AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING

AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING 2015 SPECIFICATIONS

KNOCKHARDY PUBLISHING TRANSITION METALS INTRODUCTION This Powerpoint show is one of several produced to help students understand selected topics at AS and A 2 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 A 2 topics, are available from the KNOCKHARDY SCIENCE WEBSITE at. . . www. knockhardy. org. uk/sci. htm Navigation is achieved by. . . either or clicking on the grey arrows at the foot of each page using the left and right arrow keys on the keyboard

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

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 cation Na Mg Al charge 1+ 2+ 3+ ionic radius reaction with water / p. H of chloride 0. 095 nm 0. 065 nm 0. 050 nm the greater charge density of the cation. . . the greater the polarising power the more acidic the solution and

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 cation Na Mg Al charge 1+ 2+ 3+ ionic radius reaction with water / p. H of chloride 0. 095 nm 0. 065 nm 0. 050 nm dissolves slight hydrolysis vigorous hydrolysis 7 the greater charge density of the cation. . . the greater the polarising power the more acidic the solution and

THE AQUEOUS CHEMISTRY OF IONS Theory aqueous 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 ions accept the lone pair - Lewis Acids

THE AQUEOUS CHEMISTRY OF IONS Theory aqueous 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 ions accept the lone pair - Lewis Acidity as 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 d+) it can then be removed by solvent water molecules to form H 3 O+(aq).

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 32 - , NH 3 and OH¯ ) can remove further protons. . .

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 32 - , NH 3 and OH¯ ) can remove further protons. . .

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)2(H 2 O)4](s) [M(OH)4(H 2 O)2 ]2 -(aq) OH¯ H+ [M(OH)(H 2 O)5]+(aq) [M(OH)3(H 2 O)3]¯(aq) [M(OH)5(H 2 O)]3 -(aq) OH¯ H+ [M(OH)2(H 2 O)4](s) [M(OH)4(H 2 O)2]2 -(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 3+ ions [M(H 2 O)4(OH)2](s) [M(H 2 O)3(OH)3](s) or or M(OH)2 M(OH)3

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)2(H 2 O)4](s) [M(OH)4(H 2 O)2 ]2 -(aq) OH¯ H+ [M(OH)(H 2 O)5]+(aq) [M(OH)3(H 2 O)3]¯(aq) [M(OH)5(H 2 O)]3 -(aq) OH¯ H+ [M(OH)2(H 2 O)4](s) Precipitated [M(OH)4(H 2 O)2]2 -(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 Weak bases Strong bases H 2 O NH 3, CO 32 OH¯ remove few protons remove protons until precipitation can remove all the protons

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)2(H 2 O)4](s) [M(OH)4(H 2 O)2 ]2 -(aq) OH¯ H+ [M(OH)(H 2 O)5]+(aq) [M(OH)3(H 2 O)3]¯(aq) [M(OH)5(H 2 O)]3 -(aq) OH¯ H+ [M(OH)2(H 2 O)4](s) Precipitated [M(OH)4(H 2 O)2]2 -(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) Soluble H+ [M(OH)3(H 2 O)3](s) Insoluble OH¯ [M(OH)6]3 -(aq) Soluble

STABILITY CONSTANTS Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions.

STABILITY CONSTANTS Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction [M(H 2 O)6]2+(aq) + 6 X¯(aq) [MX 6]4–(aq) + 6 H 2 O(l)

STABILITY CONSTANTS Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction the expression for the stability constant is [M(H 2 O)6]2+(aq) + 6 X¯(aq) Kstab = [MX 6]4–(aq) + 6 H 2 O(l) [ [MX 64–](aq) ] [ [M(H 2 O)6]2+(aq) ] [ X¯(aq) ]6

STABILITY CONSTANTS Definition The stability constant, Kstab, of a complex ion is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions. In the reaction the expression for the stability constant is [M(H 2 O)6]2+(aq) + 6 X¯(aq) Kstab = [MX 6]4–(aq) + 6 H 2 O(l) [ [MX 64–](aq) ] [ [M(H 2 O)6]2+(aq) ] [ X¯(aq) ]6 The concentration of X¯(aq) appears to the power of 6 because there are six of the ions in the equation. Note that the water isn’t included; it is in such overwhelming quantity that its concentration can be regarded as ‘constant’.

STABILITY CONSTANTS Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… [Co(H 2 O)6]2+(aq) + NH 3(aq) [Co(NH 3)(H 2 O)5]2+(aq) + H 2 O(l) Kstab / dm 3 mol-1 K 1 = 1. 02 x 10 -2 [Co(NH 3)(H 2 O)5]2+(aq) + NH 3(aq) [Co(NH 3)2(H 2 O)4]2+(aq) + H 2 O(l) K 2 = 3. 09 x [Co(NH 3)2(H 2 O)4]2+(aq) + NH 3(aq) [Co(NH 3)3(H 2 O)3]2+(aq) + H 2 O(l) K 3 = 1. 17 x [Co(NH 3)3(H 2 O)3]2+(aq) + NH 3(aq) [Co(NH 3)4(H 2 O)2]2+(aq) + H 2 O(l) K 4 = 2. 29 x [Co(NH 3)4(H 2 O)2]2+(aq) + NH 3(aq) [Co(NH 3)5(H 2 O)]2+(aq) 10 -2 10 -1 + H 2 O(l) K 5 = 8. 70 x 10 -1 etc

STABILITY CONSTANTS Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… [Co(H 2 O)6]2+(aq) + NH 3(aq) [Co(NH 3)(H 2 O)5]2+(aq) + H 2 O(l) Kstab / dm 3 mol-1 K 1 = 1. 02 x 10 -2 [Co(NH 3)(H 2 O)5]2+(aq) + NH 3(aq) [Co(NH 3)2(H 2 O)4]2+(aq) + H 2 O(l) K 2 = 3. 09 x [Co(NH 3)2(H 2 O)4]2+(aq) + NH 3(aq) [Co(NH 3)3(H 2 O)3]2+(aq) + H 2 O(l) K 3 = 1. 17 x [Co(NH 3)3(H 2 O)3]2+(aq) + NH 3(aq) [Co(NH 3)4(H 2 O)2]2+(aq) + H 2 O(l) K 4 = 2. 29 x [Co(NH 3)4(H 2 O)2]2+(aq) + NH 3(aq) [Co(NH 3)5(H 2 O)]2+(aq) 10 -2 10 -1 + H 2 O(l) K 5 = 8. 70 x 10 -1 etc The overall stability constant is simply the equilibrium constant for the total reaction. It is found by multiplying the individual stability constants. . . k 1 x k 2 x k 3 x k 4. . . etc Kstab or p. Kstab? For an easier comparison, the expression p. Kstab is often used… p. Kstab = -log 10 Kstab

STABILITY CONSTANTS Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… [Co(H 2 O)6]2+(aq) + NH 3(aq) [Co(NH 3)(H 2 O)5]2+(aq) + H 2 O(l) Kstab / dm 3 mol-1 K 1 = 1. 02 x 10 -2 [Co(NH 3)(H 2 O)5]2+(aq) + NH 3(aq) [Co(NH 3)2(H 2 O)4]2+(aq) + H 2 O(l) K 2 = 3. 09 x [Co(NH 3)2(H 2 O)4]2+(aq) + NH 3(aq) [Co(NH 3)3(H 2 O)3]2+(aq) + H 2 O(l) K 3 = 1. 17 x [Co(NH 3)3(H 2 O)3]2+(aq) + NH 3(aq) [Co(NH 3)4(H 2 O)2]2+(aq) + H 2 O(l) K 4 = 2. 29 x [Co(NH 3)4(H 2 O)2]2+(aq) + NH 3(aq) [Co(NH 3)5(H 2 O)]2+(aq) 10 -2 10 -1 • • + H 2 O(l) K 5 = 8. 70 x 10 -1 Summary The larger the stability constant, the further the reaction lies to the right Complex ions with large stability constants are more stable Stability constants are often given as p. Kstab Complex ions with smaller p. Kstab values are more stable etc

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 • shape Reaction types • co-ordination number • stability to oxidation or reduction ACID-BASE A-B LIGAND SUBSTITUTION LS PRECIPITATION Ppt REDOX REDOX

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

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) + 2 OH¯(aq) ——> [Co(OH)2(H 2 O)4](s) + 2 H 2 O(l) pink, octahedral blue / pink ppt. soluble in XS Na. OH ALL hexaaqua ions precipitate a hydroxide with OH¯(aq). Some re-dissolve in excess Na. OH A-B

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

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

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

REACTIONS OF COBALT(II) CO 32 - [Co(H 2 O)6]2+(aq) + CO 32 -(aq) ——> Co. CO 3(s) + 6 H 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

REACTIONS OF COBALT(II) CO 32 - [Co(H 2 O)6]2+(aq) + CO 32 -(aq) ——> Co. CO 3(s) + 6 H 2 O(l) Ppt 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) + 4 Cl¯(aq) pink, octahedral • • ——> [Co. Cl 4]2 -(aq) + 6 H 2 O(l) blue, tetrahedral LS 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

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) + 2 OH¯(aq) ——> [Cu(OH)2(H 2 O)4](s) + 2 H 2 O(l) blue, octahedral pale blue ppt. insoluble in XS Na. OH A-B

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) + 2 OH¯(aq) ——> [Cu(OH)2(H 2 O)4](s) + 2 H 2 O(l) blue, octahedral NH 3 A-B pale blue ppt. insoluble in XS Na. OH [Cu(H 2 O)6]2+(aq) + 2 NH 3(aq) ——> [Cu(OH)2(H 2 O)4](s) + 2 NH 4+(aq) A-B blue ppt. soluble in excess NH 3 then [Cu(OH)2(H 2 O)4](s) + 4 NH 3(aq) ——> [Cu(NH 3)4(H 2 O)2]2+(aq) + 2 H 2 O(l) + 2 OH¯(aq) royal blue NOTE THE FORMULA LS

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) + 2 OH¯(aq) ——> [Cu(OH)2(H 2 O)4](s) + 2 H 2 O(l) blue, octahedral NH 3 A-B pale blue ppt. insoluble in XS Na. OH [Cu(H 2 O)6]2+(aq) + 2 NH 3(aq) ——> [Cu(OH)2(H 2 O)4](s) + 2 NH 4+(aq) A-B blue ppt. soluble in excess NH 3 then [Cu(OH)2(H 2 O)4](s) + 4 NH 3(aq) ——> [Cu(NH 3)4(H 2 O)2]2+(aq) + 2 H 2 O(l) + 2 OH¯(aq) royal blue NOTE THE FORMULA CO 32 - [Cu(H 2 O)6]2+(aq) + CO 32 -(aq) ——> Cu. CO 3(s) + 6 H 2 O(l) blue ppt. LS Ppt

REACTIONS OF COPPER(II) Cl¯ [Cu(H 2 O)6]2+(aq) + 4 Cl¯(aq) ——> [Cu. Cl 4]2 -(aq) + 6 H 2 O(l) yellow, tetrahedral LS • 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¯ 2 Cu 2+(aq) + 4 I¯(aq) ——> 2 Cu. I(s) + I 2(aq) off - white ppt. • a redox reaction • used in the volumetric analysis of copper using sodium thiosulphate REDOX

REACTIONS OF COPPER(I) The aqueous chemistry of copper(I) is unstable compared to copper(0) and copper (II). Cu+(aq) + e¯ Cu 2+(aq) + e¯ subtracting ——> Cu(s) ——> Cu+(aq) 2 Cu+(aq) ——> Cu(s) + Cu 2+(aq) E° = + 0. 52 V E° = + 0. 15 V 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.

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) + 3 OH¯(aq) violet, octahedral ——> [Cr(OH)3(H 2 O)3](s) + 3 H 2 O(l) A-B green ppt. soluble in XS Na. OH As with all hydroxides the precipitate reacts with acid [Cr(OH)3(H 2 O)3](s) + 3 H+(aq) ——> [Cr(H 2 O)6]3+(aq) being a 3+ hydroxide it is AMPHOTERIC as it dissolves in excess alkali A-B [Cr(OH)3(H 2 O)3](s) A-B + 3 OH¯(aq) ——> [Cr(OH)6]3 -(aq) green, octahedral + 3 H 2 O(l)

REACTIONS OF CHROMIUM(III) CO 32 - 2 [Cr(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Cr(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 2(g) The carbonate is not precipitated but the hydroxide is. A-B high charge density of M 3+ makes the solution too acidic to form the carbonate CARBON DIOXIDE IS EVOLVED.

REACTIONS OF CHROMIUM(III) CO 32 - 2 [Cr(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Cr(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 2(g) The carbonate is not precipitated but the hydroxide is. A-B 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) + 3 NH 3(aq) ——> [Cr(OH)3(H 2 O)3](s) + 3 NH 4+(aq) A-B green ppt. soluble in XS NH 3 With EXCESS AMMONIA, the precipitate redissolves LS [Cr(OH)3(H 2 O)3](s) + 6 NH 3(aq) ——> [Cr(NH 3)6]3+(aq) + 3 H 2 O(l) + 3 OH¯(aq)

REACTIONS OF CHROMIUM(III) Oxidation In the presence of alkali, Cr(III) is unstable and can be oxidised to Cr(VI) 2 Cr 3+(aq) + 3 H 2 O 2(l) + 10 OH¯(aq) ——> 2 Cr. O 42 -(aq) + 8 H 2 O(l) green OX 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) green ——> 2 [Cr(H 2 O)6]2+(aq) blue + Zn 2+(aq) RED

REACTIONS OF CHROMIUM(VI) Occurrence Interconversion dichromate (VI) Cr 2 O 72 - orange chromate (VI) Cr. O 42 - yellow dichromate is stable in acid solution chromate is stable in alkaline solution in alkali Cr 2 O 72 -(aq) + 2 OH¯(aq) 2 Cr. O 42 -(aq) + H 2 O(l) in acid 2 Cr. O 42 -(aq) + 2 H+(aq) Cr 2 O 72 -(aq) + H 2 O(l)

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 72 -(aq) + 14 H+(aq) + 6 e¯ ——> 2 Cr 3+(aq) + 7 H 2 O(l) orange [ E° = +1. 33 V ] green • Its E° value is lower than that of Cl 2 (1. 36 V) so can be used in the presence of Cl¯ ions • Mn. O 4¯ (E° = 1. 52 V) 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 CONTENTS

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 (Mn. O 4¯) • acts as an oxidising agent in acidic or alkaline solution acidic Mn. O 4¯(aq) + 8 H+(aq) + 5 e¯ ——> Mn 2+(aq) + 4 H 2 O(l) E° = + 1. 52 V N. B. Acidify with dilute H 2 SO 4 NOT dilute HCl alkaline Mn. O 4¯(aq) + 2 H 2 O(l) + 3 e¯ ——> Mn. O 2(s) + 4 OH¯(aq) E° = + 0. 59 V

REACTIONS OF MANGANESE(II) • aqueous solutions contain the very pale ink, octahedral hexaaquamanganese(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¯ [Mn(H 2 O)6]2+(aq) + 2 OH¯(aq) ——> [Mn(OH)2(H 2 O)4](s) + 2 H 2 O(l) pale pink, octahedral off-white ppt. insoluble in XS Na. OH ALL hexaaqua ions precipitate a hydroxide with OH¯(aq). Some re-dissolve in excess Na. OH A-B

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

REACTIONS OF MANGANESE(II) CO 32 - [Mn(H 2 O)6]2+(aq) + CO 32 -(aq) ——> Mn. CO 3(s) + 6 H 2 O(l) off-white ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don’t. Ppt

REACTIONS OF MANGANESE(II) CO 32 - [Mn(H 2 O)6]2+(aq) + CO 32 -(aq) ——> Mn. CO 3(s) + 6 H 2 O(l) Ppt off-white ppt. Hexaaqua ions of metals with charge 2+ precipitate a carbonate but heaxaaqua ions with a 3+ charge don’t. Cl¯ [Mn(H 2 O)6]2+(aq) + 4 Cl¯(aq) pale pink, octahedral ——> [Mn. Cl 4]2 -(aq) + 6 H 2 O(l) tetrahedral LS • 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

VOLUMETRIC USE OF MANGANATE(VII) Potassium manganate(VII) in acidic (H 2 SO 4) solution is extremely useful for carrying out redox volumetric analysis. Mn. O 4¯(aq) + 8 H+(aq) + 5 e¯ ——> Mn 2+(aq) + 4 H 2 O(l) E° = + 1. 52 V It must be acidified with dilute sulphuric acid as Mn. O 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) Mn. O 4¯(aq) + 8 H+(aq) + 5 Fe 2+(aq) ——> Mn 2+(aq) + 5 Fe 3+(aq) + 4 H 2 O(l) this means that moles of Fe 2+ moles of Mn. O 4¯ = 5 1

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) + 2 OH¯(aq) ——> [Fe(OH)2(H 2 O)4](s) pale green + 2 H 2 O(l) A-B 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 p. H renders iron(II) unstable. Fe(OH)2(s) + OH¯(aq) ——> Fe(OH)3(s) + e¯ dirty green rusty brown OX NH 3 Iron(II) hydroxide precipitated, insoluble in excess ammonia A-B CO 32 - Off-white coloured iron(II) carbonate, Fe. CO 3, precipitated Ppt

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

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

REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯ [Fe(H 2 O)6]3+(aq) + 3 OH¯(aq) ——> [Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) yellow CO 32 - A-B rusty-brown ppt. insoluble in XS 2[Fe(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 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

REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯ [Fe(H 2 O)6]3+(aq) + 3 OH¯(aq) ——> [Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) yellow CO 32 - A-B rusty-brown ppt. insoluble in XS 2[Fe(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 2(g) rusty-brown ppt. NH 3 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 [Fe(H 2 O)6]3+(aq) + 3 NH 3(aq) ——> [Fe(OH)3(H 2 O)3](s) A-B + 3 NH 4+(aq) rusty-brown ppt. insoluble in XS

REACTIONS OF IRON(III) Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion OH¯ [Fe(H 2 O)6]3+(aq) + 3 OH¯(aq) ——> [Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) yellow CO 32 - A-B rusty-brown ppt. insoluble in XS 2[Fe(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Fe(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 2(g) rusty-brown ppt. NH 3 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 [Fe(H 2 O)6]3+(aq) + 3 NH 3(aq) ——> [Fe(OH)3(H 2 O)3](s) A-B + 3 NH 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) LS blood-red colour Very sensitive; BLOOD RED COLOUR confirms Fe(III). No reaction with Fe(II)

REACTIONS OF SILVER(I) • aqueous solutions contains the colourless, linear, diammine silver(I) ion • formed when silver halides dissolve in ammonia eg Ag. Cl(s) + 2 NH 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. Ag. Br + 2 S 2 O 32 - ——> [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

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

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 do behave as typical M 3+ ions • aqueous solutions contain the colourless, octahedral hexaaquaaluminium(III) ion OH¯ [Al(H 2 O)6]3+(aq) + 3 OH¯(aq) 3 H 2 O(l) A-B [Al(OH)3(H 2 O)3](s) + 3 H+ (aq) ——> [Al(H 2 O)6]3+(aq) being a 3+ hydroxide it is AMPHOTERIC and dissolves in excess alkali A-B [Al(OH)3(H 2 O)3](s) + 3 OH¯(aq) ——> A-B colourless, octahedral ——> [Al(OH)3(H 2 O 3](s) + white ppt. soluble in XS Na. OH As with all hydroxides the precipitate reacts with acid [Al(OH)6]3 -(aq) + 3 H 2 O(l) colourless, octahedral CO 32 - NH 3 2 [Al(H 2 O)6]3+(aq) + 3 CO 32 -(aq) ——> 2[Al(OH)3(H 2 O)3](s) + 3 H 2 O(l) + 3 CO 2(g) As with 3+ ions, the carbonate is not precipitated but the hydroxide is. A-B [Al(H 2 O)6]3+(aq) + 3 NH 3(aq) A-B ——> [Al(OH)3(H 2 O)3](s) + 3 NH 4+(aq) white ppt. insoluble in XS NH 3

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¯ 2 Cr 3+(aq) + 3 H 2 O 2(l) + 10 OH¯(aq) ——> 2 Cr. O 42 -(aq) + 8 H 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¯

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¯ 2 Cr 3+(aq) + 3 H 2 O 2(l) + 10 OH¯(aq) ——> 2 Cr. O 42 -(aq) + 8 H 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) + 2 e¯ e. g. it reduces iron(III) to iron(II) vanadium(V) to vanadium (IV) to vanadium(III)