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All right, let's look at chapter 11, problem 105.
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We have the following image, which is an electrostatic potential map for ethylene oxide, and they gave us the molecule written out, and it's a polar molecule.
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Use the electrostatic potential map to predict the geometry for how one ethyl oxide molecule interacts with another.
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And they want us to draw the structural formulas.
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And using the three -dimensional bond notation introduced in chapter 5, show the geometry.
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So we're going to have to show the geometry of the interaction.
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So unpacking this, we're going to have two steps.
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We're going to have that first step where we're going to draw the structural formula.
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So we're going to draw it out like we did in lewis dot structure back in chapter 5.
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And we're going to use vesper theory, meaning we're going to have to identify the central atom, count the valence electrons, and then connect all the bonds.
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Down here show the geometry interactions so we're going to show the interaction between two different ethylene oxide molecules so let's start with number one we're going to draw a structural formula so up here we have carbon and it's attached to two hydrogens so we're going to put one above and one below so that's our central atom now it says that there's two of these so we're just going to scoot on over put another carbon out there and this is what we're what we did in chapter 5, nothing fancy, except what we learned before with lewis dot structure.
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And we have one oxygen, so we'll put it in there.
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So we drew our central atom and we attached it.
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Now, step two was that we looked and counted all of our valence electrons.
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So we have carbon, and we have two carbons.
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So those two carbons each, when we go to the periodic table, we found that they had four valence electrons they can donate to this molecule.
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So that's eight total.
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And we have one, two, three, four hydrogens.
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And we'll notice that hydrogen's in our first row, which means each hydrogen can contribute one valence electron.
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And when we go and look at oxygen, we will find it has six valence electrons that it can contribute to the structure.
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So adding all of those together, we'll find that we have 6 plus 4 plus 8.
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That's going to be 18 valence electrons.
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And looking at the structure, we're going to try to fill out all of our electrons, use them up, to make this a stable molecule.
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So we're filling in oxygen.
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So it has 1, 2, 3, 4, 5, 6, and then this bond is 2, 8.
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So this oxygen feels 8, it meets its octet.
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That's perfect.
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It's a happy oxygen.
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Let's look at this carbon.
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It has one, two, three, four bonds.
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Each bond gives its two electrons, so we know that it's happy as well.
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Each one of these hydrogens feels one bond with two electrons, so each one of these are happy.
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Now let's look at this last carbon here.
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Oh, this carbon has one, two, three.
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So this guy is less than happy.
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He doesn't have enough electrons.
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We could probably put two here, but let's find out if we've used all 18.
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So we have 2, 4, 6, 8, 10, 12, 13, 14, 15, 16, 17, 18.
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We've already used all 18, and we weren't able to fulfill all of our octets.
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So we know that this is not the correct structure.
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And that's what we learned in chapter 5 was how to do lewis dot structures.
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So let's look at this one here.
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This actually shows us the structure on the, on the electrostatic potential map.
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It'll have an oxygen up top, bonded to a carbon, and it's in a triangle configuration...