1. AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING

2. INTRODUCTION This Powerpoint show is one of several produced to help students understand selected topics at AS and A2 level Chemistry. It is based on the requirements of the AQA and OCR specifications but is suitable for other examination boards. Individual students may use the material at home for revision purposes or it may be used for classroom teaching if an interactive white board is available. Accompanying notes on this, and the full range of AS and A2 topics, are available from the KNOCKHARDY SCIENCE WEBSITE at... www.argonet.co.uk/users/hoptonj/sci.htm Navigation is achieved by... either clicking on the grey arrows at the foot of each page orusing the left and right arrow keys on the keyboard KNOCKHARDY PUBLISHING TRANSITION METALS

3. CONTENTS Aqueous metal ions Acidity of hexaaqua ions Introduction to the reactions of complexes Reactions of cobalt Reactions of copper Reactions chromium Reactions on manganese Reactions of iron(II) Reactions of iron(III) Reactions of silver and vanadium Reactions of aluminium TRANSITION METALS

4. THE AQUEOUS CHEMISTRY OF IONS - HYDROLYSIS when salts dissolve in water the ions are stabilised this is because water molecules are polar hydrolysis can occur and the resulting solution can become acidic the acidity of the resulting solution depends on the cation present the greater the charge density of the cation, the more acidic the solution cationchargeionic radius reaction with water / pH of chloride Na 1+0.095 nm Mg 2+0.065 nm Al 3+0.050 nm the greater charge density of the cation... the greater the polarising power and the more acidic the solution

5. THE AQUEOUS CHEMISTRY OF IONS Theoryaqueous metal ions attract water molecules many have six water molecules surrounding them these are known as hexaaqua ions they are octahedral in shape water acts as a Lewis Base – a lone pair donor water forms a co-ordinate bond to the metal ion metal ions accept the lone pair - Lewis Acids

6. THE AQUEOUS CHEMISTRY OF IONS Theoryaqueous metal ions attract water molecules many have six water molecules surrounding them these are known as hexaaqua ions they are octahedral in shape water acts as a Lewis Base – a lone pair donor water forms a co-ordinate bond to the metal ion metal ions accept the lone pair - Lewis Acids Acidityas charge density increases, the cation has a greater attraction for water the attraction extends to the shared pair of electrons in water’s O-H bonds the electron pair is pulled towards the O, making the bond more polar this makes the H more acidic (more  +) it can then be removed by solvent water molecules to form H 3 O + (aq).

7. HYDROLYSIS - EQUATIONS M 2+ ions [M(H 2 O) 6 ] 2+ (aq) + H 2 O(l) [M(H 2 O) 5 (OH)] + (aq) + H 3 O + (aq) the resulting solution will now be acidic as there are more protons in the water this reaction is known as hydrolysis - the water causes the substance to split up Stronger bases (e.g. CO 3 2-, NH 3 and OH¯ ) can remove further protons...

8. HYDROLYSIS - EQUATIONS M 3+ ions [M(H 2 O) 6 ] 3+ (aq) + H 2 O(l) [M(H 2 O) 5 (OH)] 2+ (aq) + H 3 O + (aq) the resulting solution will also be acidic as there are more protons in the water this SOLUTION IS MORE ACIDIC due to the greater charge density of 3+ ions Stronger bases (e.g. CO 3 2-, NH 3 and OH¯ ) can remove further protons...

9. HYDROLYSIS OF HEXAAQUA IONS Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H 2 O) 6 ] 2+ (aq) [M(OH)(H 2 O) 5 ] + (aq) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 3 (H 2 O) 3 ]¯(aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 5 (H 2 O)] 3- (aq)[M(OH) 6 ] 4- (aq) When sufficient protons have been removed the complex becomes neutral and precipitation of a hydroxide or carbonate occurs. e.g.M 2+ ions[M(H 2 O) 4 (OH) 2 ](s) orM(OH) 2 M 3+ ions[M(H 2 O) 3 (OH) 3 ](s) orM(OH) 3 OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H +

10. HYDROLYSIS OF HEXAAQUA IONS Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H 2 O) 6 ] 2+ (aq) [M(OH)(H 2 O) 5 ] + (aq) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 3 (H 2 O) 3 ]¯(aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 5 (H 2 O)] 3- (aq)[M(OH) 6 ] 4- (aq) In some cases, if the base is strong, further protons are removed and the precipitate dissolves as soluble anionic complexes such as [M(OH) 6 ] 3- are formed. Very weak bases H 2 Oremove few protons Weak bases NH 3, CO 3 2- remove protons until precipitation Strong bases OH¯ can remove all the protons OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H + Precipitated

11. HYDROLYSIS OF HEXAAQUA IONS Lewis bases can attack the co-ordinated water molecules. Theoretically, a proton can be removed from each water molecule turning the water from a neutral molecule to a negatively charged hydroxide ion. This affects the overall charge on the complex ion. [M(H 2 O) 6 ] 2+ (aq) [M(OH)(H 2 O) 5 ] + (aq) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 2 (H 2 O) 4 ](s) [M(OH) 3 (H 2 O) 3 ]¯(aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 4 (H 2 O) 2 ] 2- (aq) [M(OH) 5 (H 2 O)] 3- (aq)[M(OH) 6 ] 4- (aq) AMPHOTERIC CHARACTER Metal ions of 3 + charge have a high charge density and their hydroxides can dissolve in both acid and alkali. [M(H 2 O) 6 ] 3+ (aq) [M(OH) 3 (H 2 O) 3 ](s) [M(OH) 6 ] 3- (aq) OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H + OH¯ H + Precipitated OH¯H+H+ Insoluble Soluble

12. REACTION TYPES The examples aim to show typical properties of transition metals and their compounds. One typical properties of transition elements is their ability to form complex ions. Complex ions consist of a central metal ion surrounded by co-ordinated ions or molecules known as ligands. This can lead to changes in... colour co-ordination number shape stability to oxidation or reduction Reaction types ACID-BASE LIGAND SUBSTITUTION PRECIPITATION REDOX A-B LS OX Ppt RED REDOX

13. REACTION TYPES The examples aim to show typical properties of transition metals and their compounds. LOOK FOR... substitution reactions of complex ions variation in oxidation state of transition metals the effect of ligands on co-ordination number and shape increased acidity of M 3+ over M 2+ due to the increased charge density differences in reactivity of M 3+ and M 2+ ions with OH¯ and NH 3 the reason why M 3+ ions don’t form carbonates amphoteric character in some metal hydroxides (Al 3+ and Cr 3+ ) the effect a ligand has on the stability of a particular oxidation state

14. REACTIONS OF COBALT(II) aqueous solutions contain the pink, octahedral hexaaquacobalt(II) ion hexaaqua ions can also be present in solid samples of the hydrated salts as a 2+ ion, the solutions are weakly acidic but protons can be removed by bases... OH¯[Co(H 2 O) 6 ] 2+ (aq) + 2OH¯ (aq) ——> [Co(OH) 2 (H 2 O) 4 ] (s) + 2H 2 O (l) pink, octahedral blue / pink ppt. soluble in XS NaOH ALL hexaaqua ions precipitate a hydroxide with OH¯(aq). Some re-dissolve in excess NaOH A-B

15. REACTIONS OF COBALT(II) NH 3 [Co(H 2 O) 6 ] 2+ (aq) + 2NH 3 (aq) ——> [Co(OH) 2 (H 2 O) 4 ] (s) + 2NH 4 + (aq) ALL hexaaqua ions precipitate a hydroxide with NH 3 (aq). It removes protons A-B

16. REACTIONS OF COBALT(II) NH 3 [Co(H 2 O) 6 ] 2+ (aq) + 2NH 3 (aq) ——> [Co(OH) 2 (H 2 O) 4 ] (s) + 2NH 4 + (aq) ALL hexaaqua ions precipitate a hydroxide with NH 3 (aq). It removes protons Some hydroxides redissolve in excess NH 3 (aq) as ammonia substitutes as a ligand. [Co(OH) 2 (H 2 O) 4 ] (s) + 6NH 3 (aq) ——> [Co(NH 3 ) 6 ] 2+ (aq) + 4H 2 O (l) + 2OH¯ (aq) A-B LS

17. REACTIONS OF COBALT(II) NH 3 [Co(H 2 O) 6 ] 2+ (aq) + 2NH 3 (aq) ——> [Co(OH) 2 (H 2 O) 4 ] (s) + 2NH 4 + (aq) ALL hexaaqua ions precipitate a hydroxide with NH 3 (aq). It removes protons Some hydroxides redissolve in excess NH 3 (aq) as ammonia substitutes as a ligand. [Co(OH) 2 (H 2 O) 4 ] (s) + 6NH 3 (aq) ——> [Co(NH 3 ) 6 ] 2+ (aq) + 4H 2 O (l) + 2OH¯ (aq) but... ammonia ligands make the Co(II) state unstable. Air oxidises Co(II) to Co(III) [Co(NH 3 ) 6 ] 2+ (aq) ——> [Co(NH 3 ) 6 ] 3+ (aq) + e¯ yellow / brown octahedral red / brown octahedral A-B LS OX

18. REACTIONS OF COBALT(II) CO 3 2- [Co(H 2 O) 6 ] 2+ (aq) + CO 3 2- (aq) ——> CoCO 3 (s) + 6H 2 O (l) mauve ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don’t. Ppt

19. REACTIONS OF COBALT(II) CO 3 2- [Co(H 2 O) 6 ] 2+ (aq) + CO 3 2- (aq) ——> CoCO 3 (s) + 6H 2 O (l) mauve ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don’t. Cl¯[Co(H 2 O) 6 ] 2+ (aq) + 4Cl¯ (aq) ——> [CoCl 4 ] 2- (aq) + 6H 2 O (l) pink, octahedral blue, tetrahedral Cl¯ ligands are larger than H 2 O Cl¯ ligands are negatively charged - H 2 O ligands are neutral the complex is more stable if tetrahedral - less repulsion between ligands adding excess water reverses the reaction LS Ppt

20. REACTIONS OF COPPER(II) Aqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH¯[Cu(H 2 O) 6 ] 2+ (aq) + 2OH¯ (aq) ——> [Cu(OH) 2 (H 2 O) 4 ] (s) + 2H 2 O (l) blue, octahedral pale blue ppt. insoluble in XS NaOH A-B

21. REACTIONS OF COPPER(II) Aqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH¯[Cu(H 2 O) 6 ] 2+ (aq) + 2OH¯ (aq) ——> [Cu(OH) 2 (H 2 O) 4 ] (s) + 2H 2 O (l) blue, octahedral pale blue ppt. insoluble in XS NaOH NH 3 [Cu(H 2 O) 6 ] 2+ (aq) + 2NH 3 (aq) ——> [Cu(OH) 2 (H 2 O) 4 ] (s) + 2NH 4 + (aq) blue ppt. soluble in excess NH 3 then [Cu(OH) 2 (H 2 O) 4 ] (s) + 4NH 3 (aq) ——> [Cu(NH 3 ) 4 (H 2 O) 2 ] 2+ (aq) + 2H 2 O (l) + 2OH¯ (aq) royal blue NOTE THE FORMULA A-B LS

22. REACTIONS OF COPPER(II) Aqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ion Most substitution reactions are similar to cobalt(II). OH¯[Cu(H 2 O) 6 ] 2+ (aq) + 2OH¯ (aq) ——> [Cu(OH) 2 (H 2 O) 4 ] (s) + 2H 2 O (l) blue, octahedral pale blue ppt. insoluble in XS NaOH NH 3 [Cu(H 2 O) 6 ] 2+ (aq) + 2NH 3 (aq) ——> [Cu(OH) 2 (H 2 O) 4 ] (s) + 2NH 4 + (aq) blue ppt. soluble in excess NH 3 then [Cu(OH) 2 (H 2 O) 4 ] (s) + 4NH 3 (aq) ——> [Cu(NH 3 ) 4 (H 2 O) 2 ] 2+ (aq) + 2H 2 O (l) + 2OH¯ (aq) royal blue NOTE THE FORMULA CO 3 2- [Cu(H 2 O) 6 ] 2+ (aq) + CO 3 2- (aq) ——> CuCO 3 (s) + 6H 2 O (l) blue ppt. A-B LS Ppt

23. REACTIONS OF COPPER(II) Cl¯[Cu(H 2 O) 6 ] 2+ (aq) + 4Cl¯ (aq) ——> [CuCl 4 ] 2- (aq) + 6H 2 O (l) yellow, tetrahedral Cl¯ ligands are larger than H 2 O and are charged the complex is more stable if the shape changes to tetrahedral adding excess water reverses the reaction I¯ 2Cu 2+ (aq) + 4I¯ (aq) ——> 2CuI (s) + I 2 (aq) off - white ppt. a redox reaction used in the volumetric analysis of copper using sodium thiosulphate LS REDOX

24. REACTIONS OF COPPER(I) The aqueous chemistry of copper(I) is unstable compared to copper(0) and copper (II). Cu + (aq) + e¯ ——> Cu (s) E° = + 0.52 V Cu 2+ (aq) + e¯ ——> Cu + (aq) E° = + 0.15 V subtracting2Cu + (aq) ——> Cu (s) + Cu 2+ (aq) E° = + 0.37 V This is an example of DISPROPORTIONATION where one species is simultaneously oxidised and reduced to more stable forms. This explains why the aqueous chemistry of copper(I) is very limited. Copper(I) can be stabilised by formation of complexes.

25. REACTIONS OF CHROMIUM(III) Chromium(III) ions are typical of M 3+ ions in this block. Aqueous solutions contain the violet, octahedral hexaaquachromium(III) ion. OH¯[Cr(H 2 O) 6 ] 3+ (aq) + 3OH¯ (aq) ——> [Cr(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) violet, octahedral green ppt. soluble in XS NaOH As with all hydroxides the precipitate reacts with acid [Cr(OH) 3 (H 2 O) 3 ] (s) + 3H + (aq) ——> [Cr(H 2 O) 6 ] 3+ (aq) being a 3+ hydroxide it is AMPHOTERIC as it dissolves in excess alkali [Cr(OH) 3 (H 2 O) 3 ] (s) + 3OH¯ (aq) ——> [Cr(OH) 6 ] 3- (aq) + 3H 2 O (l) green, octahedral A-B

26. REACTIONS OF CHROMIUM(III) CO 3 2- 2 [Cr(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Cr(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) + 3CO 2 (g) The carbonate is not precipitated but the hydroxide is. high charge density of M 3+ makes the solution too acidic to form the carbonate CARBON DIOXIDE IS EVOLVED. A-B

27. REACTIONS OF CHROMIUM(III) CO 3 2- 2 [Cr(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Cr(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) + 3CO 2 (g) The carbonate is not precipitated but the hydroxide is. high charge density of M 3+ makes the solution too acidic to form the carbonate CARBON DIOXIDE IS EVOLVED. NH 3 [Cr(H 2 O) 6 ] 3+ (aq) + 3NH 3 (aq) ——> [Cr(OH) 3 (H 2 O) 3 ] (s) + 3NH 4 + (aq) green ppt. soluble in XS NH 3 With EXCESS AMMONIA, the precipitate redissolves [Cr(OH) 3 (H 2 O) 3 ] (s) + 6NH 3 (aq) ——> [Cr(NH 3 ) 6 ] 3+ (aq) + 3H 2 O (l) + 3OH¯ (aq) LS A-B

28. REACTIONS OF CHROMIUM(III) Oxidation In the presence of alkali, Cr(III) is unstable and can be oxidised to Cr(VI) 2Cr 3+ (aq) + 3H 2 O 2 (l) + 10OH¯(aq) ——> 2CrO 4 2- (aq) + 8H 2 O(l) green yellow Acidification of the yellow chromate will produce the orange dichromate(VI) ion Reduction Chromium(III) can be reduced to the less stable chromium(II) by zinc in acid 2 [Cr(H 2 O) 6 ] 3+ (aq) + Zn(s) ——> 2 [Cr(H 2 O) 6 ] 2+ (aq) + Zn 2+ (aq) green blue OX RED

29. REACTIONS OF CHROMIUM(VI) Occurrencedichromate (VI)Cr 2 O 7 2- orange chromate (VI) CrO 4 2- yellow Interconversiondichromate is stable in acid solution chromate is stable in alkaline solution in alkali Cr 2 O 7 2- (aq) + 2OH¯ (aq) 2CrO 4 2- (aq) + H 2 O (l) in acid 2CrO 4 2- (aq) + 2H + (aq) Cr 2 O 7 2- (aq) + H 2 O (l)

30. CONTENTS OXIDATION REACTIONS OF CHROMIUM(VI) Being in the highest oxidation state (+6), chromium(VI) will be an oxidising agent. In acidic solution, dichromate is widely used in both organic (oxidation of alcohols) and inorganic chemistry. It can also be used as a volumetric reagent but with special indicators as its colour change (orange to green) makes the end point hard to observe. Cr 2 O 7 2- (aq) + 14H + (aq) + 6e¯ ——> 2Cr 3+ (aq) + 7H 2 O (l) [ E° = +1.33 V ] orange green Its E° value is lower than that of Cl 2 (1.36V) so can be used in the presence of Cl¯ ions MnO 4 ¯ (E° = 1.52V) oxidises chloride in HCl so must be acidified with sulphuric acid chromium(VI) can be reduced back to chromium(III) using zinc in acid solution

31. REACTIONS OF MANGANESE(VII) in its highest oxidation state therefore Mn(VII) will be an oxidising agent occurs in the purple, tetraoxomanganate(VII) (permanganate) ion (MnO 4 ¯) acts as an oxidising agent in acidic or alkaline solution acidic MnO 4 ¯ (aq) + 8H + (aq) + 5e¯ ——> Mn 2+ (aq) + 4H 2 O (l) E° = + 1.52 V N.B. Acidify with dilute H 2 SO 4 NOT dilute HCl alkaline MnO 4 ¯ (aq) + 2H 2 O (l) + 3e¯ ——> MnO 2 (s) + 4OH¯ (aq) E° = + 0.59 V

32. VOLUMETRIC USE OF MANGANATE(VII) Potassium manganate(VII) in acidic (H 2 SO 4 ) solution is extremely useful for carrying out redox volumetric analysis. MnO 4 ¯ (aq) + 8H + (aq) + 5e¯ ——> Mn 2+ (aq) + 4H 2 O (l) E° = + 1.52 V It must be acidified with dilute sulphuric acid as MnO 4 ¯ is powerful enough to oxidise the chloride ions in hydrochloric acid. It is used to estimate iron(II), hydrogen peroxide, ethanedioic (oxalic) acid and ethanedioate (oxalate) ions. The last two titrations are carried out above 60°C due to the slow rate of reaction. No indicator is required; the end point being the first sign of a permanent pale pink colour. Iron(II)MnO 4 ¯ (aq) + 8H + (aq) + 5Fe 2+ (aq) ——> Mn 2+ (aq) + 5Fe 3+ (aq) + 4H 2 O (l) this means that moles of Fe 2+ = 5 moles of MnO 4 ¯ 1

33. REACTIONS OF IRON(II) When iron reacts with acids it gives rise to iron(II) (ferrous) salts. Aqueous solutions of such salts contain the pale green, octahedral hexaaquairon(II) ion OH¯[Fe(H 2 O) 6 ] 2+ (aq) + 2OH¯ (aq) ——> [Fe(OH) 2 (H 2 O) 4 ] (s) + 2H 2 O (l) pale green dirty green ppt. it only re-dissolves in very conc. OH¯ but... it slowly turns a rusty brown colour due to oxidation by air to iron(III) increasing the pH renders iron(II) unstable. Fe(OH) 2 (s) + OH¯(aq) ——> Fe(OH) 3 (s) + e¯ dirty green rusty brown NH 3 Iron(II) hydroxide precipitated, insoluble in excess ammonia CO 3 2- Off-white coloured iron(II) carbonate, FeCO 3, precipitated A-B OX Ppt

34. CONTENTS REACTIONS OF IRON(II) VolumetricIron(II) can be analysed by titration with potassium manganate(VII) in acidic (H 2 SO 4 ) solution. No indicator is required. MnO 4 ¯ (aq) + 8H + (aq) + 5Fe 2+ (aq) ——> Mn 2+ (aq) + 5Fe 3+ (aq) + 4H 2 O (l) this means that moles of Fe2+ =5 moles of MnO 4 ¯ 1

35. REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯[Fe(H 2 O) 6 ] 3+ (aq) + 3OH¯ (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) yellow rusty-brown ppt. insoluble in XS A-B

36. REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯[Fe(H 2 O) 6 ] 3+ (aq) + 3OH¯ (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) yellow rusty-brown ppt. insoluble in XS CO 3 2- 2[Fe(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) + 3CO 2 (g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M 3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. A-B

37. REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯[Fe(H 2 O) 6 ] 3+ (aq) + 3OH¯ (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) yellow rusty-brown ppt. insoluble in XS CO 3 2- 2[Fe(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) + 3CO 2 (g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M 3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. NH 3 [Fe(H 2 O) 6 ] 3+ (aq) + 3NH 3 (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3NH 4 + (aq) rusty-brown ppt. insoluble in XS A-B

38. REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯[Fe(H 2 O) 6 ] 3+ (aq) + 3OH¯ (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) yellow rusty-brown ppt. insoluble in XS CO 3 2- 2[Fe(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Fe(OH) 3 (H 2 O) 3 ] (s) + 3H 2 O (l) + 3CO 2 (g) rusty-brown ppt. The carbonate is not precipitated but the hydroxide is; the high charge density of M 3+ makes the solution too acidic to form a carbonate CARBON DIOXIDE EVOLVED. NH 3 [Fe(H 2 O) 6 ] 3+ (aq) + 3NH 3 (aq) ——> [Fe(OH) 3 (H 2 O) 3 ] (s) + 3NH 4 + (aq) rusty-brown ppt. insoluble in XS SCN¯[Fe(H 2 O) 6 ] 3+ (aq) + SCN¯ (aq) ——> [Fe(SCN)(H 2 O) 5 ] 2+ (aq) + H 2 O (l) blood-red colour Very sensitive; BLOOD RED COLOUR confirms Fe(III). No reaction with Fe(II) A-B LS A-B

39. REACTIONS OF SILVER(I) aqueous solutions contains the colourless, linear, diammine silver(I) ion formed when silver halides dissolve in ammonia eg AgCl (s) + 2NH 3 (aq) ——> [Ag(NH 3 ) 2 ] + (aq) + Cl¯ (aq) [Ag(SO 3 ) 2 ] 3- Formed when silver salts are dissolved in sodium thiosulphate "hypo" solution. Important in photographic fixing. Any silver bromide not exposed to light is dissolved away leaving the black image of silver as the negative. AgBr + 2S 2 O 3 2- ——> [Ag(S 2 O 3 ) 2 ] 3- + Br¯ [Ag(CN) 2 ]¯ Formed when silver salts are dissolved in sodium or potassium cyanide the solution used for silver electroplating [Ag(NH 3 ) 2 ] + Used in Tollen’s reagent (SILVER MIRROR TEST) Tollen’s reagent is used to differentiate between aldehydes and ketones. Aldehydes produce a silver mirror on the inside of the test tube Formed when silver halides dissolve in ammonia - TEST FOR HALIDES

40. REACTIONS OF VANADIUM Reduction using zinc in acidic solution shows the various oxidation states of vanadium. Vanadium(V)VO 2+ (aq) + 2H + (aq) + e¯ ——> VO 2+ (aq) + H 2 O(l) yellow blue Vanadium(IV)VO 2+ (aq) + 2H + (aq) + e¯ ——> V 3+ (aq) + H 2 O(l) blue blue/green Vanadium(III)V 3+ (aq) + e¯ ——> V 2+ (aq) blue/green lavender

41. OXIDATION & REDUCTION - A SUMMARY Oxidation complex transition metal ions are stable in acid solution complex ions tend to be less stable in alkaline solution in alkaline conditions they form neutral hydroxides and/or anionic complexes it is easier to remove electrons from neutral or negatively charged species alkaline conditions are usually required e.g. Fe(OH) 2 (s) + OH¯ (aq) ——> Fe(OH) 3 (s) + e¯ Co(OH) 2 (s) + OH¯ (aq) ——> Co(OH) 3 (s) + e¯ 2Cr 3+ (aq) + 3H 2 O 2 (l) + 10OH¯ (aq) ——> 2CrO 4 2- (aq) + 8H 2 O (l) Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions [Co(NH 3 ) 6 ] 2+ (aq) ——> [Co(NH 3 ) 6 ] 3+ (aq) + e¯

42. OXIDATION & REDUCTION - A SUMMARY Oxidation complex transition metal ions are stable in acid solution complex ions tend to be less stable in alkaline solution in alkaline conditions they form neutral hydroxides and/or anionic complexes it is easier to remove electrons from neutral or negatively charged species alkaline conditions are usually required e.g. Fe(OH) 2 (s) + OH¯ (aq) ——> Fe(OH) 3 (s) + e¯ Co(OH) 2 (s) + OH¯ (aq) ——> Co(OH) 3 (s) + e¯ 2Cr 3+ (aq) + 3H 2 O 2 (l) + 10OH¯ (aq) ——> 2CrO 4 2- (aq) + 8H 2 O (l) Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions [Co(NH 3 ) 6 ] 2+ (aq) ——> [Co(NH 3 ) 6 ] 3+ (aq) + e¯ Reduction Zinc metal is used to reduce transition metal ions to lower oxidation states It acts in acid solution as follows...Zn (s) ——> Zn 2+ (aq) + 2e¯ e.g. it reduces iron(III) to iron(II) vanadium(V) to vanadium (IV) to vanadium(III)

43. REACTIONS OF ALUMINIUM aluminium is not a transition metal as it doesn’t make use of d orbitals BUT, due to a high charge density, aluminium ions behave as typical M 3+ ions aqueous solutions contain the colourless, octahedral hexaaquaaluminium(III) ion OH¯[Al(H 2 O) 6 ] 3+ (aq) + 3OH¯(aq) ——> [Al(OH) 3 (H 2 O 3 ](s) + 3H 2 O(l) colourless, octahedral white ppt. soluble in XS NaOH As with all hydroxides the precipitate reacts with acid [Al(OH) 3 (H 2 O) 3 ](s) + 3H + (aq) ——> [Al(H 2 O) 6 ] 3+ (aq) being a 3+ hydroxide it is AMPHOTERIC and dissolves in excess alkali [Al(OH) 3 (H 2 O) 3 ](s) + 3OH¯(aq) ——> [Al(OH) 6 ] 3- (aq) + 3H 2 O(l) colourless, octahedral CO 3 2- 2 [Al(H 2 O) 6 ] 3+ (aq) + 3CO 3 2- (aq) ——> 2[Al(OH) 3 (H 2 O) 3 ](s) + 3H 2 O(l) + 3CO 2 (g) As with 3+ ions, the carbonate is not precipitated but the hydroxide is. NH 3 [Al(H 2 O) 6 ] 3+ (aq) + 3NH 3 (aq) ——> [Al(OH) 3 (H 2 O) 3 ](s) + 3NH 4 + (aq) white ppt. insoluble in XS NH 3 A-B

44. © JONATHAN HOPTON & KNOCKHARDY PUBLISHING THE END AN INTRODUCTION TO TRANSITION METAL CHEMISTRY