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 II
Loading...
來自 Rice University 的課程
普通化学：基本概念与应用
238 個評分
從本節課中
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.
與講師見面
John Steven HutchinsonDean of Undergraduates and Professor of Chemistry
In this lecture were going to extend upon the Valence-Shell Electron Propulsion
model that we developed in the previous lecture and were going to apply it to
some specific properties of poly atomic molecules.
You'll recall that we've developed this model now by which we believe that the
geometry of the molecule was determined by separating as far as we can, the
electrons which are in pairs in the valence shell of the electron, of the
central atom. So over here, we have the methane
molecule and we've separated the four pairs of electrons which are in the
carbon-hydrogen bonds, and when we maximally separate them we get this
tetrahedral arrangement back over here. But recall, even though we don't have
four bonds in the water molecule or in the ammonium molecule We get an
arrangement or bond angles which are comparable in all three of these
molecules. What that says is we are not optimally
separating the bonds. We are optimally separating the pairs of
electrons because in all three of the mo the molecules on the screen here, there
are four pairs of valence electrons around the central atom.
Oxygen in this case, nitrogen in this case, and carbon in this case, and as a
result we wind up with a very similar geometries.
So, in our valence shell model it's important to note that bonded pairs,
bonded pairs of electrons and lone pairs of electrons have exactly the same effect
on the arrangement of the electron pairs in terms of determining what the geometry
of the molecule is going to be. Let's test this a little bit further
though by considering two other molecules.
Let's compare our ethane molecule which we have drawn back over here.
Which is C2H6. If we draw the Lewis structure for C2H6,
it's straightforward enough. Each carbon has three hydrogens
surrounding it. Each carbon, therefore has four pairs of
electrons around it. We expect a tetrahedral geometry, which
is exactly what we observe here. If we consider the Lewis structure for
C2H4, which is Ethene. That Lewis structure, recall, has a
double bond between the two carbons. But each carbon still has four pairs of
electrons around the valence shell. Here's the Lewis structure that I was
drawing. C2H4 has a double bond between the two
carbons. As a consequence, what we might want to
do is look and see, do we get a tetrahedral arrangement?
Do we get bond angles that are about 109.5?
And the answer turns out to be no. As we can see here from our data the bond
angle is actually closer to 120 degrees. In fact let's take a look at what the
molecules actually look at, look like. Here is the Ethene molecule and you can
see a couple of interesting things that distinguish it from the Ethane molecule.
One is that these bond angles are in fact, about a 120 degrees, the other is
its a perfectly planar molecule, all six atoms rest in the same plain, that's
quite different than the Ethane molecule were the molecules can only be viewed in
three dimensions, go back and look at Ethene here[SOUND] .
It's apparent that having the double bond changes the geometry here.
Let's also con consider, Ethyne, which is C2H2, the Lewis structure for that you
may recall from our earlier work, has a triple bond between the carbon atom so we
might draw it something like that. And if we look at the molecular geometry
for Ethyne, we discover the bond angle is around 180 degrees.
So, here we got about 120. Here we get about 180.
In fact, we get exactly 180. If we look at the molecular geometry of
the Ethyne molecule, you can see it here, it's a perfectly linear molecule,
straight down the line. And, you can see the bond angles, which
are there. What does that suggest?
So, in each case, we can see, there are four pairs of electrons around each
central carbon, and all three of the molecules, here.
In all three molecules, Each carbons got four pairs electrons around it.
And yet, the geometries of the molecules are very different.
It must be the case, that something has changed, and what's changed, of course,
is the appearance of a double bond or a triple bond.
Apparently, each double bond or triple bond, behaves differently, then a
separated pairs of electrons. But that actually makes a lot a sense,
because, I would not expect to be able to take this double bond and the ethene
molecule and separate these pairs of electrons.
I wouldn't be able to form a double bond if I did.
Both of these pairs must be oriented towards the, two carbons together.
The fact that we wind up with about 120 degree bond angle.
Suggest then that the three sets of electron pairs about this carbon atom
must be about optimally separated from each other.
That suggest that we treat a double bond or for that matter a triple bond not as
two pairs of electrons but rather electron as an electron domain.
So we define something new now as a region of electrons participating either
in bonding, or as lone pairs. In these arrangements, according to what
we've looked at before, a single bond counts as an electron domain to be
repelled by other electron domains, but a lone pair counts similarly the same way.
That would have been for example, in water or in ammonia.
But it's clear by looking at the data that we have over here on the side, that
we have to treat the double bond essentially as exactly the same as a
single bond, or as a lone pair, for purposes of determining the geometry of
the molecule. So what that means is, a double bond
counts as a single electron domain because I can't separate the electron
pairs in a single domain because a double bond is in fact, two pairs of electrons
in the same geometric location in the vicinity of the two carbon atoms.
And likewise, a triple bond must count as a single electron domain, as well.
This is why when sometimes we would call this Electron Domain Theory instead of
just Valence-Shell Electron Pair Theory because it tells us that we have to
examine individual electron domains and not just the numbers of bonds or the
numbers of pairs of electrons. Now, having refined the model in that
way, we'll consider one further refinement, which has to do with the fact
that the angles, in fact, are not, in fact, entirely precise, that we get the
tetrahedral geometry in each case. We'll go back and look at the water
molecule, for example. The water molecule has geometry something
like that. Two pairs of electrons about this.
This bond angle here is about 104.5 degrees.
But we remember that the tetrahedral angle [SOUND] would have predicted that
that bond angle should be 109.5 degrees. We get what is sometimes is then called
the distortion. We don't exactly get the tetrahedral
angle that we were expecting. We get a slight variation from it.
If we look at the bond angles up here in in ethane, we actually get very, very,
very close to the 109.5 degree angle but, we get at least a little bit variation
away from that when we look at the water molecule.
Notice that this number is smaller than the 109.5 that we would expect.
That appears to be the case because, the lone pairs on the opposite side of the
oxygen here, somehow or another must have bond angles away from, these OH bond that
must be larger. Now that's not a well defined quality, we
don't have a bond direction over here because there's no atom here to point to.
But what it does say is that these electron pairs here apparently are
creating a greater repulsion than the bonded pairs are creating.
So, what we have observed then here is that the angle between the bonded pairs
is smaller than the angle that we would expect between the bonded pairs and the
non-bonded pairs, even though that latter angle is not a well defined quantity.
It must be true because we've in fact sort of pinched this bond down.
Our model then becomes that the lone pairs in this case have created a greater
repulsion for each other and for the bonded pairs then the bonded pairs have
created to each other, and that's a slight modification we can use.
To make predictions about the molecular geometries.
Let's now actually apply all of this to make some predictions about some
molecular properties. What we're going to start with is by
going back and examining this concept of dipole moments that we studied earlier
that led to our concept of electronegativity.
What we observed is that there is a molecular dipole, meaning that there is a
positive end. To one end of the molecule, a negative
end to the one, other end of the molecule, that results from unequal
sharing of the electrons between the bonded pair.
Taking a particular example here of the Hydrogen Fluoride molecule, where the
Hydrogen is the positive end and the Fluoride is the negative end, we
accounted for the appearance of that dipole as resulting from the differences
in electronegativity between the hydrogen atom and the Florine atom.
The Florine atom attracts the electrons to itself more strongly, in large measure
because of it's greater core charge. And that makes it possible then for there
to be net in the molecule, a negative end of a molecule, and a positive end to the
molecule. So a molecular dipole moment exists.
When there's a positive end and a negative end in the molecule which can be
identified. We only looked at diatomic molecules when
we considered dipole moments before. So, now we're going to consider dipole
moments in polyatomic molecules. When we do that experimentally, what we
discover is, for example, that carbon tetrafluoride.
Is a non-polar molecule. Now that seems a little strange.
Because certainly it must be the case, [SOUND] if we examine the structure back
over here, that thinking about the Florine, relative to the carbons, that
the Florine is certainly much more electronegative than would anticipated to
have a negative charge to it. And we would expect each of the Florines
to have such a negative charge. And we would expect the carbon in the
middle to have something even larger in terms of a positive charge in the middle
to offset the negative charge of the Florines which surround it.
And yet, it turns out that carbon tetrafluoride is non-polar.
The question is, why does it come out to be non-polar?
And the answer has to do with the geometry of the carbon tetrafluoride.
We go back to our website from before, and examine the geometry of carbon
tetrafluoride. Let's see here, carbon tetrachloride is
here. We see that it actually looks very, very
much like, the methane molecule. Now as we stare at that molecule it's
worth asking, is there a negative end to this molecule and is there a positive
end? As we examine the data over here, it
would appear that each of the Florines is negative, but notice that every side of
this molecule, regardless of where I point it must be equally negative.
So as I've arranged it like this, this Florine here must be negative, but the
other end must also be negative. And in fact, I can't say that this end of
the molecule is negative and the other end is positive, because if I rotate the
molecule, that Florine must just be as negative as the previous Florine was.
Hence, there is no location in this molecule where I can say the end, one end
of the molecule is negative, and the other end of the molecule is positive.
What must be true is that although there are individual bonded type holes[SOUND] .
That each of those bond dipoles must cancel each other out[SOUND] .
Meaning, that net, from a distance, there does not appear to be a dipole moment
where there is a positive end to the molecule and a negative end to the
molecule. The contrast H2O, is polar.
Let's see if we can figure out why. Well, we'll recall that the molecule is
bent, it is not linear. This couple of lone pairs here that
create that bend. The angle here, is about 104.5 degrees,
we'll recall. We'll remember that oxygen is more
electronegative than the hydrogen. And as a consequence each of the
hydrogen's must have some kind of a positive change and the oxygen has a
negative charge. The result of that would be that from a
distance I can clearly see a negative end of this molecule and a positive end of
this molecule as well. As a consequence, there is a molecular
dipole moment for water. How about carbon dioxide?
How could we account for the non-polar nature of carbon dioxide, we draw the
Lewis structure for carbon dioxide. We can make a prediction of what its
molecular geometry will be, here is the Lewis structure.
Notice that there are two electron domains.
With 2 electron domains, we would predict, that carbon dioxide is a linear
molecule. Let's go and check and make sure that
that is true. We'll look at the molecular geometry of
carbon dioxide. And in fact, when we run it, we see that
it is a linear molecule here. We can see the double bonds in the
molecule as well. And what we would notice as we look at
this, is although, we would expect that each oxygen carries a negative charge,
and the carbon carries a positive charge, from a distance it's clear, that there is
no negative end of this molecule, and there is no positive end of this
molecule. As a consequence, when we are examining
molecules to determine whether they are polar or not, it's insufficient to look
at the electronegativity differences of the atoms.
Because they will generate individual bond dipoles.
But the three dimensional geometry's is what's going to determine, whether there
is a polar, a, a dipole moment in the molecule, and therefore, whether there's
any molecular polarity. This is one powerful example of how
knowing the geometry of the molecule, the arrangement of the atoms spatially in
three dimensions. Can help us make predictions about the
physical properties of individual molecules.
