In this lesson, we will learn:
- The difference between homogeneous and heterogeneous catalysts.
- To understand how transition metals can act as heterogeneous catalysts in the contact process and the internal combustion engine.
- To understand how transition metal ions can act as homogeneous catalysts in various redox reactions.
- Transition metals make good catalysts because they often have many flexible oxidation states. This means they’re often able to both reduce and oxidise other species, and chemical changes to themselves can be reversible.
There are two major categories of catalyst that relate to what state it is in, compared to the reaction:
- A homogeneous catalyst is a catalyst in the same phase as the reactants of the process. This is more common in aqueous reactions, where all species involved are in solution such as some redox reactions.
- A heterogeneous catalyst is a catalyst in a different phase to the reactants of the process being catalysed. This is common for reactions in the gas phase where a solid catalyst is used, for example.
The difference between these two types affects how the catalyst works:
- Homogeneous catalysts often take part in the reaction; they reach an intermediate and then come back to their original state, carrying the reactants through to products in doing so.
Think of a roundabout; reactants enter it and exit at a different stage as newly-formed products. The transition metal’s ability to accept and donate electrons and take up its different oxidation states is key.
- Heterogeneous catalysts usually just provide a surface for higher energy particles to interact with, an interaction called adsorption. This involves temporary interactions at the metal surface – absorbing would be getting inside the metal, adsorbing is only at the surface.
The most common examples are solid metal catalysts for gas-phase reactions; this is how catalytic converters work in car engines. We’ll see this in detail below.
- One of the most important commercial uses of transition metal catalysis is in the contact process, which uses a vanadium (V) catalyst to help produce sulfuric acid.
Vanadium’s flexible oxidation states are used to do this; V2O5, or vanadium (V) oxide oxidises sulfur dioxide to sulfur trioxide, which is an early step in making sulfuric acid.
SO2 + O2
(V2O5 cat.) → SO3
As we saw in Reactions of vanadium and other transition metals, vanadium can be reduced to vanadium (IV), another common oxidation state. This is what happens in the reaction with sulfur dioxide:
SO2 + V2O5 → SO3 + V2O4
Again, as we saw in the last lesson, vanadium (IV) is easily oxidised back to vanadium (V) by oxygen gas. We’re doing a gas phase reaction so there will be oxygen present:
V2O4 + O2 → V2O5
This vanadium (V) is capable of further catalyzing SO2 in the first equation above.
This use of vanadium in a very important industrial process is classic transition metal chemistry – vanadium can reach multiple oxidation states without harsh chemical conditions being required.
- Another important example of transition metal catalysis is in internal combustion engines as a catalytic converter.
The extremely high temperature and pressure generated in engines mean some high energy, toxic compounds get made: carbon monoxide (CO) and nitrogen monoxide (NO). Metals like platinum and rhodium can convert CO and NO fumes to less harmful CO2 and N2.
This is done using a solid metal coated surface in the engine; as a solid catalyst with gas reactants, this is heterogeneous catalysis.
- The catalytic converter is a thin surface coating on a filter-like honeycomb structure in a car engine. The NO and CO particles collide with its surface where adsorption takes place, weakening the bonds in the compounds.
- The CO and NO gases react on it to become CO2 and N2.
2CO + 2NO (Pd/Rh cat.) → 2CO2 + N2
- These products then desorb from the surface and exit the engine through the exhaust pipe.
- Remember adsorption is only a surface interaction. We only need a thin coating layer of these expensive metals. This ceramic honeycomb structure with the metal coating therefore gives maximum surface area using as little of the precious metals as possible.
- One of the major issues with surface heterogeneous catalysts is catalyst poisoning – the surface becomes contaminated or permanently blocked (maybe a very strong adsorbate) by substances that prevent the surface from catalyzing any reaction. This, obviously, reduces the performance of the catalyst. Think of the lint catcher in your dryer – the more lint it has already caught, the less space it has to catch more. You can remove that, but honeycomb catalysts will need replacing after a while.
- The best-known examples of homogeneous catalysis with transition metals is in redox reactions.
Just like with vanadium (V) oxide in the contact process, the ability to hold multiple oxidation states is very useful for this. Take the example of persulfate ions reacting with iodide to produce iodine.
S2O82- + 2I- → 2SO42- + I2
Even though persulfate is a good oxidising agent and iodide is easily oxidised, the repulsion between the ions makes this reaction very slow at room temperature.
Using Fe2+ (aq) as a catalyst increases the rate of the reaction, because the two negative species – S2O82- being reduced, I- being oxidised – will react more readily with a positive species in solution instead.
Recall the roundabout image mentioned earlier. The Fe2+/Fe3+ ions are the roundabout.
- The persulfate, S2O82- ions can jump on at the Fe2+ ‘section’, where electrons are picked up, and persulfate gets reduced to SO42-. In return, Fe2+ becomes Fe3+ (aq).
- The I- ions can enter at this stage, dropping off electrons and oxidising to form I2. This dropping off of electrons reduces the Fe3+ (aq) to Fe2+ (aq). Back in their original state, these Fe2+ (aq) ions are able to react more persulfate ions now.
With this said, it doesn’t mater if you start with Fe2+ or Fe3+ ions in solution – the reaction creates a cycle between the two states.
This type of reactivity is how some extremely important chemical processes are run.
- One more example of catalysis is autocatalysis: autocatalysis occurs when a product of a reaction catalyses the reaction itself .
A well-known example of this is permanganate (MnO42-) ions reacting with oxalic acid (C2O42-) in acidic conditions.
2MnO42- + 16H+ + 5C2O22- → 2Mn2+ + 8H2O + 10CO2
The Mn2+ ions produced by this reaction catalyse the reaction further.
Autocatalysis sometimes gives reactions a strange rate profile. They can begin slowly, speed up considerably after some time as the amount of catalyst product goes up, before slowing again as reactant runs out.