Solids Part II

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Introduction to Chemistry: Structures and Solutions

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Duke University

Introduction to Chemistry: Structures and Solutions

79 ratings

This is an introductory course for students with limited background in chemistry; basic concepts such as atomic and molecular structure, solutions, phases of matter, and quantitative problem solving will be emphasized with the goal of preparing students for further study in chemistry.

From the lesson

Intermolecular Forces and Phase Changes

<p>Welcome to week 5! This week we will examine the phases of matter more closely, looking at concepts such as intermolecular forces and thermal energy. We will also practice calculations of energetics such as those requiring enthalpies of phase transitions. Finally, we will introduce concepts in metal bonding and solid lattices. We made a few illustrative demonstrations for your viewing pleasure; I hope that these are helpful! </p><p>Remember to continue posting any questions, concerns and suggestions to the discussion forum. Also help your peers out on the forum to build a beneficial learning community together!</p>

Solids Part I11:19

Solids Part II10:46

Solids Part III12:10

Meet the Instructors

Prof. Dorian A. Canelas
Assistant Professor of the Practice
Chemistry

Greetings.

In this lecture, I will review ionic solids, and

focus on predicting relative melting points for ionic compounds.

Ionic compounds are the second type of solid that we're considering.

As you can see, I have space here for five types.

Last time we did metallic, this time we're doing ionic.

And ionic compounds have metal to non-metal

bonds, which, you'll remember we've termed ionic bonds.

Ionic bonds are non-directional.

So here's a crystal lattice for sodium chloride, and

each sodium cation is surrounded by six chloride anions.

So we could not reach in to a sample of

sodium chloride and pull out a single sodium chloride molecule.

Because of the way the lattice is arranged,

there are six chloride anions surrounding every sodium cation.

Therefore, we can't take a pair of tweezers and reach into

the sample and pick out a single molecule of sodium chloride.

It's an extended solid in three directions, and therefore, this

is called a compound and usually not called a molecule.

Ionic bonds, of course, are governed by Coulomb's Law, which is this mathematical

equation that describes the electrostatic interaction between

the positively-charged cations and the negatively-charged anions.

Remember the Coulomb's law tells us that the force of attraction or

repulsion depends on the product of this constant, the charges, divided

by the product of the shielding factor, times the distance squared.

Of course for ionic compounds, the charges, they could

be plus one and minus one, or plus one and

minus two, or plus two and minus two, just

depending on the charges of the ions in the species.

Well, those charges would appear in the numerator of Coulomb's Law.

The size of the atoms would appear in the denominator as the distance.

The smaller the atoms are, the closer together they

can get, and so the smaller the interparticle distance.

Those are the two factors that appear in Coulomb's

law that affect the properties of this type of solid.

If we wanted to melt an ionic solid and predict which solid had the higher melting

point, we would have to consider this interparticle force.

So, the ionic compound is this

three-dimensional lattice of particles that have

positive and negative charges, and those

make the ionic bonds, which are nondirectional.

But it's important to note that all ionic bonds are relatively strong forces,

and the stronger the forces, then what would

you predict is true about the melting point?

Would you predict something that has relatively strong interparticle forces

would have a relatively high melting point or a relatively low melting point?

Please go ahead and answer that now.

Thank you for submitting your answer.

The stronger the interparticle force is, the higher

the melting point because it will then be more

difficult to allow those ions to move out of

their lattice if we have a higher interparticle force.

So one of the things you can think about if you look at Coulomb's Law is the higher

the absolute value of the product of the charges,

in other words, product means we're multiplying these together.

And if we have a negative charge and

a positive charge, that's what happens in ionic compounds.

So we have to take the absolute value of that, so that we have a positive number.

The bigger that number is, the higher the melting

point because that would be relatively strong interparticle forces.

Let's review and say that a slightly different way.

So when predicting the melting points of our ionic

solids, we need to look at the ionic bond strength.

And that depends on two things.

The magnitude of the charges on the ions and the separation of the ions.

Both of those factors are found in Coulomb's Law.

Therefore, small ions, which have a small sep, separation, so

this would be coming in the equation as r, that are

highly charged, and, in other words, here I would mean

plus two, as opposed to plus one, for relatively highly charged.

Those have a higher melting point than ionic compounds which have large ions, and

remember, we, we know the periodic trend

for atomic radii, that have relatively low charges.

So for example, some compounds are large cations, like

rubidium, and it might have a relatively low charge.

I would expect that type of ionic compound to have a lower melting point

than ionic compound made from lithium, which

is very small, and has the same charge.

Here's an example.

Which of these two compounds has a

higher melting point, magnesium oxide or cesium iodide?

Please go ahead and try to answer that now.

[BLANK_AUDIO]

Thank you for submitting your answer.

Hopefully,the approach that you took to this problem is applying Coulomb's Law.

Magnesium oxide has relatively small ions when compared to cesium iodide.

And the magnesium ion is a plus 2 charge, and the oxide anion has a minus 2 charge.

Cesium iodide on the other hand has relatively large ions.

Again, that comes from the atomic radii trend.

Those large ions have plus 1 and minus 1 charges associated with them.

So if we look back up at the top,

small ions that are highly charged have higher melting points.

So I would expect magnesium oxide to have a higher melting point.

If I look up the actual melting points, I see that

magnesium oxide's melting point is almost 3,000 degrees Celsius, that's really high.

So, that was the correct answer.

But I also want you to notice that cesium

iodide's melting point is also quite high, isn't it?

Most ionic compounds have very high melting points,

even the ionic compounds that we say are relatively

low melting for being an ionic compound, because

626 degrees Celsius is still a very hot temperature.

That's hotter than your home oven probably gets

unless you look at the burner on the stove.

If you look inside your oven, you probably

can't turn it up to such a high temperature.

You'd have to buy a furnace to go to that type of temperature.

Maybe you have a furnace at home.

I shouldn't make assumptions.

Let's end this lecture with a review example.

Which of the following substances has the highest melting point?

Krypton, beryllium oxide, carbon tetrabromide,

potassium chloride or barium fluoride?

Please go ahead and answer this now.

[BLANK_AUDIO]

Thank you for submitting your answer.

In order to answer this question, we would have to

ask ourselves, what types of solids these different materials would be?

I've gotten a little bit ahead of myself,

because we haven't covered molecular solids yet, have we?

Carbon tetrabromide and krypton are both molecular solids.

Krypton's a noble gas.

And just from the name of the group, that

implies that it's a gas at room temperature and pressure,

so this probably has a relatively low melting point

below room temperature, so let's just cross that one off.

Whenever I put that on a test, students have

trouble with it, I think because of Superman and kryptonite.

People like to say krypton's high melting, but it's really not.

Carbon tetrabromide is not ionic.

So we're going to cross that one off because that one

has only intermolecular forces associated with the temporary induced dipoles.

We'll talk about that more in the next lecture.

So that leaves us some ionic compounds to

choose from, beryllium oxide, potassium chloride and barium fluoride.

Which one of those has the highest melting point?

Well potassium chloride has a plus 1 cation and a minus 1 anion.

Barium fluoride has a plus 2 cation and a minus 1 anion.

Beryllium oxide is plus 2 cation and minus 2 anion, so

just based on the charges, I would go with beryllium oxide.

But before I answer for sure, I would want to think about the relative sizes.

And I'm feeling pretty good about choosing beryllium oxide with

that additional information, because beryllium

oxide also is relatively small ions.

Fluoride here is the smallest, but oxide is almost as

small, and barium, of course, is much bigger than beryllium.

So the highest melting point is beryllium oxide.

This concludes the second part of the introduction to

solids, the examination of ionic solids and their melting points.

We learned how to predict those relative

melting points by going back to Coulomb's Law.

By using the monatomic charges and the atomic radii of the anions and cations.

[BLANK_AUDIO]

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