8. Transition metal catalysts

transition metals and their compounds make good catalysts

because they can change oxidation states gaging or losing electrons within their d orbitals

transition metals can transfer electrons

to speed up reactions

other elements aren’t used to catalyse redox reactions

because they dont have an incomplete 3d subshell and dont have variable oxidation states

heterogeneous catalyst

a catalyst that’s a different physical phase from the reactants

when heterogeneous catalysts are used

the reaction occurs on the surface of the catalyst

increasing the surface area of the catalyst

increases the number of molecules that can react at the same time, thus increasing the rate of reaction

support medium

often used to make the area of the catalyst as large as possible

during a reaction

reactants are adsorbed onto active sites on the surface of heterogeneous catalyst

impurities in the reaction mixture may also bind to the surface of a catalyst

this blocks reactants from being adsorbed

the process of impurities blocking reactants from being adsorbed

catalyst poisoning

catalyst poisoning reduces the surface area of the catalyst available to reactants

slowing down the reaction

catalyst poisoning increases the cost of chemical process

because less product can be made in a certain time or w certain amount of energy

catalysts may even need replacing or regenerating

which also costs money

catalyst poisoning can be reduced by purifying the reactants

this removes many impurities which would otherwise poison the catalyst

homogeneous catalyst

catalysts in the same physical state as reactants

usually homogeneous catalysts are aqueous catalysts

for reactions between 2 aqueous solutions

homogeneous catalysts work

by forming intermediate species

the reactants combine with the catalyst to make an intermediate species

which then reacts to form the products and reform the catalyst

this causes an enthalpy profile for a homogeneously catalysed reaction to have 2 peaks

corresponding the 2 reactions

the activation energy required to form the intermediate and product from it

is lower than the activation energy needed to make the products directly from the reactans

catalysts are always reformed

so it can carry on catalysing the reaction

Fe²⁺ catalysing the reaction between

S₂O₈^2- and I^-

redox reaction between iodide ions and peroxodisulfate S₂O₈^2- Ions

S₂O₈^2- + 2I^- → I₂ + 2SO₄^2-

The reaction occurs annoyingly slowly because both ions are negatively charged

the ions repel each other so it sunlikely they;ll collide and react

is Fe²⁺ ions are added, things are really sped up

because each stage of the reaction involves a positive and negative ion so there’s no repulsion

1. Fe²⁺ ions are oxidised to Fe³⁺ ions by S₂O₈^2- ions

2. the newly formed Fe³⁺ ions now easily oxidises I^- ions getting reduced in the process and reforming as Fe²⁺

you can test for iodine by adding start solution

the solution turning blue-black indicates iodine is present

Mn²⁺ autocatalysing the reaction

between MnO₄^- and C₂O₄^2-

is it an autocatalysing reaction

because the Mn²⁺ is a product of the reaction and acts as a catalyst for the reaction

this means that as the reaction progresses and the amount of the product increases

the reaction speeds up

there isnt many Mn²⁺ present at the beginning of the reaction to catalyse it

so at first the reaction rate is slow

during the uncatalysed part of the reaction

the activation energy is very high

because the reaction proceeds via the collision of negative ions

which require alot of energy to achive

once alittle Mn²⁺ catalyst has been made it reacts with the MnO₄^- ions

to make Mn³⁺ ions

Mn³⁺ ions are intermediate and react w C₂O₄^2- ions

to make CO₂ and reform Mn²⁺ catalyst

because Mn²⁺ autocatalyse the reaction

the reaction rate increases with time as more catalyst is made

this means a concentration time graph of the reaction looks a lil funny