Catalysts

Created by

abit lol

Cards (16)

Redox titration involving reduction and oxidation
Manganate (VII) titrations
MnO4- + 5 E- + 8 H+ -> 4 H2O + Mn2+
Manganate ions are used to find the concentration of iron (II) ions
Fe2+ -> e- + Fe3+
Full equation
8 H+ + MnO4- + 5 Fe2+ -> Mn2+ + 4 H2O + 5 Fe3+
This reaction is self-indicating
As MnO4- ion reacts with Fe2+ ions the purple colour disappears.
When all the Fe2+ ions have created a pink tinge is observed due to the excess MnO4- .
Redox titration involving reduction and oxidation
Potassium dichromate (VI) titrations
(Cr2O7)2- + 6 E- + 14 H+ -> 7 H2O + 2 Cr3+
dichromate ions are used to find the concentration of iron (II) ions
Fe2+ -> e- + Fe3+
Full equation
14 H+ + (Cr2O7)2- + 6 Fe2+ -> 2 Cr3+ + 7 H2O + 6 Fe3+
Heterogeneous catalyst exist in a different phase or state to the reactants. The reaction takes place on the surface of the catalyst The place where the reaction happens is called the active site
The three main processes of catalysts
Adsorption of the reactants
The reaction
Desorption of the products
Contact Process -> Makes sulfuric acid
Variable oxidation states are important in heterogeneous catatlysis
Catalyst: Vanadium(V) oxide, V2O5
The vanadium(v) oxide catalyst oxidises sulfur(IV) oxide, SO2
to make Sulfur (VI) oxide, SO3
SO2 + V2O5 → SO3 + V2O4
The vanadium (IV) oride, V2O4 reacts with oxygen and the vanadium (V) oxide Catalyst is regenerated
V2O4 +½ O2 → V2O5
the overall equation: SO2 + ½ O2 → SO3
Haber Process
Nitrogen gas and hydrogen gas reacts to form NH3 in the presence of Fe catalysts
N2 + 3H2 ⇌ 2NH3
the nitrogen and hydrogen adsorb onto the surface of Fe/ lron catalyst
The reaction takes place
Ammonia is desorbed from the surface.
The hydrogen can be manufactured by the reaction of methane with steam
CH4 + 2H2O → CO2 + 4H2
> Natural gas is source of methane
> However natural gas contains H2S
> The sulfur in H2S adsorbs strongly to the active site on the Fe catalyst surface.
>This poisons the iron catalyst, therefore hydrogen an nitrogen cannot adsorb onto fe catalyst.
> Therefore natural gas must be treated before use to remove the H2S.
Catalytic Converters
> These devices help to remove harmful gases from car exhausts
Catalysis: Platinum & rhodium
These catalyst are very expensive.
It is spread in a thin layer over a honey comb support which increases the surface area between the catalyst and harmful gases.
This makes effective use of the catalyst to minimises costs.
Lead poisons the Pt and Rh catalysts as it adsorbs strongly to the active sites on the Pt and Rh catalyst surface.
Cars fitted with catalytic converter must not use leaded petrol hence why we use unleaded petrol
Homogenous Catalysts are in the same phase or state as the reactants.
Decomposition of Hydrogen peroxide
Bromine is the homogenous catalyst and stays the same state as the reactants and remains unchanged
Step 1: H2O2 + BR2 -> O2 + 2Br- + 2H+
Step 2: H2O2 + 2Br- + 2H+ -> 2H2O + Br2
Overall : 2H2O2 -> 2H2O + O2
Iodide and peroxidisulfate ions
The reaction is slow because both reactants are negatively charged so they repel (high Ea).
Aqueous Fe act as homogenous catalyst and as an intermediate. They are oppositely charged and lowers the activation energy.
Step 1: (S2O8)2- + 2 Fe2+ -> 2(SO4)2- + 2 Fe3+
Step 2: 2I- + 2 Fe3+ -> 2 Fe2+ + I2
Overall: (S2O8)2- + 2I- -> 2(SO4)2- + I2
Autocatalytic Reaction between Ethanedioate and Manganate ions
The overall reaction between the manganate ion and ethanedioate ion is : 2MnO4 + 5(C2O4)2- + 16H+ → 2Mn2+ + 10CO2 + 8H2O
This is an example of autocatalysis where one of the products, Mn2+ of the reaction can catalyse the reaction
Catalysed alternative route
MnO4 + 4 Mn2+ + 8 H+ → 5 Mn3+ + 4 H2O
(C2O4)2- + 2 Mn3+ → 2 Mn2+ + 2 CO2
The initial uncatalysed reaction is slow because the reaction is a collision between two negative ions which repel each other leading to a high activation energy
The Mn2+ ions produced act as an autocatalyst and therefore the reaction starts to speed up because they bring about the alternative reaction route with lower activation energy. The reaction has a lower activation energy because the reactants are now oppositely charged. The reaction eventually slows as the concentration of MnO4 drops and the number of successful collisions decreases
Vanadium
Ammonium vanadate, NH4 VO3 contains the (VO3)- ion
> Where vanadium has a +5 oxidation state.
> This can be reduced using zinc and conc hydrochloric acid or conc sulfuric acid.
>Vanadium will be reduced from +5 to + 4 to + 3 to +2 . Each vanadium species has a different colour
>If the reaction is done in a small flask, it is stoppered with cotton wool
> This allows hydrogen (produced between the zinc and acid)
to escape. This also stops air from entering. This prevents
oxidation of the vanadium +2 species by oxygen in the air
Vanadium
When NH4VO3 solid is dissolved in an acidic solution, a yellow solution of [VO2(H2O)4]+ is produced oxidation state is +5
Vanadium is reduced from
Yellow solution [VO2(H2O)4]+ (v=+5)
to:
Blue solution [VO(H2O)5]2+ (v=+4)
to
green solution [V(H2O)6]3+ (v=+3)
to
purple solution [V(H2O)6]2+ (v=+2)
The vanadium (II) species is easily oxidised. If you remove the cotton wool from the flask and pour the solution into a test tube , it turns green due to oxygen in the air.
It is oxidised back to vanadium
Choosing the correct acid for manganate titrations
The acid is needed to supply the 8H+ ions. Some acids are not suitable as they set up alternative redox reactions and hence make the titration readings inaccurate.
Only use dilute sulfuric acid for manganate titrations. Insufficient volumes of sulfuric acid will mean the solution is not acidic enough and MnO2 will be produced instead of Mn2+.
MnO4-(aq) + 4H+(aq) + 3e- -> MnO2 (s) + 2H2O
The brown MnO2 will mask the colour change and lead to a greater (inaccurate) volume of manganate being used in the titration.