15.2 - Transition metal reactions

Redox - The oxidation number of the transition metal ion changes Acid-base reaction - One or more of the ligands gain or loses a hydrogen ion Ligand exchange - One or more of the ligands around the transition metal ion is replaced by a different ligand Coordination number change - The number of ligand changes Any one of these reactions can cause a change in colour of the complex.

Aqueous solution containing Fe2+ ions is pale green, but when exposed to the air it gradually turns yellow or brown. As oxidation number of iron increases from +2 to +3. The type and number of ligands remain unchanged in this oxidation reaction. Formula of the 2 complexes [Fe(H2O)6]2+ and [Fe(H2O)6]3+. The colour change in this reaction is best illustrated using solid samples containing these ions.

When aqueous sodium hydroxide is added to copper(II) sulfate solution. Observation is pale blue solution forms a blue precipitate. Equation: [Cu(H2O)6]2+ + 2OH- > [Cu(H2O)4(OH)2] + 2H2O Acid base reaction - 2 Hydroxide ions have removed hydrogen ions from 2 of the water ligands and converted them into water molecules. The 2 water molecules that have lost hydrogen ions are now hydroxide ligands. Term amphoteric behaviour is used to describe acid base reaction because its reversible. When acid is added to the blue precipitate, the hydrogen ions from the acid react with the hydroxide ions ligands and convert them back into water molecules.

Exactly same observations made during careful addition of aqueous ammonia instead of aqueous sodium hydroxide. [Cu(H2O)6]2+ + 2NH3 > [Cu(H2O)4(OH)2] + 2NH4+ Much easier to see this equation involves acid base reaction 2 of the water ligands transfer a hydrogen ion to the ammonia molecules. When aqueous ammonia is added to the blue precipitate formed it dissolves to form a deep blue solution [Cu(H2O)4(OH)2] + 4NH3 > [Cu(NH3)4(H2O)2]2+ + 2H2O + 2OH- Ligand exchange reaction 4 ammonia molecules replace 2 water molecules and 2 hydroxide ions. Starts light blue solution forms deep blue solution.

Reactions involve a change in ligand. Good example is reaction between copper(II) sulfate solution and conc HCl. When acid is added slowly and continuously the colour change from blue to green and finally yellow [Cu(H2O)]2+ + 4Cl- >< [CuCl4]2- + 6H2O All 6 water igands have been substituted by 4 chloride ions. Reaction is also an example of change in coordination number from 6 to 4 although change on complex change from 2+ to 2- no oxidation change. Reaction is reversible which explains the colour change observed. Hexaaqucopper (II) ion is blue and the tetrachlorocuprate (II) is yellow so the green colour is due to a mix of blue and yellow ions. Solutions at the start and end of the reaction.

When sodium hydroxide added to a solution containing hexaaquacobalt(II) ion until no further change can be seen. Observation is a pink solution forms a blue precipitate [Co(H2O)6]2+ + 2OH- > [Co(H2O)4(OH)2] + 2H2O With Copper acid base reaction 2 hydroxide ions removed hydrogen ions forms 2 of the water ligands and converted them into water molecules. The 2 water molecules have lost hydrogen ions and are now hydroxide ions. Colour of precip changes to pink.

Same observations when aqueous is used as the alkali. However when ammonia is added to excess there is an observation not made with aqueous sodium hydroxide precipitate dissolves to form a brown solution. Equation for acid-base reaction [Co(H2O)6]2+ + 2NH3 > [Co(H2O)4(OH)2] + 2NH4+ Equation for the ligand exchange reaction forming brown solution is: [Co(H2O)4(OH)2] + 6NH3 > [Co(NH3)6]2+ + 4H2O + 2OH- Brown solution changes colour because of the oxidation by the oxygen in the atmosphere. Oxidation number of cobalt increases from 2+ to 3+ and yellow [Co(NH3)6]3+ Ion forms. Resulting solution is much darker than yellow other products whose formation is different

When conc HCl is added to a solution containing the hexaaquacobalt (II) ion, pink solution gradually changes to blue. [Co(H2O)6] + 4Cl- > [CoCl4]2- + 6H2O 6 water ligands replaced by 4 chloride ions reaction is an example of change in coordination number from 6 to 4.

Reaction between sodium hydroxide is added to a solution containing the hexaaquairon (II) iron until a pale green solution forms a green precipitate. [Fe(H2O)6]2+ + 2OH- > [Fe(H2O)4(OH)2] This is an acid base reaction - the 2 hydroxide ions have removed the hydrogen ions from 2 of the water ligands and converted them into water molecules. 2 water molecules have lost hydrogen ions are now hydroxide ligands. Same observation with ammonia is used as alkali. The colour of the green precipitate changes to brown as the oxygen from the atmosphere causes oxidation forming [Fe(H2O)3(OH03] triaquatrihydroxoiron complex

Reaction with sodium hydroxide added to Hexaaquiron (III) ion observation a yellow brown solution forms a brown precipitate Product is [Fe(H2O)3(OH)3] 3H2O Acid base reaction - 3 hydroxide ions have removed hydrogen ions from 3 of the water ligands and converted them into water molecules. The 3 water molecules that have lost hydrogen ions are now hydroxide ligands. Same observation with aqueous ammonia as the alkali. Products [Fe(H2O)3(OH)3] + 3NH4+ No further reactions when an excess of either aqueous sodium hydroxide or aqueous ammonia is added and no further changes.

Some reactions of chromium is similar to those of copper cobalt and iron ions. Chromium has ions which metal has an oxidation number of +6 so reactions involving these ions must be considered differently.

Some compounds have different colours as solids and aqueous solutions Colour of a solution is dependant on concentration Presence of dissolved oxygen in an aqueous solution can affect the colour observed

Sodium hydroxide added to solution containing hexaaquachromium(III) ion until no further change. Observation is that a green or violet solution forms a green precipitate [Cr(H2O)6]3+ + 3OH- > [Cr(H2O)3(OH)3] + 3H2O Acid base reaction the 3 hydroxide ions have removed hydrogen ions from 3 of water ligands and converted them into water molecules that have lost hydrogen ions are now hydroxide ions. Same observation with aqeuous ammonia 3NH4+ formed

Green precipitate dissolves to form a green solution. Equation for this acid base reaction [Cr(H2O)3(OH)3] + OH- > [Cr(H2O)2(OH4)- + H2O If it is more concentrated further acid base reactions occur loss of water No change in colour Reaction involving hydroxide ions can be reversed by the addition of acid illustrating the reversible nature

Green precipitate is slow to dissolve but eventually a violet or purple solution forms. [Cr(H2O)6(OH)3] + 6NH3 > [Cr(NH3)6]3+ +3H2O + 3OH- Providing solutions are alkaline, oxidation is easily achieved by the addition of oxidising agent hydrogen peroxide. Reaction the solution changes from green to yellow as the chromate (VI) ion, with oxidation number +6 is formed 2[Cr(OH)6]3- + 3H2O2 > 2CrO42- + 2OH- + 8H2O

Chromate (VI) ions are stable in alkaline solution, but in acidic conditions the dichromate (VI) ion is more stable. So if acid is added there is a colour change from yellow to orange 2CrO42- + 2H+ >< CrO22- + H2O Reaction reversed by adding alkali.

When zinc metal is added to an acidic solution containing dichromate (VI) ions reductions reactions occur which the oxidation number of chromium decreases from +3 to +2 First stage of reduction involves a colour change from orange to green Cr2O72- +14H+ +3ZN > 2Cr3+ +7H2O +3Zn2+ Second stage involves a colour change from green to blue.

Oxidation of chromium(III) by hydrogen peroxide in alkaline conditions. 2Cr(OH)3 + 3H2O2 + 4OH- > 2CrO42- + 8H2O Best reducing agent is Cr(OH)3, best oxidising agent is H2O2.

Simplified equation of Chromium (VI) by zinc in acidic conditions. Cr2O72- + 14H+ + 3Zn > 2Cr3+ + 7H2O + 3Zn2+ Relevant equations are half equations are 1 and 5. Best reducing agent is Zn and best oxidising agent is Cr2O72- half equation 1 reversed to half equation 6 Since E value for half equation is more negative than E value for half equation 5, Zn is electron releasing with respect to CR2O72- hence the reaction is thermodynamically feasible.

Simplified equation for reduction of chromium(III) by zinc in acidic conditions is: 2Cr3+ + Zn > 2Cr2+ + Zn2+ Relevant half equations are 1 and 2. Best reducing agent is Zn best oxidising agent CR3+. Half equation 1 reversed to half equation 2. Since E values for half equation 1 is more negative than the E value for half equation 2, Zn is electron releasing with respect to Cr3+ reaction thermodynamically feasible.

+2, V2+, vanadium (II) , purple in aqueous solution +3, V3+, vanadium (III), green in aqueous solution +4, VO 2+, oxovanadium(IV), blue in aqueous solution +5, VO2 +, dioxovanadium (V), yellow in aqueous solution

ammonium vanadate (IV) NH4VO3. an acidic solution contains the dioxovanadium (V) ion, VO2+. When zinc is added to this solution, reduction begins and there is a gradual colour change from yellow, through blue and green, to purple as the oxidation number decreases from +5 t +2. All of these colours are of solutions no precipitates.

+5 to +4 VO2+ + 2H+ +e- > VO2+ + H2O Zn > Zn2+ + 2e- Since the E value for half equation 2 is more negative than the E value for half equation 5. Zn is electron releasing with respect to VO2+ so thermodynamically feasible +4 to +3 VO2+ + 2h+ + e- > V3+ + H2O Zn > Zn2+ + 2e- Since E value for half equation for 2 is more negative than the E value for half equation 4 the electron releasing with respect to VO2+ reaction thermodynamically feasible. +3 to +2 V3+ + e- > V3+ Zn > Zn2+ + 2e- Since e value for half equation 2 is more negative than the e value for half equation3, Zn is electron releasing with respect to V3+ so thermodynamically feasible

V2+ + 2e- > V Zn > Zn2+ + 2e- Since e value for half equation 2 is less negative than E value for half equation 1 Zn is not electron releasing with respect to V2+ reaction is not thermodynamically feasible