Lesson video

Hello, my name's Dr.

Warren.

I'm so pleased that you've decided to join me today for this lesson on properties of covalent substances.

It's part of the structure and bonding unit, and follows on from our lesson on how covalent bonds are formed.

I'm here to work with you through this lesson and support you all the way, especially through the tricky parts.

So let's get learning.

Our learning outcome for this lesson is, I can describe some different structures that non-metal atoms can form using covalent bonding and describe the properties of covalent compounds.

We've got some keywords for you, molecule, covalent bond, intermolecular forces, giant covalent structure, and polymer.

Now we're gonna look at these keywords in some sentences.

You may wish to pause the video and copy down these sentences and then refer to them later on during the lesson.

So what have we got? A molecule is a particle consisting of a fixed number of two or more non-metal atoms covalently bonded together.

A covalent bond is a strong electrostatic force of attraction between a shared pair of electrons and the nuclei of bonded atoms. Intermolecular forces are weak forces of attraction between molecules and molecular substances.

A giant covalent structure is a substance that has a large regular arrangement of atoms, all joined together by covalent bonds.

A polymer is a long chain molecule, formed by joining small molecules, or monomers, together by covalent bonds.

So if you want to copy these down, pause the video and then press play when you're ready to start.

In this lesson, we have two learning cycles.

The first is simple covalent molecules, followed by giant covalent structures.

So let's get started with our simple covalent molecules.

Non-metal elements and non-metal compounds form covalent structures.

So this is first really important point, we're dealing with non-metal elements only.

So let's have a look at an example.

Hydrogen gas exists as a diatomic molecule.

It's made from two hydrogen atoms covalently bonded together.

And there we can see a model of a hydrogen molecule, both hydrogens and non-metals.

And for an example of a compound, we've got water.

It's a covalent structure.

This time it is a molecular compound.

There are two hydrogen atoms and an oxygen atom, H2O.

Only non-metal elements are part of covalent structures.

Okay, so let's just check our understanding.

Magnesium carbonate is a covalent structure, true or false? Well done if you picked false.

That is the correct answer, but why? Is it A, covalent structures can be formed by non-metal elements as well as by non-metal compounds, or B, covalent structures can be formed by metal and non-metal compounds? Very well done if you chose A.

In a covalent structure, there are only non-metal elements or non-metal compounds.

No metal elements are present.

Okay, so let's move on to think about the size of atoms. Atoms are very small.

The size of an atom is around 1.

9 times 10 to the minus 10 metres.

That's very, very tiny.

And as you can see from the diagram, the atoms change in size and not all the same size, but they're still very small in that order of magnitude of 10 to the minus 10.

And we can see here from hydrogen is a bit smaller than oxygen.

Nitrogen's a bit bigger than oxygen.

Carbon's a bit bigger than that.

But you can see from the numbers in the table that even carbon, which is the largest atom we can see here, is 1.

7 times 10 to the minus 10 metres and that is a very, very small number.

So how do we measure the size of an atom? Well, we measure the atomic radii.

So how do we know about the size of the atoms? Well, typical atomic radii measure around 10 to the minus 10 metres.

And it's the radii we measure, but it's really difficult to actually measure the radius as such because you don't get single atoms. We know from the reactivity that the atoms react to become more stable, the full outer shell.

So what we need to do is we have to measure half the distance between the nuclei of two identical atoms. So if we take hydrogen, for example, we have our hydrogen molecule, we can measure the distance between the two hydrogen nuclei and then divide it by two, and we will get the atomic radius.

And that atomic radius is gonna be in the order of magnitude of 10 to the minus 10 metres.

So as well as the atomic radii being around that size, we also find that a typical bond length is that size as well.

Now, a bond length is the distance between the two nuclei of atoms in the bond.

And we can have a look at this example.

So if we take the O-H bond, it is the distance between the nuclei of the oxygen and the hydrogen, and that is 0.

96 times 10 to the minus 10 metres.

A second example or more examples are shown in this table where we can see some typical bond lengths.

So example, the carbon single bond, carbon carbon is 1.

5 times 10 to the minus 10, whereas the carbon double bond is 1.

3 times 10 to the minus 10.

So the key bit of learning here is that all the bond lengths measure around 10 to the minus 10 metres.

And the other point to make is simple covalent molecules are also known as simple molecular substances, and they are all very small because they are molecules.

So let's check our understanding of number.

Here are some measurements.

Which could be covalent bond lengths, A, B, C, or D? I'll give you a hint.

Two of them are correct.

So very well done if you picked A and C.

Both of those numbers are similar and they give that order of magnitude of 10 to the minus 10 metres.

So part B, that's incorrect, we haven't got enough decimal places there.

And part D, we've got to be really careful when we're using standard form because that is actually written as 10 to the power of 10 plus 10 metres, which basically is four with 10 zeros after it.

So very well done if you got those right.

Moving on.

Thinking about the properties now of some covalent substances.

And the first thing is they cannot conduct electricity.

So if we were to get a covalent compound and put it into an electric circuit, a bit like you can see here, the light bulb would not light up because there are no free moving charge carriers.

Why? All atoms are linked together by the covalent bonds, so they're tied up in shared pairs of electrons, there are no free moving charge carriers, there are no ions, there are no delocalized electrons.

So in this case, simple covalent substances cannot conduct electricity.

Another property or small covalent molecules, we normally find them as gases or liquids at room temperature.

And we'll also find that because they are gases or liquids at room temperature, they have relatively low melting and boiling point.

Now, an example where we've got a covalent substance that is a solid at room temperature is bee wax, because it melts around 75 degrees centigrade.

Now, we might think that's quite high.

Well, it is compared to the gases or liquids, but in terms of other properties of other types of structures, it is still very low.

And to give an idea of what we mean by low melting and boiling points, have a look at the melting and boiling points of some of the compounds or elements in this table.

So hydrogen, for example, which isn't a covalent compound, it's a covalent element, joined together as the H2 molecule, that has a melting point of minus 259.

1 degrees centigrade.

So it'd have to get really, really, really cold to make that into a solid state hydrogen.

Compare that with ethane minus 182.

8 and ethanol at minus 114.

5 and water, and water is very much the exception because that is zero degrees centigrade.

And you can see as well, especially with hydrogen, that element that the boiling point is very, very close to the melting point, so it only goes into the liquid state at a very, very small range.

So key learning here is that covalent molecules, simple covalent molecules, have relatively low melting and boiling points.

So why? Well, let's have a look at the structure.

So in simple covalent substances, we have weak forces of attraction between the molecules.

So if we look at this diagram, those purple oval blobs, they are representing the molecules that make up beeswax.

And in between them, what we actually have is a weak intermolecular force.

So in between those covalent molecules that are weak intermolecular forces, that's a force in between the molecules.

The strong covalent bonds are within the molecule.

So if we take our beeswax and beeswax, you've got carbons and you've got hydrogens, and I think you've got some oxygens as well, there are strong covalent bonds making that molecule, but the key point here is you have lots of them and they come together in the solid, but the forces between the molecules are really weak.

Okay.

So we're just showing this diagram, quite a simple model of intermolecular forces, and that word inter means in between molecular molecules.

So the forces in between the molecules are weak.

Okay, so let's have another look at.

Let's have a look at a different example.

Simple covalent molecules have low melting and boiling points due to their structure, and we've talked about that already.

So what we've got here is we've got an example, a different type of diagram showing hydrogen gas, and we have our hydrogen atoms joined together by a covalent bond.

And we are showing with these purple dash lines the weak intermolecular forces of attraction between the hydrogen molecules.

Within the hydrogen molecule, joining the hydrogens together, there are strong covalent bonds.

And this diagram, it's basically showing the intermolecular forces as dashed lines.

That's really important.

So they are weak forces in between the molecule.

So let's think about why it's got a low melting point.

Well, first of all, it doesn't take very much energy to break those weak intermolecular forces of attraction.

So those purple dash lines, not a lot of energy is needed.

So at low temperatures, so we saw on the previous slide how low low temperatures mean, there is enough energy for the molecules to exist in the arrangement of a liquid or a gas state.

One point that's really important, and this is something people often get wrong in exams, is covalent bonds between the atoms do not break when changing state.

So whenever we talk about a change of state with simple covalent molecules or simple covalent substances, we are talking about the weak intermolecular forces of attraction breaking between the molecules.

It's not between the atoms. So it's important to get that right.

Okay, so let's just check our understanding.

The melting point of carbon dioxide is low.

Carbon dioxide is a gas, true or false? True.

Well done if you picked that.

Now, why? Is it A, it takes a small amount of energy to break the weak intermolecular forces, so the melting point is low, or B, it takes a large amount of energy to break the strong covalent bonds, so the melting point is high? Well done if you got A.

It takes a small amount of energy to break the weak intermolecular forces, those forces in between the carbon dioxide molecules, so the melting point is low.

Excellent work if you got that right.

So now let's have a look at a different covalent structure, a polymer.

A polymer is a long-chained molecule.

And if we look at the word, poly means many, and you might think of a polygon that you would have seen in math that is a shape of many sides.

Mer means parts.

So what the word polymer means, it means a molecule made up of many parts.

So we have monomers, which are small covalent compounds, and we have lots and lots of, lots of them that react together.

So let's take for example polyethylene.

Polyethylene is an example of polymer.

It's formed from many ethene molecules.

So if we imagine that these purple circles are ethene molecules, they join together with strong covalent bonds between them to form the polymer polythene.

Okay? Polyethylene.

What you can see on that diagram is a section of the polymer because there will be hundreds of thousands of atoms making up that polymer.

So in practise, polymers are used to make things like plastic bags, and polyethylene is a polymer that is used to make shopping bags, just the standard plastic shopping bag you get at the supermarket.

So what's the structure look like? Well, you have lots of polymer molecules which are being represented by these purple lines, all packed closely together, but in between them there are intermolecular forces and it's these intermolecular forces of attraction that hold together the polymer molecules.

So most polymers exist in the arrangement of a solid at room temperature, and that is because there are many, many weak intermolecular forces.

So the attraction is stronger between long polymer molecules, meaning that they are in the arrangement of a solid state.

So compared to smaller covalent molecules, which exist in the arrangement of a gas or liquid state, a lot more energy is required to overcome the strong attractions in the polymer structure.

So we will find that most polymers are in the solid state at room temperature, or in this example here, we are given the melting point of polyethylene as 105 to 115 degrees centigrade.

So let's just quickly check our understanding.

True or false, a polymer is covalent molecule made from lots of different compounds joined together to form a long chained molecule, true or false? Okay, well done if you picked false.

Why? Well, let's have a look at the reasons.

Is it A, a polymer is a long chain molecule, formed by joining together lots of the same monomer, or B, a polymer is a giant covalent structure? Well done if you chose A, a polymer is a long chain molecule, formed by joining together lots of the same monomer.

Right, we come now to our first task.

And what we want you to do to start with in question one is circle the covalent molecules that you can see.

So just circle all the ones that are covalent molecules.

We've got ethene, calcium carbonate, iodine, ammonia, and sulphate, zinc chloride, and hydrogen sulphide.

Moving on to question two, part A, what is the covalent bond? How long is a typical covalent bond? And part C, candle wax is covalent compound, explain why it is not a conductor of electricity.

So have a go at those questions and pause the video while you do, and then we'll have a look at the answers together when you press play.

Okay, so let's have a look at the answers.

Which are covalent molecules? Right.

So we've got ethene, iodine, ammonia, and hydrogen sulphide.

Remember, covalent molecules only have atoms that are non-metal.

What is a covalent bond? Well, it is a strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms. How long is a typical covalent bond? It's about 10 to the minus 10 metres, so it's very, very small.

And for candle wax, why doesn't it conduct electricity? Well, all the atoms in a wax molecule are joined together by covalent bonds.

So there are no free moving charge carriers to conduct electricity, no ions and no delocalized electrons.

So very, very well done if you've got all those answers correct.

Okay, question three.

In terms of structure and bonding of the following simple covalent substances, explain why A, hydrogen exists as a gas at room temperature, and B, ethanol exists as a liquid at room temperature.

Pause the video while you have a go at this question and then we'll have a look at the answer together.

Okay, hydrogen exists as a gas at room temperature because it exists as a diatomic molecule, H2.

And diatomic molecules are very small.

So only a small amount of energy is needed to overcome the weak intermolecular forces or the weak forces of attraction between the hydrogen molecules, so that they have enough energy to exist in the gas state at room temperature.

So well done if you got all of those points.

When you're answering these questions, try and think through a logical structure to your answer, logical sequence.

Okay, part B, ethanol exists as a liquid at room temperature because again, ethanol exists as small molecules and only a small amount of energy is needed to overcome the weak intermolecular forces or weak forces of traction between the ethanol molecules, so they have enough energy to exist in the liquid state at room temperature.

So if you got that answer right as well, very, very well done.

Okay, so that brings us to the end of our first learning cycle on simple covalent molecules.

And now we're gonna move on to think about giant covalent structures.

So covalent structures can be formed by non-metal elements.

An example of this is diamond.

We have there a nice picture of a diamond, and then we've got the particle diagram of a diamond and it shows a giant covalent structure because it goes on and on and on.

So diamond exists as a giant covalent structure made from a large undefined number of carbon atoms. So that's important.

We have no idea how many carbon atoms are in that giant structure.

Each carbon atoms shares a pair of electrons with four other atoms, and they make this tetrahedral shape, which is really, really strong.

Covalent structures can also be formed by non-metal compounds.

And an example of this is sand or silicon dioxide.

That is a giant covalent structure as well.

But this time it's made from an undefined number of silicon and oxygen atoms. And you can see again in that particle diagram, the silicon atoms, which are in yellow, and the oxygen atoms which are in red.

And they form this large structure of undefined number of atoms. So unlike the simple covalent structures, giant covalent structures have very high melting points and boiling points.

So looking at the two examples, diamond, well, if you want to melt a diamond, you have to heat it up to 3,550 degrees centigrade, which is enormously hot.

And if you wanted to boil it and get it into the gas phase, and I'm not sure why anybody would want to do that, you'd have to heat it to almost 5,000 degrees.

Where silicon dioxide sand is a bit lower, but it's still a very high number, 1700 for the melting point and for the boiling point, 2230 degrees centigrade.

So very, very high temperatures.

And in fact, silicon dioxide, when it is melted, does become glass.

It has some glass properties.

So why? The particles can exist in the arrangement of a liquid or gas state, but many strong covalent bonds in the giant structure need to be broken and breaking all these covalent bonds requires a large amount of energy.

So to break up those structures, you have to actually break the covalent bonds, break all those shared pairs of electrons, and that is why it takes so much energy because they are very, very strong.

Okay, so, let's just check our understanding of all of this.

Which of the following statements are true about giant covalent substances? A, electrons are shared between the atoms, B, they contain a large undefined number of atoms, C, sodium chloride has a giant covalent structure, and D, they have high melting and boiling points.

So make your choices.

Okay.

So if you got A, well done, that's correct.

They're covalent, so they share electrons between atoms. Well done if you got B as well.

A large undefined number of atoms make up that giant structure.

We don't know exactly how many atoms, but there are loads and loads of them.

And finally, if you got D as well, well done.

They have high melting and boiling points.

Sodium chloride is not a giant covalent structure because it contains metal sodium ions.

So let's have a look at some of the properties, well, most covalent structures cannot conduct electricity.

There, we've got our structure of diamond, showing our carbon atom and our covalent bonds that go on throughout the structure.

All the atoms are linked together by strong covalent bonds.

So as we previously heard, there are no free moving charge carriers, no ions, no delocalized electrons, therefore it cannot conduct electricity.

Okay, so we're going to have another look at a slightly different giant covalent structure.

It's another type of carbon.

This time, graphite, and you can see a picture of some graphite there.

Each carbon in graphite is covalently bonded to three other carbon atoms, and that leaves one delocalized electron from each atom.

So very different bonding to that found in diamonds.

So we've got a diagram here.

We have got the layers, and you can see each layer has those carbon atoms with three covalent bonds coming out of it.

And in between the layers we have weak forces of attraction.

So this delocalized electron is really important because it changes the properties of the carbon.

And in fact, if we were to put a graphite rod into an electric circuit, we would find that the light bulb lit up because it does conduct electricity.

And this is because of that delocalized electron between the covalently bonded carbon atoms can carry charge through the structure.

So a big difference between the properties of diamond and graphite.

So let's have a quick check of understanding.

True or false, carbon is covalent substance, so it cannot conduct electricity, true or false? Well done if you chose false.

Now, what's the answer? What's the reason, A or B? The atoms in diamond are linked together by covalent bonds with no free moving charge carriers, or B, graphite has delocalized electrons between the layers of covalently bonded atoms. Well done if you picked B.

The reason that carbon can conduct electricity is when it's in the form of graphite, the delocalized electrons there between the layers.

Right.

This comes now to our second task, task B.

And some students are discussing why the two compounds shown have very different melting points.

So let's have a look at these diagrams. First, we have silicon dioxide.

We can see it's a giant structure.

It has a melting point of 1710 degrees centigrade, whereas water is a simple covalent molecule and it has a melting point of zero degrees centigrade.

So on the next slide, I'm going to show you the discussion the students are having.

And the question is, which students do you agree with, and can you correct any misunderstandings? So Lucas is saying there are more Si-O bonds than there are O-H bonds.

Sofia is saying, the Si-O bonds are stronger than the forces between the H2O molecules.

Izzy is saying the Si-O bonds are stronger than the O-H bonds.

And Andeep is saying the forces between the Si2O molecules are stronger than the forces between the H2O molecules.

So which students do you agree with, and can you correct any misunderstanding? So pause the video while you have a go at this question, and then we'll look at the answer together.

Okay, let's go round and see which students we agree with.

First of all, Lucas, no.

There are more Si-O bonds than O-H bonds.

Well, the O-H bonds do not break when ice melts, so it's kind of irrelevant.

It's the weak forces between the molecules that break.

So I'm afraid Lucas was not right.

What about Izzy? Well, unfortunately, Izzy was not right either.

Again, similar reason.

The O-H bonds do not break when ice melts.

It's the forces, the weak forces between the molecules that break, and that's a really important point that neither Lucas or Izzy has totally understood.

So is anybody right? Yes, Sofia is right.

Well done, Sofia, you've got this right.

The SiO bonds are stronger than the forces between the water molecules.

So it's gonna take more energy to break those bonds and it will.

So silicon dioxide will have a higher melting point.

And what about Andeep? Well, I'm afraid Andeep is not quite there either.

And what we'd say to Andeep is silicon dioxide is a giant covalent structure, so the strong covalent bonds take more energy to break than the weak forces between the water molecules.

So well done if you've got all of that right, especially if you manage to get the reasons for the misunderstandings as well.

Great work.

Right, we'll move onto our second question.

Diamond and graphite are different forms of carbon.

They both exist as giant covalent structures.

So there's three parts to this question.

A, what is a giant covalent structure? B, which is a conductor of electricity, diamond or graphite? And C, in terms of structure and bonding, explain your answer to B.

So again, pause the video while you have a go at this question.

So let's have a look at the answer.

Well, a giant covalent structure is a substance that has a large irregular arrangement of atoms all joined together by covalent bonds.

Part B, graphite, is a conductor of electricity and diamond isn't, so really well done if you got those answers correct.

So C, we're gonna give the reasons here.

So let's think about diamond first.

In diamond, each carbon atom shares a pair of electrons with four other carbon atoms. So there are no free charge carriers, so it cannot conduct electricity.

However, in graphite, each carbon atom is covalently bonded to only three other carbon atoms, leaving one delocalized electron.

And it's that delocalized electron that can carry the charge as they move between the layers of carbon atoms so it can conduct electricity.

So really, really well done if you were able to get that explanation right and linked to the structure and bonding of graphite and diamond.

Great work.

So that brings us to the end of this lesson.

So we're just going to summarise the key learning points.

First of all, a typical atomic radii and bond length are in the order of 10 to the minus 10 metres.

Non-metal elements and non-metal compounds form covalent structures.

Simple covalent substances have no free moving charge carriers, so they cannot conduct electricity, although some giant covalent compounds can.

Small molecules are usually gases or liquids with relatively low melting and boiling points.

Giant covalent structures have high melting and high boiling points because strong covalent bonds need to be broken.

I hope that you have enjoyed today's lesson and I look forward to learning with you again very soon.