 I'm glad to see that everybody made it to class, I'm sorry about all the confusion but I'm glad that we're all here and we have a big class today, that's awesome. I guess there's no rain, but there's quits. I do want to make a couple of announcements. I had some of my online students and I really tried to work with them a lot but they asked me if we could push the L2 back one day. So it was due tonight, so September 10th at 11.55. I'm going to now make it due tomorrow night, okay, so that'd be September 11th, Saturday September 11th at 11.55, okay, so it'll be due tomorrow night. I added up there for quite some time but it was due on Monday, of course, we haven't done any of Chapter 3 yet, we're going to do a little bit today, but of course we won't have it all finished by Monday, so the test is going to be the test for the in-class people, the 14-06-01, all of you guys that are watching right now. That test is going to be held Friday, okay, the online test of course is being held Wednesday and Thursday, but what we'll do is make the HAL 3 homework due the same day as the test, okay, or actually day after the test. We'll make it due next Saturday and I believe that's Saturday the 18th, okay, yeah, so that'd be the 18th. So if there's any questions, let me know after class but the test will be next Friday and that'll be over Chapters 1, 2, and 3, okay, and the HAL homework that was due tonight is now due tomorrow night and the HAL homework that was due on Monday is now due next Saturday the 18th, okay. Well, you know, I mean, if you're five HAL sections ahead or whatever, you know, no big deal. Okay, so that being said, I'd like to do a 20, a small 20-minute lecture right now before I give the quiz, you know, we can start on some of Chapter 3 stuff so you guys can get some information so you can start doing Chapter 3 homework, okay. Over the weekend, that means that you have a long weekend. So anyways, Chapter 3, electronic structure and periodic law. So essentially we're going to be talking about the periodic table and why it's set up in the way that it's set up. Okay, so there's this guy, this is a painting of this guy. Obviously it's not a picture of him, he looks quite different, a little less austere in his pictures, but in that painting he looks quite grandiose or whatever, but this guy's name is Dmitri Mendeleev. He developed the modern periodic table, so this thing that you see over here. He introduced the concept that properties of the elements are periodic functions of their atomic weights, so what does that mean, periodic functions? That means that whenever you see an atomic or an atom that is actually placed in the periodic table, what you can see is that they're placed in like columns here and all of those columns within one of those columns. So if we look at these guys here, nitrogen, phosphorus, acetane and so on, we see that all of these have similar chemical properties. So in other words, it's periodic function, so every time you go to that column, they have similar chemical properties. So we call those columns families and we'll get into that in a little bit. So in reality, the properties of the elements are actually functions of their atomic numbers. So the atomic numbers, what you see here for hydrogen is 1, so the number that's in the top right corner there, 2 for oxygen is 8, sulfur is 16. These atomic numbers describe how many protons are actually located in the nucleus. This is an expanded version of the periodic table, taking the F block, which is down here, and inserting it in the place in the periodic table where it should be, which is right after 57 and 89. And again, respect to the chemical properties. So let's talk about the elements of the periodic table, the atomic number. So let's look at carbon. Carbon is atomic number 6. Remember, the atomic number describes the number of protons in the nucleus. The symbol for carbon is C, so right there. The name, of course, all elements have a name, all elements have a symbol, all elements have a number of protons in their nucleus. And the atomic mass of carbon is, actually this should say AMU after it, because the units for atomic mass are AMU. Sometimes you'll see U. I can't remember what our book does if it's U or AMU. Does it do AMU? Yeah, I prefer AMU. I really should have put some units after that. But anyways, the atomic mass here is found on the periodic table up here, usually below the symbol of the element. And that describes what is the average mass of any carbon atom that you'll find. So again, how I implied before these columns here are known as families or groups. And they have similar properties. This is showing a picture of the different halogens. In fact, this is chlorine here. This picture here is describing bromine, and this picture here is describing iodine. Notice that they're all diatomic elements. So if you can see the picture, each one of these has composed of two atoms stuck together. And when they increase in mass, we go from gas to partial liquid to more of a salt. But the properties of these elements are very, very similar to each other. In fact, their reactive nature is what you really want to think of as being the most similar. Here's another family of elements. Magnesium here, calcium and strontium. Notice they're the nature of the way that they're... And this metal C that we will eventually describe. But notice the kind of shape and chemical or physical properties of them being similar. Those are all the 2A elements there. Again, they have similar chemical properties. So some of the common family names that I'd like you to know are group 1A. This group here, this is the alkaline metals. Group 2A is the alkaline earth metals. I want you to know these because when I say alkaline metals, I want you to know what group I'm talking about. And then group 8 or group 18 over here, group 8A or group 18. We call these the noble gases. Their name implies, I guess, nobility. This is because of their place in the periodic table and their actual electronic structure. This electronic structure that they have actually gives them a very, very stable structure, atomic structure. So in fact that they don't react with any other elements. Which is why they have the moniker who will know. So that's why you call them noble because they're always turning their nose up at all the other elements. And then we've got a group of very reactive elements called the halogens here. That's the halogens. And again, they all have similar properties. If you've ever heard of halogens, headbites, or whatever, this is where that turns. So we describe the groups and families here. But there's another periodic function to this table here. And that's the periods themselves. So we call the horizontal line here a period. So if you see we've got period 1 contains these two elements. Hydrogen and helium. Period 2 contains these eight elements. Period 3 contains these eight elements. Period 4 contains these 18. 5, 18. And then 6 you have to include these at block 2. So it gets more and more and more as you go down. There's regions of this periodic table that aren't really described on this one that we have here. This big version of it. But if you look at the slide, you can kind of see there's a difference in color in the different regions of the periodic table. Notice this purple region covers most of the periodic table. So that's almost all the elements. Those are the melds. Anything that's purple in there is a meld. So this is an element that is relatively shiny and malleable. Malleable is if I take a hammer, I can pound it into a sheet. So that's what malleable means. And they conduct both heat and electricity. Some of the metals obviously conduct that stuff better than other metals. Like I was telling some of, or a couple of you this morning, I found out that my cap conducts electricity and my tooth pretty well. Well, they just got put on last week. And I was trying to eat with silverware. And it does me every time. So now I eat with plastic forks and mess. Okay. So those are all metals. They conduct electricity, conduct heat, all that good fun stuff. Notice the other two regions of the periodic table are very, very small compared to that big metal region. Okay. The next region I would like to describe is the non-metals. This is shown in green here. Those are elements that lack metallic properties. So kind of opposite properties of the metals. So where we're going to learn about covalent and ionic bonding and that sort of business, the metals all participate in ionic bonding. There's very few exceptions where you'll find them participating in covalent bonding. Every once in a while you will. And those are usually the smaller metals up here. But for a general rule of thumb, anything you're going to see in this class, ionic bonding for metals. Okay. Non-metals, on the other hand, can participate in both ionic and covalent bonding. And it all depends on what they're bonding with. If a non-metal bonds with a non-metal, then it will participate in covalent bonding. If a non-metal binds with a metal, it will participate in ionic bonding. And the difference between ionic and covalent bonding, ionic is the transfer of electrons from one element to another. And covalent is the sharing of electrons from one element to another. We'll really get into this, into the meat of this later next week, probably on Monday. Okay, so if you don't understand, it's no big deal. Okay. But anyways, let's look at a couple more regions. The yellow here, these are the metalloids. Okay, so those are elements that have properties in between a metal and a non-metal. And then the transition metals. So these metals over here, these metals in this region here, those are called the main group metals. Okay, these metals here in this little block here, or big block I guess, those are called the transition metals. Yeah. And they have inherently different ionic bonding capabilities and properties. Okay, so that's why we kind of separate our thinking about it. And again, we'll get into more of that way. Okay, so let's try to identify the group and period of some elements. So if we look at calcium, which is C-A, oh, and I'm never going to make you memorize the names or symbols of an element. I'll always give you the periodic tables that you guys see in front of you right now that have those names and symbols. So if you don't have one yet, there should be a pack of them coming around. And you can do it. I don't think you'll need it for quiz more. Okay, if you do, I'll tell you what it is. Okay, anyways, so let's identify the group and period of calcium. So calcium is C-A, so we've got to look for it on the periodic table. Okay, there it is. So we know its atomic number, remember its atomic number is shown in the top right corner there, that's 20. The atomic number describes what? What does it describe again? The number of protons and the nucleus, okay? And a neutral atom, so any of these atoms that you see on the periodic table here has a neutral atom. Those numbers also describe the amount of electrons that atom has, right? Okay, so here calcium's got 20 electrons, 20 protons. And notice it's in the group 2A, or we call these the alkaline earth metals. And notice it's in period 4, okay? So hopefully you guys can describe that. Cool. Okay, so let's start talking about the electron configuration. So I know we went over all of that protons, neutrons, and electrons in chapter 2. So let's talk about, if we can get on that slide. Okay, electron configuration. So electron configuration is the arrangement of these electrons in atomic orbitals. So remember we describe the electrons as being circling around the nucleus, okay? So this is a kind of, not really accurate picture of what an atom looks like, but that we can kind of start to describe it. And then we'll expand our knowledge of what it looks like later when we understand more about it. So the electron configuration is the arrangement of electrons in the atomic orbitals. So there's this guy called Bohr. His model talked about circling electrons around the nucleus in atomic orbitals. Okay, so you can think about it like this. We'll hear a higher energy orbital there, higher energy orbital there, and a higher energy orbital there, okay? The higher that the energy of the orbital would be, the further the orbital is away from the nucleus. Of course, the electrons are negatively charged. The nucleus is positively charged, so they really want to be closer together. So the further away they are, the higher in energy you can think they are. The more smack they could have if they were going to hit each other. So these electrons could be raised or lowered in energy level due to energy absorbed or released. So if I pumped a bunch of energy into this atom, I could make the electrons jump up and down into these different energy levels. Okay, so this model was revised in 1929. So, well, almost 100 years ago now. And 1926, sorry. And instead of circular orbitals, the location and energy of the electrons moving around the nucleus is now defined by using three specific terms. The shell, the sub-shell, and the atomic orbitals. Okay, so this whole thing would be a shell, okay? So the location of all of these different regions, okay, all of these different electrons here. So you see, you got the shell is the big thing. The sub-shell is this 4F or 4D or 4P or 4S. The orbital, let's see, the 4P has three orbitals in it. And an electron would be one of these dots, okay? So this is just a picture of building up electron configurations, okay? So let's just, we'll do a couple of these and just stop at this slide and get ready for the quiz. But let's build up the electron configurations of the first five elements, for example, okay? So the way I would really like to draw this picture is to put an arrow pointing up and that describes the change in energy, okay? So we're going to try to draw the picture of hydrogen. Hydrogen only has one electron, okay? And its electron is in the electron shell that is the first period up there. This shell is actually the S shell. And in fact, its sub-shell is going to be the 1S. There's only one sub-shell in the S level. So you can describe it as a box like that. So I guess we should say that the sub-shells are S, then it goes to P, which is greater than an energy than S. Then it goes to D, then it goes to F, okay? And if we look here, these are all the S sub-shells in these first two groups here. Okay, then it jumps over to here. These are the P sub-shells here. And then the D sub-shells is in the middle here, the transition metals, and then the F sub-shells. We'll really be only concentrating on the S and the P's in this class, okay? And again, you may think this is all crazy right now. You don't understand anything, but by the weekend this over, you'll figure it out, I promise you. Okay? So the way to build up these electrons is to show each one of these atoms gaining one more electron each time, right? So put one electron in there to describe hydrogen. So hydrogen, if we look at the periodic table, has one proton and one electron, so we only want to put one little line. An electron is usually described by what we call a fish hook arrow, a half-sided arrow, okay? So if we're going to describe the electron configuration of helium, helium has two electrons in it, so what we'll do is put another fish hook arrow like that, okay? When the electrons are in the same orbital, remember every one of these boxes is in orbital, okay? So the whole thing here is the shell. The 1S, that's the sub-shell, is the orbital, and these lines here describe the electron. So that's helium. And what you'll notice is that only two electrons can occupy an orbital at once, okay? Okay, so the first two atoms have electrons that occupy the same orbital, okay? Let's stop there. I know this probably, we're just starting it, so it might be a little confusing. We'll do a bunch of these on Monday, okay? And I promise you you will have no problem with it. Yeah, so the left side over there is S. S, P, E, S. Yeah, so what'll happen, look, what'll happen here? Think you'll fill up your electrons going one, two. It just goes by the periodic nature of these electrons. Okay.