In this lesson we're going to introduce the ionic bonds and
how those ionic bonds are formed.
If we start out with a sodium ion, and remember that sodium is all the way
to the left of the periodic chart, it has one electron in the s shell.
And, if we take that electron away,
we have satisfied the inner core of electrons.
And as a result of stripping away that electron,
the cation given by the NA+ winds up being smaller because
of the contraction of the cloud associated with the missing of that electron.
The process that is related to the change in the Na to
the Na plus one cation requires a certain amount of energy.
And that energy is referred to as the ionization potential.
So what we're doing here is taking the sodium,
making it a cation, and it requires a certain amount of energy.
Now, we go over to the right-hand side of the periodic chart
where we have the halogens.
And, in this case,
we're looking at chlorine as a typical example of the halogens.
And the neutral chlorine atom will have a specific atomic radius.
If we add an electron, what we wind up doing, then,
is to increase the radius of the electron cloud.
The process of going from the neutral chlorine atom to the chlorine anion
winds up providing some energy to the system, and
we refer to this as the Electron affinity.
So the amount of energy that's released,
as associated with the chlorine neutral atom to chloride anion,
turns out to be on the order of about 4 electronvolts.
So the sum of these two processes,
where we take and make a sodium ion and a chlorine ion,
actually requires an amount of energy that is associated
with the difference in those two energy products.
If we look at the periodic chart, one of the things that we'll see is,
as we go down the periodic chart in a column, so
let's take a look again at the first column where we have sodium, potassium,
rubidium, as we go down that column with increased the atomic number,
what we see is the neutral atom winds up increasing.
And that is given by the symbol of a green atom.
If we ionize, and we strip an electron from any one
of those cations along that column, we find that we reduce the atom.
Going from the atom to the cation, there is a reduction in radius.
And if we look at all of those associated green and reds,
what we see is the corresponding change for the different elements,
where we have the process where we go from the neutral atom to the cation.
On the other side of the periodic chart, where our tendency is to add electrons,
what we then do is to start out with the green atom, and
as a result of creating the anion,
we wind up increasing the size of the anion over the size of the neutral atom.
Now, in addition to seeing a change associated with
the addition of one electron or the removal of one electron,
if we look at an element like iron, for example, we can see that the relative size
of the atom changes as we go through and consider different valence states of iron.
So in the first case we have iron as a neutral atom.
And over to the right, when iron is in a plus 2 configuration,
that is we have lost two electrons, there is an associated reduction.
If we go a little bit further, where we have a second
oxidation state where iron is in the 3+ state,
we see that that progressive change in the number of electrons that we have removed
winds up decreasing the size of the atom into the cations that emerge.
Now, when we look at the periodic chart, what we see is as we go down the columns,
we will see a tendency for the size of the atoms to wind up increasing.
And if we look at the elements in the first column, those are the Alkali metals,
and we look at the second column, those are the Alkaline earths.
These are the elements that are associated with the s electrons.
Now, as we go across the periodic chart,
what we find is there is an increase in the electronegativity.
That is, as we go further to the right,
we being to see the addition of electrons to the associated
neutral atom to wind up forming the anions.
We discussed in this module what happens when we consider
the elements in the group carbon, silicon, and germanium,
those that are associated with the sp3 hybridized orbitals, and
the characteristics that are associated with the diamond cubic,
crystal structures of diamond, as well as silicon and germanium, and
ultimately to their properties that are attributed to the semiconductor behavior.
Now, when we come all the way over to the far right we have that halogen column.
And that halogen column sitting right next to the inner column,
we find that it is very easy for us to add an electron to stabilize that outer orbit.
And then, when we look at the elements that are in that last column,
those are the Noble gases.
Those are the ones that are associated with stable electron configurations.
And if we look at the bottom of the periodic chart, where we have taken
out the rare earths, and now what we're considering with is the behavior of
those rare earths that are controlled by the filling of those outer f orbitals,
going from cerium all the way across for the rare earths.
In the next lesson, we're going to be describing the ionic bond.
And with the ionic bond, there's some very simple models that we can use when
we describe the ionic bond that will lead to an understanding of the development
of bond force curves, and ultimately the development of the bond energy curves.
These two behaviors we will be able to begin to understand
some of the properties and characteristics that are associated with the material
that we'll be dealing with in the remainder of this course.
Thank you.
Thomas H. Sanders, Jr.Regents' Professor
