0:00

So now we're gonna look through some of the geometries that we talk about

Â in chemistry, and we look at these based on their Lewis structures.

Â And so in this one, we've actually drawn out the Lewis structure,

Â because the key thing is that we have to look at the Lewis structure and

Â determine the number of electron groups around that central atom.

Â Here, we have two electron groups.

Â 0:22

So, therefore, we have this general form of AX2.

Â A is our central atom, in this case that's going to be carbon.

Â X represents a bonding group, and again,

Â I'm not worried about these electrons on the oxygen.

Â Because they're not actually changing anything about the geometry

Â around the carbon.

Â They're an important part of the Lewis structure, but

Â they don't do anything to the geometry of the bonds around that central carbon atom.

Â So when we have something that has this general form AX2, one central atom and

Â two bonding groups.

Â What we see is our electron pair geometry is linear and

Â our molecular geometry is linear.

Â When we start introducing structures where we have some bonding and

Â some non-bonding groups, then we'll have to look at different molecular geometries.

Â Because electron pair geometries which would be better named electron group, but

Â we're saying electron pair to be consistent with our VSEPR acronym.

Â Our electron group geometry describes the position of all the electrons.

Â When we get to molecular geometry, we know that those lone pairs are affecting

Â the geometry and affecting the angles of those bonds, but

Â we're going to describe just what it looks like with the bonding groups.

Â 1:33

Now we can look at our trigonal planar electron pair geometry.

Â And so this is what happens when we have three groups around the central atom.

Â So we're gonna look at two different examples here BF3 and SO2.

Â And we're not gonna draw the Lewis structures here, but I am gonna tell you

Â that when we look at the Lewis structures we see our central atom A is representing

Â that, and we have three bonding groups, so this has three bonding groups.

Â 1:59

And so when we have three bonding groups we know that the electron pair geometry

Â is gonna be called trigonal planar.

Â Because remember we're trying to spread those groups around evenly so

Â that they are maximizing the distance between them.

Â So it's trigonal planar.

Â For our molecular geometry it's still trigonal planar because all of our groups

Â are bonding groups.

Â And that's gonna be the case anytime the number of

Â electron groups equals the number of bonding groups.

Â So in other words, no non-bonding groups, these two will always be the same.

Â Now we're gonna look at a different example where you have SO2.

Â And here we still have three electron groups, but

Â now we have two bonding and one non-bonding group.

Â And what we see is that one non-bonding group still affects the position,

Â it still affects where these two bonding groups are, but when I

Â look at the electron pair geometry, what I see is that it still trigonal planar.

Â Because that is based solely on the number of electron groups around the central

Â atom, regardless of whether they are bonding or non-bonding.

Â Now when I look at the molecular geometry I have to remember

Â that there is a lone pair of electrons on that central atom.

Â And because of that,

Â it's going to affect the angles between these two other bonding groups.

Â So one way to think about this is to imagine that you're doing something that

Â you know you probably shouldn't be doing in the first place.

Â And you hear this little voice in your head,

Â it probably sounds a lot like your parents voice.

Â Maybe your mom or your dad's voice, and it's saying,

Â you know you shouldn't be doing this.

Â Well that's what this lone pair is.

Â We can't really see it, but it's still affecting our behavior.

Â So we can't see these lone pairs of electrons, but

Â it's affecting the position of those bonding groups.

Â And so we describe this as bent.

Â We can't go to linear because we still have to consider that we have these

Â non-binding electrons on that central atom.

Â Now we also have some lone pairs out here on our oxygens.

Â But the thing to remember there is these are kind of like the voices in somebody

Â else's head.

Â They're not affecting the geometry around our central atom or

Â sulfur in this particular example.

Â 4:14

And any time we have the AX2E model, we will have

Â electron pair geometry of trigonal planar, and molecular geometry of bent.

Â And so once we draw our Lewis structure, if we can then write this general formula

Â for our molecule, it's gonna make life a lot easier for

Â assigning those electron pair geometries, as well as those molecular geometries.

Â 4:55

We have the general form AX4, for

Â NH3 which we saw earlier we have three bonding and

Â one non-bonding.

Â And for water we saw two bonding, and two non-bonding.

Â And again I'm getting this information from our Lewis structure.

Â So the first thing I have to do before I can determine geometry

Â is actually to determine the Lewis structure of that molecule.

Â So here we have our four bonding.

Â So we have AX4.

Â For our ammonia example, we have three bonding and one non-bonding.

Â And so now, we have AX3E.

Â Remember, this represents the number of non-bonding groups.

Â 5:38

And the X is the number of bonding groups.

Â Not necessarily the number of bonds.

Â It's the number of bonding groups.

Â Double bonds and triple bonds still count as a single group.

Â Then we get over to water and we see we have AX2E2,

Â we have two bonding groups, and two non-bonding groups.

Â Now notice for all of these our electron paired geometries are all tetrahedral

Â because they all have four electron groups.

Â That's what these molecules have in common.

Â The only thing that's going to be different about them is their

Â molecular geometry.

Â So for our first example, for methane, or CH4, we have all bonding groups,

Â no non-bonding groups.

Â So our electron pair geometry and

Â molecular geometry will be exactly the same.

Â When we go to NH3, we see that we have tetrahedral and

Â our molecular geometry is trigonal pyramid.

Â Because remember we have a triangle kind of shape on the bottom.

Â And this will not be a plane like we saw in trigonal planar,

Â this atom is actually in a different plane.

Â Remember that our bond angles here are not 120.

Â Ideally our bond angles are 109.5 degrees and so

Â the bond angles here are going to be slightly less than 109.5 because

Â that lone pair of electrons is compressing those angles.

Â When we get to our water molecule,

Â again we have two lone pairs that are compressing these angles.

Â So our bond angles much less than 109.5 and less than what we saw in the ammonia.

Â And we described this as bent.

Â Now we also something that was bent when we looked at the trigonal planar geometry

Â when we talked about the AX2E.

Â 7:17

Also had a bent, molecular geometry, but remember that angle is just a little less

Â than 120 degrees, this angle is less then 109.5 degrees.

Â So both described as bent, but they are different from one another because they

Â have a different angle between those bonding groups.

Â 7:36

Now we can look at the trigonal bipyramid electron pair geometry.

Â And here we have five electron groups around our central atom.

Â Doesn't matter whether they're bonding or non-bonding, our electron pair geometry

Â is always going to be trigonal bipyramid or trigonal pyramidal.

Â Either term will be considered correct and so

Â as I go through my samples here, I have five bonding in the first one.

Â I have four bonding and one non-bonding.

Â 8:21

And three, non-bonding.

Â But I still have five groups in each of them.

Â And so as a result, their electron paired geometries are all going to be the same.

Â However, their molecular geometries are going to differ.

Â And so when we look at the first example where we have all bonding groups,

Â we know that our electron paired geometry and our molecular geometry are going to be

Â exactly the same because all of the positions are occupied by bonding groups.

Â Now trigonal bipyramid geometry is a little bit different than the other ones

Â because on those we had kind of one angle that everything was based on.

Â When we look at trigonal bipyramid we see something a little bit different.

Â We have two different angles.

Â Here we have 90 degrees from one another, so that's kind of going from these,

Â what we call our axial positions to our equatorial atoms.

Â We have 90 degrees, but if I look between two atoms here,

Â in these equatorial positions.

Â Or kinda the three around the middle their ideal bond angle is 120 degrees.

Â And where this is going to play a role is when we start looking at where lone

Â pairs are positioned in molecules.

Â Because if I take an electron from an axial position,

Â versus a equatorial position and then replace it with lone pair of electrons,

Â that's going to change which atoms are closest to that lone pair.

Â So I'm always gonna have to do it the same way.

Â So when I go to my AX4E, four bonding and one non-bonding,

Â we actually call this molecular geometry seesaw.

Â So here are the legs of our seesaw, and

Â here's kind of our seats where we could be sitting.

Â So that's where we get the seesaw shape.

Â And our lone pair actually goes into this position

Â on one of the equatorial positions because that's what maximizes its distance

Â from more of those bonding groups.

Â And that's always gonna be the way that it happens

Â with our trigonal bipyramidal electron pair geometry.

Â When I look at my next model, AX3E2.

Â 10:17

Now get to T-shape and now again I'm actually replacing another one of those

Â equatorial bonding groups with a lone pair of electrons, so now we get our T-shaped.

Â And then with two bonding groups and three non-bonding groups now have replaced all

Â of my equatorial bonding groups with non-bonding groups.

Â And as a result I get something that's linear and so

Â now I'm getting back to that 180 degree angle that I saw in linear before.

Â But, I also know that I had three lone pairs around that central atom,

Â and so this is linear.

Â Still the same bond again that we saw with linear electron pair geometry.

Â But the structure looks a little different because we have two bonding groups, but

Â we also have those three non-bonding groups.

Â 11:06

So now we can look at octahedral geometry,

Â where we have six electron groups around the central atom.

Â And when we go to octahedral, we now have 90 degree angles.

Â And it is actually the same through all of our angles in our ideal structure.

Â So trigonal pyramidal were we have some 90 and some 120 degrees,

Â here we are looking at all 90 degree angles.

Â So now what we see,

Â is we have the general form AX6 because we have six bonding groups.

Â 11:33

Our electron paired geometry, and our molecular geometry are the same,

Â we call this octahedral, because we have all bonding groups.

Â The molecular geometry is also octahedral.

Â When we go to our next example where we start replacing a bonding group with

Â a non-bonding group, here we have five bonding groups and one non-bonding.

Â So we have our general form AX5E, our electron pair geometry is octahedral.

Â And our molecular geometry is now square pyramidal or square pyramid.

Â And so, if I look at these atoms along the bottom, they form a square.

Â And it's going to form a pyramid shape coming out from that central atom.

Â So that's our square pyramid.

Â I've got my lone pair here on the bottom.

Â And because all my angles are the same, it doesn't matter where that first lone pair

Â is, unlike with the trigonal bipyramidal.

Â We had to put it in a specific type of position, axial versus equatorial.

Â Here all the angles are the same, so it doesn't really matter.

Â However, when we go to the second low pair of electrons, so

Â here we have four bonding and two non-bonding.

Â Now it does matter where those electron groups are.

Â And so what we want to do is we want to maximize the distance between the bonding

Â groups but

Â we also want to maximize the distance between those lone pairs of electrons.

Â And the way that we can do that is by putting them on opposite sides.

Â So basically in this position and this position.

Â And so by putting the lone pairs there,

Â we're maximizing the distance between the lone pairs.

Â And as much as we can maximizing the distance between the low pairs and

Â the bonding groups.

Â So here we have octahedral electron pair geometry because we still have six groups.

Â But now we have square planar molecular geometry because now the four

Â bonding groups that remain are all forming a square.

Â And they're all in the same plane.

Â 13:58

And I have 21 electrons from the I3, plus 1 more from the minus charge,

Â so I've got 22 electrons.

Â And when I start filling in, I fill in my terminal atoms first.

Â So I have completed my octet on my terminal atoms and

Â I can now put my electrons around my central atom.

Â I do have a situation where I have an expanded octet, and that's okay.

Â Because iodine is far enough down on the periodic table for that to happen.

Â Remember, it's only the second row elements that cannot have

Â an expanded octet.

Â So when I look at my structure here,

Â I've used 2, 4, 6, 8, 10, 12, 14, 16, 18, 20, 22 electrons.

Â So I know I have a reasonable Lewis structure.

Â Both of my outer atoms have an octet of electrons and

Â the central atom actually has 10 electrons around it.

Â So when I look at my electron pair geometry

Â really I'm only worried about that central atom.

Â And here I see that I have three non-bonding groups and

Â two bonding groups.

Â And so since I have a total of five groups,

Â regardless of whether they're bonding or non-bonding that tells me

Â a trigonal bipyramid, or a trigonal pyramidal geometry.

Â Remember, octahedral is six groups.

Â Tetrahedral is four.

Â Trigonal planar is three.

Â And linear, as an electron pair geometry is two.

Â And bent is actually not an electron pair geometry.

Â 15:42

So, now that I know the electron pair geometry, I can now look for

Â the molecular geometry of I3 minus.

Â And again, I'm gonna go back to my Lewis structure that I determined for I3 minus.

Â So we had our terminal iodine atoms had a complete octet.

Â And we had three lone pairs around that central iodine.

Â And we'll put brackets there with a minus charge, indicating that this is an ion.

Â And so I want to write the general form for this molecule which is A, for

Â the central atom.

Â X2, which represents the two bonding groups, then E3,

Â which represents the three non-bonding groups around that central atom.

Â And anytime we have something with the form AX2E3,

Â it's going to be linear molecular geometry.

Â