[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]
