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Chemical Science - Periodic Trends in Elemental Properties

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Principles of Chemical Science/n * Email this page/nVideo Lectures - Lecture 10/nTopics covered: /nPeriodic Trends in Elemental Properties/nInstructor: /nProf. Sylvia Ceyer/nTranscript - Lecture 10/nLast time we were talking about the energies of the various states in the multi-electron atom. Now, the question is experimentally how do we know what these energies are? How do we know these energy levels exist?/nAnd the primary technique that we have to know that is something called x-ray photoelectron spectroscopy. And x-ray photoelectron spectroscopy uses the same principle that we talked about when we talked about photo emission or the photo electric effect really explained by Einstein. And, by the way, I put on the website a very comprehensible article on Einstein's contribution to the photoelectric effect./nThis is something that just came out in a journal just very recently and I saw it. And so I urge you to take a look at it because it really is very understandable and kind of fills in some of the gaps that I left in talking about the actual problems, some of the kind of sociological effects or gaps./nAll right. So same idea, photoelectron spectroscopy and that we come in with a photon and that induced an electron to be emitted. But unlike the photoelectron spectroscopy on a solid where we essentially had just one energy level, in photoelectron spectroscopy on atoms we're going to have many energy levels as we've already seen./nAnd the photon that we're going to use, that photon is going to have an energy so that in principle it can pull out of the atom the most strongly bound electron. And if we want that, well, then that means that we're going to have to use a pretty energetic photon and that is we're going to use an x-ray photon./nSo it works in the following way. Let's take this unsuspecting neon atom. We bring in a photon, an x-ray photon. [LAUGHTER] I'll turn it down a little bit. Not too much. And that, of course, one of those electrons can be emitted. And one of the electrons that can come off is this p electron here so that the resulting ion configuration of the neon atom is 1s2 2s2 2p5./nBut also this photon can pluck off this s electron so that the resulting ion configuration is 1s2 2s1 2p6. And, of course, that photon can come in and pluck off an s electron so that the resulting ion configuration here is 1s2 2s2 2p6./nSo that one photon, in principle, has enough energy to pull off even the most strongly bond electron. And then what we do in this technique is that we measure the kinetic energies of these electrons that are emitted, and then we plot the number of electrons emitted at a particular energy. We scan the kinetic energy. And a typical plot looks like this./nThis is just the number of electrons coming off with the particular energy. And you see we see three different kinetic energies. We see electrons coming off with 383 eVs, electrons coming off with 1205 and electrons coming off with 1232 eV. Well, of course it's going to turn out here that those electrons that come off with the lowest kinetic energy are those electrons that were most strongly bound./nThere are those electrons that were in that 1s state. Why? Well, look at the energy level diagram here. Here are those 1s electrons. Here is the incident energy of the photon that we came in with. If these electrons are most strongly bond then they're going to have the highest ionization energy, the highest energy to pluck them off./nAnd, therefore, in order to conserve energy, the incident energy, the ionization energy plus that kinetic energy has to equal that incident energy. And if this is the largest energy, well, then those are the electrons that are going to have the lowest kinetic energy. So those are those 1s electrons. That's how we identify those. And then there are electrons coming off at 1205 where those are the electrons from the 2s state./nThat's a little bit higher energy. And so if this is the energy of the incident photon, it's a photon coming in as the same energy. If that's the energy of the incident photon this is then the ionization energy for those electrons in that 2s state. Well, that leftover energy then, that's the kinetic energy. And it is a little bit greater than it was for 1s because, again, those electrons are less strongly bound./nAnd then finally we have this feature here, 1232 eV. Those are the 2p electrons. These electrons are least strongly bound. So here is the incident energy of the photon. Here is the ionization energy of those electrons in the 2p state. The resulting or the leftover energy here is the energy that goes into the kinetic energy of those electrons, and that is going to be the largest kinetic energy./nBecause those electrons are least strongly bound. So in this neon atom, we've got three different states with three different energies. That is reflected in the fact that we're going to have, in the photoelectron spectrum we're going to have three different kinetic energies for the electrons coming off. So by looking at the photoelectron spectrum you can immediately tell how many energy levels you've got./nAnd, of course, as the kinetic energy goes up, well, that means that the binding energy goes up, the binding energy gets more positive meaning the electrons are less strongly bound. Or, another way to put it, as this kinetic energy goes up this way the binding energy gets more negative this way. The electrons are more strongly bound as that binding energy gets more negative./nBut from this kind of experiment here we can extract what these ionization energies are and we can extract then what the binding energies are. Because of conservation of energy, incident energy of the photon equals the ionization energy plus the kinetic energy. I can rearrange this to calculate the ionization energy. I did that for all of these three features./nHere is the incident energy of that x-ray photon. Here is the measured kinetic energy. The difference between the two is the ionization energy. I solve for that there. And then that ionization energy we know is minus the binding energy. So if the ionization energy is 21.6 eV, the binding energy here for that 2p electron is minus 21.6 eV and for 2s minus 48.4 eV and 1s minus 870.2 eV./nI want you to notice how much more strongly bound that 1s electron is. Then the electrons in the n=2 shell, look at that. That 1s electron is bound by 870 eV binding energies minus 870. Binding energy for the n=2 electrons is minus 48 and minus 21./nWhy is that more strongly bound? Well, that electron is in the n=1 shell, which is much closer to the nucleus where the attractive interaction between that positive charge on the nucleus and the electron is much greater. So this is basically the technique. This is how we calculate those binding energies, or observe those binding energies for the electrons in the multi-electron atom./nWhat I want to do now is I want to start on today's lecture. And you don't have the lecture notes for today's lecture. And there is a reason for that. Oh, and I forgot to tell you to do something very important. I'm going to start on today's lecture. And through this lecture I am going to be asking you some questions./nAnd you're going to have the opportunity to answer, but the way you're going to have the opportunity to answer is that you're going to have to make some paper airplanes. I will ask a question and then I will say launch. And what you have to do is you have to throw the paper airplane to the board. And the first paper airplane that hits the board from that section, well, that section gets a chance to answer that question. And we're going to keep score. And we've got some treats for the winning section. The TAs have some scrap paper and they're going to pass that out. And so you can be making some paper airplanes. Write the number of your section on the paper airplane./nAlso, some of you are not strategically located to his the blackboard. If that is the case, why don't you select somebody from your recitation to be the launcher. They can come down as close as these few stairs and aim at it for your section. OK?/nOK. So I am going to start talking. We're going to talk about the Periodic Table and the trends in the properties along the Periodic Table./nAnd, of course, it's the electron configurations that allow us to understand the structure of the Periodic Table. But it is important to realize that the Periodic Table was first assembled by Mendeleev and other people, not on the basis of knowledge of the electron configurations but on the basis of knowledge of the similar properties./nSimilar chemical and physical properties of the elements. So, for example, lithium, sodium, potassium here, they were all grouped in one column because they were all found to be very soft malleable metals that were very reactive. Helium, neon, argon, they were all grouped in one column because they were found to be very inert, very unreactive./nAnd, of course, now we understand the electron configurations, why these elements have the similar properties. In the case of the alkali metals, lithium, sodium, potassium, we've got that extra x electron, that unpaired electron that makes those metals very reactive. And helium, neon, argon, we've got an inner gas shell configuration that makes those elements very unreactive./nBut by the late 1800s and the early 1900s the similar chemical properties of atoms in a column of a Periodic Table, that really was strongly believed, strongly believed to an extreme. So, for example, it was known that human beings consumed sodium, salt./nIt was known that human beings, our bodies contained sodium ions and potassium ions. Well, if our body has sodium in it and potassium in it and we consume those two elements, why not a little bit of lithium? OK. About 1925 or so a soft drink came on the market./nAnd that soft drink, they wanted to have a lemony flavor to that soft drink. And so how do you get a lemony flavor. Well, you use citric acid. But citric acid isn't really soluble in water so, hey, we've got to make a salt of the citric acid. For whatever reason, a lithium citrate was made./nA lithium salt of citric acid was made. And that soft drink was none other than 7UP. And, in fact, if you look here it tells you that the lithium was touted to give healthful benefits, including an abundance of energy, enthusiasm, a clear complexion, lustrous eyes or lustrous hair and lustrous eyes./nShinning eyes. [LAUGHTER] And so this was on the market until the early 1950s when finally the anti-psychotic properties of lithium were being discovered and the side effects of lithium were being discovered. And the 7UP was then taken off the market, or at least the lithium citrate in the 7UP was removed from the market./nAnd, of course, it wasn't until the early 1970s that lithium was approved as an anti-psychotic drug, because it really is poisonous to the human body in large amounts. And so I really don't know how many individuals had side effects from drinking 7UP at that time./nBut the moral of the story is that even though we're about to talk about some of the general trends in the Periodic Table, don't put any element in your mouth just because it's in the same column as some element that you do put in your mouth. [LAUGHTER] True story. All right. One other property we're going to talk about is the ionization energy. And we've talked about it already but we haven't formally defined it. Let me do that here./nHere you see boron with its electronic structure going to boron plus. The energy change for this reaction is what we call the ionization energy. It's the product minus the reactants. That's the ionization energy. Equivalently, that ionization energy is minus the binding energy of the 2p electron in boron./nThis is a process that is removing the least strongly bound electron. When we remove the least strongly bound electron, the energy that is required to do that is called the ionization energy. And, more fully, it's the first ionization energy. It is the energy to remove the electron from this HOAO, the highest occupied atomic orbital./nOr, the least strongly bound electron. Same idea. This is highest occupied atomic orbital. Now, generally, when we say ionization energy we mean the first ionization energy. We often leave out the word Ã¬firstÃ®. So whenever you look up an ionization energy, implied in that is that it is the first ionization energy./nThe energy to remove the least strongly bound electron. The second ionization energy, well, that is the energy required to remove another electron. But now we're removing it from boron plus to make boron plus two. And, again, we're removing the next least strongly bound electron. That's the second ionization energy. Equivalently, it is minus the binding energy of the 2s electron in boron plus./nNot boron. And then there is a third ionization energy. Again, now we're taking boron plus two, pulling off the next least strongly bound electron to make boron plus three. That's the third ionization energy. And we can keep going. Boron plus three. Pull out the next electron. Boron plus four, that's the forth ionization energy./nAnd then the fifth ionization energy, taking the last electron off of boron, boron plus four to make boron plus five. That's the fifth ionization energy. So these are our definitions here of the ionization energy. Now, I'm going to switch to the side screen./nAll right. But now here comes the first question. I'm going to show you two reactions here. The first reaction is a boron plus going to boron plus two. We are removing here an s electron./nThe second reaction here is boron going to boron plus. Again, we are removing a 2s electron. And the question is, are these two energies equal. In both cases we're pulling of an s electron a 2s electron. So that's the question. Now, the way we do this is I'm going to move to the side, I'm going to say launch and then Christine and I will try to determine the first airplane to hit the board./nWe will pick it up, and we will ask that recitation section to answer. All right. Ready? Launch. Recitation one, are those energies equal? Answer? It can be a group effort./nHey, you've got a fifty-percent change here. Yes or no? No. They're right. They get a point. OK. You've got to keep score on the board./nAll right. Now why aren't they equal? Hey, big no here. They are not equal because, in this case, this is the ionization energy, the second ionization energy, as we just defined it. Boron plus going to boron plus two, that's a second ionization energy. It's minus the binding energy of the 2s electron in boron plus./nHere, this energy, what this energy is is the ionization energy of the 2s electron in boron, not boron plus. This is equal to minus the binding energy of the 2s electron in boron. This is minus the binding energy of the 2s electron in boron plus. They're not equal. The next question is which one is greater? And so wait a minute. I'm going to get out of the way here./nOK. Launch./nRecitation nine. Over there, sorry. Which one is greater, top one or bottom one? Top one. They are right. Hey, recitation nine gets a point./nYou can put the score on the board. Great. The top one is greater. The top one is greater because the effective charge of this 2s electron in boron plus is larger than the effective charge of the 2s electron in boron. The reason for that is because in boron plus we have fewer electrons. We've got less shielding of the nucleus by the electrons because we have fewer electrons./nSo that's great. But this is really very important for you to understand here. This will be on the lecture notes that I will give to you at the end of the class. Because we might run out of paper, you might want to save those. Now we're going to look at the Periodic Table here./nAnd I'm going to ask this following question here. As we go across the Periodic Table, what we want to know is how does the ionization energy change. So here is my Periodic Table. I'm going across this way. Does the ionization energy increase or does it decrease?/nNow, here is another chance for you to answer this for your recitation section. Launch. Wads don't count. [LAUGHTER] Paper airplanes. Eight. Where is eight? Back there. Hey, OK, eight./nWhat happens to the ionization energy as you go across the Periodic Table? What? Increase. Is that your final answer? You're right. All right. The ionization energy increases. Why does that ionization energy increase? Here is our example. Here is our plot./nIonization energy versus z as you go across. Here is from hydrogen to helium, that's one row, it goes up. Here's from lithium to neon, another row, it goes up. Well, the ionization energy increases because the nuclear charge increases. As you go across the Periodic Table, the nuclear charge increases. The Coulomb interaction increases./nBut there is another part to the Coulomb interaction. And, of course, that other part is the distance. But as you go across the Periodic Table, you're putting electrons into the same shell, to the same principle quantum number shell. And, in general, when you put electrons into the same principle quantum number shell, you are, on the average, the same distance away from the nucleus. And so r is the same./nZ is increasing as you go across. And so, therefore, that ionization energy increase as you go across the Periodic Table. Now, however, you can see that there are some glitches. There is a glitch here between beryllium and boron. Boron is actually a little bit lower in ionization energy that a beryllium. Can we understand that?/nWell, we sure can in terms of the electronic structure. So here is the electron configuration of beryllium, z=4. And then here is the electron configuration of boron, this is 1s, this is 2p, z=5. The bottom line is that it takes more energy to put the next electron here into this 2p state./nAnd the bottom line is that the charge on the nucleus does not increase sufficiently to fully compensate for the energy that is necessary to put that next electron in that higher energy state. So this just doesn't get large enough fast enough to compensate just for the extra energy that is needed to put that electron in the 2p state./nSo we have a little glitch between beryllium and boron. In addition, we've got a little glitch right here between nitrogen and oxygen. The oxygen ionization energy is not as large as you might expect from the increase in the z quantum number. Again, here's the nitrogen./nZ=7, the electron configuration for nitrogen. Here it is for oxygen./nAnd what we had to do for oxygen is that we had to put another electron into a 2p state, a 2p state that already had an electron in it. And there is a great deal of repulsion here between these two electrons in the same state./nAnd, again, the z here, the nuclear charge did not increase large enough or greatly enough to compensate for that electron repulsion. And so the ionization energy is just a little bit lower than it is for nitrogen. That same trend, of course, continues for the other rows in the Periodic Table. Here is the third row./nAnd there is also a glitch between magnesium and aluminum. Magnesium and aluminum are right below beryllium and boron. And a glitch between phosphorous and sulfur, they're right below nitrogen and oxygen. Next question. As we go down the Periodic Table what happens to the ionization energy?/nWait a minute. I'm getting out of the way. OK. Ready? Go. Recitation one. What happens to the ionization energy as you go down the Periodic Table? Decrease. Yes, they're right. Why does it decrease?/nIt decreases, as you go down the Periodic Table, z is increasing, that is true, but as you go down the Periodic Table you're putting electrons into shells that, on the average, are farther out from the nucleus. And so as you go down the Periodic Table it's that distance dependence that dominates the ionization energy./nSo across the Periodic Table ionization energy increases, down it decreases, across its z dependence that dominates because r is constant. As you go down, it is the r dependence that dominates. That is changing. Even though z is changing, r changes faster than z. Great. Now let's talk about another property here. And that other property is called the electron affinity./nIf I take a chlorine atom and I add an electron to it to make Cl minus, the energy change for that reaction is minus 349 kilojoules per mole. That is the chlorine ion is more stable than the chlorine atom./nThis is a reaction that releases some energy. Now, the question is there is a quantity called electron affinity. You need to define the electron affinity. Are you ready to launch? All right. Launch./nFour. Recitation four. You're up there, right? Can you define the electron affinity? What's the electron affinity equal to? I am going to write this. Electron affinity is equal to what? Yes? Answer?/nYes, that's great. But I want a mathematical definition here. EA equals? Negative ie? No. All right. You've got to do another launch. Ready? Launch./nRecitation one. What is the electron affinity equal to? Answer?/nBeep. What it is? No. Next launch. Ready?/nLaunch./nWhat is it? Nine. What is the answer? What is the electron affinity equal to?/nBeep. What?/nYes. Keep going. Not ionization energy. OK. The electron affinity here is minus this energy, right? If you add an electron to an atom there is an energy change. The electron affinity is defined as minus that energy change./nOK. Great game. We didn't get very far. However, from these scores, recitation section one does win. OK. I will see you on Monday. (less)

Channels: Lectures Chemistry (General)
Tags: Periodic Trends Elemental Properties

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