1y ago

gimmie

22 comments

For anyone not acquainted with Things I Won't Work With
And yes, what happens next is just what you think happens: you run a mixture of oxygen and fluorine through a 700-degree-heating block. "Oh, no you don't," is the common reaction of most chemists to that proposal
- "Even Streng had to give up on some of the planned experiments, though (bonus dormitat Strengus?). Sulfur compounds defeated him, because the thermodynamics were just too titanic. Hydrogen sulfide, for example, reacts with four molecules of FOOF to give sulfur hexafluoride, 2 molecules of HF and four oxygens. . .and 433 kcal, which is the kind of every-man-for-himself exotherm that you want to avoid at all cost. The sulfur chemistry of FOOF remains unexplored, so if you feel like whipping up a batch of Satan's kimchi, go right ahead."
  Holy shit I'm dying.
  
  From the peroxide-peroxides article:
  Instead of being locked in a self-storage unit with two rabid wolverines, why not three? Instead of having two liters of pyridine poured down your trousers, why not three? And so on - it's a liberating thought. It's true that adding more oxygen-oxygen bonds to a compound will eventually liberate the tiles from your floor and your windows from their frames, but that comes with the territory.
  I'm in tears from this stuff and I know fuckall about chemistry.
  
  Though I managed in chemistry, I understand that kcal is just an expression of energy and the reaction’s event total time is important in judging ‘violence’ of reaction - but for the uninitiated like myself, how do I quantify “433 kcal” compared to a more mundane reaction?
- There's also this rocket with fluorine in the fuel: https://youtu.be/KX-0Xw6kkrc
  The thing had an asbestos liner, and that's the least concerning thing about it. Oh, and they also used lithium in the mix to make a controlled metal fire. So they combined a very reactive group one metal with very reactive fluorine, and this is fine.
  
  I love this so much
  Ok, so we have a large tank of pyrophoric and highly corrosive metal at 300C which we hope to convert to a metal fire. A metal fire is scary enough. A metal fire in pure oxygen would be worse. But...we can do better.
  You have my interest yes
  Oxygen isn't the best oxidiser.
  Why is there heroic orchestra music in my mind now
- If you enjoy that series, and it has me in fits of laughter every time I read it, may I also recommend a book called Ignition! An informal history of liquid rocket propellants. It has a similar sense of awed, horrified, fascination with the goings on in the world of chemistry.
  
  You know, the fact that I clicked that link in a heartbeat makes me wonder just how many lists I'm on.
- I wish there was more of this series.
  
  https://www.science.org/action/doSearch?field1=Title&text1=%22Things+I+won%27t+work+with%22&field2=Contrib&text2=Derek+lowe&field3=AllField&text3=&publication=&Ppub=&startPage=0&sortBy=Earliest
- Super interesting compound. I wish somehow I could watch footage of strengs experiments. I bet some of those explosions would be really cool in slow mo lol
  
  Sand Won't Save You This Time, which somehow isn't in the official listing despite it being one of the better ones of "Things I Won't Work With," comes with an accompanying video of chlorine triflouride doing its thing and setting quite a few normally-not-flammable things on fire, back from the era of Youtube when Youtube was good.
You know what is scary? FOOF, ClF3 and fluorine arent the most reactive or horrifying of the oxidizers. There's much worse.
- Go on. Tell me more, and then send it to Derek Lowe.
  
  There are multiple high oxidation state metal fluorides that can be synthesized and isolated in relatively large quantities as pure substances that have a higher electron affinity than Fluorine. The most oxidizing of these is Platinum hexafluoride whose electron affinity approaches 8 electron volts. To give you an idea of how insane that is, Fluorine's electron affinity is slightly higher than 3 electron volts. i.e adding an electron to PtF6 liberates almost 3 times the energy that adding an electron to Fluorine does. It is such a strong oxidizer that it can tear an electron off Oxygen molecules to form PtF6 * O2. It was this observation that lead to the experiments that demonstrated the first noble gas compounds. PtF6 reacts with Xenon to form a similar salt. This was the first time nobel gases were proven not to be universally inert.
  Mixtures of Fluorine and lewis acids like Boron trifluoride, Antimony pentafluoride and the like can functionally act as far stromger oxidizing agents than Fluorine itself. These form superhalogen salts when they react with things or just destroy whatever it is that was unfortunate enough to be mixed with them.
  PtF6 is not the molecule with the highest electron affinity that can exist. It is just the current record holder for the highest electron affinity of something we have been able to isolate as a pure substance. Hyperhalogens, which are essentially russian nesting dolls of oxidizing agents, can approach electron affinities of 10 electron volts. And while we cannot isolate these in their pure state, their extreme affinity for electrons can be used to stabilize otherwise unobtainable positively charged ions as salts.
  Then there are molecules like datomic Beryllium monoxide that are too reactive to isolate as bulk materials that are capable of reacting with Helium and Neon which are the least reactive elements. A combination of Copper Fluoride and Sodium Fluoride in the gas phase are similarly capable of reacting with pretty much any molecule or atom we throw at them.
  Naked Hydrogen ions (protons or H+) have also never been isolated in solution or in the solid phase due to their reactivity. The H+ youve probably run across in chemistry arent actually isolated H+ ions. Theyre always bound to a ligand like water. There does not exist a neutral atom or molecule that H+ will not violently react with.
A better representation would the burning house girl meme.
With the house labelled "Everything" and the smug girl as "Fluorine".

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