If you'll remember in the previous section, we talked about molecular formulas, structural formulas, and Lewis formulas. Well, we're going to kind of take a closer look at the Lewis structures or the Lewis formulas in this section. So first of all, we need to go back and talk about valence electrons, and they're the outer shell electrons.
They're the ones that are the outer S's and P's. And they're the ones that are going to participate or are going to be involved in chemical bonding. Okay.
And if you notice, we have a periodic trend for this. Group one, and means it's the row that it's on on the periodic table. So anyone in group one, if it's the first row or for hydrogen, it's going to be 1s1.
Group two, and it also... what you'll notice is the group number, the number in the ones place, is also equal to the number of valence electrons. So we don't have to go and do all the electron configuration if we know what this periodic trend is. So if we want to draw Lewis dots, okay, we're going to use one dot to represent each valence electron. We can get that from the group number, okay, remember the number in the ones place.
of the group number and then we're going to put one dot on each side before doubling up similar to that hund's rule when we were doing the arrows where we put one in each box before we doubled up so if we have boron boron is found in group 13 so it's going to have three valence electrons And so it's only going to need three dots. And so we're going to put one dot on three different sides. And quite frankly, the three dots can go anywhere in terms of as long as you only have one on each side of the boron. What we have here is what's called the hat configuration, because it kind of looks like the boron's got a hat if we connect the dots.
But you could take the one that's on top and put it underneath the boron. But this is the classic pattern that you're going to see in chemistry. but the answer is you've got three single dots.
So let's take a look at let's take a look at oxygen and it's in group number 16 so it's going to need six dots and so if we go to put our dots around we're going to put one on each side before we double up. Sometimes you'll see the pair of dots on the left and the single on the bottom. The whole point behind this is the answer is two pairs and two singles. Those two singles are going to become important when we go to take a look at the shapes of our molecules in the next section.
So if we take a look at silicon, it's got 14, it's in group 14, so it's gonna have four valence electrons. We're gonna put four dots. And what we're going to find is we're going to put one dot on each side. And so the answer is this is four single dots. Again, that's going to influence the shape that we're going to be taking a look at in the next section.
So if you'll notice, this is a periodic trend. There's only eight. patterns and you can look right straight down the columns just got to be careful with that helium in column number 18 because helium is only big enough to hold two dots that's why it's paired up okay but if you'll notice there's only eight valence electrons there's only eight patterns And so one of the things we need to be able to do is also find the total number of valence electrons in a molecule because we can use this to verify whether we've got the right structural formula or Lewis formula. So if you'll notice, we want to know the total number of valence electrons in a molecule of carbon dioxide. Well, carbon's column number 14, oxygen's column number 16. We've got to remember to multiply two times six for the oxygen.
And so we end up with a total of 16 valence electrons, 4 plus 12. These numbers can get pretty exotic. Phosphorus is in group 15, so 2 times 5 is 10. Oxygen is in group 6, and there's 5 of them, so that's 30. And 30 plus 10 is 40. So these numbers can get pretty big, and they can get pretty exotic. But we can use these to verify whether we have a correct formula or not.
I'll show that to you in the next couple slides. So when writing Lewis structures, What we find is we draw a skeletal structure of a compound showing what atoms are bonded to each other. We put the least electronegative element in the center, so we're going to put the one that's furthest from fluorine in the center.
We're going to count the total number of valence electrons, and then for ions, we're going to add one for the negative and subtract one for the positive. See, we were just adding up valence electrons on the two previous slides. We're going to make sure they all have eight.
except for hydrogen. It only needs two. Technically, helium only needs two, but it doesn't react with anything. So it's not going to form covalent bonds. And so if our structure contains too many electrons, we form double and triple bonds on the central atom as needed.
So let's go through an example of this. So now what we're going to do is we're going to take a look at how right. the Lewis structure for nitrogen trifluoride. So the first thing we're going to do is we're going to put the least electronegative in the center.
So that's going to be the nitrogen. And I've already drawn some of this on your slide notes for you. Step number two, we're going to count the number of valence electrons.
And we are looking at that on the two previous slides. So nitrogen is group 15, fluorine is group 17. So we've got 5 plus 3 fluorine, so we've got to multiply 7 by 3, and we end up with a total of 26 valence electrons. We're going to draw single bonds between nitrogen and fluorine atoms and complete an octet on the nitrogens and fluorine atoms.
So I've already added the single bond at the bottom. Remember, we have to have single bonds or the molecule is going to fall apart. The other thing is, is remember we use the line because it's representing two electrons. So if you'll notice on each of the fluorines, they each have six dots plus a line.
Line represents two. So each fluorine thinks it owns both electrons from the line. If you take a look at nitrogen, it's got three lines, which is the equivalent of six.
plus those two unshared pairs. And so what we're going to do is we're going to check are the number of electrons in the structure equal to the number of valence electrons that we calculated in step number two. So single bonds are equal to two electrons and there's three of them. So we multiply that by three and then we can just count up the number of lone pairs around the outside of the fluorines. Fluorines, there's nine of them, plus the one pair around the nitrogen is 10. And so we have a total of 26 valence electrons.
These are equal to each other, which means the structure that we've drawn is the most likely structure for this molecule. And so notice we've got eight electrons around each of them. and we have 26 valence electrons available and we've only used 26 and eventually we're going to turn this into a three-dimensional shape and could actually build a ball and stick model for so going through that system that we did on the previous slide can get a little bit cumbersome so I'm going to try and simplify this because it's going to match up to the category of molecules that you're going to see on the next couple of pages in your or in section 7.6 or the next couple of pages in your slot notes. The one thing that I've come to realize is when you wrote those charges across the top of your periodic table, minus one for column, I'm sorry, plus one for column one, plus two for column two and so on.
If we take a look, nitrogen has a standard charge of 3 minus. As long as the bond order is zero, or what that means, as long as the molecule is in its lowest energy state, the nitrogens are going to have three bonds. In this class, we are only interested in atoms in their lowest energy state. If they're not, then we call it organic chemistry or we call it advanced inorganic chemistry. And if it's not in its lowest energy state, it has a tendency to occur because you have the overlap of the S's and the P's and the D's that we talked a little bit about with iron and copper in a couple examples.
from a couple sets of sections ago. Okay, we know that nitrogen has five valence electrons because it's in group number 15 and if we subtract the number of bonds what we end up with is we end up with the number of electrons or we end up with the number of lone pairs. This is going to be significant for section 7.6.
So what we have is we have three bonds and one unshared pair of electrons. The interesting thing is with ions, what we're going to do is we're going to, the charge goes with the central atom. So this says nitrogen has five valence electrons, but because of that positive charge, it's going to lose one.
So there are four bonds on them. the nitrogen and so we end up with zero unshared pairs of electrons. So this shape would be based on four bonds and no unshared pairs of electrons and that's going to match up to a category system we're going to use in the next set of slot notes.
If we take a look at the Lewis structure for CS2, carbon has to have four bonds because we didn't write it on the periodic table because carbon can be plus four or minus four depending on who it's bonded to but sulfur definitely has two bonds because it's got a minus two and so if we take a look there's only one answer on this where the sulfur has two bonds on it because it's got a minor like I said it's got a minus two and so the answer would be C So what we're going to do is we're going to use a process of elimination to figure some of these out. And so on these slides, what I've done is I have gone ahead and given you the Lewis structure so you can start to see what they look like, because then we're going to match up their molecular structure and their polarity. Because remember, we can have polar bonds, which means now we can have polar molecules. So HBr, again, remember hydrogen only has two valence electrons. So that's why it can only have two valence electrons.
That's why we didn't put any extra dots around the hydrogen. Bromine has to have eight. We've already seen the PCl3.
Actually, no, we haven't. We've seen the NF3, which is similar to the PCl3. Sorry about that. And then we have molecules that have double bonds in them. And we can have molecules that have triple bonds in them.
We can also have molecules that don't follow the octet rule. So here's two examples of molecules that don't follow the octet rule. We've got PCL5. If you count up, there's five bonds coming off of the phosphorus, which means it's sharing all of its valence electrons, which means that would not have a bond order of zero.
It actually would have a bond order of plus two. That's a different story for a different day. And if we take a look at boron trifluoride, what we're going to find is boron only has three valence electrons, so it doesn't have an unshared pair of electrons. And the reason that we need to understand this is this is going to influence the shape of a molecule.
The shape is going to influence the polarity of the molecule.