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Reaction Mechanisms

Reaction mechanisms

Most chemical reactions occur by mechanisms that involve more than one step. As a result,

the rate law can not be directly deduced from the stoichiometry of a balanced chemical

equation. Only elementary reactions (those which occur in one step) can give the rate law.The rate law of an elementary process does follow from the coefficients of its balanced

equation.

Among the steps in a multi-step reaction, there is always one that acts like a bottleneck. There

is usually one step that is slowest compared to the rest. This slowest step limits the overall

reaction rate and is called the rate-determining step.

The rate-determining step determines the rate law for the overall reaction.

Example:

The reaction,

NO2(g) + CO(g) --> NO(g) + CO2(g)

is found to be second order with respect to [NO2] and zero order with respect to [CO]:

rate = k[NO2]2

The following mechanism has been proposed:

Step 1: NO2(g) + NO2(g) --> NO3(g) + NO(g) (slow)

Step 2: NO3(g) + CO(g) --> NO2(g) + CO2(g) (fast)

Overall: NO2(g) + CO(g) --> NO(g) + CO2(g)

Since Step1 is the slowest step, it determines the rate of the overall reaction. Thus, the above

mechanism agrees with the observed rate law.

Here is another example:

2NO(g) + O2(g) --> 2 NO2(g)

with the following observed rate law:

rate = kobs[NO]2[O2]

Of course, one possible mechanism is a one-step reaction between two molecules of nitrogen

oxide and a molecule of oxygen. This is a termolecular step, which is very unlikely since

three way collissions are highly improbable. An alternative mechanism is provided by the

following:

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Both steps are bimolecular, which are more likely to happen than a three-way collision. The

slow step determines the overall reaction rate:

which is sort of useless since it contains a term that does not correspond to either any of the

reactants or the product.

A rate law should only consist of concentrations of reactants and/or products, no

intermediates.

Since the intermediate, dinitrogen dioxide, reacts slowly with oxygen, the reverse reaction of

Step 1 is possible. What we are allowed to do then is to assume that Step 1 will reach a

dynamic equilibrium, which will lead to a constant concentration of dinitrogen dioxide, that

is, the forward and reverse rates ofStep 1 are equal:

which allows us to express the concentration of the intermediate dinitrogen dioxide in terms

of the starting material:

Therefore allowing us to express:

as:

which then relates the observed rate constant to the rate constants of the elementary steps as

follows:

Catalysts

A catalyst is a substance that changes the speed of a chemical reaction without itself

undergoing a permanent chemical change in the process. You have seen some examples of a

catalyst. Remember the production of O2 from KClO3, a catalyst was used in this reaction

(MnO2). A catalyst provides a different pathway for a reaction to occur and this different

route involves a lower Energy of activation.

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Homogeneous versus heterogeneous catalysis. A homogeneous catalyst is present in the same

phase as the reacting molecules. A heterogeneous catalyst is in a different phase.

Research question: Why is ozone depleted faster near the polar caps of the earth. Do ice

crystals catalyze the destruction of ozone?

Enzymes (biological catalysts)

Enzymes aid biological reactions by providing sites in which a reaction between two

biological molecules can easily occur. This process is therefore very much sensitive to the

shape of the enzyme, how accessible it is to the reactant species, etc. The shape of an enzyme

is very much dependent on intermolecular forces: hydrogen-bonding, dipole-dipole

interactions and London Forces. Thus, its shape can be drastically changed via changes in

temperature. Thus, unlike industrial catalysts, the activity of enzymes normally attain a

maximum at a given temperature (normally close to the temperature of the organism which is

using this enzyme).

TYPES OF CATALYSIS

This page looks at the the different types of catalyst

(heterogeneous and homogeneous) with examples of each kind,and explanations of how they work. You will also find a description

of one example of autocatalysis - a reaction which is catalysed by

one of its products.

Note: This page doesn'tdeal with the effect of catalysts on rates

of reaction. If you don't already know about that, you might like to

follow this link first. Return to this page via the BACK button on

your browser.

Types of catalytic reactions

Catalysts can be divided into two main types - heterogeneous and

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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.

You might wonder why phasediffers 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 labelled. 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

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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 solidcatalyst with the

reactants as either liquids or gases.

Note: It is important that you remember the difference between

the two terms heterogeneousand homogeneous.

heteroimplies different(as in heterosexual). Heterogeneous

catalysis has the catalyst in a different phase from the reactants.

homoimplies the same(as in homosexual). Homogeneous

catalysis has the catalyst in the same phase as the reactants.

How the heterogeneous catalyst works (in general terms)

Most examples of heterogeneous catalysis go through the same

stages:

One or more of the reactants are adsorbedon 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. Be careful!

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.

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.

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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.

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.

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.

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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.

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.

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.

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Note: Several metals, including nickel, have the ability to absorb

hydrogen into their structure as well as adsorb it on to the surface.

In these cases, the hydrogen molecules are also converted into

atoms which can diffuse through the metal structure.

This happens with nickel if the hydrogen is under high pressures,

but I haven't been able to find any information about whether it is

also absorbed under the lower pressures usually used for these

hydrogenation reactions. I have therefore stuck with the usualexplanation in terms of adsorption.

If anyone has any firm information about this, could they contact

me via the address on the about this site page.

Catalytic converters

Catalytic converters change poisonous molecules like carbon

monoxide and various nitrogen oxides in car exhausts into more

harmless molecules like carbon dioxide and nitrogen. They use

expensive metals like platinum, palladium and rhodium as the

heterogeneous catalyst.

The metals are deposited as thin layers onto a ceramic

honeycomb. This maximises the surface area and keeps the

amount of metal used to a minimum.

Taking the reaction between carbon monoxide and nitrogen

monoxide as typical:

In the same sort of way as the previous example, the carbon

monoxide and nitrogen monoxide will be adsorbed on the surface

of the catalyst, where they react. The carbon dioxide and nitrogen

are then desorbed.

The use of vanadium(V) oxide in the Contact Process

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During the Contact Process for manufacturing sulphuric acid,

sulphur dioxide has to be converted into sulphur trioxide. This is

done by passing sulphur dioxide and oxygen over a solid

vanadium(V) oxide catalyst.

Note: The equation is written with the half in it to make the

explanation below tidier. You may well be familiar with the

equation written as twice that shown, but the present version is

perfectly acceptable. It is also shown as a one-way rather than a

reversible reaction to avoid complicating things.

This example is slightly different from the previous ones because

the gases actually react with the surface of the catalyst, temporarily

changing it. It is a good example of the ability of transition metals

and their compounds to act as catalysts because of their ability to

change their oxidation state.

Note: If you aren't sure aboutoxidation states, it might be useful

to follow this link before you go on. Use the BACK button on your

browser to return to this page.

The sulphur dioxide is oxidised to sulphur trioxide by the

vanadium(V) oxide. In the process, the vanadium(V) oxide is

reduced to vanadium(IV) oxide.

The vanadium(IV) oxide is then re-oxidised by the oxygen.

This is a good example of the way that a catalyst can be changed

during the course of a reaction. At the end of the reaction, though,

it will be chemically the same as it started.

Note: If you want more detail about the Contact Process, you will

find a full description of the conditions used and the reasons for

them by following this link.

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. The examples contain one of each of these . . .

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Examples of homogeneous catalysis

The reaction between persulphate ions and iodide ions

This is a solution reaction that you may well only meet in thecontext of catalysis, but it is a lovely example!

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.

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 toiron(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

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iron(II) ions? The reactions simply happen in a different order.

The destruction of atmospheric ozone

This is a good example of homogeneous catalysis where

everything is present as a gas.

Ozone, O3, is constantly being formed and broken up again in the

high atmosphere by the action of ultraviolet light. Ordinary oxygen

molecules absorb ultraviolet light and break into individual oxygen

atoms. These have unpaired electrons, and are known as free

radicals. They are very reactive.

The oxygen radicals can then combine with ordinary oxygen

molecules to make ozone.

Ozone can also be split up again into ordinary oxygen and anoxygen radical by absorbing ultraviolet light.

This formation and breaking up of ozone is going on all the time.

Taken together, these reactions stop a lot of harmful ultraviolet

radiation penetrating the atmosphere to reach the surface of the

Earth.

The catalytic reaction we are interested in destroys the ozone andso stops it absorbing UV in this way.

Chlorofluorocarbons (CFCs) like CF2Cl2, for example, were used

extensively in aerosols and as refrigerants. Their slow breakdown

in the atmosphere produces chlorine atoms - chlorine free radicals.

These catalyse the destruction of the ozone.

This happens in two stages. In the first, the ozone is broken up and

a new free radical is produced.

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The chlorine radical catalyst is regenerated by a second reaction.

This can happen in two ways depending on whether the ClO

radical hits an ozone molecule or an oxygen radical.

If it hits an oxygen radical (produced from one of the reactions

we've looked at previously):

Or if it hits an ozone molecule:

Because the chlorine radical keeps on being regenerated, each

one can destroy thousands of ozone molecules.

Note: If you are a UK A level student it is probable that you will

only want one of these last two equations depending on what your

syllabus says. Unfortunately, at the time of writing, different A

level syllabuses were quoting different final equations. You need

to check your current syllabus. If you don't have a copy of your

syllabus, follow this link to find out how to get one.

Autocatalysis

The oxidation of ethanedioic acid by manganate(VII) ions

In autocatalysis, the reaction is catalysed by one of its products.

One of the simplest examples of this is in the oxidation of a solutionof ethanedioic acid (oxalic acid) by an acidified solution of

potassium manganate(VII) (potassium permanganate).

The reaction is very slow at room temperature. It is used as a

titration to find the concentration of potassium manganate(VII)

solution and is usually carried out at a temperature of about 60C.

Even so, it is quite slow to start with.

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The reaction is catalysed by manganese(II) ions. There obviously

aren't any of those present before the reaction starts, and so it

starts off extremely slowly at room temperature. However, if you

look at the equation, you will find manganese(II) ions amongst the

products. More and more catalyst is produced as the reaction

proceeds and so the reaction speeds up.

You can measure this effect by plotting the concentration of one of

the reactants as time goes on. You get a graph quite unlike the

normal rate curve for a reaction.

Most reactions give a rate curve which looks like this:

Concentrations are high at the beginning and so the reaction is fast

- shown by a rapid fall in the reactant concentration. As things get

used up, the reaction slows down and eventually stops as one or

more of the reactants are completely used up.

An example of autocatalysis gives a curve like this:

You can see the slow (uncatalysed) reaction at the beginning. Ascatalyst begins to be formed in the mixture, the reaction speeds up

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- getting faster and faster as more and more catalyst is formed.

Eventually, of course, the rate falls again as things get used up.

Warning!

Don't assume that a rate curve which looks like this necessarily

shows an example of autocatalysis. There are other effects which

might produce a similar graph.

For example, if the reaction involved a solid reacting with a liquid,

there might be some sort of surface coating on the solid which the

liquid has to penetrate before the expected reaction can happen.

A more common possibility is that you have a strongly exothermic

reaction and aren't controlling the temperature properly. The heatevolved during the reaction speeds the reaction up.