And here we are again – sorry about the holdup, I was arguing with some exciting spectra that had too many protons in them, and, due to a slight hiccup with a Russian scientist inexplicably leaving out a vital detail in his paper ten years ago and me running out of drugs and momentarily forgetting what a vinyl group was a couple weeks ago, I am now closer than ever before to actually blowing myself up or writing something publishable or seriously upsetting an otherwise friendly physicist. Also just as a heads up, the Internet Archive has literally billions of things you probably aren’t supposed to be allowed to download – including the definitive standard text on rocketry and the entirety of the Miss Peregrine books and more audiobooks and concert recordings than you can shake a stick at. (Highly recommend this one.)
Today we’re going to talk about all the exciting things you can do with the xenon compounds we were talking about last time. They all have fluorine on. You actually can’t stick anything that isn’t fluorine onto xenon just from the word go, it won’t work. This makes the fluorine chemists unbearable. But once you’ve got the fluorine on, many, many, many things are possible. Remember how I was saying it is vital we keep all this stuff as dry as possible, because fluorine will react very hard with water and make hydrofluoric acid which we absolutely do not want? You’re going to hate this: the single best way of making xenon trioxide is to grab some of your xenon hexafluoride (that’s the one that makes a wonky octahedron because it doesn’t really have space for all those substituents), and dump it into water.
Don’t actually do that!
What really happens is that the XeF₆ gets put in a tube, and then a stream of nitrogen gas (which is super useful because it’s generally inert and a lot cheaper than helium) gets blown down the tube, and sweeps the XeF₆ into water without the chemist having to be in the same room. This would be a good thing, because the reaction goes like this:
XeF₆ + H₂O → XeO₃ + 6HF
Remember how I said hydrofluoric acid was quite enthusiastic? It’s not even the most exciting thing in that reaction. (Quick side note: when you dissolve carbon dioxide in water, you get carbonic acid, H2CO3, which you’ll recognize as being carbon dioxide with a random water attached to it. For a long time everyone thought carbonic acid didn’t exist except as a useful fiction for explaining why CO2 stays dissolved in water so nicely – some people at NASA made some in 1991, but for all useful purposes you can’t isolate it. Well, it turns out XeO₃ behaves similarly, because all those lone pairs on all those oxygens are really good at grabbing stray water molecules and hugging them. You can isolate xenic acid if you’re very persistent but I haven’t the foggiest why anyone would want to.) XeO₃ on its own is plenty thrilling. It makes a trigonal pyramidal molecule, which is one of the things Linus Pauling predicted about it way back in the wild west days, because VSEPR is really useful for predicting molecule shapes.
Oh – VSEPR – Valence Shell Electron Pair Repulsion. Means that whatever you’ve got around an atom center will try to be as far away from all the other things as possible, because all the electrons are picky about personal space. Lone pairs (pairs of electrons with nothing bonded to them) demand more space than actual substituents, and… VSEPR’s a whole course on it’s own, y’know? (Also didja know Linus Pauling is currently the only person to be awarded two solo Nobel Prizes in different fields? Marie Curie had two in different fields, but one of them was shared.)
Where were we? Oh yeah, once we’ve got this exciting XeO₃, which is the first of our untraceable explosives, we can use it as an oxidizing agent: it’s very strong, but quite slow.
XeO₄ is formed by bubbling ozone (which is quite toxic and tolerably unstable) through an aqueous solution of XeO₃. You have to make the solution quite basic so it strips off any extra hydrogens that might be tempted to stick around and make it back into xenic acid. Or you can take your ‘xenic acid’, dump it into a basic solution, and it will create aqueous perxenate (scream), xenon, oxygen, and water. Perxenate: XeO₆⁴⁻. If you’re still counting, that has an oxidation number of eight. Obviously we don’t want it to blow up in our faces juuuuuust yet, so just as well it’s stable-ish when you stick a barium either side of it, to ‘cancel out’ that horrendous 4- charge. Add sulfuric acid, and this happens:
Ba₂XeO₆ + 2H₂SO₄ → XeO₄ (gas) + 2BaSO₄ (precipitate, stable and boring) + 2 H₂O
XeO₄ is an explosive gas. Now, please, if you can imagine an explosive gas, could you let me know what it would look like? Hm? Because I have no clue. I can imagine a liquid or a solid exploding just fine, but a gas is completely inconceivable to me. Also it’s ludicrously unstable – dissociates back to Xe and O₂ at the drop of a hat and with a release of about 642 kJ/mol (that’s a lot).
And these are the untraceable explosives. Because if you could get them to where you wanted them without blowing your fingers off, you could detonate them and the only residue would be a little charring from a whole lot of energy being released very rapidly, and some xenon and oxygen floating around. Which are in the atmosphere anyway. (Incidentally – why does xenon tetroxide work but xenon octafluoride not? I mean, it’s the same oxidation state of xenon, right? Unfortunately the reason is absolutely as mundane as ‘there’s space for four double-bonded oxygens, but not for eight single-bonded fluorines’.)
Of course, you don’t have to replace all the fluorine on a xenon center with oxygen. Xenon oxyfluoride compounds, in order of ‘will grab any oxygen that’s floating around’ to ‘oxygen what oxygen’ are XeF₆ > XeO₃F₂ (because this one’s ridiculously stable; pictures below) > XeO₄ > XeOF₄ > XeF₄ > XeO₂F₂ > XeO₃. They all look kinda wild, but pay attention to the little dots. Those represent lone pairs, and if I was really enthusiastic here I’d draw in the full orbital. I’m not. Short version is, a lone pair is even picker about its personal space than a bond, so it shoves everything else on the molecule out of its way. If not for the lone pairs, all the substituents would be distributed evenly around the xenon center.
There are a bunch of reasons this could be, but of course we’re not sure which reason or combination of reasons it is, and… whenever you have two hypotheses, it will turn out to be something else entirely.
One of the most important principles of chemistry is steric hindrance, the concept that unless they’re bonded to each other, atoms really, really don’t like to be close together. If you’re trying to bullshit an answer in chemistry, the magic word is steric. One theory goes that on a given oxidation state of xenon, one oxygen is sterically favored over two fluorines – it takes up less space, therefore the substituents are all less antsy, therefore the molecule is more stable. (The other theory is that oxygen on a double bond ‘draws off’ less of the electron density from the xenon center, so there’s more to spare to the fluorine, so the Xe-F bond is stronger.)
And then it all just gets wildly exciting.
Brock et. al. made xenon oxydifluoride in 2007 – it’s shaped like a T, because those pesky lone pairs.
So last time we covered how you can make xenon fluoride salts, but it works with xenon oxides as well: xenon trioxide with a random metal fluoride (it’s been tried with potassium and cesium) gets you M (the metal) as a positive counter-ion to the negative [XeO₃F]⁻ – know what, I’ll just draw it for you. I love ChemDraw. It’s terribly therapeutic.
So everything about how oxygen forming a double bond to xenon and that being supergreat for displacing fluorine… that still holds. But you can also get ‘bridging oxygen’, that’s forming a single bond to xenon and another single bond to something else.
Plain old ordinary XeF₂ can be made to form a teflate (-OTeF₅: yes, that is a tellurium atom in the middle) or seflate (OSeF₅) or di-teflate or… fun times, nu? You can get surprisingly high oxidation states like this, right up to xenon pentateflate. (No more than that… because there isn’t room! What’s the magic word?) All of these involve a bunch of teflates being attached to a xenon center by a bridging oxygen. Schumacher and Schrobilgen are absolutely the dudes here; in 1984 they made xenon dioxide diteflate for the first time; I honestly don’t know if anyone else has tried that since.
In 1982, Sawyer and Schrobilgen (Schrobilgen again! Man’s everywhere) decided that attaching fluorine and oxygen to xenon was a mug’s game, and decided to give it a crack with nitrogen. Now, nitrogen doesn’t actually make particularly strong bonds (which explains why ‘my’ rocket fuels are all nitrate-based – it’s conveniently reactive) and xenon really doesn’t make strong bonds, so this should have been sort of doomed to failure, but they pulled it off with solid FXeN(SO₂F)₂, and if you can’t picture that, don’t worry, I can’t either, but it looks like this:
And you can complex it with AsF₅! (Which turns bright yellow; the only redeeming feature of inorganic chemistry is the pretty colors you get out of it.) And you can make a screaming mess of that by reacting two of it together under vacuum so that they stick together at the fluorine – it’s called µ₂-F, which means fluorine attached to two different things. That was quite a shock when I was coming straight from covalent bonds a couple of years ago.
As Smith et. al. discovered in 2007, you can also – just barely – attach Xe to nitrogen by using a Lewis acid (like AsF₆) and a nitrogenated Lewis base (like F₃S≡N – this is called thiazyl trifluoride). Which are completely different from regular acids and bases. Most of these reactions have to be kept cold – not like liquid-nitrogen cold, just about -20 or so. Peanuts, really. This gets you [F₃S≡NXeF]⁺[AsF₆]⁻. Which the very observant will realize is sticking a Xe(II) center onto both a fluorine and a thiazide. Two years later, Smith and his danger-hunting buddies got bored or drunk, and reacted their product with bromine pentafluoride, to stick an extra fluorine onto the thiazyl (they took it off the xenon, if that’s any comfort). That wasn’t risky enough, so they did it again. The second time, the charges wouldn’t balance, so they put a hydrogen onto the already dangerously overloaded thiazyl group. And then the bonds wouldn’t balance, so they had to make the sulfur-nitrogen bond only double. Gasp.
Frohn and friends made xenon bond to carbon in about 2000, but I think I’ve gone on quite long enough here.
And I really, really, really hate inorganic chemistry, just sayin’. Two weeks to go.
Thanks for reading.
I know it’s long. You haven’t seen the paper I actually enjoyed writing this semester.