Lesson video

- Hi folks, my name's Mrs. Fark.

For today's session, you're going to need a pen and a piece of paper and it's a good idea to switch your mobile devices to silent or turn them off, to help to avoid any distractions as we go through today's lesson.

If you'd like to take a moment to do that now and then press play when you're ready to resume and we'll get started with our learning for today.

So today's session is about why elements react.

We're going to be looking at why metals are reactive and why non-metals react in terms of the electrons and their outer shells.

We're also gonna take a quick look at the noble gases and see if we can explain why they are so unreactive.

So today's session you're gonna spot a few of the following keywords.

So keep an eye out for these.

We've got electron shell.

This is the orbit or pathway that an electron takes as it moves around the nucleus of an atom.

We've got ion, which is an atom that has either lost or gained electrons.

So it has become charged.

And then we have inert.

Inert is the word we use to describe substances that are unreactive.

So today's lesson on why elements react has been broken down into three parts.

We've got why metals react, why the non-metals react, and then what's going on with group zero.

So if you're ready to jump in, we'll start and take a look at the metals and why they are reactive.

Metals become increasingly reactive towards the left and bottom of the periodic table, as we can see here.

So they're increasing reactivity as we go down group one from lithium towards francium.

And as we go across the table from sort of the aluminium side on the furthest right and tin all the way across to our group one metals.

So the most reactive metal that we're able to identify in our periodic table would be francium.

Well why? Why are all of these metals reactive? Well, this is because their outer shell is not full.

Atoms want to be stable and to be stable they need to have a full outer shell of electrons.

The majority of our metal elements tend to have either one, two or three electrons in their outer shell depending on their group.

If we look at the electron configuration, for example, of our group one metals, which she got here for lithium, sodium, potassium in the image you can see that they are all in group one.

They all have one electronic in their outermost shell as indicated in blue there on the picture.

So we have lithium has a full inner shell with two electrons and then one electron in its outer shell.

Ideally to be full, it would be looking to have eight in that shell, or it could lose that one and have just the two electrons filling that very first shell.

Sodium has got two, then eight and then one.

So again, it either needs to find seven more to fill that outer shell, or lose that single blue electron in its outer shell we can see there.

And potassium has then got 2, 8, 8, and then 1 electron again in that outer most shell.

So how do metals react? What's going on? How do they become stable and fill that shell? Well, metals lose electrons during chemical reactions to achieve a full outer shell.

In doing so, our sodium atom for example, is going to lose the outer shell electrons.

You can see in this image there, we've highlighted it in a pinky colour.

So it's gonna lose the outer shell electron.

And what it then does is it forms what's called a sodium ion.

So you can see there that whole outer shell has been removed and we've lost that single electron.

This leaves just two shells for sodium, the inner one with two electrons, which is full, and the second shell, which is full with the eight electrons that are required.

Now, in doing so, we've lost negative electrons and so we've formed a positive ion.

But you have to remember, electrons have a negative charge, an atom has overall neutral charge.

If we take negativity away from it, it's going to be left with a larger number of positive protons in the nucleus, than it has negative electrons.

Excuse me.

So that's leaving a disbalance of charges across the whole atom.

So overall, our sodium ion whilst more stable, actually now doesn't have a neutral charge.

It has gained charge by losing an electron.

And we need to be able to show this because if you look at that sodium ion now, but you didn't know it was a sodium, if you hadn't seen the previous image showing that it had the outer electron, then you might mistakenly look at that and think, well, it's got 10 electrons, it could be neon.

So how do we show these ions? Well, in order to indicate that, we can see that it would still have 11 protons in its nucleus, which is what makes it sodium.

So it hasn't, the nucleus hasn't changed.

That's really important to point out.

But we have changed the number of electrons.

In doing so we now need to show it in a slightly different way.

But it's really important that when you are drawing ions, you include these two key features, and that's the square brackets that you can see around the sodium ion and the charge which it has on the outside.

So sodium, because it's lost one negative electron has become positive by one so we just write plus.

Scientists, kind of on the lazy side, if we don't have to write a number because the plus indicates that it is just plus one.

If you were looking at a metal in group two and it had lost two electrons though, you would then put plus two, okay, because it would've become positive by two.

So the positive metal ion once formed will interact with a negative ion that receives the electrons that's lost.

These electrons don't just go flying off into nowhere.

They do need to go somewhere and we're gonna look at that later.

And the interaction between the metal ion, the positive metal ion and the negative ion is what we would then call a compound.

But we don't wanna go too far into that just yet, because you'll be looking at this when you study something called bonding in future sessions.

All metals are reactive.

We're happy with that idea.

But why are some metals more reactive than others? Well, let's take an example from group one to see if we can figure it out.

We've got francium on the left and lithium on the right, and we can see that francium has a very, very large atom.

It has a huge number of electrons and it's outer shell electron is very far from the nucleus.

What effect does this have? Well, it means that francium is very much more reactive than lithium.

The outer shell electron is further from the nucleus for francium, whereas for lithium, it's much, much closer.

And remember, it's our nucleus, which shows the positive charge.

That's where our protons are.

The electrons are negative and because they're opposites, there's a force of attraction between them.

But the further that outer electron gets from the positive nucleus, the less of an effect that it's going to feel and the less attractive force there will be.

So we have a weak attractive force between the nucleus and the outer electron, partly due to the greater distance, but also to do with something called shielding.

So hopefully you can see really clearly on this image that with francium there are a large number of electrons and electron shells between that central dark blue nucleus and the very far outer electron all on its own in the outer shell.

Lithium, however you can see that only has one in a shell with two electrons on it.

And so there are far fewer electrons between the nucleus and the outer shell.

Now what this generates is something called shielding.

And so we see a huge amount of shielding of that outer shell electron in francium, whereas we see much less of it in lithium.

And so this is why francium's outer shell electron is much more readily lost.

It's able to kind of just zip off almost unnoticed.

And whereas lithium's is much harder to remove, because lithium has a much tighter hold, it has a much greater attractive force to that outer shell electron than francium does.

And so this is why francium is far more reactive than lithium.

You see much greater shielding and a the greater distance from the nucleus means that that outer shell electron is much, much more easily lost in that larger francium atom, than it would be in the smaller lithium atom.

So let's do a quick knowledge check and see what you are able to recall.

We've got true or false, metal atoms generally gain electrons to form positive ions.

Is that true or false? This is quite a tricky question.

Metal atoms generally gain electrons to form positive ions.

Well, metal atoms don't gain electrons, they lose electrons, but they do form positive ions.

So it's partly correct, but the fact that they've got gain electrons to form positive ions, that's gonna make this sentence false, okay? In order to become positive, you need to lose negative charge.

It always seems to be a bit of a brain bender, that one, it's quite difficult to get your head around the idea that you're losing something and yet becoming positive.

But that's exactly what happens in the case of our metal ions.

Well done if you got that right, that was tough.

Question number two, how many electrons would you expect to find in a full outer shell? And we've got some options.

We've got option A three, B four, C six, or D eight.

Okay, so the correct answer from those choices would in fact be eight.

I know there'll be a few die hard covers out there, but if you've only got one inner shell, then it takes two.

It does, it does.

And that would class as being full.

But in this case, with these options, the correct answer that we are looking for is eight.

Well done.

Okay, which is the most reactive metal from this list? We've got A potassium, B, aluminium, C, francium, or D calcium.

So the correct answer is C francium.

So now it's time for you guys to have a little go at practise task.

The first part of this practise task is for you to complete the paragraph by filling in the gaps.

You might wanna give this task your full attention and spend a bit of time doing it, so I suggest you pause the video and when you're ready to go through the answers, you can press play.

Good luck.

Okay, let's take a look at how you got on.

So we've got complete paragraph by filling gaps.

So let's take a look.

Metal atoms lose electrons to form positively charged ions.

Metal atoms need to lose electrons in order to gain a full outer shell.

This makes the atoms more stable.

The metal atoms that lose electrons most easily are going to be the most reactive and the most reactive metal is francium.

Well done if you've got all of those right.

Second part of this task is an exam style question.

We are asking you to explain with reference to the electronic structure why francium is more reactive than lithium and this is worth three marks.

So it's not good enough to say francium is more reactive.

You need to think about explaining why.

If you can include the word because in your answer.

Why is it that francium is more reactive than lithium? I'd like you to pause video at this point and really think about your answer.

Once you think you've got an answer that's going to give you three marks, then press play and we'll take a look at it together.

Right, let's see how you got on.

Now, in all cases where we have a long answer question, your answers don't need to be exactly the same as mine in terms of word for word responses.

But you need to have some of those general themes.

So I started my answer by looking at francium.

So francium is more reactive than lithium because it's outer shell electron is very easily lost, okay? But why is it very easily lost? We've been asked to explain so you need to go into a bit more detail.

Francium's electron is further from the nucleus and the attractive force of the nucleus is reduced by the inner electron shells creating shielding.

So we've explained why francium is more reactive.

We said that its outer shell electrons more easily lost.

And we've described francium's structure in order to support that.

In order to fully answer the question, it does say compared with lithium so I'm gonna need to say a little bit about lithium.

Lithium's outer shell electron is much closer to the nucleus and so the attractive force of the nucleus is strong and there's little shielding by other electrons, okay? And we know that means that it's less easily lost and therefore it is a less reactive metal.

Well done if you think you have managed to hit all of those marking points.

That's a really tricky question.

If you feel you need to go back and take a look at this section video again, then obviously you can re-watch it at any time in order to support you with the learning in this topic area.

So we've taken a look at metals and why the metals react.

Now it's time to see if we can explain what's going on with our non-metals.

Where are our non-metals, what are they and where do we find them? Well, non-metals are found on the periodic table towards the right hand side and their reactivity increases towards the top and right of the table.

So we can see here that astatine would be very unreactive and fluorine one of the most.

Now the reason I've highlighted the very far column in purple is because this trend does not include what's called group zero, the noble gases, okay? So when looking at non-metals, yes, the noble gases are non-metals, but they are not included in reactivity.

We're actually gonna take a look at those in a minute.

So just kind of put a pin in those for a minute.

We will come back to it.

So in terms of our least reactive, we've got things like antimony, germanium, these are gonna be our less reactive non-metals and fluorine being one of the most reactive.

So let's take a look at our non-metal structure.

Well, non-metal atoms, just like metals also require a full outer shell to become stable.

But in this case, what we tend to see is that non-metals have four or more electrons in the outer shell.

If you look here at the image, we're showing the group seven elements.

We've got fluorine at the top, two electrons that are in its inner shell and seven in its outer shell.

So in order to fill an out in its outer shell, it only needs to gain one.

This makes a lot more sense, it's gonna be a lot easier then trying to lose or give away seven.

The same can be said for chlorine with an electronic configuration of two and then eight and then seven and bromine, which is 2, 8, 8, 7.

So we can see that each of the outer shells of our group seven elements is missing an electron.

You notice they've got seven electrons highlighted there in blue.

Needs an eighth, okay to fill that outer shell and become stable.

So our full outer shell requires eight electrons, or if it's only one, if it only has one shell, it would be two in the innermost.

So where are these? What's gonna happen to these atoms when they gain some ions? So non-metal atoms gain electrons to fill the outer most shell.

So the metals were losing electrons, our non-metals are gaining electrons.

So if you take a look at our chlorine atom, which is 2, 8, 7, it's gonna gain an electron.

You can see here it's gained one and by the addition of that sort of pinky purple electron in its outer most shell there to become 2, 8, 8.

That's important to remember again, those electrons still have charge, okay? Remember our electrons are negatively charged and so we are adding negative charge.

Again, it takes a minute to get your head around it.

We're adding negative charge.

And so overall what's gonna happen to our atom is it's going to become negatively charged, okay? So our atom is gonna become negatively charged because gaining electrons results in the formation of a negative ion.

And again, we need to be very careful when we're demonstrating this to our teachers or examiners, so that we can clearly explain that we know that it's now an ion as opposed to an atom.

And we would show that essentially by using the same sort of pattern.

You can see here, you've taken a chlorine atom and added it the outer shell electron in purple.

So it's now formed an ion.

Again, this could be mistaken for one of the noble gases if you didn't have that proton number there.

And essentially to prove to everyone we know it's an ion, we draw our square brackets around the outside of the diagram and very clearly indicate that it is minus one.

It's only gained one negative electron.

So it has only gained a one negative charge, okay? So the negative metal ion is now going to want to interact with a positive metal ion and that it received that electron from, okay? So the two actually kind of work in a synchrony with one another.

The metals want to lose, the non-metals want to gain and we see an interaction between the two.

Again starting to pivot towards some bonding, which we would be looking at in a later session.

So here we have two of the group seven elements.

We have fluorine, which is the most reactive, and we have astatine which is much less reactive.

So if we take a look at the properties, then we've got fluorine which has few electron shells and so very little shielding.

And the positive nucleus therefore has a very strong attractive force to outer shell electrons.

And we can compare that with astatine, which you can see has many electron shells and leads to lots of shielding.

And so the attractive force from the nucleus on the outer shell electron is weaker.

So guess you could kind of think of fluorine's look being a bit of a sort of snatchy atom.

When a negative electron comes next to it that positive nucleus having a very short distance to the outer shell is able to grab hold of and keep hold of the outer shell electron in order to fill that shell much more readily.

This making it much more reactive.

So we see a shorter distance and less shielding makes the non-metals more reactive.

Astatine you can see, has a very large atomic radius.

Its very large distance between the nucleus and that outer shell and a huge number of electrons on the shells inside.

And so a large amount of shielding.

So the attractive force to try and pull another electron into its outer shell is much less.

So what you may have spotted is that we're seeing reverse here of the pattern that we see in the metals.

So in this case, the smaller ions with the less distance between the outer shell is actually more reactive and the larger atoms with a greater distance and greater shielding is going to be less reactive in which we saw is the opposite in our metals.

So non-metals are also able to fill their outer shells by sharing electrons.

So it's only our non-metals that are able to do this.

I'm not going to go into huge amount of detail, but ultimately it reminds me a little bit of when I used to do equipment checks in school.

Thirty children sat in front of me and there definitely wouldn't be 30 rulers out there.

And yet somehow by some sort of passing under the tables, whenever you asked to see each child's ruler, they all had one.

And this is kind of what's going on here in terms of chlorine and it's outer shell electrons.

So when someone's looking, you have the chlorine atom on the left would be able to count on and say it's got 1, 2, 3, 4, 5, 6, 7, 8 electrons.

And then if someone's looking at the atom on the right hand side, you'd be able to count round and say that that has eight electrons.

Whereas in all honesty, what they're doing is that pair of electrons in the centre there is being shared between the two of them and they're just kind of passing it to and fro as and when it needs to appear that they have them.

So this sharing only seen between our non-metals when they react with other non-metals in order to fill that outer shell.

So let's have a quick knowledge check.

We've got a true or false question.

Non-metals can only fill their outer shell by gaining electrons from metals.

Is that true or false? Okay, so the answer to this question based on the last slide is gonna be false, okay? Non-metals are able to either gain electrons from metals to form ions, or they can share electrons to fill that outer shell.

Well done if you've got that one right.

Which of the following is the most reactive non-metal? We've got choices from A, chlorine B, bromine C, fluorine, and D, iodine.

Correct answer to this one is C, fluorine.

So now it's time for a practise task.

So what we'd like you to do is make sure that you're paying these tasks your full attention, giving your very best effort.

So when you're ready to go through the answers, you can restart the video.

But for now I would like you to press pause and give it a go.

Right, let's take a look at how you got on.

So the first part of this question, this task asks us to describe two ways in which non-metals can fill their outer shell.

So again, these are kind of slightly longer responses and so doesn't need to be exactly word for word, but it should follow something like this.

Non-metal atoms can fill their outer electron shell by gaining electrons from metals or by sharing electrons with each other.

Question B on this task asks us to look at non-metals, non-metals form negative ions.

Explain why.

So they've given us a statement and they've asked us to explain.

So let's take a look at the answer.

Non-metal atoms gain negative electrons.

This means they now have more negative electrons than positive ions and so have an overall negative charge.

Okay, well done if you've got that right.

So you need to look for those key marking points, okay? You need to be talking about the idea that they're gaining electrons.

You then need to allude to the idea that electrons are negative and because they now have more of those negative charges than they do positive protons, they now have an overall negative charge.

Well done.

Okay, these are quite tricky.

Right onto the last question in this section.

Again, you would want to pause the video and give yourself an opportunity to have a go at this one.

The question is suggest why the reactivity of metals and non-metals show opposite directions.

Right, okay, so thinking about the explanation for fluorine and astatine and for lithium and francium.

So take your time.

Okay, you've got as long as you need and when you're ready to hear a model answer then press play.

Right, let's take a look how you got on.

This one's quite long.

It's almost feels like you might want to kind of plan out what you want to say in terms of, a bit like you would in English really.

So you kind of, wanna talk about the reactivity of metals, you wanna talk about the reactivity of non-metals and why they're more or less reactive.

So let's take a look at a model answer.

So we've got non-metals become less reactive down the group because the atom's ability to gain electrons weakens due to more shielding and greater distance from the nucleus.

These have become less reactive down the group.

Now we need to make a statement about our metals.

Metals however, become more reactive down the group, because the atom's ability to lose electrons is greater due to increased shielding and a greater distance from the nucleus.

This is a really tough question guys, so well done if you think you've been able to give that a really good go, all right? Okay.

So where are we at now? Well, we've taken a look at metals and why they react and we've looked at the non-metals and why they react.

It's now time to take a look at group zero.

Okay, and what's going on there? I said we'll come back to them later, later is now let's jump in and take a look.

Okay, so group zero also known as the noble gases are on the far right hand side of the periodic table as highlighted in pink just here.

They're the last our elements to be discovered, because they are inert.

So what inert means is that they are incredibly unreactive, okay? They don't require to combine with other elements.

And so they weren't.

So it's not like you would make one during an experiment and know about it.

So this is why they were the last ones found and why they their own little group on the end of the periodic table there in group zero.

Now our noble gases are so unreactive, they're actually described as being monatomic.

And what we mean by that is that they exist as single atoms. So they are so unreactive they don't even wanna hang around with each other.

And I very often think that's kind of where their namesake comes from.

They're kinda noble, very noble and very regal.

And I'm far too posh to hang around with other elements.

That was my best go at posh was, I'm sorry, I shouldn't, I won't do any more impressions.

But essentially it means that there's no requirement for them to try and connect with any other elements from within the periodic table.

But helium exists as helium atoms on its own.

It doesn't combine with other helium atoms, it doesn't combine with any of the other gas atoms. And it certainly doesn't combine with any other of the elements in the periodic table, so they are monatomic.

Well, why are they monatomic? Why are they inert? Well, if we take a look at their electron structure, hopefully by looking at these diagrams, you'll be screaming at me now.

Well, I can see, I can see why, I can see why.

It's because they have full outer shells of electrons.

If you look at helium, it just has one shell.

It has two electrons, it's full.

We look at neon, it's got two and then eight.

That's a full outer shell.

Argon has got 2, 8, 8, that's a full outer shell.

So the fact that their outer shells of electrons are full means that they are already stable.

There is no requirement for them to try and gain, or lose, or share electrons with any other elements and that's why they are unreactive or inert.

So let's do a quick little progress check then to see what you can recall about our group zero elements.

We've got true or false.

All the noble gases have eight electrons in their outer shell.

What do you think? True or false.

Okay, having just looked at their electron configurations, hopefully you've answered this one as being false, okay? They do all have full outer shells, but if you remember helium, okay, helium actually only has one shell and therefore its outer shell is full with just two.

So all the noble gases have full outer shells, but this is not always eight.

Okay.

So let's take a look at whether you're able to identify whether we have an atom, an ion, or a charge.

This is something that's quite commonly asked about and it's important to be able to demonstrate your understanding.

So we already know that the type of element is determined by the atomic number, which is the number of protons in the nucleus.

Now, through everything we've looked at today, why these elements react, we've not changed that atomical proton number.

So the elements themselves have stayed the same.

However, we have looked at the idea of forming ions.

Now an ion is an atom that has either lost or gained electrons to become charged.

And we said if you have gained electrons, then you have become positively, if you, sorry, if you have lost electrons, you have become positively charged because you now have more positive protons in your nucleus than you do negative electrons to balance those out.

And if your atom has gained electrons, then you now have more negative charges than you do positive protons in the nucleus.

And so it has become a negatively charged ion.

So it's about comparing the numbers of protons and electrons to try and identify if something is an atom or an ion.

So remember, atoms have no overall charge.

So before the loss or gain of electrons, they are neutral charge overall, they have no charge.

When we start moving electrons around, we start to form these ions.

If you have more electrons, you gain electrons, you're gonna be negatively charged.

And if you lose electrons, you're gonna be positively charged.

So what I would like you to do is have a little go at a practise task.

I'm gonna ask you to complete the table by adding the element name you're going to, which you'll look up on the periodic table using the proton number.

You're then going to look at the number of protons and electrons and decide whether it's an atom or an ion.

So is it got the same number of protons and electrons, or does it have a different number of protons and electrons? And then you're going to try and work out the charge.

So if it's an atom, it's gonna be neutral.

If it has more electrons than protons, it's going to be negative.

If it has less electrons than protons, it's going to be positive.

Okay hopefully I've explained that task clearly.

Give it a go, okay? Pause the video, take your time.

You will need a copy of the periodic table for this one.

And when you feel that you're ready to hear the answers, press play and we'll take a look at it together.

Right, okay, let's see how you got on.

So the first one we've got is a proton number of 12.

So we take 12, we're gonna look at our periodic table and we can see that an element with 12 protons is in fact magnesium.

So in this case we've got 12 positive protons and 12 negative electrons.

The number is the same so we would be looking at an atom, and because we have an atom, we know that it must have a neutral charge, okay? Some of you start to go, "Ooh, hang on, I didn't really get this, but I think I do now.

We'll do one more and then if you need to pause the video again, you can.

So we've now got a proton number of 19.

So we're gonna take that proton number, we're gonna look it up on our periodic table.

And we can see that proton number of 19 is potassium.

Now we've got 19 protons and only 18 electrons.

So there's a disbalance in the two numbers.

That means we have an ion, okay? And which way round is it? So we have a higher number of protons than electrons.

So that means that we have more positive charge and so we have a positive ion.

All right now, if you'll just said I get it now, do pause the video, okay? And have a go at the last four and then press play again in a moment, okay? For those of you who think that you've got it straight off the bat, then we will carry on.

So proton number 35 on the periodic table that is bromine.

It's got 35 protons and 36 electrons.

There's a disbalance there so it's going to be an ion.

And because it has more negative electrons, it's going to have a negative charge.

Next one down, proton number three, it's gonna be lithium.

It's got three protons, but only two electrons.

So that will be an ion and it has got more protons, so it is going to be a positively charged ion.

Well done if you're getting all of these right by the way, this is not, it does take a little bit of thinking.

We've then got proton number of eights.

Let's look up eight on the periodic table that gives us oxygen.

We've also got eight electrons.

So those two numbers are in balance.

So that's going to be an atom of oxygen.

And because it's an atom of oxygen, we can say that it is neutral charge.

And then lastly, we have a proton number of nine, which if we look up on the periodic table is fluorine.

We can see that it has nine protons, but 10 electrons, that is not the same and so it would be an ion, and because it has a greater number of negative electrons, it's going to have a negative charge, okay? Well done if you got those right now, it was a really cha challenging task, okay? But if you can do this, it shows that you are really starting to get to grips with the idea of the formation of ions and also using your periodic table.

So good job.

The second part of this task then is to have a go at a question.

The question asks, explain why the noble gases are monatomic.

Now this question would be worth three marks on a typical assessment.

So you need to try and give as much detail as you can.

I'll again suggest pausing the video at this point and pressing play when you're ready to have a look at a model answer.

Okay, let's take a look at how you got on.

So we've gotta explain why the noble gases are monatomic, or we need to say what monatomic is and then we need to say, well, why the noble gases are that.

So the atoms of the noble gases have full outer shells of electrons.

They do not need to gain, lose, or share electrons.

They are stable and unreactive.

This means they exist as single atoms and are monatomic, okay? So within that answer there, we have included the idea that monatomic means they exist as single atoms. We've gone on to explain that that is because they are stable and unreactive, and we've said they are stable and unreactive because they have full outer shell of electrons, okay? So those are our three marking points.

They don't need to be exactly in that order.

You may have started at the bottom and worked in the opposite direction, that's fine.

And it doesn't need to be word for word.

So long as you have got the general gist of each of those sentences, then you're gonna be hitting those key marking points, okay? A great effort, guys.

Well done.

I'm super impressed.

This has not been easy.

Now, that does bring us to the end of our session together on why elements react.

So let's take a look at our key takeaway points from today.

Well, we've got the idea that most elements on the periodic table are reactive.

Elements react in order to fill their outer shell of electrons, which makes them more stable.

The metals need to lose electrons to achieve a full outer shell, whereas our non-metals need to gain electrons or share them in order to fill their outer shells.

And we have the group zero elements, which are called the noble gases.

And the noble gases are inert, which means unreactive, because they already have a full outer shell of electrons.

Well, well done for all your hard work today.

It's not been the easiest of subjects, but I hope you've enjoyed it.

And don't forget, you can always take a look back at any of the sections of the lesson again, if you're feeling a little unsure or want to recover it.

Thank you so much for all your hard work and I hope you've enjoyed the session as much as I have.

Before you sign out, you might want to have a go at our exit quiz to see what you can remember.

We'd like to take this opportunity to thank you for using Oak National Academy, and we hope to see you again soon.

Bye for now.