Radioactivity can be super useful in biochemistry, and today I want to tell you about it from a biochemistry (not hard-core chemistry and definitely not a physics-y) perspective. We can substitute normal versions of atoms like phosphorus (P) for radioactive versions. These radioactive versions look the same to the other molecules, and don’t affect the labeled molecule, but they give off energy. So we can track molecules like DNA & RNA as they go about their day!

text adapted from past posts, video new 

Molecules are made up of even smaller things, atoms, which are made up of even smaller things – subatomic particles. These subatomic particles include protons, neutrons, & electrons. Protons are positively-charged and they hang out with neutrons (which are neutral) in a dense central nucleus. The electrons are negatively-charged and they whizz about the nucleus in an “electron cloud” – you never know exactly where one will be, but we can describe “orbitals” where they’re most likely to reside.⠀

If you change the number of protons you change the type of element because the # of protons defines the element – for example, carbon (C) has 6 protons, oxygen (O) has 8, nitrogen (N) has 7, and phosphorus (P) 15.⠀

BUT the number of neutrons and electrons can be different and the element is still that element. It’s like how you can gain weight or lose weight but still be you. Speaking of weight, gaining neutrons makes an element heavier (and losing them makes them lighter) and we call atoms of the same element but different numbers of neutrons isotopes. But gaining or losing electrons doesn’t really affect the weight because they’re so light – kinda like trimming your hair or growing it a little. Changes in electron number don’t change the weight, BUT they DO change the charge.⠀

Protons & electrons have equal (but opposite) charges, so an atom is neutral if # protons == # of electrons. But if they have uneven #s they *do* have charge and we call such atoms ions. if # protons > # electrons you have a net + and we call it a cation. If it’s the other way around & you have more electrons than protons, you have a net – & we call it an anion. These different forms of the atom can have very different reactivities.⠀

BUT changing # of neutrons does NOT change the charge because neutrons aren’t charged – it’s like adding or subtracting 0’s. And they react with other molecules just like the “normal version.” ⠀

BUT these are less stable than the “normal version.” If you think about the layout of an atom, it’s kinda weird – opposite charges (like that of the proton (+) & electron (-) )attract, and like charges repel… and you have a ton of positive charge concentrated together with the counterbalancing charge spread out around it. ⠀

In order to keep those protons tight together, you need some “glue” in the form of the strong nuclear force. This comes from neutrons and you want to have a good balance of protons & neutrons. Sometimes the arrangement’s kinda awkward, but you can shuffle the nucleons around a bit to get comfier without actually changing the # of protons or neutrons. ⠀

In gamma decay the arrangement of protons & neutrons changes – they shuffle around a bit to reorganize into a comfier position and, give off squirmy energy on their way to reaching that more relaxed state. This energy is given off as gamma radiation. It’s not any actual particles moving, just some energy. Like x-rays on steroids.⠀

But other times, more drastic changes are required – changes that actually change the # and/or type of subatomic particles. There are 2 major kinds of such decay – alpha decay & beta decay. ⠀

In alpha decay, which typically happens in really big, heavy nuclei, an atom gives off an alpha particle 2 protons & 2 neutrons (basically it gives off a helium nucleus)⠀

atom ➜ atom 2 places “to the left” on the periodic table + He²⁺, which can pick up a couple electrons to become elemental helium, He. (if you’re used to chemical equations where you carefully balance the total charges in your equations, nuclear decay equations often seem “wrong” because they often ignore the “normal” electrons and just focus on the stuff going on in the nucleus)⠀

an example: Uranium-238 ➜ Thorium-234 + He⠀

This isn’t very useful for radiolabeling, especially since one of the main benefits of radio labeling is that it can sneakily replace something that’s naturally there and you’re not gonna find uranium in your DNA! But you will find phosphorus… You’ll find it in places like RNA & DNA (where it’s in every letter (nucleotide)) and, while none of the protein letters (amino acids) have phosphorus in them, phosphorus *can* get incorporated into proteins after they’re made when proteins called kinases take off part of an RNA letter (ATP) and stick a phosphate group on them – this phosphorylation can change the protein’s shape & activity, and you can learn more about it here:⠀

But for now let’s look at how we can make that phosphorus “stand out.” Alpha particles are large, slow-moving decay products – easy to shield against, but what I work with is more energetic – the type of radiation I use is called beta decay. The actual details of the subatomic physics are kinda weird, but the basic idea’s fairly intuitive – if you have too many protons compared to neutrons, swap a proton for a neutron. And if your problem’s too *few* protons, do the opposite (swap a neutron for a proton). And let off charged particles to make things balance out. ⠀

Let’s look at how I use it. I use a radioisotope of phosphorus, because I can have it “replace” the normal phosphorus in the phosphate at the RNA’s 5’ end (you can also label DNA this way) – and if you want to study those kinases we talked about, you can use “hot ATP” in which the P that gets added has a radioactive version of phosphorus instead of the “normal” form ⠀

What do I mean by normal? If you were to take some random phosphorus-containing molecule and measure the mass of that phosphorus, chances are you’re going to get 31 atomic mass units (amu). 15 amu from protons (because phosphorus ALWAYS has 15 protons) and 31-15 = 16 neutrons (remember the electrons are too small to count mass-wise). That doesn’t mean you’ll *never* randomly find a phosphorus with greater or fewer than 16 neutrons, but 16 is by far the most common. So if you look up P on the periodic table you’ll see it has an average atomic mass of 30.973 amu. ⠀

There’s a reason you’re most likely to find 31P – it’s the most stable. Some elements have more than one stable form so you might chance upon a different isotope of it (like “normal” C (12C)’s friend 14C which comes in handy when scientists want to date really old stuff), but, for P, 31 is so dominant, natural P is considered “100%” 31P, with “trace amounts” of a couple others…⠀

“You” *can* add more neutrons (1 more for 32P & 2 more for 33P) BUT P’s going to get “overwhelmed” by neutrons and make some subatomic changes to get to a more stable state – accompanied by the release of radiation.⠀

The isotope I work with is 32P (usually pronounced “P thirty-two”). You’ll also see it written as P-32, Phosphorus-32. or 32P. If you compare this to “normal” P it has “too many” neutrons. If the problem is too many neutrons for how many protons you have, why not swap a neutron for a proton?⠀

When P32 decays it does so through beta-minus. It lets off something called a beta particle⠀

P32 ➜ 32S + e- + antineutrino⠀

It gains a proton!!!! And since the # of protons defines an element, it’s no longer phosphorus – now it has 16 protons and is thus sulfur! This 32S is stable and happy so it stays as is. But this proton-gaining causes you to increase charge. But physics laws tell us charge has to be conserved, so an electron is given off as well. You also give off something called a neutrino which is a weird little thing that’s electron-like in terms of being really tiny & light but it differs from electrons in that it is not charged.⠀

Sometimes, an atom’s unstable for the opposite reason – it has too *few* neutrons. Like 30-P. So it swaps a proton for a neutron instead of the other way around. And thus it has to let off some positive charge to compensate. This is called beta-plus decay. Beta-plus decay is aka positron emission because it gives off a positron – it’s like an electron in terms of tiny-ness but it’s postively-charged. And it also gives off a weird little particle thing, this time a neutrino.⠀

Different nuclear isotopes decay at different rates, which we measure as HALF-LIFE – the time it takes for 1/2 of it to decay. For P32, the one I use, it’s 14.29 days, so I have about 2 weeks before half of it will be useless. In the meantime, we can detect the radiation they give off.

The great thing about radiolabels is they’re the same size and have all the same binding & biochemical properties, just a different number of neutrons. So we can use it to, for instance, see if a protein binds to it (which you can see with gel shift assays (EMSA)) or whether it gets cut by something (which you could see by any old PAGE). You can also end-label RNA or DNA with fluorescent labels, which are easier to work with BUT these labels add bulk & change the properties. Plus, radioactivity-based methods are SUPER SENSITIVE – you can detect minuscule amounts.⠀

more on how I do it in practice:

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