Periodicity of Valence or Oxidation States Video Lecture - Chemistry Class 11 - NEET

Let's see if we can come up
with some general rules of thumb
or some general trends
for oxidation states
by looking at the
periodic table.
So first, let's just focus
on the alkali metals.
And I'll box them off.
We'll think about
hydrogen in a second.
Well, I'm going to box-- I'm
going to separate hydrogen
because it's kind
of a special case.
But if we look at the
alkali metals, the Group
1 elements right
over here, we've
already talked about
the fact they're not
too electronegative.
They have that one
valence electron.
They wouldn't mind giving
away that electron.
And so for them,
that oxidation state
might not even be a
hypothetical charge.
These are very good candidates
for actually forming
ionic bonds.
And so it's very
typical that when
these are in a molecule,
when these form bonds,
that these are the things
that are being oxidized.
They give away an electron.
So they get to-- a
typical oxidation state
for them would be positive 1.
If we go one group over right
over here to the alkaline earth
metals, two valence electrons,
still not too electronegative.
So they're likely to fully
give or partially give away
two electrons.
So if you're forced to
assign an ionic-- if you
were to say, well, none
of this partial business,
just give it all away or
take it, you would say,
well, these would
typically have an oxidation
state of positive 2.
In a hypothetical ionic
bonding situation,
they would be more likely
to give the two electrons
because they are not
too electronegative,
and it would take them a lot
to complete their valence shell
to get all the way to 8.
Now, let's go to the other
side of the periodic table
to Group 7, the halogens.
The halogens right over here,
they're quite electronegative,
sitting on the right-hand
side of the periodic table.
They're one electron
away from being satisfied
from a valence
electron point of view.
So these are typically reduced.
They typically have an
oxidation state of negative 1.
And I keep saying typically,
because these are not
going to always be the case.
There are other things
that could happen.
But this is a
typical rule of thumb
that they're likely to
want to gain an electron.
If we move over one
group to the left,
Group 6-- and that's where the
famous oxygen sits-- we already
said that oxidizing something is
doing to something what oxygen
would have done, that oxidation
is taking electrons away
from it.
So these groups are
typically oxidized.
And oxygen is a very
good oxidizing agent.
Or another way of thinking
about it is oxygen
normally takes away electrons.
These like to take
away electrons,
typically two electrons.
And so their oxidation state
is typically negative 2-- once
again, just a rule of
thumb-- or that their charge
is reduced by two electrons.
So these are typically reduced.
These are typically oxidized.
Now, we could keep going.
If we were to go right over
here to the Group 5 elements,
typical oxidation
state is negative 3.
And so you see a
general trend here.
And that general trend--
and once again, it's
not even a hard and fast rule
of thumb, even for the extremes,
but as you get closer
and closer to the middle
of the periodic table,
you have more variation
in what these typical
oxidation states could be.
Now, I mentioned that
I put hydrogen aside.
Because if you really
think about it,
hydrogen, yes, hydrogen
only has one electron.
And so you could
say, well, maybe it
wants to give away that electron
to get to zero electrons.
That could be a reasonable
configuration for hydrogen.
But you can also view hydrogen
kind of like a halogen.
So you could kind of view it
kind of like an alkali metal.
But in theory, it could
have been put here
on the periodic table as well.
You could have put hydrogen
here, because hydrogen,
in order to complete
its first shell,
it just needs one electron.
So in theory, hydrogen
could have been put there.
So hydrogen actually
could typically
could have a positive or a
negative 1 oxidation state.
And just to see an
example of that,
let's think about
a situation where
hydrogen is the oxidizing agent.
And an example of that would
be lithium hydride right
over here.
Now, in lithium hydride,
you have a situation
where hydrogen is
more electronegative.
A lithium is not
too electronegative.
It would happily give
away an electron.
And so in this
situation, hydrogen
is the one that's
oxidizing the lithium.
Lithium is reducing
the hydrogen.
Hydrogen is the one that
is hogging the electron.
So the oxidation state on the
lithium here is a positive 1.
And the oxidation state on the
hydrogen here is a negative.
So just, once
again, I really want
to make sure we
get the notation.
Lithium has been
oxidized by the hydrogen.
Hydrogen has been
reduced by the lithium.
Now, let's give an example where
hydrogen plays the other role.
Let's imagine hydroxide.
So the hydroxide anion-- so you
have a hydrogen and an oxygen.
And so essentially, you could
think of a water molecule that
loses a hydrogen proton but
keeps that hydrogen's electron.
And this has a negative charge.
This has a negative 1 charge.
But what's going
on right over here?
And actually, let
me just draw that,
because it's fun
to think about it.
So this is a situation
where oxygen typically
has-- 1, 2, 3, 4,
5, 6 electrons.
And when it's water, you
have 2 hydrogens like that.
And then you share.
And then you have
covalent bond right
over there sharing that
pair, covalent bond sharing
that right over there.
To get to hydroxide,
the oxygen essentially
nabs both of these
electrons to become--
so you get-- that
pair, that pair.
Now you have-- let me
do this in a new color.
Now, you have this pair as well.
And then you have that
other covalent bond
to the other hydrogen.
And now this hydrogen is
now just a hydrogen proton.
This one now has
a negative charge.
So this is hydroxide.
And so the whole thing
has a negative charge.
And oxygen, as we have
already talked about,
is more electronegative
than the hydrogen.
So it's hogging the electrons.
So when you look at
it right over here,
you would say, well,
look, hydrogen,
if we had to, if we were
forced to-- remember,
oxidation states is just
an intellectual tool
which we'll find useful.
If you had to
pretend this wasn't
a covalent bond, but an
ionic bond, you'd say,
OK, then maybe this hydrogen
would fully lose an electron,
so it would get an
oxidation state of plus 1.
It would be oxidized
by the oxygen.
And that the oxygen actually
has fully gained one electron.
And you could say, well,
if we're forced to,
we could say-- if we're
forced to think about this
is an ionic bond, we'll say
it fully gains two electrons.
So we'll have an oxidation
state of negative 2.
And once again,
the notation, when
you do the superscript
notation for oxidation states
and ionic charge, you write
the sign after the number.
And this is just the convention.
And now, with these two
examples, the whole point of it
is to show that
hydrogen could have
a negative 1 or a positive
1 oxidation state.
But there's also something
interesting going on here.
Notice, the oxidation
states of the molecules
here, they add up to the whole--
or the oxidation state of each
of the atoms in a
molecule, they add up
to the entire charge
of the molecule.
So if you add a positive 1
plus negative 1, you get 0.
And that makes sense because
the entire molecule lithium
hydride is neutral.
It has no charge.
Similarly, hydrogen,
plus 1 oxidation state;
oxygen, negative 2 oxidation
number or oxidation state--
you add those two together, you
have a negative 1 total charge
for the hydroxide anion,
which is exactly the charge
that we have right over there.