The Modern Alchemist: Water

The Fountain of Youth, Christmas Lectures 2012

38 footnotes:

Suggest footnote

Lecture Two: Water - The Fountain of Youth.

Water is essential to life since every reaction in our bodies takes place in it. But what makes this fluid so special? What happens when you add a lighted splint to a mixture of hydrogen and oxygen? Kaboom! But why? What makes this particular rearrangement of atoms to form water so explosive? Can we tap this energy release to provide environmentally friendly solution to our energy problems? Plants have the ability to reverse this reaction by using the energy from sunlight to release oxygen from water. We are starting to learn how to do the same.  In this lecture Dr Peter Wothers unpacks how energy lies at the heart of chemistry.

We also look at the salts contained in water. Once again we will see the startling difference between a compound and its constituent elements. Take sodium chloride - aka table salt. Sodium is a soft silvery metal that explodes with water; chlorine a deadly poisonous, choking green gas.  Both elements are lethal to us, but after they have met, a dramatic change takes place.  The sodium and chloride ions that form are essential components in our bodies. They help generate the electrical impulses that make our brains and nerves work.

Lecture Two of the Modern Alchemist demonstrates how chemistry plays a vital role in our lives.




Christmas Lecture
Dr Peter Wothers
Royal Institution, London
Filmed in:
The Theatre

The Royal Institution / BBC

Collections with this video:
CHRISTMAS LECTURES 2012 - The Modern Alchemist

Licence: © 2011 The Royal Institution



Early alchemists wrote of a mystical fountain of youth. And for centuries, kings and explorers searched for this legendary spring that would restore the youth to anyone who drank from it. Does such a place even exist? And if it does, what incredible elements might we find in its waters? Join us in the search for the fountain of youth.



Many people searched for the fountain of youth, but one of the most famous was the 16th century Spanish explorer, Juan Ponce de Leon, who spent his life trying to hang on to his good looks. And legend has it that he found the fountain in Florida.

Local people claimed that this miraculous spring could even restore frail old men to perfect health. And incredibly, the spring is still there. On the screen now, we can see the workers there filling bottles from this fountain of youth. And they've very kindly, actually, sent us a sample by air mail.

Here we have our fountain of youth. I can't wait. All right. OK. Let's have a look. "It's on the shelf." OK. [LAUGHS] All right. So actually-- yes, here it is. They've put it over here.

So here is the fountain of youth. It really has come all the way from Florida, which is quite exciting. But we need somebody to try it. Do I have a volunteer to-- oh, well, some hands went up right-- I'm afraid I can't give it to anyone down in the front, actually. You're far too young already. I mean, if I gave you this, who knows what might happen?

We need somebody slightly more mature, slightly more advanced in years. I'm looking-- on the top, actually-- [LAUGHS] to the-- oh, yes. There's a chap, right on the end-- sir, what's your name, please?


Tim. OK. And you'd like to try some of the fountain of youth?

I'll give it a go.

You'll give it a go. That's very good of you. Actually, I think we've arranged for some to be brought up to you.


Goodness me. Is it safe?

[LAUGHS] Well-- And? No, don't drink it all.


Save some for later. So, what? Are you feeling any younger?

Not yet.

Not yet?

Not at all.

No wrinkles lost yet?

Don't know.

Don't know. OK. Well, we'll come back to you later, and we'll see how it's going. But, yes. You were very good. You must be a chemistry teacher, actually. You did ask if is it safe to drink. [LAUGHS]

Really, you shouldn't be drinking in a chemistry lab. So, well done there. Thank you very much. A round of applause for Tim for trying our water.


So this year's Royal Institution Christmas Lectures are all about the elements. My name is Dr. Peter Wothers, and I'm a Fellow of the Royal Society of Chemistry. Now, the ancient Greeks thought that there were just four elements that made up everything around them. And these are the elements-- earth, air, fire, and water.

Last time, we looked at the elements in the air. And tonight, we're going to be looking at the elements in a glass of water. In fact, the water that Tim has just drunk, to be precise. And to help me, we're going to be using our periodic table. Periodic table? Look at that. Excellent. So, yes-- very good, everyone's here. Marvellous.

So we now know, of course, that we've got over 100 different elements. And so, together, we're going to see what elements we can find in our water, and maybe we're going to find something quite miraculous in our glass of water-- the water that Tim's just drunk.

OK. And-- right. I need to put on my coat. Ready for action now. So, thank you. Thank you, periodic table. Excellent. Oh, you're so well trained, really. Excellent. Now. That's perfect. Look at that.

So, our periodic table-- we saw it-- is over 100 elements. And yet, water, of course, isn't one of these elements. We know this now. But this was only first realised about 200 years ago-- just over. So, why was that? Well, partly because nobody had ever done this.





Well. That seemed to bring the house down.


That was certainly rather dramatic. But what we actually did there is just make some water. The balloon was filled with hydrogen--


Which, of course, when we light it--


Combines with the oxygen from the air--


To form water.


OK. Now, that was rather violent. And we couldn't see any of the water that we made, because as soon as it was made, all that energy that was released there vaporised it and dispersed all the droplets into the lecture theatre.

But we can do this in a more controlled way, so we can actually see some of the water being formed. And-- well, this is the apparatus that we're just bringing on now. OK. So, we're going to combine our hydrogen and oxygen in this carefully designed apparatus, and we're passing pure oxygen through this.

And Mark is just going to pass me a tube with some hydrogen.


And let's just light that. Ah, beautiful. It's very difficult to see this flame sometimes, because hydrogen burns with more or less a colourless flame. It really is burning. I think there's a few impurities in the glass there, but you can see it's definitely lit-- there we are-- it's definitely going.

I'm just going to insert this into the apparatus-- I hope I don't catch fire to [LAUGHS] our little things there. There we are. Look at that. Beautiful. Right. So, just insert this. In fact, as it's going in, the flame gets much hotter, because this is an oxygen-rich environment in this apparatus, and it's just beginning to vaporise some of the glass.

And this yellow colour that you see now is excited sodium atoms from the glass. But, I'd like everyone to keep an eye on this during the reaction. This is like our little Olympic flame. If this goes out, please let me know during the course of the reaction.

Because I'm trying to generate enough water to, maybe, taste it later. We'll have to see. We're using very pure gases here. But the important thing that we can see is that we've got our flame-- it's clearly very hot-- but we are generating water.

You can see that this is misting up now at the top here. And we're beginning to get some water, and it's dribbling down here. So this is actually showing us that water is not an element after all. It's made of these two gases-- hydrogen, oxygen.

This-- as we're forming these bonds here, it's releasing a lot of energy. This is as we're making new bonds between hydrogen atoms and oxygen atoms forming our droplets of water that are beginning to collect now.

But I'd like to show you another reaction that clearly generates a lot of energy-- a lot of heat-- as molecules are coming together, as bonds are being formed. And this one is quite miraculous. This is one of my favourite reactions, because it really does look like magic.

What we've got here is a tube just of cold water. Everyone knows what temperature water freezes at, don't we? What temperature does water freeze at?


Zero. Exactly. So, here, it's coming. This is a test tube full of cold water. Thank you very much. And what temperature is this water at?

Minus 3.

Minus 3, this one. So this is actually at minus 3. And yet, of course, we know that water should be ice at minus 3, and yet it isn't. OK? Now, watch what happens when I add just a tiny little crystal of ice. This is now freezing. OK?

So you can see that this water is turning into ice. In fact, the ice is all the way down to here now. So we know that water should be ice at minus 3 degrees, and it is now becoming ice at minus 3 degrees. OK?

In fact, let me see if I can actually turn this upside down-- can I do that? There's a little bit dribbling out, but-- there we are. It is, in fact-- yep-- solid ice there. So that's pretty impressive, I think. Let's hand that out. Thank you.


So, how can we do this? Well, we need to be very, very careful. It's like trying to balance this pen. I was trying it with a pencil earlier, and it just didn't work. But anyway, I can balance a pen on its end there.

If we're careful, we can do this. But just give it a small jolt-- of course, it falls over. Now, we can really cool down our water very slowly, very carefully. We can get it to those temperatures, if it's very, very slowly cooled, really pure water. But just give it a small jolt, it starts off this process of being how it should be.

The question now is, what do you think happens to the temperature? Do you think it stays the same? Do you think it gets hotter? Or do you think it gets colder? So we're gonna have a vote.

So, who thinks it stays the same? OK. Quite a few. Who thinks it gets colder when it freezes? Quite a few of you. And who thinks it gets hotter? It's quite-- a bit of a mix, I think. Well, let's see what happens.

I need to take this very, very carefully. And here we can see that the temperature is indeed at minus 3.7-- this is really pretty good-- in this huge tube here. And I might need some ice. But I'm just going to give this-- oh, look at that. That's just moving it to set this off.

And look what's happening to the temperature, though. The temperature has shot up. It was at minus 3.7. It's now just-- well, it's actually-- this top part is now above 0. And again, it's freezing. This is because, as the water molecules-- well, they were moving around in the liquid.

But as they're locked into place in the ice, we're forming bonds between the water molecules, and the temperature goes up. And that's really beautifully done. And it's incredibly difficult to do. And I think we should thank Fiona for preparing this all day.

Thank you to Fiona.


OK. Thank you.


So the formation of bonds releases energy. This is a very important thing. Now, actually, if we could show this in another way. If you look under your seats, you'll find a little hand warmer.

Now, for those of you-- of course-- watching at home, I'm afraid, if you look under your seats, you probably won't find a hand warmer, unless you remembered to put it there earlier, in which case, well done.


OK. Looks like people are finding their hand warmers. Your hand warmer contains a solution, and it has salts dissolved in this. But actually-- again, it's an unstable situation. There are more salts dissolved than there should be.

So if we just give this a little click-- if you haven't done this already, you just click the little bit of metal in the corner. That's it. I can hear lots of clicking going on. That's just starting this reaction now.


And it's resorted to exactly how it should be. We're forming bonds. These salts are precipitating out. And as they're forming bonds-- forming the solid here-- well, formation of bonds gives out energy, and that's what you can feel. That's what's warming up your hands now, which is rather nice.

OK. So the formation of bonds gives out energy. And that's what's taking place with our little flame here, our little Olympic flame. I think that you might just turn up the hydrogen just a bit. Ah, look at that. Lovely.

So, here we are forming bonds between hydrogen and oxygen, and that's releasing energy. And this is a really beautiful reaction here, because the only byproduct of this is water. Now, you may know, of course, when you burn petrol and other fuels, we get carbon dioxide. This is a greenhouse gas. The only byproduct here is water.

And there's certainly a lot of energy being released. So maybe this could be an answer to the energy problems of the future. Well, have a look at this.




This is Nathan Chang from Valeswood Fuel Cells Limited. Now, Nathan, can you tell us, what is this thing then? It looks like a normal road bike. It is a--


You can put it on the road, can you?

It's a road legal scooter.


But we put the fuel cell on the-- hydrogen on the--

This is-- what is it? Lanthanum nickel hydride, is that right?

That's nickel metal-based hydride, yeah.

Yeah. And it's absorbing the hydrogen--

Like a sponge.

Like a sponge, OK-- gradually releasing it. And of course, it's combining with oxygen from the air.


OK. And that's giving you the energy to power this bike.

Yeah. That fills out-- generate electricity to power the motor, power the vehicle.

OK. This is great. So, you've just driven all the way through the RI. Haven't you sort of polluted everywhere? Is there oil dripping out of everything now?

The only emission is water vapour.

It's just water vapour.

Yeah. And the hydrogen come from water.

The hydrogen itself comes from the water, as well?

Yeah. So it convert water to hydrogen. The fuel cell convert hydrogen to water again.

So why don't we see more vehicles like this now? Why don't we see more hydrogen vehicles now?

Oh, yeah. The hydrogen's still very expensive. And also, you have problem to store enough hydrogen for a bigger vehicle, say.

OK. So you'd have to change your tank, would you, if you wanted to get all the way back to Birmingham?

You need a bigger tank.

A slightly bigger tank. OK.

Yeah. That tank can power this vehicle for 70 miles.

For 70 miles?


Oh, pretty good, then. Yeah. So if we could get a better source of hydrogen, it would be cheaper, then we would see more vehicles.


Because it's non-polluting. It only produces water. Yes?


Well, give Nathan a big round of applause. Thank you very much for coming in.

Thank you.


So, we've got a bit of a challenge here. I mean, hydrogen could be the ideal fuel of the future. But the problem is, how do we get it in the first place? And one way, of course, that we can get it would be to break up the water. In fact, Nathan said that the hydrogen that they're using does come from splitting up water.

But we can see that when hydrogen and oxygen combine to form water, this gives out a lot of energy. If we want to split up our water to generate hydrogen and oxygen we need to put a lot of energy in.

Now we can show this,

This is another general thing of chemistry that, if we want to break bonds, we need to put energy into the system. And we're going to show this, first of all, by looking at the interactions between water molecules. Now, I'd like some volunteers down from the audience-- I think we've got the back row, actually. Give them all a round of applause, please.


Thank you. And I'd like you to take--


Excellent. That's it.


And keep in this space. Lovely. Hold up your balloons, so everyone can see. That's it. Now, you're water molecules, OK? And at the moment, you're bonded together to form the rigid structure of ice. So, of course, you're going to be wiggling around-- just gently vibrating. So this is what ice looks like.

But if we give them more energy, we can begin to break some of these bonds that hold the water molecules together, and we get liquid. So now you can sort of start moving around. Just sort of have a little walk. OK. This is our liquid.

In fact, on the screen now we can see this is a very complicated calculation that was carried out at the University of Cambridge right from scratch. This shows what happens when liquid water molecules get together. They're jiggling around. But they are held more or less in place with these dotted lines that you see there.

These are bonds called hydrogen bonds. This is where the oxygen of one water molecule is slightly negatively charged, and it sticks to the slightly positively charged hydrogen of the other. Keep walking around.

But what happens if we give you even more energy? OK? They start separating. And we can get them to fly out into the audience here. This is where we're making steam, all right? So, lots more energy. The water molecules are separated, and we've got steam. So thank you very much, water molecules. Please return to your seats.


So, all we've managed to achieve so far, then, is pulling the water molecules apart from each other to generate steam. But this raises a very interesting question. How much more space does the steam now take up, compared to the water? So, in other words, if I took one millilitre of water, how many millilitres of steam would I be able to get?

So this piece of apparatus here is designed to try and show us that-- how much steam we can get from one millilitre of water. And I need a volunteer for this one, please. Yes. In the white. Come down front, please. All right.


OK. And your name is?


Connor. OK, great. Now, do you know how many millilitres of steam you're going to get from one millilitre of water? Have a guess. You've got a scale here. You're going to be looking at this scale. It goes from 0 up to 100 mills. How many do you think? Oh, on the spot.


50 mill-- oh, in the middle. OK. Does anyone-- who thinks more than 50? Who thinks less than 50? More than 50? Yes?


100? Any advances on 100? Yes?


150? Well, if it goes past 100, 150's going to be passed up. I want you to watch the dial. I'm going to squirt the water in this end. Here's my syringe. So this is going to be one millilitre of water. I'm hoping that this is all nice and hot.

So one millilitre is not a lot. If you just see-- that's just one millilitre there. And you think 50, don't you? We've got all sorts of different guesses here. So, right. I'm just going to turn this tap and put this in. And I hope-- fingers crossed-- if I spurt that in, OK, close the tap-- and there we are.

Watch this thing, have to see how far it's going to go. It's-- keep going-- it's gone past your 50. It's still going. It's gone past 75. Can we turn this on, actually? Is that possible, to get this back up to temperature?


Ah, brilliant. OK. Yeah, I can see it boiling now. So, we're trying to get some heat back into this thing. How far have we gone so far?


Sorry, how many?


225. Who said 225? Oh, well don't-- OK, but we're still-- we've gone past that now. All right. OK. Still going. We've got quite a bit. Come have a look down this end, actually. Come on. Yep. So, you can see the water in there. Yep, there's still quite a lot. Oh, you missed the dial. How's he doing?


What's that?

Just gone over 300.

300. OK, good. Keep going. Yep. Still some water there. It's still going then, sir. We're now up to-- what are we up to?


450. OK. It's not quite hot enough. It's so difficult to do this. It needs to be so hot. OK, what are we up to now?


650, and we're still going. There's still quite a bit of water there. This is just our one millilitre. OK. And this is now-- what are we up to?



So, quite a few people said it was going to--


Oh, is that 1,000?


That's 1,000. OK. Now, it's still expanding. We're keep counting here. But this is actually really quite important, this expansion here. It's this expansion that drove, quite literally, the Industrial Revolution. It was the power here as water is turned into steam, driving pistons, driving our machinery. How are we up to so far?


1,375. Very precise. You're a keen scientist, I can tell.


Physicist. Ah. Almost as good-- yeah-- as a chemist.


Anyway, how are we doing now? This has come up to--


1,650. OK. We're going to have to stop this now. We're still-- there's a tiny bit there. But actually, this is going to go to over 2000 millilitres. OK. So--


You're a physicist. And you said it was at 50 mills.


OK. Well, anyway. Thank you very much. Give him a big round of applause for keeping track there. Thank you.


So, we've converted our water into steam, and we need to put quite a lot of energy in to do that. But we still haven't actually made our hydrogen. This is what we were trying to do. To do that-- to actually split apart the water molecules-- we need to put even more energy into them.

And we're going to show this now with this apparatus. So we've got some water in the tubes here, and I'm just going to plug this in. So this is connected to here, and this is a generator. So if I hop on the bike, I should be able to drive the little generator at the back.

Oh, yes. We can begin to see some bubbles. It's actually quite hard work here, trying to split up the water. So, I'm going to get somebody else to do all the hard work, I think. Actually, we'd need quite a lot of hydrogen. So, I'd like you to welcome, please, Paralympic gold medalist Mark Colbourne. Thank you very much, Mark.


Thank you for coming along, and thank you for joining us.

You're very welcome.




Some great warmth from the audience here. I mean, you've obviously done a fantastic job. Now, before you get on to this thing, tell us a little bit about yourself. I gather you had an accident paragliding, is that right?

I did, yes. May, 2009. So just over 3 and 1/2 years ago, I broke my back in a near-fatal paragliding crash in South Wales, so very lucky to survive, yeah.

And clearly, it's not affected your legs too much, since you could cycle so well.

Yes. I've been left with lower leg paralysis. So both my feet don't work. So I have to wear special ankle supports. No hamstrings firing, no bum muscles firing. So it's all [SLAP] quads.

All those. Well, it's pretty [SLAP] big quads there. So that's pretty good. Now, this is what I like to see.


Being a modern alchemist, this is some gold. Is it some real gold?

Yes, what you have here is my very own Paralympic gold medalist.

It's fantastic. Could you take it out of this--


It certainly feels pretty heavy. So, what is this-- is this solid gold?

No, it's actually 390 grammes of solid silver, and then you have 22 grammes of gold obviously coated around the outside. And proudly, actually, made in Llantrisant in South Wales.

Oh, it's fantastic. Really good. It's nice to see a bit of gold in the studio as well.


I think we've got a picture of you here just crossing the line. So how did it feel when you were crossing the line, then?

Yes. Just euphoric. It was almost like Christmas and birthdays all rolled into one. Yes.

And you got one of these as well. That's fantastic.

Yes, very much.

But have you ever split up any water molecules before? That's the real question.

No, not in this sense. Obviously, lots of sweat when I'm training. But not for this--

I think you need to give it a go. If you'd like to hop on there--

Shall I?

Then I'll look after this for you.

OK. Wonderful. Make sure he doesn't run away, OK? Because I won't be able to catch him.

Yeah. So. Now then. So I say, Mark's on this bike here. All his power is going to go into driving this little generator here. And this really is just connected to the water that we've got in these tubes. OK. Take it away, then.


That's great. So we've got two electrodes here. On the negative electrode-- well, this is where the hydrogen atoms are collecting. So, remember the hydrogen in the water's slightly positive. The negative electrode is giving them electrons forming hydrogen atoms and forming hydrogen molecules.

On the positive electrode, the oxygen atoms, which are slightly negative in the water, are having those extra electrons drift away to form oxygen atoms, and eventually oxygen molecules.

But the interesting thing here is, as Mark is pedalling away-- go on, faster.

Oh. Yeah.

The interesting thing is that we can clearly see that we're getting twice as much hydrogen gas as we are getting oxygen gas. So this clearly shows that the water is made up of twice as much hydrogen as oxygen. Oh, oh. I think he broke it, actually. Could you go a bit slower?

Is it going again? Sorry.

It's pretty tiring, isn't it? Don't you think it's pretty tiring there?

Pretty tiring.

Yeah. So, thank you very much. I think a big round of applause there for Mark.


Can I offer you a drink afterwards? Do you need a drop of water after that, you--

Yes. Definitely. Lots of water.

Do you want any of this? This is our fountain of youth.


We thought we'd put a guinea pig trying this at the top there, actually. Tim? How is it going? How is your fountain of youth?

Just the same. Don't feel any different at all.

Just the same so far? No wrinkles gone yet?

Not a change.

No? Nothing at all?

No, still there.

OK, we'll come back. Maybe we'll give it a chance later. But anyway, so thank you very much.

Thank you very much.

I'll give that one to you, actually.

OK. Cheers.

Thank you for coming on there.

Thank you.


So, we've seen, then, that we need to put energy in to split up our water. And a lot of energy is needed there. And I don't think that using Olympic cyclists-- even Britain doesn't have enough Olympic gold medal cyclists to power all the vehicles in the future using hydrogen. But there is another way of doing this.

Our plants here use the energy from sunlight to split up water. They spit out oxygen during the daytime, of course. But they're using the hydrogen to build up the molecules that they're made from. So maybe we can learn from nature.

Now, Professor Akihiko Kudo from the Tokyo University of Science has worked on a catalyst here. This is really quite remarkable stuff. This is quite cutting edge. This is a catalyst that can use the energy of light to split up water.

And the catalyst is made of-- well, actually, if we can just have our periodic tables up for a moment-- very good, periodic tables. OK. This catalyst is made up of the elements sodium-- sodium, give us a little wave-- there we are. Very good. Sodium. OK. Can you see sodium there?

And we've got tantalum right in the centre-- very good, tantalum. And oxygen up there. So this is sodium tantalate. And it's doped with lanthanum. Where's lanthanum? There's lanthanum, very good-- give us a little wave at the top there.

So this is the catalyst that he's developed. And this will convert the energy from light and use this energy to split up water. So, at ease, periodic tables. Drop down. Thank you very much. OK. So, we ready?


OK. Now, we're just going to put this on, then?


OK, now. So this is the catalyst sodium tantalate doped with lanthanum, and we are shining UV light on this, and I'm-- yes. We can-- I can see some bubbles. There are some bubbles in the upper part of the chamber there. If we just move up you can see some-- there we are. This is bubbles forming.

So, I say, this is quite remarkable. This is using light energy to catalytically split up-- so this is not changing the catalyst, but it's splitting up water into hydrogen gas and oxygen gas. And we can see these bubbles here.

Now, unfortunately, the slight snag with this one is that it's using ultraviolet light and not just visible light. The plants, of course, use visible light. But scientists all around the world are trying to work on developing a catalyst that will work very efficiently with visible light instead. OK?

And if you can do that, well, maybe somebody here from the audience will be the scientist that actually finds a catalyst that will work with visible light. And if we can do that, well, you're going to be very rich, and you will help to solve the world's energy problems. All right.

So I think it's time that we actually checked the water, after all this time making it. Now I said we had to take a lot of precautions here to ensure that this really is extra pure. This is something that you should never normally do during any science experiment. You shouldn't be drinking the products of the reaction.

But this apparatus here has been specially designed for this. We've used extra pure oxygen-- medical oxygen. We've used extra pure hydrogen here. So I am actually going to just try a few drops of this. Actually, it doesn't taste too nice, to be honest.


But it's basically just pure water. And of course, Tim was using the fountain of youth water. And if we look on the bottle of the fountain of youth water, we see that there are other minerals dissolved in it. This is, of course, because water is a very good solvent. Things dissolve in it.

And-- well, look at all the other components in our water. We've got, for instance, there's a lot of calcium-- there's quite a lot of calcium there. There's quite a bit of sodium in this as well. So sodium-- what's your symbol?


OK, Do you have N-A idea where this comes from?


Oh, thank you. Where does this symbol come from, any idea? No? No? OK. Well, I'm going to-- just give us a-- hold the sign up, please, so we can all see-- that's it. So N-A. Where does this come from? Well, actually I have a book here. Thank you.

Now, in the book we can see this is a chap-- this book is from 1557, and this man here, he's making piles of compound here. And this is actually sodium carbonate-- they called it natron or niter. And this is-- he's taking Nile water here. So this is water from the Nile.

And this is why it's called niter from the Nile water here. This corrupted into natron, and this is the word that gives us the symbol for sodium-- Na comes from the Latin version of this, natrium. But it was first discovered in water. OK, thank you.

So, periodic table, at ease. Thank you. Now, how is this detected? Because we can't see any of these substances in water, because they're only present in such small quantities. We have to use a technique-- the chemist's technique called spectroscopy.

And this looks at how energy interacts with electrons in atoms, so give them some energy, they move up. And as they drop back down again, they can give out this energy as light. And each element has its own unique characteristic-- unique colours. It has its own spectrum. It's like a rainbow bar code for each element.

And we're going to show this now with all these symbols here. So these are the symbols from Group 1 elements. Can we have Group 1 only, please? That's it. If you can put them up.

Hydrogen, well, you are, of course-- you are a component of water. You're not really in water as such, dissolved in it. So you can put your card down for the moment. And francium-- I'm afraid you're too radioactive. So there's going to be no francium in our water. So you can put your card down as well.

But these other Group 1 elements-- well, here they are here. We've actually taken some symbols here and soaked each of these symbols with salts, with compounds of the appropriate elements. And watch what happens when we light them.





Now, we can certainly see that the sodium and lithium are very different. But these ones look rather similar. But actually, they're not. If we were to look very closely at the colours of light coming down-- if we split them up using a spectroscope, we would see that they have slightly different colours.

The exact frequencies of light coming out are unique to these elements. Now, these two elements-- caesium and rubidium-- were also first discovered in water. And they were discovered by Robert Bunsen. This is Bunsen of Bunsen burner fame, of course-- not to be confused with Bunsen Honeydew from the Muppets shown here.

But Robert Bunsen took litres of mineral water, evaporated this, and he found these new elements in the water using spectroscopy. And in fact, he named these elements from the appearance of their spectra-- caesium from the sky blue lines in its spectra, and rubidium from two very distinct red lines in its spectrum.

OK. So, how do we find these elements, though? If they appear in compounds in water, not as their elements. All of you that-- all you Group 1 elements, you're actually metals. We certainly don't find metals in water. And that's because all of these metals actually react with water. And this is what we're going to show you now.

So we have here a tank, and I'm going to and a tiny little piece of sodium to the tank. Let's have this one. So, a little piece of sodium-- a little tiny piece here. I'm just going to drop it into the water. And there it is, dancing around the surface.

So it's lighter than water-- it's floating on the top-- but it's actually reacting with the water. We can see it's reacted there. It's giving out hydrogen gas. The sodium is giving up its electron to the water-- to the hydrogen in the water-- forming hydrogen gas. OK?

It's rather disappointing to see that one, though. It's not a lot there. I think we need, actually, a bigger tank and a bigger piece of sodium.


So, let's try this. OK. There's certainly a bigger piece of sodium. Now, of course, when this reaction is done at school, the teacher is always instructed to not use a piece larger than the size of a pea. And-- well, I thought I'd show you why.

So we've got a piece that is a bit larger than the size of a pea. Now then, we're going to add this piece of sodium, then, to the water. Are you ready? And step back. So it's floating on the surface there-- we could--




Oof. [LAUGHS] And there's certainly quite a lot of smoke.


[COUGHS] Well, you can certainly see why you shouldn't add a piece larger than the size of a pea.



Oh, dear, I've poisoned the audience.


It's chemistry.



Right. Now. But it's not just water that sodium can give its electron to. We can also-- it can give its electron to oxygen. So here is a piece of sodium, and I'm just going to cut this now, and chop it right down the middle. There we are.

So, this is beautiful silvery metal. So this is what sodium normally looks like. But just as it's left here, exposed to the air, it's reacting with the oxygen in the air, essentially giving up its electron to oxygen. OK?

We can see it's changing. It's actually got a sort of white crust developing over the surface. But, actually, if we just have our periodic table up for the moment. OK. And I want to focus on Group 1 again, so others down, just Group 1 up.

So we have sodium here, and the key thing here is this outermost electron that sodium has. I think we have a graphic to show this. This is the atomic structure of sodium. We have one outermost electron there. This is the thing that's very easily given in chemical reactions.

But as we come down the group, you've all got this one outermost electron. But we look at, say, caesium, right at the bottom. So here is caesium, and it has lots more electrons. But again, there's one outermost electron. And this is even more easily lost than it is for sodium.

And-- well, we have some here. And this is caesium in this vessel. And the problem was, we need to store it under argon, and we had a slight problem. We sealed it up so well we can't get it out.


So, I think-- well, the simplest thing to do is actually hit it with a hammer. OK? So-- which is what I'm going to do.




Oh! Well, that did it. Yes. This is such a reactive element that, as soon as it comes into contact with the air-- in fact, if you have the lights down, you might be able to see some of the sparks that are-- oh, yes. Look at this.

This is so reactive, that, as soon as it's coming into contact-- ooh, dear. Lots of sparks. As soon as it comes into contact with the air, it's gone off. It was a nice shiny metal. It was a sort of silvery colour. Sometimes it has a slight golden colour to it.

But as soon as it comes into contact with air, well, it's formed this black oxide. So the caesium is incredibly reactive, and it's given its electron to the oxygen. But there are other things that can take the electron away from these Group 1 elements, and one of them is contained in this-- in bleach.

Now, does anyone know what element is in this?


Chlorine. Where's chlorine? Where are you? OK, chlorine. Yes. You're in bleach. This gives this-- sort of, bleach its color. Except you're a really poisonous, nasty element, I'm afraid. You were used in World War I as a toxic gas. Not very nice. But, yes, you're very efficient, though, at taking electrons from things.

OK. Now, we're going to see some chlorine in just a moment. But before I show you that, I want to show you something else here. And this is a really, really remarkable book. This is from the archives in the Royal Institution here.

And this was from a lecture that was delivered exactly 200 years ago, in 1812. And the lectures were four lectures being part of a course on the elements of chemical philosophy, delivered by Sir Humphry Davy. OK? Exactly 200 years ago.

So Humphry Davy, he was the first person to isolate, among other elements, sodium and potassium. And he also named chlorine-- over there. And these lectures were written down during the lecture course by a young Michael Faraday, who was sitting in the audience exactly where you are now.

Davy was so impressed with that these notes that Faraday wrote up of the lectures, that he gave him a job. And he got a job here at the Royal Institution and became one of the world's most famous scientists ever. Now, I'm certainly no Humphry Davy, but who knows? Sitting in this audience there may well be the next Michael Faraday.

Now, this is what I wanted to show you, though. In this book here, experiments belonging to the lecture on chlorine, it says, Mr. Davy exhibited a specimen of chlorine gas. It was in a clean glass tube. OK. Now, we actually have that original sample of chlorine gas. Here it is.

This is the one that Davy exhibited back in 1812, which is quite remarkable. So chlorine gets its name-- Davy gave it its name-- from the Greek chloros, meaning green, because of this greeny colour that it has. We even have a sample of the original sodium.

And this is Davy's original sodium, prepared here at the Royal Institution. This is the metal, just floating, protected in the oil, because, as we've seen, it goes off in air very, very quickly. So, we've got sodium and we've got chlorine. Well, what happens when you mix the two?

Well, you get some sodium chloride. Now, they would kill me if I did this, of course, with these. So I'm not going to use the original samples of sodium and chlorine. But we do have some other samples here. Here's our sodium. Here's our chlorine.

And you can see this beautiful green colour now of the chlorine. You can see the beautiful silvery colour of the sodium. We've actually taken all the air out of this side, because, of course, we know that sodium reacts very violently with air. So there's no air here.

I'm just going to pop my goggles on, if I can do that one-handedly. There we go. And I'm now going to let the chlorine from this side into this side, and see what happens. So, let's have a look. Look at that.

This is sodium meets chlorine. And the silver mirror has disappeared, and we've got this white crust forming all the way around here. It's gone white. Looks like salt, and-- well, that's because what we've made here is salt, it's sodium chloride.

So the chlorine reacts with the sodium. The chlorine takes the electron from the sodium to form white sodium chloride. This is just the sort of thing you'd put on your chips. But chlorine itself is poisonous. It reacts by taking electrons from your body, and poisons you.

The sodium is poisonous, it would give its electron to you. But once they've reacted, we've got sodium chloride, and you eat it. OK. Thank you very much.


But there are actually different sorts of salt. Sodium chloride is just one salt. There are others. And if we have our halogens up for a moment-- OK-- all of you are very good at forming salts, especially with our Group 1. Can we have Group 1 up as well, please? OK, any mixes of you would form salts.

So, for instance, not only could we have sodium chloride. We could have, say, potassium bromide, or potassium chloride, or so on. In fact, all of you halogens-- do you know what the name halogen actually means? Does anyone know? Does anyone know anywhere? It means salt former. OK?

So you're all really good at forming salts. In fact, that's how, as a group, you get your name. So salts, though, have very different properties when they're dissolved in the water. The water itself is completely different. Sea water is not the same as normal water. And-- well, we have some seawater here to show you.

Thank you very much, periodic tables. If you go down for a moment. So, I was fortunate to visit the Dead Sea. In fact, this is me floating in the Dead Sea up here. So, while I was floating in the Dead Sea, I was thinking, well, what would it be like? What sort of things could we get to float on the Dead Sea? OK?

And-- well, I need a volunteer, actually, to help out with this. And we'll have somebody from-- yes. With the-- yes. Yes. If you'd like to come down. Lovely.


OK, good. If you'd like to face the front here. So, what's your name, please?


Katie. OK. All right. It's good. Now, we need to put some protective clothing on you. So if you just come over here and put on some protective clothing. I've got a block of metal here. Now, what do you think to this block of metal? If you would just hold this. See? What do you think?

Really heavy.

It's really, really heavy. Yes? OK? Really heavy? What do you think? How heavy?

Quite heavy, all right.

It's quite heavy, isn't it? So, this is solid metal. This is actually-- you're looking very good in those. Very fetching. Yes. OK, now. We've actually got a step here, which was a very good thing. If you'd like to stand on this step.

OK. This is some salty water. This is essentially sea water. I'm going to put on my gloves as well. And, do you think-- you're haven't held this yet. If you hold that, you just stand for a moment. [LAUGHS] Right. So, what do you think of that? Quite heavy?

Yeah! It's quite heavy, actually.

It's quite heavy, isn't it? Yes. It is quite heavy. Right. And do you think it's going to float in the Dead Sea or sink in the Dead Sea?


What do you think? 50-50. Float or sink? Ask the audience?


Ah. Float. OK. Float in the Dead--

Maybe. It's quite heavy

Maybe? Oh, covering your bets there.

It's quite heavy, though.

OK. Right. It is quite heavy, though, isn't it? So let's see, shall we? So we're going to drop it in. If you can take that side-- you hold that one-- that's it-- and you put it back in the water-- just lower it in, gently. You don't want to splash it everywhere. That's it. OK. Very good. And just let go. And--

Aw, I was wrong.

Aw. [LAUGHS] Well, it sinks. It did sink.

Yeah, it was quite heavy.

It sank pretty quickly. It was quite heavy, wasn't it? Yes. OK? Now, it did sink pretty quickly there. Now, let's see if I can just fish it out for you. I'll do this, because I've got longer gloves. And-- ooh. Not that long.


[LAUGHS] Anyway, right. OK. Out comes the magnesium here. This is a block of pure magnesium. It is actually quite heavy, isn't it? You're right. OK. There we are. Now, we've got some other water here.

So this one contains sodium salts, but this one contains caesium salts. And if we can just have our Group 1 up again for a moment? Ah. Very good. Very efficient. So, sodium right at the top there. And as we go down, we get to caesium. And so, caesium is actually heavier than sodium.

And-- now, how quickly do you think this is going to sink in this one?

Maybe a bit longer now.

Take a bit longer to sink. Yeah. OK. All right. So let's-- should we try this one, then? So I'm just gonna put this in the water. You ready? OK? Do you want to hold it there as well? Just gently let go. And-- all right-- ready? After three? One, two, three, go!




[LAUGHS] Watch out for your gloves, don't forget-- [LAUGHS] No. So it's not actually going to sink. It actually floats in the water there. So I'll just fish that out. And I think you should get a big round of applause for your help there.


Thank you. So we'll just take those off you. That's it. That's it, there. Ah, lovely job. Thank you very much indeed. Yes, the block of metal actually floats, then, in the caesium salt. So that's because caesium itself is a heavier atom than sodium is.

We can get other salts, though, from the Dead Sea, for instance. It's not just sodium chloride. In fact, in the Dead Sea, there's-- particularly, it's quite rich in another salt. If we have our halogens back, please? OK. We have bromine in the Dead Sea as well.

It's not as bromine itself. It's not as the element. It's as bromide. This is where it's taken an electron from something and formed a negative bromide ion. So we get things like potassium bromide, sodium bromide, if we evaporated our Dead Sea water.

And, actually, I can demonstrate this with some of the Dead Sea water. I've got some concentrated Dead Sea water. I actually brought this back from the Dead Sea while I was there. Just the sort of thing that you normally bring back, I suppose, if you're a chemist. Anyway, right.

Here we are. Here's my Dead Sea water. And I'm just going to add some of the bleach to that. And of course, the bleach has chlorine atoms in it. Watch what happens when we just give it a bit of a--


OK? This-- the colour that you now see-- and we haven't cheated in any way. This really is just Dead Sea water that I brought back. We just concentrated it a little bit by removing some water. This is just bleach. The colour that you see now is bromine.

So, with our halogens-- we've got bromide here, and I've just added some chlorine that was in the bleach. So what we have there-- the bromide reacts with the chlorine. Chlorine takes the electron from bromide, and forms bromine elements-- OK? And leaves chloride ions.

But, chlorine-- look behind you. Who's behind you?


Fluorine. Exactly. Now fluorine is even better at taking electrons away. And fluorine can take the electron away from chloride. OK? Fluorine-- in fact, fluorine, where are you-- I say, give us a little wave.

Fluorine, you are the most reactive non-metal in the entire periodic table, You will steal an electron from every other element in the entire periodic table, with the slight exception of the very inert noble gases sitting next to you. But everything else, you will react with. Very, very violent.

Fluorine is probably-- well, the most reactive element in the entire periodic table. Certainly, the most reactive non-metal. Even as a chemist, I had never seen any fluorine. And I thought, well, for these lectures, it would be really nice to bring some fluorine into the lectures. And I needed to find a specialist to do this.

So, would you please welcome, from the University of Leicester, Professor Eric Hope, a fluorine chemist.


Well, thank you, Eric. Thank you for coming along. So, Eric is a fluorine specialist. Fluorine is incredibly reactive. It reacts with just about everything, doesn't it?

Indeed it does.

So-- I mean, the one question that everyone wants to ask is, well, how do you store it, then? Doesn't it react with the container that you put it in?

It does react with the container you put it in. You can store it in metal containers-- you can store it in stainless steel. Or we hold it in nickel containers in Leicester and what happens is the fluorine reacts with a coating of the metal, you get nickel difluoride, a very few microns thick. Protects the rest of the nickel metal. And you get an impervious layer.

So you could store it, because it has reacted and it's--

It has reacted.

And what it's formed is pretty inert afterwards.

Indeed, yes.

But, fluorine itself is incredibly reactive, and very dangerous, isn't it?

If you control it, and handle it under appropriate conditions, then it is dangerous, but it's not that dangerous.

Oh, it's very reactive, isn't it? It is very, very reactive.

It will react with-- as you said earlier-- virtually every element in the periodic table.

Yes. When I contacted Eric, I thought, fluorine really reactive, it would be very great if we could bring some of this into the RI. But one reaction that I've always wanted to try is the reaction between fluorine-- because it is the most reactive non-metal--

I thought, I'd like to try the reaction of fluorine with the most reactive metal in the periodic table-- caesium! Yes. You are the best electron giver. Very generous. So this should be a really violent reaction. I thought, I'd love to see this.

Now, what did Eric-- what did you think of this, when I said I wanted to try the reaction of fluorine with caesium?

I thought it was the most outlandish thing I'd ever heard. And it took me a good 24 hours just to think about it, and think whether or not it was feasible or possible to actually do.

OK. So, yes. Exactly. We did wonder whether we can actually do this here. So we just have the periodic table down for a moment. We thought, how could we do this safely? We did have a practice, just to make sure we could do this safely, and show you.

In fact, the remarkable thing was-- I mean, as I say, I'm a chemist, and I've never actually seen-- I hadn't seen fluorine before, before I went up to Leicester. And remarkably, well--

I've never seen caesium before.

Eric had never seen caesium. So, we thought we had to get together, really. So, I thought I'd bring my caesium along to Leicester. And we did test things. And we're going to try and do this for you now. Now, I'm actually just going to just put-- hold this in my hand.

Just holding this in my hand, actually, just melts the caesium. So the bonds between the caesium atoms here are so weak, just holding it in my hand melts this. And there we are. Now, we've certainly got some caesium at the bottom of this tube.

Right. So we're going to try, then, the reaction between caesium and fluorine. The caesium is protected by this blanket of argon, which is heavier than air, because we've seen that caesium reacts with the air.

And I'm just going to lower this on, like so, and push this down. OK. And I think we're ready to try, then. Ah!




OK. So it's an incredibly violent reaction. We're using tiny quantities-- well, we're using more caesium here-- but a tiny quantity of fluorine. We have to just use a tiny, tiny bit. It's is in this tube here. But it did react very, very violently with the caesium to form caesium fluoride. OK.

Well, I think we need to give a big round of applause to Professor Eric Hope. Thank you very much.


So once the caesium has reacted with the fluorine-- again, the caesium gives up its electron, and it's no longer the reactive caesium metal that we started with. But of course, the fluorine-- once it grabs the electron, it's no longer the reactive fluorine that we started with. We've got fluoride ion.

And actually-- remarkably-- fluoride is something that, again, you can find in every glass of drinking water that you have. This is actually added to our tap water, because it protects our teeth and prevents decay.

Now, talking of preventing decay, I think we need to see how our fountain of youth volunteer has got on, and see whether it really has worked after all. So, Tim, how's it doing? Are you-- is it worked?


I feel a lot younger.

Tim, is that really you? Ah.


I can see what they did there.



So, Tim, do you feel any different at all?

No different at all, I'm afraid.

No different at all.

Just the same age.

OK. So, it really looks like our fountain of youth water perhaps isn't restoring Tim's youth to him. But nonetheless, water does play a very important role in our lives. Without it we'd all be dead. And this is because water actually enables chemical reactions to take place, both inside the cells in our body, and in other reactions.

And this is what we're going to demonstrate now. I have-- we've placed some magnesium powder and some silver nitrate. Now, they are touching each other at the moment. They're all mixed up very intimately, but not quite intimately enough.

What they need is to get into really close contact with each other. And we do that by adding a drop of water. OK. Now, they've given me a glass of water and a pipette, but I think I'm going to need a slightly longer pipette, please. This is quite a violent-- thank you. That's a better one. All right. Now.

OK. Now I'm going to take the lid off of this. It's sitting here quite happily at the moment. Nothing is taking place until we add a drop of water. OK. I'm adding it now.





So this is an incredibly violent reaction that takes place. But it didn't take place until we added the water. So maybe this gives us a method for slowing reactions down, if we remove the water. Well, I think-- actually, it's getting time for my dinner, and they've brought on some bananas! Great.

And these bananas have seen better days. Now, these bananas, I know, are two weeks old, actually. They've been sort of-- ugh. Hm. [LAUGHS] It's actually a banana--


OK. It's seen better days, I think. But this banana-- this banana-- this one was two weeks old. This banana here is actually six years old.

Oh! Urgh!

OK? But if I was going to eat one, I know which one I would rather have. Now, why has this one lasted so long? Actually, just-- what would you think? Have a look at this. What do you think? What does it feel like?



Scaly? Uh-huh. What do you think? What do you think it feels like?

Quite hard.

Quite hard. Yes. It is quite hard, and this is because we've removed all the water from this banana. But this has, actually, preserved it. It's stopped the reactions taking place-- the normal reactions, where things go bad-- reactions taking place in the cells. If we remove the water, they can't happen.

And, well, maybe this could be way of preserving our good looks, if we just remove the water. Well, you could. And this is what you might look like.


OK. Well, this chap might not look too good, but he is 800 years old, which is quite remarkable. And the reason he's survived looking like this is because all the water was removed when he died. He died in the north coast of Peru. It's very dry there. And removing the water has actually preserved the cells of his body.

So maybe then, that this does give us a clue to eternal life-- remove the water and you can, well, live a long time, or looking like that. But, of course, we need water for our reactions to take place, and so-- well, I certainly know which one I'm going to choose.

I think I'm going to stick to drinking the water and staying alive. Mm. Freshly synthesised water. Delicious. And I think, actually, it's time for you to all have some freshly synthesised water as well.

So that-- you may have noticed that all around the lecture theatre here is this very long tube, which is over half a kilometre of tubing filled with hydrogen and oxygen in the right proportions to make water. So we're going to synthesise some water now. So, I need the ends of the tubes please.

OK, here they come. Excellent. Thank you very much


OK. Now, the ends have some corks in, and we don't want to fire these into our audience. So we're going to fire it into this bucket of water.





No, you really don't want to say, ah. You wouldn't want this, I can assure you. All right. OK. But before I do this, though-- before we do this final thing here-- I hope you don't feel too disappointed that we haven't found the secret to eternal youth. But we have uncovered a whole host of exciting elements in that one glass of water.

Some of these elements were deadly toxic, and others were explosive metals. In the next lecture, we're going to find out how chemists are trying to extract exciting elements from the earth. These are elements that have been trapped in rocks for billions of years.

And we'll also try to solve the biggest alchemical mystery of them all-- how to turn lead into gold. But before we finish, though, I think it's time for everyone to get their freshly synthesised water. So we'll have-- ah, yes. I need move this from here.

OK. And we'll have a countdown from three when we're ready. So we're just going to aim these. And if we have the lights down. OK. So now, you might want to be looking up rather than down at us. OK? You look up at the pipe. OK, so-- three, two--






[LAUGHS] Thank you very much.


Good night. Thank you.


Related Videos

Collections containing this video: