 If we get done with some minutes to spare, then I'll let you go on and take a little extra time to study for the test, okay? So, that means that let's learn the rest of the stuff that we need for the test on Friday. So, the rest of this, the lecture of electronic structure and periodic law. And really this is, you've essentially learned everything you need to know about the periodic table. Now we're just going to learn about which types of ions combine to form compounds and how to name those compounds and how to draw their molecular formulas, okay? So, if you recall the last thing that we talked about last time was the octet rule. Elements usually react in such a way as to attain the electron configuration of the noble gas closest to them in the periodic table. So, recall we were talking about elements on the right side of the table move right to the next noble gas. So, on the right side of the table we're talking about these atoms here, right? Does anybody remember what block we call this? P-block. So, most of the P-block atoms will move to the right to take on the configuration of the noble gas closest to them on the right. So, if we look at nitrogen, oxygen, chlorine, they'll all go to neon noble gas configuration. The one element that you might confuse is that's a P-block element that will actually go to the left is aluminum, okay? So, you'll see that a lot. In fact, any of these elements here in row 3A will go to the left. But most of the P-block elements will go to the right to attain that noble gas configuration. And then elements on the left side, so like the S-block here and the first row of the P-block will move, if you will, backwards to the noble gas of the previous row. So, if we're looking at sodium, for example, or magnesium or aluminum, it's going to go backwards to attain the neon configuration by losing three, two, or one electrons, respectively, okay? So, again, the atoms are reactive in relation to the noble gases, and the noble gases are unreactive due to the filled valence shell of electrons that they have on them. So, what's happening is that the other atoms are trying to attain this non-reactive or very stable, because reactivity and stability are opposites of each other, okay? They're trying to attain this stable configuration, and in fact, that's what they do. So, when an atom goes back or forth to its noble gas configuration, it becomes an ion, okay? So, an ion is an atom or group of atoms that have a positive or negative charge. So, this is a charge species, okay? Ionic compounds are the first type of compounds we're really going to talk about in detail. And they form from a transfer of electrons from one element to another. They typically form when a metal reacts with a non-metal. If you recall, the metals are, these are the main group metals in the S block, the transition metals in the D block, and then there's some metals along the staircase on the bottom of the staircase in the P block, okay? So, those are all the metals from here all the way over, and in fact, the S block is all the metals. So, most of the periodic table, in other words, is metals, okay? There's only, in fact, a few non-metals, and most of them are either listed in red or blue here on the periodic table. Of course, there's a couple that aren't, iodine, carbon, sulfur, pardon me, are good examples of that. Okay, so, in general, if the metal becomes a positively charged ion, it's called a cation. So, becoming positively charged implies that you're losing electrons, okay? Let's look at lithium, for example, and I'm going to draw something that's called a Lewis structure of lithium. We don't have to know about Lewis structures for this exam, okay? But it really does help thinking about it. In fact, it will hit Lewis structures pretty rigorously in next chapter, okay? And for those of you who have looked at the practice exam and saw question number five, that's essentially all Lewis structures, okay, reactions of Lewis structures. But anyways, what is a Lewis structure? A Lewis structure just indicates the amount of valence electrons that the atom actually has, okay? So, recall the octet rule, the atom wants to have a valence electron. If I put one dot there, that means it has one electron, okay? If I were to draw the Lewis structure, for example, of neon, remember, it has a full octet, so I would draw it like this. It's got eight electrons around it, each one of those dots is representing one electron, okay? Lithium, so you can think of atoms attaining a full octet if they increase their electrons, their valence electrons to eight, or if they decrease them to zero. Either way, it gives it a full octet, okay? So, if we look at lithium here, you realize it's only got one valence electron, so you've got to ask yourself a question. Is it easier to give this atom seven electrons or is it easier just to take one electron away, okay? And it's always going to be easier to do the least amount of electron transfer, okay? So, if I'm comparing seven to one, those two numbers, of course one is much smaller than seven, so that's the way it's going to go, okay? So that's why elements on the left side of the periodic table tend to move backwards to the noble gas configuration that's on the period before, and that's why elements that are on the right side of the periodic table tend to move forward to the next noble gas configuration. So what you can effectively say is that lithium is in a high energy or high reactive state in its elemental form, okay? It doesn't like to have that extra electron on it, so what it wants to do is just get rid of that electron, okay? If we look at fluorine, for example, fluorine, if we look up there, it's got seven valence electrons, okay? So, okay, let's just... So, of course, fluorine wants to gain one electron instead of losing seven electrons because gaining one is much smaller amount of electron transfer than losing seven is, and what these two elements actually do is provide the electron transfer for each other, okay? So this guy wants to remove that electron, this guy wants to take it, so what you'll do is just take that and move it over there, okay? So, in fact, this is the formation of two ions, okay? So these two ions, of course, now will both have the noble gas configuration. Remember, you can have the noble gas configuration with zero electrons around you or with eight or at the electrons around you. Lithium, of course, losing that negative charge now becomes positive, right? See that? If you lose a negative, you become a positive. So this is a cation count. Of course, when fluorine gains that electron, it's gaining a negative charge, so the whole ion now becomes negatively charged and has a full octave. Notice, by transferring this electron, both of these atoms now are ions and both of them have a full octet. So both of them are much more stable than they were in their, you know, elemental form, if you will. Okay? So, of course, a negatively charged thing is called a hammer. We'll learn a lot more about this again in Chapter 4, how these electrons transfer, but I wanted to give you some sort of clue of kind of a picture mechanism of transfer, and then we can really talk about it. So anyways, the bond that results from these electrostatic attraction between oppositely charged ions, right? So what happens now is, of course, anybody who's dealt with magnets or anything like that can see that this is positive, this is negative. Remember, positive and negative charges, they want to get together to cancel those charges out. So what they're going to do is kind of just stick together and form what we call an electrostatic interaction. We're going to form this compound there, lithium fluoride. We're going to form this electrostatic interaction. This electrostatic interaction between the positive and negative charges, we call a bond. This particular bond is called an ionic bond. So this is an ionic compound. And this name or this molecular formula, actually, it's not really a molecular formula. It's called the unit formula because it doesn't really fully describe what's going on here. And I'll show you why, because it's not just a lithium ion, one lithium ion and one fluorine ion, stuck together. In fact, it's a harmless array of lithium ions and fluorine ions stuck together. But the smallest unit, the smallest repeating unit, is the one lithium and one fluorine stuck together. So we can pass this around. I don't know, I'd rather come or so. You can come up and look at it after class if you really want. That's the smallest what we call unit formula. So we just write it like this. Of course, we're not going to write lithium 500,000, fluorine 500,000. It's a 1 to 1 ratio, so we're just going to name it that. It's going to be different than a different type of bond or a different type of compound. So I just want to emphasize the difference before we talk about who they want bonds. So ionic compounds are three dimensional arrays of alternating positive and negative ions. So you see that again here. Three dimensions, alternating positive and negative ions. Okay, cool. So again, ions are electrically charged particles that result from the gain or loss of one or more electrons by the parent atom. You see lithium here, it loses its electron, becomes the lithium ion. You see fluorine here, it gains that electron, becomes the fluoride ion. The thing is with anions, you change the end of the name from E to I when they become anion. Notice here there's a little reaction equation that describes that. The one thing I want you to notice here is when lithium loses its electron, it gets small. That makes sense, right? It's like throwing out its valence shell, so it's getting smaller. Fluorine, on the other hand, is gaining electron and you can see how much bigger it gets. So that's something you're going to want to know. So again, this is another way to look at it. We can look at it through electron configurations. So we see the condensed electron or abbreviated electron configuration for sodium here, indicating that we've got the neon noble gas configuration plus one electron in the 3S sub-ship. Of course, it's easy enough to lose that 1S electron and gain the neon configuration. So that's what sodium is going to do. Sodium, having only one valence electron, when it loses it, it's going to become what? A cation or an anion? Cation. Cation, right? Because it's going to be positively charged, it's losing its electron. And what's going to be the charge of the sodium atom or ion now? Positive one. Positive one, right? Positive one. Why is it positive one? Because it's only losing one electron, okay? So ion formation in the octet rule, all atoms of a group lose the same number of electrons, okay? So all group one elements lose one electron. All group two elements lose two electrons. All group three A elements lose three electrons. All group seven A elements gain one electron. All group six A elements gain two electrons. All group five A's gain three electrons, okay? So it's this regular pattern going from negative three to positive three, wrapping around the periodic table, okay? So you can see here, starting at negative three, we go negative three, negative two, negative one, zero, one, two, three, okay? So just remember those seven numbers in sequence, and you should be able to predict any ion, okay? And again, the resulting ion has the filled octet, okay? Here you see aluminum losing its three valence electrons, becoming the AL three ion. Okay, so we talked about the main group elements, the elements in the S block and the P block. Now we're going to talk about the transition metal ions, okay? And how to, well, know what their ionic charge is, okay? So the one thing I want you to remember is that all metals, all metals are positively charged when they become ions, okay? So that applies to the main group metals, that applies to the metals in the P block, and that also applies to all of the transition metals, okay? So the transition metals are all the metals that you find in the D block here, okay? So if you look here, main group metals, main group metals, and transition metals, okay? They all become positively charged. The metalloids, don't worry about them for right now, okay? Just worry about the metals and non-metals. All metals are positive, yeah? Just remember, they're going to the left, you know? Okay, so the transition metals, unlike the main group metals, don't have a regular pattern of predicting which ions they're going to form, okay? In fact, some transition metals form more than one ion. In fact, a lot of them do. A common example is iron, they can form iron plus 2 and iron plus 3. So this element here, so it can form different ions of different charges, okay? So it's the same atom, but it's got different charges on it. Okay, well I guess we should have hit this one right before we did that transition metal stuff. But let's try to do this. What's the charge of the most probable ion of calcium? Calcium's a 2 plus, if you want to write, 2 plus, okay? So let's write it out on the board, just like this. We would say calcium 2 plus. Okay, that's how you would write it. Okay, what about strontium, SR? 2 plus. 2 plus, right? Does everybody get that? Why do you say 2 plus? Because it's in group 2, right? It's the second family, okay? What about sulfur? 2 minus, 2 minus, 2 minus, right? Sulfur you would write like this. From sulfur to sulfur 2 minus. If we were to write the Lewis dot structure, sulfur would look like this. It would want those two electrons. When it gains those two electrons, that means the two electrons are coming in, obtaining the octet field, okay? So again, this is what we're going to do. Pretty hardcore next chapter is predicting how these things work. And phosphorus, what about that one? What's the ion that? 3 minus. 3 minus. Okay, cool. Good job guys. 3 minus. Why is it 3 minus? Because it's in the third from the new noble gas, right? Okay, so which of the following pairs of atoms and ions are isoelectronic? This word isoelectronic means I have the same electron configuration as you, okay? So does argon and chlorine have the same electron configuration? Argon and chlorine minus, I guess I should say. Argon and chlorine minus? So some people are saying yes. Everybody else is just saying anything. Okay, so chlorine's here, right? If it gains one electron, what configuration does it have? Argon. Okay, so it has how many electrons essentially? 18 electrons, right? So it's got its octet field. So it gains one electron to become the argon configuration. So those two are isoelectronic. What about sodium plus? The sodium is here and neon is here. Do sodium plus and neon have the same electron configuration? Yes. Yes, yes. What about magnesium 2 plus and sodium plus? Yes, they do, right? They also have the same electron configuration as what? Neon. Neon, right? Neon. Okay, and what about oxygen 2 minus and fluorine minus? Neon. Yes, and what is the same electron configuration as that? Neon. Neon as well, or sodium plus, or magnesium 2 plus, or aluminum 3 plus. They all have the same electron configuration. How come they all have the same electron configuration? They all want to be the noble gas configuration. So that's what they're doing, is trying to get it. Okay, so of course they're going to have the same electron configuration. Okay, so you could try to do this one now. How many protons, neutrons, and electrons are in each of the following ions? We'll just go rather quickly through this one. Of course magnesium, we'll just do this one. Magnesium 24 2 plus. So how many protons does this atom have? 12. 12. How do we figure out that? Because it's got the atomic number there. If it didn't have that, you could just look up here and look at the atomic number there. I would tell you 12. How many electrons, or how many neutrons does this thing have? 12 as well. 12 as well. How do we figure that out? We look at the mass number, the number of protons, or the atomic number away, and we get the number of neutrons. How many electrons does this thing have? 10. 10 electrons, right? How do we know that? Because two of its electrons have gone away, right? So what's the noble gas configuration of this ion? Neon. Neon. Again. Okay. Neon again. Okay. So good job guys. Okay, so let's talk about some trends in the periodic table. Atomic size is a trend that we want to know. Many atoms, atomic properties correlate with the electronic structure and also with their position in the periodic table. Almost all of the reactivity does. And of course that's what this class is concerned about. So that's what we're going to talk about. Two of the things I want to talk about specifically are atomic size and ion size. Okay. So the size of an element increases moving down through a group. So in that case it means hydrogen is smaller than lithium, which is smaller than sodium, which is smaller than potassium, which is smaller than rubidium, which is smaller than cesium, which is smaller than francium. Okay. So francium is actually the biggest element of group one. Okay. So if we look at the noble gases here, which one's going to be the biggest? Helium or radon? Radon. Radon, right? Why? Because it's at the bottom of the thing. Because we're going to be the smallest noble gas. Which one's going to be the smallest? Neon. Neon. And the second smallest? Or helium. Sorry. The second smallest? Neon. Neon, right? And the third smallest? Argon. Argon. Okay. Sorry. Yeah. Neon's the second smallest. Yeah. Nine o'clock in the morning for everyone. Okay. So we understand going down the periodic table that things get bigger. Okay. What you might think is going across to the right things will get bigger too. But that's not the case. Okay. As a matter of fact, if you go from left to right, things get smaller. Very, very good. It's smaller. So the size of an element decreases from left to right across a period. Or if you want to think about it conversely, if you go this way to this way, you get bigger. Okay. So which one is the biggest element of period four? Potassium here or krypton here? Potassium. Potassium.rance-maker than krypton. Okay. So what's the biggest atom on the periodic table there? What's its atomic number? How about that? The biggest atom. The biggest atom. What's its atomic number? So where are the big ones? Are they on the left or on the right? So they're on the left, right? And the top or on the bottom? On the end. So what's the atomic number of the biggest atom? Only here like 30 people saying, and they're all saying different numbers, 87 is the biggest one. Because why is that? Because it's the furthest to the left and the furthest down, right? So it's got to be the biggest one, right? What's the smallest atom on the periodic table? Helium. Helium, because hydrogen is bigger than helium, right? Because things go smaller to the left. So helium's the biggest, uranium's the smallest, uranium's the biggest. And in fact, you can see that here. In fact, they cut off hydrogen and helium. But look how small flooring is relative to oxygen, to nitrogen, to carbon, boron, beryllium, and lithiums, enormous compared to it, OK? The reason is because the more protons you get, all of those electrons are at about the same radius from the nucleus because they're all in the same shell, OK? So the more protons you get, the more positive pull you get, and that kind of decreases the orbitals, OK? So they decrease the shell size. So the more positive nature you get to your nucleus, the kind of more it contracts that electron cloud is the way we talk about it, OK? So if we look here, it doesn't count the period 7 either, but cesium is much bigger than any other atom that's listed here. Cat ion size, an anion size, OK? We've already talked about this briefly when we looked at lithium going to lithium plus. Remember, it got smaller because it lost that electron. What happens with atoms, and in particular, you can think of the transition metals, in particular the ones that actually have multiple oxidation states or multiple charges for the cat ion, such as iron 2 plus and iron 3 plus. Remember, the more electrons you lose, the smaller you get. So we're talking about iron 2 plus and iron 3 plus. Which one has lost more electrons? Iron 3 plus, right? So if I've lost more electrons, I'm smaller, right? So which one of these two is the smaller of the two? Iron 3 plus. Why is that again? Because it lost more electrons than iron 2 plus, right? Does everybody get that? So the more electrons you lose, the smaller you get. So cat ions, of course, are smaller than the parent atom. And here's another picture that we showed earlier. Anions are a lot bigger than their parent atom. Anions have more electrons than their parent atom. I don't know, I shouldn't say. Oh yeah, then protons as well. They have more electrons than they have protons. And the excess negative charge reduces the hole on each of the individual electrons. So it gets a little bigger, OK? And you can see the relative size here of each ion to its parent atom. And you can see quite a discrepancy in the size. Sodium to sodium plus. Sodium plus is much, much smaller. We look at bromine to bromine minus. Look how much bigger it gets. This one, my iodine, almost are chlorine doubles its size. They are almost chlorine too. So they get pretty big, oxygen and relatively small. So remember these terms as well. So one thing we're just going to touch on right now is the electronegativity. Electronegativity is the amount of, it's kind of a nebulous kind of quantity. But it's very important that you understand what its consequences are mostly. But electronegativity means how closely do I hold my electrons to my nucleus? And how much do I want more electrons from other atoms? So it's kind of like the electron hoarding statistic. If you look at electronegativities, what you find is that the smaller elements are more electronegative than the bigger ones. So what you find is electronegativity increases oppositely of atomic size. So the least electronegative atoms are found in the bottom left corner. The most electronegative atoms are found in the top right corner. In fact, the noble gases, we don't really talk about their electronegativity because they're not reactive. So what we normally say is that fluorine is the most electronegative element. So fluorine looks like somebody that hoards things or hoards electrons. It wants to hoard as many electrons as it can. So I briefly mentioned some types of formulas. Let's go over the types of formulas that we can talk about. So chemical formulas show the type and number of each atom present in the smallest unit. So remember, we said lithium fluoride here looks like this. So it's a very big three-dimensional array of ions. It's not just one lithium atom and one fluorine atom. But the chemical formula shows the smallest unit of that compound. The empirical formula will show the smallest whole number ratio of elements in a compound. So the empirical formula and the unit formula are equivalent for ionic compounds. They won't be when you start talking about covalent compounds. Molecular formulas and structural formulas will come back to you because molecular formulas, again, talk about covalent compounds. And structural formulas also, see if I have to show you, friends between, which is something like this, kind of covalent compound that's just a small thing. It's just one thing. It's only one C, one, two, three, four, five, six, seven, and H, one, two, three, four, five, six, seven, eight, nine, 10, 11. So this is C7H11. That's the smallest unit of this thing. So you can see that this is not a three-dimensional array of ions like this. So you were comparing these two types of compounds. Anyways, we'll get back to that. Let's talk about nomenclature. So nomenclature is the assignment of a correct and unambiguous name to each and every chemical compound. So if we don't specify specifically the name, then people will be misconstruing what you're talking about. So we need to really kind of concentrate on the types of names that we're saying and what we're saying and what we're writing. So that's why it's important to learn the names of the different compounds. So there's two different types of naming systems. One's for ionic compounds, like these three-dimensional arrays that we've been talking about. And one's for covalent compounds like this cycloheptene molecule. So ionic compounds, metals and non-metals react to form ionic compounds. So we've got all the metals here. They react with the non-metals and form ionic compounds. So this lithium fluoride, everyone lithium, is a metal or a non-metal? Metal or a non-metal? Metal, right? Fluorine is a metal or a non-metal? Non-metal. So when they react together, they're going to form a what type of compound? Ionic compound, OK? Ionic compound. The metals are always going to be the cation. Non-metals always going to be the anion. The cations and anions arrange themselves in a regular three-dimensional repeating array called the crystal lattice. So something like this we call the crystal lattice. The formula of an ionic compound is the smallest whole number ratio of ions in that substance. So this effectively is the empirical formula. So if you recall, common oxidation numbers, I think this really gives us a good visualization of that wrapping around that I was talking about with the periodic table. So you can go from negative 3, negative 2, negative 1, 0, positive 1, positive 2, positive 3. You can see it visually right there. And again, here is common oxidation numbers all selection of all different elements in the periodic table. Again, you don't have to memorize the oxidation or the ionic charges of the transition metals because, of course, a lot of them have varying charges. And we're going to learn how to figure out what the actual charge of the transition metal ion that you're actually working with is. So main group elements, the S block and the P block. So we call those main group elements, I guess I should say. The main group elements, like we said, their ions have regular oxidation states. Oxidation states, I guess I've been using these terms interchangeably, oxidation states and the charge of the ion, those are equivalent terms. So if you hear me say oxidation state, I'm just talking about what charge is that ion at. So the oxidation state for these group 1s would be positive 1. It's just the same as the ionic charge. So they have a regular oxidation state that's a function of their group number. So group 1s have plus 1, group 2s have plus 2, group 3s have plus 3, group 7a's have minus 1, so on and so on. The total ions in each unit must be balanced when writing the formula unit. So recall here, we've got a lithium plus 1 and a fluorine minus 1. So if we add plus 1 and minus 1 like that, do we come out with? What number do we get for that equation? It's 0, right? 1 minus 1 is 0. So that's going to be the charge of the formula unit of every ion in compact. So is calcium here? Is this a metal or a non-metal? Metal, right? What about sulfur here, non-metal or metal? Non-metal, right? So will these two things combine to form an ionic compound? Yes. Yes, they will, right? Why is that? Because metals and non-metals, they combine to form ionic compounds, right? So what would be the formula for calcium combining with sulfur? What would be the formula, the unit formula of this ionic compound? It would be CAS, right? It would be CAS. How did you figure that out? Because we add plus 2 plus minus 2, and that equals what? 0, right? So all we got to do is pick these things together and say CAS. That's the unit formula of these compounds, or this compound in particular. Let's look at the combination of calcium and fluorine. Calcium's most common ion is what? What's the most common oxidation state for calcium? You'd be telling me. What is it? 2 plus, right? 2 plus. If you don't know that, you better know it by Friday, seriously, because there's a lot of this stuff on it. So 2 plus, right? How did you get that? Because it's in the group 2 element, and it's written right here on the board, right? So if we go over here, 2 plus, what's the most common oxidation state for fluorine? 1 minus, right? How did we get that? Because we looked at the board there, or it's in the groups of elements. That's minus, right? So if I added those two numbers together, would I get 0, 2 plus plus 1 minus? No. What would I have to do to this to get 0? If I put plus 2 minus, or plus to minus 1, and the value equals 0? No? What would I have to add here? No multiplying. Another fluorine? Yeah, well, wait. What would I have to add to this equation to make it equal 0? Negative 1. A negative 1. Fluorine is in the number, OK? So a negative 1, right? Now, all the fluorine people, which one is negative 1? Or plus 2, right? So which one's negative 1? Is it fluorine or calcium? Fluorine, right? So I know you guys already figured this out, right? So do we have 0 if we add plus 2, plus minus 1, plus minus 1? Yes. So what's going to be the unit formula of this ionic compound? C, A, F, 2, 2, yes, right? Straight up, right? You got it. Cool. OK, so let's talk, yeah, it's a bunch easier than what everybody's trying to make without speed, right? We don't know it yet. If you're still having trouble, just go over this, go over it, go over it, OK? Let's talk about the transition metals now fairly quickly. Remember we talked about the transition metals not having regular oxidation states. So how are we going to figure out what the oxidation state of the transition metal is? Well, let's look at a couple of compounds. So we got F, E, C, L, 3, OK? So iron is a transition metal, right? Everybody knows that because if we look for it, we find it in the D block, OK? We know the transition metals don't have regular oxidation states. So we got to figure it out, OK? So what do we know about C L? Negative 1 oxidation state. So we got 3, negative 1, right? We got negative 1 plus negative 1 plus negative 1. That equals 0? Does that equal 0? No, what does that equal? Negative 3. So what do we need to balance this equation to make it 0 over this? A plus 3. How many iron atoms do we have in this thing? Just one. So the oxidation state of this iron atom plus V plus 3. Like that, OK? So let's start naming these compounds now, OK? So this compound, well, let's go back to these compounds here. This compound is called lithium fluoride. Remember, the anion always drops the INE and adds the ID. You always name it cation first, anion second. This one is called what then? Calcium sulfide. What about this one here? Calcium fluoride, OK? No dyes, tris, anything like that, OK? What about this one? Ferris chloride. Well, no. What about this one? Iron chloride. Iron chloride is what you should be saying, right? Unless you have read more in advance, right? This, then we're not going to use that ferris chloride business. Let's just call it iron III fluoride, OK? Because it's a transition metal, if we just draw this, we might not be understanding, or if we're just talking, if I just say iron chloride, right? It's not evident which iron chloride I'm talking about, OK? So what we have to do when we're writing these names out, we got to call them iron, and then tell the oxidation state of that iron atom. And we write that in Roman numerals, iron III fluoride. So the formula between sodium and oxygen is going to be Na2O, right? That, the formula's name is going to be sodium oxide. Between lithium and bromine, it's going to be LIBR, lithium bromide. Go through the rest of them yourself, OK? And then there's the talk about the compound names. The last thing I want to talk about before we leave, so there's naming transition metals, OK? There's some transition metal complexes. You can look, there's a bunch of these names up there. I'd like you to look at them. Polyatomic ions, the same sort of thing. The polyatomic ions are a group of atoms that are put together that are charged. So I feel for 3 minus, this group of atoms is charged, OK? All has a 3 minus oxidation state. Unfortunately, you're going to have to memorize this list of polyatomic ions, OK? So you're going to have to be familiar with them before Friday, OK? It's not as difficult as it looks. I promise you. There's plenty, plenty of examples. And the last thing I want to tell you, the last thing, I still have a minute, guys. I still have one minute. It's what we're naming covalent compounds. Now, covalent compounds are compounds that are constructed between two non-metals, OK? A non-metal bonding to a non-metal. When we're talking about covalent compounds, we name them by the order of atoms that are in the compound, OK? So for example, we have this one where we have sulfur and oxygen. We call it sulfur. And in this case, we call it dioxide because there's two oxygens. Notice the difference in naming ionic and covalent compounds, OK? In covalent compounds, you put the di-tri-tetra. You indicate the number of atoms that are in that compound. Whereas when we talk about ionic compounds, we're not calling those calcium diphylora, OK? The thing you've got to know is mono, di, tri, tetra, pentate, hexaheptate, octa, nana, and deca, OK? Of course, tri is like a tricycle, right? Vi is like, I don't know, di is like, sorry, di is like, yeah, di, di, di is like di-virgin, di-virgin. Mono, mono, right? Tetra is like tetrapod. Pentas like pentagon. Hexas like a hexagon, OK? So octa is like an octopus, OK? So remember that, and that's all you need to know for the test. If there's any questions, please ask me. If you didn't sign in, the sign-in sheet is still going around. And those of you who are here to see this video, please do not forget to like, comment, and subscribe.