AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING.

AN INTRODUCTION TOAN INTRODUCTION TO

TRANSITION METALTRANSITION METALCOMPLEXESCOMPLEXES

KNOCKHARDY PUBLISHINGKNOCKHARDY PUBLISHING

INTRODUCTION

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TRANSITION METALSTRANSITION 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

TRANSITION METALSTRANSITION METALS

THE AQUEOUS CHEMISTRY OF IONS - HYDROLYSISTHE AQUEOUS CHEMISTRY OF IONS - HYDROLYSIS

when salts dissolve in water the ions are stabilisedthis is because water molecules are polarhydrolysis can occur and the resulting solution can become acidicthe acidity of the resulting solution depends on the cation presentthe greater the charge density of the cation, the more acidic the solution

cation charge ionic 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 power andthe more acidic the solution

THE AQUEOUS CHEMISTRY OF IONSTHE AQUEOUS CHEMISTRY OF IONS

Theory aqueous metal ions attract water moleculesmany have six water molecules surrounding themthese are known as hexaaqua ionsthey are octahedral in shapewater acts as a Lewis Base – a lone pair donorwater forms a co-ordinate bond to the metal ionmetal ions accept the lone pair - Lewis Acids

THE AQUEOUS CHEMISTRY OF IONSTHE AQUEOUS CHEMISTRY OF IONS

Theory aqueous metal ions attract water moleculesmany have six water molecules surrounding themthese are known as hexaaqua ionsthey are octahedral in shapewater acts as a Lewis Base – a lone pair donorwater forms a co-ordinate bond to the metal ionmetal ions accept the lone pair - Lewis Acids

Acidity as charge density increases, the cation has a greater attraction for waterthe 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 polarthis makes the H more acidic (more+)it can then be removed by solvent water molecules to form H3O+(aq).

HYDROLYSIS - HYDROLYSIS - EQUATIONSEQUATIONS

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 - HYDROLYSIS - EQUATIONSEQUATIONS

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 IONSHYDROLYSIS 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 becomesneutral 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+






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+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

Precipitated






AMPHOTERIC CHARACTERMetal 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+

OH¯

H+

OH¯

H+

OH¯

H+

OH¯

H+

Precipitated

OH¯H+

Insoluble SolubleSoluble

REACTION TYPESREACTION 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

Reactiontypes ACID-BASE

LIGAND SUBSTITUTION

PRECIPITATION

REDOX

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Ppt

RED REDOX

REACTION TYPESREACTION 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)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


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




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




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


CO32- [Co(H2O)6]2+(aq) + CO3

2-(aq) ——> CoCO3(s) + 6H2O(l)

mauve ppt.

Hexaaqua ions of metals with charge 2+ precipitate a carbonate butheaxaaqua ions with a 3+ charge don’t.

Ppt


CO32- [Co(H2O)6]2+(aq) + CO3

2-(aq) ——> CoCO3(s) + 6H2O(l)

mauve ppt.

Hexaaqua ions of metals with charge 2+ precipitate a carbonate butheaxaaqua 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

LS

Ppt

REACTIONS OF COPPER(II)REACTIONS OF COPPER(II)

Aqueous solutions of copper(II) contain the blue, octahedral hexaaquacopper(II) ionMost 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





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 blue

NOTE THE FORMULA

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A-B

LS





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 blue

NOTE THE FORMULA

CO32- [Cu(H2O)6]2+(aq) + CO3

2-(aq) ——> CuCO3(s) + 6H2O(l)

blue ppt.

A-B

A-B

LS

Ppt


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

REACTIONS OF COPPER(I)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

subtracting 2Cu+(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.

REACTIONS OF CHROMIUM(III)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


CO32- 2 [Cr(H2O)6]3+(aq) + 3CO3

2-(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


CO32- 2 [Cr(H2O)6]3+(aq) + 3CO3

2-(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)

LS

A-B

A-B


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 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

REACTIONS OF CHROMIUM(VI)REACTIONS OF CHROMIUM(VI)

Occurrence dichromate (VI) Cr2O72- orange

chromate (VI) CrO42- yellow

Interconversion dichromate is stable in acid solution

chromate is stable in alkaline solution

in alkali Cr2O72-(aq) + 2OH¯(aq) 2CrO4

2-(aq) + H2O(l)

in acid 2CrO42-(aq) + 2H+(aq) Cr2O7

2-(aq) + H2O(l)

CONTENTSCONTENTS

OXIDATION REACTIONS OF CHROMIUM(VI)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

REACTIONS OF MANGANESE(VII)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)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

REACTIONS OF IRON(II)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

NH3 Iron(II) hydroxide precipitated, insoluble in excess ammonia

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

A-B

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Ppt

CONTENTSCONTENTS

REACTIONS OF IRON(II)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+ = 5

moles of MnO4¯ 1

REACTIONS OF IRON(III)REACTIONS OF IRON(III)

Aqueous solutions contain the yellow-green, octahedral hexaaquairon(III) ion


yellow rusty-brown ppt. insoluble in XS

A-B





CO32- 2[Fe(H2O)6]3+(aq) + 3CO3

2-(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 highcharge density of M3+ makes the solution too acidic to form a carbonateCARBON DIOXIDE EVOLVED.

A-B

A-B





CO32- 2[Fe(H2O)6]3+(aq) + 3CO3


rusty-brown ppt.


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

rusty-brown ppt. insoluble in XS

A-B

A-B

A-B





CO32- 2[Fe(H2O)6]3+(aq) + 3CO3


rusty-brown ppt.


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 COLOUR confirms Fe(III). No reaction with Fe(II)

A-B

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REACTIONS OF SILVER(I)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 VANADIUMREACTIONS 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

OXIDATION & REDUCTION - OXIDATION & REDUCTION - A SUMMARYA 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¯


• 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 - OXIDATION & REDUCTION - A SUMMARYA 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¯


• 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 ALUMINIUMREACTIONS 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) + 3CO3

2-(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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© JONATHAN HOPTON & KNOCKHARDY PUBLISHING© JONATHAN HOPTON & KNOCKHARDY PUBLISHING

THE ENDTHE END

AN INTRODUCTION TOAN INTRODUCTION TO

TRANSITION METALTRANSITION METALCHEMISTRYCHEMISTRY

AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING.

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Transcript of AN INTRODUCTION TO TRANSITION METAL COMPLEXES KNOCKHARDY PUBLISHING.