Lesson video

Hello and welcome to this lesson on bonding structure and properties from the Chemistry of Carbon unit.

You might already have some experience working through the ideas of the different types of bonding and their structures and relating them to their properties.

So this might be a nice revision or recap lesson for you, but you may also have never come across these ideas before.

And so we're going to go through them in as much depth as possible to help you fully understand these core ideas from the key concepts in chemistry.

So the main outcome from this lesson is I'd like everybody that is watching to be able to explain how bonding structure and properties are linked together.

And we're going to work through some related long answer questions by the end of this lesson.

There are five key words or phrases for this lesson.

The first one is bonding, force of attraction, structure, property and charge carrier.

On the next slide are the definitions, but we are just going to go through them as we go through the lesson.

So this lesson is split into three parts.

The first part we're going to look at bonding, the different types of bonding.

We're going to talk about structure and then we're going to link it all together at the end by talking about how properties are influenced by bonding and structure.

So let's start by looking at bonding.

So by now you may know that chemical bonds are formed between atoms to increase the stability of substances.

So to become stable, a lot of elements either gain or lose or even share electrons to achieve a full outer shell.

And that full outer shell for most atoms other than for hydrogen and helium is eight electrons.

For hydrogen and helium, they have a very small shell that can only be filled up to two electrons.

Some elements are already full.

Their outer shells already have the maximum number of electrons, and so we say they're inert because they're already stable.

And that is the group 0 elements, that's helium, neon, argon, xenon, krypton, et cetera.

We're going to recap the three main types of bonding, which are covalent, ionic, and metallic.

On the right-hand side of this slide, you can see two dot-and-cross diagrams that show the bonding between sodium and chlorine in the ionic substance, sodium chloride, and the covalent bonding between carbon and oxygen in carbon dioxide.

So true or false, all atoms will chemically react to form bonds with other atoms.

True or false?

Well, that is a false statement.

And to justify your answer, here are two statements.

I'd like you to pause now.

Have a read those once you are comfortable with your answer.

Once you want to hear the answers, press play.

The answer is B.

Some atoms already have a full outer shell of electrons, so they do not react and they're already stable.

So the first type of bonding we're going to look at is covalent bonding.

Now covalent bonding is categorized by a strong force of attraction, an electrostatic force of attraction between pairs of negatively charged electrons and the positively charged nuclei of the bonded atoms.

Often this is just shortened to the covalent bond is a shared pair of electrons, but there is this electrostatic force holding them in place.

Methane, CH4, is made of one carbon atom and four hydrogen atoms.

Each of the four hydrogen atoms shares its one electron in its outermost shell.

Carbon has four electrons in its outermost shell.

Each of them achieve a full outer shell by bonding and achieves a more stable state.

So we have a shared pair of electrons between each of the carbon and hydrogen atoms.

The same is true here for our water molecule.

Water is made up of two hydrogen atoms and an oxygen atom.

And the shared pair of electrons exists between each of the hydrogen and oxygen atoms.

Hydrogen has access to two, that's its full shell and oxygen has access to two from each of the bonds and four of its remaining electrons, meaning that it has access to a total of eight.

So it has a full shell of electrons.

The next type of bonding is ionic bonding.

Now ionic bonding is between two ions.

The ionic bond itself is not the formation of the ions, although that is part of the process.

So on the left-hand side we have a potassium and a fluorine atom.

Both have an incomplete shell of outer electrons.

Potassium has one in its outermost shell.

Fluorine has seven.

To form lithium and fluoride, sorry- to form potassium and fluoride ions, potassium will kind of donate.

It will give its outermost electron to the fluorine atom leaving behind a full shell for outer shell of electrons for potassium and completing the shell for fluorine.

Together they form two separate ions and make up the structure that is potassium fluoride.

We call and ionic bond or we describe an ionic bond as the strong electrostatic force of attraction between oppositely charged ions, positive and negative ions specifically.

Now we'll look later at the structure of ionic substances because this implies there is one ion of each.

However, you may know already in giant ionic structures, there are many of each, but we'll go through that a little bit later on in this lesson.

The final type of bonding is metallic bonding, and this is slightly different to the rest.

This is categorized by what's known as a sea of delocalized electrons.

These surround central metal cations in a irregular pattern.

So we've got our positively charged metal ions known as cations, and we have our sea of delocalized electrons.

Now lithium only has one electron in its outermost shell.

So for every one nuclear- oh sorry, every one positively charged metal ion, there is going to be one electron in that sea of delocalized electrons.

Calcium has two electrons in its outermost shell.

So for every one calcium ion, there would be two negative electrons delocalized in the system.

The description of a metallic bond then is very similar to ionic bonding, so be very careful.

There is also a strong electrostatic force of attraction, but this time it's between the negatively charged electrons and the positively charged metal ions.

And like with our ionic bonds, the strength of this bond is uniform.

It's the same in all directions.

So quick question for you.

Which type of substance involves the transfer of electrons from one atom to another?

Is it metallic, ionic, covalent, or all three?

Well, the answer is ionic.

There are obviously ions that form in metallic structures, but the electrons are not being passed from one atom to another.

They're kind of delocalized around the system.

Another question for you, true or false in a covalent bond atoms share electrons?

Well, the answer to that is true.

And here are two statements for you to have a think about to justify your answer.

Pause now as you read through them and press play when you'd like to know the answer.

Well, the answer is B.

Covalent bonds involve the sharing of pairs of electrons, one electron from each atom, whereas A is incorrect.

It says there can be any number of electrons shared between atoms.

There can be two, there can be four, there can be six, but there can't be kind of any more than that really, or any less than that.

There have to be the pairs of electrons.

So here we've got a nice visual description, a nice diagram showing the similarities and differences between the three types of bonding.

So atoms bond in different ways depending on what type of atom they're also bonding to.

So in covalent bonding, there are non-metal atoms only.

Whereas in ionic bonding we generally have metals and non-metal atoms bonded together.

Be careful that's not the definition.

Often students write the definition of a covalent bond is a bond between two non-metals.

That's not enough.

We have to talk about the electrostatic forces when we're talking about bonding.

Whereas metallic bonding is slightly different.

While it is a type of bond, the difference between metallic bonding and the other types of bonding is that the electrons kind of just float.

They exist as a sea between the different atoms of metals, whereas in ionic bonding and covalent bonding, they are very much attached to atoms.

Either, you know, shared space between them or they exist only on one of the atoms.

They're transferred in the context of ionic bonding.

So some slight differences, but ultimately they all involve the movement and the sharing or transfer of electrons.

So I put that into words for you here.

We've got a nice table summarizing the differences and similarities between different types of bonding.

So covalent bonding is between atoms sharing one or more pairs of electrons, whereas in ionic bonding it's between oppositely charged ions and then for metallic bonding, it's between metal cations and delocalized electrons.

In ionic bonding and metallic bonding, the bond goes in all directions.

The force is experienced uniformly in all directions, whereas in covalent bonding, it's specifically between the two atoms that are sharing those electrons.

In covalent bonding, it's non-metals with non-metals or sometimes with metalloids.

Ionic bonding is between metals and non-metals, again, sometimes with metalloids, whereas with metallic bonding, it's only metals with other metals.

We'll learn a little bit more in a different lesson about this, but when metals form bonds with other metals, they actually form a mixture called an alloy, which is not quite the same as what happens when covalent bonds and ionic bonds form.

We form new substances called compounds.

So again, another big difference between metallic bonding and covalent and ionic bonds.

All three types of bonds are incredibly strong and depending on the different circumstances will depend on which one is the strongest.

But generally covalent bonding and covalent substances have stronger forces between them than ionic and metallic.

In the context of giant structures, which we'll talk about a little bit later, and as we talked about, the electrons are localized between two atoms in covalent bonding or in the actual atoms that formed ions for ionic bonding.

Whereas for metallic bonding, the electrons are delocalized, they can freely move within the structure.

So which of these statements are true then about ionic bonds but not metallic bonds?

I'll give you a little bit of time to read through these, so pause now and then press play when you would like to check your answers.

Well, the electrons involved are not delocalized in ionic bonds.

They are for metallic.

There are positive ions in ionic bonds, but there also are in metallic bonding.

The bond involves electrostatic forces of attraction.

That's both of them.

In fact, this also includes covalent bonding.

And the final one, the bond is made up of oppositely charged ions.

That is the only true statement for ionic bonds because there are also negatively charged ions.

Whereas in metallic bonds, the negatively charged substance is the electrons.

True or false, each sodium atom is only ionically bonded to one chlorine atom in sodium chloride?

We're going to go through this in a little bit more detail when we go through structure, but we have talked about it so far.

So true or false?

The answer is false and there are two statements here to help you justify your answer.

So have a think about these, pause if you need some time, press play when you're ready to know the answer.

Well, the answer is B.

The electrostatic attraction between ions is uniform in all directions.

That means it doesn't change in strength depending on which atoms or ions in this case are around it.

All of the closest ions are attracted equally.

It's not covalently bonded, that's a different type of bonding.

So I've got four questions for you here.

Some longer tasks are probably going to take you a short while, so you probably want to pause whilst you have a go at these questions.

Press play when you would like to find out the answers.

So we're going to go through these questions then.

So the first question asks you to describe how a lithium atom and a chlorine atom interact to form a lithium ion and a chloride ion.

Well, one electron, the one electron, on lithium's outer shell is transferred to the outer shell of chlorine.

This means that lithium forms a positively charged ion and negatively charged chloride ion forms.

Question two asks you to define a covalent bond and ionic bond and a metallic bond.

So a covalent bond is the electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms.

An ionic bond is the electrostatic force of attraction between oppositely charged ions, the positive and negative ions.

And a metallic bond is the electrostatic force of attraction between metal cations and a sea of delocalized electrons.

You need to be able to draw a dot-and-cross diagram.

So this is just a bit of recap, hopefully, if not, go and watch one of our other videos on drawing dot-and-cross diagrams.

So covalent bonds between water and oxygen.

We saw water early on as an example in the slide deck, but we've also got oxygen here which forms a double bond between its two atoms, whereas lithium chloride and magnesium oxide are two ionic structures.

So we need to draw them with their square brackets and charges to represent the number of electrons that have been lost or gained.

And the final question was to draw a labelled diagram showing metallic bonding.

So we should have at least six positively charged metal ions here.

We've used lithium, but we could have just written a plus in the center and we've drawn our sea of delocalized electrons.

So you don't need to draw this kind of wafty sort of blue outline, but just drawing the dots or dashes and labeling them as a sea of delocalized electrons is necessary for this diagram.

So that's bonding.

We're now going to talk about how that leads into structure.

So there are lots of different types of structure in chemistry.

We have simple molecular structures.

So these involve specifically covalent bonds where we have a small defined number of atoms and they make these things called molecules.

So methane, water, oxygen, these are all examples of molecular structures, but we also then have giant structures.

Now giant structures can be made up of covalent, ionic or metallic bonding.

Here we've got an example of sodium chloride.

This is a giant ionic structure and they have large undefined numbers of atoms or ions, but they do have a regular repeated structure.

The atoms in molecules then are held together by very strong covalent bonds.

These are known as intramolecular bonds.

So we can see here that's the kind of slightly greeny kind of blue line between the two hydrogen atoms.

But molecules are also held together by these weak forces of attraction.

So we call them intermolecular forces of attraction.

That's between the molecules.

They're very weak and most molecular substances are gases at room temperature because the energy required to overcome those forces is very small.

We don't need very much energy to overcome them, and so they're not very attracted to each other at room temperature.

And so they are as far apart as they possibly can be, which makes them a gas.

They are in a gas state.

So true or false molecules are held together by weak intermolecular forces of attraction.

The answer is false.

Here's two statements, pause now as you read through them and press play when you're ready to hear the answer.

Molecules are held together by weak intermolecular forces of attraction.

Don't be confused here.

Molecules aren't held together by strong bonds.

The atoms in molecules are held together by strong covalent bonds.

The molecules themselves are weakly attracted to each other and we call those intermolecular forces.

So another type of structure is a covalent structure.

So they're the two types of covalent substance, simple molecular or giant covalent.

And we've got two examples here.

We've got diamond, which you might have seen earlier on in this unit when we looked at the allotrope of carbon.

And we've got another type of giant structure called silicon dioxide or silica or quartz, which is very similar in terms of the structure to diamond.

They both have what's known as a tetrahedral structure.

So that gives us a tetrahedron where we have four-sided pyramid sort of shape around a central atom.

We've got another type of structure which is a giant ionic structure.

And so we have two examples here of the same substance.

We've got two different ways of representing them a ball-and-stick model versus a 3D space-filling model.

And we can see here that in an ionic substance we have this regular repeating pattern of positive then negative, positive then negative charge.

And so that central negative charge on both the diagrams you can see we've got on the ball-and-stick model, we've got a line being drawn to a positive and negative charges that represent an ionic bond.

But be careful because that central negative ion you can see is attracted to all of the positive ions around it.

So our ionic bond isn't just between one positive and one negative ion.

It's uniform in all directions, attracting all of the positive and negative ions around it, depending on which one around it.

And then the final structure we have is a 3D structure of metals.

They are a giant structure again where we have this regular repeating pattern, but it's just of positive metal ions with a sea of delocalized electrons.

We can see here that the positive metal ions in this lattice structure are incredibly densely and closely packed.

This means that metals have a very high density.

So I'd like you to read through these statements and select all that have giant structures based on what we've just talked about.

Pause now if you need a bit of time and press play when you're ready to hear the answer.

So we've got some structures here.

We've got an ionic substance that does have a giant structure.

They only ever have giant structures.

Metallic substances also only have giant structures.

Molecular substances don't.

They're the other type of structures.

We have simple molecular substances versus giant and then covalent lattices such as graphite.

They are also a giant structure.

Question for you, what do all giant structures have in common?

So again, pause the video now if you need to read through the statements and press play and you'd like to hear the answer.

Well, all giant structures have a regular repeated structure.

They contain a large undefined number of atoms or ions that could be held together by ionic bonding, but they might also be held together by covalent or metallic bonds.

And not all of them are technically 3D.

So graphene classes as a giant structure, but it's just a 2D plane of single atoms in a very large undefined structure.

Most are 3D structures but some are not.

So I've got a slightly longer task for you here.

First of all, I'd like you to complete the table and then I'd like you to use the phrase forces of attraction to describe the following substances.

And then I'd like you to suggest one reason why this model is a good representation of the structure and bonding and one reason why it is not.

I'm going to give you a bit of time to pause the video now and then press play when you'd like to find out the answers.

So the types of possible bonding we've got simple can only be found with covalent bonding, whereas giant bonding, I'm sorry, giant structures can be found with all three types of bonding.

The atoms within molecules are held together by strong covalent bonds.

These are intramolecular forces, but the molecules are held together by weak intermolecular forces of attraction.

And for giant metallic substances, strong electrostatic forces of attraction hold metal cations and delocalized electrons together.

And then for the final question which looks at this bonding model, it suggests one reason why this model is a good representation of the structure and bonding in ionic substances and one why it's not.

Well, what's good is it shows there's a regular repeating pattern, and it shows us that there are the positive and negative charges of the ions.

The cons are that it shows space between the ions when there isn't any space between the ions.

They are space-filling model is a much better model.

Or, annoyingly, it shows physical sticks between the ions which wouldn't realistically be there.

These are supposed to represent the bonds which are just forces holding those two ions or those many ions together.

The final section of this lesson then looks at the properties, and we're going to talk about how bonding and structure can be used to help explain properties of the four different types of structure.

So the type of bonding and structure influences physical and chemical properties.

Bonding informs structure which then informs properties.

And understanding this relationship is key to predicting and explaining the behavior of different materials.

True or false.

So what we've been working on so far today, the terms bonding and structure can be used interchangeably.

That is false.

Here are two statements.

Pause the video now if you need a bit of time to have a think about it, and then press play when you would like to hear the answer.

Well, the answer is bonding helps inform structure, which in turn helps properties.

Structure does refer to whether something is metallic, ionic or covalent.

But specifically we are referring to giant structures which all three of them can have or simple structures which only covalent can have.

So bonding doesn't refer to whether it's giant or simple structure, that structure, so they're the wrong way round, basically, in that description.

Simple molecules have relatively low melting and boiling points.

This is because of the molecules being held together by weak forces of attraction.

We touched on this a little earlier.

It doesn't take very much energy to overcome those weak forces between them.

Whereas in giant structures, the atoms or ions are held together by strong forces of attraction called bonds.

On the left, that's diamond.

They're held together by strong covalent bonds.

In the middle we've got an ionic substance that's held together by strong ionic bonds and on the right it's metals held together by strong metallic bonds.

All three of these require relatively larger amounts of energy in comparison to the simple covalent substance.

So therefore they are going to have higher melting and boiling points so that energy is required to break those bonds whereas in the simple structures, we're overcoming weak forces of attraction.

So here we're breaking bonds.

For simple substances, we have to overcome weak forces of attraction.

So which of these statements are true?

I'd like you to select all the structures below that have relatively high melting and boiling points.

Well, the answer is the bottom three.

Simple molecular substances generally have very low melting and boiling points.

Okay, so that was melting and boiling points.

Let's talk about electricity and conducting electricity now.

So to be able to conduct electricity, a structure needs what's known as a charge carrier.

A charge carrier is something like an electron or an ion that is free-moving so they move through the structure and allow charge to move through the structure.

Simple covalent molecules do not have any ions 'cause they're made up of atoms sharing electrons.

So here we've got three different types of structure.

We've got one represented as a ball-and-stick, we've got one represented as a space-filling molecule, and then on the right-hand side we've got what's known as a displayed formula for a slightly more complicated but still simple molecule because it has a defined number of atoms.

None of these can conduct electricity.

They do not have delocalized electrons.

On the other hand, giant ionic structures can conduct electricity but they have to be molten.

So that's a liquid, a hot liquid, generally because they've got high melting points, or they can be dissolved in a solvent and that's because the ions are free to move in those structures.

When giant ionic substances are solid, the ions are fixed in place.

And so the ions are fixed in place, and so that means they cannot freely move.

So they cannot conduct electricity, they cannot carry a charge or current.

Giant metallic structures on the other hand, can conduct electricity when they're solid and when they're liquid and that's because the delocalized electrons are free to move within the structure and carry a charge or current at all times.

Some giant covalent structures can conduct electricity, but most can't because they don't have any charge carriers or free-moving charge carriers, I should say.

Diamond is an excellent example of this.

All of its electrons are used in covalent bonding.

So there are no free-moving delocalized electrons.

It is not an ionic substance, so there are no free moving ions so it cannot conduct electricity, whereas graphite, as we know, does have delocalized electrons.

So that means that this structure is able to conduct electricity and therefore carry a charge or current due to the free-moving delocalized electrons.

True or false then, only metallic substances can conduct electricity.

Well, that's false.

We've seen a couple of other examples.

Can you justify your answer here?

Pause now if you need a bit of time and press play when you're ready to continue.

And the answer is B, any substance that has free-moving charge carriers is able to conduct electricity.

Let's look at another type of property then.

So diamond is one of the hardest known materials and the reason why is because all of the atoms are held together by strong covalent bonds in a rigid 3D lattice.

On the other hand, another type of pure substance, a metal.

Let's use sodium as an example or gold.

These are malleable and ductile and that's because the layers of the metal atoms, while they are held together by strong bonds, aren't fixed in a localized system.

The electrons are in a localized system for diamond.

They're held together specifically between two atoms where that covalent bond exists.

Whereas in metals, because the force is uniform in all directions, those ions are kind of able they ought to slide over the top of each other.

It still keeps the regular structure but the layers can move over each other so that makes them malleable and ductile.

Another type of property that we can talk about is solubility in the context of ionic substances.

Ionic compounds are generally soluble in water due to the nature of water molecules.

They're able to surround the ions and cause them to dissociate, kind of split them apart.

If you think about something like sodium chloride, table salt, sodium and chlorine held together by strong ionic bonds, very high melting and boiling points to separate those ions away from each other to form a liquid.

However, if you put it in water, the water can very easily move them apart from each other, dissociate them and form a solution.

On the other hand, covalent structures are generally insoluble on the other hand, in water, but they tend to be soluble in other solvents.

Namely organic solvents such as ethanol, petrol, tetrachloromethane, and various other different substances.

You don't need to know about these, but it's just a really interesting point, I think.

There are some substances, however that do dissolve in water such as sucrose or sugar.

And that's due to a special type of force of attraction between the different molecules of sugar and water due to the oxygen-hydrogen bonds.

We call this technically a hydrogen bond.

You don't need to know about it at GCSE, but for those of you that are interested and potentially go on to further chemistry, you'll learn more about hydrogen bonds in A-level chemistry.

Metallic substances are generally insoluble in most things and that's because they are not able to have their bonds disrupted by other substances.

Those are very strong bonds and metals are able to dissolve in other metals.

We've talked about this already.

When metals form metallic bonds with other metals, they don't form compounds, they actually form mixtures and we call this special type of mixture an alloy.

Most alloys end up with a mixture of properties from their constituent metals and other atoms that can be kind of added to the alloys.

There's lots of useful properties to a lot of alloys.

One example is mercury amalgams that are found in dental practices to make fillings.

Pure mercury is incredibly toxic, but alloys of mercury are not.

So there are four statements here.

I'd like you to have a read through them, pause now if needed.

And I'd like you to select all the correct statements.

Press play when you are ready to hear the answers.

So pure metals are made of layers that slide easily over each other.

When hammered, we call that malleable.

Diamond is hard due to the strong covalent bonds holding atoms in a rigid lattice structure.

Ionic molecules are, well, that's just a statement that I don't like.

Ionic substances are giant structures.

They don't form molecules.

And D, mixing metals together forms alloys and these are held together by metallic bonds.

Okay, we've got to the final set of tasks now.

So this one here gives us a table of data.

So we've got four different types of substance and we've got their melting points, and whether they conduct electricity in solid form or when they're liquid.

We've been asked to suggest what type of structure and bonding each of these different substances have using this data to explain our answer.

This generally comes up as like a six mark answer question on an exam.

It's normally a really, really good test of whether you've understood structure bonding and properties.

So pause now once you have a think back through what we've just talked about.

Have a go at answering this question, and then we'll go through the answer together.

When you're ready to, so press play when you're ready to hear that.

So substance A has a high melting point, doesn't conduct electricity at all.

That means it's probably going to be a giant covalent substance because giant substances generally have high melting and boiling points.

And if it conducted electricity in any form, then we would think it was probably a metallic or an ionic substance.

But because it can't conduct electricity at all that that means it's going to be a covalent substance because of the high melting point, a giant covalent substance.

Substance B also has a high melting point and it can conduct electricity when both solid and liquid.

Well, again, the high melting point tells us that it's likely a giant structure.

And because it can conduct electricity when it's a solid, that tells us that it's likely a metal because ionic substances can't conduct when they're in solid form.

It could be graphite.

We could be talking about a giant covalent substance here, but generally graphite and diamond for instance, they have melting points of nearly 3,600 degrees Celsius, much, much higher.

So this is unlikely to be a giant covalent substance, more likely to be giant metallic.

C also has a high melting point.

When we say high melting point, we're talking higher than kind of water, for instance, or anywhere between zero and 100 degrees Celsius, would still be considered quite a low melting point.

Anything higher than that would be considered a high melting point into the hundreds, if not thousands.

It can't conduct when it's a solid, but it can when it's a liquid.

So that tells me we've got the kind of distinctive properties here of an ionic substance, specifically a giant ionic substance.

We want to talk about structure as well as property.

And then D has a very low melting point minus 10 degrees Celsius and it can't conduct electricity at all.

That tells me because it's got a low melting point, that is likely to be a simple molecular substance and the electrical conductivity confirms that.

Okay, question number two.

The table below shows the ability of different substances to conduct electricity.

Explain the results by referring to the structure and bonding of these substances.

Again, could be another six-mark answer, or it could be a four-mark answer-style question in an assessment.

This time we've been given the names of the substances, so we need to then infer from that also what type of bonding they'll have and use the electrical conductivity to help us answer this question.

Pause the video now to have some time to answer this question.

Very similar to the last one we've just gone through.

So hopefully you've got a bit more under your belt to be able to help you with this question.

And then when you're ready to hear the answer, press play.

So copper has a giant lattice that's held together by metallic bonds and it has delocalized electrons, which are free to move through the structure to help it carry a charge or current.

Carbon dioxide is simple molecular.

The outer-shell electrons in carbon and oxygen are used in covalent bonding, so there are no delocalized electrons or free-moving ions, so it cannot conduct electricity.

Sodium chloride has a giant lattice held together by ionic bonds and in the solid state the ions are fixed in place by strong electrostatic forces.

But when it's molten, so when it's liquid, the ions are free to move and can carry a charge/current.

Well done if you had most of that in your answer.

Be careful here.

The question is asking us to explain.

So we must go through the details about how we know these various different types of bonding and structure are able to conduct electricity.

I really enjoyed going through this lesson with you.

It's one of the most important parts of GCSE chemistry, so I'm hoping that you feel a lot more confident about it.

Now what we've covered in this lesson is that there are three different types of chemical bonding, covalent, ionic, and metallic, but there are only two types of structure, simple and giant.

The type of bonding and structure are what lead to the different properties of substances.

Melting and boiling points are linked to intermolecular forces or bond strength, depending on what type of structure, and conductivity is linked to whether there are free-moving charge carriers, either delocalized electrons or free-moving ions.

Thank you very much for learning with me today.

I've thoroughly enjoyed going through this with you, and I look forward to seeing you in the next lesson.

Thank you.