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Chemistry: Lewis Dot Structures for Covalent Bonds


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  • Type: Video Tutorial
  • Length: 12:02
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  • Posted: 07/14/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry: Chemical Bonding: Fundamental Concepts (10 lessons, $16.83)
Chemistry: Lewis Dot Structures (2 lessons, $3.96)

This lesson was selected from a broader, comprehensive course, Chemistry, taught by Professor Harman, Professor Yee, and Professor Sammakia. This course and others are available from Thinkwell, Inc. The full course can be found at http://www.thinkwell.com/student/product/chemistry. The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more.

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

Gordon Yee is an associate professor of chemistry at Virginia Tech in Blacksburg, VA. He received his Ph.D. from Stanford University and completed postdoctoral work at DuPont. A widely published author, Professor Yee studies molecule-based magnetism.

Tarek Sammakia is a Professor of Chemistry at the University of Colorado at Boulder where he teaches organic chemistry to undergraduate and graduate students. He received his Ph.D. from Yale University and carried out postdoctoral research at Harvard University. He has received several national awards for his work in synthetic and mechanistic organic chemistry.

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So we saw how Lewis Dot Structures were useful for predicting the products of reactions between metals and nonmetals to give ionic products. But where Lewis Dot Structures is a really powerful idea is when we start talking about covalent compounds, or compounds in which the electrons are shared between the atoms.
To understand that, what we first have to do is talk a little bit more about the idea that an atom wants eight electrons. What we call that idea is the Octet Rule. So atoms want eight valence electrons in their outermost shell, their valence shell. So if you think about oxygen, oxygen has two 2s electrons and four 2p electrons, and it wants two more electrons to fill out the 2p shell so that it has a total of eight electrons. What's special about eight electrons? Then it has a Noble Gas configuration, and then it turns out to be special stability associated with that.
Now, there's one exception to the Octet Rule that we want to talk about now, and we're going to actually number them because there are several and I want you to keep these in mind. The first one is that hydrogen and helium can only have two electrons in their valence shell. And the reason is that when we're talking about the electron configuration for hydrogen and helium, we're talking about putting electrons into the 1s orbital, and the 1s orbital only takes two electrons. So we don't typically need to talk about the 2s or the 3s when we're talking about hydrogen, and that's why hydrogen and Helium will only want two electrons in their valence shell when they're incorporated into molecules.
All right, let's see how this idea works. Here we have two fluorine atoms, and each fluorine atom starts out with seven valence electrons, and fluorine wants to have eight. Now, fluorine, if it had been in a reaction with sodium, would just grab that electron away from sodium and be perfectly happy, but there isn't any sodium in the pot. All there is in the pot is more fluorine. So these two fluorines are going to have to get together and share in order to satisfy the Octet Rule. And they form this sort of structure where we put two dots in-between the two fluorines to indicate that there's a bond. This is a covalent bond. And in some sense, these electrons belong to both the left-hand side fluorine and the right-hand side fluorine, so that each one satisfies the Octet Rule.
If we have chlorine atoms and fluorine atoms, we're going to form pretty much the same structure except that the atomic symbol is chlorine instead of fluorine, but you can see that there's a similarity here. And finally, if we have hydrogen and fluorine, hydrogen only wants two electrons, so it's satisfied by this sort of structure. Fluorine gets its eight, and so this also is a perfectly acceptable structure.
Now, when we talk about "perfectly acceptable," basically what we're saying is, the Lewis idea, Lewis bonding theory, actually predicts how the atoms should put themselves together in order to form reasonably stable molecules. So in other words, you can't put fluorine atoms into a bottle at room temperature. You can put fluorine molecules into a bottle at room temperature but not the atoms, because if you had a bunch of atoms in the pot, they'd get together and they'd form molecules. And the molecules in particular would be dinuclear; in other words, two nuclei, two atoms, would get together to form a diatomic molecule.
We should talk a little bit about some other nomenclature. We talk about the electrons in the bond as "bonding pairs," so there's one bonding pair between these two fluorines, one bonding pair between the fluorine and the chlorine, one bonding pair between the hydrogen and the fluorine. And all these other dots around the outside, we call those "lone pairs," and they're going to be important later on.
So let's now look at what happens if we had oxygen atoms in our container at room temperature. So here are the Lewis Dot Structures for oxygen atoms. Remember that it doesn't matter where you put the dots, so long as you have the right number of dots, and the number of dots reflects the number of valence electrons for the neutral atom. So we have six and six. And if we put these two things together to form a single bond between the two oxygen atoms, you'll see that for the right-hand side oxygen, we've satisfied the Octet Rule. I'll circle the eight electrons that are in the valence shell for the oxygen on the right-hand side. But what about the oxygen on the left-hand side? It only has two, four, six electrons around it. And so the Octet Rule is violated, and that's not good--meaning that the theory predicts that this is not the right structure for the oxygen molecule. What can we do? Well, what we can do is we can take a lone pair from the oxygen on the right-hand side, the one that's octet is already satisfied, and turn it into a bonding pair. So now there are four electrons in-between the two oxygen atoms, and we say that there's a "double bond" between the two oxygens. Now the Octet Rule is satisfied for the oxygen on the left-hand side, and similarly the Octet Rule is satisfied for the oxygen on the right-hand side.
What have we got? What we've got is a prediction that there should be something different about the oxygen-oxygen bond in O[2] that's different in some sense from fluorine, F[2], which had a single bond. And we're going to talk about what that is, but for right now let's just call this a double bond and leave it at that.
Okay, so if we go to di-Nitrogen or N[2], here's where we would start with the two neutral atoms. Neither satisfies the Octet Rule. We'll move this lone pair in to form a bonding pair. Remember we're always going to be moving lone pairs to make bonding pairs whenever we don't have enough electrons to satisfy the Octet Rule. So we'll do that, and that gets to here. Now you'll see that the Octet Rule is satisfied for the left-hand side Nitrogen, but not the right-hand side. So we'll take another lone pair, now from the other side, and put it in. And now we have a molecule that satisfies the Octet Rule for the left-hand Nitrogen and the Octet Rule for the right-hand Nitrogen. Now we have three pairs that are bonding pairs and only two pairs that are lone pairs, and this is what we're going to call a "triple bond." And you can see that there isn't anything particularly special about homonuclear diatomics or the same atom. Carbon monoxide is what we call "isoelectronic," meaning that there are the same total number of valence electrons, and you can see that we end up at the exactly same Lewis Dot Structure, except instead of a nitrogen and a nitrogen, we've got a carbon and an oxygen, but we've got this triple bond. And again, we don't know what a triple bond is but the theory predicts that there should be something different about the CO bond in carbon monoxide versus, say, carbon dioxide--something you haven't looked at, but which we'll work out later on.
All right, so let me show you how you do these. Let's look at methane. Methane natural gas. The first thing you do is you add up the total number of valence electrons, so we've got four valence electrons for the carbon, four valence electrons for the hydrogen, for a total of eight valence electrons. And at this point, you have to have somebody tell you how the atoms are connected together. We're not really at a position where we can predict how the atoms are connected. In this case, carbon is the central atom. The hydrogens go all the way around, and we just have to put in our eight valence electrons. So, there we go. We've got our eight valence electrons. The Octet Rule is satisfied for carbon. The duet rule, or the rule that hydrogen wants two, is satisfied. And this is the Lewis Dot Structure for methane.
Let's look at a slightly more complicated example, Formaldehyde, which is the substance that's used to preserve biological specimens. You have two hydrogens, each of which brings in one valence electron for a total of two. We have one carbon atom with four valence electrons, and one oxygen atom with six valence electrons, 12 valence electrons total. This is the connectivity. And the first thing you should do when you have a lot of different atoms is go ahead and fill everything in as though it were a single bond. So we've put in the single bonds between hydrogen and carbon, and carbon and oxygen, and then we'll put in the rest of the electrons. So that's a total of 12. And you can see that the hydrogens are perfectly content. The oxygen is perfectly content; it has its octet. But carbon only has six electrons around it--not good enough. What we have to do is to take a lone pair, make it into a bonding pair, and so the correct Lewis Dot Structure for Formaldehyde looks like that. Still oxygen has its octet satisfied. Now carbon has its octet satisfied. This is the correct Lewis Dot Structure for Formaldehyde.
And finally, let's look at an even more complicated problem, which is the carbonate dianion. Carbon has four valence electrons. Each of the oxygens has six, so three oxygens gives you 18. And if we have two negative charges associated with this species, we have to add two more electrons, for a total of 24 valence electrons that we have to assign. And once again, someone has to tell you how the atoms are connected together. Carbon is connected to each of the three oxygens, but the three oxygens are not connected to each other. Let's put in the 24 valence electrons, starting with single bonds between the carbons and oxygens, and then go ahead and put in all the other valence electrons, and that's 24 total electrons. To remind ourselves that this thing is a dianion, it's always a good idea to put it in square brackets and put 2- on it to remind yourself that there are two extra electrons here.
All right, what's the problem with this? I mean, it actually looks pretty good. The oxygens all have their Octet Rules satisfied, so that one has eight, this one has eight, and this one has eight. But carbon doesn't have eight. Carbon only has six, and so in order to fix that what we're going to have to do is, again, turn a lone pair into a bonding pair. It's entirely arbitrary which pair we choose, so let's just go ahead and choose this one. And that gets us to the correct Lewis Dot Structure for carbonate. And you can look more carefully at this and see that in fact the Octet Rule is satisfied for each of the constituent atoms.
All right, so the Octet Rule provides a powerful way for predicting how the electrons in a molecule should be distributed. We're not at a point where we can predict how the atoms should be connected together but it did allow us to predict something--which we haven't really described yet, called a "double bond" and something else called a "triple bond"--that we're going to find out more about later. And the idea is that this Octet Rule, Lewis Dot Structure, is a really powerful predictor of how the atoms are connected together.
Chemical Bonding: Fundamental Concepts
Lewis Dot Structures
Lewis Dot Structures for Covalent Bonds Page [2 of 3]

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