>> host: Welcome back
to "Chemistry Matters."
In the last video
on this playlist we focused on
the electrons in an atom.
We learned that the Bohr model
of the atom is unique
and valuable
because it describes electrons
as existing in different
energy levels
outside of the nucleus.
We also discovered
a relationship between
the number of electrons
in an atom and that atom's
location on the periodic table.
In this episode, we'll focus on
the relationship between
the properties of an atom
and the atom's location
on the periodic table.
You'll see that the particular
period and group
an atom belongs to
is not only related
to how many energy levels
and valence electrons it has,
but how large or small
the atom is and how likely it is
to gain or lose electrons.
In order to better understand
periodic trends,
it's helpful to construct
Bohr models of the atom.
Before we get too far,
I want to remind you
of a couple of guidelines
that are important to follow
when making a Bohr model.
Although the Bohr model
focuses on placement
of electrons into different
energy levels,
it's also important to include
the number of protons
and neutrons in the nucleus
of the atom.
The interaction between
the electrons and protons
determines many atomic
properties.
Each energy level can hold
a maximum number of electrons.
This table shows the maximum
number of electrons for each
energy level.
The first energy level,
closest to the nucleus,
can only hold
up to two electrons.
The second energy level
can hold up to eight.
The third energy level
can hold up to eighteen
electrons.
And the fourth and fifth
energy levels can hold
up to 32 electrons.
Larger atoms with more electrons
can exist, but these five levels
are all you need to understand
the naturally occurring elements
found on Earth.
This Bohr model shows
the same information
in graphical form,
and it's available in our
"Chemistry Matters" tool kit.
You might be wondering why
we care about periodic trends.
Understanding periodic trends
helps us to predict
the properties of an atom,
and to determine how it
will chemically react with
other atoms to produce all sorts
of useful compounds,
like medicines, fertilizers,
and cosmetics.
Atoms are arranged
on the periodic table
in such a way that certain
properties increase or decrease
moving across a period
and down a group.
Our goal in this unit
is to identify and explain
these trends, so we can predict
the chemical reactivity
of elements.
Our students will work in teams
to make Bohr models of two
elements in the same group, 1A.
Those elements are lithium
and sodium.
So, let's go out to our
classroom to learn more
about periodic trends.
>> professor: Okay, let's get
started.
Anna and Charlie,
use your whiteboards to draw
the Bohr models of lithium
and sodium.
And Danika and Corrina,
I'll have you guys use
the magnetic models to create
the same two Bohr models.
We all need to be able to draw
the Bohr models,
but the magnetic models offer us
the advantage of keeping
the relative sizes
of the atoms correct.
When you're done,
we'll compare the electronic
structure of lithium and sodium,
and as you construct
your models, be sure you
remember to put the correct
number of protons and neutrons
in the nucleus so we can discuss
the strength of the attraction
between the protons
and the electrons.
>> student: Let's start
with the nucleus.
Lithium has three protons
and sodium has 11.
>> student: The first energy
level has two electrons,
so let's put those
in both atoms.
>> student: Now, lithium only
needs one more electron,
so that goes into the second
energy level.
>> student: Let's see.
Sodium has a total of 11
electrons, so I need to add
nine more to this model.
>> student: The second energy
level only holds
eight electrons, though.
>> student: Okay, so that means
that the last electron
goes into the third
energy level.
That's gotta be right
because sodium is in group 1A
and has one valence electron.
>> professor: Okay, did you
finish your Bohr models?
Let's have both teams hold up
their Bohr models so that we can
compare lithium and sodium.
What differences do you see?
>> student: Sodium's got
a lot more electrons.
Lithium only uses
two energy levels,
and sodium has three.
>> student: Is that why
lithium is in period two
on the periodic table
and sodium is in period three?
>> professor: Yep, Good job!
What about the nucleus?
What's different between
the two nuclei, Danika?
>> student: Lithium only has
three protons and sodium has 11,
kinda like they're electrons.
>> professor: Yes, that's right.
It's important to look
at the nucleus
because there's a strong
attraction between protons
and the electrons,
and that attraction is gonna
affect the atom's properties.
Which brings us to
the introduction of a new term:
effective nuclear charge.
A higher effective nuclear
charge means that there's gonna
be a stronger attraction
for the electrons
to the nucleus.
As you're looking
at your models, which element,
Lithium or Sodium,
has a stronger attraction
between its valence electrons
and the protons?
You need to be careful!
There's a lot to consider
when answering this question.
>> student: Sodium has more
protons than lithium.
Does that mean that its
electrons feel more
attraction to the nucleus?
>> professor: Sodium does have
more protons, but there's
a little bit more
to the story.
Let's take a look here
at the Bohr model
for a sodium atom.
Here is our Z sub eff,
which represents affective
nuclear charge.
Now, because its valence
electron is farther away
from the nucleus than lithium's,
that valence electron
actually feels less attraction
to the nucleus.
The inner electrons here,
which we refer to
as core electrons,
are also gonna repel or shield
that interaction from that one
valence electron to the nucleus,
and that's gonna make
that attraction even weaker.
>> student: So, the farther
away from the nucleus,
the weaker the attraction?
>> professor: Yes.
The higher the energy level,
the weaker the attraction,
or the smaller the effective
nuclear charge.
But remember, the distance
from the nucleus
and the shielding by the core
electrons, it lessens
the effective nuclear charge.
>> host: Now that you've been
introduced to the idea
of effective nuclear charge,
this is a good place
to introduce three atomic
properties that are of interest
to chemists:
ionization energy,
electronegativity,
and atomic radius.
Soon we'll look for
periodic trends for each
of these atomic properties.
Which often happens
in chemical reactions.
Ionization energy
is an important characteristic
of atoms because it helps
explain why some atoms lose
just one electron when bonding,
why some lose two,
or even why some atoms
don't lose electrons at all.
And is an important property
of an atom when making
chemical bonds.
We'll explore chemical bonds
in depth in unit four,
and you'll see why
electronegativity is possibly
the most important elemental
property that we will study.
We can make so many predictions
about how elements will interact
with each other
if we have a good understanding
of electronegativity.
Atomic radius gives us an idea
of the size of an atom,
and is found by measuring
the distance between the nuclei
of two touching atoms,
and then dividing that distance
in half.
There are some unusual trends
in the sizes of the atoms,
as we will soon see.
With these three properties
of an atom in mind,
let's go back to our classroom
to see how these properties
change throughout
the periodic table.
>> professor: Okay, let's look
again at our Bohr models
of Lithium and Sodium
and let's see
if we can predict how our three
atomic properties will differ.
We'll discuss the trend
of ionization energy
moving down a group
on the periodic table first,
and then you'll explore in teams
the trends of electronegativity
and of atomic radius
as we go down a group.
So, let's start by looking up
close at Bohr models of
lithium and sodium here
while we try to predict
their ionization energies.
So, who'd like to predict
what happens to the ionization
energy as we move from lithium
down to sodium
on the periodic table?
>> student: Does a high
ionization energy mean
it's hard to remove
a valence electron?
>> professor: That's right.
A high ionization energy means
that the electrons are held
more tightly by the atom,
and therefore it's harder
to remove that electron.
>> student: Sodium's valence
electron is in a higher
energy level than lithium's,
so its electron is farther away.
Electrons feel less attraction
when they're farther away,
so I think it'll be easier
to remove the valence electron
from sodium.
>> professor: Yes, that's right.
Electrons that are farther away
from the nucleus are gonna have
a lower ionization energy.
When we talked about effective
nuclear charge earlier,
remember we discussed
two factors.
We talked about the distance
from the nucleus and shielding
by the core electrons.
>> student: I remember.
The more energy levels,
the more shielding.
So sodium has more shielding
than lithium.
>> professor: Yes, good,
very good.
Now, let's combine these
two factors into one conclusion.
How does the distance
from the nucleus
and the degree of shielding
by the core electrons affect
the effective nuclear charge
and the ionization energies
of lithium and sodium?
>> student: I'll try.
Sodium's valence electron
is farther away from the nucleus
and shielded by more
core electrons than lithium,
so sodium has a lower effective
nuclear charge and lower
ionization energy.
>> professor: Excellent!
So we can conclude
that the ionization energy
is going to decrease
as we go down a group.
We're going to show these trends
on the periodic table
as we discover them,
showing arrows that represent
the trend,
and we'll put them over here
on the left-hand side
of the periodic table.
So, Let's look for patterns
that develop
as we add more arrows.
So, now what I'd like you to do
is work in teams,
and we're going to determine
the trends for electronegativity
and for atomic radius.
All right?
So, Anna and Charlie,
what I'd like you guys to do
is take a look at
electronegativity.
Danika and Corrina,
I would like you guys
to consider atomic radius.
And you may have to refer back
to these two definitions
for those properties,
but your objective
is to determine whether or not
these properties decrease
as we move down a group,
or increase
using the periodic table.
So, use your Bohr models
of lithium and sodium,
along with what you've learned
about effective nuclear charge
to help you.
So, record your conclusions,
along with supporting evidence,
and then I'll put your results
on the monitor and we can share
it with the rest of the class
once you've finished.
>> host: Now's your chance
to practice applying
your new knowledge.
You'll use the Bohr models
of Lithium and Sodium,
and what you've learned
about effective nuclear charge
to see if you can predict
whether the other two properties
we're looking at,
electronegativity
and atomic radius, increase
or decrease when moving
from top to bottom down a group
on the periodic table.
The following questions
should help you guide
your predictions:
What are the definitions
of electronegativity
and atomic radius?
Which has a higher
effective nuclear charge:
lithium or sodium?
Does a higher effective
nuclear charge increase
or decrease electronegativity?
And does a higher
effective nuclear charge
increase or decrease
atomic radius?
If you need help answering
these questions,
review our discussion
while determining the trend
for ionization energy.
The thought process
will be the same.
We'll explore trends
across a period later on.
For right now,
focus on the vertical trends.
We'll see you at the top
of the next video
on the playlist when you've
finished your predictions.