Mastering Lewis Structures and Valence Electrons

I'll be helping another YouTuber, Melissa Lucy, prepare for her next exam. And you know, I've even created a free study plan that you can download using the link in the description. That study plan has several videos or other resources that I personally recommend for you. Alright, let's learn how to draw Lewis structures properly. So today, we're basically going to be focusing on how to draw Lewis structures. I know that there's like, there's like so many questions that you're thrown. with like Lewis structures where they expect you to know one, how to draw it. And then, you know, the geometries, resonance, all of these other words that we're going to learn later on. But for now, let's at least get a good foundation of just understanding how to actually draw the structure, because that's like, I'd say that's the most important part, because if you don't have that right, then everything else is wrong. So no pressure, but let's focus on just understanding this. So we've talked about this before, looking at valence electrons, and I know I went over this a little bit. where I was just talking about one, what is an valence electron? It's the electron that's on the outermost shell of an atom. And really what I want you to memorize is everything that is highlighted here. So just knowing that this group only has one valence electron. The next one has two. We're not going to look at anything to do with, you know, transition metals for Lewis structures. And then the next one would just be, you know, this column is going to be three valence electrons, four, five, six, seven, eight, you get the point. But it's going to be really important to actually have these memorized. So some of the most common atoms that we're actually going to see that we're going to use, I'll just highlight them real quick. So definitely hydrogen. That's going to be something that's really known. We're going to see that so much. Carbon is one of the most used atoms, I would say, that we're going to use for Lewis structures. Same goes with oxygen. Same goes with, let's say, nitrogen. And then we probably are going to see some fluorine. chlorine, sulfur, and pretty much a lot of just the non-metals. Because the whole concept here with Lewis structures is the ones that we typically see are covalent Lewis structures, meaning they share electrons and it's all non-metals. All right, so if we remember that concept like going from ionic compounds to covalent compounds, ionic compounds consist of one metal and one non-metal, and then they're actually transferring electrons versus ion, actually versus covalent compounds. which are sharing electrons or, and they also consist of just non-metals. Okay, so I think we're okay with that. Let's just start off with all the different steps of how to draw a Lewis structure. So first things first, we're going to have to know every single valence electron. So like carbon has how many valence electrons and then same with hydrogen. And then we're going to add all of that together. So kind of going back. from memory, or you can look at your periodic table. How many valence electrons does carbon have? Carbon has four. Cool. I'm going to write it up here. So carbon has four. How many valence electrons does hydrogen have? Four. Perfect. And then we would multiply this by four because this subscript is telling us that there are four hydrogens in the structure. So that would give us one times four, so four. Then the next thing is just to add everything together. So I would add just four plus four to give us eight valence electrons. So first step is done. All I have to do is calculate the total number of valence electrons. We said it was eight. The next thing is identifying which one is going to be the central atom. So for the central atom, I know this one's kind of like... obvious because there's four hydrogens, so that has to surround the carbon. But just to note, carbon will typically always be, or actually it will always be the central atom if it is in the structure. Okay, so carbon, number one, that's always going to be in the center. Now, if there wasn't carbon in the Lewis structure, then it's going to be the least electronegative. So the least electronegative atom would be your central atom if carbon isn't there. Okay. So it doesn't apply for this case because hydrogen is never, and also that's actually something to note, hydrogen will never be your central atom because it could only have one valence electron, meaning it's only going to be able to create one bond. Is this making sense so far? I know I'm like throwing a lot at you. Yeah, but when you say least electronegative, like if we were comparing two and we're following that trend, then the one opposite of the trend, is that what you're saying when you say least? Yes, correct. Correct. Like if I were to have something like, I don't know, like this, and then you were asked like, oh, what's the central atom? And I know they're not going to be giving this exactly, but we'll do an example. But from there, you would see, okay, which one is the least electronegative compared to chlorine and fluorine? Well, that would be chlorine. So chlorine would be your central atom. Does that make sense? Oh, yes, because it goes up and to the right. Yes, right. Okay, right. Got it. But we'll see an example of that. For now, I just want you to kind of start to understand the steps. So choosing our central atom, that's going to be carbon. So okay, I'll put that right in the middle. Next, it's going to be surrounded by four different hydrogens. So I'll put a hydrogen, hydrogen and a hydrogen. So that's done. Step three, I do this a little bit differently. So At first, you're going to learn like the Lewis dot diagram where you'd see like, oh, carbon, we said has four balanced electrons. So you're going to draw it this way if it's just on its own. And then hydrogen only has one. So it could only make one bond. So these would combine together to then form one bond. So it would look something like. Like that. And then you'd still kind of see that whole approach here. Like if I were to actually put, to show you this, that's Cartman's Lewis dot diagram. I would do the same thing. I'm going to do this in blue for hydrogen. So just one valence electron, one valence electron, same thing here, same thing here. Essentially what's happening is they're going to form a bond because two electrons form a single bond. So instead of doing that each time, because it's time consuming, What I typically would just do is simply put a single bond already there, because this is telling us that they're going to bond regardless. Okay. But does that make sense? What I'm explaining there? Yes. Like the line? Yes. Yeah. So, and it's just because like, we're basically connecting the dots here. So it's going to form a single bond either way. So it kind of just saves you the step of drawing the Lewis dot diagram and then, you know, connecting everything. I kind of like want you to understand that they're going to form a bond regardless. So that would be the next step. And then from there, the very last thing, part of step three, is to just make sure that all of the electrons add up to what we did in step one. So in step one, we said that there are eight total valence electrons, and we want to double check that that's true. So if I were to, I'm going to draw this in red. If I were to kind of recount these electrons, I would say, well, remember, One bond consists of two electrons. So there's two so far. So two, four, six, and eight. So yes, that checks out. And because that's correct, because that checks out, we're done. That is the correct Lewis structure. So those are, I'd say, like the main three steps for a typical Lewis structure. That's what we would apply. Questions here? No, that makes sense. Perfect. Moving on to the octet roll. So. We just drew our structure, and if we saw, there's actually, there was eight valence electrons. That's what we said. So this actually follows the octet rule. And what the octet rule is saying is that the central atom prefers to have eight valence electrons surrounding it. Like it's stable, it's happy. And we always want to make that central atom happy. So we can look at this again, and I'm just looking at the central atom, so carbon. And I'm seeing that there's two valence electrons here. So. two, four, again, six, and eight. Because there are eight valence electrons surrounding that carbon, then I know, yep, this follows the octet rule. It's stable. It's happy. Once again, the structure is correct. So I'm going to keep applying this, okay? I'm going to keep applying one, the octet rule, and the steps that we went through to draw Lewis structure. So once again, let's just start with the first part where we're calculating the total number of valence electrons. How many valence electrons does phosphorus have? Five. Yes. What about chlorine? Seven. Good. And then we'd multiply by three because there are three chlorines. So 21, we'd add these together. So five plus 21 is 26 electrons. So check. Next, we have to choose our central atom. What would be our central atom in this case? Phosphorus? Yes. And then another way too, when I said like, oh, the least electronegative atom is going to be your central atom, that will help you a lot when you have like larger structures. So like when there's typically like more than one central atom, which we'll go over, but that's something that we're going to see for this case. We could have just seen that phosphorus is by itself. And the fact that there is a three subscript that tells us that. there's going to be three chlorine surrounding the phosphorus. Oh, yeah. So you can just look at it now. Yeah. Right. All right. Okay. That makes sense. Our essential is phosphorus. And then next, I'm just going to place a bond. So I said that there's three chlorine surrounding phosphorus. And then I'm already going to say, hey, this is going to add up. So there's a bond here, bond here. Now, I also want to talk about the different lone pairs. So lone pairs just mean that they are, these are lone pairs. They're electrons that are not like on a bond. So they're still on that atom. But they're not actually within a bond. So let me show you this again. Just in the very beginning, I will draw the actual like Lewis dot structure. So for chlorine, we said there are seven balanced electrons. So I'll say there's one, two, three, four, five, six, and seven. And I would do the same thing for every single one. Now, if we realize. There's always going to be one here that's being shared with whatever element it's bonded to. And then there's going to be two electrons on every single side, essentially. So there's going to be six electrons or six. Yeah, six electrons, three lone pairs, if that makes sense. Like these are lone pairs. These are the three lone pairs on the chlorine. Is this making sense so far? Okay, cool. And I would basically do this all throughout for every single one, just showing you that there are seven valence electrons on that chlorine. And then now I would look and see, okay, well, is, does phosphorus, does phosphorus have five valence electron? Well, no, it doesn't because here's one, two, three. Now I need a lone pair, five. So I could have seen it that way, or I could have done what I prefer to do, which is kind of like the faster way. So the faster way of doing this, and I'll show you again, was just to simply start with drawing the bond. Then from there I see, well, actually from there I'm going to actually draw the lone pairs on chlorine. And then something to note, every single, basically any single like halogen, same with our... like our oxygen, sulfur, and so on, those are going to also have like lone pairs like this. You're going to see a pattern. So typically with, I'm pretty much just going to say for our halogens, you're going to constantly see a pattern with this where there's going to be three different lone pairs on that atom. So that's why I already knew that, okay, I just have to put all the lone pairs in, put a bond, and I'm almost done. The last step is then to know that, well, one, this has to follow the octet rule and it doesn't right now because there's only... 1, 2, 3, 4, 5, 6 valence electrons on our central atom. It wants to have 8. So I add 2 more on phosphorus. Another thing I could have done to also check this, which is the end part, is seeing if all of the electrons add up. So I can really, I can just count this again. So I can say 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12. 13, 14, 15, 16, 17, 18, 19, 20, 21, 22, 23, 24, 25, 26. Okay, I could do that if I truly wanted to. Or I could also do this, go a little bit faster here and just saying that this is eight valence electrons. There's three here. So three times eight gives us 24 plus the additional two here. I know there's so many things that I'm throwing, but this is essentially the concept of of drawing Lewis structures. Yeah, no, that makes sense. Perfect. Okay. You get the point. And it looks something like this. I just want to leave it. Cool. Moving on. All right. So we saw that we can put lone pairs on the central atom. And then now we're going to go into just the different types of bonds. So just recall that. When two electrons combine, that forms a single bond. When four electrons combine, that forms a double bond. And then when six electrons combine, that forms a triple bond. So I'm going to do the same sort of concept here again. Let's just identify the total number of valence electrons. So I'll start with carbon. How many are there? There are four. Good. What about oxygen? Six. Perfect. And I'll multiply this by two. Giving us 12. So 4 plus 12 gives us 16 electrons. And then what's our central atom? Carbon. Yes. And that's surrounded by two oxygens. And then again, I'm just going to place the bonds already there. So I did all this so far. Now I'm just placing a bond in a bond. And then I also know that oxygen is going to have these. electrons surrounding it. Okay. These long pairs. So if we notice carbon isn't happy, okay, because it doesn't have a full octet. So it doesn't have eight valence electrons surrounding it. So far there's only one, two, three, and four. So what we have to do is we actually have to form a larger bond or a double bond. So how we do that is by actually moving. electrons to then give us that double bond. So we're not adding an additional amount of electrons because this has to give us 16. And if we notice this already has 16. So like 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16. Yep. I can't add any more electrons to that central atom. So instead I have to form a double bond. And when I do this, I'm essentially taking these electrons and moving it to form that double bond. So there's no longer three lone pairs. There's now just two. Okay. And then this is something else to note is that if this were something that was just bonded to something that was single bonded to an atom, then I would have three lone pairs. If this was double bonded, I would just have two lone pairs. If this was triple bonded, I would just have one lone pair. Okay. So definitely be aware of that just because I have seen this where you add too many electrons or too many lone pairs and you do get the answer wrong, unfortunately. Okay. So, okay, perfect. And then we can check again and just see, yes, this does obey the octet rule because here's one, two, three, four, five, six, seven, and eight. That's happy. And then we didn't add any additional electrons. It's still 16. All right, this is another one that's a little bit larger, I'd say. So for this one, we're going to have multiple central atoms. And how I would approach this is just by saying, well, there are several hydrogens, right? There's four hydrogens in this structure. But our central atom could never be hydrogen since that could only form one bond, since it only has one valence electron. So our central atoms can only be carbon and oxygen. But before I do that, I have to figure out how many valence electrons there are in total. So carbon has how many valence electrons? Four. Good. What about oxygen? Six. Perfect. And then we have, and then hydrogen we know has one. There's Don't forget about actually counting them all. So this is three plus an additional one at the end. So there are four hydrogens. One times four is four. We'll add everything up. Four plus six plus four gives us 14 electrons. And then we said that our carbon and our oxygen were our central atoms. So I'm going to draw it this way. Now, if you were curious and you weren't sure, how do I draw this? How do I know where the hydrogens go? Like what does it bond to? Well this is typically written in the form that it wants you to draw the structure. So this already is telling us that there are three hydrogens on this carbon. So I would surround this as follows and then same goes for the next one. This is already telling me oh that OH that H is going to be together. So that's how I would put a bond for everything and I'm not done yet. So then I would then look at the octet rule and see, is carbon happy? Does it have eight valence electrons on the central atom? And I would see that, yes, it does, because here's one, two, three, four, five, six, seven, eight. That checks out. Our other central atom is oxygen. And now that has to also have eight valence electrons following the octet rule. And it does not, because this only has one, two, three, and four. So instead, we would have a lone pair. two lone pairs actually, to give us eight valence electrons. Now this is complete. Questions? No, that makes sense. Perfect. Next one. Okay. So with this one, this one was actually on one of your slides or in-class assignments. And just a note, they did give you a hint. They did tell you that the, it was the oxygen was bonded to this carbon. So just a note, because I know I just said for the last one. that they typically will give you, like this is already written out the way you should bond it to. Like these hydrogens go with this carbon, this hydrogen goes with this oxygen. This one actually had to tell you that this oxygen goes with this carbon. Yeah, that's not just like, you know, you're magically going to know. No, they would have to explain that. So once again, let's just calculate how many valence electrons we have. So we know, I know you know this by now, but you're going to see this a lot. These are probably the most common atoms that we're going to see are carbon, hydrogen, and oxygen. Carbon we said has four. There are two carbons in this case, so eight. And then for hydrogens here, this is, there's three here plus this additional one. So there's one valence electron times the four hydrogens, so four. And then there's only one oxygen, so six. So I would add this all together, so eight plus. Four plus six gives us 18 valence electrons. And then now looking at this, what's our central atoms? So we're going to have two, two central atoms. What do you think are the central atoms? So carbon is one of them. Good. Can we have carbon again? Yes, we can. Perfect. Yes, we can. And if you were trying to, actually, this goes back to the rule I was talking about, where if I'm trying to decide if carbon or oxygen is like the other central atom, we would say, well, which one is the least electronegative? Oxygen is the most electronegative. Carbon is the least electronegative. So I know that the other carbon also will be our central atom. Okay, perfect. And then next, I'm just going to see that these hydrogens go to that carbon. So I'll place them around that carbon. This originally told us that this oxygen will be bonded to that carbon, so I'm just going to put that on top. And then I'll put the hydrogen next to the carbon. And then once again, I'm just going to build this just by adding single bonds all throughout. Okay, and then from there, I'm just going to double check if we have one, if we have 18 valence electrons, and if this obeys the octet rule. So for carbon, this does obey the octet rule because this has two, four, six, and eight valence electrons. Next for this carbon, it does not because this only has two, four, and six valence electrons. So what I would need to do, and something to note, was just Remember, this oxygen has three lone pairs. So if carbon isn't happy, meaning it doesn't have eight valence electrons, it's not following the optet rule, it wants to have a greater bond, so it wants to have a double bond. So I would just move one of these electrons here to form that double bond. And now this is complete. So would it be wrong if you just decided to add a lone pair to the carbon itself? Yes. Yes, it would. Why? Because here. Oh, then we're adding electrons. Yes, exactly. Then we have 20 valence electrons. Exactly. So you always want to have the total number of valence electrons. That's why that's the first step. And then another thing like, yeah, that's actually like the main reason we always want to just kind of like we're moving around the electrons. Just to get the best structure. Like that's essentially what Lewis structures are. Like we're just trying to build the best, most favorable structure. Okay. Cool. And there's exceptions, of course. So, I know. But luckily these aren't that terrible. Okay, there's going to be two main exceptions to the octet rule. So we could have an incomplete octet, meaning there's going to be less than eight valence electrons on the central atom. And I'll tell you how you can just easily tell. So the ones that will have an incomplete octet are basically going to be the atoms that don't have four valence electrons to begin with. So remember, Carbon has four valence electrons. These had three, these had two. So what I'm looking at is I'm literally just looking at the groups that do not have four valence electrons. And the reason why I just highlighted boron and beryllium is because those are the most common ones. So those are typically the exceptions to the octet rule that you will see. That's pretty much it. So I literally would say those are the main ones that they always use. Um, but just knowing that they're always going to have less than four valence electrons to begin with. That's how you know it's going to be an income incomplete octet. Okay. And I'll show you an example. So how many valence electrons does a boron have? Uh, three. Good. What about flooring? Seven. Perfect. So seven, I'm going to multiply this by three. because there's three of them. So three plus 21, that gives us 24 electrons. What's our central atom? Boron. Perfect. I'm going to surround this by three fluorines, putting our single bonds already, putting the lone pairs, knowing that halogens will always have those three lone pairs. And then now I'm just going to double check that I do have 24 electrons. So how I could easily do this? It's just by seeing that this is 2, 4, 6, 8. So this is 8. This is also 8. 8 times 3 does give us 24. So that checks out. But I know that this is an incomplete octet because there's only 2, 4, and 6 valence electrons on that central atom. But it's happy. It's totally fine. So we would just leave it like that. That is the best structure. That is the correct structure. Okay. So the exception meaning that... I get what you're saying. So it's less. Okay, cool. Yes, exactly. And then the other one is the opposite way. There's more. So it's an expanded octet. There's more than eight valence electrons on the central atom. And I'll explain what that would be. So the place that it's going to actually start is going to be the third row. So this is the third row, but specifically it's going to be here. So the third row and then below. So in your book, it's probably going to say the third row and above. And though that's correct, mentally, we don't, that doesn't make sense because this is row three, right? And we're thinking, oh, if this is row three, it would be above it, right? Three and above. And that's not how you're going to think about it, right? So I like to think about it as it's row three and like physically below everything way more sense yeah because i would totally be thinking that but are they saying that because like based on the energy like the energy levels of being like yes correct exactly that's how they're basing it yeah so but i just prefer to kind of box it up and say these are the main ones just know that these are your exceptions um so i typically say yeah three and physically on the periodic table below you know um so you The main ones that I'm actually, that I always see, and I wrote that down. So the examples that you're probably always going to see are phosphorus. Let's see, all of these, hold on. And sulfur is one of them as well. These are probably the most common ones that are going to be used. Now, something that I want to make a point here. These all have the possibility or potential to form an expanded octet. It doesn't mean that every single time they will. And I'm going to highlight back to one of the ones we just did. So phosphorus has the potential to form an expanded octet if it needs to. But as we saw previously with this example, it didn't have to, right? This had already what, two, four, six, eight, this obeyed the octet rule and that was fine. However, if I needed to, if I needed to add, like if there was five chlorines, then I would, I would see, okay, this does have that possibility of having that expanded octet if it's so needed to. So it's based on whatever atoms are attached to it. Okay. Okay. We'll do some examples there versus like, let's say boron or beryllium. They're always going to be incomplete. Like it's not saying that. They have the possibility, no, those for sure are incomplete. They're not going to obey the octet rule. So here's an example for an expanded octet rule. So let's just look at how many valence electrons we have again. So for xenon, how many valence electrons does that have? It has eight. Good. What about fluorine? Seven. Perfect. Good. And then we'd add these together. So, and I, I just found the four from here. So 28 plus eight, that would give us. 36 electrons. Okay. What's our central atom? Xenon. Good. We know this is going to be surrounded by four fluorines. So I could draw it like this for now. I'll put our single bonds. And then of course, any sort of halogen is going to have three lone pairs. So just putting them all. Perfect. And then we're going to count and see, does this have 36 electrons? So this has eight, right? Two, four, six, eight. That's eight times four of them. So eight times four gives us 32. We need more. So this right now only has 32 valence electrons. We need 36. So because of that, we need to add four electrons. So that means we're going to add two lone pairs on the central atom. So that's something to note is whenever you basically put all of the electrons on the elements that are like surrounding it or on the outside, and if you don't have enough valence electrons, then you come back to the central atom and add more. Okay. And then that's how, in this case, we do have 36. We will see that this is going to be an expanded octet because this has 2, 4, 6, 8, 10, and then 12. So this actually has 12 valence electrons surrounding it, but that's completely fine because it is part of the exception right here. Okay. We're going to keep going and we're going to see Lewis structures of ions. So now we're adding a charge. So this is basically going to be the same sort of concept where, remember that cations were actually going to lose an electron, meaning we're going to subtract an electron, and then anions, the negative charge. we're actually going to add an electron. So we're going to add. So in this case, since we have that positive charge or a cation, we're actually going to have to subtract one electron at the end. So the first thing we're going to do is just identify how many valence electrons do we have. So how many are there for nitrogen? There are five for nitrogen and one for hydrogen. We're multiplying by four though, so there's four. Good. So five plus four, and then we would subtract one based on this one plus one charge. So subtract one. And then instead of having, let's see, instead of having nine, we're actually going to have eight. And then that's something to also note, you're not going to be having like an odd number of valence electrons, because that would mean that we wouldn't have a pair, like a lone pair. So if you were to see that, there's Hmm. I'd say double check. Most likely there is a charge that you're forgetting to apply. So that's why in this case, we had to have this charge to reduce that number from nine to an actual even number of eight. What is our, what's our, actually we're here. What's the central atom? Nitrogen. Good. I'm going to surround that by four hydrogens, placing a single bond all throughout. And then I'm just going to double check if this just gives us what we wanted. So let's see, this is eight. So two, four, six, eight. Yep. That checks out the last thing. So the only difference now is that we have to actually place the entire structure in brackets and place the charge on top. This is whenever it's a charged or an ion, essentially. Whenever this is an ion, you have to put the Lewis structure in brackets and put the charge on the outside. It would be marked wrong if you didn't. Okay. Do the same thing for this next one. So now we have an anion, and I know that I'm actually going to add an electron. So go ahead and let's figure out again how many valence electrons we have. So how many valence electrons for chlorine? Chlorine has seven and oxygen has six. Good. And then multiply that by four. Okay, and then I'm going to add one more electron. So 7 plus 24 plus this one. Oh, let's see. What does that give us? So this gives us eight, 24 plus eight, 32. Okay. I'm glad you're not asking me to do that. I can't do quick math. I need my calculator for that simple stuff. But, um, I'll have you do, you know, I'm going to have you do practice problems after, but that's fine. You can use your calculator to calculate it. That's, that's not a problem. So what's our essential atom in this case? Uh, chlorine. Yes. Perfect. And then we're just going to place oxygen surrounding it and placing our bonds. You know, the oxygen is going to have three lone pairs all throughout. And then we'll double check the amount of electrons. So in this case, I have again, this is 8. 8 times 4 does give us 32. That checks out. The last step is to just place this in brackets and then our charge on the outside. So, perfect. Keep practicing those steps to drawing Lewis structures that we covered. Take notes and really just give this all you got because we're kind of in the middle of the semester right now and you know what? I know you can do this. I know you don't want to repeat this class and quite frankly, you're capable of doing this and you know it too. You just really need to believe in yourself. So keep going. Try out the next questions in the next video.

I know that there's like, there's like so many questions that you're thrown. with like Lewis structures where they expect you to know one, how to draw it. And then, you know, the geometries, resonance, all of these other words that we're going to learn later on. But for now, let's at least get a good foundation of just understanding how to actually draw the structure, because that's like, I'd say that's the most important part, because if you don't have that right, then everything else is wrong.

So no pressure, but let's focus on just understanding this. So we've talked about this before, looking at valence electrons, and I know I went over this a little bit. where I was just talking about one, what is an valence electron? It's the electron that's on the outermost shell of an atom. And really what I want you to memorize is everything that is highlighted here.

So just knowing that this group only has one valence electron. The next one has two. We're not going to look at anything to do with, you know, transition metals for Lewis structures.

And then the next one would just be, you know, this column is going to be three valence electrons, four, five, six, seven, eight, you get the point. But it's going to be really important to actually have these memorized. So some of the most common atoms that we're actually going to see that we're going to use, I'll just highlight them real quick. So definitely hydrogen. That's going to be something that's really known.

We're going to see that so much. Carbon is one of the most used atoms, I would say, that we're going to use for Lewis structures. Same goes with oxygen. Same goes with, let's say, nitrogen.

And then we probably are going to see some fluorine. chlorine, sulfur, and pretty much a lot of just the non-metals. Because the whole concept here with Lewis structures is the ones that we typically see are covalent Lewis structures, meaning they share electrons and it's all non-metals.

All right, so if we remember that concept like going from ionic compounds to covalent compounds, ionic compounds consist of one metal and one non-metal, and then they're actually transferring electrons versus ion, actually versus covalent compounds. which are sharing electrons or, and they also consist of just non-metals. Okay, so I think we're okay with that. Let's just start off with all the different steps of how to draw a Lewis structure.

So first things first, we're going to have to know every single valence electron. So like carbon has how many valence electrons and then same with hydrogen. And then we're going to add all of that together. So kind of going back. from memory, or you can look at your periodic table.

How many valence electrons does carbon have? Carbon has four. Cool. I'm going to write it up here. So carbon has four.

How many valence electrons does hydrogen have? Four. Perfect.

And then we would multiply this by four because this subscript is telling us that there are four hydrogens in the structure. So that would give us one times four, so four. Then the next thing is just to add everything together.

So I would add just four plus four to give us eight valence electrons. So first step is done. All I have to do is calculate the total number of valence electrons. We said it was eight.

The next thing is identifying which one is going to be the central atom. So for the central atom, I know this one's kind of like... obvious because there's four hydrogens, so that has to surround the carbon.

But just to note, carbon will typically always be, or actually it will always be the central atom if it is in the structure. Okay, so carbon, number one, that's always going to be in the center. Now, if there wasn't carbon in the Lewis structure, then it's going to be the least electronegative.

So the least electronegative atom would be your central atom if carbon isn't there. Okay. So it doesn't apply for this case because hydrogen is never, and also that's actually something to note, hydrogen will never be your central atom because it could only have one valence electron, meaning it's only going to be able to create one bond. Is this making sense so far? I know I'm like throwing a lot at you.

Yeah, but when you say least electronegative, like if we were comparing two and we're following that trend, then the one opposite of the trend, is that what you're saying when you say least? Yes, correct. Correct.

Like if I were to have something like, I don't know, like this, and then you were asked like, oh, what's the central atom? And I know they're not going to be giving this exactly, but we'll do an example. But from there, you would see, okay, which one is the least electronegative compared to chlorine and fluorine? Well, that would be chlorine.

So chlorine would be your central atom. Does that make sense? Oh, yes, because it goes up and to the right.

Yes, right. Okay, right. Got it. But we'll see an example of that. For now, I just want you to kind of start to understand the steps.

So choosing our central atom, that's going to be carbon. So okay, I'll put that right in the middle. Next, it's going to be surrounded by four different hydrogens. So I'll put a hydrogen, hydrogen and a hydrogen. So that's done.

Step three, I do this a little bit differently. So At first, you're going to learn like the Lewis dot diagram where you'd see like, oh, carbon, we said has four balanced electrons. So you're going to draw it this way if it's just on its own.

And then hydrogen only has one. So it could only make one bond. So these would combine together to then form one bond. So it would look something like.

Like that. And then you'd still kind of see that whole approach here. Like if I were to actually put, to show you this, that's Cartman's Lewis dot diagram.

I would do the same thing. I'm going to do this in blue for hydrogen. So just one valence electron, one valence electron, same thing here, same thing here. Essentially what's happening is they're going to form a bond because two electrons form a single bond. So instead of doing that each time, because it's time consuming, What I typically would just do is simply put a single bond already there, because this is telling us that they're going to bond regardless.

Okay. But does that make sense? What I'm explaining there? Yes. Like the line?

Yes. Yeah. So, and it's just because like, we're basically connecting the dots here. So it's going to form a single bond either way.

So it kind of just saves you the step of drawing the Lewis dot diagram and then, you know, connecting everything. I kind of like want you to understand that they're going to form a bond regardless. So that would be the next step.

And then from there, the very last thing, part of step three, is to just make sure that all of the electrons add up to what we did in step one. So in step one, we said that there are eight total valence electrons, and we want to double check that that's true. So if I were to, I'm going to draw this in red.

If I were to kind of recount these electrons, I would say, well, remember, One bond consists of two electrons. So there's two so far. So two, four, six, and eight. So yes, that checks out. And because that's correct, because that checks out, we're done.

That is the correct Lewis structure. So those are, I'd say, like the main three steps for a typical Lewis structure. That's what we would apply.

Questions here? No, that makes sense. Perfect. Moving on to the octet roll.

So. We just drew our structure, and if we saw, there's actually, there was eight valence electrons. That's what we said.

So this actually follows the octet rule. And what the octet rule is saying is that the central atom prefers to have eight valence electrons surrounding it. Like it's stable, it's happy.

And we always want to make that central atom happy. So we can look at this again, and I'm just looking at the central atom, so carbon. And I'm seeing that there's two valence electrons here. So.

two, four, again, six, and eight. Because there are eight valence electrons surrounding that carbon, then I know, yep, this follows the octet rule. It's stable. It's happy.

Once again, the structure is correct. So I'm going to keep applying this, okay? I'm going to keep applying one, the octet rule, and the steps that we went through to draw Lewis structure. So once again, let's just start with the first part where we're calculating the total number of valence electrons.

How many valence electrons does phosphorus have? Five. Yes.

What about chlorine? Seven. Good.

And then we'd multiply by three because there are three chlorines. So 21, we'd add these together. So five plus 21 is 26 electrons. So check.

Next, we have to choose our central atom. What would be our central atom in this case? Phosphorus? Yes. And then another way too, when I said like, oh, the least electronegative atom is going to be your central atom, that will help you a lot when you have like larger structures.

So like when there's typically like more than one central atom, which we'll go over, but that's something that we're going to see for this case. We could have just seen that phosphorus is by itself. And the fact that there is a three subscript that tells us that.

there's going to be three chlorine surrounding the phosphorus. Oh, yeah. So you can just look at it now.

Yeah. Right. All right. Okay. That makes sense.

Our essential is phosphorus. And then next, I'm just going to place a bond. So I said that there's three chlorine surrounding phosphorus.

And then I'm already going to say, hey, this is going to add up. So there's a bond here, bond here. Now, I also want to talk about the different lone pairs. So lone pairs just mean that they are, these are lone pairs. They're electrons that are not like on a bond.

So they're still on that atom. But they're not actually within a bond. So let me show you this again. Just in the very beginning, I will draw the actual like Lewis dot structure.

So for chlorine, we said there are seven balanced electrons. So I'll say there's one, two, three, four, five, six, and seven. And I would do the same thing for every single one. Now, if we realize. There's always going to be one here that's being shared with whatever element it's bonded to.

And then there's going to be two electrons on every single side, essentially. So there's going to be six electrons or six. Yeah, six electrons, three lone pairs, if that makes sense.

Like these are lone pairs. These are the three lone pairs on the chlorine. Is this making sense so far?

Okay, cool. And I would basically do this all throughout for every single one, just showing you that there are seven valence electrons on that chlorine. And then now I would look and see, okay, well, is, does phosphorus, does phosphorus have five valence electron?

Well, no, it doesn't because here's one, two, three. Now I need a lone pair, five. So I could have seen it that way, or I could have done what I prefer to do, which is kind of like the faster way. So the faster way of doing this, and I'll show you again, was just to simply start with drawing the bond. Then from there I see, well, actually from there I'm going to actually draw the lone pairs on chlorine.

And then something to note, every single, basically any single like halogen, same with our... like our oxygen, sulfur, and so on, those are going to also have like lone pairs like this. You're going to see a pattern.

So typically with, I'm pretty much just going to say for our halogens, you're going to constantly see a pattern with this where there's going to be three different lone pairs on that atom. So that's why I already knew that, okay, I just have to put all the lone pairs in, put a bond, and I'm almost done. The last step is then to know that, well, one, this has to follow the octet rule and it doesn't right now because there's only... 1, 2, 3, 4, 5, 6 valence electrons on our central atom.

It wants to have 8. So I add 2 more on phosphorus. Another thing I could have done to also check this, which is the end part, is seeing if all of the electrons add up. So I can really, I can just count this again. So I can say 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12. 13, 14, 15, 16, 17, 18, 19, 20, 21, 22, 23, 24, 25, 26. Okay, I could do that if I truly wanted to.

Or I could also do this, go a little bit faster here and just saying that this is eight valence electrons. There's three here. So three times eight gives us 24 plus the additional two here. I know there's so many things that I'm throwing, but this is essentially the concept of of drawing Lewis structures. Yeah, no, that makes sense.

Perfect. Okay. You get the point. And it looks something like this.

I just want to leave it. Cool. Moving on.

All right. So we saw that we can put lone pairs on the central atom. And then now we're going to go into just the different types of bonds.

So just recall that. When two electrons combine, that forms a single bond. When four electrons combine, that forms a double bond.

And then when six electrons combine, that forms a triple bond. So I'm going to do the same sort of concept here again. Let's just identify the total number of valence electrons. So I'll start with carbon.

How many are there? There are four. Good. What about oxygen?

Six. Perfect. And I'll multiply this by two.

Giving us 12. So 4 plus 12 gives us 16 electrons. And then what's our central atom? Carbon.

Yes. And that's surrounded by two oxygens. And then again, I'm just going to place the bonds already there.

So I did all this so far. Now I'm just placing a bond in a bond. And then I also know that oxygen is going to have these.

electrons surrounding it. Okay. These long pairs. So if we notice carbon isn't happy, okay, because it doesn't have a full octet. So it doesn't have eight valence electrons surrounding it.

So far there's only one, two, three, and four. So what we have to do is we actually have to form a larger bond or a double bond. So how we do that is by actually moving. electrons to then give us that double bond.

So we're not adding an additional amount of electrons because this has to give us 16. And if we notice this already has 16. So like 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16. Yep. I can't add any more electrons to that central atom. So instead I have to form a double bond. And when I do this, I'm essentially taking these electrons and moving it to form that double bond.

So there's no longer three lone pairs. There's now just two. Okay.

And then this is something else to note is that if this were something that was just bonded to something that was single bonded to an atom, then I would have three lone pairs. If this was double bonded, I would just have two lone pairs. If this was triple bonded, I would just have one lone pair.

Okay. So definitely be aware of that just because I have seen this where you add too many electrons or too many lone pairs and you do get the answer wrong, unfortunately. Okay. So, okay, perfect.

And then we can check again and just see, yes, this does obey the octet rule because here's one, two, three, four, five, six, seven, and eight. That's happy. And then we didn't add any additional electrons.

It's still 16. All right, this is another one that's a little bit larger, I'd say. So for this one, we're going to have multiple central atoms. And how I would approach this is just by saying, well, there are several hydrogens, right? There's four hydrogens in this structure. But our central atom could never be hydrogen since that could only form one bond, since it only has one valence electron.

So our central atoms can only be carbon and oxygen. But before I do that, I have to figure out how many valence electrons there are in total. So carbon has how many valence electrons?

Four. Good. What about oxygen?

Six. Perfect. And then we have, and then hydrogen we know has one. There's Don't forget about actually counting them all.

So this is three plus an additional one at the end. So there are four hydrogens. One times four is four.

We'll add everything up. Four plus six plus four gives us 14 electrons. And then we said that our carbon and our oxygen were our central atoms.

So I'm going to draw it this way. Now, if you were curious and you weren't sure, how do I draw this? How do I know where the hydrogens go? Like what does it bond to? Well this is typically written in the form that it wants you to draw the structure.

So this already is telling us that there are three hydrogens on this carbon. So I would surround this as follows and then same goes for the next one. This is already telling me oh that OH that H is going to be together.

So that's how I would put a bond for everything and I'm not done yet. So then I would then look at the octet rule and see, is carbon happy? Does it have eight valence electrons on the central atom? And I would see that, yes, it does, because here's one, two, three, four, five, six, seven, eight. That checks out.

Our other central atom is oxygen. And now that has to also have eight valence electrons following the octet rule. And it does not, because this only has one, two, three, and four.

So instead, we would have a lone pair. two lone pairs actually, to give us eight valence electrons. Now this is complete.

Questions? No, that makes sense. Perfect. Next one. Okay.

So with this one, this one was actually on one of your slides or in-class assignments. And just a note, they did give you a hint. They did tell you that the, it was the oxygen was bonded to this carbon. So just a note, because I know I just said for the last one. that they typically will give you, like this is already written out the way you should bond it to.

Like these hydrogens go with this carbon, this hydrogen goes with this oxygen. This one actually had to tell you that this oxygen goes with this carbon. Yeah, that's not just like, you know, you're magically going to know.

No, they would have to explain that. So once again, let's just calculate how many valence electrons we have. So we know, I know you know this by now, but you're going to see this a lot.

These are probably the most common atoms that we're going to see are carbon, hydrogen, and oxygen. Carbon we said has four. There are two carbons in this case, so eight. And then for hydrogens here, this is, there's three here plus this additional one. So there's one valence electron times the four hydrogens, so four.

And then there's only one oxygen, so six. So I would add this all together, so eight plus. Four plus six gives us 18 valence electrons. And then now looking at this, what's our central atoms? So we're going to have two, two central atoms.

What do you think are the central atoms? So carbon is one of them. Good.

Can we have carbon again? Yes, we can. Perfect. Yes, we can. And if you were trying to, actually, this goes back to the rule I was talking about, where if I'm trying to decide if carbon or oxygen is like the other central atom, we would say, well, which one is the least electronegative?

Oxygen is the most electronegative. Carbon is the least electronegative. So I know that the other carbon also will be our central atom. Okay, perfect.

And then next, I'm just going to see that these hydrogens go to that carbon. So I'll place them around that carbon. This originally told us that this oxygen will be bonded to that carbon, so I'm just going to put that on top. And then I'll put the hydrogen next to the carbon. And then once again, I'm just going to build this just by adding single bonds all throughout.

Okay, and then from there, I'm just going to double check if we have one, if we have 18 valence electrons, and if this obeys the octet rule. So for carbon, this does obey the octet rule because this has two, four, six, and eight valence electrons. Next for this carbon, it does not because this only has two, four, and six valence electrons. So what I would need to do, and something to note, was just Remember, this oxygen has three lone pairs.

So if carbon isn't happy, meaning it doesn't have eight valence electrons, it's not following the optet rule, it wants to have a greater bond, so it wants to have a double bond. So I would just move one of these electrons here to form that double bond. And now this is complete.

So would it be wrong if you just decided to add a lone pair to the carbon itself? Yes. Yes, it would.

Why? Because here. Oh, then we're adding electrons.

Yes, exactly. Then we have 20 valence electrons. Exactly. So you always want to have the total number of valence electrons. That's why that's the first step.

And then another thing like, yeah, that's actually like the main reason we always want to just kind of like we're moving around the electrons. Just to get the best structure. Like that's essentially what Lewis structures are. Like we're just trying to build the best, most favorable structure. Okay.

Cool. And there's exceptions, of course. So, I know. But luckily these aren't that terrible. Okay, there's going to be two main exceptions to the octet rule.

So we could have an incomplete octet, meaning there's going to be less than eight valence electrons on the central atom. And I'll tell you how you can just easily tell. So the ones that will have an incomplete octet are basically going to be the atoms that don't have four valence electrons to begin with. So remember, Carbon has four valence electrons.

These had three, these had two. So what I'm looking at is I'm literally just looking at the groups that do not have four valence electrons. And the reason why I just highlighted boron and beryllium is because those are the most common ones.

So those are typically the exceptions to the octet rule that you will see. That's pretty much it. So I literally would say those are the main ones that they always use. Um, but just knowing that they're always going to have less than four valence electrons to begin with.

That's how you know it's going to be an income incomplete octet. Okay. And I'll show you an example.

So how many valence electrons does a boron have? Uh, three. Good. What about flooring? Seven.

Perfect. So seven, I'm going to multiply this by three. because there's three of them. So three plus 21, that gives us 24 electrons. What's our central atom?

Boron. Perfect. I'm going to surround this by three fluorines, putting our single bonds already, putting the lone pairs, knowing that halogens will always have those three lone pairs.

And then now I'm just going to double check that I do have 24 electrons. So how I could easily do this? It's just by seeing that this is 2, 4, 6, 8. So this is 8. This is also 8. 8 times 3 does give us 24. So that checks out. But I know that this is an incomplete octet because there's only 2, 4, and 6 valence electrons on that central atom.

But it's happy. It's totally fine. So we would just leave it like that.

That is the best structure. That is the correct structure. Okay. So the exception meaning that...

I get what you're saying. So it's less. Okay, cool. Yes, exactly. And then the other one is the opposite way.

There's more. So it's an expanded octet. There's more than eight valence electrons on the central atom. And I'll explain what that would be. So the place that it's going to actually start is going to be the third row.

So this is the third row, but specifically it's going to be here. So the third row and then below. So in your book, it's probably going to say the third row and above.

And though that's correct, mentally, we don't, that doesn't make sense because this is row three, right? And we're thinking, oh, if this is row three, it would be above it, right? Three and above. And that's not how you're going to think about it, right? So I like to think about it as it's row three and like physically below everything way more sense yeah because i would totally be thinking that but are they saying that because like based on the energy like the energy levels of being like yes correct exactly that's how they're basing it yeah so but i just prefer to kind of box it up and say these are the main ones just know that these are your exceptions um so i typically say yeah three and physically on the periodic table below you know um so you The main ones that I'm actually, that I always see, and I wrote that down.

So the examples that you're probably always going to see are phosphorus. Let's see, all of these, hold on. And sulfur is one of them as well. These are probably the most common ones that are going to be used.

Now, something that I want to make a point here. These all have the possibility or potential to form an expanded octet. It doesn't mean that every single time they will.

And I'm going to highlight back to one of the ones we just did. So phosphorus has the potential to form an expanded octet if it needs to. But as we saw previously with this example, it didn't have to, right? This had already what, two, four, six, eight, this obeyed the octet rule and that was fine.

However, if I needed to, if I needed to add, like if there was five chlorines, then I would, I would see, okay, this does have that possibility of having that expanded octet if it's so needed to. So it's based on whatever atoms are attached to it. Okay. Okay.

We'll do some examples there versus like, let's say boron or beryllium. They're always going to be incomplete. Like it's not saying that.

They have the possibility, no, those for sure are incomplete. They're not going to obey the octet rule. So here's an example for an expanded octet rule. So let's just look at how many valence electrons we have again. So for xenon, how many valence electrons does that have?

It has eight. Good. What about fluorine?

Seven. Perfect. Good.

And then we'd add these together. So, and I, I just found the four from here. So 28 plus eight, that would give us.

36 electrons. Okay. What's our central atom?

Xenon. Good. We know this is going to be surrounded by four fluorines.

So I could draw it like this for now. I'll put our single bonds. And then of course, any sort of halogen is going to have three lone pairs.

So just putting them all. Perfect. And then we're going to count and see, does this have 36 electrons? So this has eight, right?

Two, four, six, eight. That's eight times four of them. So eight times four gives us 32. We need more. So this right now only has 32 valence electrons. We need 36. So because of that, we need to add four electrons.

So that means we're going to add two lone pairs on the central atom. So that's something to note is whenever you basically put all of the electrons on the elements that are like surrounding it or on the outside, and if you don't have enough valence electrons, then you come back to the central atom and add more. Okay.

And then that's how, in this case, we do have 36. We will see that this is going to be an expanded octet because this has 2, 4, 6, 8, 10, and then 12. So this actually has 12 valence electrons surrounding it, but that's completely fine because it is part of the exception right here. Okay. We're going to keep going and we're going to see Lewis structures of ions.

So now we're adding a charge. So this is basically going to be the same sort of concept where, remember that cations were actually going to lose an electron, meaning we're going to subtract an electron, and then anions, the negative charge. we're actually going to add an electron.

So we're going to add. So in this case, since we have that positive charge or a cation, we're actually going to have to subtract one electron at the end. So the first thing we're going to do is just identify how many valence electrons do we have. So how many are there for nitrogen?

There are five for nitrogen and one for hydrogen. We're multiplying by four though, so there's four. Good.

So five plus four, and then we would subtract one based on this one plus one charge. So subtract one. And then instead of having, let's see, instead of having nine, we're actually going to have eight.

And then that's something to also note, you're not going to be having like an odd number of valence electrons, because that would mean that we wouldn't have a pair, like a lone pair. So if you were to see that, there's Hmm. I'd say double check.

Most likely there is a charge that you're forgetting to apply. So that's why in this case, we had to have this charge to reduce that number from nine to an actual even number of eight. What is our, what's our, actually we're here.

What's the central atom? Nitrogen. Good.

I'm going to surround that by four hydrogens, placing a single bond all throughout. And then I'm just going to double check if this just gives us what we wanted. So let's see, this is eight. So two, four, six, eight.

Yep. That checks out the last thing. So the only difference now is that we have to actually place the entire structure in brackets and place the charge on top.

This is whenever it's a charged or an ion, essentially. Whenever this is an ion, you have to put the Lewis structure in brackets and put the charge on the outside. It would be marked wrong if you didn't. Okay.

Do the same thing for this next one. So now we have an anion, and I know that I'm actually going to add an electron. So go ahead and let's figure out again how many valence electrons we have. So how many valence electrons for chlorine? Chlorine has seven and oxygen has six.

Good. And then multiply that by four. Okay, and then I'm going to add one more electron. So 7 plus 24 plus this one.

Oh, let's see. What does that give us? So this gives us eight, 24 plus eight, 32. Okay. I'm glad you're not asking me to do that. I can't do quick math.

I need my calculator for that simple stuff. But, um, I'll have you do, you know, I'm going to have you do practice problems after, but that's fine. You can use your calculator to calculate it.

That's, that's not a problem. So what's our essential atom in this case? Uh, chlorine. Yes.

Perfect. And then we're just going to place oxygen surrounding it and placing our bonds. You know, the oxygen is going to have three lone pairs all throughout. And then we'll double check the amount of electrons.

So in this case, I have again, this is 8. 8 times 4 does give us 32. That checks out. The last step is to just place this in brackets and then our charge on the outside. So, perfect. Keep practicing those steps to drawing Lewis structures that we covered. Take notes and really just give this all you got because we're kind of in the middle of the semester right now and you know what?

I know you can do this. I know you don't want to repeat this class and quite frankly, you're capable of doing this and you know it too. You just really need to believe in yourself. So keep going.

Try out the next questions in the next video.

Transcript for:Mastering Lewis Structures and Valence Electrons

Transcript for:
Mastering Lewis Structures and Valence Electrons