This course is designed to cover subjects in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. This course is a precursor to the Advanced Chemistry Coursera course. Areas that are covered include atomic structure, periodic trends, compounds, reactions and stoichiometry, bonding, and thermochemistry.
6.07 VSEPR Part 2
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Del curso dictado por University of Kentucky
Chemistry
464 calificaciones
De la lección
Week 6
This unit looks more in depth at molecular compounds to see how they are bonded together by looking at Lewis structures. These structures provide information about the types of bonds (single, double, or triple) as well as the connectivity of atoms. By knowing the Lewis structure, we can also predict the three-dimensional geometry of an individual molecule.
Conoce a los instructores
Dr. Allison SoultLecturer
Chemistry
Dr. Kim WoodrumSr. Lecturer
Chemistry
So now we're gonna look through some of the geometries that we talk about
in chemistry, and we look at these based on their Lewis structures.
And so in this one, we've actually drawn out the Lewis structure,
because the key thing is that we have to look at the Lewis structure and
determine the number of electron groups around that central atom.
Here, we have two electron groups.
So, therefore, we have this general form of AX2.
A is our central atom, in this case that's going to be carbon.
X represents a bonding group, and again,
I'm not worried about these electrons on the oxygen.
Because they're not actually changing anything about the geometry
around the carbon.
They're an important part of the Lewis structure, but
they don't do anything to the geometry of the bonds around that central carbon atom.
So when we have something that has this general form AX2, one central atom and
two bonding groups.
What we see is our electron pair geometry is linear and
our molecular geometry is linear.
When we start introducing structures where we have some bonding and
some non-bonding groups, then we'll have to look at different molecular geometries.
Because electron pair geometries which would be better named electron group, but
we're saying electron pair to be consistent with our VSEPR acronym.
Our electron group geometry describes the position of all the electrons.
When we get to molecular geometry, we know that those lone pairs are affecting
the geometry and affecting the angles of those bonds, but
we're going to describe just what it looks like with the bonding groups.
Now we can look at our trigonal planar electron pair geometry.
And so this is what happens when we have three groups around the central atom.
So we're gonna look at two different examples here BF3 and SO2.
And we're not gonna draw the Lewis structures here, but I am gonna tell you
that when we look at the Lewis structures we see our central atom A is representing
that, and we have three bonding groups, so this has three bonding groups.
And so when we have three bonding groups we know that the electron pair geometry
is gonna be called trigonal planar.
Because remember we're trying to spread those groups around evenly so
that they are maximizing the distance between them.
So it's trigonal planar.
For our molecular geometry it's still trigonal planar because all of our groups
are bonding groups.
And that's gonna be the case anytime the number of
electron groups equals the number of bonding groups.
So in other words, no non-bonding groups, these two will always be the same.
Now we're gonna look at a different example where you have SO2.
And here we still have three electron groups, but
now we have two bonding and one non-bonding group.
And what we see is that one non-bonding group still affects the position,
it still affects where these two bonding groups are, but when I
look at the electron pair geometry, what I see is that it still trigonal planar.
Because that is based solely on the number of electron groups around the central
atom, regardless of whether they are bonding or non-bonding.
Now when I look at the molecular geometry I have to remember
that there is a lone pair of electrons on that central atom.
And because of that,
it's going to affect the angles between these two other bonding groups.
So one way to think about this is to imagine that you're doing something that
you know you probably shouldn't be doing in the first place.
And you hear this little voice in your head,
it probably sounds a lot like your parents voice.
Maybe your mom or your dad's voice, and it's saying,
you know you shouldn't be doing this.
Well that's what this lone pair is.
We can't really see it, but it's still affecting our behavior.
So we can't see these lone pairs of electrons, but
it's affecting the position of those bonding groups.
And so we describe this as bent.
We can't go to linear because we still have to consider that we have these
non-binding electrons on that central atom.
Now we also have some lone pairs out here on our oxygens.
But the thing to remember there is these are kind of like the voices in somebody
else's head.
They're not affecting the geometry around our central atom or
sulfur in this particular example.
And any time we have the AX2E model, we will have
electron pair geometry of trigonal planar, and molecular geometry of bent.
And so once we draw our Lewis structure, if we can then write this general formula
for our molecule, it's gonna make life a lot easier for
assigning those electron pair geometries, as well as those molecular geometries.
Now let's look at tetrahedral geometry.
Here we have four electron groups around our central atom.
For this molecule, for example CH4, we have four bonding groups.
We have the general form AX4, for
NH3 which we saw earlier we have three bonding and
one non-bonding.
And for water we saw two bonding, and two non-bonding.
And again I'm getting this information from our Lewis structure.
So the first thing I have to do before I can determine geometry
is actually to determine the Lewis structure of that molecule.
So here we have our four bonding.
So we have AX4.
For our ammonia example, we have three bonding and one non-bonding.
And so now, we have AX3E.
Remember, this represents the number of non-bonding groups.
And the X is the number of bonding groups.
Not necessarily the number of bonds.
It's the number of bonding groups.
Double bonds and triple bonds still count as a single group.
Then we get over to water and we see we have AX2E2,
we have two bonding groups, and two non-bonding groups.
Now notice for all of these our electron paired geometries are all tetrahedral
because they all have four electron groups.
That's what these molecules have in common.
The only thing that's going to be different about them is their
molecular geometry.
So for our first example, for methane, or CH4, we have all bonding groups,
no non-bonding groups.
So our electron pair geometry and
molecular geometry will be exactly the same.
When we go to NH3, we see that we have tetrahedral and
our molecular geometry is trigonal pyramid.
Because remember we have a triangle kind of shape on the bottom.
And this will not be a plane like we saw in trigonal planar,
this atom is actually in a different plane.
Remember that our bond angles here are not 120.
Ideally our bond angles are 109.5 degrees and so
the bond angles here are going to be slightly less than 109.5 because
that lone pair of electrons is compressing those angles.
When we get to our water molecule,
again we have two lone pairs that are compressing these angles.
So our bond angles much less than 109.5 and less than what we saw in the ammonia.
And we described this as bent.
Now we also something that was bent when we looked at the trigonal planar geometry
when we talked about the AX2E.
Also had a bent, molecular geometry, but remember that angle is just a little less
than 120 degrees, this angle is less then 109.5 degrees.
So both described as bent, but they are different from one another because they
have a different angle between those bonding groups.
Now we can look at the trigonal bipyramid electron pair geometry.
And here we have five electron groups around our central atom.
Doesn't matter whether they're bonding or non-bonding, our electron pair geometry
is always going to be trigonal bipyramid or trigonal pyramidal.
Either term will be considered correct and so
as I go through my samples here, I have five bonding in the first one.
I have four bonding and one non-bonding.
Here I have three bonding, two non-bonding.
And finally, two bonding.
And three, non-bonding.
But I still have five groups in each of them.
And so as a result, their electron paired geometries are all going to be the same.
However, their molecular geometries are going to differ.
And so when we look at the first example where we have all bonding groups,
we know that our electron paired geometry and our molecular geometry are going to be
exactly the same because all of the positions are occupied by bonding groups.
Now trigonal bipyramid geometry is a little bit different than the other ones
because on those we had kind of one angle that everything was based on.
When we look at trigonal bipyramid we see something a little bit different.
We have two different angles.
Here we have 90 degrees from one another, so that's kind of going from these,
what we call our axial positions to our equatorial atoms.
We have 90 degrees, but if I look between two atoms here,
in these equatorial positions.
Or kinda the three around the middle their ideal bond angle is 120 degrees.
And where this is going to play a role is when we start looking at where lone
pairs are positioned in molecules.
Because if I take an electron from an axial position,
versus a equatorial position and then replace it with lone pair of electrons,
that's going to change which atoms are closest to that lone pair.
So I'm always gonna have to do it the same way.
So when I go to my AX4E, four bonding and one non-bonding,
we actually call this molecular geometry seesaw.
So here are the legs of our seesaw, and
here's kind of our seats where we could be sitting.
So that's where we get the seesaw shape.
And our lone pair actually goes into this position
on one of the equatorial positions because that's what maximizes its distance
from more of those bonding groups.
And that's always gonna be the way that it happens
with our trigonal bipyramidal electron pair geometry.
When I look at my next model, AX3E2.
Now get to T-shape and now again I'm actually replacing another one of those
equatorial bonding groups with a lone pair of electrons, so now we get our T-shaped.
And then with two bonding groups and three non-bonding groups now have replaced all
of my equatorial bonding groups with non-bonding groups.
And as a result I get something that's linear and so
now I'm getting back to that 180 degree angle that I saw in linear before.
But, I also know that I had three lone pairs around that central atom,
and so this is linear.
Still the same bond again that we saw with linear electron pair geometry.
But the structure looks a little different because we have two bonding groups, but
we also have those three non-bonding groups.
So now we can look at octahedral geometry,
where we have six electron groups around the central atom.
And when we go to octahedral, we now have 90 degree angles.
And it is actually the same through all of our angles in our ideal structure.
So trigonal pyramidal were we have some 90 and some 120 degrees,
here we are looking at all 90 degree angles.
So now what we see,
is we have the general form AX6 because we have six bonding groups.
Our electron paired geometry, and our molecular geometry are the same,
we call this octahedral, because we have all bonding groups.
The molecular geometry is also octahedral.
When we go to our next example where we start replacing a bonding group with
a non-bonding group, here we have five bonding groups and one non-bonding.
So we have our general form AX5E, our electron pair geometry is octahedral.
And our molecular geometry is now square pyramidal or square pyramid.
And so, if I look at these atoms along the bottom, they form a square.
And it's going to form a pyramid shape coming out from that central atom.
So that's our square pyramid.
I've got my lone pair here on the bottom.
And because all my angles are the same, it doesn't matter where that first lone pair
is, unlike with the trigonal bipyramidal.
We had to put it in a specific type of position, axial versus equatorial.
Here all the angles are the same, so it doesn't really matter.
However, when we go to the second low pair of electrons, so
here we have four bonding and two non-bonding.
Now it does matter where those electron groups are.
And so what we want to do is we want to maximize the distance between the bonding
groups but
we also want to maximize the distance between those lone pairs of electrons.
And the way that we can do that is by putting them on opposite sides.
So basically in this position and this position.
And so by putting the lone pairs there,
we're maximizing the distance between the lone pairs.
And as much as we can maximizing the distance between the low pairs and
the bonding groups.
So here we have octahedral electron pair geometry because we still have six groups.
But now we have square planar molecular geometry because now the four
bonding groups that remain are all forming a square.
And they're all in the same plane.
Now let's look at an example and
see if you can determine the electron pair geometry for I3 minus.
Remember, you first have to draw the Lewis structure for this molecule.
So, our electron pair geometry for I3 minus is trigonal bipyramid.
When I look at my Lewis structure I see that I have three iodine atoms.
And I have 21 electrons from the I3, plus 1 more from the minus charge,
so I've got 22 electrons.
And when I start filling in, I fill in my terminal atoms first.
So I have completed my octet on my terminal atoms and
I can now put my electrons around my central atom.
I do have a situation where I have an expanded octet, and that's okay.
Because iodine is far enough down on the periodic table for that to happen.
Remember, it's only the second row elements that cannot have
an expanded octet.
So when I look at my structure here,
I've used 2, 4, 6, 8, 10, 12, 14, 16, 18, 20, 22 electrons.
So I know I have a reasonable Lewis structure.
Both of my outer atoms have an octet of electrons and
the central atom actually has 10 electrons around it.
So when I look at my electron pair geometry
really I'm only worried about that central atom.
And here I see that I have three non-bonding groups and
two bonding groups.
And so since I have a total of five groups,
regardless of whether they're bonding or non-bonding that tells me
a trigonal bipyramid, or a trigonal pyramidal geometry.
Remember, octahedral is six groups.
Tetrahedral is four.
Trigonal planar is three.
And linear, as an electron pair geometry is two.
And bent is actually not an electron pair geometry.
Now that we know the electron pair geometry of I3 minus,
let's find the molecular geometry.
So, now that I know the electron pair geometry, I can now look for
the molecular geometry of I3 minus.
And again, I'm gonna go back to my Lewis structure that I determined for I3 minus.
So we had our terminal iodine atoms had a complete octet.
And we had three lone pairs around that central iodine.
And we'll put brackets there with a minus charge, indicating that this is an ion.
And so I want to write the general form for this molecule which is A, for
the central atom.
X2, which represents the two bonding groups, then E3,
which represents the three non-bonding groups around that central atom.
And anytime we have something with the form AX2E3,
it's going to be linear molecular geometry.
So for a trigonal bipyramid, if that were our molecular geometry we'd have AX5.
Trigonal planar would be AX3.
Bent could be a couple of things,
AX2E or AX2E2.
T-shaped comes from our trigonal by pyramidal geometry.
So, we'd have AX3E2, and Seesaw will be AX4E.
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