This course will cover the topics of a full year, two semester General Chemistry course. We will use a free on-line textbook, Concept Development Studies in Chemistry, available via Rice’s Connexions project. The fundamental concepts in the course will be introduced via the Concept Development Approach developed at Rice University. In this approach, we will develop the concepts you need to know from experimental observations and scientific reasoning rather than simply telling you the concepts and then asking you to simply memorize or apply them. So why use this approach? One reason is that most of us are inductive learners, meaning that we like to make specific observations and then generalize from there. Many of the most significant concepts in Chemistry are counter-intuitive. When we see where those concepts come from, we can more readily accept them, explain them, and apply them. A second reason is that scientific reasoning in general and Chemistry reasoning in particular are inductive processes. This Concept Development approach illustrates those reasoning processes. A third reason is that this is simply more interesting! The structure and reactions of matter are fascinating puzzles to be solved by observation and reasoning. It is more fun intellectually when we can solve those puzzles together, rather than simply have the answers to the riddles revealed at the outset. Recommended Background: The class can be taken by someone with no prior experience in chemistry. However, some prior familiarity with the basics of chemistry is desirable as we will cover some elements only briefly. For example, a prior high school chemistry class would be helpful. Suggested Readings: Readings will be assigned from the on-line textbook “Concept Development Studies in Chemistry”, available via Rice’s Connexions project. In addition, we will suggest readings from any of the standard textbooks in General Chemistry. A particularly good free on-line resource is Dickerson, Gray, and Haight, "Chemical Principles, 3rd Edition". Links to these two texts will be available in the Introduction module.
CDS 11 Molecular Geometry and Electron Domain Theory I
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Química Geral: Desenvolvimento de Conceitos e Aplicação
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Bonding and Structures in Covalent Molecules
To understand the types of compounds which can be formed and the properties of those compounds, we have to understand how atoms bond together to form molecules. In this module, we develop a model for the bonding of non-metal atoms to non-metal atoms, called a covalent bond. The model can be used to predict which combinations of atoms are stable and which are unstable. Observations of the structures of the molecules lead to a model to understand molecular geometries and properties related to those geometries. From this, we build a foundation for understanding and predicting how molecular structure is related to molecular reactivity and function.
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John Steven HutchinsonDean of Undergraduates and Professor of Chemistry
In this lecture, we're going to begin a study of the geometries that molecules
adopt. In three dimensions, what their
structures are in three dimensions. And we're going to develop a model that
accounts for those geometries, called the electron domain theory, or otherwise
referred to as the valance shell electron pair repulsion model.
This material is from the 11th Concept Development Study on Molecular Geometry,
so I'm referring you to that, as we're going to follow it fairly closely.
Let's recall. At the end of our study of the structures
of solids, we examined the structure of solid diamond, carbon, and a network
lattice that is depicted here. We asked in passing, but never answered
the question. Why do the carbon atoms arrange
themselves in the way that they do? In order to understand that we need to
back all the way back to discussions of individual molecules and ask?
What are the three dimensional structure of molecules?
To think about that, let's go back to our simplest hydrocarbon, methane, which is
carbon surrounded by four hydrogen's that is the lowest structure for the methane
molecule. Drawn in this way we might predict that
it's experimentally we would see a player molecule in which each bond angel where a
bond angel is the angel between two adjacent bonds that we might expect that
the bond angles are all 90 degrees. In fact, experimentally when we study
methane, it turns out that the bond angle is 109.5 degrees, and that is true for
each of the bond angles between adjacent carbon-hydrogen bonds.
From that we can immediately tell that this is not a planar molecule.
If it were, each bond would have to be 90 degrees.
As it stands, we have more than 360 degrees in a circular angle.
That can't be. So this molecule doesn't exist in a
plane. What does it look like?
We'll examine that in just a moment. But let's consider a variation on methane
in which we have substituted. Two of the hydrogens, four chlorines.
Well, we might say that this would be the structure of that molecule.
If I draw it this way, and if the molecule is planar, then in fact I could
draw another version of dichloromethane in which the two chlorines appear to
adjacent from each other. Rather than across from each other.
But it turns out experimentally there's only a single molecule, which is
dichloride methane, and we even commented back in our structures about little
structures that these two types of structures are, in fact, identical to
each other. They are the same structure, even though
they appear different in two dimensions, it must be true that in three dimensions
the arrangements of the atoms are exactly the same.
We actually encountered that when we discussed water before.
And our comment about water was that it appeared to be the case that we could
draw two different drawings for water. They are both drawn here, and the thought
was then that perhaps the angle was either 90 degrees or 180 degrees.
In fact there's only one form of water, and the actual bond angle is neither 90
degrees or 180. But is instead 104.5 degrees.
Likewise, just in order to complete a set, if we examine ammonia, and ask what
the bond angles are in ammonia, it turns out that the bond angle is about a 107
degrees. And right away we notice that there's a
similarity amongst these bond angles. They are not a 120, they're not a 180,
they're not 90, but they are all very much in proximity to one another.
Question is, how do we account for what the three dimensional structures are?
We need to look at those and there are a couple of different places we can look.
One place we could look, I've highlighted here as a website that's really a
marvelous website I'm going to show in just a moment.
I'm going to be showing you some geometries based upon experimental data
from a different program. you can also actually get model kits, and
I would encourage you, that you'll understand these models better if you
have three-dimensional kits in front of you, but if you don't have access you can
use these web resources, the one I will show you in just a moment.
Here's what the geometry of methane actually looks like.
We look at this, we can see that it is in fact a three dimensional molecule, it is
not planar. If we rotate it about, it has enormous
symmetry. If we look down any one of the CH bond
axes, we can see that the other three CH bonds are arranged equally separated from
one another and each of those bond angles is 109.5.
None of them are different from the other.
If I can rotate and look down any of the CH bond angles, I wind up with
essentially exactly the same depiction. What if we push a little past methane and
ask, let's replace one of the hydrogens with a carbon.
Then we would have C2H6, and that would be Ethane.
Let's look and see what Ethane looks like.
Here's the geometry of Ethane, and we notice that the arrangement of the bonds
about the Carbon atom here at the bottom that I'm pointing to, are very much the
same as the arrangement of the bonds about the Carbon atom in Methane.
Go back and look at Methane, look at Ethane.
The arrangement of the Carbon-Hydrogens, and the other bonds about the Meth-,
about the Carbon in the middle of Meth-, of Ethane, are the same as they were in
Methane. Furthermore, the two ends of the
molecules are about the same. If I flip it over and ask about the
carbon on the other end, I immediately notice that this carbon also looks like
the bonds were arranged in the way that they were arranged around methane.
We might also push a little bit farther and ask what about a larger carbon
molecule, like you keep adding one at a time, instead I'm going to took two
Ethanes, kind of shove them together, and wind up with something we call Butane and
again, if I look at each of the Carbons. Either on the ends or if I focus on one
of these middle carbons. You can see the the arrangement of the
bonds about the, each carbon atom, is about the same.
And in fact I can tell you that the bond angles in each case are either exactly
109.5 or very, very close to it. These three dimensional geometries help
us understand questions like, what did it mean before when we said that the
dichloro methanes were the same? Let's look at now in three dimensions and
see what that looks like. We'll click on my dichloromethane
molecule here. Here is methane, with two of the
hydrogens replaced by chlorines. When we look at the molecule in three
dimensions instead of two, it's clear that each pair of carbon, of hydrogen or
each pair of bonds is like identical to each other pair of bonds.
They in fact are adjacent to each other. Now the chlorine hydrogen, I'm sorry, the
chlorine carbon bonds are longer because chlorine is a much larger element than
the hydrogen carbon bonds are, because hydrogen is a much smaller element.
But the arrangement of those bonds is essentially exactly the same as it was in
methane. It is not the case that I could ever, for
example, have, thinking back to our drawings here a moment ago.
Chlorine's on the opposite side of the molecule versus chlorine's on the same
side of the molecule. If we look at the arrangements here, it
must always be the case that the chlorines are on the same side of the
molecule, no matter which two of the hydrogens I substitute for the chlorines.
That accounts for the fact that there's only a single chloromethane molecule,
dichloromethane molecule, and we see that when we look at the three dimensions.
Let's look at a couple of more molecules, I'm actually going to go to this
molecular viewing website now and take a look at it, here.
This is what it looks like when you first pop it up.
Let's actually first pull up methane over here by looking in the menu that includes
the m's. And we find methane on the list.
And we see a molecule that in fact looks very much like the methane molecule that
we've drawn over here on the right. We can also then pull up the water
molecule, since water was one of the other molecules that we examined back
over here on our pad of paper. We made the contention that it is neither
linear, nor perpendicular, and in fact you can see the experimental geometry of
the molecule represented here. I'm going to stop it from spinning.
So that I can handle it myself and you, in fact, see that the bottom angle here
is greater than 90 degrees and much less than 120 degrees.
It lands at about 109 at about 104 degrees, part way between 90 and 120.
Let's examine ammonia to sort of complete the group that we were looking at before.
On the right side here, this is what ammonia looks like.
Again, I'm going to stop it from spinning, and pick it up, and rotate it
around. And what we notice is in fact, that the
bond angles here are comparable to the kind of bond angle structures that we
were seeing from methane over here. In fact the structures look quite
similar. If I compare methane over here to ammonia
over here, they look the same except there's a hydrogen missing above.
Actually the water looks the same as well except it looks like there are two
hydrogens missing on either side. The question we will now pursue is.
Why do these structures all appear so similar?
To do that let's examine the data carefully and, and note that what we
discovered is that in each of the structures we drew, whether it was
methane, or ethane, or butane, regardless of whether we were dealing with an H-C-H
bond angle or an H-C-C bond angle or a C-C-C bond angle, as we see here in
butane, all of the angles are very close to 109.5 degrees.
When we went and looked at water, we discovered, in fact, that water has a
bond angle about 104.5, very close to the bond angle in the hydrocarbons.
And in ammonia the bond angles are 107 degrees.
Very close to the bond angle in both of the other two.
And in fact, the question is, what is it that causes this close similarity in the
experimental bond angles? To answer that question we need to build
a model that takes advantages of whatever it is these molecules have in common.
That would cause us to think that perhaps they would have comparable bond angles.
What do they have in common? Well one thing we can say, is if we
examine their Louis Dod structures, their Lewis molecular structures, we can
definitely say that in each of these atom, in each of these molecules the
central atom, whether it's carbon, hydrogen, or nitrogen.
carbon, oxygen or nitrogen has four pairs of electrons around it.
Here I've replaced the bonds with the pairs of electrons that are the actual
covalent bond. Let's take water.
And lets take ammonia and each of these circumstances, I meant to write these as
dots so that we can see expressly, the the pairs of electrons, in each case
there are four pairs of electrons around each of the central, atoms and that's
whats noted here that in each case we've got four pairs of valence electrons.
And the angles which are produced between those pairs of valence electrons, all
appear to be about a 109.5 degrees. Why would that be?
The answer could be, in fact, that recognizing that these are electron
pairs, they'll be, they'll be, they'll repel each other.
And if they maximally repel each other they should go as far apart as they can
be. And it turns out the farthest way that
they can go apart from each other is to form the corners of something called a
tetrahedron. To figure out what a tetrahedron looks
like, actually look at a couple of images of tetrahedrons from the web.
Here is actually what a solid tetrahedron looks like.
It's got four corners on it. It has four triangular faces.
The four triangular faces are a little easier to see here, than in other
locations, you can now see the one in the back.
You can see the four corners. Imagine that we put a pair of electrons
at each corner here, so that perhaps we had four bonds coming out as we do in
methane. And imagine that we put a carbon atom
right in the middle of this structure. We would wind up with something that
looks like this. In other words, optimally separating, the
four pairs of electrons from one another. Results in an arrangement, which is
tetrahedral. Meaning that there are, that it's a four
sided equilateral triangle or a triangular pyramid here, that we refer to
as a tetrahedral and often as a consequence.
We will say that methane has a tetrahedral geometry.
Let's go back and pick up the methane again here, and see that if we were to
connect the hydrogens from the corners. We would build a four sided triangular
pyramid here, called a tetrahedron. So methane is a tetrahedral molecule.
If we were to remove one of the atoms from this tetrahedral geometry and go
back and take a look again at what amonia looks like.
Imagining that the fourth pair of electrons is still up here in that
tetrahedral geometry, the remaining bonds will form about tetrahedral angles,
correspondingly then, if our model is correct, we are separating each pair, of
electrons in the valence shell from each other.
Not just the bonds, but the pair it electrons, or the unbonded electrons as
well. The result is, that the central atom sits
at the corner or sits at the center of a tetrahedron, which is a four sided
triangular pyramid. And the bond angle which is created if we
look at that tetrahedron in the geometry before, the bond angle that gets created
turns out to be precisely a 109.5 just from geometrical considerations.
The result is we have a model that we've developed now.
Which appears to predict that we will want to maximally separate the valence
shell electron pairs. And we will call that now, the valence
shell electron pair model. In which the geometry of the molecule is
determined by maximally separating the valence shell electron pairs.
We could test that model by vari, considering a variety of different kinds
of molecules. One of the molecules we might examine,
for example, could be sulfur hexafluoride.
The Lewis structure for sulfur hexafluoride is a pretty straightforward
one. Each fluorine is directly coupled to,
bonded to, by a pair of electrons to the sulfur atom in the middle.
That means there are six pairs of electrons in the valence shell of the
sulfur atom. Clearly that violates the octet rule that
we have considered before. What is the geometry of Sulfur
Hexa-Flouride look like? We will go back and take a look at our
molecular viewer where we can actually pull down Sulfur.
Hexaflouride, hexachloride, I'm sorry. This is Sulfur Hexachloride instead of
SulfurHexaflouride but the geometry is exactly the same.
Stop this from spinning and notice that what we have done is to put the six pairs
of electrons, sort of one in the positive X, one in the negative X, one in the
positive Z, one in the negative Z. One of the positive Y and one of the
negative Y and that is the optimal way in which we can separate six pairs of
electrons. Let's consider another molecule.
Go back and look at our presentation here, consider PCL5.
How might we optimally separate five, pairs of electrons, about, the central,
phosphorous, to which it is bonded. Let's go back an actually experimentally
look at what the data suggests to us would be the case.
We'll look now for phosphorous pentachloride over here.
let's see. Phosphorous pentachloride is here, an
what we wind up with, is a molecule that looks something like a child's jack, in
the game of jacks. In this case, the maximum separation of
five pairs, will put three in an equilateral triangle on the equator of
this molecule. And one at the north pole, and one at the
south pole. The bond angles then become a 120 for
those which are in the equator. And 90 between the bonds in the equator,
and the bonds on the poles. Let's try one last one from the slideshow
back here. Which is boron trifluoride.
We might think that that would look a lot ammonia as before but in fact there's no
lone pair on the boron as there was on the nitrogen up here on ammonia.
So there's three pairs of electrons in the valence shell.
Correspondingly we would probably guess that they would separate to the corners
of. And equilateral triangle.
Let's come up and take a look and see what boron trifluoride looks like.
Boron trifluoride is here, and in fact what we discover is that it is, in fact,
an equilateral planar triangle, with all bond angles being about a 120 degrees.
What we conclude from all of this. Is that we have a pretty good model now
developed that would count for the geometries of molecules which are both
tetrahedral. As well as molecules that fit other
geometries. Dependent upon the number of valence
shell electron pairs they have. This model turns out to have great power.
But we're going to have to extend it a bit to account for a few other
considerations including what happens when we have double bonds and triple
bonds. We'll pick that up in the next lecture.
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