It’s possible my memory is faulty. I took high school chemistry in 1972 and Chemistry 101 in 1975. That is more likely than not being taught well.
It’s also possible that the approach to introductory chemistry in those days was not as comprehensive as today, or simply didn’t address isotopes as a primary topic. I seem to remember spending a huge amount of time on electron configurations, ionic and covalent bonds, the Ideal Gas Law, Avogadro’s number and the mole, oxidation and reduction reactions, endothermic and exothermic reactions, how to make hydrogen and oxygen from water. Isotopes–not so much.
At least they weren’t teaching us that atoms were little indivisible spheres.
It is (the list is 2, 8, 20, 28, 50, 82), but it’s a little surprising that it’s that magic. I suppose it hits a kind of sweet spot where the nucleus isn’t too big, but the atomic number is still high enough that there’s a lot of possible isotopes.
Well, that’s the thing. For all those topics, different isotopes are largely irrelevant. Chemistry is all about the number of electrons an element has, which is tied to the number of protons. The number of neutrons doesn’t make much difference, particularly as there’s usually one isotope that is a large majority of the naturally occurring material, as discussed above. So yeah, isotopes would get at most a passing reference in such a class.
Good point. I took A-level chemistry (AP high school, in US parlance) and some university chem (not heavy duty since my major was EE).
And I don’t recall much if any discussion of isotopes. Or radioactivity, for that matter… I don’t think I had even heard of Electron Capture decay until I came across it in my own reading later.
Where it comes up (and usually not until undergrad organic chemistry) is nuclear magnetic resonance. Then later with kinetic isotope effects. So I agree it makes sense not to have encountered it.
Every element has a neutron configuration that’s most stable. If it’s most stable then it’s least likely to decay into some other isotope or element. If it’s least likely to decay in into something else, then that will be the most abundant isotope.
So it’s not exactly right to to say “standard”, but there is a most-abundant isotope. Incidentally that’s why atomic weights are fractions. Protons and neutrons count as whole numbers for atomic weight puposes. So although there’s no such thing as .0291 of a proton or neutron of course, you can always expect a certain count of naturally occurring uranium atoms (a mole) to weigh about 238.02891 grams, because it’s not composed entirely of U-238, there’ll be a little U-235 and U-234 in there too.
So on average the molar weight of uranium can’t be exactly 238 unless you’ve taken some steps to separate out the isotopes other than U-238.
It just occured to me that there’s something wrong with that molar weight. If U atoms are mostly 238 with some small amount of 235 and a smidgen of 234, shouldn’t the molar weight be slightly less than 238 g? Or are the 92 electrons a significant factor?
I am probably the worst person to answer this but I believe that the 92 electrons plus binding energy account for the 0.02891. Because the the percentages of the other isotopes are so small (about 0.7%) they don’t bring it down much.
This raises the question that if carbon-12 is 12 Da (by definition) then where does the extra 0.05 come from in U-238. But I think my post just above answers that.
Many elements have more than one stable isotope. Not “more” stable than the others, just stable. All stable isotopes of an element will never spontaneously decay.
The relative abundance of isotopes for a given element depends on the stability of the unstable isotopes. But it also depends on decay chains, in other words, how often a specific isotope will be created as the result of the decay of another isotope. It further depends on the details of nucleosynthesis, the process in which heavy elements are created by fusion inside stars. For stable isotopes of the same element, these last two are the only mechanisms responsible for an uneven split. In other words, the relative abundance now is a function of the history of the universe.
To give an example, Wikipedia tells me that U 234 with a half life of 2.5 × 10⁵ years has a relative abundance of .005 %, but the much more stable U 236 with 2.3 × 10⁷ years only occurs in minute traces.
As a K-12 student in the 80s & 90s, my introduction to isotopes was in explaining radiocarbon dating. My takeaway was that C12 is regular carbon and C14 is the exotic isotope. That they’re all isotopes wasn’t clear to me until college.
Come to think of it, I probably also muddled the diamond/graphite/coal forms of carbon into isotopes for a while, too.
I agree with your other objections (the abundance of isotopes depends on what isotopes are initially created (by decay, stellar fusion or supernova) and how fast the isotopes decay), but the concept of “most stable” seems valid to me. The more energy it takes to make a nucleus decay the more stable it is, just as in the realm of chemical reactions the more energy required to disrupt a chemical bond, the more stable the substance is - even if both substances are inert if left alone).
A question that came to me, actually before I read this post, but it makes for a good segue:
How do we know that any given isotope is stable? Is stability something that can be derived from intranuclear forces, or is it observational - “we’ve never seen this isotope decay, so it’s stable”? And if it’s the latter, would it be possible that even the most stable isotopes will eventually decay, possibly all the way back into hydrogen after some hideously long time?
@Chronos can correct me, but I susoect that for simple nuclei, there are calculations demonstrating stability, but for more complicated ones, experiments are required. But given how many nuclei are in a macroscopic amount of an element, we can get a very good idea of whether an isotope decays from experiment (the lower bound on a proton’s life span is on the order of 10^34 years)
Protons and neutrons count as whole numbers for the atomic weight listed for any given isotope, like U-238. But for the number with a decimal on it, that’s the average mass of an atom, in AMUs. Even for a pure isotope, that’s not an integer, except for C-12 (and that’s only an integer because the AMU is defined as 1/12 the mass of that atom). Every nucleus has a binding energy, and it’s different for every one: That’s what enables nuclear reactions to release energy. (Actually, the mass of the protons and neutrons themselves are mostly binding energy, too, but that’s only relevant if you’re doing particle accelerator experiments.)
In particular, for hydrogen-1, hydrogen-2, helium-3, and I think one or two isotopes of lithium and beryllium, the only way they were made, in non-negligible quantities, is in the Big Bang, and therefore the relative abundances of those isotopes of those elements tells us a great deal about conditions during the early Universe, and in turn about the current Universe.
If you’re using energy to change an atom, that’s not a decay. A decay is what it’ll just do on its own spontaneously.
In principle we could determine it theoretically, if we knew enough about the details of the Strong Nuclear Force, and in at least some cases we can also determine theoretically that something should be stable based on various conservation laws, since there’s nothing that it could decay into. But there have been cases of elements which were previously believed to be stable, until extremely sensitive experiments detected trace amounts of decays.
For that matter, it’s possible that some of the conservation laws we know aren’t actually absolute, and then all bets are off. Most Grand Unified Models, which combine the strong and weak forces, predict that baryon number isn’t strictly completely conserved (nor lepton number), and that thus even the proton isn’t completely stable, but eventually decays to a positron and some combination of photons and neutrinos. If even the proton eventually decays, then presumably so do all nuclei. This has never been verified experimentally, but our experiments are not yet sensitive enough for it to be particularly interesting that it hasn’t been; plenty of versions of the models predict a lifetime far longer than what we could detect.