what is the “rate" of a reaction? the rate of a reaction is the speed of the reaction. it is...
TRANSCRIPT
What is the “rate" of a reaction?
• The rate of a reaction is the speed of the reaction. It is not “how much” of a product is made, but instead “how quickly” a reaction takes place.
How can we measure the rate?• If we consider a reaction
zinc + hydrochloric acid —> zinc chloride + hydrogen
• then there are two possible ways of measuring the rate: 1) measure how quickly one of the products (e.g. the
hydrogen) is made2) measure how quickly one of the reactants (e.g. the zinc) disappears
• So we could, for example, measure the volume (in ml.) of hydrogen made every 10 seconds or the loss in mass (of the zinc and hydrochloric acid as they change into hydrogen gas escaping from a beaker) every 10 seconds.
Measuring a rate of reaction
• There are several simple ways of measuring a reaction rate.
• For example, if a gas was being given off during a reaction, you could take some measurements and work out the volume being given off per second at any particular time during the reaction.
• A rate of 2 cm3 s-1 is obviously twice as fast as one of 1 cm3 s-1.
Note: Read cm3 s-1 as "cubic centimetres per second".
Measuring a rate of reaction
• However, for this more formal and mathematical look at rates of reaction, the rate is usually measured by looking at how fast the concentration of one of the reactants is falling at any one time.
• For example, suppose you had a reaction between two substances A and B. Assume that at least one of them is in a form where it is sensible to measure its concentration - for example, in solution or as a gas.
• For this reaction you could measure the rate of the reaction by finding out how fast the concentration of, say, A was falling per second.
Measuring a rate of reaction• You might, for example, find that at
the beginning of the reaction, its concentration was falling at a rate of 0.0040 mol dm-3 s-1.
• This means that every second the concentration of A was falling by 0.0040 moles per cubic decimetre. This rate will decrease during the reaction as A gets used up.
Note: Read mol dm-3 s-1 as "moles per cubic decimetre (or litre) per second".
Reactions involving collisions between two species
Two species (molecule, ion, or free radical) can only react together if they come into contact with each other.
They first have to collide, and then they may react.
Why "may react"?
It isn't enough for the two species to collide
they have to collide the right way around, and
they have to collide with enough energy for bonds to break.
ineffective collisions Reactions usually require
collisions between reactant molecules or atoms. The formation of bonds requires atoms to come close to one another. New bonds can form only if the atoms are close enough together to share electrons. Some collisions are not successful. These are called ineffective collisions. The particles simply hit and then rebound. This animation illustrates what happens in an ineffective collision.
effective collisions Collisions that lead to
products are called effective collisions. An effective collision must happen with a great enough speed, energy and force to break bonds in the colliding molecules.
The animation illustrates an effective collision between two diatomic molecules. The two product molecules formed fly outwards.
Reactions involving collisions more than two species
All three (or more) particles
would have to arrive at exactly the same point in space at the same time, with everything lined up exactly right, and having enough energy to react.
That's not likely to happen very often
The orientation of collision
As a result of the collision between the two molecules, the double bond between the two carbons is converted into a single bond. A hydrogen atom gets attached to one of the carbons and a chlorine atom to the other.
The orientation of collision
Of the collisions shown in the diagram, only collision 1 may possibly lead on to a reaction.
The orientation of collision
You may wonder why collision 2 won't work as well.
The double bond has a high concentration of negative charge around it due to the electrons
in the bonds.
The approaching chlorine atom is also slightly negative because it is more
electronegative than hydrogen.
The repulsion simply causes the molecules to bounce off each other.
Proper Orientation of the colliding molecules
The energy of the collision
Even if the species are orientated properly, you still won't get a reaction
unless the particles collide with a certain minimum energy
called the activation energy of the reaction.
Activation EnergyIt is the minimum energy required before a reaction can occur. You can show this on an energy profile for a simple over-all exothermic reaction, the energy profile looks like this:
Activation Energy
Activation Energy
You can think if the particles collide with less energy than the activation energy, nothing important happens. They bounce apart.
Only those collisions which have energies equal to or greater than the activation energy result in a reaction.
Activation EnergyAny chemical reaction results in the breaking of some bonds (needing energy) and the making of new ones (releasing energy).
Obviously some bonds have to be broken before new ones can be made.
Activation energy is involved in breaking some of the original bonds.Where collisions are relatively gentle, there isn't enough energy available to start the bond-breaking process, and so the particles don't react.
Activation Energy
The Maxwell-Boltzmann Distribution
Because of the key role of activation energy
in deciding whether a collision will result in
a reaction, it would obviously be useful to
know what sort of proportion of the particles
present have high enough energies to react
when they collide.
The Maxwell-Boltzmann Distribution
In any system, the particles present will have a
very wide range of energies. For gases, this
can be shown on a graph called the Maxwell-
Boltzmann Distribution which is a plot of the
number of particles having each particular
energy.
The Maxwell-Boltzmann Distribution
The area under the curve is a measure of the total number of particles present. The graph only applies to gases, but the conclusions that we can draw from it can also be applied to reactions involving liquids.
The Maxwell-Boltzmann Distribution and activation energy
Remember that for a reaction to happen, particles must collide with energies equal to or greater than the activation energy for the reaction:
Notice that the large majority of the particles don't have enough energy to react when they collide. To enable them to react we either have to change the shape of the curve, or move the activation energy further to the left.
You can change the shape of the curve by changing the temperature of the reaction. You can change the position of the activation energy by adding a catalyst to the reaction.
THE EFFECT OF SURFACE AREA ON REACTION RATES
This effect applies to reactions:
involving a solid and a gas, or
a solid and a liquid.
It includes cases where the solid is acting as a catalyst.
THE EFFECT OF SURFACE AREA ON REACTION RATES
The more finely divided the solid is, the faster the reaction happens.
A powdered solid will normally produce a faster reaction than if the same mass is present as a single lump.
The powdered solid has a greater surface area than the single lump.
Why normally? What exceptions
can there be?
Imagine a case of a very fine powder reacting with a gas. If the powder was in one big heap, the gas may not be able to penetrate it.
That means that its effective surface area is much the same as (or even less than) it would be if it were present in a single lump.
A small heap of fine magnesium powder tends to burn rather more slowly than a strip of magnesium ribbon, for example.
Calcium carbonate and hydrochloric acid
In the lab, powdered calcium carbonate reacts much faster with dilute hydrochloric acid than if the same mass was present as lumps of marble or limestone.
The catalytic decomposition of hydrogen
peroxide
Solid manganese(IV) oxide is often used as the catalyst. Oxygen is given off much faster if the catalyst is present as a powder than as the same mass of granules.
The explanation:You are only going to get a reaction if the particles in the gas or liquid collide with the particles in the solid.
Increasing the surface area of the solid increases the chances of collision taking place.
Imagine a reaction between magnesium metal and a dilute acid like hydrochloric acid. The reaction involves collision between magnesium atoms and hydrogen ions.
Increasing the number of collisions per second increases the rate of reaction.
Catalytic convertersCatalytic converters use metals like platinum, palladium and rhodium to convert poisonous compounds in vehicle exhausts into less harmful things. For example, a reaction which removes both carbon monoxide and an oxide of nitrogen is:
Because the exhaust gases are only in contact with the catalyst for a very short time, the reactions have to be very fast. The extremely expensive metals used as the catalyst are coated as a very thin layer onto a ceramic honeycomb structure to maximize the surface area.
The effect of particle size• Solids with a smaller
particle size (e.g. powders or small chips) react more quickly than solids with a larger particle size (e.g large chips). Here is why:
• Look at this diagram
The perimeter of the large chip is 12 units. The acid particles can only collide with the edge of the chip. However, if we break up the large chip into 9 smaller chips:
The effect of particle size
• However, if we break up the large chip into 9 smaller chips:
• then the perimeter around each chip is 4 units, but there are 9 of them so the total perimeter is 4 x 9 = 36 units. Notice how the acid in the second diagram can reach what used to be the centre of the large chip.
Reducing the size of particles increases the rate of a reaction because it increases the surface area available for collisions to take place. This increases the number of collisions. It has no effect on the energy of the particles.
THE EFFECT OF CONCENTRATION ON REACTION RATES
For many reactions involving liquids or gases, increasing the concentration of the reactants increases the rate of reaction.
In a few cases, increasing the concentration of one of the reactants may have little noticeable effect of the rate.
Don't assume that if you double the concentration of one of the reactants that you will double the rate of the reaction. It may happen like that, but the relationship may well be more complicated.
The mathematical relationship between concentration and rate of reaction is related with the orders of reaction.
Some examples to the concentration effect:
Zinc and hydrochloric acid
In the lab, zinc granules react fairly slowly with dilute hydrochloric acid, but much faster if the acid is concentrated.
The catalytic decomposition of hydrogen peroxide
Solid manganese(IV) oxide is often used as a catalyst in this reaction. Oxygen is given off much faster if the hydrogen peroxide is concentrated than if it is dilute.
The reaction between sodium thiosulphate solution and hydrochloric acid
When a dilute acid is added to sodium thiosulphate solution, a pale yellow precipitate of sulphur is formed.
As the sodium thiosulphate solution is diluted more and more, the precipitate takes longer and longer to form.
The explanation of the concentration effect:Collisions involving two particles
In order for any reaction to happen, those particles must first collide.
This is true whether both particles are in solution, or whether one is in solution and the other a solid.
If the concentration is higher, the chances of collision are greater.
The explanation of the concentration effect: Reactions involving only one particle
If a reaction only involves a single particle splitting up in some way,
then the number of collisions is irrelevant.
what matters now is how many of the particles have enough energy to react at any one time.
The explanation of the concentration effect:
Reactions involving only one particle
Suppose that at any one time 1 in a million particles have enough energy to equal or exceed the activation energy. If you had 100 million particles, 100 of them would react. If you had 200 million particles in the same volume, 200 of them would now react. The rate of reaction has doubled by doubling the concentration.
Cases where changing the concentration doesn't affect the rate of the reaction
Suppose you are using a small amount of a solid catalyst in a reaction, and a high enough
concentration of reactant in solution so that the catalyst surface was totally cluttered up with
reacting particles.
Increasing the concentration of the solution even more can't have any effect because the catalyst is
already working at its maximum capacity.
In certain multi-step reactionsSuppose you have a reaction which happens in a series of small steps. These steps are likely to have widely different
rates - some fast, some slow.
For example, suppose two reactants A and B react together in these two stages:
The overall rate of the reaction is going to be governed by how fast A splits up to make X and Y. This is described as
the rate determining step of the reaction.
In certain multi-step reactionsIf you increase the concentration
of A, you will increase the chances of this step happening
for reasons we've looked at above.
If you increase the concentration of B, that will undoubtedly speed
up the second step, but that makes hardly any difference to
the overall rate.
You can picture the second step as happening so fast already that as soon as any X is formed, it is immediately pounced on by B. That second reaction is already
"waiting around" for the first one to happen.
In certain multi-step reactions• The overall rate of reaction isn't
entirely independent of the concentration of B. If you lowered its concentration enough, you will eventually reduce the rate of the second reaction to the point where it is similar to the rate of the first. Both concentrations will matter if the concentration of B is low enough.
• However, for ordinary concentrations, you can say that (to a good approximation) the overall rate of reaction is unaffected by the concentration of B.
Orders of reaction• Orders of reaction are always found by
doing experiments. You can't deduce anything about the order of a reaction just by looking at the equation for the reaction.
• So let's suppose that you have done some experiments to find out what happens to the rate of a reaction as the concentration of one of the reactants, A, changes. Some of the simple things that you might find are:
Orders of reaction• One possibility: The rate of reaction is
proportional to the concentration of A• That means that if you double the
concentration of A, the rate doubles as well. If you increase the concentration of A by a factor of 4, the rate goes up 4 times as well.
• You can express this using symbols as:
Writing a formula in square brackets is a standard way of showing a concentration measured in moles per cubic decimetre (litre).
Orders of reaction
• You can also write this by getting rid of the proportionality sign and introducing a constant, k.
Orders of reaction
• Another possibility: The rate of reaction is proportional to the square of the concentration of A
• This means that if you doubled the concentration of A, the rate would go up 4 times (22). If you tripled the concentration of A, the rate would increase 9 times (32). In symbol terms:
Orders of reaction• Generalising this• By doing experiments involving a reaction between
A and B, you would find that the rate of the reaction was related to the concentrations of A and B in this way:
• This is called the rate equation for the reaction.• The concentrations of A and B have to be raised to
some power to show how they affect the rate of the reaction. These powers are called the orders of reaction with respect to A and B.
Orders of reaction
• If the order of reaction with respect to A is 0 (zero), this means that the concentration of A doesn't affect the rate of reaction. Mathematically, any number raised to the power of zero (x0) is equal to 1. That means that that particular term disappears from the rate equation.
overall order of the reaction
• The overall order of the reaction is found by adding up the individual orders. For example, if the reaction is first order with respect to both A and B (a = 1 and b = 1),
• R= k[A] [B]
• the overall order is 2. We call this an overall second order reaction.
Some examples
• Each of these examples involves a reaction between A and B, and each rate equation comes from doing some experiments to find out how the concentrations of A and B affect the rate of reaction.
Example 1
• Rate = k[A] [B]
•In this case, the order of reaction with respect to both A and B is 1. The overall order of reaction is 2 - found by adding up the individual orders.
Note: Where the order is 1 with respect to one of the reactants, the "1" isn't written into the equation. [A] means [A]1.
Example 2
• This reaction is zero order with respect to A because the concentration of A doesn't affect the rate of the reaction.
• The order with respect to B is 2 - it's a second order reaction with respect to B.
• The reaction is also second order overall (because 0 + 2 = 2).
•Rate = k [B]2
Example 3
• This reaction is first order with respect to A and zero order with respect to B, because the concentration of B doesn't affect the rate of the reaction.
• The reaction is first order overall (because 1 + 0 = 1).
Rate = k[A]
Orders of reaction
• What if you have some other number of reactants?
•It doesn't matter how many reactants there are.
• The concentration of each reactant will occur in the rate equation, raised to some power.
• Those powers are the individual orders of reaction.
•The overall order of the reaction is found by adding them all up.
THE EFFECT OF PRESSURE ON REACTION RATES
Up to know we dealt with concentration effect in the reactions that occure in aqueous solutions.
Changing the concentration of a gas is achieved by changing its pressure.
Increasing the pressure on a reaction involving reacting gases increases the rate of reaction. Changing the pressure on a reaction which involves only solids or liquids has no effect on the rate.
An example to pressure effect on reaction rate
• In the manufacture of ammonia by the Haber Process, the rate of reaction between the hydrogen and the nitrogen is increased by the use of very high pressures.
• In fact, the main reason for using high pressures is to improve the percentage of ammonia in the equilibrium mixture, but there is a useful effect on rate of reaction as well.
The explanation of the relationship between pressure and concentration• Increasing the pressure of a gas is exactly
the same as increasing its concentration. If you have a given mass of gas, the way you increase its pressure is to squeeze it into a smaller volume. If you have the same mass in a smaller volume, then its concentration is higher.
• You can also show this relationship mathematically if you have come across the ideal gas equation:
The explanation of the relationship between pressure and concentration
Rearranging PV = nRT gives:
Because "RT" is constant as long as the temperature is constant, this shows that the pressure is directly proportional to the concentration. If you double one, you will also double the other.
The effect of increasing the pressure on the rate of reaction
Collisions involving two particles
• The same argument applies whether the reaction involves collision between two different particles or two of the same particle.
• In order for any reaction to happen, those particles must first collide. This is true whether both particles are in the gas state, or whether one is a gas and the other a solid. If the pressure is higher, the chances of collision are greater.
The effect of increasing the pressure on the rate of reaction
Collisions involving two particles
The effect of increasing the pressure on the rate of reaction
• Reactions involving only one particle
• If a reaction only involves a single particle splitting up in some way, then the number of collisions is irrelevant. What matters now is how many of the particles have enough energy to react at any one time.
THE EFFECT OF TEMPERATURE ON REACTION RATES
• As you increase the temperature the rate of reaction increases. As a rough approximation, for many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature.
• You have to be careful not to take this too literally. It doesn't apply to all reactions. Even where it is approximately true, it may be that the rate doubles every 9°C or 11°C or whatever. The number of degrees needed to double the rate will also change gradually as the temperature increases.
THE EFFECT OF TEMPERATURE ON REACTION RATES
• Some reactions are virtually instantaneous - for example, a precipitation reaction involving the coming together of ions in solution to make an insoluble solid, or the reaction between hydrogen ions from an acid and hydroxide ions from an alkali in solution. So heating one of these won't make any noticeable difference to the rate of the reaction.
• Almost any other reaction you care to name will happen faster if you heat it - either in the lab, or in industry.
The explanation of the temperature effectIncreasing the collision frequency
• Particles can only react when they collide. If you heat a substance, the particles move faster and so collide more frequently. That will speed up the rate of reaction.
• That seems a fairly straightforward explanation until you look at the numbers!
• It turns out that the frequency of two-particle collisions in gases is proportional to the square root of the kelvin temperature.
The explanation of the temperature effectIncreasing the collision frequency
• That's an increase of 1.7% for a 10° rise. The rate of reaction will probably have doubled for that increase in temperature - in other words, an increase of about 100%. The effect of increasing collision frequency on the rate of the reaction is very minor. The important effect is quite different . . .
•If you increase the temperature from 293 K to 303 K (20°C to 30°C), you will increase the collision frequency by a factor of:
Temperature effects on rates and activation energy diagram
• This illustration shows what happens to an exothermic reaction when the temperature is changed.
• The dotted blue curve shows the energy for a reaction mixture that is heated. The reactants are "part way" up the energy barrier because they are "hot".
• The dotted magenta curve shows what cooling does to the reactant energy. The energy goes down and the reaction happens with more difficulty.
Temperature effects on rates and activation energy diagram
NOTE: The energies of reactants and products have changed. They both have different energies because they were either heated or cooled. The heat of reaction is the slightly different. The relative amounts of reactants and products are slightly different because of the temperature changes.
The key importance of activation energy• Collisions only result in a
reaction if the particles collide with enough energy to get the reaction started. This minimum energy required is called the activation energy for the reaction.
• Only those particles represented by the area to the right of the activation energy will react when they collide. The great majority don't have enough energy, and will simply bounce apart.
The key importance of activation energy
• To speed up the reaction, you need to increase the number of the very energetic particles - those with energies equal to or greater than the activation energy. Increasing the temperature has exactly that effect - it changes the shape of the graph.
The key importance of activation energy• If you now mark the
position of the activation energy, you can see that although the curve hasn't moved very much overall, there has been such a large increase in the number of the very energetic particles that many more now collide with enough energy to react.
The key importance of activation energy
• Remember that the area under a curve gives a count of the number of particles. On the last diagram, the area under the higher temperature curve to the right of the activation energy looks to have at least doubled - therefore at least doubling the rate of the reaction.
Summary of the temperature effect
Increasing the temperature increases reaction rates because of the disproportionately large increase in the number of high energy collisions. It is only these collisions (possessing at least the activation energy for the reaction) which result in a reaction.
THE EFFECT OF CATALYSTS ON REACTION RATES
A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the reaction. When the reaction has finished, you would have exactly the same mass of catalyst as you had at the beginning.
Some examples of catalysts
reaction catalyst Decomposition of hydrogen peroxide
manganese(IV) oxide, MnO2
Nitration of benzene concentrated sulphuric acid
Manufacture of ammonia by the Haber Process
iron
Conversion of SO2 into SO3 during the Contact Process to make sulphuric acid
vanadium(V) oxide, V2O5
Hydrogenation of a C=C double bond nickel
Catalysts and activation energy• To increase the rate of a
reaction you need to increase the number of successful collisions. One possible way of doing this is to provide an alternative way for the reaction to happen which has a lower activation energy.
• In other words, to move the activation energy on the graph like this:
Catalysts and activation energy
• Adding a catalyst has exactly this effect on activation energy. A catalyst provides an alternative route for the reaction. That alternative route has a lower activation energy. Showing this on an energy profile:
Catalysts and activation energy
Catalysts and activation energy
• Be very careful if you are asked about this in an exam. The correct form of words is
• "A catalyst provides an alternative route for the reaction with a lower activation energy."
• It does not "lower the activation energy of the reaction". There is a subtle difference between the two statements.
Catalysts and activation energy with a simple analogy.
• Suppose you have a mountain between two valleys so that the only way for people to get from one valley to the other is over the mountain. Only the most active people will manage to get from one valley to the other.
• Now suppose a tunnel is cut through the mountain. Many more people will now manage to get from one valley to the other by this easier route. You could say that the tunnel route has a lower activation energy than going over the mountain.
• But you haven't lowered the mountain! The tunnel has provided an alternative route but hasn't lowered the original one. The original mountain is still there, and some people will still choose to climb it.
Catalysts and activation energy with a simple analogy.
In the chemistry case, if particles collide with enough energy they can still react in exactly the same way as if the catalyst wasn't there. It is simply that the majority of particles will react via the easier catalysed route.
TYPES OF CATALYSIS
• heterogeneous and homogeneous
• autocatalysis - a reaction which is catalysed by one of its products.
Types of catalytic reactions
• Catalysts can be divided into two main types
• heterogeneous and homogeneous.
• In a heterogeneous reaction, the catalyst is in a different phase from the reactants.
• In a homogeneous reaction, the catalyst is in the same phase as the reactants.
What is a phase?
• If you look at a mixture and can see a boundary between two of the components, those substances are in different phases.
• A mixture containing a solid and a liquid consists of two phases.
• A mixture of various chemicals in a single solution consists of only one phase, because you can't see any boundary between them.
What is a phase?
What is a phase?• You might wonder why
phase differs from the term physical state (solid, liquid or gas). It includes solids, liquids and gases, but is actually a bit more general.
• It can also apply to two liquids (oil and water, for example) which don't dissolve in each other. You could see the boundary between the two liquids.
If you want to be fussy about things, the diagrams actually show more phases than are labeled. Each, for example, also has the glass beaker as a solid phase. All probably have a gas above the liquid - that's another phase. We don't count these extra phases because they aren't a part of the reaction.
Heterogeneous catalysis
• This involves the use of a catalyst in a different phase from the reactants.
• Typical examples involve a solid catalyst with the reactants as either liquids or gases.
Note: It is important that you remember the difference between the two terms heterogeneous and homogeneous.
hetero implies different (as in heterosexual). Heterogeneous catalysis has the catalyst in a different phase from the reactants.
homo implies the same (as in homosexual). Homogeneous catalysis
has the catalyst in the same phase as the reactants.
How the heterogeneous catalyst works • One or more of the reactants are adsorbed on to
the surface of the catalyst at active sites.• Adsorption is where something sticks to a
surface. It isn't the same as absorption where one substance is taken up within the structure of another.
• An active site is a part of the surface which is particularly good at adsorbing things and helping them to react.
• There is some sort of interaction between the surface of the catalyst and the reactant molecules which makes them more reactive.
• This might involve an actual reaction with the surface, or some weakening of the bonds in the attached molecules.
How the heterogeneous catalyst works• The reaction happens.• At this stage, both of the reactant molecules
might be attached to the surface, or one might be attached and hit by the other one moving freely in the gas or liquid.
• The product molecules are desorbed.• Desorption simply means that the product
molecules break away. This leaves the active site available for a new set of molecules to attach to and react.
How the heterogeneous catalyst works
• A good catalyst needs to adsorb the reactant molecules strongly enough for them to react, but not so strongly that the product molecules stick more or less permanently to the surface.
• Silver, for example, isn't a good catalyst because it doesn't form strong enough attachments with reactant molecules. Tungsten, on the other hand, isn't a good catalyst because it adsorbs too strongly.
• Metals like platinum and nickel make good catalysts because they adsorb strongly enough to hold and activate the reactants, but not so strongly that the products can't break away.
Examples of heterogeneous catalysis
• The hydrogenation of a carbon-carbon double bond
• The simplest example of this is the reaction between ethene and hydrogen in the presence of a nickel catalyst.
• In practice, this is a pointless reaction, because you are converting the extremely useful ethene into the relatively useless ethane. However, the same reaction will happen with any compound containing a carbon-carbon double bond.
• One important industrial use is in the hydrogenation of vegetable oils to make margarine, which also involves reacting a carbon-carbon double bond in the vegetable oil with hydrogen in the presence of a nickel catalyst
The hydrogenation of a carbon-carbon double bond
• Ethene molecules are adsorbed on the surface of the nickel.
• The double bond between the carbon atoms breaks and the electrons are used to bond it to the nickel surface.
The hydrogenation of a carbon-carbon double bond
• Hydrogen molecules are also adsorbed on to the surface of the nickel.
• When this happens, the hydrogen molecules are broken into atoms.
• These can move around on the surface of the nickel.
The hydrogenation of a carbon-carbon double bond
• If a hydrogen atom diffuses close to one of the bonded carbons, the bond between the carbon and the nickel is replaced by one between the carbon and hydrogen.
•That end of the original ethene now breaks free of the surface, and •eventually the same thing will happen at the other end.
The hydrogenation of a carbon-carbon double bond
• As before, one of the hydrogen atoms forms a bond with the carbon, and that end also breaks free.
• There is now space on the surface of the nickel for new reactant molecules to go through the whole process again.
Homogeneous catalysis
• This has the catalyst in the same phase as the reactants.
• Typically everything will be present as a gas or contained in a single liquid phase.
Examples of homogeneous catalysis
• The reaction between persulphate ions and iodide ions
• Persulphate ions (peroxodisulphate ions), S2O82-, are very
powerful oxidising agents. Iodide ions are very easily oxidised to iodine. And yet the reaction between them in solution in water is very slow.
• If you look at the equation, it is easy to see why that is:
• The reaction needs a collision between two negative ions. Repulsion is going to get seriously in the way of that!
• The catalysed reaction avoids that problem completely. The catalyst can be either iron(II) or iron(III) ions which are added to the same solution. This is another good example of the use of transition metal compounds as catalysts because of their ability to change oxidation state.
The reaction between persulphate ions and iodide ions
• For the sake of argument, we'll take the catalyst to be iron(II) ions. As you will see shortly, it doesn't actually matter whether you use iron(II) or iron(III) ions.
• The persulphate ions oxidise the iron(II) ions to iron(III) ions. In the process the persulphate ions are reduced to sulphate ions.
• The iron(III) ions are strong enough oxidising agents to oxidise iodide ions to iodine. In the process, they are reduced back to iron(II) ions again.
• Both of these individual stages in the overall reaction involve collision between positive and negative ions. This will be much more likely to be successful than collision between two negative ions in the uncatalysed reaction.
• What happens if you use iron(III) ions as the catalyst instead of iron(II) ions? The reactions simply happen in a different order.