1. KNOCKHARDY PUBLISHING
AN INTRODUCTION TO
TRANSITION METAL
COMPLEXES
2008 SPECIFICATIONS
KNOCKHARDY PUBLISHING

2. KNOCKHARDY PUBLISHING
TRANSITION METALS
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...
Navigation is achieved by...
either clicking on the grey arrows at the foot of each page
or using the left and right arrow keys on the keyboard

3. 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 on manganese
Reactions of iron(II)
Reactions of iron(III)
Reactions of silver and vanadium
Reactions of aluminium

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
cation charge ionic radius reaction with water / pH of chloride
Na nm
Mg nm
Al 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 - 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 charge ionic radius reaction with water / pH of chloride
Na nm dissolves
Mg nm slight hydrolysis
Al nm vigorous hydrolysis
the greater charge density of the cation...
the greater the polarising power and
the more acidic the solution

6. 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 ion
metal ions accept the lone pair - Lewis Acids

7. 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 ion
metal ions accept the lone pair - Lewis Acids
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 H3O+(aq).

8. HYDROLYSIS - EQUATIONS
M2+ ions [M(H2O)6]2+(aq) H2O(l) [M(H2O)5(OH)]+(aq) H3O+(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. CO32- , NH3 and OH¯ ) can remove further protons...

9. HYDROLYSIS - EQUATIONS
M3+ ions [M(H2O)6]3+(aq) H2O(l) [M(H2O)5(OH)]2+(aq) H3O+(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. CO32- , NH3 and OH¯ ) can remove further protons...

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(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s)
[M(OH)2(H2O)4](s) [M(OH)3(H2O)3]¯(aq) [M(OH)4(H2O)2]2-(aq)
[M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]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. M2+ ions [M(H2O)4(OH)2](s) or M(OH)2
M3+ ions [M(H2O)3(OH)3](s) or M(OH)3
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+

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(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s)
[M(OH)2(H2O)4](s) [M(OH)3(H2O)3]¯(aq) [M(OH)4(H2O)2]2-(aq)
[M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]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. M2+ ions [M(H2O)4(OH)2](s) or M(OH)2
M3+ ions [M(H2O)3(OH)3](s) or M(OH)3
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+

12. 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(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s)
[M(OH)2(H2O)4](s) [M(OH)3(H2O)3]¯(aq) [M(OH)4(H2O)2]2-(aq)
[M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]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 H2O remove few protons
Weak bases NH3, CO32- remove protons until precipitation
Strong bases OH¯ can remove all the protons
OH¯ H+
OH¯ H+
Precipitated
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+

13. 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(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s)
[M(OH)2(H2O)4](s) [M(OH)3(H2O)3]¯(aq) [M(OH)4(H2O)2]2-(aq)
[M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]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 H2O remove few protons
Weak bases NH3, CO32- remove protons until precipitation
Strong bases OH¯ can remove all the protons
OH¯ H+
OH¯ H+
Precipitated
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+

14. 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(H2O)6]2+(aq) [M(OH)(H2O)5]+(aq) [M(OH)2(H2O)4](s)
[M(OH)2(H2O)4](s) [M(OH)3(H2O)3]¯(aq) [M(OH)4(H2O)2]2-(aq)
[M(OH)4(H2O)2]2-(aq) [M(OH)5(H2O)]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(H2O)6]3+(aq) [M(OH)3(H2O)3](s) [M(OH)6]3-(aq)
OH¯ H+
OH¯ H+
Precipitated
OH¯ H+
OH¯ H+
OH¯ H+
OH¯ H+ H+
OH¯
Soluble
Insoluble
Soluble

15. 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.

16. 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(H2O)6]2+(aq) + 6X¯(aq) [MX6]4–(aq) H2O(l)

17. 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(H2O)6]2+(aq) + 6X¯(aq) [MX6]4–(aq) H2O(l)
the expression for the
stability constant is Kstab = [ [MX64–](aq) ]
[ [M(H2O)6]2+(aq) ] [ X¯(aq) ]6

18. 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(H2O)6]2+(aq) + 6X¯(aq) [MX6]4–(aq) H2O(l)
the expression for the
stability constant is Kstab = [ [MX64–](aq) ]
[ [M(H2O)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’.

19. STABILITY CONSTANTS
Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… Kstab / dm3 mol-1
[Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) H2O(l) K1 = x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) H2O(l) K2 = x 10-2
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) H2O(l) K3 = x 10-1
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) H2O(l) K4 = x 10-1
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) H2O(l) K5 = x etc

20. STABILITY CONSTANTS
Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… Kstab / dm3 mol-1
[Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) H2O(l) K1 = x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) H2O(l) K2 = x 10-2
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) H2O(l) K3 = x 10-1
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) H2O(l) K4 = x 10-1
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) H2O(l) K5 = x etc
The overall stability constant is simply the equilibrium constant for the total reaction.
It is found by multiplying the individual stability constants k1 x k2 x k3 x k etc
Kstab or pKstab? For an easier comparison,
the expression pKstab is often used… pKstab = -log10Kstab

21. STABILITY CONSTANTS
Because ligand exchange involves a series of equilibria, each step in the process has a different stability constant… Kstab / dm3 mol-1
[Co(H2O)6]2+(aq) + NH3(aq) [Co(NH3)(H2O)5]2+(aq) H2O(l) K1 = x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq) [Co(NH3)2(H2O)4]2+(aq) H2O(l) K2 = x 10-2
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq) [Co(NH3)3(H2O)3]2+(aq) H2O(l) K3 = x 10-1
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq) [Co(NH3)4(H2O)2]2+(aq) H2O(l) K4 = x 10-1
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq) [Co(NH3)5(H2O)]2+(aq) H2O(l) K5 = x etc
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 pKstab
• Complex ions with smaller pKstab values are more stable

22. REACTION TYPES • shape • stability to oxidation or reduction
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
Ppt
RED OX
REDOX

23. REACTION TYPES LOOK FOR...
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 M3+ over M2+ due to the increased charge density
differences in reactivity of M3+ and M2+ ions with OH¯ and NH3
the reason why M3+ ions don’t form carbonates
amphoteric character in some metal hydroxides (Al3+ and Cr3+)
the effect a ligand has on the stability of a particular oxidation state

24. 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(H2O)6]2+(aq) OH¯(aq) ——> [Co(OH)2(H2O)4](s) H2O(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

25. REACTIONS OF COBALT(II)
NH3 [Co(H2O)6]2+(aq) + 2NH3(aq) ——> [Co(OH)2(H2O)4](s) + 2NH4+(aq)
ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons
A-B

26. REACTIONS OF COBALT(II)
NH3 [Co(H2O)6]2+(aq) + 2NH3(aq) ——> [Co(OH)2(H2O)4](s) + 2NH4+(aq)
ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons
Some hydroxides redissolve in excess NH3(aq) as ammonia substitutes as a ligand.
[Co(OH)2(H2O)4](s) + 6NH3(aq) ——> [Co(NH3)6]2+(aq) + 4H2O(l) + 2OH¯(aq)
A-B LS

27. REACTIONS OF COBALT(II)
NH3 [Co(H2O)6]2+(aq) + 2NH3(aq) ——> [Co(OH)2(H2O)4](s) + 2NH4+(aq)
ALL hexaaqua ions precipitate a hydroxide with NH3 (aq). It removes protons
Some hydroxides redissolve in excess NH3(aq) as ammonia substitutes as a ligand.
[Co(OH)2(H2O)4](s) + 6NH3(aq) ——> [Co(NH3)6]2+(aq) + 4H2O(l) + 2OH¯(aq)
but ... ammonia ligands make the Co(II) state unstable. Air oxidises Co(II) to Co(III)
[Co(NH3)6]2+(aq) ——> [Co(NH3)6]3+(aq) e¯
yellow / brown octahedral red / brown octahedral
A-B LS OX

28. REACTIONS OF COBALT(II)
CO32- [Co(H2O)6]2+(aq) CO32-(aq) ——> CoCO3(s) H2O(l)
mauve ppt.
Hexaaqua ions of metals with charge 2+ precipitate a carbonate but
heaxaaqua ions with a 3+ charge don’t.
Ppt

29. REACTIONS OF COBALT(II)
CO32- [Co(H2O)6]2+(aq) CO32-(aq) ——> CoCO3(s) H2O(l)
mauve ppt.
Hexaaqua ions of metals with charge 2+ precipitate a carbonate but
heaxaaqua ions with a 3+ charge don’t.
Cl¯ [Co(H2O)6]2+(aq) + 4Cl¯(aq) ——> [CoCl4]2-(aq) H2O(l) pink, octahedral blue, tetrahedral
• Cl¯ ligands are larger than H2O
• Cl¯ ligands are negatively charged - H2O ligands are neutral
• the complex is more stable if tetrahedral - less repulsion between ligands
• adding excess water reverses the reaction
Ppt LS

30. 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(H2O)6]2+(aq) OH¯(aq) ——> [Cu(OH)2(H2O)4](s) H2O(l)
blue, octahedral pale blue ppt.
insoluble in XS NaOH
A-B

31. 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(H2O)6]2+(aq) OH¯(aq) ——> [Cu(OH)2(H2O)4](s) H2O(l)
blue, octahedral pale blue ppt.
insoluble in XS NaOH
NH3 [Cu(H2O)6]2+(aq) NH3(aq) ——> [Cu(OH)2(H2O)4](s) NH4+(aq)
blue ppt. soluble in excess NH3
then [Cu(OH)2(H2O)4](s) + 4NH3(aq) ——> [Cu(NH3)4(H2O)2]2+(aq) + 2H2O(l) + 2OH¯(aq) royal blue NOTE THE FORMULA
A-B
A-B LS

32. 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(H2O)6]2+(aq) OH¯(aq) ——> [Cu(OH)2(H2O)4](s) H2O(l)
blue, octahedral pale blue ppt.
insoluble in XS NaOH
NH3 [Cu(H2O)6]2+(aq) NH3(aq) ——> [Cu(OH)2(H2O)4](s) NH4+(aq)
blue ppt. soluble in excess NH3
then [Cu(OH)2(H2O)4](s) + 4NH3(aq) ——> [Cu(NH3)4(H2O)2]2+(aq) + 2H2O(l) + 2OH¯(aq) royal blue NOTE THE FORMULA
CO32- [Cu(H2O)6]2+(aq) CO32-(aq) ——> CuCO3(s) H2O(l)
blue ppt.
A-B
A-B LS
Ppt

33. REACTIONS OF COPPER(II)
Cl¯ [Cu(H2O)6]2+(aq) + 4Cl¯(aq) ——> [CuCl4]2-(aq) H2O(l) yellow, tetrahedral
• Cl¯ ligands are larger than H2O and are charged
• the complex is more stable if the shape changes to tetrahedral
• adding excess water reverses the reaction
I¯ 2Cu2+(aq) I¯(aq) ——> 2CuI(s) I2(aq)
off - white ppt.
• a redox reaction
• used in the volumetric analysis of copper using sodium thiosulphate LS
REDOX

34. 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° = V
Cu2+(aq) e¯ ——> Cu+(aq) E° = V
subtracting 2Cu+(aq) ——> Cu(s) Cu2+(aq) E° = 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.

35. REACTIONS OF CHROMIUM(III)
Chromium(III) ions are typical of M3+ ions in this block.
Aqueous solutions contain the violet, octahedral hexaaquachromium(III) ion.
OH¯ [Cr(H2O)6]3+(aq) OH¯(aq) ——> [Cr(OH)3(H2O)3](s) H2O(l)
violet, octahedral green ppt. soluble in XS NaOH
As with all hydroxides the precipitate reacts with acid
[Cr(OH)3(H2O)3](s) H+(aq) ——> [Cr(H2O)6]3+(aq)
being a 3+ hydroxide it is AMPHOTERIC as it dissolves in excess alkali
[Cr(OH)3(H2O)3](s) OH¯(aq) ——> [Cr(OH)6]3-(aq) H2O(l)
green, octahedral
A-B
A-B
A-B

36. REACTIONS OF CHROMIUM(III)
CO [Cr(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Cr(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g)
The carbonate is not precipitated but the hydroxide is.
high charge density of M3+ makes the solution too acidic to form the carbonate
CARBON DIOXIDE IS EVOLVED.
A-B

37. REACTIONS OF CHROMIUM(III)
CO [Cr(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Cr(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g)
The carbonate is not precipitated but the hydroxide is.
high charge density of M3+ makes the solution too acidic to form the carbonate
CARBON DIOXIDE IS EVOLVED.
NH3 [Cr(H2O)6]3+(aq) NH3(aq) ——> [Cr(OH)3(H2O)3](s) + 3NH4+(aq)
green ppt. soluble in XS NH3
With EXCESS AMMONIA, the precipitate redissolves
[Cr(OH)3(H2O)3](s) + 6NH3(aq) ——> [Cr(NH3)6]3+(aq) + 3H2O(l) + 3OH¯(aq)
A-B
A-B LS

38. REACTIONS OF CHROMIUM(III)
Oxidation In the presence of alkali, Cr(III) is unstable and can be oxidised to Cr(VI)
2Cr3+(aq) + 3H2O2(l) OH¯(aq) ——> 2CrO42-(aq) H2O(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(H2O)6]3+(aq) Zn(s) ——> 2 [Cr(H2O)6]2+(aq) Zn2+(aq)
green blue OX
RED

39. REACTIONS OF CHROMIUM(VI)
Occurrence dichromate (VI) Cr2O72- orange
chromate (VI) CrO yellow
Interconversion dichromate is stable in acid solution
chromate is stable in alkaline solution
in alkali Cr2O72-(aq) + 2OH¯(aq) CrO42-(aq) H2O(l)
in acid CrO42-(aq) H+(aq) Cr2O72-(aq) H2O(l)

40. 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.
Cr2O72-(aq) H+(aq) e¯ ——> 2Cr3+(aq) H2O(l) [ E° = V ]
orange green
• Its E° value is lower than that of Cl2 (1.36V) so can be used in the presence of Cl¯ ions
• MnO4¯ (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
CONTENTS

41. 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 (MnO4¯)
• acts as an oxidising agent in acidic or alkaline solution
acidic MnO4¯(aq) H+(aq) + 5e¯ ——> Mn2+(aq) H2O(l) E° = V
N.B. Acidify with dilute H2SO4 NOT dilute HCl
alkaline MnO4¯(aq) H2O(l) + 3e¯ ——> MnO2(s) + 4OH¯(aq) E° = V

42. VOLUMETRIC USE OF MANGANATE(VII)
Potassium manganate(VII) in acidic (H2SO4) solution is extremely useful for carrying out redox volumetric analysis.
MnO4¯(aq) H+(aq) + 5e¯ ——> Mn2+(aq) H2O(l) E° = V
It must be acidified with dilute sulphuric acid as MnO4¯ 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) MnO4¯(aq) H+(aq) + 5Fe2+(aq) ——> Mn2+(aq) Fe3+(aq) + 4H2O(l)
this means that moles of Fe = 5
moles of MnO4¯

43. 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(H2O)6]2+(aq) OH¯(aq) ——> [Fe(OH)2(H2O)4](s) H2O(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
NH3 Iron(II) hydroxide precipitated, insoluble in excess ammonia
CO32- Off-white coloured iron(II) carbonate, FeCO3, precipitated
A-B OX
A-B
Ppt

44. REACTIONS OF IRON(II)
Volumetric Iron(II) can be analysed by titration with potassium manganate(VII)
in acidic (H2SO4) solution. No indicator is required.
MnO4¯(aq) H+(aq) + 5Fe2+(aq) ——> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
this means that moles of Fe = 5
moles of MnO4¯ 1
CONTENTS

45. REACTIONS OF IRON(III)
Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion
OH¯ [Fe(H2O)6]3+(aq) OH¯(aq) ——> [Fe(OH)3(H2O)3](s) H2O(l)
yellow rusty-brown ppt. insoluble in XS
A-B

46. REACTIONS OF IRON(III)
Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion
OH¯ [Fe(H2O)6]3+(aq) OH¯(aq) ——> [Fe(OH)3(H2O)3](s) H2O(l)
yellow rusty-brown ppt. insoluble in XS
CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) H2O(l) + 3CO2(g) rusty-brown ppt.
The carbonate is not precipitated but the hydroxide is; the high
charge density of M3+ makes the solution too acidic to form a carbonate
CARBON DIOXIDE EVOLVED.
A-B
A-B

47. REACTIONS OF IRON(III)
Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion
OH¯ [Fe(H2O)6]3+(aq) OH¯(aq) ——> [Fe(OH)3(H2O)3](s) H2O(l)
yellow rusty-brown ppt. insoluble in XS
CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) H2O(l) + 3CO2(g) rusty-brown ppt.
The carbonate is not precipitated but the hydroxide is; the high
charge density of M3+ makes the solution too acidic to form a carbonate
CARBON DIOXIDE EVOLVED.
NH3 [Fe(H2O)6]3+(aq) NH3(aq) ——> [Fe(OH)3(H2O)3](s) NH4+(aq)
rusty-brown ppt. insoluble in XS
A-B
A-B
A-B

48. REACTIONS OF IRON(III)
Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion
OH¯ [Fe(H2O)6]3+(aq) OH¯(aq) ——> [Fe(OH)3(H2O)3](s) H2O(l)
yellow rusty-brown ppt. insoluble in XS
CO32- 2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) H2O(l) + 3CO2(g) rusty-brown ppt.
The carbonate is not precipitated but the hydroxide is; the high
charge density of M3+ makes the solution too acidic to form a carbonate
CARBON DIOXIDE EVOLVED.
NH3 [Fe(H2O)6]3+(aq) NH3(aq) ——> [Fe(OH)3(H2O)3](s) NH4+(aq)
rusty-brown ppt. insoluble in XS
SCN¯ [Fe(H2O)6]3+(aq) SCN¯(aq) ——> [Fe(SCN)(H2O)5]2+(aq) + H2O(l)
blood-red colour
Very sensitive; BLOOD RED COLOUR confirms Fe(III). No reaction with Fe(II)
A-B
A-B
A-B LS

49. REACTIONS OF SILVER(I)
• aqueous solutions contains the colourless, linear, diammine silver(I) ion
• formed when silver halides dissolve in ammonia
eg AgCl(s) NH3(aq) ——> [Ag(NH3)2]+(aq) Cl¯(aq)
[Ag(SO3)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 S2O32- ——> [Ag(S2O3)2] Br¯
[Ag(CN)2]¯ Formed when silver salts are dissolved in sodium or potassium cyanide
the solution used for silver electroplating
[Ag(NH3)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

50. REACTIONS OF VANADIUM
Reduction using zinc in acidic solution shows the various oxidation states of vanadium.
Vanadium(V) VO2+(aq) H+(aq) e¯ ——> VO2+(aq) H2O(l)
yellow blue
Vanadium(IV) VO2+(aq) H+(aq) e¯ ——> V3+(aq) H2O(l)
blue blue/green
Vanadium(III) V3+(aq) e¯ ——> V2+(aq)
blue/green lavender

51. OXIDATION & REDUCTION - A SUMMARY
• 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¯
2Cr3+(aq) + 3H2O2(l) OH¯(aq) ——> 2CrO42-(aq) H2O(l)
• Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions
[Co(NH3)6]2+(aq) ——> [Co(NH3)6]3+(aq) e¯

52. OXIDATION & REDUCTION - A SUMMARY
• 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¯
2Cr3+(aq) + 3H2O2(l) OH¯(aq) ——> 2CrO42-(aq) H2O(l)
• Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions
[Co(NH3)6]2+(aq) ——> [Co(NH3)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) ——> Zn2+(aq) e¯
e.g. it reduces iron(III) to iron(II)
vanadium(V) to vanadium (IV) to vanadium(III)

53. 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 M3+ ions
• aqueous solutions contain the colourless, octahedral hexaaquaaluminium(III) ion
OH¯ [Al(H2O)6]3+(aq) OH¯(aq) ——> [Al(OH)3(H2O3](s) H2O(l)
colourless, octahedral white ppt. soluble in XS NaOH
As with all hydroxides the precipitate reacts with acid
[Al(OH)3(H2O)3](s) H+ (aq) ——> [Al(H2O)6]3+(aq)
being a 3+ hydroxide it is AMPHOTERIC and dissolves in excess alkali
[Al(OH)3(H2O)3](s) OH¯(aq) ——> [Al(OH)6]3-(aq) H2O(l)
colourless, octahedral
CO [Al(H2O)6]3+(aq) CO32-(aq) ——> 2[Al(OH)3(H2O)3](s) H2O(l) + 3CO2(g)
As with 3+ ions, the carbonate is not precipitated but the hydroxide is.
NH3 [Al(H2O)6]3+(aq) NH3(aq) ——> [Al(OH)3(H2O)3](s) + 3NH4+(aq) white ppt. insoluble in XS NH3
A-B
A-B
A-B
A-B
A-B

54. © 2009 JONATHAN HOPTON & KNOCKHARDY PUBLISHING
AN INTRODUCTION TO
TRANSITION METAL
COMPLEXES
THE END
© JONATHAN HOPTON & KNOCKHARDY PUBLISHING