Subject: Can someone please explain?
Posted by: Mixamatosis
Date: Jan 21 17
I've read that it's dangerous to mix ammonia and bleach. Variously I've read that it can produce deadly cyanide gas, chlorine gas (which is said to be bad for you) and even explosions.
However swimming pools are kept fit for use with chlorine, and our urine contains ammonia but then we may clean toilets with bleach. Also many cleaning products contain either ammonia or bleach and it would be easy to use them unthinkingly in combination.
How is it that people aren't generally harmed by these dangers when swimming in swimming pools or doing daily cleaning, or are we being harmed at low level and is the harm cumulative?
Posted by: Mixamatosis
Date: Jan 21 17
I've read that it's dangerous to mix ammonia and bleach. Variously I've read that it can produce deadly cyanide gas, chlorine gas (which is said to be bad for you) and even explosions.
However swimming pools are kept fit for use with chlorine, and our urine contains ammonia but then we may clean toilets with bleach. Also many cleaning products contain either ammonia or bleach and it would be easy to use them unthinkingly in combination.
How is it that people aren't generally harmed by these dangers when swimming in swimming pools or doing daily cleaning, or are we being harmed at low level and is the harm cumulative?
526 replies. On page 25 of 27 pages. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27
brm50diboll 
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So when discussing radioactive decay, we should mention the various modes of decay and what factors influence which mode is followed. Most people have at least heard of alpha, beta, and gamma rays even if they don't know what they are. But there are actually several more modes of decay beyond these three. I think it useful to discuss these modes of decay. Let's begin with alpha decay. Alpha rays are actually Helium-4 nuclei (2 protons and two neutrons bound together *without* the electrons, and thus with a charge of +2) ejected at high speed from a parent radioactive nucleus. They quickly acquire the missing two electrons to form helium, which is why helium exists in significant quantities underground - it is produced by the decay of natural radioactive elements in the earth's crust, primarily uranium and thorium. Alpha rays are not very penetrating, only passing through a few cm of air, a sheet of paper, and the top layer of skin. However, if an alpha source is taken internally, it can do extreme damage, especially to the bone marrow and its production of blood cells. Alpha rays are produced primarily by the decay of heavy isotopes beyond lead in the periodic table, isotopes with approximately a 1.4 to 1 neutron-to-proton ratio. As previously stated, natural thorium and uranium are the main sources of alpha rays, but they are also produced by many of the more highly radioactive intermediate decay products of those two elements, such as Ra-226, Rn-222, and Po-210 as examples. They are also produced by several of the artificial elements beyond uranium, particularly many isotopes from neptunium to californium. Alpha rays were utilized by Lord Rutherford in his famous gold foil experiment that led to the discovery of atomic nuclei. Because alpha rays are matter, they not part of the electromagnetic spectrum, which is radiation composed of photons. They were called alpha rays because they were the first type of ionizing radiation discovered. Alpha decay has the effect of producing a daughter nucleus which has an atomic number which is two less than the parent nucleus, and a mass number which is four less than the parent nucleus. For example, uranium-238 decays by alpha emission to thorium-234. The atomic number has dropped by two (U is 92, Th is 90), and the mass number has decreased by four, from 238 to 234. It should be noted in the this example, like many other examples of alpha decay I could have chosen, the daughter isotope is in itself radioactive and decays further. In fact, thorium-234 is much more unstable (shorter half life) than uranium-238, which makes it considerably rarer. In general, the shorter the half life of the isotope, the smaller amount of it exists in nature, even though it is produced in nature by the decay of heavier isotopes. This leads to so-called "decay series" of isotopes, with the three most prominent naturally occurring decay series beginning with U-238, U-235, and Th-232, all of which are alpha emitters, although many of the isotopes they produce further down in the decay series produce different types of radiation other than alpha rays. A topic for another day. Reply #481. Mar 15 21, 2:16 PM |
| brm50diboll
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Now onto beta decay. Beta decay occurs when a neutron in a radioactive nucleus, under the effect of the weak nuclear force, changes into a proton, an electron, and an electron antineutrino. The electron and the electron antineutrino are both ejected from the nucleus at very high speed, but the proton remains in the nucleus. The simplest example of beta decay is in the decay of the free neutron itself which I described when I was discussing the action of the weak nuclear force. But in the case of nuclei that contain protons, what properties of such nuclei make them decay by the beta decay mode? That is a complicated answer that deserves an essay all to itself. But for now, nuclei that contain an unusually high neutron-to-proton ratio tend to undergo beta decay. What is "unusually high"? Well, that's the sticking point that I'm going to gloss over today and fill in the details someday later. For the lighter nuclei (hydrogen through calcium), the neutron-to-proton ratio is about 1:1. After that, the neutron-to-proton ratio slowly creeps up until it gets to about 1.4:1 for the heavier elements in the region of lead in the periodic table. Whenever that ratio is too high for the region of the periodic table that the given nuclide is in, beta decay tends to occur. Let's consider a lighter element, one of the most famous radioactive nuclides in it: radiocarbon, or carbon-14, which is used by scientists to determine the age of relatively recent fossils (35,000 years ago or more recently). Natural carbon is actually stable, not radioactive, and consists of about 98.9% carbon-12, which has a nice 1:1 neutron-to-proton ratio, and 1.1% carbon-13, which is also stable despite a neutron-to-proton ratio of 1.16:1. But radiocarbon is radioactive and decays by beta decay mode, and has a neutron-to-proton ratio of 1.33:1, too high for the early part of the periodic table. Radiocarbon has a half-life of 5730 years and *does* occur naturally, but only in *extremely tiny* proportions, about one part per *trillion* of natural carbon. Carbon-14 is produced by the action of cosmic rays in our atmosphere, and it gets into the food chain, so all living organisms contain a very tiny amount of this radiocarbon. Despite being so scarce, it is easily detectable by radiation counters. When C-14 decays, one of its eight neutrons changes into a proton. In doing so, an electron and an electron antineutrino are ejected at very high speeds. The electron antineutrino is essentially a "ghost" particle that doesn't do much, but the high speed electron that is ejected does some damage and is known as the *beta ray*. Beta rays are more penetrative than alpha rays, and will go through several inches of flesh or sheets of paper. As for the nucleus, what has happened is essentially a neutron has changed into a proton. What this means is that beta decay does not change the mass number at all, but it *raises* the atomic number by one. Carbon-14 has eight neutrons and six protons. After it beta decays, it changes to nitrogen-14, which has seven protons (note the increase by one) and seven neutrons (note the decrease by one). N-14 is stable, unlike the C-14 it came from. In fact, nitrogen-14 is the primary isotope of nitrogen in nature and has a 1:1 neutron-to-proton ratio. What beta decay does is reduce the neutron-to-proton ratio, moving the nucleus towards stability. In all cases of beta decay, this is what happens. Here's a few other relatively famous examples of beta decay: K-40 changing into Ca-40, Co-60 changing into Ni-60, and Pa-234 changing into U-234. I think I will go into some of those gritty details of what exactly makes a neutron-to-proton ratio "too high" next time. Reply #482. Apr 08 21, 5:13 PM |
| brm50diboll
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A full discussion of the factors that determine which isotopes are radioactive and, if so, what their modes of decay are require a deep look at quantum mechanics and the energy levels of atomic nuclei. But I'm not going to get into those. Nevertheless, even at a qualitative level, interesting things can be noted. For one, even numbers have greater stability than odd numbers. Even numbers of what, you may ask? Even numbers of protons and even numbers of neutrons. In addition, there are so-called "magic numbers" associated with stability in nuclei. Let's look at the even number effect first. There are four possibilities for an isotope: Even-even, Even-odd, Odd-even, and Odd-odd. Let's look at each of these possibilities one at a time. First, there is the one with the greatest associated stability: Even-even. This means that the number of protons in the nucleus is even, and so is the number of neutrons in the nucleus. For the lighter elements in the periodic table, the neutron-to-proton ratio also tends to be very close to 1:1 with one very big exception worth discussing: beryllium-8, which, despite having 4 protons and four neutrons, is not only not stable, but, in fact, *extremely* unstable, with an incredibly short half-life of about 10^(-16) seconds. Beryllium-8 almost immediately fractures into two helium-4 nuclei. This big exception is due to the unusual stability of He-4, as 4 is the first "magic number" of nucleons. Aside from Be-8, all the other early 1:1 neutron-to-proton ratio Even-even nuclei are stable: C-12, O-16, Ne-20, Mg-24, Si-28, S-32, Ar-36, and Ca-40. Ca-40 is last on this list, because after calcium, the neutron-to-proton ratio of stable isotopes begins to slowly drift upwards from 1:1. Even atomic numbers (even numbers of protons) tend to have multiple stable isotopes, unlike odd atomic numbers (odd numbers of protons), which usually have only one or two stable isotopes for those odd elements. Tin (Sn), element number 50, has ten (10) stable isotopes as 50 is another of the "magic numbers". Even for heavier even elements that do not have stable isotopes at a 1:1 neutron-to-proton ratio, they still typically have multiple stable isotopes at higher ratios. Iron-56, with 26 protons and 30 neutrons, is an Even-even isotope that is stable and well-represented in nature. It can be argued, but I won't explain it here, that Fe-56 is the "most stable" isotope of all. Back to beryllium for a moment: we are fortunate as humans that the Even-odd isotope Be-9 happens to be stable (the only stable isotope of beryllium), otherwise, due to the unusual previously mentioned extreme instability of Be-8, beryllium would not exist in nature at all. A very high percentage of *all* stable isotopes happen to be Even-even. Let's look at the next category: Even-odd, which means an even number of protons, but an odd number of neutrons. For the lighter elements, there are quite a few of these that are stable (though much less than the Even-evens), but the number of neutrons is always *greater* than the number of protons, even if only by a little bit, with only one exception I will save for discussion later. So, as an example, we previously mentioned the very stable Even-even isotope O-16, which makes up over 99% of natural oxygen. But oxygen has two other stable isotopes that exist in nature besides O-16, though they are much rarer. One is O-18, another Even-even. But O-17 is also stable, though very rare, and it is an Even-odd, with 8 protons and 9 neutrons. Note the number of neutrons does slightly exceed the number of protons, as is always the case for stable Even-odd isotopes. By contrast, the isotope O-15, with 8 protons and 7 neutrons, is radioactive and does not exist in nature, though it can be artificially created. Only one stable Even-odd isotope where the number of neutrons is *less* than the number of protons exists, and I will discuss that very special case out of order later. Next we have the Odd-even isotopes. Elements with odd atomic numbers typically have only one or two (at most) stable isotopes, unlike the even atomic number elements, which may have many more. A general rule for the Odd-even isotopes is again that the number of neutrons is higher than the number of protons in stable nuclei. For the lighter nuclei, the neutron number exceeds the proton number by *exactly* one in most cases. However, there is one interesting early Odd-even isotope worth discussing as special exception: hydrogen-1 (protium), with 1 proton but 0 neutrons, which happens to be the most common isotope of *all* in nature. It actually has *fewer* neutrons than protons and has a neutron-to-proton ratio of 0. The reason it exists is that with only one proton in the nucleus, there is no electric repulsion in the nucleus for the strong nuclear force to have to overcome, so H-1 can get away with having no neutrons at all. But once you add that second proton, now proton-proton repulsion makes having neutrons necessary for stability. Let's now circle back to that *one* exception I mentioned when I was discussing Even-odd: it was helium-3, with 2 protons but only 1 neutron, and aside from H-1, it is the *only* stable isotope where the number of neutrons is *less* than the number of protons. He-3 is indeed stable, but it is vanishingly rare in nature, only a few parts per million as compared to its much more familiar "big brother" He-4. After He-3, we never again see any stable isotopes where the number of neutrons is *less* than the number of protons. So F-19 (9 protons and 10 neutrons) exists and is stable (and happens to be the only stable isotope of fluorine, which is not surprising for an Odd-even), but F-17 (9 protons and 8 neutrons) is highly radioactive and does not exist in nature, though it can be artificially created. There are a few Odd-even isotopes that are stable for *two* isotopes of the same element. Examples include Cl-35 and Cl-37, K-39 and K-41, Cu-63 and Cu-65, and several other Odd-evens scattered through the periodic table. But we never see *three* or more stable Odd-evens for the same element, unlike the Even-evens, where multiple stable isotopes of the same element occurs frequently. And in the case of technetium, element number 43, despite being a relatively light element completely surrounded in the periodic table by other stable elements, poor technetium has *no* stable isotopes. For technetium, even its Odd-even isotopes of Tc-97 and Tc-99 are radioactive, so poor technetium makes a "hole" in the periodic table underneath manganese. Now we arrive at the most unstable combination of all: Odd-odd. There are very, very few stable Odd-odd isotopes, and the few that do exist are all at the beginning of the periodic table: H-2, Li-6, B-10, and N-14. And that's it. The next one in that sequence would be F-18, but F-18 is highly radioactive and does not exist in nature, and all naturally occurring fluorine is the Odd-even F-19 previously mentioned. Even if we look at that short list of four stable isotopes I mentioned, only N-14 is actually the major isotope of that element in nature. For the others, the Odd-even isotopes H-1, Li-7, and B-11 are more common in nature than their Odd-odd counterparts. So that makes nitrogen quite interesting, actually. It is the only element in the periodic table dominated in nature by an Odd-odd isotope. N-15 exists and is stable, but rarer than the Odd-odd N-14. Very odd, that nitrogen. But have you really looked hard enough, Brian? Are you *sure* no more stable Odd-odd isotopes exist in the rest of the periodic table? Really, really sure? Well, now that you mention it and are pressing the point, there is lanthanum-138 to consider. La-138 is *almost* stable. Almost, but not quite. It does exist naturally though and makes up 0.09% of natural lanthanum (the stable Odd-even isotope La-139 makes up the other 99.91% of natural lanthanum.) Lanthanum, element number 57, is deep enough into the periodic table that the ideal neutron-to-proton ratio is now significantly above 1:1. For La-138, it is indeed Odd-odd with 57 protons and 81 neutrons and may be mistaken for being stable. Except that it isn't quite stable, it has the extremely long half-life of 1.02 × 10^11 years, or 102 *billion* years, which is much longer than the age of the universe itself (13.7 billion years), so of all the La-138 that was ever created in our universe, only a tiny fraction has decayed. The radioactivity level of La-138 is so low it will escape notice altogether unless very carefully measured and does not present any significant danger to anyone exposed to it, and can be treated as practically stable. So a big fat asterisk is in order for La-138, which is Odd-odd: * To conclude, we see that isotopes with neutron-to-proton ratios of less than one are exceedingly rare and unstable and generally do not exist in nature except for the special cases of H-1 and He-3 I discussed separately. However, many of them can be created artificially in cyclotrons and when they are created, it brings up an interesting question: How do radioactive isotopes that have "too few" neutrons as compared to protons decay. We had previously said that "too many" decayed by the beta decay mode. So what about the "too few" case? I will deal with that situation next time. Reply #483. Apr 24 21, 1:27 PM |
| brm50diboll
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Time for another interlude. I have been fascinated these past several weeks by the ongoing eruption of the volcano in Iceland. Iceland is unusual geologically as it is both a seafloor spreading center (Mid-Atlantic Ridge, the boundary between the North American and Eurasian tectonic plates) and a geologic hotspot. This unusual combination results in Iceland being the only part of the Mid-Atlantic Ridge which is above sea level. The current eruption, unlike the one several years ago which shut off flights to Europe for weeks, has very little ash and instead produces rivers of lava which spectators flock to see and they are able to do so fairly safely. Not exactly Pompeii in 79 AD. Anyway:  https://m.youtube.com/watch?v=75IrgdA5W3sThe map of the growing lava flow is interesting. If the eruption continues long enough, the flows will eventually break through a pass and head to the sea. Lava reaching the ocean is always an interesting sight to see. Reply #484. May 10 21, 9:29 PM |
| brm50diboll
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Another interlude: There is a total lunar eclipse visible in the Western US in the predawn hours tomorrow, May 26. It is also visible in the entire Pacific basin region, but the time will depend on the time zones (and which side of the International Date Line one is on). I hope it's not too cloudy here tomorrow as I am planning to make an effort to see it despite it being so early in the morning here. https://www.npr.org/2021/05/24/999832188/how-to-watch-the-super-flower-blood-moon-lunar-eclipse-this-weekReply #485. May 25 21, 9:02 AM |
Catreona
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I've been rereading "Asimov's Guide to Science" which, though several decades old at this point, nonetheless contains remarkably clear, understandable explanations of the topics covered. Your explanations remind me of Asimov's, Brian. Your students are indeed fortunate. I hope they appreciate you. Reply #486. May 29 21, 3:28 PM |
| brm50diboll
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Thank you. What I have said in this thread regarding astronomy is very similar to what I have presented to my Astronomy classes, although my frequent use of videos and pictures in class helped even more. My actual presentations contain a good deal more humor than in this thread, however. Teenagers seem to like that. However, my more recent discussion of subatomic physics here is well above the level of even the AP Physics classes I have taught. That material is considerably above the college freshman level physics which AP Physics actually is, so I've never actually presented it in the form I'm doing in this thread in a classroom setting. It is material that particularly fascinates me - something I wish I could teach but know I never will. I slip in bits and pieces when I can. All science interests me, but some branches of it more than others. And since, as Galileo said (loosely paraphrasing), mathematics is the language of science, I have an especially intense interest in mathematics I have not really explored in this thread - not yet, anyway. If you think my discussion of quarks and leptons was confusing, just wait until I get started on complex variables. I could devote several posts just to applications of : e^(ix) = cos x + isin x, demonstrating results like ln (-1) = pi × i. (Just for starters) So much I wish I could talk about, but I just don't have the time. I don't know if I'll ever even get to mathematics, though. I still have a lot more subatomic particle physics left to talk about, and I'm moving through that much slower than I'd like. Please don't think my lack of discussion of chemistry and biology indicates a lack of interest in those areas, though. "A person's strongest dreams are about what he *can't* do." (From another thread of mine.) I've done biology and chemistry constantly. Next topic: What happens when an isotope with a neutron-to-proton ratio that is too *low* undergoes radioactive decay. Been wanting to get to it, but stuff keeps happening. Reply #487. May 29 21, 9:10 PM |
| Catreona
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Please, don't rush into Math on my account! :P Due to my disabilities and the consequent somewhat disorganized nature of my high school level education, I never even made it all the way through Algebra II. Never got to Trig, which I regret but can't see how to remedy at this point. It's probably my unsteadiness in Algebra that accounts for my ineptness with regard to Chemistry. Again, I have to thank that fantastic explainer, Dr. Asimov, for making Chemistry even remotely understandable. Physics, Astronomy and Cosmology OTOH I enjoy vastly, if as an eternal amateur. Haven't had any trouble following you there. You've talked about conservation laws, not so much about the various symetries. No doubt you'll get to those in time. It's a pity you don't have the time; you'd write killer quizzes! Reply #488. May 30 21, 8:12 PM |
| brm50diboll
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Thank you. Popular authors of physics (to some people that may sound like an oxymoron) are quite readable. I'm a fan of Brian Greene myself. The real details of the theories discussed involve some very high-powered math, but the books avoid most of that, as I try to do in this thread. Reply #489. May 30 21, 9:05 PM |
| Catreona
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I very much enjoy Brian Green's books. Besides Physics, they also taught me all I need to know about The Simpsons. Reply #490. May 31 21, 6:14 PM |
| brm50diboll
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Haha! The Simpsons! It helps to inject a little humor in scientific discussions. There's over 30 years of material in The Simpsons (though, to my tastes, most of the good stuff was in the 90s.) Reply #491. May 31 21, 6:24 PM |
Quiz_Beagle
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The 'Science of Discworld' books were very useful to me in explaining many concepts that had eluded me during formal education... Reply #492. May 31 21, 7:44 PM |
| Quiz_Beagle
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...and I'd also like to add my thanks for your posts! Reply #493. May 31 21, 7:54 PM |
| brm50diboll
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Thank you. Reply #494. Jun 01 21, 3:24 AM |
| Catreona
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There's a book, "The Science of 'The Hitchhiker's Guide to the Galaxy'" that's pretty good. Quiz_Beagle, isn't Discworld the one that's turtles all the way down? Never have read any of those books. Maybe they need to go on my already tottering 'to read' stack. Reply #495. Jun 02 21, 7:32 PM |
| brm50diboll
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Now to the topic of the mode of decay that occurs in isotopes where the neutron-to-proton ratio is too low for stability. Previously, in the case where it was too high, I mentioned that beta decay was the mode of decay and that beta decay has the effect of lowering the neutron-to-proton ratio. It is quite reasonable to assume a sort of "reverse beta decay" would have the effect of raising the neutron-to-proton ratio and therefore would be the mode of decay seen in the case under discussion today, and that assumption would be correct, but it is a little more complicated than that. For one, while there are numerous naturally existing radioactive isotopes where the neutron-to-proton ratio is too high and therefore undergo beta decay, it turns out a low neutron-to-proton ratio is extremely rare in nature. Such isotopes can be readily made artificially, of course, but, even then, what is seen is *two* distinct modes of decay for such cases: electron capture and positron emission. Both are forms of "reverse beta decay". In regular beta decay, a high-speed electron is emitted (the beta ray). In electron capture, an electron is *absorbed* into the nucleus, merging with a proton to turn it into a neutron. In order to balance lepton number, an electron neutrino is created and released from the nucleus. This neutrino creation is often overlooked because is very difficult to detect because neutrinos interact so poorly with other particles. Electron capture is generally seen in nuclei of low atomic numbers that have low neutron-to-proton ratios. Conversely, there is positron emission, found primarily in nuclei of high atomic numbers with low neutron-to-proton ratios. Positrons are the antiparticles of electrons. Positrons have the same mass as electrons, but a positive charge. Although stable, positrons are extremely rare in our universe because they attract electrons (which are plentiful) and when they collide, like all antimatter-matter reactions, they annihilate each other with the production of gamma ray photons (gamma rays to be discussed later). One might expect a similar reaction between neutrinos and antineutrinos, but since both are neutral and do not attract each other, neutrinos rarely collide with antineutrinos and so antineutrinos are actually fairly common in our universe, unlike positrons (although antineutrinos, like neutrinos, are difficult to detect.) Anyway, in positron emission, a proton in a nucleus changes into a neutron and releases a positron at high speeds (a sort of "anti beta ray"). To balance lepton number, an electron neutrino is also released, as in electron capture. Why does electron capture typically occur in low atomic number nuclei while positron emission occurs in high atomic number nuclei? It has to do with available mass-energy in the nucleus. Electron capture creates one new particle, the electron neutrino. Positron emission creates two new particles, the positron and the electron neutrino. When mass-energy is limited the electron capture mode is preferred because creation of particles requires mass-energy from the nucleus, and one particle requires less mass-energy to produce than two particles. Both modes of decay have the effect of turning a proton into a neutron, this raising the neutron-to-proton ratio towards stability. Here's some examples: The artificially created isotope titanium-44 has exactly a 1:1 neutron-to-proton ratio. But because titanium is after calcium, a 1:1 neutron-to-proton ratio here is *too low* for stability. So Ti-44 is radioactive with a half-life of 60 years and decays by electron capture to form Sc-44, which is also radioactive with a half-life of about 4 hours and decays by positron emission to form Ca-44, which is stable. Note that in both electron capture and positron emission, the effect of the decay is to produce a daughter isotope that has the same mass number, but an atomic number that has been *reduced* by one, the exact opposite of what occurs in beta decay. Because naturally occurring isotopes that have low neutron-to-proton ratios are quite rare, electron capture and positron emission are much rarer in nature than alpha, beta, or gamma decays. However, because such isotopes are readily synthesized in cyclotrons, these modes of decay are easily studied. Positron emission, in particular, has important medical applications. A PET scan, which is used to study physiology (such as brain metabolism), involves injecting an isotope which contains a positron emitter in a molecule used by a particular organ or tissue, then detecting the positrons it produces to create an image. Due to quantum mechanical considerations I would rather not get into here, not all radioactive decay occurs by only one mode. Sometimes a radioactive isotope may decay by more than one mode, with the probability of a particular mode being calculable. This leads to the rare but very interesting case where an Odd-odd isotope (notoriously unstable in most cases, as previously discussed, may decay by *either* beta decay or electron capture, two completely opposite processes. Consider the famous radioactive isotope potassium-40, which is Odd-odd. Its neighbors K-39 and K-41, which are both Odd-even, are stable, but poor K-40, because it is Odd-odd, is radioactive, though very weakly so. K-40 has such weak radioactivity that it has a very long half-life of 1.251 billion years, long enough for it to exist naturally, and, in fact, K-40 makes up 0.012% of naturally occurring potassium. All living things contain some potassium, an element necessary for life, and so we all are slightly radioactive for several reasons, one of them being K-40. Bananas, being rich in potassium, have detectable radioactivity by most Geiger counters. But what is especially interesting to me about K-40 is that it may decay either by beta decay or by its completely opposite process of electron capture. 89.28% of the time, it undergoes beta decay to form the stable isotope Ca-40 (Even-even, the last stable isotope with a 1:1 neutron-to-proton ratio). But 10.72% of the time, it undergoes electron capture to form the stable isotope Ar-40 (also stable, also Even-even, with a neutron-to-proton ratio slightly above 1:1). The presence of argon gas trapped in rocks is due to this mode of K-40 decay and is very useful in determining the age of rock formations in geology. Next time I'll look at gamma decay. Reply #496. Jun 04 21, 6:30 PM |
| brm50diboll
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Gamma rays are composed of photons - those chargeless, massless particles of pure energy I've discussed before. Gamma rays are generally considered the part of the electromagnetic spectrum that has the highest energy, highest frequency, and shortest wavelength of all electromagnetic radiation. This is not *quite* correct. Gamma rays cover a band of energies, and the low end of gamma ray energies overlaps a bit with the high end of X-ray energies. In fact, if we chose a gamma photon and an X-ray photon from that overlap that had exactly the same energy, they would be completely indistinguishable from each other in every way. Then what is the difference? The difference is how they are created. X-rays are emitted when electrons fall into inner energy levels. Gamma rays are emitted when there is a transition between *nuclear* energy levels. Nuclear energy levels are not as easy to diagram as electron energy levels, but they certainly exist, and transitions between high energy nuclear states and low energy nuclear states release gamma ray photons. These transitions are a form of radioactive decay. Like electronic energy levels where electrons normally occupy the lowest states available to them (the so-called "ground state"), in the nucleus, the quarks are generally found in their corresponding "ground states" most of the time. When they are in a higher energy state ("excited state"), that is unstable and, with a particular half-life characteristic of such an excited state nucleus, they decay to the ground state by release of a gamma ray photon. Since photons are bosons rather than fermions, they can be created or destroyed without violating certain laws that fermions must obey. However, since Einstein showed that matter and energy are interchangeable according to his famous equation E=mc^2, the release of a gamma ray from a nucleus results in a small decrease in the mass of that nucleus *without* altering the number of protons or neutrons in the nucleus *in any way*. Both the atomic number and the mass number of a nucleus is unchanged by gamma decay. Gamma rays are much more energetic than alpha and beta rays (reminder: alpha and beta rays, previously discussed, are radiation but *not* part of the electromagnetic spectrum, unlike gamma rays). Gamma rays pass through flesh, even bones, quite easily, doing damage as they pass. It takes several inches of lead to stop gamma rays. And no, "The Incredible Hulk" is not scientifically accurate. If Dr. David Bruce Banner were exposed to an "accidental overdose of gamma radiation", he would not transform into The Hulk. Most likely, he would get radiation poisoning and die. Not as interesting dramatically, unfortunately. But if a nucleus does not change either its atomic number or its mass number when it undergoes gamma decay, then how do we distinguish between the "before" and "after" nuclear states? And also, what causes gamma decay? If a nucleus is in its ground state (as is usually the case), it is denoted simply by its atomic and mass numbers, as we have been doing when I write something like C-14, which is a carbon (atomic number 6) isotope with 6 protons and 8 neutrons (mass number 14). But, in the case where a nucleus happens to not be in the ground state but in an excited state, then the designation "m" is added to indicate the excited state. This the distinction between, for example, Tc-99 and Tc-99m, both of which are radioactive, but are quite different from each other. I've mentioned technetium before, an artificial element peculiar in that it appears fairly early in the periodic table (#43) and is completely surrounded by naturally occurring elements. Traces of Tc do exist due to low level nuclear processes occurring in ores of molybdenum or ruthenium (neighboring elements), but, for practical purposes, almost all the technetium on earth (and there is actually quite a bit of it now - tons, in fact) is manmade. There are no stable isotopes of Tc, but the most stable happens to be Tc-97. Tc-98, and Tc-99 are also relatively stable with fairly long half-lives. But I want to focus on the difference between Tc-99 and Tc-99m. "Meta" states can be created when a radioactive isotope decays by alpha or beta decay *without* releasing sufficient energy to take the daughter isotope to its ground nuclear state. Then that daughter nucleus is in an excited state. Technetium isotopes can be obtained in commercial cyclotrons fairly easily by bombarding ordinary molybdenum with neutrons. The new molybdenum isotopes with extra neutrons tend to be radioactive and quickly decay (in minutes, hours, or days) by beta decay into technetium isotopes, which are also radioactive, but much slower to decay, and thus can be extracted from the molybdenum source. (A shout out here to my old Texas A&M prof Dr. Schmitt, a nuclear chemist, who inspired me to read further on these topics - wherever you are, sir). Anyway, depending on exactly how it was synthesized, some newly created Tc-99 ends up as just that, but some ends up in the excited state as Tc-99m, which is quite different. Let's look at the difference. Ordinary Tc-99 has a half-life of 211,100 years, quite long, though vastly too short to have persisted from the original formation of the earth 4.5 billion years ago. Nevertheless, with such a long half-life, kilogram quantities of Tc-99 can and have been synthesized, and none of that is going to disappear in a human life span. (By comparison, Tc-97, the "most stable" isotope of technetium, has a half-life of 4.21 million years, still too short to be from primordial earth, but long enough to be essential permanent compared to the human life span). Tc-99 decays by beta decay into stable Ru-99, but very slowly, given its half-life of over 200,000 years. By contrast, Tc-99m has a half-life of only 6.01 *hours*, far too short to isolate in macroscopic quantities, but only as a tracer. Hospitals and medical clinics use Tc-99m for various "technetium scans", including scans of cardiac function. These are very useful. A freshly prepared compound containing a few atoms of Tc-99m is injected into a patient and then a scan is made from the gamma rays it produces as it decays into ordinary Tc-99. Because only a few atoms were used, the Tc-99m involved does not produce enough radiation to be dangerous, and, a few hours later, it is all gone, turned into Tc-99, which, although also radioactive, is of such low radioactivity compared to Tc-99m that it is negligible, and that Tc-99 is mostly cleared out of the body anyway in a few days and any atoms that do remain are unlikely to decay in the course of the lifespan of the patient, given the over 200,000 year half-life of Tc-99. Gamma ray decay is useful for a variety of medical nuclear imaging and radiotherapy applications (such as treating cancer), not just Tc-99m. Although natural uranium and natural thorium are not, in and of themselves, gamma emitters, many of the radioactive daughter isotopes in their decay sequences are meta states and emit gamma radiation. Consequently, uranium and thorium ores and even the processed metals themselves, do produce a small amount of gamma rays as a result of these daughter isotopes. Gamma rays are also produced by nuclear reactions in celestial objects, such as stars, supernovas, quasars, etc. Space is full of gamma rays and we have even designed "gamma ray telescopes" to image celestial objects by their gamma ray emission. The Earth's atmosphere absorbs those gamma rays, so the gamma ray telescopes are in orbit. The level of gamma radiation in space is sufficiently low that astronauts, even after several months in space, do not absorb enough of it to be significantly harmed (loss of calcium from bones due to weightlessness is actually a much more significant problem for astronauts than gamma rays.) Nevertheless, since there is no effective shielding from gamma rays, we do not know what *years* of exposure to gamma rays in space might do. Jeff Bezos and Elon Musk might want to consider gamma rays before they book themselves or other billionaires on flights to Mars. Reply #497. Jul 09 21, 9:28 PM |
| brm50diboll
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I've been away too long from this thread. There are other modes of radioactive decay beyond those I've discussed in the last several entries, but they are quite rare, and generally involve select specially created artificial isotopes. However, for completion's sake, an overview of some of these modes is worth mentioning. Consider the following: proton emission, neutron emission, and spontaneous fission. Proton emission has the effect of moving the daughter nucleus to a higher neutron-proton ratio, which positron emission also does, and, in fact, positron emission is almost always energetically favored over proton emission and therefore isotopes with low neutron-proton ratios will generally decay by positron emission rather than proton emission except in rare special circumstances. It turns out that the light atomic weight and unusual stability of the main stable and naturally occurring isotope of helium, helium-4, is responsible for some of these exceptions. Natural stable lithium consists of a mixture of the two stable isotopes lithium-6 (odd-odd, but very early in the periodic table, as previously discussed) and lithium-7 (odd-even). No other isotopes of lithium exist in nature and the artificially created ones are difficult to synthesize and quite short-lived. Nevertheless, lithium-5 is known to be possible to create, though it is incredibly short-lived. Not surprising, as it has a neutron-proton ratio of less than one. Positron emission in this case would only create helium-5, which, because of the nearby presence of the stable helium-4, is just as unstable as lithium-5 is (half-life of Li-5 is only 3.70 × 10^(-22) s. Hard to believe such a short half-life can even be measured). Li-5 decays by proton emission almost instantly, forming stable He-4 and H-1. And the flip side, neutron emission, is also brought out by the unusual stability of He-4 in a similar case. Normally, radioactive isotopes with high neutron-proton ratios decay by beta decay. The artificially created and extremely unstable isotope helium-5 is an exception. Beta decay of He-5 would only convert it to its flip side, Li-5 (just discussed). So He-5, (half-life 7.00 × 10^(-22) s) decays by neutron emission to stable He-4 and a free neutron. There are other examples of proton and neutron emission, but they are quite similar to the cases I just mentioned (very rare and unstable artificially created isotopes). Spontaneous fission is a mode of decay occasionally seen at the opposite end of the periodic table, the heavy transuranic elements. Often it is only one of several modes of decay seen for these isotopes and it is usually a minority mode of decay in such cases. Spontaneous fission breaks a nucleus into two smaller nuclei (usually unequal) and a few free neutrons as well. Californium-252, one of the heaviest nuclei achievable by the neutron capture process in nuclear fission reactors, has a half-life of 2.645 years and generally decays by ordinary alpha decay. At least it decays by alpha decay 96.91% of the time. But the other 3.09% of the time Cf-252 decays by spontaneous fission, creating a variety of daughter isotopes. As a lead-in to my next topic, I will mention that creating rare and highly unstable isotopes is not really the most efficient means of observing proton emission, neutron emission, or fission. These processes can be induced much more efficiently with *nuclear reactions*, not radioactivity. And nuclear reactions will be my topic next time, whenever that may be. Reply #498. Aug 25 21, 5:54 PM |
| brm50diboll
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It definitely has been quite awhile since I've contributed to this thread. Not from lack of interest, mind you, but rather, from the strain of actively working on multiple Monthly badge upgrades which has been extraordinarily time consuming, especially as I've moved to the higher Tiers in some of those badges. It is tough to stay in the Top few of the monthly standings in multiple categories without playing a whole lot of games. Anyway, I had said I would discuss nuclear reactions and I do intend to do that. It's too bad I've avoided the discussion of medical science in this thread, however. I have done so because, despite all precautions, if I went there, it is likely I could be accused of violating the "do not give medical advice" admonition here on FunTrivia. Too bad. There is a great deal of material in medical science worthy of discussion and explanation. But, as an icebreaker from my long absence from this thread, here's a little clip: https://tinyurl.com/35b5uyjv Anyway, here's a piece of blatant medical advice for you: Do not do what I'm about to describe: An interesting example of a nuclear reaction that generates free neutrons is to mix beryllium metal with polonium-210. This combination creates free neutrons through the following reaction: Be-9 + Po-210 --- C-12 + n Beryllium-9 is the only natural isotope of beryllium and is stable though beryllium does have significant chemical toxicity. Polonium-210 is not the longest-lived isotope of polonium (Po-209 is), but it is present in nature, unlike Po-209. However, it is vanishingly rare. Po-210 has only a 138 day half-life and is a powerful alpha emitter (forming stable Pb-206). It was discovered by the great Marie Curie (who also discovered radium) and named by her for her country of birth - Poland. Because of its short half-life, it is extremely rare even in its natural source, uranium ore, and it was truly a magnificent accomplishment for Curie to isolate such a tiny amount of polonium from even a large amount of pitchblende. No wonder she won the Nobel Prize twice, once for Chemistry and once for Physics. Of course, years of sorting through radioactive materials did give her cancer, so there's that. Although Po-210 does exist naturally in vanishingly small amounts, it can be synthesized in significant amounts from ordinary bismuth by another nuclear reaction, but I don't wish to digress. Anyway, the Po-210 releases a high-energy alpha particle (He-4 nucleus), some of which strikes the Be-9 to produce C-12 and a free neutron. Why isn't C-13 (which is also stable) the product here? It has to do with conservation of energy and momentum concerns. The release of the free neutron is energetically favored here. This mixture of polonium and beryllium is a very good way to initiate free neutron production, which can then be greatly amplified by a subsequent nuclear fission reaction, but that's a topic for another time. Reply #499. Oct 19 21, 6:18 PM |
| brm50diboll
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The correct reaction should have been: Be-9 + Po-210 --- C-12 + n + Pb-206 Neglected to put in the Pb-206, which is the product of the alpha decay of Po-210. It is the alpha ray (He-4 nucleus) that actually strikes the Be-9 nucleus to produce the free neutron. Reply #500. Oct 20 21, 12:21 PM |
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