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Practice English Speaking&Listening with: Investigating the Periodic Table with Experiments - with Peter Wothers

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[MUSIC PLAYING]

[APPLAUSE]

Well, thank you very much.

It's wonderful to be here again.

Now, yes, this is the International Year

of the Periodic Table.

Currently, there are 118 elements, different elements.

And the periodic table is trying to bring some sense

to-- some order to this vast number of elements

that we have.

The song, of course, was in pretty random order.

So there was no sense behind that.

And this is what the periodic table is all about.

And as you've heard, as I'm sure you all

knew before, this year, 2019, is the International

Year of the Periodic Table.

And that's because it is the 150th anniversary

of the publication-- not actually of the very first

periodic table, but the first one

by this chap, the Russian scientist Dimitri Mendeleev.

So he first published his version in 1869.

And this is shown here.

It's in Russian.

So it's rather difficult read.

It says, "Chemie" at the top there.

But this is his table.

It's slightly different from our modern version.

In this one-- well, in our modern version

we see that groups of similar elements

with similar properties are arranged in vertical columns,

as we shall see in a moment.

In this one, they go across.

So it's slightly different.

And he changed this later.

We're more familiar with this version.

The printed one, that I just showed you,

actually is one that's in the collection of St. Catherine's

College in Cambridge.

And this is a display we put on the grass, in our main court

there, when we had an exhibition.

And the exhibition is going to be coming to London.

It's coming to the Royal Society of Chemists,

in Burlington House, in August.

So you can go and see, not only Mendeleev's version,

but the very, very, very first one,

from seven years earlier, which is really

rather interesting because it wraps around a cylinder.

So it's quite different.

This one here, people don't quite realise that this

isn't some sort of photoshopped version.

This is actually set out on the grass.

And they're all different sizes and all placed

in exactly the right position.

So that when you stand in just the right place,

order comes to the elements there.

And it looks like that.

So this was really a photograph that was taken.

This is inside the exhibition.

I say, do go and see these things

when they come to the Royal Society of Chemistry,

to Burlington House.

There's all sorts of different forms

of periodic tables on display.

In fact the very-- the long chart in the tall cabinet

there is the very, very, very first periodic table,

which is really quite exciting and very different

from the ones.

But hopefully after today, you'll

understand what makes a periodic table.

And they can look quite different, indeed.

And this is one that was donated to the college, commissioned

for us for this year, for the International

Year of the Periodic Table.

It's absolutely fantastic.

It's made out of silver.

And it really is a periodic table.

But it looks very different from the versions

that we're going to be looking at today

and we're familiar with.

But nonetheless, it still works.

So go and check this out if you get a chance in August.

Well, we're going to be sticking with our more familiar form.

But we need to understand exactly what it

is that makes a periodic table and how it works.

And in order to do this, we need to understand

the very structure of the atom itself,

of what all atoms are made of.

Now as far as a chemist is concerned,

atoms are made of just three things.

We have positively charged protons, in the heart

of the atom, in the nucleus.

And sometimes-- well, almost always, apart from one atom,

we have neutrons, which have no charge at all.

And then whizzing around the outside,

we have negatively charged electrons.

And so long as it's a neutral atom,

we have the same number of positively charged protons

and negatively charged electrons.

Physicists, of course, love smashing things up even further

and breaking everything.

But we're not going to be concerned

of what we could smash those things up

into, into quarks and things, all sorts of nonsense

like this.

We've just got to stick with our lovely protons, neutrons,

and electrons.

So let's go then to the very, very first atom.

And this is hydrogen. So where is hydrogen?

Right at the top there, very good.

So you are atom number one, the first atom

in the periodic table, the first element.

And what makes you unique and different from every other one

is that in the heart of the nucleus,

you have just one proton.

OK.

That's what makes a hydrogen atom.

And in the neutral atom, it's balanced by the electron.

One electron whizzing around there.

So this is what we see.

When we go to the next atom, well, this

is the element helium.

So right over here--

very good, helium.

OK, so there's helium.

Well, helium now has two protons in the nucleus.

It happens to have-- most of the helium that we come across also

happens to have two neutrons.

We're not going to be interested in neutrons.

They increase the mass of the atom, of the nucleus there.

But they don't really alter its chemical properties.

But what makes helium special is the fact

that unlike hydrogen, with only one proton,

helium has two protons in the heart of the atom.

Balanced by, in the neutral atom,

two electrons whizzing around.

And then we come to atom number three, which

is all the way back over here again, which is lithium.

Oh, already there.

Look at that.

OK, very quick.

So element number three here, lithium, three

protons in the nucleus there.

OK, now we're going to move across.

So element number four is beryllium.

So four protons, four electrons.

Then we come back over to this side,

with boron, five protons, five electrons.

Carbon, six.

Then we have-- very good, carbon--

nitrogen, with seven; oxygen, eight.

So each time we're increasing by one proton.

Fluorine here, with nine.

And then neon, 10.

And then we come all the way back to this side.

And we come to the element sodium, number 11, very good.

OK, well, so far, so good.

Now, I'm not going to go through all 118.

That would be just a little bit tedious.

But this is the underlying order.

It is the order of their atomic number.

This is the number of these positively charged protons

in the heart of the atom.

But there's more to a periodic table

than just having them arranged in order.

We also need to group the elements

that have similar chemical properties

in these vertical groups.

This is really important.

And this is what any periodic table must achieve.

So we've got this underlying order

as we move from one element to the next,

where always each atom is getting heavier.

And we're adding more protons and electrons.

But we also need this order, with the similar properties.

And this is what we're going to look at now.

So let's have a look at group 1.

So these are the purple ones.

So if you're a purple or got a purple card, hold up your card.

OK, well done, very good.

So these are the group 1 elements.

Now, what they all have in common in terms of their atomic

structure-- and we will see some of their chemical properties

in a moment--

but what they all have in common is

that they all have just one electron

in their outermost shell.

Of course, hydrogen only has one electron anyway.

But then all of the others have one in their outermost shell.

So let's look at lithium again.

Lithium has, well, three electrons

in total, three protons.

But there's one in the outermost shell there.

When we come to sodium, more electrons.

But there is still one in the outermost shell.

It's relatively easy to lose that one electron.

And this gives it similar properties

to lithium, and in some ways even to hydrogen.

Then keep on going, potassium.

So the nucleus in the heart of the atom

there is getting much bigger and heavier.

The atoms are getting heavier.

But there's still one electron right in outside there,

whizzing around, which gives them all similar properties.

Let's go to rubidium.

Getting pretty massive now, loads of electrons.

But there's still the outermost one.

And we get to cesium, really crazy, craziness going on now.

But there is still that one electron whizzing around.

It's quite away from the nucleus.

It means it's quite easy to remove that one electron.

That is what gives it its chemical properties.

And we'll look at these in a moment.

But they will have something similar there.

There is, of course, francium, the heaviest element

in this group.

I'm not going to show francium because there

would be too many electrons.

But also because actually francium, I'm afraid,

falls apart quite easily.

This is where-- well, for the very heaviest of elements,

the nuclei are getting so large that basically bits fall off,

essentially.

This is actually radioactive decay.

What happens is, for instance, we

could lose two protons, two neutrons.

This is the heart of a helium atom, essentially.

So an alpha particle is a helios--

a helium nucleus.

This can ping out.

And then it's changed into a different element.

So this is why we're not going to be looking at the very, very

heaviest elements in the groups there

because these are unstable.

They are radioactive.

So cesium is one that we can look at.

It is stable.

But it has this one, all-important electron

whizzing around there.

So that's our group 1.

So here they all are again.

So you see group one again.

So all the purple ones, all have one electron

in the outermost shell.

Then we move to group 2.

So group 2, the orange ones, very good.

You're very busy, aren't you?

You got two to do, yes.

OK.

So group 2 here, all have two electrons

in the outermost shell.

And this gives them similar properties to each other,

as we shall see, but different from those from group 1.

OK, we're going to move around, all the way to here,

now to the brown cards.

So these are group 13 or sometimes called group 3.

And there are good reasons for this.

They all have three electrons in their outermost shell

that they can use.

OK, we come to group 14, the dark grey ones,

with four electrons, very good.

Group 15, the light blue, very good.

All beautifully aligned, all with five electrons

in the outermost shell.

16, the nice yellow ones, very good.

Six electrons in their outermost shell.

Sulphur is a little bit the wrong way up.

But it doesn't matter because it works either way.

And then we have group 17.

They are light, these ones.

These are the so-called halogens.

We'll come back to those in a moment.

And then right on the end, the final ones here.

We have group 18, the so-called noble gases.

OK, we'll come to those in a moment as well.

Let's go back for a moment to the halogens,

so the green ones.

So just show us again our halogens here.

These are one of my favourite groups.

These are some of the most reactive elements

in the whole periodic table.

So fluorine.

In fact, when this was first put into the periodic table,

into Mendeleev's periodic table, the element

hadn't even been prepared yet.

They knew it was there.

It was in certain salts, so particularly calcium fluoride.

And also the acid was known, hydrofluoric acid.

But nobody had yet isolated fluorine.

In fact, Humphry Davy, here, tried to do this,

nearly killed himself.

Other chemists did kill themselves,

trying to isolate fluorine.

It is so reactive.

Underneath fluorine, we have chlorine.

Chlorine is a toxic, poisonous gas.

And it gets its name from the greeny colour.

So "chloro" here, "chloros" means greeny-yellow.

It's the same as-- you might have heard of chlorophyll.

Well, that's not because it contains chlorine.

It's because it means, well, greeny--

greeney leaf, chlorophyll.

So it's the colour there that gets its name.

Bromine, aargh, bromine, one of my favourite elements.

It's really evil looking.

It's a liquid, a dark brown, orangey liquid.

And it gives off this vapour, really heavy, choking,

horrible smell.

In fact, it's name, "bromos" comes from the Greek

meaning "stinky" or "smelly."

So that's how it gets its name.

Iodine, underneath, gets its name from the beautiful violet

colour of this substance.

And this actually-- iodine and chlorine

were named by Humphrey Davy, here in the Royal Institution.

Astatine gets its name from being unstable there.

And tennessine, from the state.

So quite a nice, fun group.

I thought I might show you one of these elements

now, despite the fact that some of them are very, very toxic.

I'm going to show you some iodine.

And remember, this gets its colour--

name from this beautiful violet colour.

So we're going to make some iodine.

Now, at the back of the lectern--

this is why you had to come in through the sides--

I've put a very explosive compound.

This is called nitrogen triiodide.

And this compound has nitrogen bonded to iodine.

But these are very, very weak bonds.

And I'm hoping that when I just touch this compound--

not with my finger as you shall see--

but when I touch this, it should explode very violently.

And this is because these weak nitrogen iodine bonds all

rearrange during the explosion and form the very, very

strong bonds that we have in nitrogen molecules.

So nitrogen, like the nitrogen that we breathe in,

are molecules of nitrogen where two atoms are very strongly

bonded together to form very stable molecules.

And it is this that drives, I hope will drive this explosion.

It's the formation of the nitrogen gas.

But we can't see the nitrogen gas that's going to be formed.

It's invisible.

It's in the air around us cause.

We can't see this.

But we should be able to see the other byproduct

of this reaction, which is the iodine.

So you need to look very carefully.

And you should see a purple cloud if it all goes well.

Now this reaction is a little bit loud.

OK?

And it is quite a sharp explosion.

So I will be wearing these ear defenders.

And this is always a good sign when you see the lecturer ear

defenders, you really do need to just cover your ears.

Now, you won't need to stick your fingers in your ears.

Just put your hands over them.

And this will stop the shock wave.

OK, which could, otherwise if you're close,

particularly close-- damage your ears.

So just cover them.

You don't need to cover them really tightly.

Just put your hands over your ears.

OK, so let's give it a go.

So I should put these on.

OK, and just check that everyone's

got their ears covered.

So just cover your ears.

OK, and--

[POPPING]

There we are.

[APPLAUSE]

So I hope you saw the nice, violet cloud.

I seem to have changed the screen a little bit.

Never mind.

So iodine is, of course, something

that you'll be familiar with.

Maybe you've got some at home in a first aid box.

It's tincture of iodine.

So if you have a cut, you can use it as a disinfectant.

So when you have the solution, it has this sort of browny

colour.

This is the colour that we see on the screen.

That will actually disappear gradually

during the course of the lecture.

It'll probably be gone by the end of the lecture, which

is quite interesting.

OK, anyway, so that's one of the elements from group 17.

This group, altogether with these similar chemical

properties are known as the halogens.

And this word actually means the salt makers.

Because they very easily readily form salt. I mean,

we know, for instance, that our common table

salt, the proper name is?

Sodium chloride.

Exactly, sodium chloride.

And that is then made with some sodium from this side

with the chlorine from that side.

So these are all capable of forming salts.

The group as a whole are known as the salt makers.

OK, now that's the halogens, group 17.

Let's have a look at a couple of the noble gases.

So where are your noble gases?

So very good.

OK, so these are all the noble gases.

Now this is quite a remarkable group, OK?

This is the last group of elements to be discovered.

None of these, none of these are in Mendeleev's table.

OK.

Actually, the helium had been suggested.

And sort of somebody thought there

might be this strange element in the sun.

Mendeleev thought he was mad and so didn't include it

and had no way to include it in the periodic table, anyway.

This whole group was discovered after his table.

The first of these elements to be discovered was argon.

And it's actually the third most abundant gas in the atmosphere.

Remarkable, OK.

So the first most abundant is nitrogen, then oxygen,

and then, well, it's not carbon dioxide

which some people think.

It's actually argon.

There's 1% argon in the atmosphere.

And yet people knew about nitrogen. People

knew about oxygen for 100 years before they noticed the argon.

And the reason for this, actually,

is because it doesn't do anything.

In fact, argon doesn't combine with any other element

at all in the periodic table and neither does helium,

neither does neon, OK?

Some of the other ones just begin to.

So some of krypton, xenon, radon,

they have been made into compounds

with some of the most reactive elements such as our fluorine

and maybe oxygen. But otherwise, these are very inert.

And actually, this is what gave the name to argon.

Argon comes from the Greek word "argos,"

which is, of course, a well-known shop.

OK.

And argos, well, I don't know.

I presume they knew this when they named their shop.

It means lazy or inactive, OK, inert, OK, so not not

doing any work.

So it's just an easy shop to do your shopping with.

So this is because of this, the fact

that it didn't react with anything else.

It was a completely new element when it was discovered.

Nothing like this had been seen before.

OK, so now remember as we go down, all of these elements

have similar properties.

So all of our gases here have similar properties.

They are all gases, as far as we know.

We don't know about oganesson since there were only about

five atoms ever made anyway.

But the rest of them are certainly all gases.

And I have a balloon with, well, one of them here.

And this is our helium.

OK, now, remarkably, I say these are so inert, OK, these even

don't bond to each other.

And this is quite remarkable.

So with all of the other elements, so most of them

are actually nice, shiny metals, like this piece of an aircraft.

I'm sure they won't miss it.

It's just the turbine blade.

This is actually made of titanium.

But it's a solid.

And this is because the atoms of titanium

are all bonding to each other, OK,

to form a very, very high-melting-point solid here.

But helium, none of these elements, in fact,

do bond to each other.

And they exist as individual atoms.

And if ever you have a helium balloon,

this is pretty much the only time

you will ever come across individual atoms,

so individual atoms in this balloon here.

OK, but as we see, they are lighter than air.

Of course, helium is element number two.

It's only got two protons, two neutrons, and two electrons,

so lighter than air.

Well, I have some of the next element underneath it,

some neon down here.

So this is the element neon.

Now, there's a rather remarkable thing.

And that is that whenever we have equal volumes of gases--

and so these are clearly, pretty much

equal volumes, the same size--

so long as they're at the same pressure and temperature,

they actually end up containing equal numbers of particles.

In this case, what it means is there

is the same number of atoms of helium in here

as we have atoms of neon in here.

These are, of course, lighter than air.

What about our neon?

Well, it's probably the balloon itself that is sinking here.

But remember, as we go down the group,

we are increasing the number of protons, neutrons,

and electrons, so each atom is getting heavier.

Now, I can't have a balloon of oganesson

because there were only five atoms or six or so ever made

and they disappeared like that.

Radon, I can't have a balloon full of that

because it's very radioactive.

So that wouldn't be good.

The heaviest one that I can have is the xenon.

And this is a balloon of xenon.

OK.

So it's gone down a little bit over the course

of the lectures, but pretty much the same number of atoms,

maybe slightly less than we have in here and in here.

But what about the mass of these?

So it really is actually, quite heavy, OK, for a gas.

As I said, it's the same number of atoms.

So this really wouldn't be a very fun balloon

to take your parties.

But this shows very nicely, though,

that as we go down the group, what's happening,

they're all still gases.

But because we're increasing the number of protons, neutrons,

and electrons as we go from one period down to the next,

one road to the next, OK, the atoms themselves

are getting heavier, very important.

OK, well, thank you very much, noble gases.

You've been very good holding your signs up.

Thank you.

Now, we mustn't forget the other elements in our periodic table.

We have a lot in the middle, the dark blue cards,

these are the so-called transition metals.

Do you want to give us a wave, transition metals?

So we have all sorts of things like titanium

that we saw in the blade there.

We have iron, copper, nickel, all sorts

of nice, fun metals here, gold, very beautiful mercury there.

OK, so these are our transition metals.

And then we also have the light grey.

Now, you're right at the back there.

So, yeah, I'm afraid that's rather important.

Some of these are becoming more and more important

now, such as neodymium.

So we have our magnets contained in those.

Very important.

Very strong magnets.

But these elements here, they ought

to be in the main body of the periodic table.

So there's a little place marker there.

The asterisk and the little cross

shows where they should go, these so-called lanthanides

and actinides.

But it would make the periodic table very long.

And so this is why we've put you out the way at the back there,

OK?

But you ought to be in the main body of the periodic table.

OK.

So these ought to be in there.

But it would just make it very long, too long for the screen.

OK.

So that's a brief overview of the elements

in the periodic table.

Now, what we want to do is look at the chemical properties

of some of these.

I'm going to start off with, well, the first two elements.

And I think Chris has prepared a balloon of hydrogen for me

and a balloon of helium so we could compare these two.

OK.

Chris, you seem to have used the same colour balloons.

I think I deliberately told you to use different colours so I

know which one is which.

OK, well, one of these apparently

is filled with hydrogen, one of them is filled with helium.

I have no idea which one's which.

Ah, there is a difference between them.

So remember whichever one it is, the helium actually

has individual atoms in here.

The hydrogen, well, hydrogen forms molecules

like with every other element in the product table

apart from the noble gases.

They all stick to each other.

Hydrogen forms molecules where two hydrogen atoms bond

together.

So one of these is filled with hydrogen molecules,

one is filled with helium atoms.

But actually, the hydrogen is still lighter.

OK.

Hydrogen is the lightest element.

But I don't know which one is which.

So, oh, dear, never mind.

Anyway, but I do know a way to find out.

We can use a test, a chemical test.

OK.

I'm going to use my little candle here, my lighted

splint, to see if I can see which one is which.

So let's try this first balloon, then.

So it's either going to be hydrogen

or it's going to be helium.

You don't need to cover your ears, unless you really,

really, really, really, really don't like any bangs at all.

But otherwise, you'll be fine.

It should be fine, even if it is the real explosive one.

We don't know.

We'll see.

OK, well, let's give it a go.

Now, watch very carefully what happens with my flame.

OK, we'll see which one this is.

OK.

So, well, all that happened there is, well, it actually

made my flame go out.

It didn't burst into flames.

It must be the helium.

So the helium remember is incredibly inert, unreactive.

It does not react with the atmosphere.

There's no explosion.

All that happened was the balloon burst.

And it even put my flame out.

So it extinguishes flame because, well, it

doesn't do anything.

OK, it's completely inert.

So that means we hope that this one, unless Chris has played

a trick on me and they're both helium,

we hope that this one is hydrogen.

So let's try the hydrogen one, then.

So let's see if this is hydrogen.

[POPPING] Oh, that was nice.

[APPLAUSE]

So that was very nice, then.

They clearly have very different, very different

chemical properties.

Our hydrogen, our hydrogen molecules

were explosive with the air.

So they were combining with the oxygen.

The helium didn't do anything whatsoever.

So very different chemical properties, these two

first elements.

But, of course, when the hydrogen

does react, the flame that you saw there,

it's the hydrogen gas there reacting

with the oxygen from the air to form water,

to form hydrogen oxide, H2O.

OK.

And, well, hydrogen was first appreciated as an element

by this chap here and his colleagues.

This is Lavoisier, the French chemist.

And he first drew up a list of elements,

a modern list of elements.

Unlike the Greeks, who thought, for instance,

you may heard that they thought that things could be broken

down into the principles of gases

like air, fire, earth, and water.

Well, he drew up, and his colleagues,

the first list of elements in their modern forms.

And it's in the book here, his treaty published in 1789.

And this is the English version from 1790.

And here is the list.

It's in the original French here.

We're going to look at this in a little bit more detail.

But one of the elements here, which

is from the English version, is hydrogen. OK.

So he recognises.

In fact, he even gave it its name--

hydrogen. And the name here derives from the Greek.

Well, of course, I'm sure you know that hydroelectric means

electricity generated using water from a dam.

So "hydro" means water.

This means "water maker."

Because as you just saw, our hydrogen

easily reacts with the oxygen from the air to form water.

So that's its unique property, in a sense.

It can easily form water.

It's the only element that can form water

when it combines with oxygen. So Lavoisier named

this because of that reaction.

But it was a little bit wrong here because two

of the other so-called elements that are in his list

both got light and what he's called caloric, which is heat.

It's another word for that.

And this is because he thought that these were also

contained in different substances

and they could be released.

And so they must therefore be there.

Now Lavoisier made some very, very careful measurements.

He measured how much of one reagent

combines with another during a chemical reaction.

And this is one of the very, very important things.

He measured how much hydrogen combines with oxygen,

for instance.

OK, and this became very, very important later.

And we're going to try this in a moment.

And we'll also see why he thought

that light and heat should be in the list of elements.

I say, we would not include them now.

We know that they are forms of energy instead.

Well, I have here some iron wool.

OK, this is very finely divided iron.

And we're going to combine this with the oxygen from the air.

So it's a chemical reaction.

In a sense, we're going to be burning the iron wool here.

And we're going to be seeing how the mass changes.

So I've set my balance here to be zero.

We're going to see if it goes down, if it stays the same,

or if it goes up.

OK, so we have three possibilities.

So actually, maybe we'll have a vote.

OK, we'll have a vote to see who's right.

So who thinks that the mass will stay

the same during the reaction?

No one.

Who thinks the mass is going to go down when it burns?

That's quite a few hands there.

OK, quite a few hands.

Who thinks it's going to go up?

Oh, even more hands.

I think if we had to ask the audience,

we would say it's going to go up.

Well, let's try it.

We need to do the experiment.

And this is the sort of thing that Lavoisier did.

So I'm just going to start this, start the reaction.

And actually, maybe we could have the lights

down for this perhaps.

So I'm just going to start the reaction here.

And you can see the beautiful light that's being given out.

And this is why Lavoisier thought that light is contained

either in the iron or in the oxygen.

He thought it was in the oxygen. The mass has actually

gone down.

It's almost a gramme lighter.

So if you said it goes down, you were temporarily right.

Because it's going back up again.

If anyone said it stays the same,

well, it's almost back to where we started now.

In fact, now it's back where we started.

But actually, now, it's continuing to get heavier.

So if you did say it's going to get heavier,

you were quite right.

OK, now why is this?

Well, I say, we can certainly see the light.

And I can feel the heat very nicely here.

And it's probably this heat that's given out.

And it decreases the density of the air that's trapped in here.

That's probably why the mass went down initially

just as it became a little bit less dense, the air there.

But it is actually going up.

And why is this?

Well, none of the iron has disappeared.

It's still all there.

However, it's now combining or combined

with the oxygen from the air.

And so we haven't lost any iron.

What we've actually done is gained some oxygen.

And we formed iron oxide.

So the mass, of course, is heavier

because it still has the iron.

And now it's also got oxygen built into this structure

that we have here, this solid.

So it's over three grammes heavier now

due to the combining mass of the oxygen that's now

trapped into the solid here.

But the light and heat were given out.

That's why Lavoisier included them in his list of elements.

Also in his list, well, we see oxygen.

He named the word "oxygen" as well.

And azote, well, that was the name

that he gave to nitrogen gas.

It comes from the Greek meaning "no life,"

because it doesn't support life.

It is a suffocating gas.

We don't call it that any longer.

And that's because, actually, pretty much

every gas is poisonous or toxic or suffocating

other than oxygen. So it's not a unique name.

So this is why the name was eventually changed.

But rather than using Lavoisier's list

and looking at that further, I'm going

to use the list from another of my heroes from this time.

And this was the chemist, well, the author Jane Marcet, here.

She published a series of books that

were really very, very good.

And they got a lot of people excited and interested

in science and, particularly, in chemistry.

And so the book that I'm particularly interested in,

this Conversations on Chemistry.

And this really did have quite a huge impact.

It went through many, many editions.

And this is why we're going to be

looking at this because we can see how her list of elements

changes.

It went through many editions.

But it also got many people started on science.

And actually one of the most famous people

that they got interested in science thanks to her book

was a, well, a young bookbinder's apprentice.

OK, and this is Michael Faraday, of all people.

He probably should have been binding the book.

But at any rate, he was reading the book

and doing the little experiments in here.

And it got him excited, interested in science.

He then approached Humphry Davy, made notes of his lectures,

and got a job here in this building.

OK, and he became one of Britain's most famous

scientists.

But it all started thanks to this lady.

They became good friends throughout their lives.

They corresponded with each other.

She often came to this very room attending his lectures

to hear him talk about science.

And then she would update her books with all the new science

there.

So absolutely fantastic.

But let's have a look at her list

of elements from the very first edition from 1806.

And here it is.

So it's very similar to Lavoisier's list.

Also at the top we see light and caloric

and oxygen, then nitrogen, hydrogen.

So it starts off exactly the same.

This is because, well, this is her list of elements.

There are a few more elements in here

than Lavoisier had because it's a few years later.

It's 20 years later.

But there are some things that aren't elements, that are not

elements that are in both Lavoisier's list

and in this list here.

And these are, we see potash and soda.

Now I have some soda crystals here.

I bought these from Sainsbury's.

OK, it's just washing soda.

What this actually is is just sodium carbonate.

And it dissolves in water quite nicely.

And I could use this to wash my clothes if I wanted.

It's an alkaline solution.

If I add some of this is, it's a very dark green.

This shows that it's a neutral solution.

If I pour some in here, it goes to this beautiful blue-purple

colour here to show that it's the alkaline,

the opposite of an acid.

Now why is this in the list?

Well, it's in the list because Lavoisier could not

break down this substance.

He could heat it up to very high temperatures,

and it would lose some carbon dioxide,

but it couldn't be broken down any further.

They couldn't get the real elements out of this.

So as far as Lavoisier was concerned,

this was a simple substance that couldn't be broken down.

He suspected that maybe there would be some element in here,

but he couldn't get it.

Similarly for the potash that's mentioned here.

This is potassium carbonate.

This couldn't be broken down either.

OK.

And we also have some other ones here,

so there's lime magnesia, strontites, and barytes.

Well, again, these are substances

that could not be broken down.

The lime, what it actually is is what

you get if you try to break it down

by applying a strong heat, some calcium carbonate.

Now calcium carbonate is commonly known as chalk

or it's a marble.

OK, so calcium carbonate if you heat it up very strongly what

you end up with is lime.

And I've got some here.

Now, we're going to look at this.

It's quite an interesting substance.

So this was included in Marcet's list of simple substances.

It's, you know, quite hard.

It's not hot.

It's just been lying around here.

But it can be something really quite fun.

And I need a volunteer.

Oh, your hand went up very quickly in the grey.

Do you want to come down, please?

OK.

Right.

Now, OK, very good.

Thank you.

And you might to see if you can put those over your goggles, so

extra, extra goggles.

This is good, extra, extra ones.

Now we got a tiny little watering can there,

if you'd like to pick up that without spilling

too much water.

And what I'd like you to do is to pour water over this,

just flood this.

And then just step back a bit, just over to there, OK?

Thank you very much.

So if you just pour water all over it,

give it a good flooding all over, both pieces.

That's good.

That's very good.

OK, and just step back.

Thank you.

OK, so far-- now, it was cold and the water was cold.

And a violent reaction is taking place.

In fact, this reaction is so violent

that this is steam being given off.

And the reason I had to ask you just to step back

there is because actually little bits

can ping off of this sometimes.

It really just depends.

So now it's much flakier.

I'll add a bit more just to get it reacting there.

So the substance was lime.

Well, lime is sometimes also used for calcium

carbonate itself.

But this substance, really, is also known as quick lime.

And the quick here doesn't mean it's fast.

Because it isn't really very fast.

"Quick" is an old word meaning "living."

So, for instance, "quicksilver" means "living silver."

It's mercury.

We have the "quick" of our fingernails, the living

tissue there.

This phrase the "quick and the dead."

It doesn't mean the fast-moving and dead people.

It means the living and the dead.

So this is a living rock, this living lime

here, OK, very, very violently reacts with the water.

A round of applause there.

Shall I take your goggles there?

Yeah, definitely [INAUDIBLE].

Thank you.

The chemical reaction taking place there,

the calcium oxide that we have reacts with the water,

and it forms calcium hydroxide.

This doesn't dissolve very much in water.

This is sometimes called an earth.

It does give an alkaline solution.

These were eventually called alkaline earths.

OK, but they couldn't be broken down any further at the time.

So they were simple substances.

They were, as far as Lavoisier was concerned, elements,

simple substances.

But then what happened, well, somebody

did actually manage to isolate some elements, not only

from this, also from our soda.

And this was Humphry Davy here in this very building.

OK.

He isolated from the soda, he isolated a new metal

that he called sodium because it was in the soda.

From this substance here, he isolated a metal

that he called calcium.

OK.

Now, so these elements now appear in the next edition

of Marcet's book.

OK, so these new metals that really are elements,

they feature in here.

But she needed a way of trying to classify these.

Because this is some 50 years or so or more,

60 years before the first periodic table.

So this list is getting quite long of elements.

And she thought, well, how can I divide it up sensibly.

And so what she wanted to do is to classify

things by how they react with oxygen. OK.

So because most of the elements or most everything--

and, of course, the noble gases weren't known at this time--

almost everything combines with oxygen.

So this was a nice, convenient way

to classify the elements, how they react with oxygen.

So in the, well, second edition, third edition, she says here,

"these are metallic bodies that form alkalies."

And so this is now the metal potassium

that forms potash and sodium that forms soda.

So these were the metals that were newly isolated.

We shall look at these in a moment.

OK.

Ah, we also have the metal calcium

that forms calcium oxide here.

Initially, Davy, when he first isolated these and proposed

the names, he named magnesium.

He called it magnium for various reasons.

Eventually, this became magnesium.

So these are the new metals that were isolated.

Now, this idea of how the elements react with oxygen

is something that also Mendeleev picked up

on in his periodic table.

So remember, the very first one was published in 1869.

This was in the first volume of his textbook.

Well, in that one, actually the groups of similar elements

went horizontally across the page.

In the next version that he put together two years later, which

came in volume two of his textbook,

he now has the groups of similar elements going vertically.

But what I wanted to draw attention to

was what's right at the top there.

He calls these vertical arrangement groups of elements.

It's in Russian.

It says "groupa," groupa one, two, three, and so on.

But he also gives these generalised formulae,

the formula of the compounds when they react with oxygen.

Now, so their general form is that we

have two atoms of the element with one atom of oxygen,

then one to one.

Or, as he writes here, two atoms of the element

react with two atoms of the oxygen, and so on.

So there are these different trends

as we move across the periodic table, so

according to how they react with oxygen.

And this is what we're going to look at now.

So we'll start off with group one.

In fact, we've already seen one of the elements

and how it reacts with oxygen. This is, of course, the element

hydrogen. And we know the formula of what

we get when hydrogen reacts with oxygen. We get H2O.

So two atoms of the element, two atoms of hydrogen

combine with one of oxygen. But it isn't just hydrogen

that this is the case.

Other elements can also form similar compounds

with similar formulae.

I'm going to show now reaction of lithium.

So Chris here is just putting some oxygen into the gas jar.

And I'm going to take a little piece of lithium.

So this is lithium foil.

So I can easily cut this with the scissors here.

So I'm going to cut a little piece of foil.

Actually, I'll take a little snip, another little snip

off there we can use, in a moment,

for a different reaction.

So I'm going to put my lithium on this.

Well, this is called a deflagrating spoon.

It's meant to have a spoon on the bottom.

This one doesn't, but never mind.

Anyway, you'll see why in a moment.

So I'm just going to put this, hang this on the bottom here.

So that's my lithium wrapped around the bottom there.

So we've got oxygen in the jar.

And I'm just going to get this burning.

Now, lithium compounds when they're introduced into flame

often give a really nice, brilliant red colour.

So you may see a little bit of red colour, but then,

hopefully, you'll see the chemical reaction taking place.

So an incredibly vigorous reaction taking place there.

In fact, this one is so vigorous that it's actually now

the iron that's burning in the oxygen there.

That glowing thing is the iron.

And the end of the iron has just fallen off.

And my rod has got much smaller.

I shall give that to Chris, thank you.

But the formula of the oxide that we're making is Li2O.

So like H2O, we have Li2O, so two atoms

of lithium with one of the oxygen there.

OK, if I dissolved some of that smoke in water,

it is also an alkali.

It easily dissolves.

It's the opposite of an acid.

But I can make that same alkali by reacting some of the lithium

metal with water.

And this is what I shall show you now.

So remember, I snipped off an extra little piece here.

So I'm just going to add this to the water.

And we can see some interesting things.

So let's add this to the water.

There it is.

Now, remarkably, this floats on the surface of the water.

And this is because, well, this is element number three.

It's only got three protons, three electrons,

and some neutrons.

It's actually light enough, less dense than the water,

floats on top.

But it doesn't mean to say it would make a good material

to make a boat out of.

Because, well, actually, it's now disappeared.

It's reacted with the water to form lithium hydroxide.

OK, so we formed lithium hydroxide and hydrogen gas.

OK, it completely disappeared.

Well, actually, it's not just lithium that does this.

Some of the other elements do.

These also react with water to form alkaline potassium

and sodium lithium isn't in this list

because it hadn't been discovered when Marcet

wrote this edition of the book.

Let's try the next element underneath lithium.

And this is the element sodium.

So we're going to try some sodium with water.

So I'm just going to take a tiny little piece here of sodium

and add this to the water I hope.

There it is.

And it's actually fizzing around on the surface of the water.

Oh, there's a little flash of flame there.

This is a remarkable substance.

I mean, nothing like this had ever been seen before.

When Davy first isolated this here,

he thought, is this even really a metal?

Metals aren't lighter than water.

They shouldn't float.

They don't react with water like this.

This is something that had never been seen before.

But all of our group one elements,

apart from hydrogen, which is a gas,

they all react with water in this same way.

Now, that was a little bit disappointing.

We couldn't really see what was going on there.

Should we try a larger piece?

What do you think?

Should we try a larger piece?

Yeah.

OK, right.

OK, so Chris, I think, has a special piece

of apparatus designed to do this reaction safely.

OK.

OK, so this is a special tank designed

to do this reaction safely.

I'll just ask you, for this one, if you can just

move back just a little bit.

I think you'll need to do that.

Good, thank you.

So this is made of super strong polycarbonate.

There's water in the bottom of the tank.

And there's actually three different lids on top of this.

And this is to make sure that this

is safe to do the experiment.

Because we're going to use a larger piece of sodium.

So I have some larger pieces here.

OK.

And here is my larger piece.

It definitely is larger.

Now, I need to be slightly careful when I'm adding this.

Because, as I say, there are three different lids.

Which way?

Is it like that?

OK.

OK, we're coming in this way, are we?

All right.

Always slightly nerve racking, this one.

OK.

OK, we're ready with this?

So I need to make sure I get this through, I say,

three lids and three holes.

And I need to get this aligned.

OK, ready.

[POPPING]

[LAUGHING]

OK, Chris?

[APPLAUSE]

Well.

OK.

Now, I particularly like doing that experiment.

I think it's a really important one

to do because this is one of the elements that is often

shown at schools.

In schools, it's in the GCSE curriculum.

The teacher tries the little piece on the water

there, as you saw.

And you can hardly see anything.

And then it's, "go on, use a bigger bit."

And then they say, OK, I'll use a bigger bit.

And I tell you, I have heard so many,

so many cases where what's happened there is the sodium

has been on the water and has exploded and then showered

burning, molten sodium everywhere, which

sounds hilarious, unless you're one

of the students getting coated in burning, molten sodium.

And so it really is very important to show this.

And curiously, in the A-level descriptions,

you say that lithium just fizzes quietly.

Sodium fizzes but there is no flame.

Potassium goes with the flame.

As you saw, it's nonsense.

We certainly get a flame there from the sodium.

But remarkably, even though this reaction

has been known ever since Davy's time, this explosion,

it was thought that it was exploding

because of the hydrogen that's given out during this reaction.

But actually it isn't.

And the reason for this has only recently,

within the last few years, been discovered.

What's happening during the chemical reaction

is the sodium, remember that outermost

one electron that was whizzing around

with all of our group one, that electron is easily lost.

And if an atom loses its electron,

it becomes positively charged.

And what happens here, the charge on our molten sodium

is becoming more and more positive.

It's becoming more and more and more and more positive.

And of course, we know that like charges repel each other.

And so what happens when the charge builds up too much,

it pings apart.

It's called a Coulombic explosion

of a charge explosion.

That's the reason that we get the bang there.

This has only recently been studied.

So with high-speed photography, you

can see the sodium literally forcing itself apart.

And this has been modelled using theoretical computers

or theoretical modelling to show that this

is what's happened there.

But this is all relatively new.

OK, so, well, let's go to our group two elements.

So these ones, remember, all have two electrons

in their outermost shell.

I'm not going to show you any beryllium

at the top of the chart here.

Because the compounds of beryllium are all very toxic.

So I can't show you any beryllium.

But I can show you some of the other ones.

In fact, we're going to look at the element magnesium.

And I have a little piece of magnesium here.

And this reacts with the oxygen from the air.

You get a fantastic, brilliant white light.

The white smoke here is magnesium oxide.

OK, but the formula for the magnesium oxide that we're

making there is, well, now it's one magnesium to one oxygen.

Or using Mendeleev's system there,

two atoms of the element, so two magnesium

combined with two oxygens.

Magnesium is in group two.

OK, so should we try a bigger piece of magnesium?

Yeah.

Yeah, exactly.

OK, I have a nice, big piece of magnesium.

This is--

[LAUGHING]

This is magnesium.

It's really, very, very light.

So this is pure magnesium.

So I'll just put this in the flame.

OK, I'll heat it up on the end there.

So while we're waiting for that one to get going,

well, we saw the lithium and the sodium react very violently

with water.

What about magnesium?

So if I take a little piece of the magnesium ribbon,

so here is the magnesium.

If I add this to some water, well, it sinks.

In fact, it's only our group one metals,

only the lightest of those.

So it's lithium, sodium, potassium.

These are the only metals that are less dense than water.

So even all the ones from group two sink.

But actually, even this is not reacting very quickly.

It is very, very, very, very, very slowly reacting.

If I heated up the water to a high temperature,

it would react.

But if I pick the element underneath magnesium,

well, this is calcium.

This one does react a little bit more vigorously.

So if I take a little piece of calcium.

So here is some calcium, adding this to the water.

Again, the calcium sinks.

But instantly, we're getting bubbles

of hydrogen gas given out here.

The chemical reaction is the calcium

reacts with the water to form calcium hydroxide.

Well, that's the same substance that we made here

on the reaction with our calcium oxide with water.

So we're making calcium hydroxide.

That's what's making the water go

very cloudy here at the moment.

Because it doesn't really dissolve very well.

But we're also getting hydrogen gas there given out.

OK.

Well, what about my magnesium?

It doesn't seem to be reacting.

And actually, that's probably a good thing.

It would be a bad day if this caught fire.

But it's because it's too big.

OK, it's too big to react.

This is because at the moment, it's

taking all of the heat from this Bunsen

as I was heating it there, and the heat is being spread out.

And it's being slowly absorbed into the block

here and dissipated from the block as well.

So actually, I can still pick this up.

It's not hot enough yet.

It can't get hot enough for the reaction

to actually start for this to combine with oxygen. OK.

It's just too big.

OK.

And so, yeah, what we need actually

is the very thin magnesium ribbon

to get a nice, quick reaction here.

So this is just too large there to react.

In fact, this was chopped off from an even larger piece,

an ingot of magnesium that was made by pouring

molten magnesium into a mould.

And this would actually melt before it

reacts with the oxygen from the air

and burst into flames there.

OK.

So yes, perfectly safe to heat that one up.

Phew.

But I still feel relieved, nonetheless.

Now, let's go back to Mendeleev's table for a moment

and we'll see one of the real genius things of Mendeleev

life was, and this is probably why

we're celebrating the anniversary of his table

and not the true first table.

This is because of some of these gaps

that he has in his table, these horizontal lines.

Well, some of them, so, for instance, the line that we see

here, this is where he thought, maybe the element yttrium

ought to go here.

So underneath, he's put "YT?"

Well, YT isn't the symbol that we now use.

It's just Y. But he was quite right,

yttrium ought to go in that space there.

OK.

So that one was a known element.

But some of them, so these two were actually

elements that were not known at the time.

And so Mendeleev said, well, he thinks

there should be some elements here.

And he even predicted the properties of these elements

and the compounds of these elements.

And he was spot on.

And this is what got him recognised.

And this is probably why his system, everyone paid attention

to this.

They thought, well, if we can predict these properties so

well, there must be something in his system.

And so this is why we're celebrating his system today.

So let's have a look at one of these.

So we have the symbol EB.

And this is, well, sort of underneath B.

We have B, then Al, and then Eb, El.

So Eb, it stands for ekaboron.

"Eka" was the Sanskrit meaning "one."

It's one space below the boron there,

if we ignore the aluminium.

OK, and the El is one space underneath the aluminium.

So this element Eb was discovered shortly afterwards.

And it was named by the discoverer of the element

scandium.

Mendeleev predicted its atomic weight to be 44.

It turns out to be 45.

But importantly, he predicted the formula

of the oxide of this and how it behaves.

So he said the formula should be,

well, two atoms of the element with three atoms of oxygen.

And when scandium oxide was discovered,

it was found out to be, well, two atoms of scandium

with three of oxygen there.

So all of Mendeleev's predictions were spot on.

OK, well, so where is scandium?

In Mendeleev's system there, he actually mixed up

some of the elements from group three with those from group 13.

This is because, well, both of these groups

have three electrons in their outermost shell

that are available to form bonds with.

And so this is why they're sometimes,

in his system at least, grouped together.

OK, but they ought to be, in this modern form,

separated like this.

OK, well, let's have a look at some scandium.

So I have a little bit of scandium here.

It's very expensive because, actually, it is very rare.

This is why Mendeleev didn't know about it.

Because it hadn't been discovered because it

is quite rare.

So it's quite expensive.

But I thought I should get some to show you

because I'm sure you want to see how it burns.

Who wouldn't?

I had never seen scandium burning before.

So let's give this a go.

So I have some very finely-divided scandium here.

And if I heat it up, it should burn with a brilliant white

flame-- oh, that was quite exciting, wasn't it--

with a brilliant white flame, forming scandium oxide.

So a very violent reaction there,

but indeed it is two atoms of scandium

with three atoms of oxygen. OK.

Right, well, let's pick one of the elements from group three.

So we have boron at the top.

And underneath boron, we have aluminium.

Can we do the same thing with aluminium?

So I have some aluminium.

Now, who thinks the aluminium is going to burn very easily?

Who thinks it's not going to burn very easily?

Well, let's try.

So I have some aluminium foil.

Well, of course, it doesn't burst into flames.

We wrap our turkeys with it and we put them into the oven.

Well, I don't, because I'm vegetarian.

Anyway, but it doesn't burn.

And actually, part of the reason--

it is forming the oxide.

And the formula for the oxide is, again,

two atoms of the element with three atoms of oxygen.

But that's part of the reason why it doesn't burn,

because it's actually being protected

by a layer of this aluminium oxide, which stops it bursting

into flames, which is probably a good thing when we're using it.

Or, of course, with our saucepans,

if we have a large aluminium saucepan,

it doesn't burn instantly.

And again, it's because it's quite large like this one.

But it's also protected with a fine layer of the aluminium

oxide.

Well, but it is forming the oxide.

And so this probably has quite a bit of the oxide on it now.

So we can form this oxide.

And I'm going to try this in a different way.

The formula for aluminium oxide is

two atoms of aluminium with three of oxygen.

And I can make it, but I need a really finely divided

aluminium.

Remember, my big block of magnesium didn't really react.

If I have very finely divided aluminium, it should.

OK.

But I'm going to do this in a slightly different way.

What I've got in the flowerpot here

is aluminium powder mixed with iron oxide.

OK.

Now, the aluminium really does want

to combine with the oxygen. It is a very vigorous reaction

when it forms the aluminium oxide.

So what's going to happen here?

The aluminium powder is going to steal the oxygen from the iron

oxide, and this is going to form aluminium oxide.

That's the driving force for this reaction.

But the other byproduct for this is, well,

if the iron oxide has had all the oxygen taken away,

we're going to be left with iron.

So we should be making some iron.

This is another one.

I'm just going to ask you just to move back just a little bit,

if you wouldn't mind.

Thank you.

It would in your best interests.

Right,

OK.

So, OK, maybe just a little bit more.

You can come back afterwards.

That's great.

That's lovely.

OK.

Thank you.

So now so sticking out of this, I've

got a little piece of magnesium ribbon.

So initially I'm going to light that.

And that will then get the reaction started.

So you'll see a white flame with a white, bright light

from the magnesium and some of the magnesium oxide smoke.

So let's just start the reaction.

Here we are.

So that's the magnesium going.

That's the bright light of the magnesium forming the smoke,

magnesium oxide.

That eventually will get down to the mix

and start properly-- there it is.

OK, that's good.

And could you hit the lights down for a second, please?

This reaction generated so much heat.

And that heat came from the reaction

of the aluminium reacting with the oxygen from the iron oxide

that it actually produced molten iron.

So this reaction is known as the thermite reaction.

OK?

And it generates so much heat here,

it actually liquefies the iron that's made.

And so that was molten iron coming out the bottom.

It's still red hot.

This is actually very useful.

This could be used to make liquid iron, molten iron, when

it's needed in rather inaccessible areas.

So for instance, in the construction of railways,

you can use this to make molten iron

to weld the tracks together.

OK.

Incredibly violent reaction, leaves us with molten iron.

But it's driven by the strong bonds

between the aluminium and the oxygen

there, forming aluminium oxide.

OK, well, let's keep on going through our periodic table.

Let's go to the elements from group four and group 14.

And we have, for instance, the elements titanium, zirconium,

in group four.

But in group 14, we have carbon.

So all of these have four electrons

in their outermost shell.

I'm going to try a little reaction with some carbon.

So Chris is going to put some oxygen into the jar again.

I have, well, this is some impure carbon.

It's actually charcoal.

So this is just what you might use in a barbecue.

So I'm going to put this on my little spoon here.

Now we all know how hard it is to start a barbecue.

OK, so, you know, you're heating up your charcoal,

trying to get it to go, fanning it or whatever.

Well, what you really need is some oxygen. OK, so

if you have some oxygen to hand, preferably liquid oxygen,

always very good.

If you use that, you get a much more vigorous reaction

taking place here.

OK, so it goes very easily.

This is combining with plenty of oxygen.

We're going to form carbon dioxide,

so one carbon with two oxygens.

Or as Mendeleev was looking, so he said, how many oxygens

combine with two of the atom?

So with two carbons, we can get four oxygen atoms there.

And carbon is in group four.

Now, remember Marcet's system.

She says, well, what sort of compounds

are we making when they do combine with oxygen?

Well, unlike all the ones that we've

seen so far, which actually form alkaline solutions,

this now forms very weak acid solution.

So if we dissolve the gas here, carbon dioxide,

in water, or fizzy water, it's weakly acidic We can also

called it carbonic acid.

We get stronger acids if we keep moving

across the periodic table.

So from group four or 14, going to group five or 15,

and on top of that, we have phosphorus.

And phosphorus reacts with oxygen. We get an oxide.

This has the formula two oxygens with three

or up to five oxygen atoms.

And if we dissolve the smoke in water,

we end up with an acid, phosphoric acid.

So I'm going to show you this now.

I have a little piece of phosphorus here.

And I'm going to add this to a flask here.

The flask is full of air.

And there's some hot sand in the bottom.

If we have perhaps the lights down for this, please.

So I'm just going to drop my phosphorus in.

So it's on the surface of the sand.

And it's instantly reacted.

And it's burst into flames.

And it's giving out a fantastic light.

In fact, the name phosphorus means "the light bringer."

OK, it's giving out a lot of lovely light here.

So that's how it gets its name.

But the smoke that has been safely

absorbed by the apparatus here, well, this

is phosphorus and the oxides here.

And we can have P2O5.

It forms phosphoric acid when it reacts with water.

So this is making an acid.

Two phosphorus, though, with five oxygen.

Well, let's keep on going.

After phosphorus in group five, we have Sulphur in group six.

I have some Sulphur in here.

And again, Chris is just going to put some oxygen

into the jar.

OK, and I'm going to pick up a little piece of Sulphur.

So here is some Sulphur.

Now, Sulphur, an old name for Sulphur,

in fact, the name that occurs in the Bible for this

is brimstone.

And this is because, well, you can

find lumps of this lying around, especially near volcanoes.

So it's thought to be a mineral.

But the name brimstone means "burning stone" or "the stone

that burns."

So if I just start this to react with the oxygen there.

And again, if we can have the lights down, perhaps, for this.

So I think it's now going.

And if I lower this in, we get a fantastic blue light, OK,

as the Sulphur is burning.

OK, that's rather nice, huh?

And again, the smoke that's being

absorbed with the solution at the bottom there is an acid.

But the formula for this is, well, we

can have one Sulphur with two or three oxygens.

Or as Mendeleev, when he was considering

how many oxygens, [INAUDIBLE] to be two atoms of the elements.

So two sulfurs would combine with up to six atoms of oxygen.

And Sulphur is in group six of the periodic table.

So if we look at Mendeleev's system again,

OK, he was saying, how many atoms of oxygen

combine as we go across the periodic table?

This was all because of the number of electrons

that these atoms have in their outermost shell that

can be used to form bonds.

So we saw, for instance, hydrogen and lithium

both react with oxygen with two of the metal with one oxygen,

whereas, from group two, we have two

or, well, we'd now say that the formula for magnesium oxide

is MgO or CaO for calcium oxide.

As we keep on going across, we're

increasing the number of oxygens that are reacting.

We got as far as Sulphur.

That was from group 16.

We could have had chromium from group 6.

It forms an oxide, chromium trioxide.

We could have continued, and I didn't.

We could have elements from group seven

and even from group eight.

And they still fit the trend.

The reason I didn't show these, of course, well, chlorine

is very poisonous.

And these compounds are very poisonous and also

highly explosive.

Similarly, for the osmium tetraoxide that we have there,

it is rather poisonous.

And the xenon tetraoxide, well, it's

not poisonous because all it does is explode very violently

and forms the harmless gases oxygen and xenon.

But it's because it's very difficult

to make this compound.

And it is highly explosive.

That's why I couldn't show you any of that.

But the trends continue throughout the periodic table.

OK, what I want to do now is--

now, I did say that I'd try and explain

why it is that our noble gases exist as single atoms.

OK.

And in order to do this, we need to use the bonding machine

that we have here.

So Chris is just going to bring this on.

And we're going to look at, well,

why is it that some of our elements, in fact, most of them

are metals and form very strong bonds between the atoms,

whereas for our noble gases, they exist as individual atoms.

But it's all to do with how many electrons do these atoms have

that are available to form bonds with our neighbours?

So when I have a sample of metal,

such as, a sample of lithium, for instance,

this is my titanium, of course.

And we have many, many countless atoms here.

OK, and they're all bonding together.

But they are bonding together using

just the outermost electrons that they have available.

Now, if we have lots of lithium atoms, each lithium atom,

it can only use its one outermost electron

to form bonds to its neighbouring atoms.

OK, now, these electrons do form bonds.

We can think of these as the negatively charged electrons

going in between the nuclei of neighbouring atoms.

And this helps glue them together.

That is our bond.

So we have our electrons in between the nuclei here

to give us our bond.

But we only have one electron for any of our group one

elements to form these bonds.

And this means that we don't need much energy to separate

the atoms again afterwards.

OK, and this is what this graph shows.

So this graph shows how much energy

is needed to separate a given number of atoms

of any of the elements in their normal form.

So to separate lithium, we don't need an awful lot of energy.

If we move across from lithium to beryllium,

because each beryllium atom has two electrons that it can use,

well, these help to hold all the nuclei even stronger together.

So we form stronger bonds.

So we need more energy to separate

a load of beryllium atoms from this solid

than we do for lithium.

If we keep on going, we get into boron,

and we have three electrons per atom to form all our bonds.

And this means the bonds are even stronger now.

So we'd need more energy to separate the boron.

And if we keep on going, we get to carbon.

And carbon here with four electrons per atom

forms really strong bonds.

We would say that each carbon atom forms four bonds around it

to the other atoms that are near it.

So this is, well, a form of carbon.

This is actually graphite.

And it turns out that the carbon has the highest boiling point

and melting point of all of the elements with perhaps one

exception.

So this is because it has just the right number

electrons to form really good bonds with its neighbouring

atoms.

Now, I have a little demonstration

that I'm going to show you here.

But I need another volunteer.

Oh, your hand went up very quickly, yes.

Would you like to just come round to the front then,

please?

OK.

Excellent, very good.

OK, do you like carbon?

Yes, you do, good.

Which is your favourite form of carbon?

Do you like charcoal or do you like graphite

or do you like diamonds?

Diamonds.

Diamonds, there's no hesitation there.

Right, well, OK, if you come around this way, please,

I have a diamond in my pocket.

This is a real diamond.

And it is several thousand pounds worth.

So please don't break it.

So if you just hold that for me.

That's right.

That's good.

OK, now so that is a real diamond.

A diamond does actually have--

I say, well, diamond and graphite, they

have some the highest melting points, boiling points.

OK.

But also because the bonds are so strong, this means--

you might be able to break it, so don't try.

Because it is several thousand pounds.

But it is the hardest substance.

And so this means that actually that diamond there,

you could scratch anything with that.

OK, you could scratch glass, you could scratch metal,

nothing is harder than that.

It doesn't mean to say it's indestructible.

It isn't.

So this is something you can try later afterwards.

I shall leave the diamond for people to have a go with.

But there's another really interesting property, which

is what I want to show you now.

And this is that diamond is also the best conductor of heat.

Of all substances, there's nothing better.

And this is because of the really rigid bonds

that we have between the atoms here.

Now, what I have is just a little piece of ice.

Now, at the moment, how does this feel?

Does this feel hot, does it feel cold,

or is it room temperature?

In the middle.

In the middle, exactly, room temperature.

Now, what I want you to do is just hold it like that.

That's it, like, just hold it, that's it.

Just there, lovely, like that, yes.

Now, just push gently, push gently onto my ice.

What-- oh.

Now, what did you feel?

Did you notice anything?

It went cold.

It went instantly cold.

And we'll just try this again.

Just try pushing down.

So instantly, it goes cold, and you cut through.

You're cutting through not because it is hard,

which it is.

OK, but you're cutting through because it's

using your body heat.

Let's have another little piece of ice here.

Let's have an ice cube.

So just push against that.

So just gently pushing, and you can feel it instantly

slicing into this.

And it gets very cold, doesn't it?

And that's because it is the best conductor of heat.

It's better than any other substance.

So, of course, we would normally--

I'll just take that one back, thank you.

So we would normally, of course, use metal for our cooking pans

at home for our saucepans.

OK, because they are very good conductors of heat.

But actually, the best substance you could use would be diamond.

So I am predicting now that in the future,

you will get diamond frying pans.

OK, it'd be great.

Because, you know, it would be wonderful and see-through.

You won't be able to scratch it with anything, OK, apart

from maybe another diamond.

And it's a really good conductor of heat,

the best conductor of heat.

So a round of applause, thank you very much for your help

there.

[APPLAUSE]

Now, as I said, I will take this outside afterwards,

and people can have a go for themselves.

Please, do be careful with this.

I say, it isn't indestructible.

I don't want to test it.

It is extremely expensive.

This is a synthetic diamond.

It's made by a process called chemical vapour deposition

where the carbon atoms are laid down sort of one at a time from

the gas phase there to form the diamond structure.

So it's an artificial diamond.

But it is a diamond.

It is expensive.

But it is quite remarkable.

So do that, have to go with that.

So that has the strongest bonding

between any of the atoms of the same element.

But surely, if we keep on going across our periodic table,

we can be adding more electrons and the bonds

should get even stronger.

Well, actually, they don't.

OK.

Well, as we move from carbon to the next element, to nitrogen,

the bonds get a little bit weaker.

They're still pretty strong.

Remember, the formation of, in this case,

nitrogen molecules with two atoms of nitrogen, these

are still pretty strong.

OK.

But they're not as strong as the bonds

that we have between carbon atoms in our graphite

and in our diamond.

And this is because, well, now the extra electrons

that we have, essentially, do not go in between the nuclei.

They go more outside the regions here.

And this begins to pull the atoms apart again.

OK.

If we go from nitrogen to oxygen,

the bonds get weaker still.

We have one extra electron per atom.

And again, these electrons, now, are in regions outside the two

nuclei here.

They are beginning to pull the atoms apart.

We say that these electrons have to go into antibonding levels.

The ones in the middle here holding it all together,

these are bonding electrons.

We say these ones here are going into antibonding levels.

And if we keep on going to fluorine,

they're getting weaker still.

So almost pulling it all apart.

And then finally, if we went to neon, well,

neon has just the right number of electrons that any bonds

that we had would be completely more than cancelled out

by these electrons in these antibonding levels.

There's no bond at all.

That's why all of our noble gases

exist as individual atoms.

OK, it's because we cannot have any bonds because we have too

many electrons.

We'd form equal numbers of bonding

and antibonding electrons-- no net bonding.

This same pattern, when their bonds get stronger and stronger

and stronger and then weaker and weaker and weaker,

we see, not only as we go across the first row

of the periodic table, but also as we go across the next row,

so the elements underneath, we see again the same spike

in the middle for the element underneath carbon.

This is the element silicon that goes down again.

We see the same pattern in the next row

where we see germanium in group 14,

OK, with the strongest bond, and then

get weaker and weaker again in individual atoms

from our noble gases.

In fact, remarkably, we even see the same pattern

because the same reasons.

As we go across the row with the transition metals,

this is our last one.

This is quite fun.

Now, I'm actually-- well, before I show this,

I'll just show you a different way of seeing this.

So we see the same pattern here.

So the bonds with potassium, with sodium, with lithium

are all very, very weak, then really strong bonds

in the middle with carbon, and then individual atoms

down at the other end.

But I can show how weak the bonds are with my group one

elements.

Where's my goggles?

What have I done?

Oh, there, thank you, yes.

All right, I shall just put these on.

I'm going good to show you how weak the bonds are.

So remember the graph here says how much energy is

needed to separate these atoms?

Well, what I have here is lots of atoms of potassium.

So this is a lump of potassium metal.

And what I'm going to do is just heat up this lump.

Now, there's no oxygen in my flask

to react with the potassium.

So we've completely evacuated the flask.

So it's empty, apart from the lump

of potassium at the bottom.

But I'm hoping that we should be able to see a reaction when--

so just heating this up, just warming this,

I'm hoping that it should be quite easy to separate

the atoms of potassium.

So this is what I'm trying to do.

So if I heat this up--

Hm.

Which is quite fun.

So this has now made a potassium mirror.

So because it's quite easy to separate the atoms just

by heating this lump of metal up,

OK, it easily boils and turns into potassium gas,

which is now condensed on the cold flask.

There is nothing for it to react with, OK,

so it instantly just coats the whole of the flask

with this beautiful potassium mirror.

If there was something for it to react with, it would react.

In fact, if I opened the tap and let some oxygen in,

it will gradually react with some of this.

As you see, it's already losing its colour there,

and it's forming the white potassium oxide there.

This is reacting with some of the oxygen from the air.

So, but a beautiful experiment there.

So easy to pull these atoms apart.

But as I say, as we move across one

of the last full rows of the periodic table,

we see the same patterns.

It's easy to separate the atoms of caesium.

I could have done that, but cesium is even more reactive,

and it's more expensive.

So I could've made a nice cesium mirror.

And as I go across, we're putting electrons

into nice bonding levels.

We get really, really strong bonds in the middle

with the element W. What's W?

It's tungsten.

Oh, good, yes, very good.

OK, so tungsten, actually, when I

said carbon has the highest meltable boiling point, perhaps

with one exception.

Well, this one exception is tungsten.

And this is why tungsten used to be

used in the filaments with light bulbs.

Because you could heat it up to really high temperatures and it

didn't melt or boil.

So you'd get it white hot to give out the light.

Not very efficient, of course, because it's also very hot.

It's white hot.

This is why now we have much better light bulbs that

are energy efficient, of course, these light-emitting diodes.

So tungsten-- the highest melting point, boiling point.

Perhaps, you know, with carbon.

It's very difficult to measure.

But if we keep on adding electrons,

well, these end up with the bonds getting weaker.

And this is because, essentially, we

are beginning to put electrons into antibonding levels.

And mercury is a liquid.

OK.

And it's very easy to turn it into a very poisonous mercury

vapour.

Now, this actually has one other interesting consequence.

And this is with the density of these elements.

So as I move from one atom to the next, remember,

as I'm going across from one atom to the next,

we're always increasing the number of protons by one

and increasing maybe the number of neutrons

and the number of electrons.

So the atoms are always getting heavier

as we go from one to the next.

But what about the density?

So the density of an element depends

on the mass of each atom, but also on how closely

and how strongly they're all packed together.

So tungsten has the strongest bonding here.

So actually, we see a slightly interesting curve

with the densities.

So even though the masses of the atoms are getting heavier,

this is showing the density of the solid.

There's a slight lag.

So it turns out that, actually, osmium and iridium

are the most dense ones.

There's still pretty good bonding,

and the atoms are getting heavier.

But even though the atoms continue

to get heavier because the bonding isn't so good,

the density goes down again.

It turns out that gold, Au, and tungsten, W,

have about the same densities.

And this is really pretty heavy.

And I need a very strong, maybe someone

wants to volunteer maybe their dads,

if there's somebody really strong.

Oh, I see, somebody's hand went up there.

That's good.

So right, if you could come down to the front, please.

We'll see how strong you are.

This could be embarrassing, couldn't it?

Anyway, so what I'd like you to do, please,

is just pick up the bar, the close bar to me,

one hand, no sliding, one hand, no sliding.

Well, there we are.

Actually, it's very easy, wasn't it?

So this actually-- oh, don't do that one yet.

OK, this is incredibly light.

This is solid.

It is solid.

It's actually made of magnesium.

This is even lighter than aluminium.

It's one of the very early elements in the periodic table,

not many protons, neutrons, electrons per atom.

Not particularly strong bonds, they're quite large atoms.

It's really very, very light.

This one, we have more protons, more neutrons, more electrons,

better bonds, smaller atoms.

Can you try and pick that one up for me,

please, one-handed, no sliding.

[LAUGHING]

Oh, go on then, you can use two hands.

[LAUGHING]

Actually, it is possible with one hand.

Shall I demonstrate?

So you can just about-- well, if you're good,

you can-- like that.

There we are.

Give that a go.

Yeah.

Now you know how to do it.

It is really very heavy.

Now, usually--

[LAUGHING]

OK, you should try this for yourselves later.

We're going to take both bars outside.

Now, this was completely unfair of me, by the way.

Because, A, I've practised doing that.

But, B, we've made these the same sides

as a pretty standard gold bar.

OK, so this is what a gold bar feels like,

because tungsten and gold have the same densities.

OK, and it really is quite heavy.

So, I mean, holding this, if you will.

We'll have a go later with everyone.

So it really is quite heavy.

That's the same.

That is what a gold bar feels like.

You couldn't run away with many of those, could you?

OK.

But the other interesting thing is, in all the films

that you see, apart from the fact

that they're clearly not gold bars because they really

are very, very heavy.

Remembering one of the James Bond films,

he has it in his pocket and it slides down.

I don't think you could do that very easily.

But they always put the gold bars the wrong way up

in the films.

So they ought to be this way up so you can easily pick them up

with one hand.

All right, you can do this one now, go on.

Yes, yes.

There we are.

Round of applause there.

[APPLAUSE]

And this way it really is not a sensible way to have them.

So that's why they're meant to be that way up.

So thank you very much for that help, then.

[APPLAUSE]

OK.

I say, tungsten and gold have pretty much the same densities,

which is why some very naughty people have actually

drilled out the gold bar, some gold

bars, and put in tungsten instead, and then

covered it with gold on the other end, and you can't tell.

Because it has the same density, only by chopping into it can

you see.

You can't X-ray it.

You just can't see through it.

So I'm not giving you ideas here.

But anyway, so, tungsten, gold--

the same density.

So do have a go with that afterwards.

We will take this outside for you

so that it won't be too busy in here.

Well, that just about brings me to the end of the lecture.

I hope you've enjoyed this.

I hope you've had a fantastic time.

I think, before we go, we should maybe

look at the most abundant element in the universe--

element number one.

This is hydrogen. So where are you?

Up there again.

So we're going to look at hydrogen one

last time before we go.

So I think Chris has-- do you have another hydrogen balloon?

That's great.

Ah, wonderful.

[OOHING]

OK, so we do have one more hydrogen balloon.

I should just say, so there's no oxygen inside the balloon.

OK, so it's not going to be a very, very, very loud bang.

So you should be OK.

You don't need to cover your ears.

If you don't like loud bangs, please do, of course.

But you will be OK.

But you should, I'm hoping that you should all be

leaving with a nice warm glow.

OK, so I hope you've enjoyed the lecture.

Thank you very much for coming.

So thank you very much indeed.

[POPPING]

[OOHING]

Thank you.

[APPLAUSE]

The Description of Investigating the Periodic Table with Experiments - with Peter Wothers