- [Voiceover] Ionic bonds are the bonds
that hold together ionic compounds.
So basically it's what holds together cations and anions.
An example of a compound that's
held together with ionic bonds
is sodium chloride, also known as table salt.
So here we have a close-up picture
of some really nice crystals of sodium chloride.
And this is something that you could try at home.
You can take any table salt that you might have,
dissolve it up in some water,
and then let that water slowly evaporate.
And if you're lucky, you might get some
beautiful symmetric crystals like these.
For me personally at least,
growing crystals that look beautiful
is one of the most fun things about chemistry.
You can see when we look at the
close-up shapes of these crystals
that they have some very beautiful symmetry.
That symmetry tells us a little bit
about the structure of these compounds
on a molecular level.
If we zoom in on these crystals,
we can imagine, actually we don't have to imagine,
you can look at these with different kinds of instruments
like x-ray crystallography,
and you can look at the crystal lattice
and get information about
how the different ions are arranged in the solids.
So the way that the ions are arranged
determines a lot of things
about the properties of these compounds.
So these ionic bonds and how the ions are arranged
tell us a lot about the solubility of the compound.
Solubility.
And other properties like melting or boiling points.
And it even can be related back to things
like how hard a particular ionic solid is.
So the ionic bonds here, in the sodium chloride,
are the ones that hold together
our sodium ions and our chloride ions.
So our sodium plus and our chloride minus.
And the strength of an ionic bond
is related to the electrostatic force.
The electrostatic force between them.
And I'm going to abbreviate the electrostatic force
as F subscript e.
So this is the force between two charged species.
And it's equal to some constant k...
Times the two charges that are interacting
divided by the distance between the two charges squared.
So here, q1 and q2 are the charges,
and in the case of sodium chloride for example,
q1 and q2 would be...
q1 might be one plus, from our sodium ion,
and q2 might be one minus, from our chloride ion.
And we could also just switch those two,
we could say chloride is q1 and sodium is q2,
and that wouldn't change what we get from this equation.
And then r2 here, is the distance between the ions,
and we usually approximate it as saying
it's the sum of the ionic radii...
for the two ions we're looking at.
So we can use Coulomb's law here
to explain some properties that are related
to the strengths of ionic bonds.
And so the example we're gonna go through today
is going to be that of melting point.
So we're gonna look at some melting point trends
and try to relate them to the
different variables in Coulomb's law.
So the first thing we'll look at,
the first two compounds we'll compare
are sodium fluoride...
and magnesium oxide.
Sodium fluoride has a melting point of 993 degrees Celsius,
and magnesium oxide has a melting point
of 2852 degrees Celsius.
The other information we know about these two compounds,
if you look up the ionic radii,
it turns out that sodium fluoride,
the distance between the ions
is about the same as magnesium oxide.
They're not exactly the same, but they're pretty close,
so if we were to say that r is
approximately the same for these two,
then we can explain the difference
in melting points using the charges.
Since melting point is a measure of,
basically how much energy do you need
to add to these compounds to break apart your ions,
we would expect melting point to go up,
to increase, as Fe increases.
As the force between the ions increases,
we would expect to have to add more energy
to break those ions apart.
And we can see that in our first example.
Magnesium oxide, if we look at the charges on the ions,
magnesium is two plus, and oxide is two minus.
In sodium fluoride, sodium is one plus,
and fluoride is one minus.
So we would expect, assuming that r is about the same,
this q1 times q2 is four times bigger
in magnesium oxide versus sodium fluoride.
So q1 and q2, the product of q1 and q2...
is higher for magnesium oxide,
and that's why we would expect
the melting point to be higher.
We can also look at sodium chloride versus sodium fluoride.
And in this case, let's look at,
well, I don't know, maybe this is kind of artificial,
the boiling point, the melting point,
sorry, the melting point of sodium chloride
is 801 degrees Celsius...
and the melting point of sodium fluoride is,
like we said early, 993 degrees Celsius.
And so this time, the charges are the same on our ions,
our q1 and q2 is one plus for the sodium in both compounds,
and one minus for the chloride and the fluoride.
So q1 times q2 didn't change for these two compounds,
but since we changed the anion from fluoride to chloride,
we increased r here,
and increasing r in the denominator makes
the electrostatic force goes down.
Another way we could put it is that since r decreases,
as we go from sodium chloride to sodium fluoride,
the melting point goes up.
So in each of these pairs,
the compound that has the higher melting point
is the one that also...
has the higher electrostatic forces,
and that's either because the charges are higher,
q1 and q2 are higher,
or because the distance between the ions went down.
So these are some examples for how we can relate
the properties of ionic compounds
to the electrostatic force using Coulomb's law
between the cation and the anion.