Lesson video

Hello and welcome to this lesson all about state changes and how we can use models to explain them.

My name is Mr. Gundry and I'll be going through this lesson with you today.

I can't wait to get started.

This is a very important point in chemistry and it's something that students often get wrong in assessments because they don't fully understand how the models are used and their advantages and disadvantages, but also how we can use them to explain various concepts in chemistry, so I'm really looking forward to going through this with you today.

So in today's lesson then we are going to be aiming to explain why there is a difference in melting and boiling points between simple and giant covalent substances and we're gonna use models to show the shapes and structures of the covalent substances to help us explain these changes in state.

There are some keywords that will be useful through this lesson.

They are simple molecular, model, intermolecular force, and giant covalent.

On this next slide, you can see the definitions, but we're just gonna go through them as we get to them in this lesson.

This lesson is split into two parts.

We're gonna be looking at simple molecular substances to begin with and then we're gonna go through giant covalent later.

So, as already stated, there are two types of covalent substances.

We have simple molecular and giant covalent.

Simple molecular are very different to giant covalent substances in that simple molecular substances are made up of distinct discreet units.

So examples include water made up of two hydrogen and one oxygen atom, carbon dioxide made up of one carbon and two oxygen atoms, and nitrogen made up of two nitrogen atoms. These all exist as distinct units and so you can see that they will go around as just for water, for instance, two hydrogens and an oxygen atom.

And if we have a large amount of water, we have a large number of these molecules.

These are very simple molecules.

There are some slightly larger molecules that also come under the definition or type of structure such as simple molecular, and that could include things like glucose, which has 6 carbon atoms, 12 hydrogens, and 6 oxygens all in one distinct molecule.

But we can also have macromolecular substances which are also still technically simple molecules which contain huge numbers of atoms, such as proteins.

These do not have a regular lattice structure, a regular repeating unit, and so they would not be classed as giant covalent.

We'll cover that in a little bit more detail later on.

So those diagrams that you just saw on the previous slide they are actually a type of model, and scientists use models to represent the different parts of the natural world normally that are too difficult to observe or explain directly.

So you may be familiar with the classic building of a model volcano to show how a volcano erupts.

We generally wouldn't send school-age students to go and watch a volcano eruption 'cause that would be considered unsafe.

In chemistry, we're often interested in the models of the kind of submicroscopic levels, so things smaller than things we can see in the microscope, so they include things like atoms or molecules, and so it can be quite difficult for us to visualise this.

And so we have, very fortunately, a set of resources that we can use to model the behaviour of atoms. The easiest one in the modern era to use is digital 3D representations, so here we can see a model that I made online of water.

In front of me here, hopefully you can see this, I've got a physical model of water which incorporates the same colour scheme where I have this central oxygen atom and two hydrogen atoms connected either side.

We can use these models to help us visualise the arrangement of atoms in different chemicals and later on we'll use them to help us understand the physical properties.

So here are a couple more ways that we can model the atoms and their bonding in specifically water.

So on the left-hand side we can see we have a physical 3D model and this is showing that the atoms are these kind of hard spheres connected by these physical connectors.

These are supposed to represent the covalent bonds.

Now these are, it's not a bad model, it's not a great model, though, because as hopefully you know, atoms are held together by forces of attraction, not by physical connectors.

So this is a limitation of this model, they aren't actually physical connectors between these two atoms. Another limitation is that these atoms are quite far apart.

The atoms tend to be more kind of overlapping in their electron structure.

But a good thing about this model is that it is 3D because all atoms are three dimensional.

A lot of diagrams show two-dimensional structures, which we can see on this screen here which we'll talk about in a second.

The digital 3D model is very similar.

It represents the bonding between the 3D atoms as sticks, but we can also see the 2D displayed formula where we're displaying all the bonds between the hydrogen and oxygen atoms also shows these sticks between the atoms, which just aren't there, but it's just a a nice visual representation where the sticks represent the covalent bond, that shared pair of electrons.

And then the top here we have this dot-and-cross diagram where we have this overlap between the atoms that it shows the shared pair of electrons between the two atoms in the covalent bond, a much better representation of the bond itself, but still it has its limitations as a model because now we've lost the 3D nature of the diagram and we've got a 2D drawing.

We've lost the rest of the electrons, at least there are some electrons.

However, in the 3D models and the displayed formula, we have no showing of electrons, so again another limitation of those models there as well.

So on the screen you can see there are three models.

I'd like you to have a look through them and I'd like you to have a think about which type of model you think best shows the 3D arrangement of atoms in a carbon dioxide molecule.

So the answer is B, this is a ball-and-stick diagram, that's the kind of unofficial official name we give to this 3D structure.

And we can see that there are double bonds holding those carbon and oxygen atoms together.

It's very similar to the other two, but A has a 2D drawing, as does C.

It's just a nicer way to visualise the 3D structure.

So that was just a nice recap or a reintroduction to models in chemistry.

How can we use them to help us explain the properties of substances then? Well, simple molecular substances have low melting and boiling points.

This means that they are mostly gases or liquids at room temperature.

Because simple molecular substances are discreet kind of units, they are made up of a specific number of atoms, normally a very small number, we can see here with the example of hydrogen that they don't really have a way to connect together very well.

So I'm talking lots of molecules coming together.

How do they hold themselves together? Well, they're held together by these weak intermolecular forces of attraction.

Intermolecular meaning between the molecules.

We mustn't get confused between the actual strong covalent bonds holding those atoms together inside the molecules.

So the hydrogen atoms here are held together by very strong bonds between the two atoms, but then when they come close to other molecules of hydrogen, there's a very, very weak force of attraction.

We call the weak intermolecular force.

So for simple molecular substances to melt or boil, so that's to go from the solid to the liquid state or the liquid to the gas state, those weak forces have to be overcome, so we have to use enough energy to force these molecules away from each other.

Bear in mind that the strong forces, the bonds between the atoms, do not break when state changes happen.

We're gonna comment on that a little bit later as to why.

So I'd like you to match these two terms, strong covalent bond and weak intermolecular force, to the correct positions on the diagram.

Give you a bit of time.

So A is pointing towards the dash line between oxygen on one molecule and hydrogen on the other, and B is pointed at the physical line, the kind of solid line between the hydrogen and the oxygen on just one molecule.

So the strong covalent bond is B that's between the atoms in the molecule, and then the weak intermolecular force of attraction is between the atoms of different molecules.

So not very much energy is required to overcome these forces and we're going to say that it is relatively low energy because we're comparing it to other types of structure.

So all the giant structures, ionic, metallic, and covalent, require far more energy to change state.

And so, compared to those, relatively low energy is required for simple molecular substances.

So melting and boiling, as well as condensing and freezing, are all physical changes.

These end up with the same chemical substances at the end of them, just in a different state.

They're not chemical changes.

Chemical changes are much different.

They require the breaking and the making of bonds.

We'll look at that on the next slide.

Here it's important to note that when we are kind of taking these particles further away from each other, which is what happens as substances go from solid to liquid to gas, they start to vibrate more, they start to move more, they push away from each other.

As they do this, they're overcoming those forces of attraction.

When the opposite happens, so when they start to condense back into a liquid from a gas and then they start to freeze into a solid, these intermolecular forces start to reform and that's what brings the atoms and molecules closer together.

So to break covalent bonds, we need large amounts of energy.

Generally, when bonds are broken, we're talking about chemical reactions.

So here we can see that we've got two molecules of hydrogen reacting with one molecule of oxygen to make two molecules of water.

There are no hydrogen-to-hydrogen bonds in our products, which means those hydrogen-to-hydrogen bonds in the reactants on the left-hand side of this reaction must have broken to form these new bonds.

And the same for our oxygen bonds.

Those oxygen-to-oxygen bonds don't exist in our products.

So large amounts of energy are required to overcome these and far less energy is required to just change the state, so to overcome those weak intermolecular forces of attraction.

So there are some statements here on this page, I'd like you to have a read through them and decide which correctly describe the melting process of simple molecular substances.

I'll give you a bit of time to read through those, and so press pause now and then press play when you are ready to continue.

So the correct statement here is the intermolecular forces between the molecules must be overcome.

Remember, the covalent bonds within the molecule do not break when a state change happens and new chemical bonds form only when a chemical reaction takes place.

State changes are physical changes, not chemical changes, so we don't form new chemicals in the process.

So here's a slightly longer extended task for you now, this is part one and there's another question coming in a second.

Some students are discussing here why water has a higher boiling point than methane.

I'd like you to discuss who is correct and who has misunderstood.

Pause now whilst you have a read through those and press play when you're ready to continue.

So only Laura was originally correct.

So she says that water must have stronger intermolecular forces.

Aisha is wrong, unfortunately.

She says, "Water's boiling point is higher because it's lighter than methane." It's nothing to do with its density as to whether it has a higher or lower boiling point, it's to do with those strength of intermolecular forces.

Alex says that "water boils at a higher temperature because it's a liquid and methane is already a gas." Well, again, he's got the right idea, methane will become a gas before water does, but that's not because it's already a gas.

Again, it's the link to those intermolecular forces.

And Jacob originally said that "more energy is needed to break the bonds in water compared to methane." Remember, it's not the bonds that break, it's the intermolecular forces between molecules that must be overcome, so there's a slight correction here because Jacob was very close to being correct but fell into the classic mistake that students fall for when writing answers to this question.

So here's a table for you to complete where I'd like you to draw some dot-and-cross diagrams for methane and for carbon dioxide.

And then for each of the three types of model, I'd like you to say what represents the covalent bond.

Okay, you probably need a bit of time for that, so pause here as you complete this task and when you're ready to move on, press play.

So on the screen you can see the two dot-and-cross diagrams between the carbon and the hydrogen atoms. For methane, there is just one shared pair of electrons.

One electron comes from each hydrogen atom and then one electron from carbon.

Carbon had a total of four electrons on its outer shell to begin with, and now it has eight.

Hydrogen only had one and now it has two, so both now have a full stable electron configuration.

For carbon dioxide, we have two double pairs of electrons between carbon and each oxygen atom.

Again totaling eight for carbon, but also eight for oxygen now, so both have a stable electron configuration.

For the space-filling model where we have these kind of large, or enlarged, so should I say, is this overlap between the atoms where there is the electron pairs.

And for the ball and stick, it's the actual sticks that represent the covalent bonding, so that pair of electrons.

And then in the dot-and-cross diagram, it's actually the dot and cross together that's the bonding pair of electrons, and that sits in the overlap between the atoms. So now we're gonna look at some giant covalent substances.

I've got some more models like I had my water molecule here, we've got some more models that we're gonna look at together in a second.

And we're gonna talk specifically about some more carbon-based structures, but we are gonna remind ourselves that there are other giant covalent substances as well.

So the two main types of giant covalent that we've looked at are diamond and graphite, but there's also silicon dioxide, which has a very similar structure to diamond.

It has instead of just carbon atoms, it's got silicon and oxygen atoms. For every one silicon atom, there are two oxygen atoms, and they kind of sit in this repeating pattern also in a tetrahedral structure like they do for diamond.

In all three of these examples, every single atom is bonded to other atoms in a giant network of repeating structure.

We don't have individual units anymore like we do for simple molecular substances, and so the properties of these structures are going to be much more different.

On the screen but also at hand I have these models that I've built, so this one is of diamond.

So I'll just give you some time to have a look at that.

It's a bit confusing to look at when it's using these physical models, I prefer the 2D diagrams, but the 3D diagrams offer something different.

And here we have our kind of layers of graphite, so you can see we have these kind of honeycomb hexagonal rings and they exist on different layers.

So physical models can be built to represent the lattice structure of giant covalent substances where we have these bonds between the atoms. In these models, we're using these grey connectors to represent these strong covalent bonds.

And then for our graphite structure, we have these weak forces of attraction between the layers, and I've used some purple connectors.

I don't know whether you can see that on these models but it's easier to see on the screen, some purple connectors between the layers.

So very quickly then, can you match up the labels to the correct positions on the diagram? Where are the strong covalent bonds, is it A or B? Or are they the weak forces of attraction at A or B? Well, I'm hoping, as we've very quickly just gone through this, you managed to get that one quite quickly.

So strong covalent bonds are B, they're between the atoms in the structure.

And then we've got the weak forces of attraction at A, they are the dash lines between the layers.

This is where we were talking earlier about how when bonds break, we are normally talking about chemical reactions.

However, for giant covalent substances to melt or boil, the strong covalent bonds need to be broken.

And that is because all of the atoms are connected in a giant network.

Unlike for simple molecules like for our water molecules where we have these intermolecular forces between other molecules that can be easily broken and overcome and pulled apart from each other, in our giant network, we have to break every single one of these bonds to get this to change state.

So that requires a significant amount of energy, relatively higher amounts of energy than for simple molecules.

So very quickly, why does diamond have such a high melting point? Have a read through these answers, pause now, and then when you're ready to move on, press play.

So the answer is B, because all atoms are held together by strong covalent bonds.

There are no weak forces of attraction in diamond.

I know there are in graphite, but it's not the weak forces that dictate the kind of melting point or boiling point.

It's the atoms, all the atoms have to be broken apart.

We don't kind of separate the layers and the layers then become different states.

Every single atom has to be pulled away from its neighbours, so we have to break every single covalent bond.

And it's nothing to do with its ability to conduct electricity or heat, although they are some useful properties which are covered in another lesson.

We're not gonna talk about those in today's lesson.

So another question for you, a little bit more reading here.

So how does the melting process of giant covalent structures like diamonds differ from that of simple molecular substances? So I'm gonna give you some time to read through this, so pause now and when you're ready to move on, press play.

So giant covalent structures melt at higher temperatures due to strong covalent bonds, so A is incorrect.

The weak forces of attraction are broken through various means, but it's not when they melt or boil.

That's not what's happening specifically for giant covalent substances.

Remember, only some have these weak forces of attraction.

Diamond doesn't have any, so it's not B.

D states that simple molecular substances do not melt, but rather break down into individual atoms. That is not true.

When simple molecular substances melt, the molecules kind of overcome forces of attraction between them, the actual bonds between atoms remain intact.

So the answer is C, then.

Giant covalent substances require the breaking of strong covalent bonds to melt, needing much higher temperatures.

Now the reality of the situation is that most giant covalent substances when they melt don't melt in the way that you would expect.

Now for simple molecular substances, when these guys melt, they're still these distinct units.

It's just how close they are together and how strongly attracted they are together that dictates whether they are in the solid, the liquid or the gas state.

However, for things like diamond to melt, all the bonds need to be broken.

Which therefore means it's not diamond anymore, because diamond is this very specific arrangement of carbon atoms. And when diamonds melt, they form a very similar kind of soup-like substance of carbon that graphite and coal and other carbon allotropes also form when they melt.

So under standard pressure, diamond doesn't really actually melt at all anyway because the conditions required for melting require quite high pressures.

And so, actually, most likely it will happen at the very high temperature of 3,600 degrees Celsius, or give or take a few degrees.

The diamond actually will sublimate and it will turn directly into gaseous carbon atoms and particles.

So carbon is unlikely at standard kind of conditions to actually turn into a liquid anyway.

It's more likely to sublimate and turn directly into a gas.

So extremely high temperatures and pressures will be required for us to be able to melt diamond, to get it to a liquid, and we're talking way above 4,000 degrees Celsius.

And the pressure's extremely high too.

If we were able to actually melt diamond, we would have a very kind of soup-like, dense, carbon-rich fluid where the carbon atoms would still be bonded to other carbon atoms, but there would not be this regular arrangement anymore and they wouldn't all be connected to each other.

Going slightly above and beyond here.

So this is an example phase diagram that we can see that as we increase temperature on the X axis or as we increase pressure on the Y axis, we can see how we can manipulate carbon into various different states.

And so we can see that graphite exists as a solid at low temperatures and low pressures, but at low temperatures, if we increase the pressure, we could potentially turn graphite into diamond structure.

But again, if we then increase the temperature, we're likely to turn it into a vapour at low pressures.

So, true or false, giant covalent structures like diamond will melt in the same way as simple molecular substances.

True or false? Well, this is a false statement.

Here are two justifications as to why it's a false statement.

I'd like you to pick the correct one, so pause now whilst you read through them and press play when you're ready to continue.

Well, the correct answer is that diamond requires extremely high pressures to sublimate or, if possible, melt under high pressures.

So when diamonds melt, they form a molten diamond is incorrect.

Molten diamond is not a thing.

So what happens when you cool the either vapour of carbon atoms or this kinda soup-like, dense, carbon-rich fluid if you've managed to make it? Well, as you cool it down, if you're cooling it down at standard temperature and pressure, so room pressure, you're probably gonna get graphite or some kind of amorphous carbon structure, like coal.

You're not gonna get diamond, which is a very highly-ordered structure.

You would need incredibly high pressures to be able to get that to happen.

In reality, though, if we were to heat diamond and graphite just in the room, so in a lab, for instance, then if we were even able to get it to those high temperatures, we would find that our diamond and graphite were burning far, far sooner than we would even be able to get them to melt.

And that is because these are carbon-rich structures.

Carbon will combust with oxygen in an oxidation reaction to form things like carbon dioxide, and this will happen at a much lower temperature, about 850 degrees Celsius, far, far lower than the actual melting point of the structure.

Graphite is a little bit more resistant to combustion in its bulk form than diamond, and that's one of the reasons why we use it as a refractory material, so we use it to line ovens and kilns as a way to protect them from the heat in some industrial processes.

Due to the delocalized electrons within graphite, it's a very good thermal conductor, so it can pass the heat quite easily through its structure.

However, under the right conditions, it will still combust.

And so if we were to grind it down into a powder, powdered kind of carbon-based structures are incredibly flammable and can actually be quite explosive when ignited, when there is a good supply of air that contains oxygen.

So, true or false, both diamond and graphite will burn when exposed to high temperatures in the presence of oxygen instead of melting.

What do you think? Well, the answer is true.

Here are two answers that could justify the answer as to why it's true.

I'd like you to pause the video as you read through these and then press play when you're ready to continue.

So the correct answer is A, the melting point of diamond and graphite is much higher than the temperature at which they combust.

Whilst graphite is an excellent conductor of heat, it doesn't burn very easily in an oxygen-rich environment as a bulk form.

If we powdered it down, if we ground it down into kind of tiny little particles, then yes, it would actually burn very easily.

But in a large chunk it won't.

So we get now to the final set of tasks for this lesson.

So here is a paragraph of text.

You can see in the square brackets there are some choice words for you to complete the sentence.

And then for question two I'd like you to explain why diamond has these properties and then I'd like you to explain why graphite has those properties.

Once you've had a go at that, you can press play on the video, but for now pause the video as you complete these tasks.

So, compared to the simple molecular substances, giant covalent structures like diamond and graphite require much higher temperatures to melt or boil because all the strong covalent bonds need to be broken.

This is in contrast to simple molecules where melting and boiling involve overcoming weaker intermolecular forces.

Diamond is very hard as it has a rigid structure of carbon atoms bonded to four other carbon atoms. These strong covalent bonds need a relatively large amount of energy to break them, giving diamond a high melting point.

Diamond has no delocalized electrons or ions, so cannot conduct electricity.

Graphite is made up of layers of hexagonal rings of carbon atoms that are covalently bonded to three other carbon atoms. These layers are held together by weak forces of attraction and require a relatively small amount of energy to overcome.

This makes graphite quite soft.

To melt graphite, all of the strong covalent bonds need to be broken, requiring a large amount of energy.

This is a high melting point.

Graphite can conduct electricity as it has free-moving delocalized electrons which can carry a charge or current.

So, in summary, molecular modelling represents the bonding and structure of substances.

Diamond is very hard and has a high melting point and does not conduct electricity where every atom is bonded to another.

In graphite, each atom is bonded to three others to form layers of hexagonal rings and each atom contributes a free electron.

Simple molecules are held together by intermolecular forces, not chemical bonds.

There are chemical bonds within the actual molecule itself holding the atoms together, but between molecules it's just weak forces.

And the way that giant covalent structures melt and boil is very different to simple molecules, as we saw earlier on in the lesson.

Thank you so much for learning with me today.

I've had a great time and I hope you have too.

I hope it's been very informative and I look forward to seeing you in the next lesson.