With all this talk of Nobel Prizes (still excited!), I thought it would be good to talk about noble gases – because being a noble gas is basically the atomic equivalent of being a Nobel laureate! Typically, noble gases like neon are non-reactive (inert) because they already “have it all” (a full valence shell), but give them some energy and they can really *shine*!
The periodic table is like a “menu” of all the known elements – everything from carbon and hydrogen (which we’ve looked at a lot) to ones like neon and argon, which encounter in our posts, we typically have not. On the far right of the table, the last column, you find these and 4 other “noble gases” which have a complete complement of outer electrons. Atoms share and give and take electrons all to become like a noble gas. But what if you are the noble gas? You were born “perfect”. Your electronic bucket list is empty. So now what? Typically, not much of anything… You don’t encounter noble gases very much in biochemistry because of this. But you do encounter them all the time when you’re out and about (although most newer lights use different methodology)
Atoms are the basic units of elements and they’re made up of 3 key subatomic particles: protons, neutrons, and electrons. Protons are little positively charged things – they hang out with neutral neutrons in a dense central atomic nucleus and then negatively-charged electrons (which have an equal but opposite charge despite being itty-bitty-er) whizz around them in an electron cloud – and the ones on the edge of the cloud are called valence electrons, and they’re “special” (though fundamentally they’re all the same (the chemical equivalent of the whole “celebrities are just like us” thing)
You see, the difference between elements (e.g. carbon (C) vs nitrogen (N) vs neon (Ne)) is the # of protons they have. Carbon has 6 protons, so we say it has an atomic number of 6, or Z=6. Neon has 10 protons, so atomic number 10, Z=10. Etc. etc. etc. But we pay more attention to the electrons because, while the protons are fixed in number and stuck in place, electrons can move around and interact with other molecules. You can never know exactly where an electron will be – but there are places they most like to hang out, which we call “orbitals.”
At the physical chemistry (p-chem) level, it’s not actually like this, but it can be really helpful to think of electron orbitals as “shells” or onion layers, and the periodic table is organized in a way that makes this shell model quite convenient. As you look down a column (these columns are frequently called groups or families), you add a layer but keep the same number of electrons in the outermost shell in the element’s neutral form – and, conveniently, that # is the same as the column # (i.e. carbon is in the 4th column (of the main groups of the table) & it has 4 valence electrons in its neutral form whereas nitrogen, in the 5th column has 5, & neon, in the 8th column has 8). These outermost electrons are called valence electrons – they’re the most energetic and, being furthest away from the positive pull of the nucleus, they’re “least loyal” to the atom they come from, and more liable to interact with electrons of neighboring atoms and even leave altogether.
Why would they want to do that? For reasons outside the scope of this post, atoms are usually most stable (and thus happiest) if they have a full “outer shell” like the elements in the last column (the noble gases) do – for most of the elements we commonly deal with in biochemistry, this means having an octet – 8 electrons in the outer shell. An exception is hydrogen, which only wants 2 because its outer shell is an inner shell for everyone else (except helium). Hydrogen only has a single proton, so you can’t expect it to reign in a ton of electrons! Instead it only wants 2, like the noble gas helium (He).
Atoms can share pairs of electrons to form “covalent bonds” (1 pair for a single bond, 2 pairs for a double, and 3 pairs for a triple). So, H, with its single electron can join up with another H, with its single electron to satisfy them both. And N, with its 5 valence electrons can join up with another N through a triple bond that makes them both get to 8. So hydrogen and nitrogen gases exist in the diatomic (2 atom) form – H₂ and N₂ (oxygen is this way too)
Atoms can also try to get to their magic numbers by gaining and losing electrons. If a neutral molecule loses an electron, # of protons > # electrons, so it becomes positively-charged (cationic) and if a neutral molecule gains an electron, # of protons < # of neutrons, so it becomes negatively-charged (anionic). “Ionic” just refers to a charged thing and things usually don’t really want to be charged. So there’s this kind of compromise atoms have to make about getting that full outer shell vs becoming charged and what decision they make has to do with things like how close they are to full and how well they can handle the charge.
If they do decide to go charged, they commonly also get “help” from other molecules – since opposite charges attract, cations are attracted to anions & vice versa – so even when Na (which has one electron it doesn’t want in a shell all to itself) gives up an electron to Cl (which just needs one to complete its shell) to give you Na⁺ & Cl⁻, those ions hang out together through an “ionic bond” which is really just a strong attraction – unlike covalent bonds which involve electron sharing (i.e. orbitals unite!)
So, atoms are sharing and giving and taking all to become like a noble gas. But what if you are the noble gas? You don’t need help from other atoms to help make you happy.
The noble gases are, in order of increasing bigness: Helium (He), Z=2; Neon (Ne), Z=10; Argon (Ar), Z=18; Krypton (Kr), Z=36; Xenon (Xe), Z=54; Radon (Rn), Z=86. They’re sometimes called “rare gases” and you can understand why if you try to find them in the atmosphere. In % of atmosphere terms, they range in abundance from ~0.000009% (Xenon) to ~0.934% (Argon).
And the real specialness of these noble gases comes from their independent spirit and preference for going it alone. Molecules don’t really *like* hanging out with other molecules, because forming bonds “ties them down.” It’s that whole entropy concept we talked about http://bit.ly/thermodynamicsfreeenergy
Basically, there’s this concept called entropy, which is commonly described as randomness or disorder, and the 2nd law of thermodynamics says that Nature likes this disorder – so molecules react in ways that leads to increasing entropy and thus the entropy of the universe is constantly increasing.
Entropy can be a kinda confusing concept, especially since we’ve spent weeks and weeks talking about how molecules join together in ways that build nice orderly molecules like proteins and DNA – but that joining takes energy and is paid for entropically in other ways – like the splitting of 1 ATP molecule into multiple molecules – and “enthalpically” by having more stable bonds.
Another way to think of entropy is the number of “microstates” something could occupy. The more tied up something is, the fewer possible ways it can move, so the less freedom it has to explore different micro states – think of all the ways you could move your arm if it’s unconstrained versus in a sling. And if it can’t occupy many micro states, chances are you’re more likely to find it in just a couple orientations – it’s less random. Similarly, when molecules are bound together they have limited rotation so there are less different rotational “micro states.” In addition to rotating different ways, molecules can hang out in different locations.
Imagine you’re trying to predict where each gas molecule in a room will be – the more possible places it *could* be the harder the prediction will be. Simplifying things for now and not accounting for rotation and stuff, just think of trying to predict where in a room a gas molecule will be (by the door, the window?)
If you have a elemental gas like neon that likes to hang out as a single atom (it’s monatomic), each neon molecule can whizz around and if you were to put a tracker on one of them you’d have a really hard time predicting where it’d end up – because it’s too random. Now try putting trackers on each neon atom… yeah, no…. But if you have a gas that likes to hang out in pairs (diatomically), like oxygen or hydrogen or nitrogen, you still have a lot of disorder, but you only need 1/2 as many trackers because if you know where 1 atom is, you know where its partner is – kinda like how if you know where your pinky finger is, your ring finger’s gotta be nearby.
Even if your arm isn’t in a sling, you need energy to move it around. And molecules need energy to move around too. It takes a lot of energy to be a gas, but a lot more energy if you’re a gas of a “sticky” molecule. In a gas, molecules are free to roam around independently from other molecules (for the most part). But in a solid, the molecules are close together and they’re restricted by not having enough energy to break free and escape, so solids have lower entropy than gases (liquids are in between) and if you add energy to a solid you can get it to become a liquid (melt) and then a gas (vaporize).
If the molecules in a liquid are attracted to one another, the molecules don’t just need enough energy to “fly,” they also need enough energy to break off the (weak) intermolecular attractions on their way out and not get pulled back in. The stickier the molecules are to one another, the more energy these transitions will require and thus the higher the melting and boiling points will be. This is why water is a liquid at room temp.
Noble gases – since they aren’t really attracted to anything as is – have a really easy time with this. So we find them as gases at room temp and those gases are typically considered “inert” or “nonreactive” because, having a complete outer shell of valence electrons they’re satisfied and not on the lookout for better opportunities. So, unlike other gases like oxygen & hydrogen, which can unpair to form unstable, highly reactive molecules, you can stick them places without having to worry about them reacting with other things and causing problems. So you can use them as space “filler” in things like welding setups where oxygen would be dangerous and you can use helium (which is both really light and non-reactive) in balloons & blimps.
This typical inertness isn’t to say that you can’t get a noble gas to do anything – you just need a lot of energy to do so.
The reason the outer electrons are stuck where they are is that, as negative things, they’re being pulled in by the positive draw of the protons in the nucleus (opposites attracting and all). But the electrons are kinda like energetic dogs on leashes with the nucleus acting as the poor walker tasked with reigning them in. If the electron “dogs” tug hard enough they can stretch the “leash” & move further away – you add enough energy you can get electrons to venture out into another, higher energy, orbital. And if you add even more energy you can “break the leash” and strip electrons off altogether.
As you go down the periodic table, it becomes easier and easier to remove an electron, because they’re getting further and further away from the pull of the protons thanks to all those extra shells acting as “shielding” – so it takes less energy to break off, so we say that these lower-down elements have “lower ionization energies” (e.g. it takes 2372.3 E/kL mol-1 to remove the first electron from He, but “only” about half that much (1037.1) for Rn). I say “only,” but noble gases as a whole have really high ionization energies compared to other elements because they have those full shells and if they lose an electron they won’t anymore.
When you remove electrons (assuming you started with a neutral molecule), you leave the reluctant giver with more protons than electrons, so it’s positively-charged (cationic) – it’ll travel towards a negatively-charged electrode (basically just a charged metal plate). And the electrons, attracted to the positively-charged electrode, flow the opposite direction, towards the negatively-charged electrode. And when you have flow of charged particles you have electricity! So, by ionizing the gas you can get it to conduct electricity, and we call ionized gas “plasma” – yep, just like plasma TVs or plasma in stars.
So you can fill a tube with a noble gas like neon or argon, with a positive electrode on one end of the tube & a negative electrode on the other, zap it with electricity to ionize the gas and that ionized gas (aka plasma) will conduct electricity to complete your light circuit. But where does the light itself come from?
If an electron absorbs enough energy to get to an “excited state” but not enough to break all the way off, it can still feel the protons’ pull and, having spent the energy needed to resist, it’s liable to fall back down to its “ground state” and when it does this, it releases the energy it had absorbed to get up there. A lot of times this release happens through “boring” stuff like bumping into other molecules who absorb the energy, or letting off some heat. BUT if the difference in energy between the ground state and the excited state is *just right*, the energy can be released like a little packet of energy that travels in waves, is called a photon, and is what makes light.
Light is ElectroMagnetic Radiation (EMR) – when people think of “light” they usually think of “visible light,” but that’s just a teeny little piece of the EMR spectrum squeezed in between infrared to the left (lower energy, longer wavelength) and ultraviolet to the right (higher energy, shorter wavelength). Light can be thought of as little packets of energy (photons) traveling as waves through space. The only thing different between the light you see and the light you don’t see is the amount of energy in their photons.
Since the differences in energy between different ground and excited states is different for different molecules, the *just right* amount is different for different molecules, so the photons released by different molecules have different amounts of energy. And the more energy a photon has the faster it’ll move – but all light travels (linearly) at the same speed (the speed of light), so in order to not race ahead of their peers, more energetic photons have to wiggle up and down more, so the peaks of their waves come closer together – we say they have a higher frequency (f) and shorter wavelength (λ).
The photons released this way by neon are orangey-red. Argon gives you a bluish purpley color. To get other colors you can mix and match gases or, as is more common, you can paint the walls of the tube with fluorescent paints. The molecules in these paints will absorb the photons released by the noble gas – this will excite them – then they’ll fall down and release light. So you see the color of the dye’s emission wavelength instead of the noble gas’. And this gives you a lot more options, since in fluorescence, the light that is released is of a different wavelength (emission wavelength) than the light initially absorbed (excitation wavelength).
This can happen because different molecules have different electronic orbital layouts and electrons can be at different “vibrational levels” and stuff so the same amount of energy can get electrons to places where they’ll fall different amounts from. Plus, some of the energy absorbed is lost as heat, etc. So the emission wavelength is lower energy, longer wavelength (red-shifted) compared to the excitation wavelength and the difference between the 2 is called the Stokes shift. http://bit.ly/fretandfluorescence
Neon lights have kinda fallen out of favor (except for their vintage appeal), largely replaced by LED (Light Emitting Diode) lights. In a diode you have 2 next-to-each-other “versions” of a material (often aluminum-gallium-arsenide (AlGaAs)) that, while mostly neutral, is “doped” with imperfections – atoms of other elements that interrupt that nice AlGaAs structure and have different numbers of electrons. And each side is hooked up to an electrode. The material on one side has a sprinkling of atoms with “extra” electrons. Electrons are negatively charged, so we say this is N-type material. And next to it you put a material that, while mostly neutral, has a sprinkle of atoms which are “missing” electrons, thus providing electron “holes”. We call this material P-type because it has the effect of positive charge. And when you provide electricity, hooking up the negative electrode to the N-side and the positive electrode to the P-side, the electrons can get repelled away from that negativity to seek out those positive holes. When they do, they become thermodynamically happier, so they fall to a lower energy level and, if the gap between the excited and un-excited states is an optimal distance, they release light. Tada!
That LED stuff is outside my field of “expertise” but I found this good explanation if people want to learn more: https://electronics.howstuffworks.com/led.htm