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How to get Lewis structures from electron configuration, and how valence electrons are involved in forming bonds. Created by Sal Khan.
Video transcript
Voiceover: Now that we know a little bit about electron configurations, I want to start thinking about Valence electrons which are the electrons that are associated with an atom. They are the ones that are most likely to react with other atoms or most likely to form bonds or be taken away or shared in some way with other atoms. We've already kind of talked a little bit about this as we talked about the trends through the periodic table but now we'll do a little bit more formally and we'll use what we call Lewis dot structures to visualize this. As I said this tend to be the highest energy and furthest out. You might be saying well aren't the highest energy the furthest out or aren't the furthest out, the highest energy electrons. Well that is true if we're talking about the s-block, if we're talking about the s-block right over here or if we're talking about the p-block right, the p-block right over here but that is not true. That is not necessarily true for the transition metals for the non-main block elements. That's why I'm not using an absolute. I'm saying this it tends to be the highest energy for the said electrons. Let's think this through a little bit and let's try to draw some Lewis dot structure so we could start with the simplest of elements. We could start with hydrogen. Hydrogen has one electron. Its electron configuration is one, S1 that one electron is its one Valence electron. That's the electron that hydrogen will use when it reacts. It could be swiped away from hydrogen where hydrogen will just become a positive ion or maybe it will share it with another one, it can share this electron with another atom and that atom will share an electron with it and then hydrogen maybe could pretend to have a more stable configuration like helium. That's its one Valence electron. We would do that just as a dot around the hydrogen. Now what about the other group one elements? Let's say sodium. Sodium, what's its electron configuration. Well I use a short hand where we could say, well it's base configuration is the same as neon. It has a base configuration of neon. Neon is one S two, two S two, two P six, that's what this represents and then to get to sodium, you would then have three S one. How many Valence electrons does sodium have? Well its highest energy, furthest out electron or I say the electron that's in a non-stable shell. That's in a shell that hasn't been stabilized. It hasn't gotten to its fully stable state. There's only one electron in that situation right over here, the three S one electron. Sodium as well, you could depict like that. It only has one Valence electron that's the electron that could be swiped away from it or that somehow could be involved in a covalent bond somehow. Now let's do things with more, more Valence electrons than hydrogen or sodium but the important thing to realize and actually for the example of hydrogen sodium is that all of these group one elements are going to have one Valence electron. They're going to have one electron that they tend to use when they are either getting lost to form an ion or that they might be able to use to form a covalent bond. Now let's think about helium and helium's an interesting character because all of the rest of the noble gases have eight Valence electrons which makes them very stable but helium only has two Valence electrons. The reason why it's included here is because helium is also very stable because for that first shell, you only need two electrons to fill full, to fill stable. Helium has two Valence electrons, its electron configuration is one S two. Once again the reason why it's out here with the noble gases is because it's very stable and very inert like the noble gases that's why we now use those helium for balloons instead of hydrogen. It's not going to blow up like the Hindenburg but you might say, well, if it has two Valence electrons maybe it should be in group two because wouldn't all of the group two elements have two Valence electrons? That actually would be a very reasonable argument and we've seen that already. One can make a very reasonable argument to put helium in group two for that reason. All of the elements in group two are going to have two Valence electrons. Now, let's jump to one of the most interesting and versatile elements in the periodic table, the one that really forms the basis of life as we know it, and that's carbon. I encourage you to pause this video and based on what we've just talked about, think about how many Valence electrons carbon has and what its Lewis dot structure could look like. Well carbon's electron configuration is going to be the same as helium plus you're going to have two S two and then two P two. How many electrons does it have in its outer most shell that has not been completed yet? Well it has these four, two plus two. We could depict them as one, two, three, four Valence electrons. Why is this interesting? Well we can now think about especially if we see carbon's Valence electrons and we see hydrogen is this Lewis dot structure, we can begin to predict what types of molecules carbon and hydrogen could form together. For example, carbon would like to get to A, it would like to pretend like it has electrons so it feels more stable like the noble gas neon and hydrogen would like to at least feel like it has two electrons in its outer most shell so it can feel more stable like helium. If these are carbon atoms and if these are hydrogen atoms, I'll do the hydrogen orange. Actually let me do it this way. Hydrogen, hydrogen, hydrogen and actually let me do it the way I was doing it first. You could imagine something like this where carbon could bond just based on what we've learned about Valence electrons and Lewis dot structures. You could say, well, I would predict that maybe a molecule like this could form where a carbon shares it's four Valence electrons with four different hydrogens and in exchange it shares an electron from each of those four hydrogens and so the carbon can feel like it has eight Valence electrons. Each of the hydrogens can feel like it has two Valence electrons. If you did this, if you say, "Well there should be some molecule out there" "in nature that seems pretty stable like this." You would be absolutely correct. This is methane and the way that this would be depicted with the Lewis dot structure is this right what I did over here is less conventional notation. Each of these electron pair so that electron pair would be represented as a covalent bond. This would be represented as a covalent bond. That would be represented as covalent bond. That would be represented as a covalent bond. Each of these bonds or the sharing of essentially two electrons, the two electrons. Carbon can feel like it has two, four, six, eight electrons even it's sharing. Each of the hydrogens can feel like they have two electrons which gets it into a more stable state. In any of the elements, in carbon's group, they are all going to have four Valence electrons. For example 10, even though neutral 10 is going to have 50 electrons, the Valence electrons, the ones that are going to react are going to be the one, two, three, four in its outer shell. One, two, three, four and so you might predict well maybe it could form bonds not too different or it might react in a similar way in some ways to say something like carbon and/or maybe silicon could react something similar to say carbon. People even think, there could be life forms in other planets that aren't based on carbon but actually are based on silicon because silicon would have similar types of bond that it can form similar types of structures to carbon for this exact reason. Now what about, I kind of said that you have your transition metals, you d-block right over here and actually your f-blocks are going to be thrown in here as well and these are special cases. These get a little bit more involved because as we already learned that once you're in the fourth period, Let's say we want to do the electron configuration. of say iron. Iron's electron configuration, we could start with argon as a base and then we're going to go four S two and then we're now in the d-block but we're not going to fill the four D suborbital. We're not going to back up and back fill into the three d-suborbital. It's one, two, three, four, five, six. It's three D six. This is where this gets a little bit more vigorous. What are the highest energy electrons? Well those are these D electrons right over here. What are the furthest out? Well they're the ones in the fourth sub shell, these four S two. That's why iron's reactivity is a little bit at least based on just this superficial electron configuration. It's a little bit harder to predict. Iron is known to lose one electron, known to lose two electrons, known to lose three electrons and so those could be some combination of these highest energy electrons, both the ones that are furthest out and the ones that are highest energy. Iron becomes all the transition elements, figuring out the Valence electrons, the electrons that are most slightly right becomes a little bit hard to predict. Some people, some even textbooks will say, "Oh, all the transition metals" "have two Valence electrons" "because they all get the four S two" "and then they're back filling.” They would say, "Okay," "these are the two Valence electrons" "for all of these transition metals." Well that doesn't hold up for all of them because you even have special cases like copper and chromium that only go four S one and then start filling three D depending on the circumstances. Sometimes it does it otherwise but even for the other transition elements like say iron is not necessarily the case so these are the one, the only two electrons that are going to react. You might have some of your D electrons, your three D electrons which are high energy might also be involved in reaction. Might be taken away or might form a bond somehow.