So earlier we looked at covalent bonding,
and specifically we looked at an example-- I think we did.
If not, you're perhaps familiar with methane.
Methane is a this molecule, CH4.
It's carbon with the four little hydrogens bonded to it.
And each of those little straight lines
was a covalent bond.
And so this was I was fine, in fact.
And even if we draw this type of depiction here with little
electron dots, and then we say, let's add in some other
something else, like hydrogen. And I'm just colour coding this
so you can get a sense for who's contributing the electron.
You do that, and then you see, well, that's great.
We got eight electrons, and that made us happy,
made the carbon happy.
Because that was octet-- stability.
It had this inert gas, or sorry, noble gas electron
configuration.
But we need to be a little more careful.
It would be useful to us, I think,
to explore this in more detail.
So this is really a two dimensional depiction-- two
dimensional.
And again, it works fine.
We're just adding up to 8.
But if we look at the electron configuration
for carbon in a little more detail,
we know it's 1s2, 2s2, 2p2.
and so that was where we went and we said,
OK, there's four valence electrons,
so let's draw in these four electrons.
But what we didn't do is differentiate
between 2s and 2p.
We just drew these four white dots.
And in fact, in a way what we did
when we sketched this dot diagram here was we
just pretended they were all kind of the same.
And really, there's actually a formal name for that model,
if you will.
And that's what we're going to get into.
So what we've got is we've got 2s,
and we've got the-- I should draw that in a different
colour, just for effect here.
So then we've got 2s, and we've got 2p.
That's 2p.
And so we said, well, we need to make four bonds.
Why don't we just combine them together?
So we'll take the three p states, and the one s state.
So we get s, and we've got p.
And how many ps are there?
There's three of them.
We say, let's just blend them all together
and start filling up.
So then we've got the four electrons from carbon,
and we can add one each from hydrogen,
and then get our four bonds.
But to do that, we had to treat them the same.
In fact, just to drive home, I'll just draw them all
in the same colour box-- not that the colour of the box
matters.
Of course, you'll appreciate that,
but we treat them all the same.
So it's a blend.
It's like if you ride a bike to work,
and you could ride a mountain bike,
you could ride a road bike, a racing
kind of bike, high pressure tyres, and all that.
Or a lot of people ride a blend of the two.
It's a hybrid bicycle.
You get two species of fruit or something like that,
and you blend them together to make
something delicious and nutritious, well,
it's a hybrid.
So this is a hybrid orbiter.
It's a hybrid orbiter.
And specifically, this one here, that's taken 1s
and 3p p is called sp3 hybridised.
So it's sp3 hybridised.
And so now we've got four bonds that are equivalent.
There's nothing different about these.
It's not like two of them are s, and two of them are p.
There's four equivalent bonds.
And in fact, we can extend that out
into predicting well, if there's just four of them,
it would look exactly the same.
And then we want to go out into three dimension-- we would
expect that there should be nothing
really different between one of bonds
and any of the other ones.
So we take our carbon, and we draw these bonds out here.
And I'm going to try to depict them in three dimensions, here.
OK, so we need three dimensions, going off to the hydrogen.
Where this one was going-- this guy here
is going back into the page, and these two
are kind of coming out.
It forms this very common structure in chemistry.
And if I make a little surface on that like this,
and dot that in at the back, this is a four sided solid.
Four sided, it's called a tetrahedra.
So that geometric solid is called a tetrahedra,
or this is a tetrahedral arrangement
of hydrogens around the carbon.
Another way that sometimes you can
see this spatially and oriented in space
is to take one of my famous spheres here,
and-- or infamous, I don't know which--
but I guess I'm drawing this thing.
There you go, look at that.
Beautiful, as always.
And I'm going to position a little carbon or something,
right at the very centre.
And then each of the bonds is equivalent,
so they should all go the same distance.
So when we go up to here, perhaps, one
would go out this way, and this way,
and the other guy would go off into the back.
But they're all exactly the same bond length,
and each of these bonded atoms to the central atom
is at exactly the same distance from the central one.
So that's another way of looking at it,
and that geometric solid that I drew
there is another way that it's often drawn.
If I draw it without the stuff, the little bonds,
it's perhaps a little bit more clear
you can see the four sides.
This surface over here, one surface.
This right side faces another.
And then we've got that back face.
Now, look at the bottom face.
1, 2, 3, 4.
Tetrahedra bottom.
OK, so let's let that geometry-- and it's really actually a very
important geometry, in chemistry, in the world ,
because it shows up in a lot of things.
In fact, carbon in diamond is also
based on this tetrahedral arrangement.
And again, it's because it's sp3 hybridised.
It's just this time in diamond, carbon is bonded to four
other carbons, rather than four hydrogens.
And anyway, again this symmetry shows up
in other ceramic crystal structures, as well,
so it's a really important one to add to understand.
Historically, actually, the structure was known,
the tetrahedral arrangement was known before people had
developed a hybridised model.
And that was developed afterwards, to explain this.
But we have the electricity of seeing them together
at this point.
What we'll do in a separate topics
is look at the structure of diamond
in a little more detail, and look
at some other ceramic structures that
have this tetrahedral symmetry to them.