Ionic, Covalent and Metallic Bonds

Hey gang, this is mister Spencer. Welcome to another center chemistry podcast. What we're going to do today is we're going to talk a little bit more in depth about different kinds of chemical bonds and also the kinds of properties that compounds that are made out of these kind of bonds have. So let's just go into just a quick review. We've been talking about this quite a bit. We've talked about what makes atoms happy. And except for hydrogen and helium, all of those other atoms in the periodic table follow what's called the octet rule, which is the idea that they want to have 8 valence electrons. Not just 8 electrons, but 8 valence electrons, or electrons that are in that outer energy shell that are in charge of all of those chemical reactions and things like that. Noble gases, those are already happy. They already have their 8 valence electrons, and that's why they don't react with anything else. But all other atoms form bonds in order to get those 8 valence electrons. So chemical bonding is important because they are trying to get those 8 valence electrons. Now there's different ways of doing that and we've talked about this before. You have your ionic bonds. Those are your, well, that's kind of the nice way of saying it is give and take so some atoms will give away electrons, some will take them away. But what you end up with is charged particles. And those charged particles, those are called ions, now when you, when you lose electrons, okay, that's going to give you positive ions or what are called cations. If you gain electrons, that's going to give you negatively charged ions called anions. And we'll talk more about that later. If you have another kind of bond, those are your covalent bonds. Those are your nice bonds. Those are the ones that watch Sesame Street because they share. And another one that we're going to talk about briefly, but that you need to know about our metallic bonds. And these are interesting because these have what's called a sea of electrons. Of valence electrons actually where they're just sharing all of these valence electrons. And you get some interesting properties that come from that. So let's go and talk more about ionic bonds. Now, remember, these are your give or take bonds, these are the ones where you have some electrons that are being given from one atom and are to another atom. Wow, I can't talk today. But here we have our famous salt. This is sodium atom. This is a chlorine. Now, if you look at the sodium atom and we've talked about this before, you see that there's different energy levels. And that first energy level, you have two electrons. In the second energy level, you have 8, and then the third you can have up to 18. But in this case, sodium just has one, okay? So there's one valence electron. Chlorine over here. That has two electrons in its first energy level. 8 electrons in its second energy level. And 7 in its third. So what's going to happen here is sodium is going to say, hey, you know what? I don't need, well, I have this one valence electron. I either have to give it away or I have to get 7. So it's going to be a whole lot easier for sodium to just give away that one valence electron. And you'll notice that when it does that, right here, it has a complete shell. So it's still going to follow the it's still going to follow the octet rule. Chlorine, on the other hand, is going to say, hey, I've got my 7 valence electrons. I either need to get one or I need to give away 7 and it's a whole lot easier to get one. So what happens is sodium donates or gives up its electron to chlorine. So when that happens, I'm losing an electron, here with sodium, so this loses an electron, and now I have one more proton in the nucleus than I do electrons in the electron cloud. So that means I'm going to have an extra positive charge. So right here I have a charge of plus one. Chlorine on the other hand, it's just gained an extra electron. It has an extra negative charge now. So it gains an electron and now has a charge of negative one. So what's going to happen here is you are going to have a difference of charge. And positive and negative attract each other opposites attract. So that's how these stick together. And that brings us to what's called crystal structure. Crystal structure is just how these ionic bonds stack up. Let's just do a quick review on this. Remember on the periodic table, you've got this pattern that we've talked about now, here's your alkali metals. And all of those atoms there, they all have one valence electron. So when they form ions, what happens is they give up that extra electron. So these all have a charge of plus one. And if I could write that would help out a lot. These right here, your alkali earth metals, those all have two valence electrons. So it's going to be a lot easier for them to donate two electrons than it is for them to gain 6. So these all have a charge of plus two. Over here, these are your noble gases. These all have a charge of zero. Man, I'm having a hard time writing today, okay? These are your halogens. And those all have 7 valence electrons, so what they're going to do is they are going to gain an electron. So they all have a charge of minus one. And right here, this is your oxygen group. Those all have 6 valence electrons, so they're all going to gain two electrons and end up with a charge of minus two. And there's other ones that can have charges of plus three and minus three and things like that. But we're not going to talk too much about that. So here's what happens. Let's take the example of sodium chloride, where you have a sodium ion with a charge of plus one, and a chlorine ion with a charge of negative one. So because the total charge needs to be zero, you have a ratio here of one to one, so one chlorine or one sodium atom for every chlorine atom or chlorine ion, sorry. And they would stack up kind of like this. You could also have calcium and calcium whereas calcium calcium is like right. Right there. And oxygen, oxygen's up here, okay, hopefully that makes sense. But calcium has a calcium ions have a charge of plus two, oxygen atoms have a charge of negative two. So they are also going to form crystal structures with a ratio of one calcium to every one oxygen. Now, if we have some different charges, then you're going to end up with sodium. That's right over here. And chlorine, where's chlorine at chlorine is right down here, okay? So chlorine has a charge of negative one, sodium has a charge of plus one. Oxygen, we already have that. So if we have a sodium ion with a charge of plus one, and an oxygen ion with a charge of minus two, that means we're going to end up with a crystal structure that has a ratio of two sodiums to every one oxygen. And we could also change that around if you had calcium with a charge of plus two and chlorine with a charge of negative one. There, you're just going to have a ratio of one to two in that structure. And we'll talk more about that later. All right, now some things with ionic bonds, we did a lab the other day where you were investigating this. And you found that ionic bonds have high melting points. You take sodium salt, okay? If you were to go throw that in your oven, there's no way you're going to be able to get that to melt. It's just not going to happen because the melting point or the point where something goes from a solid to a liquid, at least for salt is 801°C. So that's pretty hot. Also with ionic bonds, they are a solid at room temperature. Ionic compounds being liquids or gases when they're at room temperature. And the reason is because these ionic bonds are this attraction between positives and negative charges is so strong that it holds those it holds those molecules or those particles together. They're also very brittle or hard, okay? If you've ever used salt to put on your to put on your sidewalk when it's when it's snowy out or something like that or icy. You can see that those chunks, those chunks of salt are pretty hard. And also one of the fun things is when you take ionic compounds and put them in water, they break up into those ions. And those electrical charges are able to move around and it conducts electricity. All right, which is why water on its own doesn't collect conduct electricity, but it does when you put when you dissolve things in it, there's a great MythBusters about what happens when you drop electrical appliances in bathtub water. You should watch it sometime. All right, so that's enough with ionic properties. Now we have covalent bonds. And these are your happy ones. These are the nice ones. Okay, they share they share their electrons. So if you have a pair of electrons being shared, you have a single bond. If you have two pairs of electrons being shared, you have a double bond and three pairs of electrons being shared. You have a triple bond. And triple bonds are much stronger than single bonds. All right, so this is kind of what happens here with water, H2O. You have water or I'm sorry, you have a hydrogen atom that just has one electron and you have oxygen, which has 6 electrons, so you get two hydrogen atoms that share their electrons with oxygen, so now you can say hydrogen is following the duet rule. It has its two electrons. This one is doing the same, and oxygen here now has its 8 valence electrons. But that brings us to what's called electronegativity. When atoms form bonds, especially when they form covalent bonds. Atoms pull on electrons with different amounts of strength. Kind of think of it like you have a mean big brother or for some reason you have to share your lunch money with the school bully. And you guys are sitting in line and you're looking at that and you have your $2 or whatever. And you say, you know what, I really want tater tots today. I really, that really sounds good, okay? And the school bully says, you know what, Spencer? We're not having tater tots today. We're having pie. And I say, I want tater tots. And he says, we're having pie. And guess what we end up with? We have pie. Okay? We both get to eat, but the stronger person gets to choose what we're going to eat. So think of it like this. Well, actually, you know what it is. The trend for electronegativity right here, this is fluorine. Okay? This is the most electronegative. And right down here is francium. Is that FR? I think that's FR. That's the least electronegative. So as you move this way and this way, on the periodic table, that pull for electrons becomes greater and greater. So when you're up here in the periodic table, those atoms are going to pull those electrons in that bond with greater force than ones that are in this area. They still have that Bond, so they're still sharing electrons, but it's not as much. And if I was smart, actually, let's see. I had to fumble around for this. Okay. All right. Oxygen has an electronegativity of 3.61. Hydrogen has an electronegativity of 2.30. So you've got your electrons here, got your electrons that are in this bond between oxygen and hydrogen. But you can see that oxygen here is the bully. Okay, oxygen's the one that's going to be pulling on those electrons with greater force. So those electrons are pulled this way in that Bond. So there's still a bond there, but oxygen hogs those electrons a little more than hydrogen does. All right, so that is that leads us to what's called polarity. And polarity is just how well those electrons are shared. So. Yeah, how electrons are shared. In covalent bonds. In nonpolar covalent bonds, what you have is an equal sharing. Of electrons. In polar covalent bonds, you have an unequal sharing. So that's kind of what happened with the water molecule example. So now with things that are with compounds that are covalent, they also have certain properties, which are quite different than ionic bonds. So some examples of some covalent chemicals are compounds. Things like butter, okay? Oh, a lot of greenhouse gases and smog, those tend to be those tend to be covalent compounds, sugar, and other things like that. Those are those are covalent compounds. Okay, so covalent compounds have a very low melting point, butter is going to melt a lot faster than salt does. They tend to be gases or liquidy. You can have solids, but you know that there's a big difference between sugar, butter and salt, okay? With how hard they are. And that's also they tend to be soft. And when you put these in water, they don't conduct electricity very well. So covalent properties tend to be the opposite of ionic properties. And the last one that we're going to talk about is just metallic bonds. The idea here is that whenever you have a metal, all of those atoms share their valence electrons. So the valence electrons are just zip around all over the place. And just go and kind of wherever they want and all the atoms are happy because they don't have to they have their valence electrons that they're sharing with everybody else, which is nice because this allows for some pretty unique qualities. They are metals because of the sea of electrons that they have are very good conductors of electricity and heat. With electricity, those electrons are able to flow through that metal structure very easily, which is good for us. We have a lot of copper wire in your house that makes it possible to live the way that we do. Heat is also able to go through metals very well. Metals are also malleable, meaning you can hit them with a hammer, you can bend them, you can flatten them, just like with aluminum foil. They are ductile, meaning you can draw them out into wires and they're shiny too. And shiny is good. So those are metallic bonds. And that is it. So we've talked today about a little more in depth about ionic bonds. We've talked a little more in depth about covalent bonds and also went over briefly. Metallic bons. So let me know if you have any questions.