>> host: Welcome back
to "Chemistry Matters."
In our last episode,
you were asked to make
predictions about the patterns
of electronegativity
and atomic radius.
So, let's go back
to our classroom and see
how your predictions compare
to those of our students.
>> professor: Let's start
with electronegativity,
which is the attraction an atom
has for additional electrons.
On the monitor here, you can see
a photo of the first team's
conclusions.
Can you tell us what you found
regarding the trend
for electronegativity
as you move down a group?
>> student: We looked at
our Bohr models of lithium
and sodium to figure out what
happens to electronegativity.
>> student: We know that when
effective nuclear charge
decreases, the attraction
between electrons
and the nucleus decreases.
>> student: The effective
nuclear charge of sodium
is less than lithium's
because the valence electron
is farther away from the nucleus
and is shielded
by more core electrons.
>> professor: Wow, that's
very impressive.
You guys will have a great
handle on the important
vocabulary,
effective nuclear charge
and shielding.
So, let's add this
electronegativity trend
to our periodic table.
And now, let's see what
Danika and Corrina have found
regarding the atomic radius
as we move down a group
on the periodic table.
>> student: Once we read the
definition of atomic radius,
it was pretty easy to see
that it increases
going down a group.
>> student: Right!
Lithium is in period two,
so it only has
two energy levels.
Sodium is in period three,
so it has three energy levels.
That extra energy level
makes sodium bigger
than lithium.
>> student: We felt like we
didn't have to use
effective nuclear charge.
>> professor: That's
exactly right.
Great job, guys.
The extra energy level
is the most important reason
why sodium is larger
than lithium.
But there are a couple
of other factors that we need
to consider that make it work
to be a larger atom.
There's more shielding,
so the effective nuclear charge
is gonna be weaker,
and that lesser attraction
between the valence electron
and the protons is gonna mean
it's not held as closely,
or as tightly, to the nucleus.
And so, you also have additional
electron-electron repulsions
that are going on out
in the electron cloud.
So, as the electrons try to move
away from each other,
they're also gonna contribute
to a larger atomic size.
So, let's add atomic radius
to our periodic table,
and we can look at the trend
there as we go down a group,
and then we'll be ready
to tackle the direction
of these trends as we move
left to right across a period.
>> host: Now that we've
determined the periodic trends
for ionization energy,
electronegativity
and atomic radius as we move
down a group,
let's figure out
what the trends are
as we move across a period.
This time we'll need to compare
two elements in the same period,
so we'll compare the Bohr models
of lithium and Beryllium,
both in period two.
You've already seen
the Bohr model of Lithium
from our students,
and you might have drawn
or constructed one yourself.
Now, we need to construct
a Bohr model of Beryllium
before we study periodic trends
across a period.
So, back to our classroom
to finish up with periodic
trends, moving left to right
across a period.
>> professor: So, let's
get started.
Take a couple of minutes
to draw a Bohr model
of Beryllium so that we can
compare its structure
to Lithium.
>> student: That's easy!
You just add one more electron
to the Bohr model for Lithium.
>> professor: That's right,
but don't forget that there's
also another proton
in the nucleus of Beryllium.
And that extra proton
is going to increase
the effective nuclear charge.
>> student: I'm just gonna copy
the Bohr model for Lithium
and add another electron
to the energy level too.
>> student: Okay, but you need
to put +4 in the nucleus
because Beryllium
has four protons,
and add one more neutron
to match the mass number.
>> student: Let me see.
Beryllium's got four electrons,
so we need to put two electrons
into the first energy level
and two electrons
in the second energy level.
>> professor: Okay, if
you're done, please hold up
your models
so that we can compare
the electronic structures
of lithium and beryllium.
And it looks like you're all
becoming experts at making
Bohr models.
Let's look at a close-up
of Lithium and Beryllium
so we can compare their
atomic structures.
We'll show the models
side by side, just like they
are on the Periodic Table.
And so, as you look at this,
What are some differences
you notice between
the Bohr model of lithium
compared to beryllium?
>> student: The main difference
I see is that Beryllium's got
one more electron and proton.
>> professor: That is
the main difference.
Let's remember how we approached
determining the periodic trends
When we were focusing on
going down a group.
We focused on differences
in effective nuclear charge
between the elements.
So, which of the two
do you think has a stronger
effective nuclear charge?
>> student: Last time
the elements we compared
were in different periods
and had a different number
of energy levels.
Lithium and Beryllium each
have two energy levels,
so there won't be
a difference in shielding.
>> professor: Yes, you've done
a very good job of ruling
out shielding as a factor
because when we're
talking about elements
that are in the same period
that's very important.
And so, with that in mind,
which one do you think
has a higher
effective nuclear charge?
>> student: That makes it easy.
Beryllium has more protons,
so they feel a stronger pull
towards the nucleus.
>> professor: Very good.
Effective nuclear charge
actually increases
with each added proton
to the nucleus,
so keep that in mind
as we move across the period
when you're trying to figure out
the trends.
So, I'd like you to brainstorm
with your partner
for a few minutes,
and see if you can predict
whether ionization energy,
electronegativity,
and atomic radius
either increase or decrease
as we move left to right
across a period.
But a word of caution.
Atomic radius can be tricky,
so be careful with that one.
>> student: Let's do
ionization energy first.
That's the energy you need
to remove a valence electron.
The stronger the attraction
between the valence electron
and the nucleus, the higher
the ionization energy.
>> student: In that case,
attraction gets stronger
left to right.
So, ionization energy
must get higher left to right.
>> student: Let's figure out
electronegativity.
Isn't the effective nuclear
charge that's going to attract
extra electrons?
>> student: Yeah,
sounds good to me.
So, we'll say that
electronegativity increases
moving from left to right
across a period.
>> student: Atomic size
is supposed to be tricky.
Let's see if we can
figure it out.
>> student: Don't you think
atoms get larger across a period
because you're adding electrons?
>> student: I guess we'll
find out.
>> professor: All right, let's
see what everyone predicted
for the periodic trends
across a period.
So, make sure that you explain
your predictions.
And we'll start of with
ionization energy.
What did you guys find?
>> student: Uh, we figured out
ionization energy.
It increases across a period.
>> professor: Very good.
So, let's remember to justify
our conclusions though.
>> student: We wrote
everything down.
Let's see...
ionization energy increases
across a period
because the attraction
between the nucleus
and the electrons increases,
making it harder
to remove an electron.
>> student: And the attraction,
or effective nuclear charge
increases because there are more
protons being added
as you move across a period.
>> professor: That's a very
good explanation.
Both of you gave
a great explanation there.
So, let's add that trend
to our periodic table.
Okay, as we're moving
left to right we can see
here that ionization energy
will increase as we move
left to right across a period.
So, now let's hear
what Danika and Corrina
had to say regarding the trend
for electronegativity.
>> student: We think that
electronegativity also increases
across a period.
Which is pretty close
to the definition
for effective nuclear charge.
So we think
that electronegativity
will increase or decrease
in the same direction
as effective nuclear charge.
>> professor: Yeah, that's good.
I like that way of thinking
about electronegativity
because effective nuclear charge
does increase across a period
and so does electronegativity.
So, can you remember why
they're both increasing?
>> student: Because the number
of protons is increasing
when we move to the right
across a period,
and that makes attractions
between the electrons
and nucleus stronger.
>> professor: Absolutely.
Okay, so let's add
increasing electronegativity
across a period
to our periodic table now.
And, we're just about done
exploring periodic trends.
So, who'd like to share
their prediction on how
atomic radius changes
across a period?
And, I'm gonna give you a hint.
The trends that we've discussed
so far as we move
across a period,
they're gonna change
because the number of protons
is increasing.
So, think about how an increase
in the number of protons
found in the nucleus
might affect the atomic size.
>> student: We were thinking
that the atoms get bigger
going across a period
just because we're adding
more protons, neutrons,
and electrons.
>> professor: Okay, so I'm
really glad that you
suggested that, because
that's a common misconception.
Because that's kinda what
our intuition tells us, right?
We're moving left--
left to right across a period,
you would think that the atoms
are getting larger,
although what we're doing--
we are adding the electrons
that-that's true, but you also
have to keep in mind
as we go left to right
and the atomic number increases,
that means we're adding
additional protons as well.
So, we need to consider that
when we're looking at
effective nuclear charge,
and how it affects the size.
So, what are you thinking, Anna?
>> student: I think that
the effective nuclear charge
increases across a period,
but I don't know how that would
affect the size of an atom.
>> professor: Let me show you
something that might help
visualize how increasing
the effective nuclear charge
will change the atomic radius.
So, as we add an extra proton
to the nucleus--
okay, as we move from lithium
to beryllium,
this extra proton here
is gonna have a greater pull.
So, the-the electrons
will actually pull in,
sorta condensing the electron
cloud, therefore it'll make
the atom actually smaller.
>> student: Does that mean
that the atoms keep getting
smaller across a period?
>> professor: It sure does.
Okay? So, let's take a look
at period two
on the periodic table
as we move left to right,
and you'll notice
looking at the period,
that we can predict that neon
would probably have the smallest
atomic radius
of all the elements
in the second period.
Now, let's take a look
at all the trends you've drawn.
Do you notice any patterns?
>> student: Yeah!
Ionization Energy
and Electronegativity always
trend in the same direction,
and Atomic Radius always moves
in the opposite direction.
>> professor: Very good
observation, okay?
So, we're finished exploring
periodic trends,
but you'll probably want
to write these trends down
so that you can refer to them
later when we cover
additional units.
They are a very important
predictor of chemical
reactivity.
>> host: By now,
you have a good understanding
of periodic trends.
If not, feel free to review
this segment until
you're comfortable
with the information.
If you're ready,
join us for our next video
in the unit three playlist
about Electron Configuration.
I'll see you there!