> LICK BASIC FOUNDING BLOCK OF REALITY
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I really like this kind of lateral thinking marvelous ideas!
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there’s a lot of salty stuff there that would be fun to lick before my tongue fell off!
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Can you really lick a gas?
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----sticks out tongue as far as possible----retracts tongue----smacks----
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Hmm, elemental Fluorine seems like it should be in the “See you on the other side” category, since your tongue would probably literally be on fire, as well as the rest of you, and most everything else nearby.
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Any gas can stop being a gas if you’re determined enough!
…It might not be a good idea to lick it then, though.
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True, which is why this post needs to specify temperature and pressure.
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Radon and Oganesson are the fun little gotchas for anyone who gets complacent licking the rest of the noble gases
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Fun fact: Oganesson was in my first IFComp entry.
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Now you just need to make an IF with these mechanics…
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Says a lot about this community, doesn’t it?
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>LICK TOPIC
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Would the world’s largest sample of Oganesson even be enough to give you a detectable dose of radiation? Sure, the stuff is rediculously unstable, but that instability makes just synthesizing a single atom of the stuff an engineering marvel and google informs me only 5 atoms of the stuff have ever been made by humans. Assuming an average of 1 decay per nucleon lost* in the process of decaying to, presumably lead, we’re looking at less than 100 decays per atom and I’m guessing it takes astronomical numbers of decays to match background levels.
*No idea how well this figure matches reality, since alpha decay loses 4 nucleons, beta decay doesn’t change nucleon count, and I have no idea how gamma decay affects an atom.
Also, now I’m curious what the highest boiling/melting/sublimation points among elements that are gaseous at standard temperature and pressure are… I believe Gallium has the lowest melting point of solid at STP elements, being solid at STP, but liquid at human body temperature.
And yeah, I don’t want to be in the same building as elemental fluorine, or most of its compounds… when professional chemists nope out of working with a chemical, you know it ain’t something to mess with… though considering what I’ve heard about Flourine compounds, I have to wonder what makes sodium flouride so safe compared to other flourine compounds.
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Where did I put my chart of nuclides…
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Salt!
The thing that makes halogens (Greek for “salt-makers”) so absurdly reactive is that they really, really want electrons. Fluorine is the most electronegative (electron-attracting) element in the whole table!
The alkali metals, on the other hand, have electrons they’re not especially attached to. So when a halogen reacts with an alkali metal, it rips that electron right off, and now both of them are happy—a halogen ion with an extra electron, and an alkali ion missing an electron.
The electric charge then holds these ions together in an ionic bond, forming a salt. When fluorine is in a compound with something like oxygen, on the other hand, they’re trying to form a covalent bond (sharing one or more electrons between them), and this is a much less stable arrangement, because the fluorine isn’t satisfied with sharing. It wants the electron all to itself, and will take any opportunity to seize it.
So if you want to keep a halogen and/or an alkali metal stable, put it in salt form. It’s best for everyone that way.
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Wait, so why isn’t Group 6 more reactive than the Halogens? Surely since group 6 is missing two electrons, it wants more electrons, right?
(I literally don’t know the answer - I’ve learnt about everything you just said, except that alkali metals are highest in the reactivity series…)
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Great question! I’ll do my best to answer, but this is rapidly running up against the limits of my chemistry knowledge. I’m sure I’ve gotten some details wrong here (possibly fundamentally), so I welcome all corrections.
First of all, group 16 (the chalcogens, “ore-makers”) is pretty darn reactive too! Electronegativity is a fairly good indicator of how much an element will want to react with other elements, and while fluorine is the absolute highest, oxygen is the second-highest. (Chlorine is third.) It’s not quite as destructive as fluorine, but oxygen also wants to react with pretty much anything available. When it happens slowly, we call it corrosion or rust; when it happens quickly, we call it fire. And almost every substance is vulnerable to it under the right conditions!
In fact, oxygen is so good at reacting that a whole lot of our biological machinery is dedicated, not to using oxygen for power (that part’s easy), but to repairing all the damage caused by having that oxygen around. When that damage accumulates faster than it can be repaired, that’s a major component of aging! About 2.4 billion years ago, some organisms evolved to spew out oxygen without it killing them in the process, and the change to our atmosphere obliterated 80% of all life on earth. Ever since then, everything living in the air (instead of around deep undersea vents) has needed a way to deal with oxygen—or it will, effectively, burn to death from exposure to the atmosphere.
But, why is fluorine more reactive than oxygen? The short version is that atoms “want” to have every available orbital filled with electrons. How does an atom “want” anything? The same way a rubber ball “wants” to fall to the ground: “want” is a shorthand for “if that happens, there’s less potential energy in the system”. (And “potential energy” itself is a shorthand for a bunch of complicated quantum mechanics that are beyond my understanding.)
But this is only one type of potential energy! What if, instead of dropping a rubber ball, I drop a steel ball—and it’s sitting right underneath a powerful magnet? It “wants” to fall, to minimize the potential energy due to gravity, but it also “wants” to stay in place, to minimize the potential energy due to magnetism. Depending on the relative strength of these two forces, it might fall, or it might not.
So while an oxygen atom “wants” to have two extra electrons, it also, separately, “wants” to have the same number of electrons and protons—electrons repel other electrons, so they need the charge of the protons to keep them close. Since oxygen has eight protons, it “wants” to have ten electrons to fill the orbitals, but it also “wants” to have no more than eight electrons to balance out its charge. And the relative strength of these wants is what determines how it will behave.
Fluorine, meanwhile, “wants” ten electrons to fill its orbitals, and “wants” to have no more than nine electrons to balance out its charge. This means it’s less “unhappy” to have ten electrons than oxygen is: it only needs to unbalance its charge a little bit (becoming a F- ion), while oxygen needs to unbalance its charge a lot (becoming a O2- ion). This makes it easier for the first desire, “wanting” ten electrons, to win out.
This is also why chlorine is less reactive than fluorine. Chlorine “wants” to have 18 electrons to fill out its orbitals, but the desire to have 18 electrons is weaker than the desire to have 10 (since it’s filling a shell farther away from the nucleus).
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So,
Carbon shares with everyone because a full shell would overdose it on negativity.
N2 is happy being covalent because being neutral is more appealing than having 3 extra electrons to fill their valence shell and so they’re happy to share.
O2 easily sets fires because they’re conflicted over being neutral or having a full shell and are always looking for a better arrangement.
F2 is explosive because both parties are greedy and want their shared electrons all for themselves.
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Can I lick it? (Yes, you can!)
Can I lick it? (No, you really shouldn’t man.)
Can I lick it? (Yeah, but you might lose your tongue again)
Well, I’m gonna (Go on then, your funeral)
Can I lick it?
To all the people who thought they was intelligent
Before this, did you really know them elements?
Comprehend the table, for it’s evident
The fact that they’re small’s not an impediment
Lappin’ left and right like a degenerate
This verse stands here as my testament
'Cause I feel the urge to lick, and it’s imminent
Can I lick it?
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Back in the day, I can remember running a reaction that required hydrofluoric acid (HF). Nasty stuff, spacesuit stuff! The acid would condense on the upper levels of the fume hood and then rain down on everything below. Yikes!
Don’t miss those back in the days…
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