>> Hi everyone.
Now that you have a basic
understanding of the atom,
atomic mass, and isotopes,
I would like to talk
to you a little bit more
about a very important aspect
of the atom and those
are actually called
valence electrons.
And the valence electrons
will actually get us
into a bonding model, which
you may already have heard of,
called the Lewis structure.
So let's get started.
So, here is a simple
model of an atom.
This is called actually
the Bohr Model.
Your nucleus is in the center.
You have your protons
and neutrons.
And then you have these red
dots represent your electrons.
Notice that there're
circles here.
And a very simplistic
way of thinking
about these electrons is they
kind of orbit the nucleus,
kind of like how the
planets orbit the sun.
It's not exactly right.
But it's a kind of a simple
easy way to think about it,
which is totally fine.
But anyway, so you look at these
electrons, and notice how some
of these electrons are closer
to the nucleus and some
of the electrons in this
ring are further away
from the nucleus.
So the main point here
is that electrons are
at different distances
from the nucleus.
The further away the
electrons are from the nucleus,
the higher energy
those electrons have.
And the electrons that
are in the furthest region
from the nucleus are called
the valence electrons.
So if you look at this example
of oxygen, there's 2, 4, 5,
6 that are in this final ring.
And those are the
valence electrons.
These two in the middle, they're
called your core electrons.
They're not part of
that valence electron.
And the valence electrons
are actually really important
because they are involved in
all your chemical bonding.
You can actually predict your
valence electrons in a very,
very, very, very simple way.
Okay, so here is
your periodic table
without all the elements listed.
But you have your you have
your group 1A, your group 2A,
your group 3A, 4A,
5A, 6A, 7A, 8A.
Notice I'm ignoring the
transition elements.
So in your main group, these
are called your main group.
One through eight
are your main group.
Your group 1A has
two valence electron.
Your group 2A has two
valence electrons.
Your group 3, here, has
three valence electrons.
Four, group 4 has four
valence electrons.
Group 5 has five
valence electrons; 6,
six valence electrons; 7,
seven valence electrons.
And 8, eight valence electrons.
Very simple.
And it's really nice that
you could just actually look
at the numbers on the
periodic table by their groups
and know their valence
electrons.
Okay, so I'm going to go ahead
and show you with this one
with actually the elements
in the periodic table.
So again, you're going to ignore
the transition elements not just
because you can't
predict them as easily
as you can the main group.
Okay, again, the [inaudible]
that you can use
the group number
or you can actually count, even.
You can go 1, 2,
3, 4, like that.
So for example, oxygen
is in 6A and, therefore,
it has six valence electrons.
Another example would
be magnesium.
Magnesium has two valence
electrons because it's
in the group number 2A.
But helium is an exception.
It only has two valence
electrons.
And you might be wondering why.
Don't worry about too much
but if you actually count
across the row, hydrogen
has one proton, right.
Helium has two protons.
What are the maximum number
of protons helium must have?
Two to keep it charge neutral.
Otherwise, if you had, you
know, eight valence electrons,
you would have like 8 minus 2
will give you 6 negative charge.
So no, you can't have a
negative charge on your helium.
It needs to be neutral.
You have two protons; therefore,
you must have two
electrons, maximum.
So now we have the valence
electrons out of the way,
you have a general understanding
of it, I want you to kind
of get an idea of what
models of the atoms are.
And we use models to explain not
just what the material is made
of, but also how it is going
to behave, how it's going
to change, basically how
to predict how things
are going to react.
So an example is hydrogen gas.
They actually used this
in 1937 for a blimp.
It was a cross-Atlantic
blimp ride.
And as it is landing,
it actually blew up.
And this is because hydrogen
gas is really reactive.
Now if you compare that to
helium, which is a noble gas,
it's more of an inert
gas and not very reactive
and this is what we now
fill our blimps with.
And keep in mind again
that chemistry is all
about the electrons.
So if you were to
compare the electrons --
Here we go.
So for example, if you compare
the electrons of hydrogen,
you have the same number of
valence electrons as hydrogen,
lithium, sodium, potassium,
whatever, all these guys here
in group 1 have one
valence electron.
All of these elements
are highly reactive.
They are very similar in terms
of their reactivity as hydrogen.
And the same thing if
you look at helium,
all of these guys here,
helium, neon, argon,
all these guys here
are your noble gases.
They're all inert and they
have very similar properties.
And they also have the same
number of valence electrons.
So again, keep in mind,
chemistry is all
about electrons.
We can actually use a
very simplistic modeling
which one type is called
Lewis Bonding Theory,
and we can use this to describe
how atoms bond together.
And atoms bond because
it results
in a more stable
electron configuration
or a more stable
electron arrangement.
The atoms bond together
by either transferring
or sharing electrons
so that all atoms
of obtain an outer shell
with eight electrons.
That's the important part
here, eight electrons.
And that means they will look
like a noble gas because except
for helium, all your other
noble gases have eight
valence electrons.
And this is called
your Octet Rule
because of the eight electrons,
Note, there are some exceptions
to this rule, okay, and the
key to remembering is to try
to get an electron
configuration like a noble gas,
the closest noble gas.
So for example say helium
is your closest noble gas,
your exceptions to that
would be lithium, beryllium,
hydrogen, and of course helium.
All these will only
have two electrons,
doublet because they're more
likely they going to try to look
like helium since helium is
the closest noble gas to them.
We can represent using
the Lewis structures,
we can represent our
valence electrons with dots.
So for example, lithium
is in column 1 or group 1.
So it has one valence
electron, put one dot.
Beryllium would have two.
And you can keep going forward.
And notice that for
example carbon, each dot,
each valence electron goes
around the symbol for carbon.
And then once you've
gone around all the way,
then you can put another dot.
So nitrogen would be another
dot here, where it's paired.
And then finally you get to
neon where you have your octet.
So again, it's stable
because that's your noble gas.
It has eight valence electrons.
Again, helium only has
two but it's still stable.
So here's your first example.
So we're going to write a
Lewis structure for water.
Your first step, okay,
you're going to write
out how many electrons you have,
okay, valence electrons, right.
So each hydrogen has
one valence electron.
You have two hydrogens, so
that's two plus six, right,
so two times each one plus six
for you oxygen gives you
eight valence electrons.
Your second step is just to
write out a basic structure.
Okay, you're going to
put oxygen in the middle,
and then the two
hydrogens on the side.
Okay, your skeleton structure
should be symmetrical.
Step three, put bonds
between them.
So each of these bonds you end
up using two electrons
for each bond.
So this would be two,
four electrons are used
which means you have
four remaining electrons.
Hydrogen can only have
two electrons total due
to the Duplet Rule.
So you can put the remaining
electrons around your oxygen.
So that would give
you 2, 4, 6, 8.
So your octet rule is
satisfied on oxygen
and your hydrogens
have doublets.
Okay, so that's just
a simple example.
And to note, you can
actually have double bonds.
In the water example you
only have a single bond
but you may need to use a
double bond if, for example,
you cannot satisfy
the octet rule.
So say you have oxygen, the
O2, O2 oxygen in the air,
each oxygen only gas
six valence electrons.
So if you were just to
bond it with a single bond,
which is like again,
the two dots
or you could put a single
line, either one is fine,
you would have 12 electrons.
And this guy on the
left, 2, 4, 6,
only has six valence
electrons surrounding it.
So the octet rule
is not satisfied.
To make up for that, you can
actually take one of the pairs
on the other oxygen and put it
in the middle, you have two dots
or you put the lines,
they mean the same thing.
What happens then is
you now have 2, 4, 6,
8 around one oxygen and 2, 4,
6, 8 around the other oxygen.
So now both oxygens have been
satisfied with the Octet Rule.
You can actually
do the same thing
with a triple bond
for the same reason.
So for example you have
nitrogen in the air N2,
each N atom has five
valence electrons.
So in the end, you're only going
to have ten valence
electrons here.
It's not enough to
get an Octet Rule,
to satisfy an Octet Rule
for either nitrogen.
So what you can do is you can
actually take two of these pairs
of electrons and put
them in the middle,
so now they're being
shared by both nitrogens.
So you're going to have 2, 4,
6 bonds actually being shared
or six electrons being shared
where you can draw those lines.
And so that's called a triple
bond because you have 1, 2,
3, triple, triple bond.
But now if you count,
you have 2, 4, 6,
8 valence electrons
around each nitrogen.
And again, you're Octet
Rule is now satisfied.
Okay, for example, you can
write a Lewis structure
for CO2, carbon dioxide.
Step one, count your
valence electrons.
I have 16 electrons.
I got that from four from the
carbon plus I have two oxygens,
two times the six valence
electrons for oxygen equals 16.
Draw my structure out.
This should be symmetrical.
Okay, you can draw just
the bond between them.
I have four electrons used.
I have 12 electrons remaining.
So I'm just going to go
ahead and start putting dots
around the oxygens to
give them the Octet Rule.
The Octet Rule is
satisfied for the oxygens.
And I'm out of electrons.
So notice I used them all up.
I have 2, 4, 6, 8,
10, 12, 14, 16,
so used of all the electrons,
the valence electrons.
But the carbon, the Octet
Rule is not satisfied.
So what I can do is actually
convert the lone pairs
out on the sides of the
oxygen into a double bond.
And now everyone has the
Octet Rule satisfied.
Carbon -- So for the
oxygens, they have 2, 4, 6, 8.
And the carbon has 2, 4, 6, 8.
The final thing you should
be aware of is resonance.
So Lewis structures often do
not accurately represent the
electron distribution
in a molecule.
And real molecules are a hybrid
of all possible Lewis
structures.
And resonance stabilizes
the module.
So what that means for
you is that if you look
at these double bonds,
here I have single bonds,
notice how these are all
oxygens, so it's the same atom.
What actually occurs is
that this double bond
that this bond is actually, if
you draw the resonance, okay,
just to show you, you
put these arrows here
between all the molecules
and the double bond could be
on any of the oxygens.
Right. Why pick one.
It could be any of them, really.
Right. And so what this really
means is that the structure,
the size of the bond
is somewhere
in between a double
bond and a single bond,
but the way we represent that is
through these showing all
the possible structures,
we call this resonance.
Okay, everyone, that's
it for now.
I'll see you next time.
