Transcript transition metal complexes

AN INTRODUCTION TO
TRANSITION METAL
COMPLEXES
KNOCKHARDY PUBLISHING
TRANSITION METALS
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
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 lone pair donor
water forms a co-ordinate bond to the metal ion
metal ions accept the lone pair
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)2(H2O)4](s)
[M(OH)4(H2O)2
]2-(aq)
OH¯
H+
OH¯
H+
OH¯
H+
[M(OH)(H2O)5]+(aq)
[M(OH)3(H2O)3]¯(aq)
[M(OH)5(H2
O)]3-(aq)
OH¯
H+
OH¯
H+
OH¯
H+
[M(OH)2(H2O)4](s)
[M(OH)4(H2O)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.
M2+ ions
M3+ ions
[M(H2O)4(OH)2](s)
[M(H2O)3(OH)3](s)
or
or
M(OH)2
M(OH)3
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
• co-ordination number
• stability to oxidation or reduction
Reaction
types
LIGAND SUBSTITUTION
LS
PRECIPITATION
Ppt
REDOX
RED
OX
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
differences in reactivity of M3+ and M2+ ions with OH¯ and NH3
the effect a ligand has on the stability of a particular oxidation state
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) + 2OH¯(aq) ——> [Cu(OH)2(H2O)4](s) + 2H2O(l)
blue, octahedral
pale blue ppt.
insoluble in XS NaOH
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(H2O)6]2+(aq) + 2OH¯(aq) ——> [Cu(OH)2(H2O)4](s) + 2H2O(l)
blue, octahedral
NH3
A-B
pale blue ppt.
insoluble in XS NaOH
[Cu(H2O)6]2+(aq) + 2NH3(aq) ——> [Cu(OH)2(H2O)4](s) + 2NH4+(aq)
A-B
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
LS
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) + 8H+(aq) + 5e¯ ——> Mn2+(aq) + 4H2O(l)
E° = + 1.52 V
N.B. Acidify with dilute H2SO4 NOT dilute HCl
alkaline
MnO4¯(aq) + 2H2O(l) + 3e¯ ——> MnO2(s) + 4OH¯(aq)
E° = + 0.59 V
VOLUMETRIC USE OF MANGANATE(VII)
Potassium manganate(VII) in acidic (H2SO4) solution is extremely useful for carrying out
redox volumetric analysis.
MnO4¯(aq) + 8H+(aq) + 5e¯ ——> Mn2+(aq) + 4H2O(l)
E° = + 1.52 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) + 8H+(aq) + 5Fe2+(aq) ——> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
this means that
moles of Fe2+
moles of MnO4¯
=
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(H2O)6]2+(aq) + 2OH¯(aq) ——> [Fe(OH)2(H2O)4](s)
pale green
+ 2H2O(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 pH renders iron(II) unstable.
NH3
Fe(OH)2(s) + OH¯(aq) ——>
Fe(OH)3(s) + e¯
dirty green
rusty brown
Iron(II) hydroxide precipitated, insoluble in excess ammonia
OX
A-B
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) + 8H+(aq) + 5Fe2+(aq) ——> Mn2+(aq) + 5Fe3+(aq) + 4H2O(l)
this means that
moles of Fe2+
moles of MnO4¯
CONTENTS
=
5
1
REACTIONS OF IRON(III)
Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion
OH¯
[Fe(H2O)6]3+(aq) + 3OH¯(aq) ——> [Fe(OH)3(H2O)3](s) + 3H2O(l)
yellow
NH3
A-B
rusty-brown ppt. insoluble in XS
[Fe(H2O)6]3+(aq) + 3NH3(aq) ——> [Fe(OH)3(H2O)3](s)
+
3NH4+(aq)
rusty-brown ppt. insoluble in XS
A-B
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¯
2Cr3+(aq) + 3H2O2(l) + 10OH¯(aq) ——> 2CrO42-(aq) + 8H2O(l)
• Solutions of cobalt(II) can be oxidised by air under ammoniacal conditions
[Co(NH3)6]2+(aq) ——> [Co(NH3)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¯
2Cr3+(aq) + 3H2O2(l) + 10OH¯(aq) ——> 2CrO42-(aq) + 8H2O(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) + 2e¯
e.g. it reduces
iron(III) to iron(II)
vanadium(V) to vanadium (IV) to vanadium(III)