This lecture marks actually an exciting transition in our study of chemistry. The material we're going to be studying here is from the 7 Concept Development Study on Covalent Bonding and Electron Pair Sharing. This transition actually is the point at whcih we move away from studying atoms and begin to study molecules. You'll recall at the outset of the semster, we demonstarted that all matter is made up of atoms, which combined in small integer ratios to form molecules, but then we've spent several weeks really focused just on the properties and structures of atoms, not really paying much attention to molecules at all. But the number of pure substances made up of atoms of just one type, just the elements themselves, is quite small. There are only 90 naturally occurring elements. There are only 118 elements that we know all together. By contrast, the number of compounds that are known, both inorganic and organic compounds, actually according to a number I looked up just this morning, is 71 million in the chemical registry. Those compounds are made up of molecules. So it's important for us now to begin the transition to studying of molecules. Chemistry, really, is the study of molecules. Their structures, their properties, and the reactions that they undergo as the atoms in those molecules rearrange to form new molecules, with new properties. and new attributes that might be of use to us. So up to this point what do we know about molecules? The answer right now is about all we know are molecular formulae. Remember we developed a means to study molecular formulae, to determine them by weighing out the masses of the constituent elements in each compound. And by knowing the relative masses of each element for example, carbon to hydrogen on atomic mass scale carbon is 12.01 and hydrogen is 1.01. We can use those to determine the number of carbon atoms and the number of hydrogen atoms in individual molecules. So we wind up with molecular formula that looks like this. Here's a few of them. A single carbon atom combines with four hydrogen atoms. Two carbon atoms could combine with six hydrogen atoms. Eight carbon atoms could combine with 18 hydrogen atoms. These are molecules which are known. The first is methane. The second is ethane, the third is octane. These are common hydrocarbons that we use, for example, for fuels, but when we look through the registry of all of the different kinds of compounds formed from carbon and hydrogen what we discover is that some types of compounds are not present. For example, very uncommon to try to find a molecule that has four carbon atoms and one hydrogen atom, even though it's very common to find a molecule that has one carbon atom and four hydrogen atoms. The reverse ratio doesn't seem to work. And although we can take eight carbon atoms and 18 hydrogen atoms... Turns out we can't take 2 carbon atoms and 20 hydrogen atoms. Some combinations of atoms just don't form. So at the outset for our understanding of molecules, what we need to understand is what is it which determines what molecules can form and what kinds of molecules cannot form? What ratios are permissible, what ratios are not permissible, and what's the underlying reason for that. What we will discover is that we need to understand how it is the atoms actually combine to form molecules. What is it that holds them together? And what does it mean to form a chemical bond between two atoms inside of a molecule? To understand that, we need to understand the new property of atoms first. This is a property we actually haven't spent that much time talking about. Which is this valence of an atom. The valence of an atom, is the number of other atoms that it typically likes to form, eh, to bond with, the number of bonds it likes to form, as it forms molecules. So we'll take this sort of the short hand, and reminder ourselves that the valence of an atom is the number of bonds that it most commonly forms. Over here, actually, we have such molecules, CH4 for example, that show us that carbon typically forms 4 bonds, for example by combing with 4 hydrogen atoms. So we will say that it's valence of the carbon atom is equal to 4. Just an experimental observation of the number of bonds that a molecule might like to take. That's because we see it from methane. Let's try nitrogen, which forms the common molecule ammonia. Which, it turns out, has molecular formula NH3. So nitrogen apparently forms 3 bonds. So the nitrogen valence, the valence of a nitrogen atom, is three. Similarly of course, we know that oxygen commonly combines with hydrogen to form H2O. So oxygen forms two bonds. So the oxygen valence is 2. Carbon is 4, Nitrogen is 3, Oxygen is 2. We can do this for a whole set of different compound, I'm sorry different el atoms to determine what the common valences are for various atoms. I tabulated them here for elements on the second and third row of the periodic table and what we see here is a couple of really interesting patterns reflective of what we call the Periodic Law before. The first of these patterns, of course you'll notice, is that the common valences are similar for atoms in the same group of the Periodic table. The second thing that you'll note, is that the valencers are also, a periodic function. They are a periodic function. Going across a row of the periodic table notice how it rises from one to four and back down to zero and then starts over again, starting one to four and then back down to zero as well. That is a really interesting pattern and to understand that pattern we want to compare it to some other patterns that we've observed. In particular, having to do with the number of valence electrons that each atom has. Well remember, that's a familiar concept from before, the number of valence electrons is the number of electrons in its outer most shell. In this particular tato-, table, what I've shown you is a set of data, where we're comparing two different kinds of data. The first, is the valence Which you'll recall, the valence is the number of bonds that the atom likes to form, it's not equal to the of valence electrons. The number of valence electrons is the number of electrons which are in the outer shell of the atom. And here if we look at these we notice that there are interesting patterns as well. Most notably, if we stare at these data for just a little while we actually see one particularly fascinating thing. Let's sum up the number of valence electrons plus the valence for each atom. For lithium that's two, for berillium four. For boron, six, for carbon, eight. We might start thinking that this is an interesting pattern of two, four, six, eight. But now the pattern changes in extremely important way. Notice for nitrogen, 3 plus 5 is 8, for oxygen 2 plus 6 is 8, for, for fluorine 1 plus 7 is 8, and for neon Zero plus eight is eight, and we see a fascinating fact here that for the elements in group four through eight, carbon through neon, the sum of the valence of the atom and the number of valence electrons in the atom is always eight. That observation is actually summarized in one of the most famous rules in all of chemistry called the octet rule, and here it's just an experimental observation. It is the observation that the common valance of an atom plus the number of valence electrons in that atom equals eight in groups four through eight. Now just remember again groups 4 through 8 are these elements here, carbon going through neon. Turns out this rule works well in the next row of the periodic table, going from silicon to argon, as well. It does not work for the elements in group 1 through 3. So remember, the octet rule is an experimental observation based upon the atoms in groups 4 through 8. Now eight is an interesting number. We've seen the number eight before. Eight is in fact the maximum number of electrons that could fit into a valence shell. For example, going from neon to sodium, we knew we had to put a new electron in sodium into the next valence shell because the valence shell of neon was filled with its eight valence electrons. Do we think that this is a, a coincidence here that the octet rule tells us that the valence plus the number of valence electrons is eight and we know that the maximum number of electrons that can fit into a valence shell is eight? That combination is a fascinating combination. Because it tells us something interesting here. If we realize that the maximum number of valence electrons is eight and we subtract from that the number of valence electrons then there's a number of vacancies. For example, let's take nitrogen where the number of valence electrons is 5. The maximum number of electrons that would fit into the shell about nitrogen is 8. 8 minus 5 is 3, which is exactly equal to the valence of nitrogen. So what we have discovered is that the valence of each atom is equal to the number of vacancies in the valence shell of that atom. Let's think about that nitrogen example again here. Nitrogen likes to form three bonds. Its valence is 3. It has 5 valence electrons. The maximum number of electrons in the valence shell lis 8. 8 minus 5 is 3. The vacancies are 3, and the number of bonds it wants to form is 3. That suggests a really interesting model, where we say that the number of bonds the atom wants to form is equal to the number of vacancies in its valence shell. If that's the case, then the model that we developed is the following. It appears to be the case that each atom, and again we're talking about groups 4 through 8, forms a number of bonds, which will allow it to add enough electrons to fill its valence shell. Thinking again about nitrogen, it, it has three vacancies. It forms three bonds. Apparently it is bonding in such a way that it can complete an octet of electrons by filling it's valence shell. by bonding to three other atoms. This process is actually what we're going to call a covalent bond, and it is described as a shared pair of electrons. Why shared? Let's think briefly for a minute about, flourine in particular, which we know has seven valence electrons. And let's remember as well, that fluorine forms the molecule F2, we know that fluorine is a diatomic molecule, so that means that somehow or another, fluorine atoms are going to bond together. If one fluorine atom's attempting to bond with its seven valence electrons, To another fluorine atom with its 7 famous electrons, the fluorine on the left cannot satisfy its valence by stealing electrons from the fluorine on the right because then the fluorine on the right would not have satisfied its valence. Apparently the way in which the electrons get satisfied, the valence get satisfied is they must share one pair of electrons in order to satisfy this octet rule, in which the number of bonds formed is equal to the valence of these atoms. So when now we have A slight variation on the Octet Rule. The Octet Rule before was just an experimental observation, but now we have a model of the Octet Rule, which suggest that atoms will bond to satisfy their valance by forming enough covalent bonds, meaning sharing a pair of electrons, such that they will each windup with eight electrons in the valence shell. Remember this rule applies only to groups fourth to eighth because experimentally we only saw that to be true back in the data doesn't work for group 1 through 3. So what we can say here then is that atoms have the tendency to bond to complete an octet of valence electrons. That is the model that we have developed here. Of course, hydrogen is an important exception to this. We'll forget helium for the time being since it forms no compounds. Hydrogen has only a single electron. That first shell has only room for two electrons, as having, in spite the fact that there are only two elements in that row of the periodic table, hydrogen and helium, before we have to move to new shell with lithium. So the separate rule here, separate from the octet rule, is that hydrogen atoms bond to form two electrons in their valence shell. What that gives rise to then, is a new model that we're going to call the Lewis model structure for bonding. It is not just a model for bonding, which is what I've written here. It is by and large the single most important model for bonding in all of chemistry. Indeed, all of organic chemistry is based upon the Lewis model for bonding. And it is based upon the two things that we have discovered in this lecture, the octet rule and covalent bonding. And it allows us to develop an interesting model for things called Lewis Structures, which is just a representation of these 2 rules, that atoms tend to bond in such a way as to form a cov, an octet of electrons. And they do so by sharing electrons with other atoms. So to do that, we'll actually draw a diagram following Lewis to say carbon has 4 valence electrons. And we will write them as little dots surrounding the carbon atom. Separated from one another. And the advantage to drawing them in this way is that if each, bond is going to form, and carbon wants to form four bonds. Remember, it has four valence electrons, and its valence is also equal to four. And if it is to share electrons to complete that valence, then in fact it cannot share electrons which are already being paired. So we draw this structure in this way. Hydrogen you'll recall has a single valence electron, and it only wants to form one other bond, because there's only one vacancy in it's shell. Then we can wind up drawing the structure for ch 4 the most common compound of carbon and hydrogen, or the simplest compound of carbon and hydrogen, in which we form 4 covalent bonds by sharing hydrogen's electron with carbon for each of the previously unpaired electrons, forming four pairs of electrons corresponding to the four bonds. More commonly, this structure is written as lines. Each line representing a pair of electrons. And either of these two structures are what we refer to as the Lewis Structure or the molecular structure of the methane molecule. So, we figured out two things. One, why does carbon form four bonds? The answer, because it has 4 valence electrons and therefor 4 vacancies in its valence shell. How does it form those bonds? It forms those bonds by sharing pairs of electrons with other atoms. And therefore, we've figured out what the structure of the molecule is, CH4. In the next lecture, we're going to use this powerful methodology to draw a lot of Lewis structures for many different molecules. And in the lecture after that, we'll actually begin drawing structures that involve atoms other than carbon and hydrogen.
