AN INTRODUCTION TO
TRANSITION METAL
COMPLEXES
KNOCKHARDY PUBLISHING
2008
SPECIFICATIONS
KNOCKHARDY PUBLISHING
TRANSITION METALS
INTRODUCTION
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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
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 / pH 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 / pH 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 ion
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 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).
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...
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...
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
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
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)
Precipitated
[M(OH)4(H2O)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
H2O
NH3, CO32OH¯
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(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)
Precipitated
[M(OH)4(H2O)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
H2O
NH3, CO32OH¯
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(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)
Precipitated
[M(OH)4(H2O)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(H2O)6]3+(aq)
Soluble
H+
OH¯
[M(OH)3(H2O)3](s)
Insoluble
[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(H2O)6]2+(aq) + 6X¯(aq)
[MX6]4–(aq) + 6H2O(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(H2O)6]2+(aq) + 6X¯(aq)
Kstab =
[MX6]4–(aq) + 6H2O(l)
[ [MX64–](aq) ]
[ [M(H2O)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(H2O)6]2+(aq) + 6X¯(aq)
Kstab =
[MX6]4–(aq) + 6H2O(l)
[ [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’.
STABILITY CONSTANTS
Because ligand exchange involves a series of equilibria, each step in the process has a
different stability constant…
[Co(H2O)6]2+(aq) + NH3(aq)
[Co(NH3)(H2O)5]2+(aq)
+ H2O(l)
Kstab / dm3 mol-1
K1 = 1.02 x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq)
10-2
[Co(NH3)2(H2O)4]2+(aq)
+ H2O(l)
K2 = 3.09 x
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)
10-1
[Co(NH3)3(H2O)3]2+(aq)
+ H2O(l)
K3 = 1.17 x
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)
10-1
[Co(NH3)4(H2O)2]2+(aq)
+ H2O(l)
K4 = 2.29 x
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq)
[Co(NH3)5(H2O)]2+(aq)
+ H2O(l)
K5 = 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(H2O)6]2+(aq) + NH3(aq)
[Co(NH3)(H2O)5]2+(aq)
+ H2O(l)
Kstab / dm3 mol-1
K1 = 1.02 x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq)
10-2
[Co(NH3)2(H2O)4]2+(aq)
+ H2O(l)
K2 = 3.09 x
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)
10-1
[Co(NH3)3(H2O)3]2+(aq)
+ H2O(l)
K3 = 1.17 x
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)
10-1
[Co(NH3)4(H2O)2]2+(aq)
+ H2O(l)
K4 = 2.29 x
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq)
[Co(NH3)5(H2O)]2+(aq)
+ H2O(l)
K5 = 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... k1 x k2 x k3 x k4 ... etc
Kstab or pKstab?
For an easier comparison,
the expression pKstab is often used…
pKstab = -log10Kstab
STABILITY CONSTANTS
Because ligand exchange involves a series of equilibria, each step in the process has a
different stability constant…
[Co(H2O)6]2+(aq) + NH3(aq)
[Co(NH3)(H2O)5]2+(aq)
+ H2O(l)
Kstab / dm3 mol-1
K1 = 1.02 x 10-2
[Co(NH3)(H2O)5]2+(aq) + NH3(aq)
10-2
[Co(NH3)2(H2O)4]2+(aq)
+ H2O(l)
K2 = 3.09 x
[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)
10-1
[Co(NH3)3(H2O)3]2+(aq)
+ H2O(l)
K3 = 1.17 x
[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)
10-1
[Co(NH3)4(H2O)2]2+(aq)
+ H2O(l)
K4 = 2.29 x
[Co(NH3)4(H2O)2]2+(aq) + NH3(aq)
[Co(NH3)5(H2O)]2+(aq)
•
•
•
•
+ H2O(l)
K5 = 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 pKstab
Complex ions with smaller pKstab 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
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
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
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) + 2OH¯(aq) ——> [Co(OH)2(H2O)4](s) + 2H2O(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
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
REACTIONS OF COBALT(II)
NH3
[Co(H2O)6]2+(aq) + 2NH3(aq) ——> [Co(OH)2(H2O)4](s) + 2NH4+(aq)
A-B
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)
LS
REACTIONS OF COBALT(II)
NH3
[Co(H2O)6]2+(aq) + 2NH3(aq) ——> [Co(OH)2(H2O)4](s) + 2NH4+(aq)
A-B
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)
LS
but ... ammonia ligands make the Co(II) state unstable. Air oxidises Co(II) to Co(III)
[Co(NH3)6]2+(aq)
yellow / brown octahedral
——>
[Co(NH3)6]3+(aq)
+
red / brown octahedral
e¯
OX
REACTIONS OF COBALT(II)
CO32-
[Co(H2O)6]2+(aq)
+ CO32-(aq)
——>
CoCO3(s) + 6H2O(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)
CO32-
[Co(H2O)6]2+(aq)
+ CO32-(aq)
——>
CoCO3(s) + 6H2O(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(H2O)6]2+(aq) + 4Cl¯(aq)
——> [CoCl4]2-(aq) + 6H2O(l)
pink, octahedral
blue, tetrahedral
LS
• 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
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 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
CO32-
[Cu(H2O)6]2+(aq)
+ CO32-(aq)
——>
CuCO3(s) + 6H2O(l)
blue ppt.
LS
Ppt
REACTIONS OF COPPER(II)
Cl¯
[Cu(H2O)6]2+(aq) + 4Cl¯(aq)
——>
[CuCl4]2-(aq) + 6H2O(l)
yellow, tetrahedral
LS
• 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) +
4I¯(aq)
——>
2CuI(s) +
I2(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(s)
Cu+(aq)
E° = + 0.52 V
E° = + 0.15 V
2Cu+(aq) ——> Cu(s) +
Cu2+(aq)
E° = + 0.37 V
Cu+(aq) + e¯
Cu2+(aq) + e¯
subtracting
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 M3+ ions in this block.
Aqueous solutions contain the violet, octahedral hexaaquachromium(III) ion.
OH¯
[Cr(H2O)6]3+(aq) + 3OH¯(aq)
violet, octahedral
——>
[Cr(OH)3(H2O)3](s) + 3H2O(l)
A-B
green ppt. soluble in XS NaOH
As with all hydroxides the precipitate reacts with acid
[Cr(OH)3(H2O)3](s) + 3H+(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)
+ 3OH¯(aq) ——> [Cr(OH)6]3-(aq)
green, octahedral
+
3H2O(l)
A-B
A-B
REACTIONS OF CHROMIUM(III)
CO32-
2 [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.
A-B
high charge density of M3+ makes the solution too acidic to form the carbonate
CARBON DIOXIDE IS EVOLVED.
REACTIONS OF CHROMIUM(III)
CO32-
2 [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.
A-B
high charge density of M3+ makes the solution too acidic to form the carbonate
CARBON DIOXIDE IS EVOLVED.
NH3
[Cr(H2O)6]3+(aq) + 3NH3(aq)
——> [Cr(OH)3(H2O)3](s) + 3NH4+(aq)
A-B
green ppt. soluble in XS NH3
With EXCESS AMMONIA, the precipitate redissolves
LS
[Cr(OH)3(H2O)3](s) + 6NH3(aq) ——> [Cr(NH3)6]3+(aq) + 3H2O(l) + 3OH¯(aq)
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) + 10OH¯(aq) ——> 2CrO42-(aq) + 8H2O(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(H2O)6]3+(aq) + Zn(s)
green
——>
2 [Cr(H2O)6]2+(aq)
blue
+
Zn2+(aq)
RED
REACTIONS OF CHROMIUM(VI)
Occurrence
Interconversion
dichromate (VI)
Cr2O72-
orange
chromate (VI)
CrO42-
yellow
dichromate is stable in acid solution
chromate is stable in alkaline solution
in alkali
Cr2O72-(aq) + 2OH¯(aq)
2CrO42-(aq) + H2O(l)
in acid
2CrO42-(aq) + 2H+(aq)
Cr2O72-(aq) +
H2O(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.
Cr2O72-(aq) + 14H+(aq) + 6e¯ ——> 2Cr3+(aq) + 7H2O(l)
orange
[ E° = +1.33 V ]
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
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.
Fe(OH)2(s) + OH¯(aq) ——>
Fe(OH)3(s) + e¯
dirty green
rusty brown
OX
NH3
Iron(II) hydroxide precipitated, insoluble in excess ammonia
A-B
CO32-
Off-white coloured iron(II) carbonate, FeCO3, precipitated
Ppt
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
rusty-brown ppt. insoluble in XS
A-B
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
CO32-
A-B
rusty-brown ppt. insoluble in XS
2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) + 3H2O(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
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
CO32-
A-B
rusty-brown ppt. insoluble in XS
2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g)
rusty-brown ppt.
NH3
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
[Fe(H2O)6]3+(aq) + 3NH3(aq) ——> [Fe(OH)3(H2O)3](s)
A-B
+
3NH4+(aq)
rusty-brown ppt. insoluble in XS
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
CO32-
A-B
rusty-brown ppt. insoluble in XS
2[Fe(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Fe(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g)
rusty-brown ppt.
NH3
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
[Fe(H2O)6]3+(aq) + 3NH3(aq) ——> [Fe(OH)3(H2O)3](s)
A-B
+
3NH4+(aq)
rusty-brown ppt. insoluble in XS
SCN¯
[Fe(H2O)6]3+(aq) + SCN¯(aq) ——> [Fe(SCN)(H2O)5]2+(aq) + H2O(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
AgCl(s)
+
2NH3(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 + 2S2O32- ——>
[Ag(S2O3)2]3-
+ 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
REACTIONS OF VANADIUM
Reduction using zinc in acidic solution shows the various oxidation states of vanadium.
Vanadium(V)
VO2+(aq)
+
2H+(aq)
+
e¯
——>
yellow
Vanadium(IV)
VO2+(aq)
V3+(aq)
+
H2O(l)
blue
+
2H+(aq)
+
e¯
blue
Vanadium(III)
VO2+(aq)
——>
V3+(aq)
blue/green
+
blue/green
e¯
——>
V2+(aq)
lavender
+
H2O(l)
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)
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) + 3OH¯(aq)
——>
3H2O(l)
A-B
[Al(OH)3(H2O)3](s) + 3H+ (aq) ——> [Al(H2O)6]3+(aq)
being a 3+ hydroxide it is AMPHOTERIC and dissolves in excess alkali
A-B
[Al(OH)3(H2O)3](s) + 3OH¯(aq) ——>
A-B
colourless, octahedral
[Al(OH)3(H2O3](s)
+
white ppt. soluble in XS NaOH
As with all hydroxides the precipitate reacts with acid
[Al(OH)6]3-(aq)
+
3H2O(l)
colourless, octahedral
CO32-
2 [Al(H2O)6]3+(aq) + 3CO32-(aq) ——> 2[Al(OH)3(H2O)3](s) + 3H2O(l) + 3CO2(g)
As with 3+ ions, the carbonate is not precipitated but the hydroxide is.
A-B
NH3
[Al(H2O)6]3+(aq) + 3NH3(aq)
——>
[Al(OH)3(H2O)3](s) + 3NH4+(aq)
white ppt. insoluble in XS NH3
A-B
AN INTRODUCTION TO
TRANSITION METAL
COMPLEXES
THE END
© 2009 JONATHAN HOPTON & KNOCKHARDY PUBLISHING