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  • - [Sal] 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 tetrafluoride.

  • 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.E. for short,

  • valence electrons.

  • So first let's 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 does silicon have,

  • and then how many valence electrons

  • does each of the fluorines have

  • if they were just free atoms and neutral,

  • and then multiply that times four,

  • 'cause 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,

  • its outer shell is 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 gonna have four plus 28 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 wanna 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 elements,

  • fluorine 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 fluorines 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.

  • This is another two electrons 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.

  • 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 fluorine.

  • Two, 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 wanted to account

  • for all of the valence electrons.

  • We wanna represent them somehow in this Lewis diagram.

  • The next thing to check for is how satisfied

  • the various atoms are relative to 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 wanna 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 should be the 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 assigned 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 total in 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 atom.

  • 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.