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>> Hi guys.
I'm Jackie.
I'm a post-grad student here at Curtin.
We hear you've been having some issues with bonding
and polarity, so we're here at the new Curtin Resources
and Chemistry Precinct, where we're going to be talking
to Professor Mark Buntine, who's an expert on that kind of stuff,
and hopefully that will help some problems
that you guys have been having.
>> Good day, guys.
My name is Professor Mark Buntine,
and I work at the chemistry department at Curtin University.
One of my jobs is to teach first-year chemistry students
about theories of chemical bonding,
and are recording this today to help you guys a little bit
in the bonding theories that you need to understand
in your year-12 curriculum.
Many students get confused with bonding theories in chemistry
because they think that the different theories are
indeed different.
But the important point to understand is that all
of the bonding theories are underpinned
with the principles of electrostatics.
We have positively charged nuclei consisting of protons
and neutrons, and we have negatively charged electrons.
And it's the interaction between the positive
and the negative-charged species that underpin all
of the bonding theories that you need to understand
in your year-12 curriculum.
I'm going to start today by talking
about some covalent bonding, and I will then go into some
of the other theories.
And you will see, as we progress,
that it's the same underlying electrostatic principles
that will apply to everything.
The most commonly found form of bonding
and the most commonly description
that we have is called covalent bonding.
And you would know that term in terms
of sharing of electron density.
One of the analogies that I like to use is to think
of the nuclei, the atoms, as being the bricks
in a brick wall, and the electrons in a covalent sense
of the cement or the mortar that holds those bricks together.
If you don't have the cement and you just stack the bricks up,
they're not very stable and they will fall apart.
So we have a covalent bond consisting
of positively charged nuclei
with the negatively charged electrons shared between them,
holding the nuclei together.
And the overall stability of the molecules comes from the balance
between the attraction between the negative electrons
and the positive nuclei, coupled with the repulsion
of the positive nuclei and the repulsion
of the negative electrons themselves.
Those three terms coupled together give the overall
molecular stability, but it's the attraction
between the electrons and the nuclei
that drive the overall molecular shape and structure.
In most cases, the sharing of electron density between atoms
in a molecule is not even.
That is because most chemical bonds involve different atoms
to make up that constituent bond.
As chemists, we can understand how the sharing
of electron density is unequal
by invoking a concept knock as electronegativity.
Electronegativity is a concept that gives an indication
of the ability of an atom to draw electron density to itself.
And in the electronegativity concept, the scale increases
from the lower left to the upper right of the periodic table.
Therefore, fluorine is the most electronegative atom
in the periodic table, and elements in the first period,
including sodium and the like, are the least electronegative.
As a result, in a small molecule like hydrogen fluoride or HF,
the fluorine atom draws the electron density towards itself,
away from the hydrogen, which is electronegative.
As such, more of the electron density is on the fluorine atom,
and we have a slightly negative end of the molecule
and a slightly positive end of the molecule
where the electron density is less.
This is the concept of a polar bond, which I'll talk
about again a little bit later on.
So what happens when the unequal sharing of the electrons
between atoms in a molecule becomes extreme?
For example, the halogens on the right-hand side
of the periodic table are extremely electronegative
and will accept electron density readily.
On the left-hand side of the periodic table,
the alkali metals and the alkali earths are electropositive,
and they are perfectly happy to give up their electron density.
Under these circumstances, rather than talking
about the unequal sharing of electron density,
it becomes more appropriate to talk about the formal transfer
of an electron from the electropositive
to the electronegative species.
So for example, a sodium atom gives up an electron
to become a sodium cation.
And a halogen, for example, a chlorine atom, readily accepts
that electron to become a chloride and iron.
This alternating three-dimensional array
of positive and negative charges is described as ionic bonding.
And because we come back to the idea
of the electrostatic attraction between positive charges
and negative charges, we have very strong attractions
between the cations and anions.
And so, the bonds between the atoms are very strong,
and that is manifest as the physical property
of ionic compounds, in terms of having very high melting points.
So far, I have been focusing this presentation
on covalent bonding theories for individual molecules.
If you recall back to my earlier discussion of ionic bonding,
we a have a three-dimensional network
of alternating positive and negative charges.
For example, the sodium chloride.
Well, in some cases, we can have covalent bonding that extends
into networks between an individual molecule.
And some examples that you can see,
if you look into your textbooks,
would be extended carbon covalent bonding networks.
And carbon is a really interesting case.
If the carbon bonds
in two-dimensional shapes, we get graphite.
And so, we have strong covalent bonding within a sheet,
but no covalent bonding between sheets.
And therefore, the physician manifestation of graphite is
that those sheets can sheer away, and we can use a pencil,
for example, to write on the page.
However, if the arrangement of electron density
around each carbon atom is tetrahedral,
that means there are four regions of electron density,
then we can have a three-dimensional continuous
covalent bonding network for carbon.
And under those circumstances,
we have several examples available to us.
One of them is to have diamond.
Diamond will just have this three-dimensional
covalent network.
And diamond is very strong.
It's very hard to overcome those bonds.
More recently, we've found other variations
of carbon-only bonding, and that can be a combination
of tetrahedral and trigonal planar bonding
around the carbon atoms.
And we can get carbon nanotubes, and we can get buckyballs
and other sorts of structures
that you'll see in your textbooks.
Another example of extended three-dimensional covalent
bonding networks can involve silicone and oxygen atoms,
where, for example, if we have a tetrahedral arrangement
of four oxygen atoms around each silicon atoms,
we get this extended network of silica,
which is essentially stained.
A very hard material.
High melting point.
Very, very strong.
Finally, I'd like to talk about metallic bonding.
It's no different, in an underlying sense,
to all of the other bonding theories,
in that it involves the interactions
of negatively charges electrons with positively charged nuclei.
If you imagine taking a whole of molecules together
and squeezing them, the energy levels
where they share their valence electrons all get compressed.
And in the middle, those compression of energy levels
for the valence electrons, not the core electrons.
They stay very tightly attached to each atom.
But the energy levels for all
of the valence electrons becomes smeared out.
The word we use in chemistry is we say
that they are delocalised.
And delocalised energy levels for the electrons means
that they can move about.
And a typical model that you will have seen
in your chemistry books for a metal is
to have the atoms swimming in a sea, a sea of electrons.
The electrons can move freely between one atom and another
because the energy levels between each
of the atoms have all been smeared together
to form effectively a continuum.
And this bonding arrangement,
where the electrons can move freely about,
gives metals some unique properties.
One of them is that they're obviously high electrical
conductors, because the electrons can flow
through the metal.
Another one is that many metals are useful,
effective thermal conductors.
And finally, we are able to shape metals.
Metalworkers are able to hammer and change the shape of metals.
They tend to ductile, because by moving the atoms around in space
within this sea of electron densities,
we're not changing the overall energy levels of the system.
So a lot of the unique properties of metals
that make them different from other materials come
from the overall bonding arrangements, but the bonding
for metals is intrinsically no different
to all the other bonding arrangements.
We have negative electrons interacting
with positively charged nuclei.
And that's really all one needs to know
to understand all the different types
of bonding available to you in chemistry.
We can use the idea of sharing of electrons between atoms,
that was the whole concept of covalent bonding
that I mentioned earlier, to talk about predicting the shape
and structure of molecules.
And a very common and easy-to-use theory
to address this issue is known as Lewis structures,
developed by a chemist by the name
of Lewis almost a hundred years ago.
The whole idea of Lewis structures is
to pair electrons up, and then to have those electrons
as either bonding pairs or lone pairs in a molecule.
And there are a series of steps that you need to think
about when trying to develop a Lewis structure for a molecule.
Step one. Count how many valence electrons they are
for each atom in the molecule.
By valence electrons, I mean the outer electrons.
We don't count the core electrons.
Let's think of a carbon atom, for example.
You would all know from year 11 that the electron configuration
for carbon is 1s2, 2s2, 2p2.
It is it the 2s and the 2p electrons, the outer electrons,
that have valence electrons available to share
with other atoms in a molecule,
and they're the electrons we want to focus on.
So let's count the number of valence electrons.
Having determined how many valence electrons there are
for each atom in the molecule,
we want to determine how many pairs
of valence electrons there are, and we can then start
to arrange those pairs of electrons around the atoms,
in order to predict a molecular shape.
I'm going to go through an example now
to show you how all of these steps work.
So let's consider sulphur dioxide, SO2.
A sulphur atom has six valence electrons.
An oxygen atom has six valence electrons.
And we've had two oxygen atoms in the molecule,
so we have a total of 18 valence electrons available
for this molecule.
We are not counting the core electrons
that stay tightly abound around each of the sulphur
and the two oxygen atoms.
A total of 18 valence electrons tells us immediately
that we have nine pairs of valence electrons to account
for in this particular molecule.
Having determined that we need to account for nine pairs
of valence electrons in the molecule, we need to start
to fill those electrons in around the atoms.
A useful starting point is to think
about the least electronegative atom being
in the centre of the molecule.
The exception to that rule is when we have hydrogen atoms.
They, of course, will always be on the outside and a molecule
like methane, carbon would be in the centre,
even though carbon is more electronegative.
But in most cases, the central atom will be the least
electronegative atom.
From what we've seen earlier in this presentation,
that is the sulphur atom.
So let's put sulphur in the middle
and the oxygens on the outside.
What we now want to do is to start to account
for the nine pairs of valence electrons that we have
by drawing single bonds between each
of the atoms in the molecule.
We've accounted now for one pair and a second pair,
so we have seven pairs
of valence electrons left to account for.
A rule of thumb.
You will learn this from practice.
You can't get around this by just doing it once
and applying a simple set of algorithms.
But terminal oxygen atoms usually have either three lone
pairs or two lone pairs around them.
Usually two lone pairs when oxygen's involved
in a double bond, or three lone pairs when oxygen's involved
in a single bond as a terminal atom.
So because we have, as a starting point,
two sulphur oxygen single bonds, let's put three lone pairs
around each electronegative oxygen atom.
Now we have accounted for one, two, three, four, five, six,
seven, eight pairs of valence electrons.
We have a total of nine to account for,
so we will put the ninth pair on the central sulphur atom.
When we look at this structure, we need to see whether
or not we have satisfied or violated the octet rule.
The octet rule tells us that we should have a total of no more
than eight electrons around each atom.
We can see in the oxygen we have two, four, six, eight.
So the octet's okay for the oxygen.
Over here for this oxygen, two, four, six, eight electrons.
The octet's okay for that.
For the sulphur, we have two, four, six.
That doesn't seem to be something that we want to worry
about too much for the moment,
because the peripheral atoms are okay.
But we might have to come back and revisit it.
The next step is to consider the formal charges on the molecule.
The formal charge on an atom is the difference
between the number of valence electrons
that that atom would have in a hypothetical isolated state
and the number of electrons available to the atom
in the shared structure.
So let's consider the sulphur atom first.
Sulphur would have six valence electrons if it was isolated.
And if we look at this structure on here on the board,
we have two electrons in this lone pair, and we share one
from this bond and one from that bond.
So there are actually a total of four electrons
around the sulphur atom in this structure, so we can subtract
from that four valence electrons,
and that would tell us that we have a formal charge
of plus two.
So we would draw plus two on here.
If we consider the oxygen atom,
each of them are going to be the same.
We start off with six valence electrons being the isolated
atom to begin with.
And then, if we look here, we have two, four, six,
and we're sharing one in this bond.
Seven. So six minus seven means
that we will have valence electrons.
We will have a formal charge of minus one on this oxygen atom,
and minus one on this oxygen atom.
What we want to do in this Lewis structure is minimise the number
of formal charges on the molecule.
This particular structure, we have three formal charges.
And instinctively, that would suggest
to me that that's too many.
So how might we address the issue?
So our current structure, the sulphur dioxide,
has three formal charges, minus one on each
of the terminal oxygen atoms, and plus two
on the central carbon atom.
If we were to consider moving one of these lone pairs
in to share with a bond between the central sulphur
and the two terminal oxygen atoms,
we would now have a structure that looks
like it's two sulphur oxygen double bonds, two lone pairs
around the terminal oxygens, and now,
a single lone pair remaining on the central sulphur.
If we look at the formal charges now,
let's start with the sulphur.
We know we have six valence electrons
in an isolated sulphur atom.
Now we have two, one from this bond,
one from that bond, which is four.
One from that bond, one from that bond, which is six.
Six minus six is zero.
We have a neutral sulphur.
Let's look at the oxygen atoms.
We have two, four, one from this bond.
One from that bond, which is six.
Six minus six is zero, so we have a structure
that has no formal charges on it.
We have minimised the number of formal charges.
And this is the predicted Lewis Dot Structure
for sulphur dioxide.
The Lewis structure model
of chemical bonding is really useful for being able
to determine how many pairs of valence electrons we have
in a molecule and the overall geometric arrangement
of molecules on a two-dimensional page.
As I mentioned earlier, one of the important points to check
for in a valid Lewis dot structure,
apart from minimising a number of formal charges,
is to see whether an octet has been satisfied.
And we can see with our terminal oxygen atoms that we have two,
four, six, eight total number of electrons
around each oxygen atom.
So we have satisfied the octet.
In your year 12 curriculum, you only need
to worry between the octet.
So if we look at the central sulphur,
we can see that there are actually more
than eight electrons.
The octet theory doesn't always apply,
and that can confuse many students.
So this particular molecular example is one
that you're unlikely to see in an examination context,
but the example I went through shows you how to go
about allocating all of those nine pairs
of valence electrons available to the molecule.
One, two, three, four, five, six, seven,
eight, and the lone pair.
Nine. What this theory doesn't show you
or this Lewis structure theory doesn't show you is the
three-dimensional arrangement
of electron densities around a molecule.
And to do that, we invoke a model known as VSEPR,
which is an acronym for Valence Shell Electron Pair Repulsion.
And what the VSEPR model requires you to do is
to first draw a Lewis structure for a molecule, and then look
for the number of regions of electron density
around the central molecule.
Not the number of bonds.
Not the type of bond, but the number
of regions of electron density.
So if we look at our SO2 example,
the central sulphur atom has a region of electron density
for one sulphur oxygen bonding region.
It does not matter whether it's a single or a double bond.
It's one region of electron density.
We have a second region of electron density
for the other sulphur oxygen bond.
And then, finally, we have a third region of electron density
for the lone pair on the central sulphur atom.
So the key thing about VSEPR is to count the number of regions
of electron density around the central atom.
And now, we can look at a simple table of information to work
out what will be the overall geometric arrangement
of electron density in three dimensions
around the central atom.
If we had two regions of electron density,
we would have a liner arrangement of electrons.
Three regions of electron density, which is what we have
for our sulphur dioxide molecule.
We have a trigonal planar arrangement.
Four regions of electron density and we would say
that the arrangement of electrons
around the central atom is tetrahedral.
If we had five regions, we have a funny word here.
Trigonal bi-pyramidal.
You would all know that a pyramid
in Egypt has a square base, so a pyramid has a square base
and it comes up to a point.
That's called a pyramid.
If you have a triangle base and it comes up to a pyramid,
that's called a trigonal pyramid.
And a trigonal bi-pyramidal structure means
that we have a trigonal pyramid of electron density
above the central atom
and another one below the central atom.
Finally, if we have six regions of electron density
around our central atom, we would have what's known
as a octahedral arrangement of electron density.
In our sulphur dioxide example,
we found that there were three regions of electron density,
so we would say the geometric arrangement of electrons
around that sulphur atom is trigonal planar.
Trigonal planar.
However, the shape of the molecule is not trigonal planar.
This non-bonding pair of electrons, this lone pair here,
contributes to the arrangement of the electron density,
but it does not contribute to the overall shape
of the molecule, and the shape
of the molecule is therefore shape equals bent, B-E-N-T.
it doesn't sound very technical, but that's the word we use
to describe the shape of sulphur dioxide.
We discount the lone pairs for the shape of the molecule,
but the lone pairs are critically important
for determining the geometric arrangement of electron density
around the central atom.
Having determined the shape of a molecule, it's important
to be able to predict some
of the overall molecular properties of a molecule.
For example, is that molecule polar or non-polar?
I'll show in the next video that the answer
to that question plays a fundamental role on some
of the intermolecular forces, and therefore,
the physical properties of that molecule.
So let's build up to understanding what it means
to be a polar or a non-polar molecule.
We come back to the concept of electronegative
that I talked about earlier.
We have electronegative elements in the upper right
of the periodic table, and the left-hand side
of the periodic table has less electronegative elements.
So if we were to draw a covalent representation of the HF bond,
we would have an electropositive H atom
and an electronegative F atom.
That unequal sharing of electron density that I talked
about earlier can be represented by a simple diagram I've drawn
in green here, which is a bond moment.
B-O-N-D moment.
What the bond moment representation is showing us is
that the arrow is pointing to the negative end of the bond,
and to help us remember that, we draw a little vertical line
at the opposite end of the arrow to try
and represent a plus sign.
So we've got a bond moment to represent unequal sharing
of electron density, pointing towards the negative element.
In our sulphur oxygen bonds that we talked about earlier,
we would have bond moments going
from the least electronegative sulphur
to the more electronegative oxygen, with the positive sign
at the more positive end.
So let's consider carbon dioxide.
I've drawn the lowest structure
for carbon dioxide on the board here.
We can see that we have one, two, three, four,
a total of eight electrons around this oxygen,
so the octet's satisfied.
The octet is satisfied on that oxygen, and we have one, two,
three, four pairs, eight electrons
around the central carbon atom.
No lone pairs on the carbon atom.
However, each of these bonds is polar.
Carbon is less electronegative than oxygen,
so we can draw a bond moment
for that particular carbon oxygen bond,
as I've represented on the board.
The other carbon oxygen bond has a bond moment that's equal
and opposite in magnitude.
So these two bond moments will cancel out, and I'll come back
to that concept in a moment.
If I was to draw my sulphur dioxide molecule
from earlier on, we can see, again,
that we have a bond moment
for one sulphur oxygen bond, pointing as shown.
And the bond moment for the other sulphur oxygen bond
as shown.
These two bond moments do not cancel each other
out in a vector sense.
They are both pointing in the same direction.
So, overall-- sorry about that.
Overall, if I change pen,
we find that the bond moments here cancel.
[Writing on Board] Bond moments cancel.
And the molecule in CO2 has no net dipole moment.
It is said to be non-polar, because it has no dipole moment.
In our SO2 example, these two individual bond moments are
pointing generally in the same direction,
so we have a net dipole moment.
Dipole moment.
And the molecule is said to be polar.
>> Hi guys.
Hopefully that helped you out.
If you still have any issues with chemistry in general,
just go to the links that we've given you,
or ask a Chem teacher.
They're probably your best resource.
Good luck with exams,
and hopefully we'll see you here at Curtin.