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Video transcript

in this video we're going to think about constructing Lewis diagrams which you've probably seen before they're nice ways of visualizing how the atoms in a molecule are bonded to each other and what other lone pairs of valence electrons various atoms might have and so let's just start with an example then we'll come up with some rules for trying to draw these Lewis diagrams so the first example that we will look at is silicon tetrafluoride and tetrafluoride is just a fancy way of saying four fluorines so tetra fluoride now the first step is to say well what are the electrons that are of interest to us and if we're talking about the electrons that are likely to react we're talking about the valence electrons so v period EP read for short valence electrons so first was just think about how many total valence electrons are involved in silicon tetrafluoride well to think about that we could think about how many valence electrons the silicon have and then how many valence electrons does each of the fluorines have if they were just three atoms in neutral and then multiply that times four because you have four fluorines so let's get out our periodic table of elements and then you can see here that silicon its outer shell is the third shell and in that third shell it has one two three four valence electrons so silicon here has four valence electrons and then to that we're going to add the valence electrons from the four fluorines a free neutral fluorine atom it's outer shells the second shell and in that outer shell it has one two three four five six seven electrons so each of these fluorines has seven valence electrons but there are four of them so one silicon tetrafluoride molecule is going to have four plus twenty eight valence electrons so this is going to be a total of 32 now the next step is to think about how might these be configured and as a general rule of thumb we'd want to put the least electronegative atom that is not hydrogen at the center and we've talked about this before but you can even see from the periodic table of alum Florine is actually the most electronegative element and so we would at least try to put silicon at the center and make fluorine a terminal atom something on the outside so let's try to do that so let's put silicon in the center and then we have to put the four four in some place let's just put one fluorine there one fluorine there one fluorine there and one fluorine there now the next step is well let's just say for simplicity that we just have single bonds between the silicon and each of the fluorines so let's do that so one bond a bond a bond a bond now each of these covalent bonds each of these lines in our Lewis diagram they represent two electrons so for example this one right over here that I'm doing in yellow that represents two electrons that are shared by this fluorine and this silicon this represents another two electrons that is shared between this fluorine and the silicon there's another two electrons disrupt that's shared between this fluorine and this silicon and this is another two electrons shared between that fluorine and the silicon so so far how many electrons have we accounted for well each of these represent two electrons so two four six eight electrons so if we subtract eight from this we are left with 24 electrons to account for 24 valence electrons so now our general rule of thumb would be try to put those on those terminal atoms with the goal of getting those terminal atoms to having eight valence electrons and in general we try to get the octet rule for any atom except for hydrogen hydrogen you just need to get to two in that outer shell but fluorine you want to get it to eight it already has two that it can share so it needs six more so let's add that two four six let's do that again for this fluorine two four six do it again for this fluorine two four six and then last but not least for this 14 to four and six now how many more electrons are now accounted for well six in this fluorine six in this fluorine six in this fluorine six in this fluorine so six times four we've now accounted for 24 more electrons we've now used up all of the valence electrons now that's good because we want to account for all of the valence electrons we want to represent them somehow in this Lewis diagram the next thing to check for is how satisfied the various elect the various atoms are relative to the octet rule we've already seen that the fluorines are feeling pretty good they each have six electrons that are not in a bond and then they're able to share two electrons that are in a bond so each of them can kind of feel like they have eight outer electrons eight valence electrons hanging out with them and then the silicon is able to share in four bonds each of those bonds have two electrons so the silicon is also feeling good about the octet rule so I would feel very confident in this being the Lewis diagram sometimes called the lewis structure for silicon tetrafluoride so just to hit the point home on what we just did I will give you these steps but hopefully you find them pretty intuitive that's why I didn't want to show you from the beginning but as you see step one was find the total number of valence electrons we did that that's the four from silicon and then the 28 from the fluorines it says add an electron for every negative charge subtract an electron for every positive charge we didn't have to do that in this example because it's a neutral molecule then it says decide the central atom which is should be the least electronegative except for hydrogen that's why we picked silicon because fluorine is the most electronegative atom and then we drew the bonds we saw that the bonds accounted for eight electrons and we subtracted those electrons from the total in step one and that's just to keep track of the number of valence electrons that we are accounting for and then we had 24 left over and then the next step it says assign the valence electrons to the terminal atoms that's where we assign these extra lone pair electrons to the various fluorines giving them an extra six each so that they were all able to fulfill the octet rule and then we subtracted that from the total really just to account to make sure that we're using all of our electrons it says it right here subtract the electrons from the Tolan step two and then we saw that all of our electrons were accounted for but then in step four it says if necessary assign any leftover electrons to the central we didn't have to do that in this example if the central atom has an octet or exceeds an octet you are usually done in this case it had an octet so we felt done and it finally says if a central atom does not have an octet create multiple bonds once again in this example we were able to stay pretty simple with just single bonds but in future examples we're going to see where we might have to do some of these more nuanced steps so I will leave you there and I'll see you in the next example