Thank you very much.
This is a lecture on a much simpler topic than you had just before.
Catalysis concerns a rate of a chemical process
and the rate of a chemical process is concerned with the breaking and formation of new bonds.
And if you look at the progress of a chemical reaction, it’s always associated with changes in energy.
So whenever you form a new bond,
you have to break bonds and you are forming new bonds
and you have to overcome an energy barrier which is the energy of activation.
The rate depends exponentially on the activation energy.
The higher the barrier the slower the rate.
A catalyst offers an alternative way for this reaction by forming intermediates which is shown here by this dashed line.
And in this way the overall rate becomes faster.
If the catalyst is in the same state as the molecules to react then it is homogeneous catalysis.
And you will hear more about homogeneous catalysis during this conference.
In chemical industry and for most practical applications the catalyst is in a solid state.
So the interaction of the molecules with a solid surface come into play and this is called heterogeneous catalysis.
The principle of heterogeneous catalysis is shown here.
You have a solid surface.
The topmost atoms of the solid have three valences.
They can interact with particles arriving from the gas phase, make bonds
which are called chemisorptions bonds and then this piece on the surface,
these are then intermediates, are forming new molecules which can be released back into the gas phase.
And under steady state conditions we have a continuous flow of reacting molecules
and we have a steady state rate of the reaction depending on the external parameters,
the partial pressures, the temperature and the nature of the catalyst.
And it is the task of the chemical engineer to optimise it.
So heterogeneous catalysis is a basis of the chemical industry.
And one of the most important heterogeneous catalyst reactions was started in industrial production exactly 100 years ago.
It’s a Haber-Bosch process, the formation of ammonia from the elements nitrogen and hydrogen.
And this is one of the plants.
The next figure shows the variation of the world population and the variation of the production of ammonia over time.
And you see there is a close parallelity.
There is an estimate that 1/3 of the work population would starve without this reaction
because this reaction is the basis for the formation of fertilisers.
As I said it was started in industrial production at BASF just 100 years ago
but it took about 7 decades until this reaction was understood in its details.
And the mechanism of the reaction has to do with the complicated structure of the catalyst.
That’s a catalyst which was developed by Mitesh at BASF
and it’s essentially an iron oxide catalyst with small amounts of other elements like potassium, aluminium and calcium.
That’s the surface. These parts here are the oxide particles of aluminium and calcium.
They form a framework which prevents the small iron particles from sintering at the high temperatures of reactions.
While the catalytic active spots are essentially consisting of metallic iron covered with a monolayer of potassium and oxygen.
The potassium is a so called electronic promoter.
It enhances the activity of the iron particle to dissociate the nitrogen molecule which is a rate limiting step.
So you see how complicated a real catalyst is.
And that’s why it took so long to understand the mechanism of the reaction which is summarised here.
If you start with nitrogen and hydrogen and form ammonia, that’s the overall enthalpy of the reaction.
If you would do it without a catalyst just by splitting nitrogen and hydrogen,
you would need an enormous amount of energy to form the atoms
which then like in a cascade could recombine to the formation of the ammonia.
The catalyst offers an alternative path by splitting the nitrogen and hydrogen molecules.
This is a decisive step here forming particles of nitrogen and hydrogen
which can recombine to NH, NH2, NH3 which finally dissolves.
So this is the mechanism of this reaction.
It looks very complicated and that’s why I will not explain to you in detail the mechanism of this reaction
but go to a much simpler reaction where we can also see the principles underlying catalytic reactions.
That’s a reaction which is taking place in the exhaust catalyst of your car
where we have an oxide carrier with small particles of platinum or palladium
on which the toxic molecules, carbon monoxide and nitric oxides and hydrocarbons are transformed into less toxic materials.
Like for example here the carbon monoxide is oxidised to CO2.
And these are the plains exposed to these toxic gases.
And as I mentioned there’s usually platinum or related elements.
If you look closer at such a catalytic particle... This is shown here.
This is an oxide support onto which small particles, in this case of ruthenium, are deposited.
And that’s a high resolution electron micrograph showing the atomic configuration of this metallic atoms.
And this is a ball model of this catalyst particle showing the different crystal plains
which are exposed and defects like this edges or corners.
And so you see that the surface itself is not homogeneous, it’s heterogeneous.
It exposes different structural elements with different reactivity.
So to get inside into the elementary steps one has to look for a model system.
And this model system is just looking for individual crystal plains, single crystal plains with the tools
which are offered by surface physics.
And that’s why it’s called the surface science approach.
So we can use the single crystal surfaces and look at the atomic details of the structure and of the dynamics.
So we have seen that the decisive steps sometimes or very often is the dissociation of a molecule.
How can the surface dissociate a molecule without much energy demand?
This is shown here. You have such a molecule, a diatomic molecule approaches a surface.
This is the energy balance. If the molecule would be attached to the surface itself, it would find only a weak binding.
If it dissociates a molecule in the gas phase, we need dissociation energy.
But the atoms themselves will make a strong bond.
And it’s the crossing of these 2 lines which marks the activation energy for overcoming dissociation.
In the 2-dimensional diagrams it shows the variation of the energy
as a function of the separation X of the molecule from the surface and the separation Y between the 2 atoms.
So the molecule is absorbed weakly.
It overcomes dissociation barrier and then it separates in this way.
And that’s the secret of the catalyst.
We are forming dissociated species on the surface which are then strongly held on the surface and can react further.
So the secret of the catalyst consists in overcoming this barrier
and make on the other hand the bonding to the products not too strong because they have to react further.
And that’s why we are given reactions as usually always only a single material which is optimised for this purpose.
You can look at the atomic structure of such a surface with a scanning tunnelling microscope.
That’s a platinum 111 surface, these are platinum atoms.
And if you expose it at low temperatures to a small amount of oxygen,
we find pairs of dots which are oxygen atoms resulting from the dissociation of the oxygen molecules.
They are not on nearest neighbour sites.
They are always separated by 5 to 8 angstroms from each other.
And this is related to the fact that upon dissociation energy is released which has to be transferred to the solid.
And this takes some time. And from the distance and the energy we can estimate this time.
It’s in the order of about 3 times 10^-13 seconds.
That’s a typical time scale for energy exchange between particles on the surface and the solid.
This is at low temperature. So here the atoms are fixed.
If we increase the temperature the atoms, can charge or start to jump to neighbouring sites.
And the next figure shows a snap shot from oxygen atoms on a ruthenium surface at room temperature.
And a snap shot means we can also follow the motion in real time.
That’s a motion of the atoms on the surface in real time.
The oxygen atoms jump from one site to another one.
And this is like a Brownian motion in 2 dimensions.
If you look closer, you find that the atoms have residence times on certain sites depending on the neighbourhood.
If they come closer to each other they stay a little bit longer.
And this is a consequence of the fact that the atoms are feeling each other.
There is an interaction through the electronic structure of the substrate.
And there is a weak attraction between pairs.
And if we increase now the coverage, the consequence is that we have no longer random distribution
but we have a separation into 2 phases.
A quasi-2-dimensional solid phase and a 2-dimensional gas phase like with solid/gas equilibrium.
There’s a 2-dimensional solid/gas equilibrium.
And the atoms condense and evaporate again.
So there’s a real equilibrium situation showing that at higher coverage
and low enough temperature we have a tendency for formation of ordered 2-dimensional structures.
And these structures can be analysed by diffraction techniques like x-ray diffraction
so that we can really determine the actual structure of this chemical overlayers.
And for our problem of oxygen and CO interaction on platinum type metals
I show you just the result for CO and oxygen absorbed on a rhodium surface.
The CO molecules prefer to occupy this on top sides, that’s a side view here, and form densely packed unit cells.
While the oxygen atoms form open meshes, prefer to stay in the 3-fold hollow sides.
And that’s more or less the saturation coverage with oxygen.
So if you expose the surface to both species, oxygen and CO, within these unit cells
there is still space for the absorption of carbon monoxide.
So we form this mixed phase. While on the other hand, if we have too much of the carbon monoxide on the surface,
there is no space available for oxygen dissociation.
So that CO molecules block sites for oxygen absorption.
And that’s why under steady state conditions the surface will usually be covered by CO and the reaction cannot proceed.
One of the reactants is poisoning the reaction itself.
It’s like in biological objects where are blocking the sites by carbon monoxide so that you can no longer take up oxygen.
We have to increase the temperature.
Usually part of this carbon monoxide desorbs into the gas phase so that oxygen can be absorbed.
That’s why the catalyst of your car doesn’t work at room temperature.
You have to go to elevated temperature
in order to have part of the surface under steady state conditions always free for oxygen absorption.
And we can put this together into the mechanism of the reaction showing
that the CO molecules have a tendency for formation of densely packed phases.
And if the concentration is too high, oxygen cannot be absorbed.
While at the other hand oxygen, if it’s absorbed, forms this open structures.
In between the CO molecules can be absorbed.
These are the individual steps.
And this is the energy diagram.
You start with carbon monoxide and oxygen.
We absorb both on the surface.
That’s the energy of absorption.
By overcoming a small activation barrier of about 100 kilojoule carbon monoxide and oxygen can recombine,
form the CO2 which is only very weakly held and is immediately release into the gas phase.
This was the result with a perfect single crystal surface.
But as I have shown you before a real surface has also crystal defects, structural defects.
And they may be very important for the activity of a real catalyst.
Now this can be seen here with this example.
If we expose this ruthenium single crystal surface to NO,
-and this line here shows a mono atomic step running across the surface- after 6 minutes of exposure we get new structures.
And these triangles are the nitrogen atoms.
The oxygen atoms are faster that move across the terraces.
So the steps obviously are the active sites for dissociation of the molecules.
After 2 hours all the nitrogen atoms have spread across the surface.
And that’s not the whole story.
These terraces are the sinks for the molecules arriving from the gas phase.
So this is a cartoon showing that the NO molecules are absorbed on the terraces.
They move to the steps where they dissociate.
And from there they can move back onto the terraces.
This was with ruthenium.
Ruthenium is the only metal among the platinum group metals which is not very active in this carbon monoxide oxidation.
It’s in fact a very poor catalyst.
And there was a long standing controversy because it turned out that at high pressures ruthenium becomes active
while at low pressure it’s inactive.
This was considered to be a manifestation of the so called pressure gap.
The answer has to do with the state of the surface itself.
And if you look at the activity of the surface at high pressures at room temperature,
we still have this monolayer of oxygen atoms which makes the surface completely inactive.
But if we increase the exposure to oxygen at elevated temperature, it becomes active.
And this is the rate of CO2 formation at elevated pressures as a function of the concentration of oxygen on the surface.
Up to 1 monolayer is completely inactive.
If there is more than 1 monolayer, it becomes active.
So obviously the surface itself is transformed into another state.
And this new state under reaction conditions is a transformation into ruthenium dioxide, the 110 plain.
This is a chemisorpted state on ruthenium.
This is a new plain of ruthenium dioxide.
This is the active phase. And its structure is shown here.
It exposes oxygen atoms on bridge positons and coordinatively unsaturated ruthenium atoms
which can be occupied by oxygen or CO molecules.
This is a vibration spectrum which is just a finger print of the 2 states of oxygen on the surface.
And we can also follow the reaction between CO and oxygen by looking at the variation of the vibrational spectrum with exposure.
We start with an oxygen covered surface with a 2O and expose it to carbon monoxide at room temperature.
The oxygen is removed. It has reacted off away and it’s replaced by carbon monoxide.
And if we have now the carbon monoxide covered surface, expose it to oxygen, the reverse takes place.
So it’s a very simple example of so called Langmuir-Hinshelwood mechanism.
Both reactants compete for free sites on the surface.
And if they are on neighbouring sites they can form the product CO2.
Theory has in the meantime advanced very much thanks to the development of density functional theory
which was invented by Walter Kohn who is also here.
And my colleagues have calculated the energetic changes
if we have 2 neighbouring sites occupied by carbon monoxide and oxygen atoms and bring them together.
This is an energy diagram. And you see there is an energy barrier for formation of CO2 of about .9eV.
So we can really calculate now from first principles the energy barrier for the formation of CO2.
We can go even one step further and evaluate all the rates under steady state conditions.
The rate in this case of course depends on the partial pressure of oxygen, of CO and on temperature.
So if you keep 2 of these parameters constant
and measure the rate of CO2 formation as a function of CO pressure at fixed oxygen and temperature,
you see that the rate increases to a maximum and then decays by further increase of CO pressure.
These are absolute rates without any adjustable parameter.
And this is as a result of theory. And you see because rates are depending exponentially on energetics,
the accuracy is excellent because we have very good agreement between the position of the maximum and the absolutely weights.
So that’s what theory can provide nowadays for the understanding of catalysis.
Let’s go back to heterogeneous catalysis.
Under steady state reaction rate and steady state conditions
we will have a constant flow and constant temperature, also constant reaction rate.
However this is not always the case.
There are cases where the situation looks very different.
And this is one example. Again this reaction on the platinum 110 surface.
At this point here we change the oxygen pressure from 2.0 to 2.7 times 10^-4 milibars.
The other 2 parameters are constant.
So it starts to increase slowly and then starts to oscillate.
Although we keep everything constant, it oscillates over long periods if you like.
And this is really surprising.
But if we look around in nature there are many, many processes which are oscillatory.
They are not constant in time.
And even in population dynamics there is an example.
The number of furs from hares and lynxes supplied to the economy as a function of time varies periodically.
And you see there is also a phase relation between the number of hares and lynxes.
Whenever the number of hares have a maximum the number of lynxes increases and then the hares decay
and with some phase lack also the populations of lynxes disappear.
There is a simple qualitative explanation.
If there are many hares available, there is enough food for the lynxes
so they can increase their population and if there have increases in population,
they eat more than they’re regrowing and as a consequence the number of hares is decaying
and with some delays the lynxes are starving.
This effect was modelled by mathematicians many years ago.
It’s a so called Lotka-Volterra model describing the rate of the concentration change of the hares and of the lynxes
in terms of 2 coupled non-linear ordinary differential equations.
Non-linear because both vary levels are here.
And if these equations are solved numerically under certain conditions, the result is indeed oscillatory as we have seen.
So that’s the secret behind these effects, the non-linear character of the underlying differential equations.
And that’s why this field is also called non-linear dynamics to describe the complexity of nature.
This is not the whole story.
If we have these populations varying periodically there must be some synchronisation between the different parts.
And this synchronisation has to take place through diffusion.
So we have to couple reaction and diffusion.
And as a consequence we will have the formation of structures, of patterns on the surface,
so called reaction diffusion patterns which are on a mesoscopic scale.
They are not on the atomic scale.
The length scale is determined by the coupling between reaction rate and diffusion.
And there is one very famous example in chemistry, the so called Belousov-Zhabotinsky reaction.
Very complicated reaction.
In the petri dish you show the formation of these patterns on a length scale of centimetres.
It’s just a coupling of reaction and diffusion.
With our reaction we can look at these concentration patterns
by looking at the distribution of CO and oxygen absorption with a technique
which is photo imaging electron microscopy which just probes the local dipole moments.
So we can really see where is the oxygen, where is the carbon monoxide on mesoscopic scale.
And this is the result. These are the target patterns on the surface, these elliptical structures.
Elliptical because this is the direction of the faster diffusion of that sub-particles.
Just reflects the influence of the diffusion.
And we can follow this pattern formation also in times they develop.
You see these elliptical waves expand and are superceded by a periodic variation of the background.
So there’s oscillation of the reaction rate reflected by this periodic variation of the coverage of oxygen and carbon monoxide.
Carbon monoxide is bright, oxygen is dark.
On a length scale of microns.
Again it’s the mesoscopic length scale which is responsible for this kind of pattern formation.
And it’s very sensitively affected by the external parameters.
So if you change this parameter a little bit you find other types of patterns.
And one very typical pattern is this kind of spiral which is shown here which we find very often.
And these spirals also propagate. Dark areas are oxygen covered. Bright areas are CO covered.
They propagate with the speed of several micrometres per second.
While in these dark areas, these are the catalytic active areas about 10 reaction steps per second,
per active site on the surface takes place. Completely different time scale.
And the different length scale of the spirals depends on the structure of the core of the spirals
which is usually determined by local defects.
So look. Also in the section of 2 spirals we often have this formation here.
Spirals are found in many, many different areas.
Galaxies, hurricanes but also in biology.
And this is a recent example from my colleague Petra Schwille with a cell
where these waves are propagating also again on the length scale of micrometres.
This kind of propagating waves are thought to be responsible for the division of a cell into 2 halves.
And spirals have been observed many, many years ago.
This is an example from the Stone Age, 3,000 years ago on the small island of Malta.
You can find this stone here with these typical protrusions as we have seen before.
Now back to our system. We know everything about the reaction.
We know also diffusion so we can really model and calculate the properties of this system.
And this has been done here, computer simulation starting with a random distribution of carbon monoxide and oxygen.
And as a function of time now we see under these conditions the formation of spirals.
As we have seen in the experiment.
If you change the conditions a little bit, these spirals will change their behaviour
and also in the computer simulations.
The spirals break up and form new structures which are disordered, which are called spiral turbulence.
This is what is called chaos.
2-dimensional space time chaos in this very simple surface reaction.
And we can also observe it in experiment.
This is an experimental observation of the break-up of spirals.
We get a chaotic regime with this very simple reaction.
And this is really surprising because we have a very simple reaction.
2 small molecules, a single crystal surface, well defined conditions
and nevertheless it shows this rather complex behaviour
as a consequence of the underlying non-linear partial differential equations describing the elementary steps.
And this is obviously the situation with our world as a whole.
There we find also very often chaos.
And we are very often at the edge of chaos which can be controlled by feedback considerations.
So we can consider this system also as a model for our world as it has been done many years ago by an artist.
And this artist was Vincent Van Gogh.
That’s his view of the world showing the spirals, the target patterns, the chaos.
And his conclusion was for the, ‘great things are not done by mere impulse but by a series of small things brought together’.
This is exactly what is done with our surface reaction.
Thank you very much.