A gas- Uranium Hexafluoride.
Of course you do.
Of course, it wouldn’t work if you used just U238, just like you can’t turn water into whiskey no matter how much you distill it. Which is where that fraction of a percent of U235 in natural uranium comes in.
After you’re done, you’ve got some amount of enriched uranium, with a high concentration of 235, and many times that amount of depleted uranium, which is even more overwhelmingly 238 than the natural stuff. Of course, even depleted uranium still has some 235 and other isotopes in it, just low enough that it’s not worth trying to get the rest of them out. From the point of view of the nuclear program, depleted uranium is a waste product, but since we end up with a bunch of the stuff anyway and since it’s a very high-density material, and density is useful for projectiles, it gets used to make some kinds of munitions. It also has some other properties that happen to be useful for armor-piercing specifically, but even absent those, it’d probably still be used for something.
Which leads to a lot of depleted uranium dust in war zones, which is now believed to be responsible for a variety of health problems in veterans (mostly due to the heavy metal toxicity, not the radioactivity).
You know about the periodic table, in which similarities between elements exhibit a repeating pattern. This is chemistry, which is all about the shells of electrons, that is, the outer parts of atoms.
But nuclei care more about their internal arrangements, and nuclear physics has its own periodic table of sorts (I’m taking poetic liberties here), a diagram of proton and neutron numbers. If you check out the linked Wikipedia article (about relatively stable big nuclei) you will see this diagram referenced again and again in different ways. It’s almost like another chemical universe, in my view at least.
Yeah, it’s a lot more complicated to study in part because we don’t understand the strong nuclear force nearly as well as we do the electromagnetic force, in part because the strong nuclear force has a very short effective range, and in part because the particles are spread pretty uniformly within the nucleus, rather than having a nearly-point-source nucleus surrounded by far-spread electrons like for electron chemistry.
OK, help me correct my understanding. My elementary chemistry lessons wanted me to think that each element has a “standard” number of neutrons, and that variants of that were isotopes. For example, hydrogen has no neutrons; that is the baseline “standard” element. Add a neutron and you get deuterium, an isotope. It seems that we go to great pains to distinguish the two, rather than saying “Here are two isotopes of hydrogen.” Is this an oversimplification so as not to confuse high school students or college freshmen?
I think so. I would definitely say “here are two isotopes of hydrogen”.
Certainly, an elementary chemist (like a high school teacher) would paint a picture where one isotope is “standard” and the others are weird and irrelevant… for their purposes (which is elementry chemistry, entirely about electrons and bonds and such).
But it’s not really as simple as that. There are elements in which the abundance difference between two natural-occuring isotopes is a single-digit percentage. You’re almost as likely to find an atom of isotope “a” as isotope “b” in any sample.
Generally, the differences in chemistry behavior are slight, almost to be point of negligible. (But with a few exceptions.)
Either you were not taught well, or your memory could be faulty?
The identity of an element is determined by the number of protons in the nucleus.
The atomic number… as established by Mosely:
But all elements except hydrogen have some number of neutrons in the nucleus, otherwise they wouldn’t be stable: the positive electric charge would blow them apart except that the strong nuclear force holds them together. We don’t understand the strong nuclear force very well, mind you, but it clearly needs neutrons to work.
Many elements have a few different variants (isotopes) with the same number of protons but different numbers of neutrons. Some of these are stable and some aren’t, the unstable ones decay radioactively.
There are cases where there is only one proton/neutron mix which is stable: Fluorine 19, for example. It’s probably just a matter of convention whether we call that an ‘isotope’ or not?
In general, stable elements have one (or maybe two) configurations that are stable (hydrogen and deuterium, lithium-6 and lithium-7), while most in heavier elements that have multiple configurations (‘isotopes’) the off-nominal isotopes are unstable, often radically so. This is actually the basis for radiometric dating, where the ratios between isotopes can be used to establish the age (within statistical bounds) at which the body containing them stopped absorbing material from the environment; the most common radiometric dating method is radiocarbon dating where organic materials absorbed atmospheric carbon dioxide with a percentage of 14C produced by upper atmospheric interactions with cosmic ray particle compared with the stable 12C and 13C primordial isotopes, and as the 14C decays (half-life of 5,700±30 years) the ratio goes down.
Stranger
Including hydrogen.
And deuterium (2H) is naturally occuring and stable.
And technically “zero” is a number, so even protium (“regular” hydrogen) has a number of neutrons. (But I’ll concede that’s probably nitpicky beyond reasonable limits.)
Think of it like shape. You might refer to a piece of paper, say, as being cut into a shape, but every piece of paper still has a shape. It’s just that for some of them, the shape is a typical, boring rectangle. Just so, every atom is an isotope; some of them are just more boring than others.
It’s not quite as simple as strong force versus electromagnetism, though that’s very relevant for large nuclei. For instance, in a hypothetical nucleus of Helium-2 (that is to say, nothing but two protons), the strong force would be considerably stronger than the electromagnetic… except that in that configuration, the strong force would also be repulsive. Meanwhile, helium-3 (two protons and one neutron) is stable, while hydrogen-3 (one proton and two neutrons) isn’t, even though the former has repulsive electrical forces and the latter does not.
Very roughly speaking, the strong force “likes” to have an even number of protons and neutrons, and the same number of each. That’s why the helium nucleus, or alpha particle, is remarkably stable, and why most of the lighter elements have close to the same numbers of both. But the strong force is short-ranged enough to mostly only be relevant between particles directly adjacent to each other, and in a large nucleus, they won’t all be directly adjacent, which allows electromagnetism to be more relevant than for smaller nuclei. And the electromagnetic force would “like” to have as few of the particles be charged as possible, with as many neutral particles in between as buffers. So for large nuclei like uranium, the stablest isotopes have significantly more neutrons than protons.
It can be more. Lead, for instance, has four stable isotopes. Natural lead is roughly a quarter each Pb-206 and Pb-207, half Pb-208, and a bit more than a percent of of Pb-204.
There’s always some isotope of any given element that’s the most common, and usually that one is over 99% of what’s found in nature, but as seen with lead, there are exceptions where the abundances are comparable to each other. Another notable one is chlorine, where about 3/4 is Cl-37, and 1/4 Cl-35.
Tin has ten stable isotopes, all with an abundance between 0.3% and 32.6%.
And then there’s poor Technetium, which despite being smack in the middle of the periodic table (i.e., not one of those elements with giant nuclei), has zero stable isotopes. It just got… unlucky with its proton count (43). Promethium, too (61).
There are 251 stable isotopes among the first 82 elements, so the average is a bit closer to 3. Though without checking, the median might be more like 2, since it looks like there are a handful of special elements that drive up the average.
Technetium’s problem, fundamentally, is that it has too many stable neighbors. There’s a general principle that there’s at most one stable isotope of any given mass number, because of all of the isotopes with that mass number, one will have the lowest actual mass (not quite the same thing as mass number, since binding energy varies), and anything else with that mass number can turn into that through beta decay or something similar (electron capture, positron emission, etc.). And in quantum mechanics, if something can be in two different energy levels, it will eventually make it to the lowest one, no matter how high the barriers between them.
Well, it’s a little of each. It has a notably un-magical proton count. And it’s odd (less stable), which itself implies that its neighbors will be even (more stable). So at least in part, “stable elements have unstable neighbors” and “unstable elements have stable neighbors” are one and the same.
Most elements get several chances at having a stable isotope. It just needs to get lucky for one of them, where it has the lowest binding energy for all possible decay chains. Technetium got its chances, too… and blew it on all of them. Without checking, I’d suspect that at least some of its isotopes have lower energy than some of the isotopes of its neighbors. But what matters is the isotopes in the decay chain. If you have the highest energy in all of those pairings, then you’re screwed.
Look around the table. If you can’t spot the sucker unstable element, you are the sucker unstable element.
Here is a wikilist of all the elements and their isotopes: It is in German, but that should not be a problem, as it is in the form of a table with the right units which show what the number refers to:
It starts with the first ten elements up to neon (including the neutron, which is unstable when not bound). For more elements you have to click your way though the periodic table.
I don’t see what else you could call it. There are several isotopes of F ranging from 14F to 31F (according to the list above) with varying (in)stability, but why not call the only stable one an isotope? Isotope just means “iso” = same and “topos” = place (in the periodic table, which is determined by the number of protons, not neutrons). It says nothing about stability, and when several isotopes are stable (like lead and tin) or when all are unstable (poor technetium and promethium were mentioned by Dr.Strangelove) we call them isotopes too. So why not call the only stable one among several unstable ones isotope too? They are all in the same (iso) place (tope).
I ran some stats on Wikipedia’s list of the 251 stable isotopes. Mean is 3.14 stable isotopes per element, median is 2.00, and the mode is 1. Here are the number of elements with a given stable isotope count:
| Isotopes | Elements |
|---|---|
| 10 | 1 |
| 7 | 5 |
| 6 | 7 |
| 5 | 11 |
| 4 | 9 |
| 3 | 5 |
| 2 | 16 |
| 1 | 26 |
So a bare majority of elements do have 1-2 stable isotopes, with 1 in the lead. But there’s an interesting bump at 5 and a dip at 3. Also a bit strange that tin is so far out in the lead with 10.
Maybe 50 is a magic number of protons? I dunno.