In the previous lecture we developed a model for chemical bonding based upon a sharing of electron pairs between atoms. And we developed a methodology for drawing molecular structures, that might help us understand how molecules form and what types of molecules form. In this lecture we're really going to flush that out in detail. Remember,our model is to design these things called Lewis Structures. And the Lewis Structures are based upon two principles. One, is the octet rule. That is that the atoms tend to bond in such a way that they can have a complete set of eight electrons in their valence shell. Meaning that we take the number of valence electrons that they have, and the number of electrons that they need to complete an octet will be equal to the number of bonds that they form in general. And that they get those extra electrons by sharing electrons in a covalent bond. Let's illustrate this with a couple of the molecules that we actually began the last lecture with, but had not gone back to analyze. For example C2H6, which goes by the name ethane. We might actually be able to use our lowest model to figure out how ethane bonds. Remember that we have drawn a carbon atom with four valance electrons. We know carbon atoms have four valance electrons, but we've drawn it in this way to emphasize that there are four electrons which could be shared in an electron pair with four other atoms. In C2H6, of course, there are two such carbon atoms. And of course there are six hydrogen atoms out there, each one of which has a single valence electron. It should be relatively clear that the two carbon atoms are going to have to bond with each other. They can't bond strictly through hydrogens and not to each other, so they're going to form a bond here. By sharing this pair of electrons here, in the normal way that we are now drawing for a covalent bond. That leaves, of course, an unshared electron on each corner here for the carbon, so there are three unshared, unpaired electrons. But there are six hydrogens, and notice that there are six. Places where we might share hydrogens with those carbon atoms, so when we complete all of the shared pairs of electrons, each hydrogen bringing in its own electron, we wind up with a model then, for the ethane molecule And notice that in this model we have, let's see, around each carbon atom, if we just sort of count up the carbon electrons in the valent shell of each carbon, there are two, four, six, eight, that's the octet we were looking for. And around this carbon atom as well, there are two, four, six, eight. But we notice as well that there is a shared pair. And so we have then here a covalent bond between the two carbon atoms. And in addition we have an octet for each of the two carbon atoms. And correspondingly, at least according to our model, it should not surprise us That this particular molecule, which is called ethane, is in fact a stable molecular structure and a stable compound. Let's do another molecule then. One of the other ones that we considered was in fact C3 H8, a molicule which we call propane. And, not belaboring the point too much, when we draw the four carbon atoms, or three carbon atoms adjcent to each other, we clearly wind up with a chain of three carbon atoms, and one, two, three, four, five, six, seven, eight unshared, unpaired electrons around those carbon atoms which is exactly right for the number of hydrogens we have so correspondingly it should not surprise us then to draw hydrogens around the outside here each sharing an electron from itself and from the carbon into a pair of electrons. Notice that each of the carbon atoms now has an octet of electrons, consistent with our experimental observation of the octet rule, and we wind up then with a structure in which every carbon atom has satisfied the octet rule. Let me make a note. That typically chemists will not draw all of these little pairs of dots but rather will quickly replace those pairs of dots with a straight line. When you see that straight line, you have to get used to believing, or, or seeing rather, that each straight line corresponds to. A shared pair of electrons, and is therefore a covalent bond. Let's do one more of these. The next one in the series, obviously, is going to be c4h8. Those are the ones that we listed before. I'm sorry c4h10. This is a molecule referred to as butane. And by now, you can imagine how this is going to go. The first four carbons, the first carbon with its four unshared, unpaired electrons. The next carbon comes in, the next carbon comes in, the next carbon comes in. We wind up with bonds amongst the four. Again, hydrogens share those unpaired, unshared electrons on each carbon to complete an octet on each carbon. And a pair of electrons for each hydrogen. Notice again, each carbon now has an octet of electrons. And again this molecule is consistent with our molecular structures, it's consistent with the idea of covalent bonding, and of the octet rule, and so it doesn't surprise us then that butane is a stable molecule. Notice there's a pattern developing here, we had c 2 h 6 c 3 h 8 c 4 h 10. Every time I add a carbon I'm adding two hydrogens. That seems clear if I compare these structures because I've basically inserted another carbon. Into this chain here, and with the carbon it binds to a carbon on either side, it has two locations for additional hydrogens. So if we look at these, we can actually see a formula, Cn carbon atoms, H2N plus two carbon atoms/ Try this out, if n is two then h six that's right, if n is three h eight that's right, if n is four h ten that's right. Any molecule that has this particular molecular formula for whatever integer of n we choose in there Is likely to be a stable molecule. We wind up then with our first prediction arising out of this model we have developed. That the Lewis structure model works very well to explain the stability of molecules with the formula Cn H2n plus 2. It also works well for accounting for the fact that we don't see some other molecules. In the previous lecture I mentioned for example, c2h20. Well, if I tried to draw c connected to another c, there's no place to put 20 hydrogens. Correspondingly, this would not be stable. Imagine slightly differently that I tried a c2h5. If I try to draw that one, with a carbon next to it, now I have five hydrogens which could come in to play but there remains over here a single unpaired, unshared electron. Which will probably wind up going and looking for something to react with, so as to complete the octet on that carbon. Correspondingly, this is also not a stable moleculant species. So are model is actually pretty darn good for making predictions about what's going to work and what's not going to work. But, turns out there's some other molecules out there that do not fit this general form of C2, HN 2N plus 2. In fact, there's a lot of molecules that don't fit that. These work, but these are not the only things that work. Lots of other things will work as well. I've illustrated several of the molecules here. Sort of make note of them. C two h four, that clearly doesn't fit the c n two n h two n plus two model, that's c n h two n. Similarly here, or even here. Notice that is c n h n. So there are other possibilities as well. How should we account for the bonding in those kinds of structures? One way to do that is to actually go back and look at some experimental data. Because the experimental data is going to tell us that these three molecules have rather different properties than the previous molecules we have looked at. Lets look at three molecules that seem to sort of go into a set. So for example here, notice what I have here is C2H6, that fits the C n H 2n plus 2 model, that's one of the ones we just worked out. Here's C2H4 with two fewer hydrogens, and CH2 with two fewer hydrogens yet. Neither of the last two molecules fit this general molecular formula. Notice that the properties of the bonds forming those molecules are very different. We looked at two different kinds of properties here. We have something we are going to call the bond strength and something we are going to call the bond length back over here. The bond strength and the bond length. The bond strength refers to how much energy is required to pull apart the bond, to pull the atoms apart in that bond. And the bond length is how far apart are those atoms. Pm is picometers, that's ten to the minus 12 meters or ten to the minus three nanometers. So these are the measurements that can be made experimentally notice that in these three molecules, the strength- in terms of how much energy is required to break the Carbon-Carbon bond, and the length- how far apart the Carbon-Carbon atoms are, are very different in these three molecules. That suggests that the bonding is very different in these three molecules. So now lets attempt to draw the molecular structure for these three molecules as well. Of from the two new ones, at any rate. Let's start with C2H4. Well, here's the carbon with its four electrons around the outside. Here's the other carbon with its four electrons around the outside. Sharing a pair of electrons between the two carbon atoms here there are, I'm sorry I wrote six but I should have said four, C2 H4 is the molecule we're looking at. If I add the hydrogen's. Bunch of different ways I can do this but I'll do it like this. Notice that we wind up with two Unshared, unpaired electrons up here, which might imply that this is not a stable molecular structure, but it is. This molecule is actually called ethene, or sometimes called ethylene, was its older name. We'll call it ethene. It's actually a stable compound. So how could it possibly be stable with these two unshared, unpaired electrons? And notice, neither carbon seems to have satisfied its valence. It has not satisfied the octet rule. Each carbon has formed only three bonds, the valence of carbon is four. Each carbon has only seven valence electrons. We expected that it wanted eight. The answer to this question is fairly readily apparent. If this electron could be shared With the carbon and the other electron can be shared with the carbon, I could draw a structure in which, the two carbon atoms are sharing two pairs of electrons instead of just one. The structure in which I have drawn two lines between the carbons, representing two bonds between the carbons, and I call this thing a double bond. But I wouldn't want to do that just on the basis of the octet rule because the octet rule was based entirely upon bonding to hydrogen atoms. So what would be the justification for thinking that we could do such a thing? The answer is this data here look at the difference in the strength of the bond in C2H6 which had a single bond in it versus the strength of the bond C2H4 that has a double bond in it. Likewise apparently by having a double bond the bond length is shorter. When we have actually two bonds in there compared to having just one bond. How about this other molecule, C2H2, which is the third one in our table here. The C2H2, again, we're going to have the carbon with three electro- -- four electrons around the outside of it, the other carbon With four electrons around the outside of it. We'll share one pair of electrons as we are used to, we'll add one hydrogen to each carbon since there are only two hydrogens to go around. In this structure, there are lots of unshared unpaired electrons, so it seems that this should be unstable, and sure enough the carbons also don't seem to have satisfied their valence. Nor do they have octets but perhaps, if we take these electrons that are unshared and unpaired and share and pair them, we could draw a structure in which there are three paired and shared electrons. Correspondingly, we might draw this as a triple bond as follows and in fact this molecule is variously known either as ethyne or much more commonly acetylene, a very common and well known material and in fact acetylene, although it is a highly reactive molecule it is a stable compound. You can purchase this in bottles and use it as a fuel, or use it for welding torches and so forth. And agan, what is the evidance that we can form such a triple bond relitive to the double bond, the answer is that our data are consistant with this, that if I look at the energy of the triple bond, it is greater than the energy of the double bond, and the bond length of the triple bond is shorter than the bond length of the single bond. And notice as we go all the way across, that in fact the bond strengths get larger as we go from, the single bond to the triple bond and the bond lengths get shorter as we go. From the single bond to the triple bond. So we now have a model in which we can account for double bonding and triple bonding and we see data that suggests that there are lots of different ways in which the atoms can bond. Now, so far all we've done is to get carbon atoms and hydrogen atoms and bond those together. Clearly our octet rule was based upon adding in fluorine, oxygen, nitrogen amongst other atoms and we'll take that up in the next lecture.
