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Lewis Electron Dot Diagram 26,469 views

Teacher/Instructor Kendal Orenstein
Kendal Orenstein

Rutger's University
M.Ed., Columbia Teachers College

Kendal founded an academic coaching company in Washington D.C. and teaches in local area schools. In her spare time she loves to explore new places.

Lewis Electron Dot Diagrams are used to visually depict bonding by representing valence electrons as dots surrounding an elemental symbol. These dots can be on any of the four sides of the symbol, each side representing a different orbital (1 s orbital and 3 p orbitals).

Alright so let's talk about covalent bonding in Lewis Dot Diagram. Lewis Dot Diagrams are illustrations of how the elements in a covalent bond come together to form a new, a structure or a molecule. So there lots of ways they can share those valence electrons, they can come together within a single bond sharing only 2 valence electrons, they can come together in a double bond sharing 4 valence electrons or they can come together in a triple bond sharing 6 valence electrons between the 2 atoms. Alright there's also something called a coordinate covalent bond, typically when you think of 2 elements coming together to share valence electrons where 1 is coming from one atom the other is coming from the other atom. But sometimes there's 2 valence electrons will come from a single atom. The single atom will just give up two its electrons and share, that's called a coordinate covalent bond.

And the main thing we also want to get through is the octet rule, we want to say all atoms have a goal of having 8 electrons around it. Don't forget 8 electrons are the same thing as having electron configuration as a noble gas which is exactly when we want stable good nice. So all atoms are going to want 8 valence electrons around it, and the 2 exceptions mainly are going to be hydrogen which is just 2 tiny to have that many around it, it only wants 2 and boron which is also pretty small. That guy only wants 6 but everybody else wants 8. Alright so let's take this into action, alright when we're putting this up but we have an example of NF3 or nitrogen trifluoride we're going to put the least electronegativeatom in the center. Because that guy is going to be sharing a lot of its electrons and we know electro negativity, something that's highly electronegative is going to want to hog all these electrons so it doesn't want to share. So because fluorine is highly electronegative we're going to put nitrogen in the middle, it's going to be our essential atom and fluorine is going to be surrounding the nitrogen.

Okay the next thing we're going to do is we're going to say okay how many electrons are we working with in this particular compound? Well nitrogen is in group 5 so it has 5, fluorine is in group 7 and there's 3 of them so it's 7 times 3 which is 21 so we have 26 electrons in this whole, that we're working with in this whole molecule. So we know they're are bonded together on some way and we're going to denote that by the lines connecting them. In each line we're going to say there's 2 electrons, so we just use 2, 4, 6 electrons those are already used up. So now we have 20 electrons to work with. Okay but the most electronegative atom is going to want to be the guy who hogs all the electrons around it. So that's fluorine, so we're going to start giving it, giving up the electrons of fluorine so we're going to say okay we have 20 of them 2, 4, 6, 8, 10, 12, 14, 16, 18 so we the 18 and these guys have 8 around it 2, 4, 6 and 2 are shared here that's 8 so that's good 2, 4, 6, 2 are shared here that's 8 that's perfect 2, 4, 6, 8 that's perfect.

We have 2 electrons left nitrogen 2, 4, 6 because don't forget they're either shared between them and the 2 left over making everybody has 8 around it fantastic, awesome. So this is what nitrogen trifluoride looks like in a Lewis Dot Diagram. Let's look at carbon disulphide, okay the most electronegative atom in this case is sulphur so we're going to make carbon a central item and have sulphur surrounding it. Let's figure out how many valence electrons we're working with here. We have 4 from carbon because it's in group 4 plus 6 from sulphur because it's in group 6 times 2 is 12. 4 plus 12 is 16, so 16 electrons in this whole thing.

Alright so they're going to be bonded somehow like this side, okay so we just used up 4 electrons 2 from here and 2 from here so we now have 12 electrons left. Okay so the most electronegative atoms are going to get these electrons. So it give me a 12, so we're going to say they're going to go on sulphur 2, 4, 6, 8, 10, 12 fantastic. This guy has 8 around it 2, 4, 6, 8 from this it's happy, this guy has 2, 4, 6, 8 around it's happy but carbon, poor carbon only has 4 and it needs 2 more pairs. So what's going to happen? Well sulphur is going to be nice enough to actually give up its electrons and say "don't worry carbon I'll take care of you I will bond with you." So now we have what we call a double bond. Carbon 2, 4, 6, 8 and everybody is happy okay good awesome, we're good here the third use of double bonds and notice that when sulphur gave up its electrons that's the coordinate covalent bond.

Let's look at a polyatomic ion cyanide, polyatomic ions are covalently bound together so they also have shapes. So in this case there's only 2 so it doesn't matter which one is a central atom or which one is more electronegative. Carbon has 4 electrons so can I say 4 because it's in group 4, nitrogen is in group 5 it has 5 electrons plus we have an extra one that's given here. So we're going to have plus that extra one, so we have 10 electrons total. Alright that were connected somehow so we use 2 so we have 8 left and they're going to go on the nitrogen first 2, 4, 6, and we have 2 left over 8 but again that shouldn't have it carbon is not, what's going to happen? They're going to coordinate a covalent bond here carbon now has 6 nitrogen is going to say okay I'm going to give it up and here's an example of a triple bond. This is a maximum amount of electrons that can be shared between 2 atoms.

And we also know this is a polyatomic ion, we add an extra electron in there so we're going to denote that, we're going to put it in brackets and say I added an extra electron to this structure to make it look like this. Okay so this is essentially all the different ways you can, they can combine in single, double and triple bonding. But there're also like exceptions of this so let's talk about those. There are things called "resonance structure" and a resonate structure is when one or more valid LDD can be drawn. So let's look at this and show you what I'm talking about, we're going to do it the same exact way we did the other ones, we're going to put our most electronegative atom on the outside and our least electronegative atom on the inside. So sulphur is our least electronegative atom and so it's going to go via our central atom and the oxygen is going to be surrounding it.

Okay fair enough sulphur has 6 it's in group 6, oxygen has 6 it's in group 6 there's 3 of them so it's 18. So we have, what is that 24 electrons to work with. They can be connected 2, 4, 6 I used 6 of them I have 18 left. Okay so they're going to go on oxygen because they're the most electronegative 2, 4, 6, 8, 10, 12, 14, 16, 18 fantastic all these oxygens are totally happy and they have 8 electrons around they're fine. Poor sulphur only has 6 so what are we going to do? Well we can easily say this guy is going to give it up and share, but why that oxygen, why does that oxygen have to share and not this one? Well that's the whole thing and that's it's in structures right so it can be that one or it can be that one or it can be the other one and the way to know all 3 structures we have to draw all of them and put a double arrow in both of them saying these all actually exist.

This double bond actually exist in all 3 of these places, this guy will share sometimes, this guy will sometimes and this guy will share sometimes. Electrons are constantly moving so they're able to actually share all these electrons. Resonance structures are the changing of the bonding not those structure of the actual thing. I'm not moving any elements around I'm just changing the bonding okay. One last thing is radical that's an odd number of electrons that would just have, one they'd just have an odd number so instead of everything being paired there's just going to be one left over. That's extremely highly reactive you're probably not going to come across these very often they're very, very reactive when you think of like free radicals you've probably heard that before that is something that you've heard in a that or just to create like cancer gets created by free radical and things like that so you want to destroy them they're not good.

Expanded octets are things that can have actually more than the 8, they can have anything that has above period 3 means that they have this inner shell electrons that they have available. So they can actually have more than 8 because some can have up to 12 of valence electrons PCl5 is an example it can have the phosphorus one will actually have 10 if we were to draw it, it's sharing 10 electrons with the chlorine ion because phosphorus is available to do that because it actually have access to the d orbital. So that pretty much sums up with our diagrams in a nut shell.