5.3Unit 5. Transition Metals and Organic Nitrogen Chemistry
Transition Metal Reactions — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions
10 min read 12 parts Lesson 4 of 8 1 quiz
Lesson summary
5.3 Transition Metal Reactions — IAL EDEXCEL MODULAR (CHEMISTRY) notes: Describe the redox chemistry of vanadium and chromium, reactions of aqua ions with bases, and catalysis by transition metals.Oxidation states — vanadium shows +2, +3, +4 and +5, each with a distinct colour.Ox…
- Electrode potentials
- Strongest agents
- Common confusion
- Zinc as reducing agent
- Reduction to vanadium metal
- Acidic reduction by zinc
Section 1 of 12Vanadium oxidation states and electrode potentials
- Oxidation states
- vanadium shows +2, +3, +4 and +5, each with a distinct colour.
| Oxidation state | Species | Colour |
|---|---|---|
| +5 | VO2+ | yellow |
| +4 | VO2+ | blue |
| +3 | V3+ | green |
| +2 | V2+ | violet |
Electrode potentials
- V2+ + 2e- ⇌ V is -1.18 V and V3+ + e- ⇌ V2+ is -0.26 V.
- VO2+ + 2H+ + e- ⇌ V3+ + H2O is +0.34 V.
- VO2+ + 2H+ + e- ⇌ VO2+ + H2O is +1.00 V.
Strongest agents
- VO2+ is the strongest oxidising agent and V (most negative E°) is the strongest reducing agent.
- (most positive E°)
Common confusion
- Common confusion the two oxo ions VO2+ and VO2+ must not be mixed up.
Section 2 of 12Reducing vanadium(V) with zinc
1
Dissolve ammonium vanadate(V) in acid to give yellow VO2+.
2
Add zinc, which reduces it to blue VO2+ via a green mixture.
3
Further reduction gives green V3+.
4
Finally violet V2+ forms, which air readily oxidises.
Reducing vanadium(V) with zinc
Zinc as reducing agent
- Zn2+ + 2e- ⇌ Zn is -0.76 V, so zinc reduces vanadium from +5 to +4, +4 to +3 and +3 to +2 because each E°cell is positive.
- 2VO2+ + 4H+ + Zn → 2VO2+ + Zn2+ + 2H2O.
- 2VO2+ + 4H+ + Zn → 2V3+ + Zn2+ + 2H2O.
- 2V3+ + Zn → 2V2+ + Zn2+.
Reduction to vanadium metal
- V2+ → V by Zn is not feasible because E°cell is negative.
Method
- Method the same half-equation method predicts which oxidising agents can raise vanadium to a higher oxidation state.
- The half-equation with the more positive E° is the reduction and the other is reversed.
- Electrons are balanced before the half-equations are added.
Section 3 of 12Chromium oxidation states and redox
Chromium oxidation states and redox
Oxidation states
- +2 (blue Cr2+), +3 (green [Cr(H2O)6]3+) and +6 (orange Cr2O72-, yellow CrO42-).
Strongest agents
- Cr2O72- (+1.33 V) is the strongest oxidising agent and zinc (-0.76 V) is the strongest reducing agent in this set.
Acidic reduction by zinc
- Cr2O72- + 14H+ + 3Zn → 2Cr3+ + 7H2O + 3Zn2+ .
- (orange to green)
- Further reduction in acid gives blue Cr2+: 2Cr3+ + Zn → 2Cr2+ + Zn2+.
- Cr2+ is re-oxidised by air.
Alkaline oxidation by hydrogen peroxide
- 2Cr(OH)3 + 4OH- + 3H2O2 → 2CrO42- + 8H2O.
- Cr(III) is oxidised to yellow chromate(VI) in alkaline conditions.
Reaction schemes
- Reaction schemes require the formula and state, colour, reagent and conditions for each conversion.
- Green [Cr(H2O)6]3+ + 3OH- → grey-green [Cr(H2O)3(OH)3](s) + 3H2O.
Section 4 of 12Chromate(VI) and dichromate(VI) equilibrium
Chromate(VI) and dichromate(VI) equilibrium
Equilibrium
- 2CrO42- + 2H+ ⇌ Cr2O72- + H2O.
Adding acid
- Adding acid shifts the equilibrium to dichromate and the colour changes from yellow to orange.
Adding alkali
- Adding alkali removes H+, shifts the equilibrium to chromate and gives yellow.
Type of reaction
- Type of reaction it is an acid-base reaction, not redox, because chromium is +6 in both ions.
Stability
- Stability chromate is stable in alkaline solution and dichromate is more stable in acidic solution.
Section 5 of 12Aqueous ions with sodium hydroxide
| Ion | Aqua ion colour | With NaOH(aq) | With excess NaOH(aq) |
|---|---|---|---|
| Cr3+ | green solution | green precipitate Cr(OH)3(H2O)3 | green solution [Cr(OH)6]3- |
| Fe2+ | green solution | green precipitate Fe(OH)2(H2O)4 | no change |
| Fe3+ | yellow-brown solution | brown precipitate Fe(OH)3(H2O)3 | no change |
| Co2+ | pink solution | blue precipitate Co(OH)2(H2O)4 | no change |
| Mn2+ | pale pink solution | pale brown precipitate Mn(OH)2(H2O)4 | no change |
| Ni2+ | green solution | green precipitate Ni(OH)2(H2O)4 | no change |
| Cu2+ | blue solution | blue precipitate Cu(OH)2(H2O)4 | no change |
| Zn2+ | colourless solution | white precipitate Zn(OH)2(H2O)4 | colourless solution [Zn(OH)4(H2O)2]2- |
- Formula of the precipitate
- the number of OH- ligands equals the charge on the metal ion.
Section 6 of 12Aqueous ions with ammonia
| Ion | With NH3(aq) | With excess NH3(aq) |
|---|---|---|
| Cr3+ | green precipitate | purple solution [Cr(NH3)6]3+ |
| Fe2+ | green precipitate | no change |
| Fe3+ | brown precipitate | no change |
| Co2+ | blue precipitate | yellow solution [Co(NH3)6]2+ |
| Mn2+ | pale brown precipitate | no change |
| Ni2+ | green precipitate | dark blue solution [Ni(NH3)6]2+ |
| Cu2+ | blue precipitate | dark blue solution [Cu(NH3)4(H2O)2]2+ |
| Zn2+ | white precipitate | colourless solution [Zn(NH3)4]2+ |
Excess ammonia
- Excess ammonia only Cr3+, Co2+, Ni2+, Cu2+ and Zn2+ dissolve by ligand substitution.
Cobalt(II) ammine
- Cobalt(II) ammine the yellow [Co(NH3)6]2+ darkens to brown in air as it is oxidised to cobalt(III).
Ligand exchange
- NH3 and H2O are similar in size and uncharged, so the co-ordination number stays the same except for zinc, which changes from six to four.
Section 7 of 12Reactions of aqua ions with bases
Limited OH− or NH3
- Limited OH− or NH3 deprotonates water ligands to give neutral hydroxide precipitates.
- [Cu(H2O)6]2+ + 2OH- → [Cu(H2O)4(OH)2] + 2H2O.
- [Fe(H2O)6]3+ + 3NH3 → [Fe(H2O)3(OH)3] + 3NH4+.
- It looks like ligand substitution but is an acid-base reaction in which the aqua ion acts as an acid.
Acidity of aqua ions
- [M(H2O)6]3+ ions are more acidic than [M(H2O)6]2+ ions because the higher charge density polarises the water ligands more.
- [Fe(H2O)6]3+ ⇌ [Fe(H2O)5(OH)]2+ + H+.
Excess OH−
- Excess OH− further deprotonation redissolves Cr(III) and Zn hydroxides.
- [Cr(H2O)3(OH)3] + 3OH- → [Cr(OH)6]3- + 3H2O.
- The hydroxide also reacts with acid, [Cr(H2O)3(OH)3] + 3H+ → [Cr(H2O)6]3+, so it is amphoteric.
Excess NH3
- NH3 acts as a Lewis base and substitutes ligands.
- [Cu(H2O)4(OH)2] + 4NH3 → [Cu(NH3)4(H2O)2]2+ + 2H2O + 2OH-.
- [Zn(H2O)6]2+ + 4NH3 → [Zn(NH3)4]2+ + 6H2O.
Section 8 of 12Heterogeneous catalysis
Heterogeneous catalysis
Heterogeneous catalyst
- Heterogeneous catalyst in a different phase from the reactants, usually a solid with gaseous or dissolved reactants, reacting at active sites on its surface.
Surface adsorption theory
- Surface adsorption theory three stages.
- Adsorption of reactants at active sites.
- Reaction after the bonds in the adsorbed reactants are weakened.
- Desorption of the products from the surface.
Benefits of adsorption
- Benefits of adsorption bonds weaken or molecules are held in a reactive orientation,
- and reactant concentration at the surface is higher so collision frequency is greater.
Strength of adsorption
- Strength of adsorption must be intermediate.
- Tungsten adsorbs too strongly so products cannot leave.
- Silver adsorbs too weakly so too few reactants attach.
- Nickel and platinum have about the right strength.
Surface area
- Surface area a support medium maximises surface area and minimises cost, for example rhodium on a ceramic support.
Advantages
- Advantages easy to filter off and suited to continuous processes.
Poisoning
- Poisoning impurities adsorb on active sites or bond strongly so they do not desorb, reducing effectiveness.
Section 9 of 12Contact process and catalytic converters
1
Contact process
- S + O2 → SO2, then 2SO2 + O2 ⇌ 2SO3 over a vanadium(V) oxide catalyst.
- SO2 + V2O5 → SO3 + V2O4, with vanadium reduced from +5 to +4.
- 2V2O4 + O2 → 2V2O5, with vanadium oxidised back to +5.
- Variable oxidation states are used in heterogeneous catalysis.
2
Catalytic converters
- Catalytic converters finely divided platinum and rhodium on a ceramic honeycomb support in car exhausts.
- They remove CO (toxic), NO2 (respiratory irritant causing acid rain) and unburnt hydrocarbons.
- 2NO + 2CO → N2 + 2CO2.
- CO and NO adsorb, their bonds weaken, they react, and CO2 and N2 desorb.
- A coating on a support and a large surface area reduce the cost of the precious metals.
3
Adsorption and absorption
- Adsorption and absorption adsorption happens only at a surface while absorption is distribution throughout a substance.
Section 10 of 12Homogeneous catalysis and autocatalysis
Homogeneous catalysis and autocatalysis
Homogeneous catalyst
- Homogeneous catalyst in the same phase as the reactants, usually all in aqueous solution, and works through an intermediate with a specific formula.
Transition metal ions as catalysts
- Transition metal ions as catalysts variable oxidation states let them accept and release electrons, acting as both oxidising and reducing agents.
Iron(II) in the iodide-peroxodisulfate reaction
- S2O82- + 2I- → 2SO42- + I2 is slow because two negative ions repel.
- Fe2+ reduces peroxodisulfate: S2O82- + 2Fe2+ → 2SO42- + 2Fe3+.
- Fe3+ oxidises iodide: 2I- + 2Fe3+ → I2 + 2Fe2+, regenerating the catalyst.
- Each step involves oppositely charged ions, so the activation energy is lower.
- Starch gives a blue-black colour showing iodine formation.
Autocatalysis
- Autocatalysis a reaction is speeded up by one of its products.
- The concentration-time graph becomes steeper as the reaction proceeds, so the rate rises at first.
Manganese(II) in MnO4− + C2O42-
- 2MnO4- + 16H+ + 5C2O42- → 2Mn2+ + 10CO2 + 8H2O.
- Mn2+ is absent at the start so the reaction begins slowly.
- 4Mn2+ + MnO4- + 8H+ → 5Mn3+ + 4H2O.
- 2Mn3+ + C2O42- → 2CO2 + 2Mn2+.
- The curve is S-shaped, slow, then fast, then slowing as reactants run out; a colorimeter can follow it.
Section 11 of 12Key terms
| Term | Meaning |
|---|---|
| Variable oxidation state | The ability of transition metals to form ions in several oxidation states because 3d and 4s energies are close. |
| Amphoteric hydroxide | A hydroxide that reacts with both acids and bases, such as Cr(OH)3. |
| Heterogeneous catalyst | A catalyst in a different phase from the reactants, acting at active sites on its surface. |
| Homogeneous catalyst | A catalyst in the same phase as the reactants, working through an intermediate. |
| Autocatalysis | A reaction in which a product acts as a catalyst for the reaction. |
| Adsorption | Attachment of molecules to the surface of a solid. |
| Catalyst poisoning | Loss of activity when impurities occupy the active sites of a catalyst. |
Section 12 of 12Exam technique
1
Treating ligand exchange with ammonia as a change of oxidation state — the metal oxidation state is unchanged.
2
Mixing up VO₂⁺ (+5, yellow) and VO²⁺ (+4, blue).
3
Saying chromate-dichromate interconversion is redox — chromium stays +6, so it is an acid-base equilibrium.
4
Confusing adsorption with absorption — adsorption is bonding at a surface only.
5
Writing the hydroxide precipitates as simple M(OH)n — they are neutral aqua-hydroxide complexes such as [Cu(H₂O)₄(OH)₂].
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