Transition Metal Catalysts (Oxford AQA International A Level Chemistry)

Transition Metal Catalysts

• Catalysts play an important part in industry allowing reactions to be carried out more quickly and at lower temperatures

• Transition elements and their compounds are often used as catalysts because they can form ions with more than one stable oxidation state

  • This is due to vacant d orbitals

• There are two types of catalysts, heterogeneous and homogeneous catalysts

• A heterogenous catalyst is in a different physical state (phase) from the reactants

  • The reaction occurs at active sites on the surface of the catalyst

  • For example, vanadium(V) oxide, V₂O₅, is a solid catalyst used in the Contact process for making sulfuric acid

  • Another example is the use of solid iron, Fe, in the Haber process for making ammonia

• A homogeneous catalyst is in the same physical state (phase) as the reactants

  • An example of a homogeneous catalyst is the role of iron(II) ions, Fe²⁺, in the reaction between iodide ions, I^-, and peroxydisulfate ions, S₂O₈^2-

• Transition elements are often used as catalysts due to their ability to form ions with more than one stable oxidation state, and the fact that they contain vacant d orbitals

Heterogeneous Catalysts

• Some transition metals are precious metals which means that they can be very expensive

• In order to minimise the cost and maximise the efficiency of the catalyst the following measures can be taken:

  • Increasing the surface area of the catalyst

  • Coating an inert surface medium with the catalyst to avoid using large amounts of the catalyst

• This is achieved by spreading the catalyst over a hollow matrix such as a honeycomb-like structure

Diagram of a catalytic converter

• Catalytic converters are used in car exhaust boxes to reduce air pollution

  • They usually consist of a mixture of finely divided platinum and rhodium supported on a ceramic base

The Contact Process

• The manufacture of sulfuric acid is a very important piece of industrial chemistry that makes use of heterogeneous catalysis

• The first step of the process is roasting sulfur in air to produce sulfur dioxide

S (s) + O₂ (g)   →  SO₂ (g)

• The second step is an equilibrium reaction which is catalysed by vanadium(V) oxide, V₂O₅

2SO₂ (g) + O₂ (g) ⇌   2SO₃(g)

• The vanadium(V) oxide catalyst converts sulfur dioxide into sulfur trioxide and is reduced to vanadium(IV) oxide

SO₂ (g) + V₂O₅ (s) →  V₂O₄ + SO₃ (g)

• The vanadium(V) oxide is then re-generated by reaction with oxygen, fulfilling its role as a catalyst

O₂ (g) + 2V₂O₄ (s) →  2V₂O₅ (s)

• Overall, the vanadium(V) oxide has acted as a catalyst because it has:

  • Not been "used up" as they are reformed

  • Increased the rate of reaction

  • Provided an alternative reaction pathway

• This reaction shows that a variable oxidation state can also be utilised in heterogenous catalysis

The Haber Process

• The industrial production of ammonia takes place using the Haber process

• The main reaction is:

N₂ (g)  +  3H₂ (g)    ⇌    2NH₃ (g)

• Iron pellets are used as the catalyst to provide a large surface area

• In the reaction the surface of the iron attracts electrons in the hydrogen and nitrogen

  • These molecules form temporary loose attachments to the surface

  • This is called adsorption

How heterogeneous catalysts work

Limitations of heterogeneous catalysts

• The operating conditions of car exhausts as well as industrial reactions means that catalysts do not last forever

• Tiny quantities of impurities gradually become lodged on the surface of the catalyst binding with the active sites and preventing further catalytic reaction

  • This is know as 'poisoning' of the catalyst

• Other heavy metal impurities such as lead can do this

  • This is one of the reasons why lead additives were taken out of petrol as catalytic converters would soon fail due to the lead

  • The other reason was environmental

• As the catalyst has a very thin coating on the ceramic support it will be gradually be lost over time

  • This will reduce the efficiency of the catalyst / catalytic converter over time

• Catalytic converters do not work well at low temperatures

  • They need about 15 minutes of the engine running and warming up the exhaust box before they are effective

  • So, short journeys do almost nothing for pollution reduction

Homogeneous Catalysts

• Transition element ions can adopt more than one stable oxidation state

• This means that they can accept and lose electrons easily to go from one oxidation state to another

• Therefore, they can catalyse redox reactions by acting as both oxidising agents and reducing agents

Iron as a reducing and oxidising agent

• Iron is often used as a catalyst due to its ability to form Fe²⁺ and Fe³⁺ ions

• Fe²⁺ ions act as a reducing agent:

  • They reduce another species

  • They are oxidised

Fe²⁺ → Fe³⁺ + e^-

• The Fe³⁺ formed can then act as an oxidising agent:

  • They oxidise another species

  • They get reduced to reform the Fe²⁺ ion

Fe³⁺ + e^- → Fe²⁺

The iodide / peroxodisulfate reaction

• Iron(II) ions catalyse the reaction between iodide ions and peroxodisulfate ions, S₂O₈^2-

S₂O₈^2- + 2I^-  → I₂ + 2SO₄^2-

• The overall reaction is energetically favourable but slow

  • This is because the repulsion of two negative ions means that they are unlikely to collide successfully

• However, if iron(II) ions are added to the reaction, the rate is much quicker

• The Fe²⁺ ions can:

  • Reduce the peroxodisulfate ions to sulfate ions

  • Produce Fe³⁺ ions

S₂O₈^2- + 2Fe²⁺  →  2SO₄^2-   + 2Fe³⁺  

• The Fe³⁺ ions produced can:

  • Oxidise the iodide ions to iodine

  • Reform Fe²⁺ ions

2I^- + 2Fe³⁺  →  I₂   +  2Fe²⁺  

• Overall, the Fe²⁺ ions have acted as a catalyst because they have:

  • Not been "used up" as they are reformed

  • Increased the rate of reaction

  • Provided an alternative reaction pathway

Iodide / peroxodisulfate energy level diagram

Autocatalysis

• Autocatalysis is where a reaction speeds up as a catalytic product forms

  • This product acts as a catalyst for the original reaction

• The rate graph of concentration versus time has an usual shape

• The gradient becomes steeper during the course of the reaction

  • This indicates that the rate of reaction is increasing

  • Normally, the rate of reaction decreases as reactants are used up

Concentration versus time for an autocatalytic reaction

Example of an autocatalytic reaction

• The reaction between manganate(VII) ions and oxalate (ethandioate) ions is an example of an autocatalysis reaction

• The half equations for the reaction are:

MnO₄^– (aq) + 8H⁺ (aq) + 5e^– → Mn²⁺ (aq) + 4H₂O (l) 

C₂O₄^2– (aq) → 2CO₂ (g) + 2e^–

• These combine to give the overall reaction equation:

5C₂O₄^2– (aq) + 2MnO₄^– (aq) + 16H⁺ (aq) → 2Mn²⁺ (aq) + 8H₂O (l) + 10CO₂ (g)

• One product of the reaction is manganese(II) ions, Mn²⁺

• These Mn²⁺ ions are the catalyst for an alternative reaction involving the ethanedioate ions

  • The Mn²⁺ ions react with manganate(VII) ions to produce Mn³⁺ ions

  • The Mn³⁺ ions react with ethanedioate ions to produce carbon dioxide and reform the Mn²⁺ ions

4Mn²⁺ (aq)  +  MnO₄^– (aq) + 8H⁺ (aq)   →   5Mn³⁺ (aq)  + 4H₂O (aq)

2Mn³⁺ (aq)  +  C₂O₄^2- (aq)  →  2CO₂ (g) +  2Mn²⁺ (aq)

• The Mn²⁺ ions are not present at the start of the reaction

  • They are formed during the reaction and speed the reaction up

  • They are re-generated during the reactions from Mn²⁺ to Mn³⁺ and back to Mn²⁺

• This reaction can be followed with a colorimeter

  • The colour intensity of the purple manganate(VII) decreases over time

  • The rate of the colour change can be measured

  • The rate of reaction will increase over time

  • This is because the purple manganate(VII) ion are used up by the Mn²⁺ ions