
II. Manganese Test Revisited
III. An Exercise in Periodic Law
IV. Site News


I. Action of Enzymes on Ethylene Glycol

A.) Background
Ethylene glycol is commonly used in antifreeze. Its ultimate toxicologic destiny in humans and other animals is oxalate, which precipitates in the presence of Ca++ ions forming calcium oxalate and causing severe organ damage (Amdur, et al., 1991). Before it meets its metabolic fate as oxalate, however, ethylene glycol is converted in the liver to glycoaldehyde and then glycolic acid (1991). It is glycolic acid that the following experiment tries to detect.
There's nothing ground-breaking about the following experiment; it's just something cobbled together both from memory and from available sources (noted, of course, where used). Ethylene glycol is poisonous and must never be ingested or left where someone or something could ingest it.
It's common biochemical procedure to take some cells, either animal or plant, and extract the enzymes for experiments. The cells are ruptured (a high-speed blender is a common choice) and the contents which were once neatly compartmentalized or otherwise kept away from each other form a slurry in which everything is thrown together; enzymes find themselves in contact with their inhibitors, other enzymes1, cell membrane (or cell wall) pieces, structural proteins, salts, metabolic intermediates, and more. It's always best, therefore, to keep the homogenate as cool as possible so the enzyme in question isn't ruined before we can experiment with it. An ice bath works well for storing any containers that have your enzyme preparation in them.
Some compounds are poisonous because they "fit" into the active site of an enzyme which is made primarily to take something else. The compound may either inhibit an enzyme necessary for normal functioning of an organism, or else the compound may be made into something much more toxic than it was. Ethylene glycol is in this latter category; it is not in the diet of any normal animal (and shouldn't be!), but in the unfortunate event that it finds its way in there, its shape and other molecular characteristics make it fit in several enzymes to the point that they'll act on it and chemically change it-- to the ultimate detriment and possible death of whatever organism is involved.

B.) Experiment

Supplies: Fresh liver; Ethylene glycol; Test tubes; High-speed blender; Centrifuge; Ice bath; Beaker; Droppers; Spot Plate.
Optional: Spectrophotometer.

General Procedure: The following should be done with pre-chilled glassware if possible; better yet, conduct the homogenization in a very cold room.
First, the liver is homogenized in a high-speed blender and poured into centrifuge tubes. Second, the homogenate is spun long enough in the centrifuge, preferably in a refrigerated room, to get the larger solids out of suspension. Third, the suspension is poured off into a beaker, which is kept in an ice bath to keep it cool.
This experiment is just qualitative, so we're not going to need micropipettors. First, add about three drops of chilled 1:10 liver homogenate in water to a clean spot plate well. Add one drop of 1:10 ethylene glycol in water and mix gently. Let this stand for anywhere from 30 seconds to 5 minutes. Try different intervals before adding the next reagent. The idea is to develop enough product that the colorimetic reagent will react with it and yield a strong color to the unaided eye in a spot plate well.
According to the method of Snell and Snell (1953) we'll add an equal volume of 1:9 sulfuric acid in water to acidify the mixture in preparation for the glycolic acid assay. This should also severely degrade or even halt the action of the enzyme(s) on the ethylene glycol, but we needn't worry about whether it's complete (we're not doing a quantitative assay). Next, add 5 mg of coarse magnesium metal filings and let the mixture stand for about an hour (1953).
The colorimetric reagent itself is a 0.01% solution of 2,7-dihydroxynaphthalene in concentrated sulfuric acid; this will not produce false results with acetic, tartaric, benzoic, succinic, citric, or unreacted oxalic acid (Snell and Snell,1953), though it seems wise to run a "blank" in which no ethylene glycol has been added.
The procedure of Snell and Snell (1953) then recommends adding 2 mL of the dihydroxynaphthalene reagent if the assay mixture is about 0.2 mL; so we'll use 10 parts reagent solution for every 1 part assay mixture. The resulting liquid is then heated in a boiling water bath for about 20 minutes; afterward, 4 mL of 1:18 sulfuric acid in water is added and the mixture is stirred until the heat subsides.
A deep red color announces the presence of glycolic acid (one of the enzymatic breakdown products of ethylene glycol; more specifically, a product of a product).
Try the experiment on ethylene glycol itself and see if there's a reaction-- there shouldn't be, unless glycolic acid is present as an impurity or additive. If 2,7-dihydroxynaphthalene is unavailable, try adding some calcium ions to the solution (there will already be some from the ruptured cells, but you could add an excess) and looking for a white precipitate under a compound microscope. If there is little or no white precipitate without ethylene glycol but a definite precipitate after the homogenate has been allowed to stand for some time with the glycol, this is a sure sign that the enzymes have been converting the glycol into oxalate.



II. Manganese test revisited

Ed Jurczak recently wrote in with a question about the qualitative test for manganese in minerals:
On [Orsino C.] Smith's Page 67, the manganese test (Ammonium Phosphate test) is supposed to give a pink precipitate, which I can't seem to get... it's white; although when the precipitate is treated with soda and KNO3 on charcoal it does give a pretty good bluish-green bead (which is what you stated in the Feb-March CR-Scientific newsletter). Why isn't the precipitate from the Ammonium Phosphate test turning pink?
I wrote back, trying to figure out why no pink precipitate was appearing. My response started with grabbing at straws: was it possible that the manganese had become oxidized from Mn++ to the colorless Mn+++? Even as I suggested this as a possibility, I considered it unlikely because Mn(III) is unstable and disproportionates into Mn(II) and Mn(IV) pretty quickly in solution. Mn(III) would not have been present at the outset, either; Ed said he was using MnSO4 in place of a mineral sample.2
Another possibility I suggested was to make sure the precipitate was being viewed in a bright, full-spectrum light. Perhaps the color really was there but wasn't showing up well in whatever lighting he was using? It turned out this wasn't far from what Ed discovered himself on further inspection. He wrote back to me and reported the following:
Finally figured out the pink precipitate from the manganese problem. I've been getting the pink precipitate all along. I put the solution under the microscope and found it. It's kinda (the crystal precipitate) got the shape of a fishing fly or dragon fly, and it's definitely pink. The picture makes it look brown, but it is pink.
Thanks for sharing your results, Ed. It's nice to hear the bead test gave good results, as well.
Following is the photo which Ed submitted, showing the cluster of crystals.

Photo submitted by Ed J.


III. An Exercise in Periodic Law...?

It would be easy to step on some philosophical toes if one thought aloud while trying to explain to oneself "why" something happens. Shaky indeed is the philosophical ground one treads when one dares to use observed facts to form a theory which in turn will be used to predict what will happen next time. For perhaps more philosophical debate than you want to read about3, do a web search on "inductivism" or "inductivist reasoning". Read some of Karl Popper's writings; or, for a downright entertaining journey into the debate, read David Deutsch's The Fabric of Reality.
Acting as if none of the above makes any difference at all, we're going to attempt to use Periodic Law to predict something. We'll have a look at solubilities of minerals... this should be easy, right? That's what I thought when I was first trying to write this article.
Let's suppose we have two minerals to test; one is essentially pure MgSO4 with some water of hydration (i.e., epsomite), while the other is essentially pure CaSO4 (gypsum). Repeated experiments will show that gypsum dissolves with much greater difficulty in water at a given temperature than does epsomite. If we look at barite, BaSO4, we find that it's even more resistant to dissolving in water-- to the point that it essentially doesn't, even when that water is acidic enough to dissolve the gypsum4. If we're dealing with sulfates as is the case here, we might say the solubility decreases when going from Mg toward Ba; that is, lower placement in Group II means lower solubility. A quick look, however, suggests that the hydroxides of the Group II metals follow a reverse trend. According to the 69th Edition of the Handbook of Chemistry and Physics (CRC Press, 1988), solubilities in water of some Group II hydroxides are as follows:
Mg(OH)2 0.004 g / 100 mL at 100 C
Ca(OH)2 0.077 g / 100 mL at 100 C 5
Sr(OH)2 21.83 g / 100 mL at 100 C
Ba(OH)2 94.7 g / 100 mL at 78 C (above this temperature the substance melts)

Why does this happen? Perhaps a little bit of knowledge really is a dangerous thing... having had many years pass since General Chem I & II, my memory gets vague when pressed to say how electronegativity, atomic mass, atomic radius (etc.) relate to these trends, or how they affect something like a carbonate's solubility in acid... or, for example, how these things relate to crystal structure. As one who fancies himself an experimentalist rather than a theorist, I'm not discouraged - perhaps the answer can be found with... more experiments!
As an experiment, then, we obtain four minerals: calcite (CaCO3), witherite (BaCO3), magnesite (MgCO3), and strontianite (SrCO3). We attempt to dissolve each of them in water, then in dilute HCl. We of course expect them to follow a neat trend in solubility according to the placement of the metal cations in the periodic table. In other words, the solubilities should be either increasing or decreasing in order of Mg, Ca, Sr, and Ba. Magnesite starts out dissolving quite readily in dilute HCl, followed by calcite; everything looks fine when we try strontianite, which dissolves slightly less easily in turn than calcite. Is this a trend? Witherite is next, and it's... a little more soluble than strontianite.
In case you wanted to know, there's no easy explanation of this behavior for experimentalists like me. I couldn't look at the Periodic system and say with confidence why this "trend" occurs (rather, why the trend is broken) in this way. Fortunately, I stumbled onto a very informative page by Jim Clark at Chemguide UK: http://www.chemguide.co.uk/inorganic/group2/solubility.html. Prepare to delve into enthalpy and entropy while the explanation is developed, but it's worth a look. A link near the bottom of the second page, "To an attempt to explain these trends", puts into words what so many textbooks simply gloss over. According to Clark (2002), it's probably better that first-year chem courses simply ignore the question altogether. Even at the undergrad degree level, I don't recall any fellow students who had a real grasp of this; I sure didn't. Enthalpy and entropy may seem dry, even boring subjects (at least they did in college), but if they have this kind of effect on things, they're worth another look.



IV. Site News

More glassware is now available in the on-line catalog, including 500 mL distilling apparatus and the elusive glass retort in 250 and 500 ML sizes. There are also more types and sizes of clamps, stands, and support rings now available.




That concludes this issue of the CR-Scientific newsletter.
Until next time, stay safe and have fun.


Email: sales_AT_crscientific_DOT_com

(replace the "_AT_" with an @ symbol and the "_DOT_" with a period. In order to thwart spam-bots, I'm no longer providing a direct link or easily-harvested address.)

Notes:


1 It's not all that difficult to guess that antagonistic enzymes could be a problem. Suppose you're doing an assay for something that splits apart a molecule into two products, but the homogenate contains another enzyme that puts those two "products" right back together into the original molecule. The hope here is that one or the other is favored under the conditions encountered in the homogenate (hard to predict; the cell contents are dumped everywhere!), or at least that you can separate the antagonistic enzymes from one another. It could also be that your assay method chemically "ties up" the products of the first enzyme so they can't be put back together by the second one.

2 Note that using a high-purity reagent instead of a mineral sample is worth trying if you're attempting an assay for the first time. You then needn't worry about impurities and other elements making the test results difficult to interpret, as might happen if you've picked a mineral specimen that had a couple of different metal cations (and in more than one oxidation state, perhaps).

3 Well, maybe you're like me and really do enjoy spending hours reading about things like the inductivist-positivist debate. Even though I like Popper's writings and agree with many of his points, I find there are few things better suited to treating insomnia than reading books on the philosophy of scientific method.

4 Even the hydrochloric acid of the human stomach won't attack barite; here's a useful property for medical science. Have you ever tried one of those "tasty" barium sulfate drinks before going to the X-ray room? Normally, it would be hazardous or fatal to ingest barium compounds, but the sulfate can make it through the human digestive system almost completely intact. The fraction of barium sulfate that goes into solution is negligible. Barium sulfate will, however, start going into solution when heated with concentrated sulfuric acid.

5 The book shows Ca(OH)2 as having lower solubility at 100 C than at 0 C. I suspect a typographical error or a transposition of values- as far as I ever knew, only gases were less soluble in hot water than in cold. Books seem to be unable to agree on consistent solubility values for compounds anyway.


Works Cited / Suggested Reading:

Amdur, Mary, J. Doull, and C. Klaassen, eds. Casarett and Doull's Toxicology: The Basic Science of Poisons. New York: Pergamon Press, 1991.

Anger, V., and Feigl, F. Spot Tests in Inorganic Analysis. Amsterdam: Elsevier, 1972.

Chamot, Emile, and Mason, Clyde. Handbook of Chemical Microscopy. London: Chapman & Hall, 1940.

Clark, Jim. "Problems in Explaining the Solubility of Group 2 Compounds", web article at http://www.chemguide.co.uk/inorganic/group2/problems.html. 2000.

Feigl, F. Spot Tests in Inorganic Analysis. Amsterdam: Elsevier, 1958.

Feigl, Fritz. Spot Tests in Organic Analysis. Amsterdam: Elsevier, 1956.

Hillebrand, W., and Lundell, G. Applied Inorganic Analysis. New York: John Wiley and Sons, 1929.

Merck & Co., Inc. The Merck Index, Tenth Edition. Rahway, New Jersey, 1983.

Snell, F., and Snell, C. Colorimetric Methods of Analysis: Volume III - Organic I. New York: D. Van Nostrand, 1953.





While the information in this newsletter is thought to be accurate to the best of the writer's current knowledge, it is not guaranteed to be free of errors or to be suitable for any particular use. The procedures and experiments outlined within can be dangerous or even fatal if carried out improperly. If you choose to attempt any of them, you proceed entirely at your own risk.

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