Chemistry Exam > Chemistry Videos > Inorganic Chemistry > Common Mistakes in VSEPR Theory - Chemical Bonding
Common Mistakes in VSEPR Theory - Chemical Bonding Video Lecture | Inorganic Chemistry
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FAQs on Common Mistakes in VSEPR Theory - Chemical Bonding Video Lecture - Inorganic Chemistry
| 1. What is VSEPR theory and how does it relate to chemical bonding? |  |
Ans. VSEPR theory (Valence Shell Electron Pair Repulsion theory) is a model used to predict the shape of molecules based on the repulsion between electron pairs in the valence shell of atoms. It explains the three-dimensional arrangement of atoms in a molecule and how this arrangement affects the chemical properties and reactivity of the molecule.
| 2. What are some common mistakes made when applying VSEPR theory? | |
Ans. Some common mistakes made when applying VSEPR theory are:
1. Neglecting lone pairs: Many students overlook the presence of lone pairs of electrons in a molecule and only consider the positions of bonded atoms. However, lone pairs also affect the molecular geometry and should be taken into account.
2. Incorrect counting of electron pairs: Some students may miscount the number of electron pairs, leading to incorrect predictions of molecular shapes. It is important to correctly identify all the electron pairs, including both bonding pairs and lone pairs.
3. Ignoring multiple bonds: Double or triple bonds are often overlooked when determining the molecular geometry. Each bond, whether single, double, or triple, contributes to the overall shape of the molecule.
4. Confusing electron pair geometry with molecular geometry: Electron pair geometry refers to the arrangement of all electron pairs, including both bonding pairs and lone pairs. Molecular geometry, on the other hand, considers only the positions of the atoms. It is crucial to differentiate between the two and apply the correct geometry when predicting molecular shapes.
5. Overgeneralizing bond angles: While VSEPR theory provides a good estimate of bond angles, it is important to recognize that bond angles can deviate from the ideal values due to various factors such as lone pairs and electronegativity differences. Overgeneralizing bond angles can lead to inaccurate predictions of molecular shapes.
| 3. How does VSEPR theory explain the shapes of molecules? | |
Ans. VSEPR theory explains the shapes of molecules based on the repulsion between electron pairs in the valence shell of atoms. According to the theory, electron pairs, whether bonding pairs or lone pairs, repel each other and strive to be as far apart as possible. This results in specific arrangements of atoms in a molecule, which determine its shape.
The number of electron pairs around the central atom determines the electron pair geometry. For example, if there are two electron pairs, the electron pair geometry will be linear. If there are three electron pairs, the electron pair geometry will be trigonal planar, and so on.
The molecular geometry, on the other hand, considers only the positions of the atoms and ignores the lone pairs. It is derived from the electron pair geometry by removing the lone pairs from the equation. For example, if there are two bonding pairs and one lone pair, the molecular geometry will be bent or angular.
By applying VSEPR theory, one can predict the shapes of molecules and understand how the arrangement of atoms affects the chemical properties and reactivity of the molecule.
| 4. Can VSEPR theory accurately predict the bond angles in all molecules? | |
Ans. While VSEPR theory provides a good estimate of bond angles in many molecules, there are cases where it may not accurately predict the exact bond angles. This is because bond angles can be influenced by various factors, including lone pairs, electronegativity differences, and molecular distortions.
Lone pairs of electrons exert a stronger repulsion compared to bonding pairs, causing a deviation from the ideal bond angles predicted by VSEPR theory. For example, in a water molecule (H2O), the ideal bond angle is 109.5 degrees, but due to the presence of two lone pairs on the oxygen atom, the actual bond angle is approximately 104.5 degrees.
Electronegativity differences between atoms can also affect bond angles. For example, in a molecule like ammonia (NH3), the nitrogen atom is more electronegative than the hydrogen atoms, resulting in a smaller bond angle of approximately 107 degrees compared to the ideal tetrahedral angle of 109.5 degrees.
Additionally, molecular distortions caused by factors such as steric hindrance or repulsion between bulky groups can lead to deviations from the ideal bond angles.
Therefore, while VSEPR theory provides a useful framework for predicting molecular shapes and bond angles, it is important to consider these additional factors that can influence the actual angles observed in specific molecules.
| 5. How can one avoid common mistakes when applying VSEPR theory? | |
Ans. To avoid common mistakes when applying VSEPR theory, it is important to:
1. Consider all electron pairs: Recognize and consider the presence of both bonding pairs and lone pairs of electrons in a molecule. Lone pairs have a significant influence on molecular geometry and should not be overlooked.
2. Count electron pairs accurately: Ensure that the correct number of electron pairs is counted when determining the electron pair geometry and molecular geometry of a molecule. Miscounting can lead to incorrect predictions of molecular shapes.
3. Account for multiple bonds: Take into account the presence of double or triple bonds when predicting molecular shapes. Each bond, regardless of its multiplicity, contributes to the overall shape of the molecule.
4. Differentiate electron pair geometry and molecular geometry: Understand the distinction between electron pair geometry (which considers all electron pairs) and molecular geometry (which considers only the positions of the atoms). Apply the correct geometry when predicting molecular shapes.
5. Recognize deviations from ideal bond angles: Understand that bond angles can deviate from the ideal values due to factors such as lone pairs and electronegativity differences. Avoid overgeneralizing bond angles and consider these factors when predicting bond angles in specific molecules.
Text Transcript from Video
Now if you're just learning VSEPR, it can
be really hard and chances are you might be
making some very common mistakes, that I'm
going to talk about in this video.
Okay.
Here's how it's going to work.
I'm going to go through a bunch of common
mistakes that people make, and for each mistake
I'm going to say why it's a mistake, I'm going
to say what's wrong about it and I'm going
to show you what the right answer is.
I want you to pay close attention to each
one of these mistakes even if you're not making
them now; because understanding the mistakes
is one of the best ways to learn and make
sure that you really understand something.
Okay.
So here's the first example.
NH4 -- I'm sorry -- NH3, for ammonia, what
does it look like in three dimensions?
What's it's VSEPR shape?
Okay.
So one of the biggest mistakes is people look
at this molecule and they think that it's
trigonal planar.
Okay.
A trigonal planar molecule looks like this.
It's got these three atoms around the outside.
And look, all these atoms are flat in a plane.
They're all straight in this plane, and that's
why you call it planar for this plane.
Okay.
Now I think a lot of people make this mistake
because they look at the structure of NH3
and they say okay.
There's one atom in the center and then there
are three atoms around the outside, these
three atoms arranged on the outside.
So here's the atom in the center, and here
are the three atoms around the outside.
I think that's why they come up with this
wrong idea here, but here's why it's wrong.
Okay.
Because check out this.
Look at this unshared electron pair here.
It's very important.
When people think this is trigonal planar,
they forget to take into account the unshared
electron pair.
Okay.
So there are actually four things that need
to be arranged around this [inaudible].
Sure there are three atoms, but then there's
also unshared an electron pair.
So this structure, NH3, actually looks like
this, where you have the three atoms, but
then we have an unshared electron pair here.
Each one of these four things wants to push
against the other and they want to be as far
apart as possible.
Okay.
So this unshared electron pair it's pushing
the atoms down, and that's why it gives it
this shape.
We call this shape trigonal pyramidal because
it looks kind of like a pyramid.
You can particularly see that if I hold it
on its side, you can see that all of the atoms
are bent down this direction, and that's very
different from the trigonal planar where they're
all flat on this plane.
Trigonal pyramidal is bent down because this
unshared electron pair up here is pushing
these atoms down.
Sometimes when you're talking about VSEPR
shapes, you need to know the angles between
the bonds.
So if I have my Lewis structure here, the
angles between any of these bonds are about
107 degrees.
Okay.
Now finally, if we're saying that NH3 is not
a trigonal planar molecule, what is?
Here are two examples.
This one and this one are both trigonal planar.
Look what they have in common.
They have one central atom and then they have
three atoms around the outside, but they don't
have any unshared electron pairs.
Okay.
It's this unshared electron pair here on nitrogen.
That is what makes this trigonal pyrimidal
instead of trigonal planer.
So always keep an eye out for these unshared
electron pairs.
If something doesn't have it, it's trigonal
planer in this example.
If it does, it's going to be trigonal pyrimidal.
Okay.
Let's look at some more.
Okay.
SH2.
A big common mistake that people make is they
look at this molecule and they say it has
this shape.
It's a linear molecule.
And I think they do that because they see
one atom in the middle and two atoms on either
sides.
So they say okay.
A linear molecule.
It makes sense.
Here is one atom in the middle and two atoms
on either side; but the mistake they're making
here is they're forgetting to look at this
and this, which are the unshared electrons
pairs.
These unshared electron pairs want to arrange
themselves around the central atom just the
way these other atoms do.
Okay.
So the molecule is actually going to look
like this.
Here are the two atoms, but now here are the
two unshared electron pairs, okay, and they
are pushing against each other and they're
pushing against these atoms.
All four of these things want to be as far
apart as possible.
So it will push these two atoms into this
sort of bent shape.
And we actually call this molecule -- we call
it a bent molecule.
It's these unshared electron pairs on the
central atom on S here that are bending these
other two atoms down.
Okay.
Now something that I think confuses people
is the way this Lewis structure is drawn.
They say but it's drawn as if all the atoms
are in a row.
Okay.
Well, there are a lot of correct ways to draw
a Lewis structure.
You don't always have to draw a Lewis structure
to show what the actual shape of the molecule
would be like.
So don't get thrown off by something like
this.
But we could draw a Lewis structure like this
that shows the shape a little bit more accurately.
And if we were doing this, we might want to
show the angle between these bonds here.
And this angle would be about 105 degrees.
So a bent molecule like this with two unshared
electron pairs and two atoms will have about
105 degrees between them.
Okay.
Now, as I did before, let's talk about what
molecules would look like that really were
linear, okay, if this isn't a linear molecule.
These are two examples of linear molecules.
Let me move this up here.
And what you see here is they have one central
atom and two atoms on either side but there
are no unshared electron pairs on the central
atom.
So all they have is these two atoms on either
side and no unshared electron pairs.
That's how you know that they're actually
going to be linear.
Let's look at one more example.
This example, like the previous one, people
look at it.
This is ozone, 03.
You know, it's up in the atmosphere.
It makes the ozone layer, O3.
They look at it and, just like before, they
see it and they think I bet this is a linear
molecule, three atoms, that they're all in
a row like this.
But as I keep saying, you're probably so sick
of hearing me say it, is look at this unshared
electron pair.
Okay.
This unshared electron pair is going to push
on these other two atoms and it's going to
make a shape like this.
Here are the two atoms.
Here's the unshared electron pair.
And so this is also going to make what we
call a bent molecule because of the push from
the unshared electron pairs here.
Now it turns out that the angle is a little
bit different when you only have one unshared
electron pair pushing.
We saw a minute ago -- where did I put that?
-- that when you have two unshared electron
pairs pushing, you have an angle of about
105 degrees.
It turns out that if I were to draw this Lewis
structure so that we could actually see the
shape a little bit better -- hang in here
with me as I draw it -- the angle between
these two bonds would be about 116 degrees,
kind of between 116 and 118, a little less
than 120.
So this angle is going to be a little bit
bigger than when we have two unshared electrons
pairs.
Okay.
But either way keep in mind this unshared
electron pair here makes it a bent molecule.
It's not a linear molecule.
Okay.
So like the big common theme here is that
it's all about unshared electron pairs.
Don't just look at the atoms that are surrounding
the central atom, also have a good look at
the unshared electron pairs that are on the
central atom and then figure out how they're
going to be pushing the other atoms in this
molecule.
The last thing that I want to say is I want
to give a shout out to the importance of Lewis
structures.
As you can see, VSEPR relies on Lewis structure.
Okay.
You got to look at the Lewis structure to
figure out what the shape would be.
So that means that if the Lewis structure
is wrong, the VSEPR shape is going to be wrong
too.
So even before you start thinking about what
the VSPER shape is make sure that your Lewis
structure is right or everything is going
to be thrown off.
There are a bunch of videos that you can find
that talk about how to draw Lewis structures
and then how to check your work and make sure
the Lewis structures are correct.
So always do that even before you start thinking
about the VSEPR shape, but remember don't
forget about the unshared electron pairs because
that is probably what's going to throw you
off.
be really hard and chances are you might be
making some very common mistakes, that I'm
going to talk about in this video.
Okay.
Here's how it's going to work.
I'm going to go through a bunch of common
mistakes that people make, and for each mistake
I'm going to say why it's a mistake, I'm going
to say what's wrong about it and I'm going
to show you what the right answer is.
I want you to pay close attention to each
one of these mistakes even if you're not making
them now; because understanding the mistakes
is one of the best ways to learn and make
sure that you really understand something.
Okay.
So here's the first example.
NH4 -- I'm sorry -- NH3, for ammonia, what
does it look like in three dimensions?
What's it's VSEPR shape?
Okay.
So one of the biggest mistakes is people look
at this molecule and they think that it's
trigonal planar.
Okay.
A trigonal planar molecule looks like this.
It's got these three atoms around the outside.
And look, all these atoms are flat in a plane.
They're all straight in this plane, and that's
why you call it planar for this plane.
Okay.
Now I think a lot of people make this mistake
because they look at the structure of NH3
and they say okay.
There's one atom in the center and then there
are three atoms around the outside, these
three atoms arranged on the outside.
So here's the atom in the center, and here
are the three atoms around the outside.
I think that's why they come up with this
wrong idea here, but here's why it's wrong.
Okay.
Because check out this.
Look at this unshared electron pair here.
It's very important.
When people think this is trigonal planar,
they forget to take into account the unshared
electron pair.
Okay.
So there are actually four things that need
to be arranged around this [inaudible].
Sure there are three atoms, but then there's
also unshared an electron pair.
So this structure, NH3, actually looks like
this, where you have the three atoms, but
then we have an unshared electron pair here.
Each one of these four things wants to push
against the other and they want to be as far
apart as possible.
Okay.
So this unshared electron pair it's pushing
the atoms down, and that's why it gives it
this shape.
We call this shape trigonal pyramidal because
it looks kind of like a pyramid.
You can particularly see that if I hold it
on its side, you can see that all of the atoms
are bent down this direction, and that's very
different from the trigonal planar where they're
all flat on this plane.
Trigonal pyramidal is bent down because this
unshared electron pair up here is pushing
these atoms down.
Sometimes when you're talking about VSEPR
shapes, you need to know the angles between
the bonds.
So if I have my Lewis structure here, the
angles between any of these bonds are about
107 degrees.
Okay.
Now finally, if we're saying that NH3 is not
a trigonal planar molecule, what is?
Here are two examples.
This one and this one are both trigonal planar.
Look what they have in common.
They have one central atom and then they have
three atoms around the outside, but they don't
have any unshared electron pairs.
Okay.
It's this unshared electron pair here on nitrogen.
That is what makes this trigonal pyrimidal
instead of trigonal planer.
So always keep an eye out for these unshared
electron pairs.
If something doesn't have it, it's trigonal
planer in this example.
If it does, it's going to be trigonal pyrimidal.
Okay.
Let's look at some more.
Okay.
SH2.
A big common mistake that people make is they
look at this molecule and they say it has
this shape.
It's a linear molecule.
And I think they do that because they see
one atom in the middle and two atoms on either
sides.
So they say okay.
A linear molecule.
It makes sense.
Here is one atom in the middle and two atoms
on either side; but the mistake they're making
here is they're forgetting to look at this
and this, which are the unshared electrons
pairs.
These unshared electron pairs want to arrange
themselves around the central atom just the
way these other atoms do.
Okay.
So the molecule is actually going to look
like this.
Here are the two atoms, but now here are the
two unshared electron pairs, okay, and they
are pushing against each other and they're
pushing against these atoms.
All four of these things want to be as far
apart as possible.
So it will push these two atoms into this
sort of bent shape.
And we actually call this molecule -- we call
it a bent molecule.
It's these unshared electron pairs on the
central atom on S here that are bending these
other two atoms down.
Okay.
Now something that I think confuses people
is the way this Lewis structure is drawn.
They say but it's drawn as if all the atoms
are in a row.
Okay.
Well, there are a lot of correct ways to draw
a Lewis structure.
You don't always have to draw a Lewis structure
to show what the actual shape of the molecule
would be like.
So don't get thrown off by something like
this.
But we could draw a Lewis structure like this
that shows the shape a little bit more accurately.
And if we were doing this, we might want to
show the angle between these bonds here.
And this angle would be about 105 degrees.
So a bent molecule like this with two unshared
electron pairs and two atoms will have about
105 degrees between them.
Okay.
Now, as I did before, let's talk about what
molecules would look like that really were
linear, okay, if this isn't a linear molecule.
These are two examples of linear molecules.
Let me move this up here.
And what you see here is they have one central
atom and two atoms on either side but there
are no unshared electron pairs on the central
atom.
So all they have is these two atoms on either
side and no unshared electron pairs.
That's how you know that they're actually
going to be linear.
Let's look at one more example.
This example, like the previous one, people
look at it.
This is ozone, 03.
You know, it's up in the atmosphere.
It makes the ozone layer, O3.
They look at it and, just like before, they
see it and they think I bet this is a linear
molecule, three atoms, that they're all in
a row like this.
But as I keep saying, you're probably so sick
of hearing me say it, is look at this unshared
electron pair.
Okay.
This unshared electron pair is going to push
on these other two atoms and it's going to
make a shape like this.
Here are the two atoms.
Here's the unshared electron pair.
And so this is also going to make what we
call a bent molecule because of the push from
the unshared electron pairs here.
Now it turns out that the angle is a little
bit different when you only have one unshared
electron pair pushing.
We saw a minute ago -- where did I put that?
-- that when you have two unshared electron
pairs pushing, you have an angle of about
105 degrees.
It turns out that if I were to draw this Lewis
structure so that we could actually see the
shape a little bit better -- hang in here
with me as I draw it -- the angle between
these two bonds would be about 116 degrees,
kind of between 116 and 118, a little less
than 120.
So this angle is going to be a little bit
bigger than when we have two unshared electrons
pairs.
Okay.
But either way keep in mind this unshared
electron pair here makes it a bent molecule.
It's not a linear molecule.
Okay.
So like the big common theme here is that
it's all about unshared electron pairs.
Don't just look at the atoms that are surrounding
the central atom, also have a good look at
the unshared electron pairs that are on the
central atom and then figure out how they're
going to be pushing the other atoms in this
molecule.
The last thing that I want to say is I want
to give a shout out to the importance of Lewis
structures.
As you can see, VSEPR relies on Lewis structure.
Okay.
You got to look at the Lewis structure to
figure out what the shape would be.
So that means that if the Lewis structure
is wrong, the VSEPR shape is going to be wrong
too.
So even before you start thinking about what
the VSPER shape is make sure that your Lewis
structure is right or everything is going
to be thrown off.
There are a bunch of videos that you can find
that talk about how to draw Lewis structures
and then how to check your work and make sure
the Lewis structures are correct.
So always do that even before you start thinking
about the VSEPR shape, but remember don't
forget about the unshared electron pairs because
that is probably what's going to throw you
off.
About this Video
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Oct 27, 2024 Last updated |
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