MCAT General Chemistry High Yield: VSEPR Theory

Added: 2026-05-12

[00:00:01]This is a video we've been waiting to make for a very very very long time.

[00:00:06]We are going to talk about VPR theory.

[00:00:10]VPR theory is veilance.

[00:00:14]Actually you tell me veence what?

[00:00:17]>> Veence shell e >> electron.

[00:00:23]Next repulsion theory pr theory. Sorry pair repulsion theory.

[00:00:34]You get a lot from this. You get a lot from just the title.

[00:00:39]Veence shell electrons used for what?

[00:00:45]Bonding. Very good.

[00:00:48]Bonding.

[00:00:50]Meaning that we're primarily going to be talking about what type of compounds here are. We going to be talking about salts, ions, molecules. What are we talking about? Talking about molecules mostly.

[00:01:02]Electrons. What are they?

[00:01:06]If you're taking the MCAT, you have to be able to define an electron for me.

[00:01:10]This is the fundamental negatively charged particle that exists in probability densities known as orbitals.

[00:01:36]that surround the nucleus of the atom and contribute to chemical bonds.

[00:01:53]You can't read that doesn't really matter as long as you know what an electron is. Electron pair means we're not talking about what means we're not talking about radicals.

[00:02:06]We're not talking about radicals. We're not talking about lone electrons.

[00:02:11]We're not talking about electron transfer repulsion.

[00:02:20]Electrostatic forces rule all and like forces alhamdulillah. And like forces repel or sorry like charges repel.

[00:02:48]like charges repel.

[00:02:53]Yes, these like charges will repel to a point where they are as far away as possible and they stay there.

[00:03:21]Let us take a very very very simple example.

[00:03:35]Uh there's no real simple examples. Let's take the example of a carboation.

[00:03:51]Yeah, the carboation I'm going to draw it in this plane right here. Right. The p orbital is going to go like this.

[00:03:59]up and down like that. So if you imagine the white board is the plane at which the p orbital is coming up and down that means that one of the orbitals will come this way. One of the orbitals will come out at you and the next orbital will go into the board.

[00:04:20]What is the degree angle between these orbitals?

[00:04:26]120° right? Why?

[00:04:32]Why?

[00:04:36]>> That's like the most they can hold each other.

[00:04:38]>> It's the furthest away that they can possibly be. Because if you pulled these two orbitals closer to one another, or let's say you pulled these two closer to one another, if this was 119 and this was 121, and this is whatever it is, right? 120.

[00:04:55]Now what hap if you pulled these closer together by one degree and pulled those further apart by one degree now these guys are too uncomfortably close by one degree which means that this guy is going to push that guy away and this guy's going to push that guy away and you have to imagine this for the rest of this lecture I want you to imagine this that there is a force coming out of this orbital going this way and going this way repelling the things around it and there's a force coming out of this orbital going this way and going this way and one going this way and one going this way and they are all pushing pushing one another away with the same amount of strength because they are the same number of electron pairs in the same position in the same molecule at the same time. They're all pushing one another away. So now when you bring those two forces together, what are they going to do? They're going to bounce off of one another and naturally this is going to go back to being 120° and 120°.

[00:05:45]Yes.

[00:05:47]Make sense?

[00:05:50]>> Now let's get messy.

[00:05:56]Everyone knows now the electron repulsion theory, right?

[00:06:00]Electron parapulion theory. Okay, very good.

[00:06:07]That's a sound you like to hear.

[00:06:16]Means I'm doing my job.

[00:06:20]I'm going to presume I will make the presumption that you know how to draw a Louis dot diagram.

[00:06:28]We know that if you don't watch a video about it. I don't think I've like gone over it crazy before because it's a general chemistry concept. I also don't love teaching it. I'm going to presume that you know how to draw one. If you don't, go learn. Okay, let's continue.

[00:06:44]Let us take something with four electron domains. Actually before we do that let's very quickly talk about hybridization.

[00:07:01]Hybridization is what? It is the mixing of electron orbitals into hybrid orbitals when bonding I guess when molecular bonding occurs.

[00:07:28]This is why it's important to know how to draw a loose dot diagram. If you were to draw a Louis dot diagram of something like carbon tetraflloride versus carbon dioxide versus methanol versus acetic acid.

[00:08:02]Sorry, that's not versus formic acid.

[00:08:07]You have to know where all the electrons go on all these is why drawing le dot diagrams is important. And you also have to know how these electron orbitals are going to look in each of these things. So you need to know the hybridization of this guy and that guy. What the and that guy and that guy and that guy.

[00:08:33]Hybridization is counted off by what?

[00:08:36]Electron domains.

[00:08:43]What counts as one? All of these following things count as one electron domain.

[00:08:49]Any type of bond and a lone pair.

[00:08:58]That's what counts. Any type of bond and a lone pair?

[00:09:02]Yes. You have the following. Cool.

[00:09:12]To continue.

[00:09:20]And now you guys are finally going to realize why I draw my structures the way that I do by the end of this lecture.

[00:09:26]What's the hybridization here?

[00:09:30]How do we figure this out? We count the number of electron domains. I have a trick for hybridization. This is my hybridization trick. Everyone loves it.

[00:09:35]Every single time I teach this, someone's like, "Oh my god, never do that before." Okay, take how many s orbitals do we have inside of a specific energy level?

[00:09:44]One. It's one s orbital. In the p orbitals, there are three. The d orbitals there are five. F orbitals there are seven. 1 3 5 seven. Right? 1 2 3 4. Throw down those four letters.

[00:09:57]Right. How many electron domains do we have here?

[00:10:01]Four. One 2 3 four. One two three four.

[00:10:06]Yeah. All four of those letters are being taken up. That means they're all contributing. That's sp3. You have an s and three p orbitals is sp3. Let's go on to the next one. Take the carboation.

[00:10:16]The humble carboation. Not very humble actually very very arrogant.

[00:10:24]How many electron domains?

[00:10:28]Three. Yeah.

[00:10:31]One, two, three. How many of these are being used? Two p orbitals and s orbital. Right. That means this is sp what? Two. The two p orbitals go up here. What is this guy?

[00:10:44]It is the free empty p orbital that runs perpendicular to the carboation. We draw it like this.

[00:10:53]Make sense?

[00:10:56]What about this guy?

[00:11:03]How many electron domains?

[00:11:07]Huh?

[00:11:08]Two.

[00:11:10]SP. Yeah. And those extra two are the p orbitals that are involved in the triple bond. That's how we get the two pie bonds of the triple bond. Make sense?

[00:11:24]Let's keep this in mind.

[00:11:45]Uh meme's not strong enough.

[00:11:53]We're going to need it.

[00:12:00]VPR theory basically states at a specific hybridization and for a specific structure.

[00:12:12]What predictions can I make about the shape of this thing? Right? Because I know how many electron domains there are. I know where the electron domains are. I know the way that they're going to repel. And I know that they're going to repel to give each other the most space possible. Let's start with the simplest one. I want to take a carbon that has four electron domains. I don't care what's attached to it. I want a carbon with four electron domains.

[00:12:36]Right? These four electron domains indicate to me that this hybridization of this thing is what? SP3. It's the most simple thing that it can be. It doesn't seem simple, but it is. These four things are going to repel one another in a way that they are all as far away from one another as possible.

[00:12:57]The best way this is, listen, this is about visualizing chemistry. I'm going to leave a link to a little website that you can use. It's the FET Colorado. You guys can look it up right now. Actually, you can keep along with this lecture if you'd like on a second screen. FET Colorado.

[00:13:12]Go to their VPR molecule simulator.

[00:13:18]You can keep up with this video. You guys, one of you guys should open it up.

[00:13:21]You open it.

[00:13:24]Let me open it for you. Hold on.

[00:13:30]Want to make sure you're using the right one because I have two.

[00:13:33]Have your laptop.

[00:13:50]molecule shapes.

[00:13:58]Okay, beautiful.

[00:14:13]If you have this open, which looks like this, can't take this, which looks like this. If you have this open, I want you to follow along. Okay, you have the thing with four bonds attached to it. What's going to happen is that I want you to imagine that every single one of these is going to push everything else away from it first.

[00:14:39]Let's take the guy at the top.

[00:14:42]This guy is going to exert its forces moving downward and say, "Everyone get away from me." Right? What that's going to do is it's going to take the priority and everything's going to get pushed down like this, like that. Because it's going to push down and push down. What's going to happen next? The guy in the middle is going to push the other two away. This way and this way because now these guys are way too close to it.

[00:15:08]Correct? Let's make it even more dramatic.

[00:15:15]Now this guy in the middle is going to push these guys away as far as possible.

[00:15:20]However, what's going to happen is that is it going to be able to push them up?

[00:15:24]No. It has to push them sideways. So, what's going to end up happening is that you have one guy all the way up top.

[00:15:30]Then you have a guy on the bottom who's pushing the other two sideways along the same axis. One of them is going to end up in the front over here and one of them ends up in the back of that axis.

[00:15:42]And this is how we get the tetrahedral shape.

[00:15:47]Make sense? This guy forcing all of these guys down pushes them onto the bottom molecule. You can see it on your screen, right? If I could pull it up right here, I would.

[00:15:59]Actually, you know what? I think I can.

[00:16:04]Big brain moment.

[00:16:07]Didn't go to four years of medical school for nothing.

[00:16:22]Oh my golly G. I'm gonna do the thing.

[00:16:28]Wow, they have Rutherfords gathering here. That's crazy. They have some really crazy interactions.

[00:16:39]You can see right here, this is your shape.

[00:16:44]And if you move it around, can I can they see it?

[00:16:47]And if you move it around, let's remove the uh bond angles. You can see that this guy is pushing all three of these downwards. And if you line it up the way that I did on the board, you'll see one of them comes out, one of them goes to the back. This is a tetrahedral structure.

[00:17:03]Now, here's the trick that anyone who's teaching VPR theory doesn't tell you.

[00:17:07]Correct? If you were to take any of those, pause for one moment. Take any of these and replace it with an electron pair. What's going to happen to the shape?

[00:17:22]Huh?

[00:17:25]Nothing's going to happen. Watch. Do you see how the bond is exerting the force downward on everything? Take this.

[00:17:35]Let's say that for some central atom, some central atom X, I want to take this and make it a pair of electrons.

[00:17:47]Does that change the dynamics at all?

[00:17:49]No. No, because that pair of electrons is still pushing down with the same amount of force as it was when it was in a bond. Because the bond is just a pair of electrons.

[00:18:05]>> Technically, technically, yes, on a to a small degree, but the shape doesn't change. So now that is actually the same. So the shape is still in quote unquote tetrahedral form, but one of the things has been plucked off. So now what's going to happen is that everything on the bottom remains in that tetrahedral shape like it was before, but there's nothing on the top. So now you have a ceiling on top. This is called your parameal.

[00:18:37]Why? Because it looks like a triangle pyramid, right? So what you guys are going to do in the interactive sim is you're going to remove one of the bonds.

[00:18:45]How do you do that? You hit the little X on the single bond and then add a lone pair.

[00:18:51]And what that does is it gives you the parameal structure.

[00:18:58]So for every single one of these I'm going to write down a reference molecule. The reference molecule for tetrahedral is what?

[00:19:08]Methane.

[00:19:13]What is this?

[00:19:16]You guys have seen it plenty of times before. NH3. Very good. Someone's going to say, Yousef, how do I know that that's a Louis dot structure? I'm assuming that you know that it's going to take me too much time.

[00:19:29]Which is why whenever I draw NH3, you'll see sometimes I just put those dots right there. Right? This is your tragonal parameal.

[00:19:41]This is also known as the 4 + 0 and the 3 + 1.

[00:19:48]The first number refers to the number of bonds. The second number refers to the number of electron pairs.

[00:19:56]Next after the 3+ 1 is going to be the 2 plus two. Correct?

[00:20:02]All right.

[00:20:04]That electron pair up there, right, is now going to get invaded. It's going to get it space invaded a little bit.

[00:20:14]This one up the top, let's switch switch out number letter X again. I want you guys on your SIM, remove one of the single bonds, put another lone pair.

[00:20:24]You will see that the removed bond probably comes off from here.

[00:20:28]And what you do is that these guys actually fold in a bit closer to one another. Why does this happen? Well, it's because the lone pairs actually curve up a little bit. They take up less space. So, and you're not restricted by the bind the bonding principles and the way that bonds actually have to form.

[00:20:44]So, they flip up. And what happens that the lone pair comes up like this. You have a lone pair like this and a lone pair like that. And then you have two bonds like this. This is known as the bent formation.

[00:20:57]What is this? Water.

[00:21:09]bent 2 plus 2.

[00:21:15]Once again, remove the bond, put the lone pair, and you are able to see how that forms that molecule shape. Can they see molecule shape?

[00:21:31]Cool. What's next?

[00:21:34]So, these are your fours.

[00:21:38]Let's go over your less than fours. By the way, the reason I'm doing it like this is because you'll realize what is the hybridization of all of these?

[00:21:53]Never changed. Why didn't it change?

[00:21:55]Because every time I took away a bond, I replaced it with a pair, which is the same.

[00:22:02]Make sense?

[00:22:06]Mahome, make sense? Good.

[00:22:13]Should we use the pink marker?

[00:22:15]Let's do it.

[00:22:28]Hassan, what's next? Should I go up or should I go down?

[00:22:32]Four to three or four to five? I think five is a little complicated. No, let's do three.

[00:22:39]3 plus 0 is because now we can't do 3 plus one because that lands us in four.

[00:22:45]Now the totals are going to add up to three, right? 3 plus 0 is really the only one you're going to come across.

[00:22:51]This is something like the carboation like we spoke about before. You're going to have three things. Let's say let's say for harm's sake just just to mess around. All three of them are going to end start out like this.

[00:23:06]Number one, we want them to be flat because now you can take the three things and just move them into a flat plane all the way all the way away from one another. These guys are going to push down. The central guy's going to push down. What's going to happen is that you're going to have this and it's going to push down all the way over here. And then these guys are going to push away from one another and then you're going to get this and they're going to be in a perfect circle.

[00:23:26]That circle has what number of degrees inside of it? Total 360 and 60. Which means that the total here is going to be 360 divided by three domains which is 120 each. There's going to be 120 degrees between each and every single one of these. But you can see by doing what? So, I'm going to show you. Tell me when they can see.

[00:23:51]>> You're going to take away these lone pairs and put on one of the bonds. And now you can see that flat molecule. This is known as a tragonal planer geometry.

[00:23:58]Bond angles right there. Right? This is what I mean. Look, when I drew the carboation on the screen earlier, I said I'm going to draw it in the plane of the board with the p orbital like this. So, you can imagine here. Here is my flat tragonal planer like that. I put it flat against here. And then what I do, let me move it properly.

[00:24:21]What I do is that I draw my p orbital up and down through that molecule.

[00:24:27]Up and down through that molecule. Make sense? So what I'm showing you there is what is this?

[00:24:40]Do you guys see how? Show me one more time.

[00:24:45]How that is that that right there and then the p orbital.

[00:24:51]Why is there a p orbital? Because there's three domains. There's supposed to be four. One of them is free. The p orbital goes perpendicular to those things. Absolutely perpendicular. Once again, as far away as it can be. This is the carboation.

[00:25:04]Okay.

[00:25:11]This is known as what? Trigonal planer.

[00:25:24]I'm not going to do 2 plus one probably because you're not going to see it.

[00:25:29]But what you are going to see and what's going to get you points on the discrete portion is five and six. Why do we stop at six? Can't really go beyond that.

[00:25:40]There's no like room to go beyond that.

[00:25:43]By the way, a lot of times the MCAT will ask you for the hybridization of something and it will go sp3 or it'll go sp2 sp3 sp3d sp3d2.

[00:25:55]For the sake of the mcat, I want you to know that yes, that's possible. But the leading theory these days is that we are unsure of whether or not the d orbital actually hybridizes. We don't truly know that.

[00:26:08]Unless the theory has changed since I've been in college, which maybe it has.

[00:26:12]It's been a while.

[00:26:15]I was telling my friends about this um this sweater that I have at home. It's a Supreme sweatshirt. I got it when I was in my second year of college. And I was like, "Yeah, uh this is I got this in my sophomore year of college six years ago." And I was like, "Oh my god."

[00:26:32]All right.

[00:26:37]to continue on to five.

[00:26:41]Five is just four with an addition clearly. But let's look at what four looks like again.

[00:26:51]This is what four looks like. Correct?

[00:26:54]Where is the most likely place that you're going to put a fifth thing? What looks like the most open area here? On the bottom, right? Take your model, put four back on with no loan pairs.

[00:27:07]When you place a fifth, you'll realize that the fifth sort of sneaks its way in the bottom and pushes everything up. So now what you're doing is you're adding this guy. These three are way too close.

[00:27:17]So it pushes them away back towards the flat axis. Because what's happening is you're pushing from the top and you're pushing from the bottom. And these level out to be between the two of them. So you basically have two poles. You have a north pole and a south pole. And then you have the flat axis between them.

[00:27:37]This, if you were to cut it in half, looks like two of these.

[00:27:45]If you stack two of these on top of one another, this is what that would look like. This is why the five plus 0 shape is known as what? By parameal to show you on the screen. Here's three and four and five.

[00:28:05]Let's remove the bond angles.

[00:28:09]You will see that you have two poles.

[00:28:10]This is so annoying to get to work with me.

[00:28:14]Come on, man.

[00:28:20]Hold on.

[00:28:23]Beautiful.

[00:28:25]Can they see You will see you have a top pole, a bottom pole, and the flat axis between them. And those two, those three are separated by 120 degree angles on this flat axis. And then 180 degrees between the top and the bottom. Correct? And then between the top axis and the middle axis, between the top pole and the middle axis, there's 90 degrees. So important to mention here between A and B there are 90° between A and C there are 120 degrees and between B and D there are 180 degrees right this is important to note this is known as theal parameal what's an example anyone know any I don't know one off the top of my head let me look one PF5 make sense now.

[00:29:52]If I remove one and I add a lone pair, the lone pair basically takes the it takes the place of one of these things on the axis.

[00:30:01]Remove this. Add a pair. The repulsion doesn't change. Yes. So what that does is now you have two things that are 120 degrees separated and one thing two things that between them that are 180.

[00:30:13]This creates a shape known as a seessaw.

[00:30:18]because it looks like it's standing. If you were to draw it this way, that looks like a seessaw like that.

[00:30:28]This is literally called a seesaw shape.

[00:30:30]This is a 4 + one. From here on, I'm just going to write them out and give you the number because the examples don't really matter.

[00:30:43]You can see the shape here.

[00:30:48]Seesaw which has that lone pair at the top and then 120° separated axes on the bottom and those 180°ree poles that we had in the 5 plus 0.

[00:31:04]What happens when you remove one more and add another lone pair?

[00:31:09]Now what can happen is that it really does start to resemble three cuz what you do is if you remove another one of these and add this everyone realizes that they can just rearrange to be on one single axis. So what happens is you get that make sense. This once again makes the trigonal planer shape because once you have three domains, everything can just be on the same axis.

[00:31:39]Technically that's five domains. Yes.

[00:31:41]But it's three bonding domains.

[00:31:44]You guys following? Make sense? I think following along with the interactive like molecules helps a lot.

[00:31:53]Let's very quickly talk about six.

[00:31:55]You're hardly ever going to see it.

[00:31:58]Not seven. No, there is no seven as far as I know when I'm not a theoretical chemist.

[00:32:14]Xenon tetraflloride is the commonly quoted example of this and the actual molecule here looks like well let's before take an example of something that actually has six bonds to its name a central atom X has two poles a pole going up and a pole going down as you begin to assort the four things in the way that they are going to be assorted two of them are going to end up on this plane to make them 90° separated from one another then one of them is going to go into the back and one of them is going to go into the front coming directly out at you and directly out into the board. These are all separated by how many degrees? 90. Every single one of them. What that's going to do is it's going to form a little pyramid up here with a square base and a pyramid down here with a square base which means it's square parameal. Well, sorry, sorry, not square ocahedral. It's like an a little ocahedron. Looks like the molecule literally looks like this.

[00:33:22]This is the ocahedral shape.

[00:33:26]6 plus O.

[00:33:32]Let's take a look at that on the lab.

[00:33:37]This one's kind of easy to understand because once you see the the actual shape of the molecule, it's not that bad, right? A square base with two poles.

[00:33:50]What happens when we do 5 + 1?

[00:33:54]Well, there's no real place for any of these to go.

[00:34:00]So, you just kind of replace this.

[00:34:04]And now you have a square base with only one pole. And that can make a pyramid.

[00:34:08]And that's going to make square parameal.

[00:34:21]You guys get that? When you take away a second one, where are you taking from?

[00:34:24]Now you understand the sort of repulsions that are taking place. Where can you take from here?

[00:34:30]Take from the bottom. take from the pole because that allows you to conserve the repulsion in the middle. So what is that? That is square planer.

[00:34:41]>> Yeah, very good. You guys are getting the hang of it.

[00:34:47]And then from here you go uh trigonal planer bent and then linear. Correct.

[00:34:54]Very good.

[00:34:57]This is SP2.

[00:35:01]This is SP3D.

[00:35:04]This is SP3D.

[00:35:08]And these are all SP3D2 technically.

[00:35:18]Any questions?

[00:35:20]Yes.

[00:35:23]T-shape. Oh, what was T-shape?

[00:35:36]Yeah, that's the that's the um thing after this one. I didn't go that far down. You hardly ever see it. You're really because here Oh, sorry. I said that that was traal uh trigonal planer.

[00:35:48]My bad. You're absolutely correct. When you take away one of these, what actually happens is that you have repulsion from both sides, the poles and the axis. So they actually just sort of do this and they form a little T. It's complicated, but also hardly hardly hardly ever going to see it. Let's look at it on the diagram right there.

[00:36:13]So you can see that if we take away three of the lone pairs, they make a little >> they make a little T- shape.

[00:36:22]Yeah, very symmetrical.

[00:36:28]That was shapes.

[00:36:33]Are we done? No, we're not.

[00:36:37]What's the practical implication of all of this? You need to know a couple of things about each of these shapes and these hybridizations.

[00:36:48]Number one, you guys ready?

[00:37:03]When I have something like a carbonal, what is the hybridization of that carbon?

[00:37:11]SP2. Correct? Let's take R and R1.

[00:37:16]So that it's different. two different things and I have a base B.

[00:37:22]This is SP2 trigonal planer 3 + 0. Correct?

[00:37:32]You guys now know what all of that means. Correct?

[00:37:35]When I add this base B, what am I adding? I'm adding a new electron domain.

[00:37:47]I'm going to draw this out explicitly now instead of the way I did before.

[00:37:55]That's that right. We took something that was sp2 trigonal planer 3 plus 0.

[00:38:00]And now that carbon is sp 3 tetrahedral 4 plus 0.

[00:38:11]Here's the thing. Whenever you go from something to tetrahedral, you have the chance of what?

[00:38:20]Kirality.

[00:38:23]Because remember, one of the definitions of kirality is that kirality can only happen when something is sp 3 or tetrahedral.

[00:38:32]Now that you see sp3 tetrahedral, kirality is possible. Which means that in this product you're going to have a receiving mixture. This is the implication of knowing the SCPR theory and hybridization that you know where the kirality comes from. Let's look at another example of when this exact thing happens in the SN1 reaction where you have the bromine which leaves and gives you this carboation. This carboation is sp 2.

[00:39:08]Why? There's only two electron domains.

[00:39:11]There's an implicit hydrogen.

[00:39:13]There's a hydrogen that never gets drawn. Yeah. Implicit hydrogen. So, sp2.

[00:39:19]What else? Huh?

[00:39:22]Not bent. Implicit hydrogen. Remember trigonal planer 3 plus zero. 3 plus one would be a carbanion.

[00:39:36]Sp2 3+ 0 any different than the carbonial?

[00:39:40]No.

[00:39:42]in terms of reactivity.

[00:39:52]Why does this make sense?

[00:39:56]Lis acids both l acids. Electron density being withdrawn from the carbon here by the oxygen creates a partial positive which mimics the positive density here in base conditions and in acid conditions. These function similarly this is highle chemistry guys. This is really really high level MCAT stuff.

[00:40:17]This is the way that a very analytical thinker is thinking of the MCAT.

[00:40:26]Cool. SP2 planer 3 plus 0 gets attacked by solvent which turns it into on that sp3 tetrahedral yeah sp3 tetrahedral coming from sp2 formation of sp3 means kyality that's why the one makes a recemic mixture we made sp3 from something less than it. Make sense?

[00:41:00]>> Yes. Very good.

[00:41:03]You will often hear me refer to something in carbonial chemistry as the TCAI. If you go back in my videos, specifically chapter, believe seven or eight of organic chemistry, you'll hear me talk about the TCAI, that when you have something like a carbonal and you protetonate the carbonal, which you guys have heard me say a hundred times, and then you take a weak base and you attack this thing, you can get something like this.

[00:41:44]You had something like this. Correct. I said that this is known as the TCAI.

[00:41:52]TCAI stands for tetrahedral carbonial addition intermediate because you added something to the carbonal and got a tetrahedral structure. And the reason we call it tetrahedral is because we want to make sure that we add in our kyality.

[00:42:14]Make sense?

[00:42:17]What's the practical implication of this? This is all great theoretical organic chemistry, right?

[00:42:23]Where can someone point out to me a place in the body that this is relevant?

[00:42:38]Where do you where in the body do you have the attack of a carbonal and turning it into something tetrahedral? This is a little complicated.

[00:42:57]This guy attacks there. And what that does, come on on screen.

[00:43:09]Remember, it creates chirality right here. Why does it create chirality?

[00:43:12]Because that guy went from being sp2 to being sp three. Connect one, two, three.

[00:43:17]And then the implicit hydrogen is number four. The anomeic carbon is chyro. You got a recemic mixture of both alpha and beta animer because of the fact that you're attacking an sp2 carbon making sp3. This is the tetra before it gets proteinated. That is the tetrahedral or before it gets deproinated. The tetrahedral carbonial addition intermediate. Make sense? Very good.

[00:43:40]This is very high level chemistry at least for our our sake.

[00:43:49]Sometimes they're kind of rude.

[00:43:54]What's a coordinate coalent bond?

[00:44:01]>> Oops. What's a coordinate coalent bond?

[00:44:12]A coalent bond normally has a shared electron pair between the two of them where they sort of both introduce things into the bond. A coordinate coalent bond is where one atom donates a pair something like that that Oh, sorry. Yeah. Yeah. Yeah. That because that came from that.

[00:44:53]Yeah, this is iron held in the hemoglobin molecule suspended inside of hemoglobin.

[00:45:10]Is that not the same thing?

[00:45:14]right?

[00:45:17]Where the oxygen donates its electrons into the iron. What is it donating into?

[00:45:21]It's donating into the d orbital of the iron. There's a d orbital that floats around the iron. It's donating into that. They want you to be able to know maybe they don't want you to know the exact geometry, but if you were to see that the iron was suspended by these four amino acids around it, you would need to know that the hybridization is greater than sp3. And speed and not hybridization specifically for coalent stuff but the electron geometry that you have electron geometry around here and then something donates into the d orbital. So the d orbital has play it is at play there.

[00:45:52]Make sense?

[00:45:53]Very good.

[00:45:55]Any questions?

[00:46:02]One quick note on p orbitals.

[00:46:11]The p orbital of a carboation is free to be attacked. Feel free please. Yeah.

[00:46:20]Yeah. Go for it.

[00:46:23]The p orbital of the carboation is free to be attacked because it is open and it moves up and down across here. The other time where p orbitals are used are inside of double bonds. Correct? Just this doesn't really have to do with vscpr theory as much as it does about p orbital dynamics.

[00:46:38]When something forms a double bond, what's actually happening is that of course these carbons are sp2, correct?

[00:46:49]We know this. So what's happening is that the carbon has three hybrid orbitals going out like that. And then this carbon has three hybrid orbitals going out like this. And the hydrogen has its s orbital that's conjugating or bonding into the p orbitals like that. And now we have this p orbital that goes up and down like this and this p orbital that goes up and down like this because these are both sp two which means that they have free p orbitals. They are using those p orbitals to do what? Are those p orbitals empty? No, they're not empty.

[00:47:23]Those p orbitals are full of electrons.

[00:47:25]So what's happening? The p orbitals are binding over the linearity of the molecule. So this p orbital is actually throwing itself this way and this one is throwing itself that way and these are throwing themselves this way.

[00:47:44]So that creates a bonding and an anti-bonding domain of the p orbital.

[00:47:47]Don't ask me what that means. I actually have no clue. Right? And that arches above and below the actual molecule. So this is what we mean when we say this.

[00:47:59]When we say this, we actually mean that.

[00:48:01]So this right here, this right here is known as the what?

[00:48:08]The sigma bond.

[00:48:13]The sigma bond. But this guy A and B together are known as the pi bond.

[00:48:23]Don't get confused. It's not both of them. These are not two separate pi bonds. It's one pi bond. Now here's the question.

[00:48:35]Yeah, Rob, what happens now?

[00:48:39]Very good. Let's take a look.

[00:48:44]I like how Khadijah just said they go they go like this. Let's take a look.

[00:48:51]Carbon carbon, right? And now these are sp, which means they're flat, correct?

[00:48:58]So, we're going to throw a orbital this way. The sp hybridized orbitals and the sp hybridized orbitals this way. The hydrogens's are going to bond into the flat domain of the sp orbital. And now we have a pair of p orbitals that go like this.

[00:49:13]And I don't mind messing up my markers because they're mine and not anyone else's.

[00:49:18]And then another set of p orbitals are going to go, if that's up and down, this is going to go in and out of the board.

[00:49:23]Correct? So now our next set of poras is going to go like that.

[00:49:30]Yeah. Watch this.

[00:49:34]These guys bond over the top and the bottom.

[00:49:45]And then those guys bond over the front and the back.

[00:49:52]So these guys bond over the back and these guys bond over the front.

[00:50:08]That's what we mean when we say that.

[00:50:10]Here's the thing. The MCAT will actually ask you what's more difficult to break, a sigma bond or a pi bond?

[00:50:19]Huh? sigma bond. Why? It's more locked in. It's more in line. Remember that this is the sigma bond.

[00:50:27]The pi bond is this.

[00:50:31]That's the pi bond. So, the strength of the sigma bond is greater than the pi bond. However, the strength of the triple bond is greater than the double bond is greater than the single bond.

[00:50:44]And you were like, you just told me that the sigma bond is stronger than than the pi bond. So why is the double stronger than the single? Because you're forgetting that this is a sigma bond and this is a sigma and a pi bond and this is a sigma and two pi bonds. So just because the sigma bond is stronger than the pi bond, if you add something to something else, it's going to get larger. So the energy to cut through both of these is bigger than this, even though the energy to cut through this is less than that. Make sense? You guys get that? This is almost like let's say that the power to cut through this is 10 and this one's five. Now it's 15 and now it's 20. So the power to cut through these bonds is greater and the triple bond is definitely more stable, but the sigma bond in and of itself is more stable than the pi bond. Make sense? Very good question. What are some things that lock molecules into shape?

[00:51:39]Pi bonds, number one. Number two, what else?

[00:51:43]Rings.

[00:51:45]Rings will lock molecules into shape. If you have a five membered ring, can you have a double bond inside of it? You could. It's just going to add ring strain, right? What's another thing that locks a molecule into shape? Resonance.

[00:51:57]Don't forget that. Don't forget that when you have something like this, tried desperately as it may to be tetrahedral, this will always be tragonal planer, right? Why? because it's resonating its electrons and resonance can only happen when what? This is a very very defining question.

[00:52:24]Resonance can only happen when something is sp2. And since resonance happens with that nitrogen that means it must be sp2 and everything that's sp2 is trigonal planer.

[00:52:36]You guys understand?

[00:52:46]Nice pen.

[00:52:50]>> No, I know. But like your pen settings are nice.

[00:52:53]>> I just changed it ago.

[00:53:12]It's a pretty diagram.

[00:53:17]Any questions?

[00:53:20]Guys want to go home.

[00:53:24]Thank you guys so much for watching.

[00:53:25]I'll be around for some questions after the fact. Uh if I Let's get them out of here. Please support the channel in the description below. I'll see you guys next time.