Chapter 18: The Group 18 Elements

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Welcome back to the Deep Dive.

Today we're tackling one of chemistry's real head -scratchers, the noble gases.

For ages, everyone thought they were just, you know, completely inert, didn't react with anything, but it turns out their story is way more interesting than that.

There's some surprising chemistry hiding there, especially with one of them.

Think of this as your step -by -step guide through this chapter will break down the big ideas, make sense of it all.

It really is a fascinating bit of science history.

I mean, finding these elements back in the late 1800s, it didn't just add boxes to the periodic table.

It actually forced chemists to rethink bonding itself.

We went from calling them rare, which wasn't always accurate, to inert and now noble.

And noble kind of fits, doesn't it?

Reserved, but not completely unresponsive.

So our goal today is to walk you through that, especially xenon's chemistry.

How do we get them to react?

What do the compounds look like?

How are they used?

We'll try and paint a picture, you know, make the structures clear even without diagrams in front of you.

Exactly.

We'll start with the basics, like why are they usually so unreactive?

Then we'll get into the really cool stuff.

The compounds they do form, the uses, and the clever ways scientists actually make them.

Okay, ready to dive into group 18?

Let's unpack this.

All right, first things first.

Group 18, we've got helium, neon, argon, krypton, xenon, and radon.

They're all monatomic gases, so just single atoms floating around.

And those names we mentioned, rare gases.

Well, argon's nearly 1 % of the air we breathe.

That's hardly rare, is it?

Not at all.

More common than keros, actually.

And inert gases, well, we know that's not quite right either, because compounds do exist.

So the noble gases is where we landed.

It suggests that, you know, they're a bit standoffish chemically, but not impossible to persuade.

So what makes them so stable, so noble?

It all boils down to their electrons.

They have a completely full outer electron shell.

All the spots are taken, basically.

For most, that's an NSNPO configuration, 2s electrons, 6p electrons, just perfectly stable.

And because of that full shell, they have really high ionization energies.

Meaning, it's hard to rip an electron away from them.

Exactly.

It takes a ton of energy.

And the flip side is, they have negative electron affinities.

They don't want to gain an electron either, because a new electron would have to start a whole new shell, higher up in energy.

Not favorable.

Okay, that makes sense.

And where do we find them?

You mentioned argon in the air.

Right.

Helium is number two in the whole universe, after hydrogen.

But on Earth, pretty scarce in the atmosphere.

It's just too light.

It literally escapes Earth's gravity.

Argon and neon, though, they stick around.

They're surprisingly abundant in our air.

Then there's radon.

That one's a bit different.

It comes from radioactive decay of heavier elements in the ground.

And it's radioactive itself.

It actually contributes a fair bit to our natural background radiation exposure.

Right.

I've heard about radon testing in basements.

Exactly.

Now, if you look at the group as a whole, going down from helium to radon, you see trends.

The atoms get bigger, obviously.

Radii increase.

And as they get bigger, the melting and boiling points go up, too, because the weak forces between atoms get a bit stronger.

But here's the crucial part.

That first ionization energy, it decreases as you go down.

Ah, so it gets easier to remove an electron from the bigger ones.

Precisely.

The outer electrons are further from the nucleus, less tightly held.

And that's the key to why the heavier ones, especially xenon, start to show some chemistry.

Okay, so they mostly keep to themselves, but xenon.

Xenon is the exception.

That's where it gets really interesting.

That's the one.

Xenon forms a whole range of compounds, particularly with elements that really pull hard on electrons like fluorine and oxygen.

What kind of compounds are we talking about?

We see xenon in oxidation states like plus two, plus four, and plus six, typically.

It forms bonds like XEF, XCO, even Skeif N, XCH, XE.

It can even bond to metals or act as a ligand in coordination complexes.

It's surprisingly versatile.

Krypton does a little bit, but much less.

And radon.

Well, it should be more reactive than xenon based on trends, but it's so radioactive, studying it is extremely difficult.

So the xenon fluorides are the big ones.

Definitely the classics.

You can react to xenon gas directly with fluorine gas, depending on the conditions, temperature, pressure, ratio of gases.

You can get ZS, ZF, or S.

And what do they look like molecule wise?

Okay, visualize this.

EZF is linear, xenon in the middle, two fluorines sticking straight out opposite sides.

SXEF, XEF is square planar, xenon in the center, four fluorines at the corners of a square, all in the same plane.

XEF is trickier.

In the gas phase, it's fluxional.

Its shape is constantly distorting, kind of wobbling.

In the solid state, it's more like a network of ions.

Well, okay.

Not simple shapes then for the hexafluoride.

Not at all.

And then you have oxides, like COOs and COO.

Both are known, but be careful, they're extremely unstable.

Explosive, actually.

Yanks.

They're also personates, which have the iso -ri on inside.

And oxafluorides mixing oxygen and fluorine, like ZORs, that's key -shaped, or isoero, which is a square pyramid shape.

We even have some hydrides, like HXEH, and these things called clathrates.

Clathrates?

What are those?

Imagine a cage made of water molecules, like ice, but with gaps.

The noble gas atoms, argon, krypton, or xenon, can get trapped inside these cages.

They're not chemically bonded, just physically stuck.

It's actually a useful way to handle radioactive krypton or xenon isotopes safely.

That is clever.

So it's amazing this whole family was missed for so long.

Mendeleev didn't even predict them.

Nope.

No gaps left for them.

Their lack of reactivity meant they didn't fit the patterns he saw.

The first clue was helium, seen in the Sun's spectrum in 1868, way before they found it on Earth.

Hence the name helium from Helios the Sun.

Exactly.

And the others followed.

Neon for new, argon for inactive, krypton for hidden, xenon for strange.

Radon?

Well, that's just named after radium, its parent element in decay.

And where do we get them from today?

You said helium escapes Earth.

Right.

Most terrestrial helium comes from radioactive alpha decay deep underground.

It gets trapped in natural gas deposits.

North America is a big source.

The others, neon, argon, krypton, xenon, we get from the air.

We liquefy air and then use fractional distillation, separating them by their different boiling points.

It's a low temperature process.

And in their pure form, just single atoms.

But you mentioned dimers in the liquid.

Yeah.

If you cool them enough to liquefy, the atoms can weakly attract each other through dispersion forces, forming temporary pairs or dimers.

That's why their boiling points are so low.

Those forces are really weak.

And speaking of cold, liquid helium is wild.

Cool helium -4 below about 2 .2 Kelvin, and it becomes helium -2, a superfluid.

It's a superfluid.

Yeah.

It flows with zero viscosity, zero friction.

It can creep up the walls of its container, flow through microscopic cracks.

It's bizarre stuff.

Okay.

Mine's slightly blown.

So, beyond the weird physics,

what are the practical uses for us?

Oh, loads.

Helium, we talked about balloons, obviously safer than hydrogen, but the really big use is cryogenics.

Its super low boiling point makes it essential for cooling superconducting magnets in MRI machines and NMR spectrometers.

Huge for medicine and research.

Plus, it's used to create inert atmospheres for growing perfect silicon crystals for electronics and in breathing mixtures for deep sea divers to prevent the bends.

Right, the helium -oxygen mix.

And argon.

Argon's a workhorse, too.

Great inert atmosphere for welding, preventing metals from reacting with oxygen at high heat.

Also used in incandescent light bulbs to protect the filament.

And because it doesn't conduct heat well, it's used as an insulating gas between panes and double glazed windows.

Keeps the heat in or out.

Useful stuff.

What about xenon?

It's rare, right?

And more expensive, yeah.

But it has niche uses.

It actually works as an anesthetic, though cost is a factor.

More importantly, its isotopes are amazing for medical imaging.

Techniques like hyperpolarized XCNMR give incredibly detailed images of lungs, blood flow, even the brain.

Things traditional MRI struggles with.

That sounds incredibly powerful.

And radon, we mentioned, is mostly a hazard.

Pretty much, yeah.

The main concern is indoor accumulation from natural sources in the ground.

Long -term exposure isn't good.

Okay.

But back to the good uses, lighting.

Absolutely.

Their unreactivity is perfect here.

Neon signs are the obvious one.

That classic glow.

But they're also in fluorescent lamps, powerful xenon flash lamps and cameras, and crucially in lasers.

Helium -neon lasers, argon ion lasers, krypton ion lasers.

Very important tools.

The basic idea is you zap the gas electricity, excite the atoms, and when they relax back down, they emit light at specific wavelengths.

So how did we even start making these compounds?

You said they were thought to be inert.

The breakthrough moment was 1962.

Neil Bartlett.

He was working with this really powerful oxidizing agent, platinum hexafluoride, PTF.

He found it could rip an electron off an oxygen molecule.

Then he realized something crucial.

The energy needed to remove an electron from euro was almost identical to the energy needed to remove one from xenon.

So you thought maybe PTF could react with xenon too.

Exactly.

He tried it and boom.

He made the first noble gas compound, xenon hexafluoroplatinate.

It just blew the field wide open.

Suddenly, inert wasn't true anymore.

That's brilliant.

So that paved the way for making things like Jaxa.

It did.

The direct reaction of xenon and fluorine is now the standard way.

You just mix the gases and control the conditions, more fluorine, higher pressure.

Specific temperatures generally lead to the higher fluorides like ZF and AVF.

For ZF, you might use something like 300 degrees Celsius, 60 atmospheres pressure, and a 1 to 20 ratio of xenon to fluorine.

I read about a simpler way for SS, the window sill method.

Ha.

Yes.

It actually works.

You just need a sealed dry glass bulb with xenon and fluorine, and then you leave it in the sun.

The UV light breaks the F -euro molecules into reactive F atoms and they attack the xenon.

You get nice ecureo crystals forming right there.

Pretty neat.

Very cool.

And exeostructure is still kind of debated.

Well, its gas phase structure is definitely dynamic, fluxional because of that lone pair on xenon causing distortions.

Solid state, more complex, involving XU ions bridged by fluoride ions.

It's structurally related to some other complex halogen -containing ions, actually.

So once you have these fluorides, what do they do chemically?

They're powerful oxidizing agents, first and foremost.

They readily grab electrons.

Their reactions often mirror other reactive halogen compounds.

Think redox reactions and monathesis swapping atoms around.

Can you give an example?

Sure.

XCF reacts with water carefully to make CS, swapping fluorines for oxygens, that's monathesis.

Or ASEF reacting with water gives you back xenon plus oxygen gas, that's redox.

The xenon gets reduced, water gets oxidized.

They also react with Lewis acids, things that accept electron pairs.

XCF plus antimony pentafluoride, SBFOs, gives you the salt.

XBFO, it forms a contortion.

And Lewis bases too, things that donate electrons.

XCF can react with a fluoride ion, F, which is a Lewis base, to form the XES anion.

That one has a cool pentagonal planar shape, five fluorines in a pentagon around xenon, with lone pairs above and below.

There are even bigger ones, like X and the XCO known.

Wow.

So these fluorides are like starting points for other xenon compounds.

Exactly.

They're gateways.

You can use them to make compounds with agian bonds, for example.

You might use nucleophilic attack or react them with a Lewis acid first, then add something else.

Lots of synthetic strategies branch out from the fluorides.

Okay, what about the pure oxides, like DASHO?

You said they're unstable?

Highly unstable, yes.

They're endergonic.

You have to put energy in to make them.

They don't form spontaneously from xenon and oxygen.

That tells you they're not thermodynamically happy.

You generally make them by hydrolyzing the fluorides.

XF plus water gives ZO, handle with care.

It's a powerful oxidizer and yeah, explosive.

And the personates,

ZOHEL.

How do those form?

If you have the Z -oxyanion HXOL in a basic solution, it can disproportionate.

That means some of it gets oxidized to Z of the personate ion and some gets reduced all the way back down to elemental xenon gas.

So Z basically reacts with itself to form Z and ZO.

And personates are more stable.

Relatively speaking, yes.

They're usually white crystalline solids.

The ZO ion itself has a nice symmetrical octahedral shape.

Xenon in the middle, six oxygens around it at the corners of an octahedron.

The Isapor theory predicts that nicely.

If you take a personate salt, like barium personate, and react it with strong sulfuric acid, you can generate XEO gas.

But again, very unstable, explosive stuff.

And the oxofluorides are mixtures.

ZOO, ZOL.

Right.

Combining both oxygen and fluorine bonded to xenon.

Like AZOS, T -shaped,

AZOEVOEVER square pyramidal.

There's even an AZOEVO ion, pentagonal pyramidal, formed if you dissolve fluoride ions in AZOEVO.

Okay.

Shifting gears a bit.

Xenon insertion compounds.

And this missing xenon idea.

Tell me more about that.

Right.

The missing xenon problem.

We find way less xenon in our atmosphere than we expect based on the amounts of other noble gases, like krypton or argon, relative to their cosmic abundance.

It's off by a factor of 20.

So where did it go?

The leading theory is that, deep inside the Earth, under incredibly high pressures and could form stable insertion compounds, maybe substituting for silicon in silicate rocks, effectively locking it away in the Earth's mantle or core.

It's a fascinating possibility explaining the atmospheric depletion.

That's wild, trapped in rocks deep underground.

Potentially.

And related to insertion, we have these HEY compounds, noble gas hydrides like HXCCL, HXEOH, even HXEOXE, the smallest molecule with two xenon atoms.

They're usually made by UV photolysis of precursors frozen in solid noble gas matrices at very low temperatures.

Highly unstable.

But they exist.

And actual xenon carbon bonds.

Organoxenon.

Yep.

That was another big step, first reported in 89.

We can now make compounds with direct ICSI bonds, often starting from X1 or Hesero.

One method involves reacting them with organobranes.

The xenon kicks out the boron.

Are they stable?

The Z1s not very.

They tend to decompose above maybe negative 40 Celsius.

But there's this cool hexioion two xenon atoms bonded together, which has the longest bond known between identical main group elements.

Really stretched out.

For sero -organic compounds, stability improves if the carbon is part of an aromatic ring, especially one with electron -withdrawing groups like fluorine attached.

That helps stabilize the XE bond.

One more really weird one.

Noble gases as ligands.

Attached to metals.

Yeah, coordination compounds.

Known since the 70s, actually.

The first stable one characterized was Algorizur, a gold ion surrounded by four xenon atoms in a square plane.

Stable up to about neck at 78 Celsius.

Made in a special solvent mixture.

How does that even bond?

It's a weak interaction, but real.

We also see transient ones, like metal carbonols with a noble gas attached, MCOE where M is chromium, molybdenum or tungsten, and E is argon, krypton or xenon.

These are usually trapped and studied using matrix isolation, freezing them in solid argon at like 10 Kelvin.

And there are trends in stability.

Definitely.

Tungsten forms the most stable ones among those metals, and xenon forms stronger bonds than krypton, which is stronger than argon.

Matches the polarizability trend.

The bonding involves interactions between the noble gas P orbitals and orbitals on the CO groups.

Subtle stuff.

Some rhodium complexes have even been made in supercritical xenon.

Amazing.

So just quickly, what about krypton and radon chemistry?

It's more limited.

Much more limited for krypton.

Mostly just KR -ero.

Radon should be more reactive due to lower ionization energy, and there's evidence for things like R in fosalts, but stugging it is hampered by its intense radioactivity.

It's just too dangerous to work with easily.

And KR -eros itself.

Krypton de fluoride.

Made using electric discharge in a cold KRF mix.

Like ECS, it's linear, but it's colorless, highly reactive, and only stable at very low temperatures.

Much argon hydrofluoride.

Made by UV or radiation in solid argon.

Stable only up to about 27 Kelvin.

A real achievement.

And the clathrates we mentioned are KRA, shrapton water, or quinol cages.

Helium and neon are too small.

They just slip out.

There's even speculation that clathrates might explain why krypton and xenon seem depleted on Saturn's moon titans.

And noble gases inside fluorines.

Yeah, endohedral fluorines.

Literally trapping he or knee atoms inside the CRO buckyball cage.

Not chemically bonded, just physically encapsulating.

Really helium.

Still basically unreactive.

Pretty much the king of unreactive.

Apart from being trapped or some fleeting ions seen in experiments, no stable neutral helium compounds are known.

There are theoretical predictions, like for Hebio, but nothing isolated.

It really holds onto its electrons tightly.

So wrapping this up, we went from thinking noble gases did nothing to finding this whole rich complex chemistry, especially for xenon.

It really challenges that old inert label.

Absolutely.

It shows how understanding atomic properties, ionization, energy size, electron affinity helps predict and explain reactivity, even where you don't expect it.

We've seen xenon form fluorides, oxides, hydrides, bonds to carbon, bonds to metals.

It's incredibly diverse.

And it highlights the clever techniques chemists use low temps, high pressures, matrix isolation to push the boundaries and study these fascinating, sometimes fleeting molecules.

Yeah, it's a great story of scientific discovery.

And that missing xenon idea makes you wonder, doesn't it?

What other weird compounds might be out there?

Maybe in extreme places, other planets, deep earth, things we haven't even imagined yet.

It definitely leaves the door open.

There's always more to find out.

Well, that's all the time we have for this Deep Dive.

Huge thanks for walking us through that.

Until next time, from all of us here at the Deep Dive and the Last Minute Lecture Team, thank you for listening.