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CATHERINE DRENNAN: All right.
Let's just take 10 more seconds.
OK.
So let's look at this one.
So first, you want to notice
that you have a plus 1 here.
So you've lost an electron.
And then you want to think
about what happens when you
start filling the deorbitals.
So when you start
filling the deorbitals,
then the energy changes and
4s and 3d switch an energy.
And so we could write
this one either way.
We could have put the 3d
first and the 4s second.
But importantly now, because
of that switch in energy,
the electron that is lost is
lost from the 4s over here.
So this has to do with the fact
that electron configuration
of neutral atoms and ions
are different and especially
with this 4s-3d switch, that the
3d orbitals drop, change energy
when you start to fill them.
And there's really very similar
energy between 4s and 3d.
And that leads to
some of the exceptions
that you're responsible
to know that there
can be subtle things that
switch the energy a little bit.
So because they're
so close in energy,
you have this half-filled
and full-filled thing
where you can pull
an electron from 4s
and put it in 3d to make
3d5 or to make 3d10.
So they're very close in energy.
And that leads to some of
these interesting features.
OK.
So today's lecture, we're
moving on to the periodic table.
But we're actually talking
about a lot of the things
that we just talked about.
So today actually
turns out to be
an awesome review for some
of the material that's
on the exams.
So that worked out really well.
In past years, this
material was on exam 1.
Exam 1 was later.
And so if you get old exams from
other people, not the ones that
are posted, the
ones that are posted
are mostly old exams, except
we were placed questions
for material not covered on
this exam with material that
is covered.
So it's not 100% an old
exam from this class
because we've never had
an exam this early before.
So there were no good examples.
So if you get old exams
from other people,
do not freak out
when you look at it,
and like, oh my goodness.
Somehow I haven't learned this.
There are today's
lecture and also
Friday's lecture we're
typically on exam 1.
So just keep that in mind.
So use our practice
exams and then you
will not have that problem.
All right.
So moving on to
the periodic table.
This is very exciting for me.
And so today, we're going
to talk about trends
in the periodic table.
We're going to finish
that up on Friday
after the exam, which will
be a clicker competition.
And then we're going
to go on to talk
about bonding of the elements
in the periodic table.
So that's where we're headed.
So the periodic table, here
is one of them that's up here.
So this was originally
put together a while ago.
And it turned out to
be amazingly accurate.
And this really describes
all of the elements.
So this is kind of like the
artist's paintbox for a chemist
or wordsmiths words.
These are all the ingredients
that go into making everything.
Some of these elements
are incredibly dangerous
and they're used
to make explosives.
They're used to make bombs.
Other elements here are
found in the human body
and allow us to live.
All materials, whether
it's a desk, a pointer,
a bottle of water, everything
is made up of elements.
So this is one of
the reasons why
chemistry is so cool because
we think about the elements.
And elements are
made of everything.
So we think about everything
that makes up everything.
And that's pretty neat.
So just to kind of
give you a flavor
of the joy of the periodic
table and introduce you
to the elements that make
up this periodic table,
I feel like we should
think about this in music.
[MUSIC - THEY MIGHT BE GIANTS,
 "MEET THE ELEMENTS"]
[SINGING] Iron is a metal.
You see it every day.
Oxygen, eventually,
will make it rust away.
Carbon in its
ordinary form is coal.
Crush it together and
diamonds are born.
Come on, come on and
meet the elements.
May I introduce you to
our friends, the elements?
Like a box of paints that are
mixed to make every shade,
they either combine to
make a chemical compound
or stand alone as they are.
Neon's a gas that lights up
the sign for a pizza place.
The coins that you pay with
are copper, nickel, and zinc.
Silicon and oxygen make
concrete bricks and glass.
Now add some gold and silver
for some pizza place class.
Come on, come on and
meet the elements.
I think you should check out
the ones they call the elements.
Like a box of paints that are
mixed to make every shade,
they either combine to
make a chemical compound
or stand alone as they are.
OK.
So you get the sense of this.
The song is quite accurate.
It has lots of information
[INAUDIBLE] them.
And it points out
some other things
like elephants are
made of elements.
And we're made of elephants--
oh no, wait-- elements.
No, it's a really fun, fun song.
And it really, I
think, expresses
why It's so important to
learn about the properties
of the elements
and all the things
that you can do with them.
So when it was
originally put together,
it was put together
based on sorting elements
by their properties, such as
ones over here in column 1
are soft and reactive metals.
And it was observed that
the elements over here
are pretty inert.
So they were all
grouped together.
And later, we have pretty
much kept this grouping.
But now, it's really grouped
by the electron configurations,
which is one of the things
you need to know for the exam,
how to write these
electron configurations.
And these reactive
metals, it turns out,
they only have one
valence electron.
So they like to react because
they want to have a noble gas
configuration, so they're very
reactive whereas these others
that were not reactive, have
filled electron configurations.
So they don't want
any extra electrons
and they don't
want to lose them.
They don't want to get any.
They're very happy as they are.
So they're inert.
So now these groupings
make a lot of sense
in terms of the
electron configurations.
Now, it doesn't
tell you everything.
So if you know one element
is very safe to consume,
that doesn't necessarily
mean something right next
to it is just as good.
And if we consider over here,
we consider lithium, sodium,
and potassium,
sodium and potassium
are ions that are really
important in the human body.
And you have to make sure that
if you're exercising a lot,
that you keep up the
amounts that you're getting.
So they're very important ions.
And they will often hang around
and serve as counter-ions
to other molecules in your body.
You have citrate in your body.
So you could have sodium citrate
where the sodium is hanging out
or potassium citrate.
But that doesn't
necessarily mean
that other things will
work as well, like lithium,
for example.
But a while ago when 7-Up soda
was first put on the market,
they thought well, sodium and
potassium are a little boring.
Let's sort of make things
a little more exciting
and use lithium citrate instead.
It's right there in the same
column of the periodic table.
And citrate makes
things taste lemony,
which is a lovely taste.
And we'll have lithium as
the counter-ion to that.
And so they said this
dispels hangovers.
It takes the ouch out of grouch.
And does anyone know what
lithium is used for today?
So it's often used
for people who
have bipolar disorders
or manic depressive.
So it really did take
the ouch out of grouch.
But it's not something
that you should just
put in a consumable soda.
So it has somewhat
different properties,
even though it's part
of that same group.
So this is a lesson
that I feel like we
keep learning over and over
again with other things,
that you have to be a
little more careful.
Just because it's hanging
out near its friends,
doesn't mean it's going
to be exactly the same.
All right.
So the periodic table
is an amazing thing.
Let's think about the trends
in the periodic table.
So we're going to first
do ionization energy.
And we've already talked
about ionization energy.
So this is awesome
because it turns out
to be a really good
review for the exam.
So ionization energy,
again, is the minimum energy
it's going to take to remove
an electron from an atom.
And if we just talk about-- just
say IE for ionization energy--
we're going to assume
it's the first ionization
energy unless it is specified.
And we saw before that
ionization energy is opposite
in sign to the binding energy.
And so here we have the
binding energy of an electron.
And we know that this
is a multi-electron atom
because it depends on n and l.
If it was just hydrogen or
one other one-electron atom,
then anything with n, all
those orbitals, are degenerate.
But if you have putting
in multi-electrons,
then it matters whether you're
talking about not just n,
but l matters too, whether
it's an s orbital or p orbital.
So we've seen this before.
But now let's talk more about
different ionization energies.
So let's look at boron and
think about the first ionization
energy.
And this is the energy to move
an electron from the highest
occupied atomic orbital.
That's what that stands for.
And it's written
out in your notes.
And so what is the highest
occupied orbital in this case?
Just yell it out.
2p.
So let's look at removing
an electron from 2p.
If we do that, we
go to boron plus.
And we have 1s2,
2s2, an electron.
And this process, the energy
involved in this process,
is the ionization energy--
the first ionization energy.
It's also equal to the binding
energy of the 2p electron.
And again, the signs
are opposite here.
So now second ionization
energy, we just keep going.
The next highest occupied
atomic orbital is 2s.
So if we remove, we get
boron plus 2 1s2, 2s1
and an electron.
And now the energy difference is
due to these second ionization
energies.
So we say IE2.
And that is equal to the binding
energy of the 2s electron in B
plus because that's
what we're removing.
We're moving a 2s
electron from B plus here.
So we can keep going.
We can go to the third
ionization energy.
And now we're also going to be
removing an electron from 2s.
And when we remove the
electron, it only had one.
So now we have a boron
plus 3 1s2 and an electron.
The energy difference is the
third ionization energy--
IE sub 3.
And this is equal to
the binding energy.
Or the difference in sign
is the binding energy of 2s
in the plus 2 system.
So now if we look at this
little table over here that's
in your handout or this
little chart in your handout,
you can see that there is
quite a bit of difference
between these different
ionization energies.
So we were talking about boron.
So here we have the
first ionization energy,
second ionization energy,
third ionization energy,
and fourth ionization energy.
And so there can be
quite a bit of difference
in the magnitude of
these ionization energies
or how hard it is to pull
off successive electrons.
And so here are some of the
other ones you see when you're
going here with boron, you
remove the first one here,
the second one is about
three times harder.
We're jumping from p to s.
Not too much
difference within 2s.
But once we get to
helium here, 1s2, that's
really hard to pull off
another electron here.
So this fourth
one is really big.
And then we can look
at these other trends.
Beryllium here, we have the
2s and then we go to a 1s.
And then we have just the one
electron over here for lithium.
And then when we come up
here, it's a lot harder.
So we can look at these
tables and realize
these are not going to
be necessarily the same.
There can be big jumps
in ionization energy.
And I'm going to come back
to all of this and sodium
and potassium in a little bit.
But first, let's just stick
with boron for a minute
and think more about the
different kinds of ionization
there.
So now let's just consider
taking a 2s electron,
but from two different
types of boron-- boron
plus and regular boron.
So in this first
case here, we're
going to take one of
these two s electrons.
And now we have a
difference in energy.
This is the second
ionization energy.
The first one removed
the electron from p.
So we saw this before.
We're moving one of
the electrons from 2s.
And so this is the
second ionization energy.
It's also equal to the binding
energy of the 2s electron
in boron plus.
Now we're going to
take a 2s electron.
But we're going to
do it from boron.
So the p electron
is still there.
So we go from 2s2
to 2s1 over here.
And this energy difference
is an ionization energy
for a 2s electron.
And that's equal to
the binding energy
of the 2s electron in boron.
So do you think these
energies are going
to be the same or different?
Are they equal?
No.
So I showed you that
little chart and that
made you probably
think that there is
going to be some differences.
No, they're not equal.
So why are they not equal?
Well, when you have boron plus,
you have lost an electron.
So you have less electrons
available to shield.
So you have less shielding
in boron plus than in boron.
And if there's less
shielding, you're
going to have a higher Z
effective, less shielding.
They'll feel more of the
force of the positive charge
of the nucleus.
And therefore, it's
going to take more energy
to pull it off.
So it's going to be
more tightly bound.
It's going to be held in because
of this less shielding, higher
Z effective.
All right.
So now let's try a
clicker question.
10 more seconds.
OK.
So most of you did
not like answer 1.
But does someone
want to explain this?
And do you want to just walk up?
Someone want to
give an answer why?
OK, over there.
AUDIENCE: OK.
So if we're choosing between 2
and 3, the answer for number 2,
the 3p orbital has
two electrons in it.
And so the electrons, by
nature, kind of repulse
each other, right?
So it's a little easier
to pop one of those two
out than if there was only one
electron in there by itself
and you're trying
to pull it out.
So yeah, you'd pick 2 over 3.
CATHERINE DRENNAN: OK, yeah.
So actually, I don't know
if you can take the answer.
It's a little hard to read it
with the colors on top of it.
But here, you have
this plus system here.
So you've removed
this extra electron.
So there should be, you feel,
a higher Z effective here,
which will mean that it's harder
to sort of pull things off.
And let's see this one.
There's no way to take
the answer down, right?
So this one-- oh yeah.
OK there, that's better.
I can see this more.
So this one here, you're pulling
one from the s orbitals here.
The p orbital's
easier to pull it off.
It takes less energy
from p than from s.
OK.
So let's continue.
We'll come back to
some of these ideas
as we go along
because now, we're
going to think about
how these trends go
across the periodic table.
So across a row, ionization
energy is going to increase.
And the reason for this
is that Z is increasing.
So we're having more and more
protons, a bigger Z effective.
You're also adding
electrons though.
But n, the shell,
remains the same.
So if Z is increasing,
n is remaining the same,
what do you predict
about Z effective?
Is it going to increase,
decrease, or stay the same?
It's going to increase.
So Z effective
will also increase.
And if you had a case
that every single time you
had total shielding of
that added electron,
then it might stay the same.
But you're not going to have
this case, this extreme case
of total shielding.
So if Z increased,
Z effective is also
going to increase
as you go across.
And because n is
staying the same,
you have more or less the
same amount of distance
from the nucleus.
So you just have this
stronger Z effective
and it's holding on
to the electrons.
Now, if you go down a
column, the ionization energy
decreases.
So in this case, you're
also increasing Z.
But now you're
increasing n as well.
And so when you increase n,
you have your 3s and you go
to your 4's and your 5's.
And so now, those other
orbitals are way far away
and you have a much bigger
effective radius here.
The Z is getting bigger.
But it's not really
reaching as strongly out.
So here, the effect
of increasing n
is making a much bigger
difference than increasing Z.
So going across, we have this
increase in ionization energy
because Z effective
is increasing
while n is staying the
same or Z is increasing
while n is staying
the same, which means
the effective is increasing.
And going down, it's really n
that dominates that pattern.
So you have a decrease
because you're
going to higher and higher n.
So let's look at some of those.
And we can go back and look at
it what I showed you before.
I said I'd get back to
sodium and potassium here.
So if we consider all these,
if we remove one electron,
then we're going to go to
a noble gas configuration.
So with our first ionization
when we're over here,
we're going to go.
And so when we do
that, then we say,
why are these numbers different
for the second ionization?
We have a noble gas
configuration after we've
lost one electron in each case.
But then we can say, OK,
well helium is up here,
then neon, then argon.
So the ionization
energy is decreasing
as we go down here
because n is increasing.
So we see that trend
in our plot over here.
There's a couple other things
that we can see in this plot.
So we also see that for boron,
this fourth ionization energy
is really big.
And it's bigger than beryllium's
third, which is bigger
than lithium's second.
So let's think about
why that's the case.
And that is another
clicker question.
OK.
Let's just do 10 more seconds.
Oops.
All right.
I was actually expecting
a lower number for this.
That's awesome.
Right.
So it turns out 1 is true.
But all of these
other ones are also
going to be the same
because they've just
lost more electrons.
So all of them have
the same configuration.
So that doesn't explain
what's going on here.
And this is also true.
But binding energies
are always negative.
That does not explain anything.
So the thing that explains
the trend is this one.
Even though they all now
have the same configuration,
it's going to be a
lot harder to pull off
the electron from the one that
has the biggest Z effective
because that's going to
be bound more tightly.
Great.
So you're getting the hang
of these types of questions.
All right.
So those are some of the trends.
And, of course, when
there are trends,
there is always glitches.
These aren't really exceptions.
They are more glitches.
And we can rationalize
them pretty easily.
So again the trend, increasing
ionization energy across,
decreasing iron energy
down, the increase
across as the Z effective
increase, and down
is the increase in n.
But when you actually look
at ionization energies, which
are often reported in
kilojoules per mole, versus Z,
you see that it's not just
kind of a straight line here.
And if we put the elements on
here that these correspond to,
we see 1s1, 1s2, a
drop to 2s, and then
we're doing another
2s, a drop to 2p,
and then so on as
you go up along.
So let's look at some of
these little glitches.
Why isn't this a
straighter line here?
So I'm now going to blow up
this region on this slide here.
And I can just put up
this diagram again.
And you can see that it's true.
So we're talking about the
first ionization energies here.
We see lithium is lower
then it goes higher then
it goes down again.
So that's that little
trend over here.
So why is this the case?
So the ionization
energy for beryllium
is a bit higher than the
ionization energy for boron.
And so it turns out
that this glitch then
is we're going from
the 2s to the 2p.
And 2p, It's easier to
pull off that electron.
So that's why you have this
lower ionization energy.
We have another
glitch over here.
Now we're just within p.
So what's going on there?
And it's very small.
It's a very small
little difference.
But here, the ionization
energy for nitrogen
is bigger than for
oxygen. So it's easier
to pull off an
electron from oxygen.
And if you draw out
your diagram here,
is nitrogen where we
have obeyed Hund's rules
and we put everything
in parallel.
But for oxygen,
we have one extra.
So we had to pair the electron.
So it turns out
it's a little easier
to steal this 2p
electron because it's
the first one paired.
And I kind of think about
that as, again, sort
of the bus where everyone sits.
You can sit two people per seat.
One person sits down, no one
else wants to sit next to them
until all the seats are taken.
And sometimes when you're
sitting in the seat,
you're really, really, happy
when that person gets up
who's sitting next to you.
Maybe there's another
seat available.
They move over to another seat.
So it's often easier to
eject the second person
from the seat.
There's a little bit of
repulsion going on there.
Everyone's working.
They're moving their
arms as they're
doing their chemistry
homework, at least
the buses I'm on anyway.
[LAUGHTER]
So that's why there's
a glitch there.
All right.
So this is all well and good.
We have our trends.
But I always like to think
about how do we know any of this
is really true?
How do you actually measure
these ionization energies?
And so we're just going to talk
about one method for measuring
these for a minute.
So photoelectron
spectroscopy, PES,
is used to determine
ionization values.
And so you can have
some energy that you
will use to excite
something like neon, which
is gas which lights up a
sign for a pizza place,
and you can inject
an electron from it
that has a certain
amount of kinetic energy.
And what you actually
measure in this technique
is the velocity of the electron.
But from velocity, as you know,
you can get kinetic energy.
And from kinetic energy, we
can get ionization energy.
So let's look at this experiment
and think about the electrons
being ejected.
So we have, again, our
neon configuration.
And we'll lose one
electron from p here.
And it will have a velocity
and a kinetic energy.
We can also think about
losing an electron from s.
And we're just going to lose
one electron per shell here.
And we can lose an electron
from the 2s and the 1s.
And all of those should
have distinct velocities
and distinct kinetic energies.
So if we measure velocity,
calculate kinetic energy,
then we can find the
ionization energy
if we knew the energy that
we used to excite the neon.
So the incident energy
equals ionization energy
plus kinetic energy.
Or rewritten,
ionization energy equals
the incident energy
or initial energy
minus the kinetic energy.
So we can use this to
calculate ionization energies.
And this should look
awfully familiar to you.
It's very similar
to something that
will be an exam 1 where we're
talking about using photons
and shooting them
at metal surfaces
and ejecting electrons
that have kinetic energy
if the energy used to hit the
metal is greater in energy
than the work
function and the extra
comes off in kinetic energy.
This is basically all
the same idea here.
All right.
So in this particular
case, you would
measure three different
velocities or three
different kinetic energies.
And now we can think about
what those should probably
correspond to using our
chemistry knowledge.
And some calculations here.
So we have these three
different kinetic energies.
We know the energy
of the incident.
So we can do some math.
And when we subtract those,
we get one energy of 22.
And then this kinetic
energy is less.
So we're going to get a higher
ionization energy of 48.
And this is really small.
And so now we get an
ionization energy of 870.
And so we might not
necessarily know which
orbitals these correspond to.
But if we think
about it, you should
have the lowest ionization
energy to take an electron out
of 2p, next would
be to 2s, and then
the hardest electron to
eject would be from the 1s.
And these are pretty
similar to each other.
But this is a much
bigger number over here.
And so that's kind of
consistent with what we know.
All right.
So this is how you measure it.
And again, this is a
multi-electron system.
And so then the energy is going
to depend on the two quantum
numbers.
It depends on l and n.
It matters what specific
orbital you're talking about.
So let's just think about
another problem here.
Suppose you had five really
distinct kinetic energies.
Assume that a very
distinct kinetic energy
means a different subshell.
And so then we want
to think about what
are the possible elements
in the periodic table that
could produce a spectrum
with these five very,
very distinct kinetic energies?
And so the way you
think about this
is you want to find what
elements are going to have five
different kinds of orbitals.
And so we can list the first
set here-- 1s, 2s, 2p, 3s, 3p,
that's five.
And then you need to know
from the periodic table
where are the elements
where you're filling the 3p.
And those are over here.
So again, we talked and these
problems aren't on the exam.
But on the exam,
you need to know
that this is 1s-- you're
filling 1s, you're filling 2s,
you're filling 2p,
you're filling 3s,
you're filling 3p-- that
you need to interpret.
You'll be given
a periodic table.
But you need to be able
to know what orbitals
are being filled in
the different parts
of the periodic table.
So let's just try a
practice with that.
All right.
Let's just take 10 more seconds.
All right.
So we might need to work
on the sort of counting.
So again, you want to think
about you have 1s, 2s, 2p, 3s,
3p, 4s, 3d, and 4p.
So I think if we get the
counting down we'll be good.
But again, you need to
look at the periodic table
and know what's getting filled.
All right.
So let's move on and talk
about electron affinity.
And maybe we can squeeze in some
electronegativity at the end.
These are very related
topics and pretty fast.
All right.
So electron affinity-- the
ability to gain electrons.
So what we're
talking about here is
how likely atom X is to grab
an electron and become X minus.
So we often think
about halogens when
we're talking about
this like chlorine.
So we have Cl plus an
electron, Cl minus.
And here the change
in energy associated
with gaining that electron is
minus 349 kilojoules per mole.
Energy is released.
And that means that the ion is
more stable than the parent.
So chloride is very
happy to become Cl minus.
And so you think about
energy being released,
if you think about a kid--
my husband's out of town,
so I was watching our
six-year-old daughter
this weekend.
And she was racing around
like a crazy person
until she collapsed in a heap.
So energy is released.
And she became a more
stable six-year-old.
So that's what's happening
with chloride as well-- more
or less.
So here, the electron affinity
is minus the change in energy.
So if we talked about the
electron affinity of chloride
for the electron
to become Cl minus,
you would say that was plus
349 kilojoules per mole.
So unlike ionization energy,
which is always what?
Positive or negative--
ionization energy?
Always positive.
Electron affinity can
be positive or negative.
And that tells you
something about how much
it wants to gain electrons.
So nitrogen plus an
electron, going to N minus,
has a positive energy value
here and has a negative electron
affinity.
So N minus-- the minus one ion
is less stable than its parent.
So it is not as happy
as chloride to do this.
So trends in ionization.
Usually you have an
increase going across
and a decrease going down.
And let's just
consider noble gases
and what you think about them.
So we'll do one final
clicker question.
It should be pretty fast.
OK, ten seconds.
OK, yup.
They are, in fact, negative.
And so we can think
about this over here.
Noble gases have negative
electron affinities.
Noble gases are very
happy the way they are.
If you had to add
another electron to them,
you would need to make
a new subshell there,
which they don't want to do.
And so halogens,
on the other hand,
which are right next door,
have highest electron affinity.
So if you're over here, they
want to gain an electron
and become a noble gas.
Noble gases want to
stay the way they are.
So the increase trend
ends right before you
get to the noble gases.
They're in their own category.
OK.
So we're going to end with that.
And we'll continue with
electronegativity on Friday.
