Chad's General Chemistry Videos
- 19.1 Oxidation Reduction Reactions and Oxidation States
- 19.2 Balancing Oxidation Reduction Reactions
- 19.3 Galvanic Cells
- 19.4 Standard Cell Potentials aka emf or Voltage
- 19.5 Nonstandard Cell Potentials the Nernst Equation
- 19.6 Reduction Potentials and the Relationship between Cell Potential, Delta G, and the Equilibrium Constant
- 19.7 Electrolytic Cells
- 19.8 Electrolysis Calculations
All right so, your next chapter you're really going to get into all the different structures and Lewis dot structures. We'll tie that into the next chapter with all the different geometries and stuff like that. We'll probably get just a little bit ahead of where you are in lecture, but before we talk about that we really ought to just talk really generally about the three different types of bonding that goes on. So, you've got ionic bonding, covalent bonding, metallic bonding. So ionic bonding, we talked about that a little bit earlier, let's just pretend that a bond is like a pencil. And I just, sorry, how about an electron is like your pencil and I just stole your electron. It's mine. Because I took your electron what charge do you now have? Positive charge? What charge do I know how because I took it? Negative. And so you're about to get pissed. You're like "Dude I'm trying to take notes, you stole my pencil." But then you realize that I'm negative and you're positive, and you're like "Oh but I guess, Chad's kind of cute actually," and there's an ionic bond going on. That's kind of what happens with an ionic bond. So the nonmetal steals the electron from the metal, creating a difference in charge. It's a transfer of electrons. But that leads to a bond. In this case an electrostatic bond, or ionic. It involves a, if I can spell it correctly, transfer of electrons. On the other hand, in a covalent bond, what's going on in a covalent bond? Not a transfer of electrons. It's two nonmetals. And it's a, what's going on with the electrons? Yeah, so it's a sharing of electrons. And as you also said, you most notably recognize it because it's a nonmetal with another nonmetal. So, when you see a bond between two nonmetals, that's a covalent bond. Whereas ionic bond, what kind of elements are we looking for in ionic bonding? Yeah. Metal with a nonmetal. Cool. The last type, metallic bonding, we're not going to talk a whole lot about metallic bonding but in metallic bonding, you will see pretty much a metal in a metal. And it may just be all a pure metal. You know. If I have pure iron, pure iron you have tons and tons of iron atoms. And the atom the iron atoms are all held together by metallic bonds. Now the one thing I will say about metallic compounds, is there's a number of properties of metallic compounds that you want to know. And so in this case, it's not completely just like, we can't really call it just a sharing of electrons, or we can't really call it a transfer of electrons. But there is something we can do to describe this. And we call it a sea of electrons. So it turns out the valence electrons in metals are not held by anybody or any kind of, you know, any specific nucleus all that tightly. And so they're full of a sea of loosely held electrons. And that is what makes metals electrically conductive. So some things we should know metals are conductive both electrically conductive and they conduct heat really well, as well. So what else. Anyway remember the other properties of metals that you should remember? Shiny. What's the word we use though? Good. So lustrous or they have a luster associated with them. What else? So... not all of them do. But many metals, like the alkali metals, will oxidize easily. But not a general property of all metals. They also can be pounded into sheets. Anybody remember that? Malleable. And they can also form wires, we say be drawn into wires very well. Starts with a D. Ductile. So, cool! So, yeah they can be drawn into thin wires, that's what ductile means. So, malleable means metals can be pounded into sheets. Luster means they are typically shiny, conductive, both of electricity and heat. But these are general properties we associate with metals and you just want to memorize this. Oh yeah, those four properties, yeah, those are what we talk about with metals. Whereas if we look at the properties we talk about for ionic and covalent compounds. So if we look at ionic compounds, ionic compounds we usually look at super high, not always, but fairly high melting points and boiling points for ionic compounds. If you ever tried to melt salt? Like, not dissolved in water, I mean just like straight up melt it. Yeah. Good luck. Like you try and you just put pure salt in, you know, a pot on your stove. You're stove doesn't get hot enough to melt it, it's super high melting points. Well salt is an ionic compound, and ionic compounds generally have high melting points and boiling points. So they're also typically, fairly, brittle. Typically fairly brittle. If you look at a typical ionic compound, you've got cations next to anions, next to cations, next to anions, next to cations. And then in the next layer you'll have an anion next to a cation, next to an anion, next to a cation, next to an anion, and they alternate. And so, if you look, a cation will always be surrounded by anions. An anion will always be surrounded by cations. And so everybody loves their neighbors. I love all my neighbors! I'm attracted all of them. So however, if you were to like hit a big ol' block of salt with a hammer, what's going to happen? It's not going to dissolve. It's brittle. But it will just shatter, right? Get a big salt crystal, like the ones you get for salt licks and horses, and you hit it with a sledgehammer, it's going to shatter into a million pieces. And the reason is they're brittle. The reason they're brittle is when you hit it with a hammer, what you cause is stress and the layers might shift. And so notice if these layers shift just a little bit; I'm just going to shift this middle layer over ever so slightly. I'm not going to draw the third layer here but, though if I shifted the top layer from the middle layer, now how do these two layers feel about each other? Good. This one doesn't like, they don't like each other! And you see a separation. That's why when you hit it you also get nice clean areas where the salt fractured and stuff like that, because it was an entire layer all the way down the line where they repelled each other and it fractured. Cool. But that's why compounds are typically bristol- brittle. Any stress, like a hammer, puts a little shifting in the layers. Cool. Now covalent compounds; we got two types of covalent compounds. Molecular is by far the most common and the other type, what we'll call network covalent solids. Now we typically think of, you know, ionic compounds as forming crystals. Whereas, you know, most covalent compounds, the molecular ones, we don't usually think of those as big as crystals you know. At least not at room temperature. But network covalent solids, which are covalent compounds, do exist as crystals. Anybody name for me a network covalent solid? Bling, bling? Diamonds. So what's a diamond made of? Pure carbon. And carbon, being a nonmetal, all the carbon atoms are bonded to each other with covalent bonds. And so it's covalent but it does form a crystal. A crystalline solid. In this case we call it a network covalent solid. The other one you might see, SiO2 -if I can spell it again- quartz. So silicon oxygen -and silicon is kind of like a semi-metal here, metalloid- but they still consider it covalent bonding, so in this case, it also forms a crystal quartz. So those are your two big examples of network covalent solids. Just about any other example of all nonmetals is going to be a molecular compound, one that actually forms molecules. So notice ionic compounds? They don't form molecules. Network covalent solids? They don't form molecules. Only molecular compounds form molecules. So if I said "Hey I got a diamond molecule." You'd be like, "You're a liar, there's no such thing." If I said, "I have a sodium chloride molecule." You'd be like, "You're a liar, there's no such thing." But if I told you I have a water molecule, well water H2O, is a molecule and I could have a water molecule. In all likelihood, I've probably got billions, millions, trillions or a mole of water molecules, or some huge number of them, but it is a molecular compound. Nonmetal, nonmetal. And they'll exist as individual molecules, not big old crystals. At least at room temperature usually. Your molecular compounds are typically associated with much lower melting points and boiling points. And that's suffice it to say here. When I say lower, obviously that means relative. Well relative to ionic compounds or relative to network covalent solids. Anybody ever try and melt a diamond lately? Yeah, I wouldn't recommend it first of all. But even if you try you're probably not going to get there. So diamond has crazy high melting point. I don't even know how high, it's in the thousands of degrees. So same thing with ionic compounds. NaCl? You want to melt salt? Like 800 degrees Celsius. Good luck. Your stove at home is not going to get there. Anybody ever try and melt ice? Yeah, notice it's already melted at room temperature right? You've got to heat water above zero degrees Celsius and it will melt. Molecular compounds have much, much, much lower melting points and boiling points. So and we'll talk about this more in one of the later chapters. But for the next exam, the idea here is that there's relatively weak forces holding the water molecules together. So in a molec- in the case of ice or in any molecular compound. Whereas what holds the atoms in this case together, in this case the ions, is actual ionic bonds. What actually holds all the atoms together here are actual covalent bonds, rather than the weak forces. And so when you melt or boil something you have to break apart those forces. Well here they're weak. Great! They're easy. It doesn't take as much heat to do it. Here the bonds are really strong and takes a lot of energy to actually break covalent or ionic bonds. Cool. Any questions on just the kind of three different types of compounds and characteristics of each? Cool. By the way, network covalent solids? These are your two examples. Memorize that. Know that, oh yeah, diamond, quartz, carbon, silicon dioxide, those are my two common examples of network covalent solids. There's a couple of examples out there but none you're probably likely to see. But those are worth knowing. I'll talk for a minute about lattice energy. So we'll talk very briefly once more about ionic compounds. Lattice energy is associated with ionic compounds. But we'll probably spend the rest of the night after that talking exclusively about molecular compounds. So lattice energy. So technical definition of lattice energy is probably easiest to describe with a chemical reaction. So it's for ionic compounds, and it's the amount of energy it takes to take a solid ionic compound and break it up into separate ions. And so it's a measure of how strong the ionic bond holding that crystal is together. Because to break it up, this is the energy it takes. Now there's one thing I haven't written here and it's the one thing that students most commonly forget. It's not just breaking it up into ions, but you have to remember the phase of those ions. Anybody remember? They got to be gaseous ions. So the textbook definition for lattice energy is the energy it takes to break an ionic compound up into its gaseous ions. So really important technical definition. You may get a question that just says "Which of the following is the definition of lattice energy?" You got to come up with the right definition and it's usually the gaseous part that is the distinguishing factor. Or you may get a question that just says, "Which of the following chemical reactions describes the lattice energy of NaCl?" Well it would be exactly this. If that said aqueous, liquid, or solid, those are all wrong. Has to be gas for the products. Okay so that's our textbook definition of lattice energy. But you also have to be able to compare lattice energies, and there are two things that affect the strength of lattice energy. So, the first is charge. And the second is size. What would I be looking- So charge has a more significant effect. Size a little bit less or so. And so the first thing you look for in one is the difference in charge in comparing ionic compounds. What kind of charge would I be looking for to look for a high lattice energy? What's that? Well we're always going to have one positive ion and one negative ion. But notice again this lattice energy is the energy it takes to break these apart into separate ions. And so it's a measure of how strong the ionic bonds are. Well how do I make ionic bonds stronger? Well have higher charges. Notice if I was a positive one charge and you were negative one versus I was positive one and he was negative five, I'd look over at him going "Psh forget you, he's negative five! I got a bigger attraction here going on." And so higher charges. When I say higher I mean larger in magnitude. More positive and more negative, in both the cations and anions. Plus two is better than plus one. Negative two is better than negative one. Higher charges lead to higher lattice energies. Because it's a stronger ionic bond, it will take more energy to break them apart. And so in this case again, higher charge, then you get a higher- I'm just going to abbreviate it- L.E. A higher charge, higher lattice energy. Now size also comes into play here. And again this is a measure, lattice energy, of the strength of the ionic bond. Well again what kind of bonds are stronger? Longer or shorter? Shorter. Then what do I want about the size of my ions? I want smaller ions. And so in this case smaller ions lead to a higher lattice energy, as well. And again which one of these has have more significant impact typically? Charge. That's why I listed it first. And so the comparisons you might see on the test, so a typical question might just start listing for you some different ions. I'm going to mount a periodic table in here. So if you look kind of how they're showing up on the periodic table here. So the first thing you want to look at if I give you four compounds, and I say "Which one has the highest lattice energy?" Find all the charges. What's the charge on the sodium ion here? Plus one. And the fluoride ion? Awesome. Sodium is plus one in fact in all of these, right? What about chloride here? Minus 1 as well. Bromide? Cool. What about magnesium? And oxygen? Cool. Based on charge alone who's going to have the highest lattice energy? Awesome. In this case, if I asked you which one had the highest lattice energy you'd go with MgO. Now let's say though I give you just a slightly different question here, and I give you a very similar set of elements, but instead of MgO I put sodium iodide in there instead. So we got sodium fluoride, sodium chloride, sodium bromide, sodium iodide. Difference in charge? Yeah, there's all same charges. You're going to have to go to size here. Plus one. Minus one. Plus one. Minus one. Plus one. Minus one. Plus one. Minus one. So again for lattice energy, charge has a more significant impact. But if all the charges on the anions and cations are the same, then you've got to go to size by default. And so if I'm looking for the largest or highest lattice energy again, what kind of size are we looking for? Smaller size. Which of these are the smallest size ions. NaF. And so in this case, NaF, would have the highest lattice energy. Cool. Those are typically the two aces as presented. Either you give a hands-down winner based on charge, or if all the ion charges are the same then you got to pick one based on the smallest size.
In this lesson you will learn:
-What hybrid orbitals are and the types of hybridization (sp, sp2, sp3 hybridization, etc.)
-How to identify an atom's hybridization from the Lewis structure
-The relationship between an atom's hybridization and its bond angles
|5||Trigonal Bipyramidal||90o, 120o, 180o||sp3d|
The hybridization and bond angles of an atom can be determined from the number of electron domains surrounding the atom. An electron domain is either an atom it is bonded to (it does not matter whether it is a single, double, or triple bond; it counts as one electron domain) or a non-bonding pair ("lone pair") of electrons.
In the molecule above the carbon atom has 3 electron domains as it is bonded to 3 other atoms but has no lone pairs. Therefore, it is sp2 hybridized and has bond angles of 120o. Incidentally, the oxygen atom also has 3 electrons domains (bonded to 1 atom and has 2 lone pairs) and is also sp2 hybridized.