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

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An introduction to transition metal complexes

AN INTRODUCTION TO

TRANSITION METAL

COMPLEXES

2008 SPECIFICATIONS

KNOCKHARDY PUBLISHING


An introduction to transition metal complexes

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...

www.knockhardy.org.uk/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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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 powerand

the more acidic the solution


An introduction to transition metal complexes

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 nmdissolves 7

Mg 2+0.065 nmslight hydrolysis

Al 3+0.050 nmvigorous hydrolysis

the greater charge density of the cation...

the greater the polarising powerand

the more acidic the solution


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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 d+)

it can then be removed by solvent water molecules to form H3O+(aq).


An introduction to transition metal complexes

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...


An introduction to transition metal complexes

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...


An introduction to transition metal complexes

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) orM(OH)2

M3+ ions[M(H2O)3(OH)3](s) orM(OH)3

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+


An introduction to transition metal complexes

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) orM(OH)2

M3+ ions[M(H2O)3(OH)3](s) orM(OH)3

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+


An introduction to transition metal complexes

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 H2Oremove 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+


An introduction to transition metal complexes

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 H2Oremove 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+


An introduction to transition metal complexes

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

Soluble

Insoluble


An introduction to transition metal complexes

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.


An introduction to transition metal complexes

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)


An introduction to transition metal complexes

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)

the expression for the

stability constant is Kstab = [ [MX64–](aq) ]

[ [M(H2O)6]2+(aq) ] [ X¯(aq) ]6


An introduction to transition metal complexes

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)

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’.


An introduction to transition metal complexes

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 = 1.02 x 10-2

[Co(NH3)(H2O)5]2+(aq) + NH3(aq)[Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2

[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)[Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1

[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)[Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1

[Co(NH3)4(H2O)2]2+(aq) + NH3(aq)[Co(NH3)5(H2O)]2+(aq) + H2O(l) K5 = 8.70 x 10-1 etc


An introduction to transition metal complexes

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 = 1.02 x 10-2

[Co(NH3)(H2O)5]2+(aq) + NH3(aq)[Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2

[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)[Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1

[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)[Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1

[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


An introduction to transition metal complexes

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 = 1.02 x 10-2

[Co(NH3)(H2O)5]2+(aq) + NH3(aq)[Co(NH3)2(H2O)4]2+(aq) + H2O(l) K2 = 3.09 x 10-2

[Co(NH3)2(H2O)4]2+(aq) + NH3(aq)[Co(NH3)3(H2O)3]2+(aq) + H2O(l) K3 = 1.17 x 10-1

[Co(NH3)3(H2O)3]2+(aq) + NH3(aq)[Co(NH3)4(H2O)2]2+(aq) + H2O(l) K4 = 2.29 x 10-1

[Co(NH3)4(H2O)2]2+(aq) + NH3(aq)[Co(NH3)5(H2O)]2+(aq) + H2O(l) K5 = 8.70 x 10-1 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


An introduction to transition metal complexes

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

typesACID-BASE

LIGAND SUBSTITUTION

PRECIPITATION

REDOX

A-B

LS

Ppt

RED

OX

REDOX


An introduction to transition metal complexes

REACTION TYPES

The examples aim to show typical properties of transition metals and their compounds.

LOOKFOR...

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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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.

Cl¯[Co(H2O)6]2+(aq) + 4Cl¯(aq) ——> [CoCl4]2-(aq) + 6H2O(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


An introduction to transition metal complexes

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


An introduction to transition metal complexes

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

NH3[Cu(H2O)6]2+(aq) + 2NH3(aq) ——> [Cu(OH)2(H2O)4](s) + 2NH4+(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 blueNOTE THE FORMULA

A-B

A-B

LS


An introduction to transition metal complexes

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

NH3[Cu(H2O)6]2+(aq) + 2NH3(aq) ——> [Cu(OH)2(H2O)4](s) + 2NH4+(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 blueNOTE THE FORMULA

CO32-[Cu(H2O)6]2+(aq) + CO32-(aq) ——> CuCO3(s) + 6H2O(l)

blue ppt.

A-B

A-B

LS

Ppt


An introduction to transition metal complexes

REACTIONS OF COPPER(II)

Cl¯[Cu(H2O)6]2+(aq) + 4Cl¯(aq) ——> [CuCl4]2-(aq) + 6H2O(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) + 4I¯(aq) ——> 2CuI(s) + I2(aq)

off - white ppt.

• a redox reaction

• used in the volumetric analysis of copper using sodium thiosulphate

LS

REDOX


An introduction to transition metal complexes

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

Cu2+(aq) + e¯ ——> Cu+(aq)E° = + 0.15 V

subtracting2Cu+(aq) ——> Cu(s) + Cu2+(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.


An introduction to transition metal 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) ——> [Cr(OH)3(H2O)3](s) + 3H2O(l)

violet, octahedral 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) + 3H2O(l)

green, octahedral

A-B

A-B

A-B


An introduction to transition metal complexes

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.

high charge density of M3+ makes the solution too acidic to form the carbonate

CARBON DIOXIDE IS EVOLVED.

A-B


An introduction to transition metal complexes

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.

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)

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


An introduction to transition metal complexes

REACTIONS OF CHROMIUM(III)

OxidationIn 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)

greenyellow

Acidification of the yellow chromate will produce the orange dichromate(VI) ion

ReductionChromium(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


An introduction to transition metal complexes

REACTIONS OF CHROMIUM(VI)

Occurrencedichromate (VI)Cr2O72- orange

chromate (VI) CrO42- yellow

Interconversiondichromate 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)


An introduction to transition metal complexes

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) [ E° = +1.33 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


An introduction to transition metal complexes

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

acidicMnO4¯(aq) + 8H+(aq) + 5e¯ ——> Mn2+(aq) + 4H2O(l) E° = + 1.52 V

N.B. Acidify with dilute H2SO4 NOT dilute HCl

alkalineMnO4¯(aq) + 2H2O(l) + 3e¯ ——> MnO2(s) + 4OH¯(aq) E° = + 0.59 V


An introduction to transition metal complexes

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+ = 5

moles of MnO4¯ 1


An introduction to transition metal complexes

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) + 2H2O(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

NH3Iron(II) hydroxide precipitated, insoluble in excess ammonia

CO32-Off-white coloured iron(II) carbonate, FeCO3, precipitated

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An introduction to transition metal complexes

REACTIONS OF IRON(II)

VolumetricIron(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+ =5

moles of MnO4¯ 1

CONTENTS


An introduction to transition metal complexes

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

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An introduction to transition metal complexes

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

CO32-2[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.

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An introduction to transition metal complexes

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

CO32-2[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.

NH3[Fe(H2O)6]3+(aq) + 3NH3(aq) ——> [Fe(OH)3(H2O)3](s) + 3NH4+(aq)

rusty-brown ppt. insoluble in XS

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An introduction to transition metal complexes

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

CO32-2[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.

NH3[Fe(H2O)6]3+(aq) + 3NH3(aq) ——> [Fe(OH)3(H2O)3](s) + 3NH4+(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 COLOURconfirms Fe(III). No reaction with Fe(II)

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An introduction to transition metal complexes

REACTIONS OF SILVER(I)

• aqueous solutions contains the colourless, linear, diammine silver(I)ion

• formed when silver halides dissolve in ammonia

egAgCl(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


An introduction to transition metal complexes

REACTIONS OF VANADIUM

Reduction using zinc in acidic solution shows the various oxidation states of vanadium.

Vanadium(V)VO2+(aq) + 2H+(aq) + e¯ ——> VO2+(aq) + H2O(l)

yellow blue

Vanadium(IV)VO2+(aq) + 2H+(aq) + e¯ ——> V3+(aq) + H2O(l)

blue blue/green

Vanadium(III)V3+(aq) + e¯ ——> V2+(aq)

blue/green lavender


An introduction to transition metal complexes

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¯


An introduction to transition metal complexes

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)


An introduction to transition metal complexes

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) ——> [Al(OH)3(H2O3](s) + 3H2O(l)

colourless, octahedral white ppt. soluble in XS NaOH

As with all hydroxides the precipitate reacts with acid

[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

[Al(OH)3(H2O)3](s) + 3OH¯(aq) ——> [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.

NH3[Al(H2O)6]3+(aq) + 3NH3(aq) ——> [Al(OH)3(H2O)3](s) + 3NH4+(aq)white ppt. insoluble in XS NH3

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An introduction to transition metal complexes

AN INTRODUCTION TO

TRANSITION METAL

COMPLEXES

THE END

© 2009 JONATHAN HOPTON & KNOCKHARDY PUBLISHING


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