[MUSIC] In this lecture I'll review the driving forces for ion formation, and show how this was translated into a set of rules for assigning oxidation states to atoms in all kinds of compounds. Finally, we'll wrap up by talking about some of the common polyatomic ions. Look at a periodic table now, to remind yourself of a couple of periodic trends that have been discussed in this course. One of the things you might remember is that the metals have relatively small ionization energy. They tend to lose electrons and become cations because of this. Nonmetals on the other hand, have relatively low energy open orbitals. Which means they have large electronegativity. They tend to gain electrons, often by taking those electrons away from metals, and the nonmetals tend to form anions. The driving force, once again, for these processes. Are that, the system wants to lower its energy. So the electrons can go to lower energy by leaving metals and joining nonmetals. Another driving force that you may have heard about is that atoms of the main group elements want to become isoelectronic with the noble gases. What do I mean by an atom becoming isoelectronic with a noble gas? Well let's take a look at the periodic table and see. This periodic table has drawn hydrogen above the alkaline metals. I don't actually think that's the best place for hydrogen. Sometimes it does loose an electron, like in alkaline metals, but other times it doesn't. So the first thing I'm going to do to make the periodic table more the way I personally like it is move hydrogen. So you see hydrogen there is right above lithium. I'm just going to dissolve it away. Move it over here. You know, I like to put it kind of in the middle. Sometimes you see it above lithium. And sometimes it's above fluorine. But, really it sometimes act like a metal and sometimes like a nonmetal. So let's just put it in the middle. Shall we? Okay, there it is. Now, let's look at the main group elements. It's the main group elements that like to be isoelectronic with the noble gasses. So what does this word isoelectronic mean? That means it has the same electron configuration, [SOUND] and we haven't talked about electron configurations. We'll get to that later. But for now, let's just say, it wants to have the same number of electrons as a noble gas. [SOUND] So, for the main group elements, that includes group one and two, remember? Does not include the transition metal. So we're not going to talk about those right now. I'm just going to cross 'em out. And I'm going to cross out the. Lanthanide and actinide are inter transition metals. But, it does include also group 13 through 18. Okay, so that's what I'm focusing on right now. Which of the noble gases over here, is closest to the number of electrons to lithium? Well helium is, right? Helium has two, so lithium, can lose one of the three electrons, and become isoelectronic with helium. And that's what lithium likes to do. In fact, everything in group one, all of the alkaline metals, prefer to lose one electron, they lose that electron quite easily. It's a very, it's a very low ionization energy electron if you will. And they become isoelectronic with noble gases by losing one electron. So sodium becomes isoelectronic with neon when it's sodium one plus. The potassium cation is isoelectronic with the argon. You see how that works? What do you think is true about the alkaline earth metals, that's group two. So beryllium is the, the smallest alkaline earth metal. Again, it wants to be isoelectronic with helium, so what does it need to do? Well, beryllium needs to lose two electrons and that's what beryllium does most of the time. So when beryllium loses two electrons it has a plus two charge. Its most common oxidation state is plus two, and it becomes isoelectronic with helium. Calcium, for example, loses two electrons. So neutral calcium has 20. The most prevalent calcium cation has only 18 and is isoelectronic with argon. Do you see how that works? Now the transition metals can gain or lose different numbers of electrons so they're a little more interesting. Let's focus on things that usually do the same thing. Let's jump over and look at the halogens that have fluorine at the top, so now we're looking at group 17. What's the easiest way for fluorine. Let's in fact, let's use green to outline these. What's the easiest way for fluorine, chlorine, bromine and iodine to be isoelectronic with the noble gases. Let's take iodine as an example. What does iodine need to do to become isoelectronic with xenon. Well it has 53 electrons and xenon has 54, so it needs to gain one electron and that would make it a minus one anion, wouldn't it? It becomes the iodide anion and that's what the halogens most commonly do. I'm not going to worry too much about what the metalloids do. So I'm going to draw a little line like this, okay. And I'm also not going to worry too much about what the late metals do. These metals down here were ten in nadar. You see ten in nadar. Those are called the late metals. So we're not going to worry about those right now either. because they also like to do strange things. But let's focus now that we've done the halogens let's do the next nonmetal column over. Let's do the column that has oxygen sulfur and selenium in it. And not worry about tellurium down here because it's a metalloid. Well oxygen would like to be isoelectronic with neon. Sulfur would like to be isoelectronic with argon and selenium would like to gain two electrons to become isoelectronic with krypton. You see that? So, all of those three elements tend to have the most common charge and their monatomic anion is minus two. Monatomic meaning one atom, right? [SOUND] So the monatomic anion common charges are drawn over here. Minus one and minus two for those two groups. And the monatomic cations, I'll just put in red over here. [SOUND] monatomic meaning one atom. Have a plus one for the alkaline metals. A plus two charge for the alkaline earth metals. And then actually we can jump across the transition metals and find aluminum right here, which I sort of crossed off but didn't mean to. Aluminum tends to have a charge of plus three. It likes to lose all three of it's electrons to become isoelectronic with neon. Okay, so that's some common charges that occur. The driving force is that the nonmetals like to gain electrons because of their high electronegativity. The cations prefer to lose their electrons because of their high energy electrons, and in both cases, electrons are able to go to lower energy. Okay. So, electrons can leave the metal, join the nonmetal. The metal becomes a cation. The nonmetal becomes an anion. All of this is played out in this rule, list of rules I'm about to give you for oxidation states. So, to treat compounds in a systematic way. Conventions are used to assign electron ownership for particular atoms in the compound. There have been two major systems devised for doing this and these are human constructs. One of the systems is called oxidation states and that's what we will be talking about today. There's another system called formal charge, which I actually really like, that gets used quite a bit in organic chemistry. But, today, we're going to learn about oxidation state. Because, that's the system of accounting for electrons, that is used when you name ionic compounds. So again, this is a human construct, we decided to account this way. So this is just an electron accounting system. And in the assignment of oxidation states what we do is we're assuming that all bonds are ionic. Even bonds that are not ionic we assume are ionic. And in an ionic bond. All of the bonding electrons go to the more electronegative atom, right? And an ionic bond, the electrons leave, the highest energy electrons leave the less electronegative element, which is usually a metal. There could be a nonmetal that has relatively low electronegativity and they join the atom that is more electronegative. So this is, this system of oxydation states is treating most bonds as ionic bonds. The only bonds that get treated as covalent bonds are bonds where the two atoms are the same type, you'll see what I mean in a minute. Here's an interesting figure that shows the electronegativities of some of the main group elements on the polyne scale. We can use the difference in electronegativity to determine which atoms would gain the electrons in any particular compound. First, let's zero in on this figure. So make sure that you understand what's going on. Fluorine is the most electronegative element, and it has an electronegativity of 4.0. If you compare that to a metal, for example, caesium, here is a large alkaline metal, and it's electronegativity is only 0.7. All right? So, fluorine is much more likely to take electrons away from another atom than cesium is. Let's do a model compound. How about, what would happen if calcium and sulfur got together to make a compound. What would the oxidation states, or the charges on those monatomic ions be? So calcium is a metal? Let's just do this is red, here's calcium. In order for it to be isoelectronic with a noble gas, it wants to lose two of it's electron to become isoelectronic with argon. Argon has 18 electrons. Calcium normally has 20 electrons. So it loses its two electrons, and it gains or, it doesn't gain anything, it loses its two electrons to become charged with plus two. And sulfur's over here. Sulfur has a higher electronegativity. Sulfur's electronegativity is 2.5. So it's going to take the two electrons away from calcium and not give them back. So the sulfur ends up with a minus two charge. Let's write that down over here so we can see it clearly. I've got calcium with a positive two charge and sulfur with a minus two charge. These numbers, two plus and two minus, sorry, my minus is a little bit crooked there, those are the oxidation states on those atoms. [SOUND] And if I wanted to write what the compound is, the molecular formula, the molecular formula is just. CaS, for calcium sulfide. I don't need to write the charges, do I? But we know that the oxidation state here of the calcium is plus two and the oxidation state of the sulfur is minus two, okay. Because in this ionic compound, the sulfur has taken both bonding electrons from the calcium. One of the things we need to be really careful about when we do oxidation state assignment, is that we need to make sure that all the oxidation states together add up to the net charge on the species. So here, the net charge on calcium sulfide is zero. It has no charge overall. If I take the oxidation states and add them up, they should add up to zero. So I have plus two for calcium and I'll add that with the oxidation state on the sulfur, which is minus two and that does indeed add up to zero. So I'm in good shape, works out just the way we need it do. Putting a little check there. All right, so oxidation states you need to make sure you balance the books properly. Some of you in the United States might recognize this guy, and I apologize if you don't follow the news of the United States you probably don't know who this is. But you don't want to do what this guy did. He did not keep the books balanced. So don't be like Bernie Madoff. You have to make sure that when you are assigning oxidation states, the sum of the oxidation states of all of the atoms in the species that you are assigning oxidation states to. Add up to the overall charge on that species. So, for ions, the sum of the oxidation state should equal the overall charge on the ion, and you'll see what I mean when I do some examples. And for neutral molecules, which is like calcium sulfide, which we just did. The sum of all the oxidation states of the various atoms in the molecule should equal zero. That's always a good final check to do, to make sure that you're assigning oxidation states properly. This is not a particularly difficult thing to do, because there's a list of rules. You just follow the list of rules and you should get the right answer. You might want to actually print these rules out, and there's a copy of these rules in the section of useful information on the main course page. But here I'll go through the rules one at a time and we can talk about why the rules are the way they are. The first rule is that if you have pure element all of the atoms have an oxidation state of zero. So this could be for example all kinds of different things, it could be metallic sodium, if you just had sodium with a solid next to it. The oxidation state of the sodium there is zero. There's all sorts of diatomic elements. For example, there is diatomic chlorine. In diatomic chlorine, the oxidation state of the chlorine is zero. This is the only case where the bond is not treated as ionic. Here we're assigning one electron from the bond to each atom. The next rule is that if you have a compound that contains fluorine and a different type of element. Then fluorine has an oxidation state of minus 1. That's because fluorine is the most electronegative element. So no matter what else fluorine bonds with, it's going to take an electron away from that other element to make the fluoride anion which is fluorine with a minus one charge. The next rule deals with metals giving their electrons away. So the alkali metals tend to be plus 1. The alkaline earth metals tend to be plus two. And aluminum tends to give up all three of its electrons to be plus three. So look on your periodic table and you'll see that the group that has lithium, sodium, and potassium in it. Rubidium, right? Those tend to give their electrons away quite easily. Beryllium, magnesium, calcium likes to be plus two and aluminum likes to give all three electrons away because those species have very low electronegativity so if they bond with any of the nonmetals they give their electrons away. So the first three rules are pretty firm, right? Elemental form equals zero, fluorine turns into fluoride if it's in any kind of compound, and then these metals that are on the far left side of periodic table tend to give their electrons to anything else on the periodic table, so that all makes sense. Now we get to the rules where things can have different oxidation states depending on their situation. For example, hydrogen which you remember I moved from being on the left side of the periodic table to more in the middle. Hydrogen is usually plus one. It does frequently give up its electron. What we would call that? Lemme draw a picture of that, here's a simple Bohr model of the hydrogen right, we've been doing. If it loses it's electron say minus electron, what do I have left? What is that? What would you call that? I'd call that a proton wouldn't you? Okay, so hydrogen is usually H plus which is a proton, but if the hydrogen, it encounters a metal like, let's say, a neutral sodium, the hydrogen can take the electron away from the sodium because hydrogen has an open orbital at lower energy. And in that case, it becomes hydride. So, H1 minus is hydride and that's hydrogen with two electrons. I'm drawing two dots for the two electrons. So, hydrogen is plus one sometimes and it's minus one sometimes. Oxygen is usually minus two. But it can be minus one. For example, it's minus one in hydrogen peroxide. You have to follow these rules in this order, all right? So you always want to go from top to bottom when you're doing these rules. The final rule is that halogens are usually minus one. Halogens tend to be fairly electronegative. And so they take electronics away from moth-uh, most other species on the periodic table.
