Section 13.9, 13.10

Welcome to sections 13.9 and 13.10. All right, gentle people, in this lecture, what we're going to do is we're going to cover Lewis dot structures. Now I'm going to show you the way that I approach Lewis dot structures. You might have learned things in high school, and if you feel comfortable doing that, that's fine. Now the way that I'm going to show you I know works. It's a little bit slow and tedious, but it always manages to get you the right answer. So give this a try and when you get more comfortable, you can start taking slight shortcuts or incorporate what you learned in high school. Now when we do Lewis dot structures, they're most applicable for covalent bonding. The idea behind this is we are trying to draw a representation of how covalent bonds form. What a Lewis dot structure tells you is who is connected to who. how they're bonded or how many electrons they're sharing, and the number of lone pair electrons. It doesn't give you any other information. It doesn't tell you shape. It doesn't tell you angle. None of that can be derived from Lewistock's structure. So let's go ahead and think about what's happening here. What I'm going to do is I'm going to say that molecules are going to share electrons when they form a covalent bond. In the case of hydrogen, we have two electrons being shared. When they share a pair of electrons, we are going to say that that is a covalent bond, and I'm going to represent that by a dash. Now, we can share more than two electrons, and usually these will be in multiples of two. So we can share four electrons, forming the double bond, and we can share six electrons, forming the triple bond. And we can depict... the sharing of electrons in the same way. For example, if oxygen is going to share four electrons, we can draw two horizontal lines representing a double bond. And of course a triple bond would be three horizontal lines. Now before you get too crazy using multiple bonds, double and triple bonds, you should note that there are only certain elements on the periodic table that go ahead and form double and triple bonds. One of the elements has to be carbon, nitrogen, oxygen, phosphorus, or sulfur. If you don't have one of these elements involved, then it's most likely that you're not forming a double or triple bond. Now I know a lot of science fiction likes to say that carbon is the basis of life on planet earth and silicon would be right under carbon on our periodic table and that there should be silicon-based life form outside of our universe and this is kind of a misnomer and i would say that this is very unlikely what people have to understand is that carbon likes to form double bonds and there are many compounds with it forming double and triple bonds the first stable silicon silicon bond was made in the 1970s and was explosive when it got in contact with air. So the idea that silicon-based life form parallels the carbon-based life form here on planet Earth is really far-fetched. Now when I go ahead and share electrons between two atoms, I'm not going to share every electron that atom has available to it. So there are going to be other electrons that do not partake in the sharing and forming of a bond. These are going to be called the lone pair electrons. Now, while they're not part of bonding, they are important because they are stereochemically active. And what I mean by that is they take up space around my molecule. And this is going to be important in the properties of how that molecule behaves. Now, in general, what we're going to do is follow what's called the octet rule. And that is... we are going to try to bond and place things around our atoms such that each atom receives eight electrons around it. Now there are going to be exceptions and I'm going to talk about the exceptions to the octet rule. The first exception that I want to talk about is hydrogen. Hydrogen is not going to obey the octet rule and it only needs to have two electrons around it for it to be satisfied. So how we go about forming the Lewis dot structures is we are only going to go ahead and consider the valence electrons. The valence electrons are what's going to form the covalent bond. As we see in the last lecture, the core electrons are deep inside the atom and very, very rarely will try to move away from the nucleus because they are so stable. So the valence electrons are going to determine our bonding. The way that you're going to arrange your molecule is you're going to put the most electropositive element at the center. Or what I mean is put the least electronegative atom. in the center, and then you're going to arrange all the other atoms radiating around the central atom. This is typically how you're going to arrange things. Because hydrogen only wants two electrons, they tend to be terminal atoms and are found at the ends of my molecule. Once you arrange all the atoms, then you want to distribute all the remaining electrons such that you satisfy the octet. If you have to invoke multiple bonds, do such knowing that you still have to satisfy the octet. So these are the general rules. I think the best way to learn Lewis dot structures is to actually practice with them. So what I'm going to do is I'm going to draw the Lewis dot structure to these molecules. I'm going to show you the way that I approach it so you guys can get some practice. So what I like to do when I draw Lewis dot structures is I like to make a table. The first thing I'm going to do is I'm going to put the molecule that I'm interested in. Then across the table, I'm going to count the number of valence electrons and the number of that particular atom. Going down, I'm going to write what atoms are involved in making this molecule. So for Cl2, it is made out of Cl atoms. If I look on the periodic table, the valence electrons on chlorine is seven. You guys can go ahead and count across from left to right and what you will see is that it is the seventh atom in the third row. Now the number of chlorines I have is I have two chlorines in my molecule. So what this means is I need to draw a Lewis dot structure with 14 electrons. That will be the complete picture. So the first thing I'll do is I'll put chlorine. I will bond it to another chlorine. And then what I'm going to do is I'm going to fulfill the octet. So this chlorine on the left-hand side, 1, 2, 3, 4, 5, 6 electrons, plus 7 and 8 that comes from the bond. Now let's do the chlorine on the right-hand side. 1, 2, 3, 4. 5, 6, and again, 7 and 8 come from the bond. So this is my first proposal for the structure Cl2. Now if I count the number of electrons that I used here, what I see is I have 2, 4, 6, 8, 10, 12, 14 electrons. So this picture uses 14 electrons. And that's what I had to make my picture out of. I calculated that I needed a picture with 14 electrons. So this is the correct structure of Cl2. So let's go to our next molecule. H2O is made out of hydrogen atoms. It's made out of oxygen atoms. Hydrogen has a valence of 1 because it is in the first column. Oxygen has a valence of 6 because it is the sixth element in the second row. I have two hydrogens and one oxygen. I do my multiplication out, and I add the total number of valence electrons. So I need an eight electron picture. So what I'm going to do is I'm going to put oxygen in the center, because I know hydrogens are terminal atoms. They go on the ends of my molecule. I'm going to fill the octet. However, hydrogen doesn't get a full octet. it is an exception. It needs only two electrons. Because it's part of a bond, the bond satisfies those two electrons. Now oxygen, I'm going to put two, four, and six and eight are these bonds right here. So now everything that needs its octet fulfilled has its octet fulfilled. If I count the total number of electrons I used, two, four, six, eight. This is an 8 electron picture. And remember, I calculated I needed to get an 8 electron picture, so this would be a correct Lewis dot structure for water. Now you could have drawn things a little differently. Your book might have drawn it something like this with a bend in it. But again, Lewis dot structures doesn't tell you angles of things. It doesn't tell you who is across from who. It just tells you who is connected to who, and if it's a single, double, triple bond, and the number of lone pairs. So either of these would be an acceptable Lewis dot structure. So let's go on to our next molecule, O2. Now this one is only made out of oxygen. Oxygen has a valence of 6. I have two oxygens in this molecule. This means I need a 12 electron picture. So the first thing I'm going to do is I'm going to connect my two oxygens together. I'm going to fill the octet to each oxygen. And I'm going to see if this proposal meets my criteria. I'm going to count the number of electrons I use. 2, 4, 6, 8, 10, 12, 14. This is a 14 electron picture in this first proposal. What you guys will notice is I've used too many electrons in this Lewis dot structure. My constraint... was that i had to draw a picture with 12 electrons if you use too many electrons or you have a picture that used too many electrons your picture is wrong so this proposal is incorrect now if you use too many electrons what you have to do is you have to start double bonding things so you're going to look in your molecule and you're going to ask yourself what atoms are allowed to double bond. And just to remind you, these are the atoms that are allowed to do double and triple bonds. What you will see is I have oxygen here, so I can double bond my oxygens. So I'm going to redraw my picture. I know I used too many electrons, so I'm going to double bond these two oxygens. Next, I'm going to fulfill the octet to each oxygen. Two, four. and then 6 and 8 come from the double bond. I'm going to do the same with the other oxygen, 2, 4, and again 6 and 8 from the double bond. So this has all the oxygens have its octet fulfilled. There are 8 electrons around each oxygen. Now let's count the number of electrons I used in this picture. 2, 4, 6, 8, 10, 12. This is a 12 electron picture. This is the correct Lewis dot structure of oxygen. I use the correct number of electrons and oxygen has its octet fulfilled. All right, let's do the polyatomic ion NH4 plus. So valence and number. So I have an N, I have an H, And in my table, I'm going to include the charge. So nitrogen is the fifth element across in the second row. So five valence electrons. There's only one of them. Hydrogen has one valence electron. There are four hydrogens in this molecule. Now what I have to deal with is the charge of the molecule. The charge of the molecule is going to change the number of valence electrons. If I have a positive charge, that means I've removed one electron from my system. So a positive charge is going to be one minus. If it was a negative charge, that means I'm adding electrons or I'm adding valence electrons, then it would be a positive one. Now I just have a charge of one, so I'm just going to put one for the number that's associated with my charge. So five times one. 5 electrons, 4 electrons from 4 times 1, and finally negative 1 for my positive charge. So in total, I have 8 electrons that I need to use in my picture. I'm going to put nitrogen in the center, and I'm going to radiate out my hydrogens because hydrogens are terminal. Hydrogens... are happy because they all have two electrons around it. Nitrogen wants its octet fulfilled, 2, 4, 6, 8. So I don't have to put any lone pair electrons because it has its octet fulfilled in this. Now if I go ahead and count the total number of electrons used here, what I'll see is this is an eight electron picture, which is how many electrons I had to use to generate this Lewis dot structure. The last thing that you're going to do with charged molecules is you are going to put them in a bracket and put the total charge. So in this case, I'm going to put a plus outside of these brackets that surround the NH4. All right, gentle people, I want you guys to go ahead and try this out. Show me the Lewis dot structure for O3. And once you show me that Lewis dot structure for O3, what I want you to do is look at the central oxygen atom. and tell me how many lone pairs i have all right gentle people o3 is the molecule of interest that i have so i'm going to put valence and number ozone is only made out of oxygen and we saw that this has six valence electrons there are three ozone atoms and so i have 18 electrons that i want to distribute over these three oxygen atoms. So let's go ahead and connect everything together. And I'm going to just go ahead and make sure everyone's octet is fulfilled. 2, 4, 6, and 8 from the bond. 2, 4, 6, and the other bond makes 8. And then lastly, 2, 4, 6, and the bond makes 8. So now I have my first proposal. Let's see if I used enough electrons or used too much. So I got 2, 4, 6, 8, 10, 12, 14, 16, 18, 20. This is a 20 electron picture, so I used too many electrons for my first proposal. Oxygen is allowed to double bond, so I'm going to go ahead and double bond my first oxygen to my middle oxygen now i'm going to redo my octet so this first oxygen two four six and eight from the double bond the central oxygen has three bonds so that's two four six and then one more lone pair to complete its octet the last oxygen looks like the oxygen in the above proposal two four six and eight from the bomb So now I have all my oxygens with a complete octet. Let's go ahead and count the number of electrons. 2, 4, 6, 8, 10, 12, 14, 16, 18. This is an 18 electron picture, and that's exactly how many electrons that I needed to use. So this is a correct Lewis dot structure for ozone. We'll look on our central oxygen, and we have only one lone pair on there. So to answer this question, there is only one lone pair on that central oxygen. Well, gentle people, I hope that made sense. And remember to stay safe, Chem 1A.

You might have learned things in high school, and if you feel comfortable doing that, that's fine. Now the way that I'm going to show you I know works. It's a little bit slow and tedious, but it always manages to get you the right answer.

So give this a try and when you get more comfortable, you can start taking slight shortcuts or incorporate what you learned in high school. Now when we do Lewis dot structures, they're most applicable for covalent bonding. The idea behind this is we are trying to draw a representation of how covalent bonds form. What a Lewis dot structure tells you is who is connected to who. how they're bonded or how many electrons they're sharing, and the number of lone pair electrons.

It doesn't give you any other information. It doesn't tell you shape. It doesn't tell you angle.

None of that can be derived from Lewistock's structure. So let's go ahead and think about what's happening here. What I'm going to do is I'm going to say that molecules are going to share electrons when they form a covalent bond. In the case of hydrogen, we have two electrons being shared. When they share a pair of electrons, we are going to say that that is a covalent bond, and I'm going to represent that by a dash.

Now, we can share more than two electrons, and usually these will be in multiples of two. So we can share four electrons, forming the double bond, and we can share six electrons, forming the triple bond. And we can depict... the sharing of electrons in the same way. For example, if oxygen is going to share four electrons, we can draw two horizontal lines representing a double bond.

And of course a triple bond would be three horizontal lines. Now before you get too crazy using multiple bonds, double and triple bonds, you should note that there are only certain elements on the periodic table that go ahead and form double and triple bonds. One of the elements has to be carbon, nitrogen, oxygen, phosphorus, or sulfur. If you don't have one of these elements involved, then it's most likely that you're not forming a double or triple bond. Now I know a lot of science fiction likes to say that carbon is the basis of life on planet earth and silicon would be right under carbon on our periodic table and that there should be silicon-based life form outside of our universe and this is kind of a misnomer and i would say that this is very unlikely what people have to understand is that carbon likes to form double bonds and there are many compounds with it forming double and triple bonds the first stable silicon silicon bond was made in the 1970s and was explosive when it got in contact with air.

So the idea that silicon-based life form parallels the carbon-based life form here on planet Earth is really far-fetched. Now when I go ahead and share electrons between two atoms, I'm not going to share every electron that atom has available to it. So there are going to be other electrons that do not partake in the sharing and forming of a bond. These are going to be called the lone pair electrons.

Now, while they're not part of bonding, they are important because they are stereochemically active. And what I mean by that is they take up space around my molecule. And this is going to be important in the properties of how that molecule behaves.

Now, in general, what we're going to do is follow what's called the octet rule. And that is... we are going to try to bond and place things around our atoms such that each atom receives eight electrons around it.

Now there are going to be exceptions and I'm going to talk about the exceptions to the octet rule. The first exception that I want to talk about is hydrogen. Hydrogen is not going to obey the octet rule and it only needs to have two electrons around it for it to be satisfied.

So how we go about forming the Lewis dot structures is we are only going to go ahead and consider the valence electrons. The valence electrons are what's going to form the covalent bond. As we see in the last lecture, the core electrons are deep inside the atom and very, very rarely will try to move away from the nucleus because they are so stable. So the valence electrons are going to determine our bonding. The way that you're going to arrange your molecule is you're going to put the most electropositive element at the center.

Or what I mean is put the least electronegative atom. in the center, and then you're going to arrange all the other atoms radiating around the central atom. This is typically how you're going to arrange things.

Because hydrogen only wants two electrons, they tend to be terminal atoms and are found at the ends of my molecule. Once you arrange all the atoms, then you want to distribute all the remaining electrons such that you satisfy the octet. If you have to invoke multiple bonds, do such knowing that you still have to satisfy the octet. So these are the general rules.

I think the best way to learn Lewis dot structures is to actually practice with them. So what I'm going to do is I'm going to draw the Lewis dot structure to these molecules. I'm going to show you the way that I approach it so you guys can get some practice.

So what I like to do when I draw Lewis dot structures is I like to make a table. The first thing I'm going to do is I'm going to put the molecule that I'm interested in. Then across the table, I'm going to count the number of valence electrons and the number of that particular atom. Going down, I'm going to write what atoms are involved in making this molecule. So for Cl2, it is made out of Cl atoms.

If I look on the periodic table, the valence electrons on chlorine is seven. You guys can go ahead and count across from left to right and what you will see is that it is the seventh atom in the third row. Now the number of chlorines I have is I have two chlorines in my molecule.

So what this means is I need to draw a Lewis dot structure with 14 electrons. That will be the complete picture. So the first thing I'll do is I'll put chlorine. I will bond it to another chlorine.

And then what I'm going to do is I'm going to fulfill the octet. So this chlorine on the left-hand side, 1, 2, 3, 4, 5, 6 electrons, plus 7 and 8 that comes from the bond. Now let's do the chlorine on the right-hand side.

1, 2, 3, 4. 5, 6, and again, 7 and 8 come from the bond. So this is my first proposal for the structure Cl2. Now if I count the number of electrons that I used here, what I see is I have 2, 4, 6, 8, 10, 12, 14 electrons. So this picture uses 14 electrons.

And that's what I had to make my picture out of. I calculated that I needed a picture with 14 electrons. So this is the correct structure of Cl2. So let's go to our next molecule.

H2O is made out of hydrogen atoms. It's made out of oxygen atoms. Hydrogen has a valence of 1 because it is in the first column. Oxygen has a valence of 6 because it is the sixth element in the second row. I have two hydrogens and one oxygen.

I do my multiplication out, and I add the total number of valence electrons. So I need an eight electron picture. So what I'm going to do is I'm going to put oxygen in the center, because I know hydrogens are terminal atoms.

They go on the ends of my molecule. I'm going to fill the octet. However, hydrogen doesn't get a full octet.

it is an exception. It needs only two electrons. Because it's part of a bond, the bond satisfies those two electrons. Now oxygen, I'm going to put two, four, and six and eight are these bonds right here.

So now everything that needs its octet fulfilled has its octet fulfilled. If I count the total number of electrons I used, two, four, six, eight. This is an 8 electron picture.

And remember, I calculated I needed to get an 8 electron picture, so this would be a correct Lewis dot structure for water. Now you could have drawn things a little differently. Your book might have drawn it something like this with a bend in it. But again, Lewis dot structures doesn't tell you angles of things. It doesn't tell you who is across from who.

It just tells you who is connected to who, and if it's a single, double, triple bond, and the number of lone pairs. So either of these would be an acceptable Lewis dot structure. So let's go on to our next molecule, O2.

Now this one is only made out of oxygen. Oxygen has a valence of 6. I have two oxygens in this molecule. This means I need a 12 electron picture.

So the first thing I'm going to do is I'm going to connect my two oxygens together. I'm going to fill the octet to each oxygen. And I'm going to see if this proposal meets my criteria. I'm going to count the number of electrons I use.

2, 4, 6, 8, 10, 12, 14. This is a 14 electron picture in this first proposal. What you guys will notice is I've used too many electrons in this Lewis dot structure. My constraint...

was that i had to draw a picture with 12 electrons if you use too many electrons or you have a picture that used too many electrons your picture is wrong so this proposal is incorrect now if you use too many electrons what you have to do is you have to start double bonding things so you're going to look in your molecule and you're going to ask yourself what atoms are allowed to double bond. And just to remind you, these are the atoms that are allowed to do double and triple bonds. What you will see is I have oxygen here, so I can double bond my oxygens.

So I'm going to redraw my picture. I know I used too many electrons, so I'm going to double bond these two oxygens. Next, I'm going to fulfill the octet to each oxygen.

Two, four. and then 6 and 8 come from the double bond. I'm going to do the same with the other oxygen, 2, 4, and again 6 and 8 from the double bond. So this has all the oxygens have its octet fulfilled.

There are 8 electrons around each oxygen. Now let's count the number of electrons I used in this picture. 2, 4, 6, 8, 10, 12. This is a 12 electron picture. This is the correct Lewis dot structure of oxygen. I use the correct number of electrons and oxygen has its octet fulfilled.

All right, let's do the polyatomic ion NH4 plus. So valence and number. So I have an N, I have an H, And in my table, I'm going to include the charge.

So nitrogen is the fifth element across in the second row. So five valence electrons. There's only one of them. Hydrogen has one valence electron.

There are four hydrogens in this molecule. Now what I have to deal with is the charge of the molecule. The charge of the molecule is going to change the number of valence electrons.

If I have a positive charge, that means I've removed one electron from my system. So a positive charge is going to be one minus. If it was a negative charge, that means I'm adding electrons or I'm adding valence electrons, then it would be a positive one. Now I just have a charge of one, so I'm just going to put one for the number that's associated with my charge.

So five times one. 5 electrons, 4 electrons from 4 times 1, and finally negative 1 for my positive charge. So in total, I have 8 electrons that I need to use in my picture. I'm going to put nitrogen in the center, and I'm going to radiate out my hydrogens because hydrogens are terminal.

Hydrogens... are happy because they all have two electrons around it. Nitrogen wants its octet fulfilled, 2, 4, 6, 8. So I don't have to put any lone pair electrons because it has its octet fulfilled in this. Now if I go ahead and count the total number of electrons used here, what I'll see is this is an eight electron picture, which is how many electrons I had to use to generate this Lewis dot structure.

The last thing that you're going to do with charged molecules is you are going to put them in a bracket and put the total charge. So in this case, I'm going to put a plus outside of these brackets that surround the NH4. All right, gentle people, I want you guys to go ahead and try this out. Show me the Lewis dot structure for O3.

And once you show me that Lewis dot structure for O3, what I want you to do is look at the central oxygen atom. and tell me how many lone pairs i have all right gentle people o3 is the molecule of interest that i have so i'm going to put valence and number ozone is only made out of oxygen and we saw that this has six valence electrons there are three ozone atoms and so i have 18 electrons that i want to distribute over these three oxygen atoms. So let's go ahead and connect everything together. And I'm going to just go ahead and make sure everyone's octet is fulfilled.

2, 4, 6, and 8 from the bond. 2, 4, 6, and the other bond makes 8. And then lastly, 2, 4, 6, and the bond makes 8. So now I have my first proposal. Let's see if I used enough electrons or used too much.

So I got 2, 4, 6, 8, 10, 12, 14, 16, 18, 20. This is a 20 electron picture, so I used too many electrons for my first proposal. Oxygen is allowed to double bond, so I'm going to go ahead and double bond my first oxygen to my middle oxygen now i'm going to redo my octet so this first oxygen two four six and eight from the double bond the central oxygen has three bonds so that's two four six and then one more lone pair to complete its octet the last oxygen looks like the oxygen in the above proposal two four six and eight from the bomb So now I have all my oxygens with a complete octet. Let's go ahead and count the number of electrons. 2, 4, 6, 8, 10, 12, 14, 16, 18. This is an 18 electron picture, and that's exactly how many electrons that I needed to use. So this is a correct Lewis dot structure for ozone.

We'll look on our central oxygen, and we have only one lone pair on there. So to answer this question, there is only one lone pair on that central oxygen. Well, gentle people, I hope that made sense.

And remember to stay safe, Chem 1A.

Transcript for:Section 13.9, 13.10

Transcript for:
Section 13.9, 13.10