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3.2.5-6 Inorganic Chemistry Y2 (2) (Catalysts (Heterogenous catalyst:…
3.2.5-6 Inorganic Chemistry Y2 (2)
Variable oxidation states
oxidation of lower oxidation states of transition metal ions tend to happen in alkaline solutions as there is a tendency to form negative ions
Vanadium species in oxidation states IV, III and II are formed by the reduction of vanadate(V) ions by zinc in acidic solution.
The redox potential for a transition metal ion changing from a higher to a lower oxidation state is influenced by pH and by the ligand.
The redox titration of Fe²⁺ and C₂O₄²⁻ with MnO₄⁻
calculations of titrations and similar redox reactions
only the lower oxidation states of transition metals actually exist as simple ions; higher oxidation state ions are covalently bonded to something else so exists as a compound ion
potassium manganate(VII) reaction:
construct half equation for reduction of Mn(VII) to Mn(II)
MnO₄⁻(aq) --> Mn²⁺
balance oxygens by adding water
MnO₄⁻(aq) --> Mn²⁺(aq) + 4H₂O(l)
balance hydrogens by adding H+ ions:
MnO₄⁻(aq) + 8H⁺(aq) --> Mn²⁺(aq) + 4H₂O(l)
balance charge by adding electrons
MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ --> Mn²⁺(aq) + 4H₂O(l)
construct half equation for reduction of iron(II) to iron(III)
Fe²⁺(aq) --> Fe³⁺(aq) + e⁻
balance both half equations so they have the same number of electrons
MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ --> Mn²⁺(aq) + 4H₂O(l)
5Fe²⁺(aq) --> 5Fe³⁺(aq) + 5e⁻
add the half equations together and cancel the electrons:
MnO₄⁻(aq) + 8H⁺(aq) + 5Fe²⁺(aq) --> Mn²⁺(aq) +5Fe³⁺(aq) + 4H₂O(l)
potassium manganate(VII) acts as an oxidising agent in acidic solution; oxidation number of manganese falls from +7 to +2
Catalysts
Fe²⁺ ions act as homogenous catalysts in the reaction between I⁻ and S₂O₈²⁻:
overall equation: S₂O₈²⁻(aq) + 2I⁻(aq) --> 2SO₄²⁻(aq) + I₂(aq)
S₂O₈²⁻(aq) + 2Fe²⁺(aq) --> 2SO₄²⁻(aq) + 2Fe³⁺(aq)
2Fe³⁺(aq) + 2I⁻(aq) --> 2Fe²⁺(aq) + I₂(aq)
uncatalyzed reaction is between two negative ions that repel each other so require a high activation energy; catalysed reaction involves reactions between oppositely charged ions so increase the reaction rate
autocatalysis: when one of the products of the reaction is a catalyst for the reaction; reaction starts slowly then as concentration of product builds up the rate of reaction increases
Mn²⁺ ions autocatalyse the reaction between C₂O₄²⁻ and MnO₄⁻:
overall equation: 2MnO₄⁻(aq) + 16H⁺(aq) + 5C₂O₄²⁻(aq) --> 2Mn²⁺(aq) +8H₂O(l) + 10CO₂(g)
4Mn²⁺(aq) + MnO₄⁻(aq) + 8H⁺(aq) --> 5Mn³⁺(aq) + 4H₂O(l)
2Mn3+(aq) + C₂O₄²⁻(aq) --> 2CO₂(g) + 2Mn²⁺(aq)
Transition metals and their compounds can act as heterogeneous and homogeneous catalysts.
Heterogenous catalyst:
Heterogeneous catalysts can become poisoned by impurities that block the active sites and consequently have reduced efficiency; this has a cost implication.
usually present as solids, and reactants are gases or liquids
A heterogeneous catalyst is in a different phase from the reactants and the reaction occurs at active sites on the surface.
increase SA and spread the catalyst on an inert support medium to increase SA:mass ratio
Homogenous catalyst:
A homogeneous catalyst is in the same phase as the reactants.
When catalysts and reactants are in the same phase, the reaction proceeds through an intermediate species.
explain the importance of variable oxidation states in catalysis
V₂O₅ acts as a heterogeneous catalyst in the Contact process:
overall equation: 2SO₂ + O₂ <=> 2SO₃
SO₂ + V₂O₅ --> SO₃ + V₂O₄
2V₂O₄ + O₂ --> 2V₂O₅
Fe is used as a heterogeneous catalyst in the Haber process:
N₂(g) + 3H₂(g) --> NH₃(g)
Reactions of ions in aqueous solution:
The acidity of [M(H₂O)₆]³⁺ is greater than that of [M(H₂O)₆]²⁺ ; M³⁺(aq) ions are smaller and have a greater charge so are more strongly polarising and attract electrons from the water ligands more; more O-H bonds in water are weakened; complex ion more readily releases H⁺ ions into solution making it acidic (hydrolysis):
[M(H₂O)₆]³⁺(aq) <=> [M(H₂O)₅(OH)]²⁺(aq) + H⁺(aq)
Some metal hydroxides show amphoteric character by dissolving in both acids and bases (eg hydroxides of Al³⁺).
M²⁺(aq) ions
add CO₃²⁻:
[M(H₂O)₆]²⁺(aq) + CO₃²⁻ --> MCO₃(s) + 6H₂O(l)
less acidity so carbonate ion can remove protons and metal carbonate can be formed
iron(II): green ppt
copper(II): blue-green ppt
add NH₃:
[M(H₂O)₆]²⁺(aq) + 2NH₃(aq) --> [M(H₂O)₄(OH)₂] (s) + 2NH₄⁺
iron(II): green gelatinous ppt --> oxidised to brown [Fe(H₂O)₃(OH)₃]
copper(II): pale blue ppt --> (if excess) deep blue solution of [Cu[NH₃)₄(H₂O)₂]²⁺
add OH⁻:
[M(H₂O)₆]²⁺(aq) + OH⁻(aq) --> [M(H₂O)₅(OH)]⁺(aq) + H₂O(l)
[M(H₂O)₅(OH)]⁺(aq) + OH⁻(aq) --> [M(H₂O)₄(OH)₂] (s) + H₂O(l)
iron(II): green gelatinous ppt --> oxidised to brown [Fe(H₂O)₃(OH)₃]
copper(II): pale blue ppt
[Fe(H₂O)₆]²⁺ pale green
[Cu(H₂O)₆]²⁺ pale blue
M³⁺(aq) ions
add NH₃:
[M(H₂O)₆]³⁺(aq) + 3NH₃(aq) --> [M(H₂O)₃(OH)₃] (s) + 3NH₄⁺
iron(III): brown gelatinous ppt
aluminium(III): white ppt
[Al(H₂O)₆]³⁺ colourless
[Fe(H₂O)₆]³⁺ purple/yellow/brown
add OH⁻:
[M(H₂O)₆]³⁺(aq) + OH⁻(aq) --> [M(H₂O)₅(OH)]²⁺ (aq) + H₂O(l)
[M(H₂O)₅(OH)]²⁺(aq) + OH⁻(aq) --> [M(H₂O)₄(OH)₂]⁺ (aq) + H₂O(l)
[M(H₂O)₄(OH)₂]⁺(aq) + OH⁻(aq) --> [M(H₂O)₃(OH)₃] (s) + H₂O(l)
iron(III): brown gelatinous ppt
aluminium(III): white ppt --> (if excess) colourless solution of [Al(OH)₄]⁻
add CO₃²⁻-:
2[M(H₂O)]₆(aq) + 3CO₃²⁻(aq) --> 2[Fe(H₂O)₃(OH)₃] (aq) + 3CO₂(g) + 3H₂O(l)
greater acidity of M³⁺ ions prevents metal carbonate from being formed; effervescence due to CO₂ being released
iron(III): brown gelatinous ppt
aluminium(III): white ppt
Lewis acid: electron pair acceptors
Lewis base: electron pair donors
Sammer Sheikh