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APPLIED INORGANIC CHEMISTRY FOR CHEMICAL ENGINEERS

Transition Metal Chemistry

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Elements are divided into four categories

Periodic Table

Main-group elements(S-Block) Transition metals

Main-group elements (P-Block)

1. Main-group elements

2. Transition metals

3. Lanthanides

4. Actinides

( )

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Lanthanides

Actinides

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Transition metals vs. Main-group elements

Main‐group elements

Transition metals

Main‐group metals 

• malleable and ductile

• conduct heat and electricity• form positive ions

Transition metals

• more electronegative than the main group metals

• more likely to form covalent compounds

• easily form complexesCisplatin

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There is some controversy about the classification of the elements

i.e. Zinc (Zn), Cadmium (Cd) and Mercury (Hg)     e‐ configuration  [  ]ns2 n‐1d10

• form stable compounds with neutral molecules

IUPAC ‐ A transition metal is "an element whose atom has an incomplete d sub‐shell, or which can give rise to cations with an incomplete d sub‐shell.” 

Electron configuration of Transition-metal ions

The relationship between the electron configurations of transition‐metal

elements and their ions is complex.

Example

Consider the chemistry of cobalt which forms complexes that contain

either Co2+or Co3+ ions.

Co:

Co2+:

Co has 27 electrons

[Ar] has 18 electrons[Ar]

[Ar]

4s2 3d7

3d7

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Co3+:

In general, electrons are removed from the valence shell s orbitals before

they are removed from valence d orbitals when transition metals are

ionized.

[Ar] 3d6

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How do we determine the electronic configuration of the central metal ion

in any complex?

• Try to recognise all the entities making up the complex

• Need knowing whether the ligands are neutral or anionic

• Then you can determine the oxidation state of the metal ion.

A simple procedure exists for the M(II) case …same as M(+2) or M2+

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Ti V Cr Mn Fe Co Ni Cu

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Cross off the first 2 gives you total No. of valence electrons left

2 3 4 5 6 7 8 9

EXAMPLES

Elements Configuration Oxidized elements Configuration

Sc [Ar]4s23d1 Sc(III) [Ar]Sc [Ar]4s 3d Sc(III) [Ar]

V [Ar]4s23d3 V(II) [Ar]3d3

Cr [Ar]4s13d5 Cr(III) [Ar]3d3

Fe  [Ar]4s23d6 Fe(II) [Ar]3d6

Ni [Ar]4s23d8 Ni(II) [Ar]3d8

Cu [Ar]4s13d10 Cu(I) [Ar]3d10

Zn [Ar]4s23d10 Zn(II) [Ar}3d10

…variety of oxidation states !!

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Evaluating the oxidation state

[CoCl(NO2)(NH3)4]+

Neutralzero charge

Net charge on complex ion (+1)

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x = +3

X + (- 2) + 0 = +1

Co3+

zero charge

X - 2 = +1

Why do these elements exhibit a variety of oxidation states?

Because of the closeness of the 3d and 4s energy states

Sc +3

Oxidation states and their relative stabilities

Sc +3

Ti +1 +2 +3 +4

V +1 +2 +3 +4 +5

Cr +1 +2 +3 +4 +5 +6

Mn +1 +2 +3 +4 +5 +6 +7

Fe +1 +2 +3 +4 +5 +6

C +1 +2 +3 +4 +5

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The most prevalent oxidation numbers are shown in green. 

Co +1 +2 +3 +4 +5

Ni +1 +2 +3 +4

Cu +1 +2 +3

Zn +2

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An increase in the No. of oxidation states from Sc to Mn.

All seven oxidation states are exhibited by Mn.

There is a decrease in the No. of oxidation states from Mn to Zn.

WHY?

Because the pairing of d-electrons occurs after Mn (Hund's rule)

which in turn decreases the number of available unpaired electrons

and hence, the number of oxidation states.

Sc +3

Ti +1 +2 +3 +4

V +1 +2 +3 +4 +5

Cr +1 +2 +3 +4 +5 +6

Mn +1 +2 +3 +4 +5 +6 +7

Fe +1 +2 +3 +4 +5 +6

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The stability of higher oxidation states decreases in moving from Sc

to Zn.

Mn(VII) and Fe(VI) are powerful oxidizing agents and the higher

oxidation states of Co, Ni and Zn are unknown.

Co +1 +2 +3 +4 +5

Ni +1 +2 +3 +4

Cu +1 +2 +3

Zn +2

The relative stability of +2 state with respect to higher oxidation

states increases in moving from left to right. On the other hand +3

state becomes less stable from left to right.

Why? Sc +3

This is justifiable since it will be increasingly difficult to remove the

third electron from the d-orbital.

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Ti V Cr Mn Fe Co Ni Cu

Example

Ti +1 +2 +3 +4

V +1 +2 +3 +4 +5

Cr +1 +2 +3 +4 +5 +6

Mn +1 +2 +3 +4 +5 +6 +7

Fe +1 +2 +3 +4 +5 +6

Co +1 +2 +3 +4 +5

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Ti V Cr Mn Fe Co Ni Cu

M = [Ar]4s23dx

M+2 = [Ar]3dx loss of the two s electrons

M+3 = [Ar]3dx-1 more difficult

Ni +1 +2 +3 +4

Cu +1 +2 +3

Zn +2

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• Oxidized by HCl or H2SO4 to form blue Cr2+ ion

• Cr2+ oxidized by O2 in air to form green Cr3+

Chromium

A i t 1

• Cr also found in +6 state as in CrO42− and

Cr2O72− are strong oxidizer

Write down balance equations that show the

two reactions

Assignment 1

Cr2O7 are strong oxidizer

Use balanced equations to show that CrO42−

and Cr2O72− are strong oxidizing agents

Assignment 2

Assignment 1

Solution

Cr + H SO Cr SO + H

2 Cr(s) + 4 HCl(aq) 2 CrCl2(aq) + 2H2(g)

Cr(S) + H2SO4(aq) Cr2SO4(aq) + H2(g)

2CrCl2(aq) + O2(g) Cr2O2Cl2(aq) + Cl2(g)

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• Fe exists in solution in +2 or +3 state

• Elemental Fe reacts with non-oxidizing acids to

Iron

form Fe2+, which oxidizes in air (O2) to Fe3+

• Brown water running from a faucet is caused by

insoluble Fe2O3

• Fe3+ soluble in acidic solution, but forms a

hydrated oxide as red-brown gel in basic

solutionsolution

Assignment 3

Fe2O3

Complete and balance the following equation

+ HCl

Coordination Chemistry

A coordination compound (complex), contains a central metal atom

(or ion) surrounded by a number of oppositely charged ions or neutral

molecules (possessing lone pairs of electrons) which are known as

ligands.

If a ligand is capable of forming more than

one bond with the central metal atom or ion,

then ring structures are produced which are

known as metal chelates

the ring forming groups are described as

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the ring forming groups are described as

chelating agents or polydentate ligands.

The coordination number of the central metal atom or ion is the total

number of sites occupied by ligands.

Note: a bidentate ligand uses 2 sites, a tridentate 3 sites etc.

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Ligands

molecular formula

Lewis base/ligand

Lewis acid

donor atom

coordination numberformula base/ligand acid atom number

[Zn(CN)4]2- CN- Zn2+ C 4

[PtCl6]2- Cl- Pt4+ Cl 6

[Ni(NH3)6]2+ :NH3 Ni2+ N 6

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Mono-dentate

Multidentate ligands

Abbreviation Name Formula

en Ethylenediamine

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ox2- Oxalato

EDTA4- Ethylenediamine-tetraacetanato

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Chelating ligands bond to metal

Five or six atoms rings are common

forms rings – chelate rings

Coordination numbers and geometries

(i.e. including metal)

Li

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Linear

Square planar Tetrahedral Octahedral

Nomenclature of Coordination Compounds

• The basic protocol in coordination nomenclature is to name the ligands

attached to the metal as prefixes before the metal name.

• Some common ligands and their names are listed above.

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As is the case with ionic compounds, the name of the cation appears

first; the anion is named last.

Ligands are listed alphabetically before the metal. Prefixes denoting

the number of a particular ligand are ignored when alphabetizing.

Example

[Co(NH3)5Cl]Cl2 Pentaamminechorocobalt(III) chloride

cation anion

5 NH3 

ligandsCl‐

ligands

cobalt in +3 oxidation  states

The names of anionic ligands end in “o”; the endings of the

names of neutral ligands are not changed.

Prefixes tell the number of a type of ligand in the complex.

If th f th li d it lf h h fiIf the name of the ligand itself has such a prefix,

alternatives like bis-, tris-, etc., are used.

[Co(NH2CH2CH2NH2)2Cl2]+ dichlorobis(ethylenediammine)cobalt(III)

cationExample

2 Cl‐

ligands 2 en ligands with 2 NH2 groups

cobalt in +3 oxidation  states

en = ethylenediammine

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If the complex is an anion, its ending is changed to -ate.

The oxidation number of the metal is listed as a roman numeral

in parentheses immediately after the name of the metal.

Example

Na2[MoOCl4]

Exercise 1

Name the following coordination complexes:

(i) Cr(NH3)Cl3

(ii) Pt(en)Cl2(ii) Pt(en)Cl2

(iii) [Pt(ox)2]2-

Exercise 2

Give the chemical formular for the following coordination complexes:

(i) Tris(acetylacetanato)iron(III)

(ii) Hexabromoplatinate(2-)

(iii) Potassium diamminetetrabromocobaltate(III)

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(i) Cr(NH3)Cl3

Ammine

Solutions

chromium ammine chloro(III) tri

trichlorochromium(III)Ammine

(ii) Pt(en)Cl2

Dichloro

Platinum ethylenediammine chloro(II) di

trichlorochromium(III)

ethylenediammineplatinum(II)

(iii) [Pt(ox)2]2-

Dioxalato

Platinate oxalato(II) di

y p ( )

platinate(II)

(i) Tris(acetylacetanato)iron(III)

Fe(acac)3

Solutions

Fe acac3+ ( )3Fe(acac)3

(ii) Hexabromoplatinate(2-)

[PtBr6]2-

(ii) P t i di i t t b b lt t (III)

Pt Br [ ]2-6

(ii) Potassium diamminetetrabromocobaltate(III)

K[Co(NH3)2Br4]

K NH3 Br Co( )2 43+

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Isomers

Primarily in coordination numbers 4 and 6.

Arrangement of ligands in space and also the ligands themselves.

Types

Ionization isomers Isomers can produce different ions in solution

e.g. [PtCl2(NH3)4]Br2 [PtBr2(NH3)4]Cl2

Polymerization isomers

Same empirical formula or stoichiometry, but different molar mass.

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Different compounds with similar formula

[Co(NH3)3 (NO2)3 ]° ( n = 1)

[Co(NH3)6 ]3+ [Co(NO2)6 ]3− ( n = 2)

[Co(NH3)4 (NO2)2 ]+ [Co(NH3)2 (NO2 )4]− ( n = 2)

e.g.

[MXx Bb ]n

Hydration isomers exist for crystals of complexes containing water

molecules

exist in three different crystalline

Hydration isomers

exist in three different crystalline

forms, in which the number of

water molecules directly attached

to the Cr 3+ ion differs

[Cr(H2O)4 Cl2]Cl·2H2O dark green

e.g. CrCl3·6H2O

[Cr(H2O)5 Cl]Cl2·H2O light green

[Cr(H2O)6 ]Cl3 gray-blue

In each case, the coordination number of the chromium cation is 6

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Coordination isomers

[Co(NH3)6]3+ [Cr(CN)6]-3 and [Cr(NH3)6]+3 [Co(CN)6]-3

In compounds, both cation and anion are complex, the distribution of

ligands can vary, giving rise to isomers.

[ ( 3)6] [ ( )6] [ ( 3)6] [ ( )6]

Linkage isomers

e.g. Nitro and nitrito (a) [Co(NO2)(NH3)5]2+

Yellow

How the ligands arrangethemselves and attach to thecentral metal

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N or O coordinationpossible

(b) [Co(ONO)(NH3)5]2+Red

Geometric isomers

Formula is the same but the

arrangement in 3‐D space is

different.

e.g. square planar molecules give cis

and trans isomers.

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For hexacoordinate systems

Purple Green

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For M(X)3(Y)3 systems (e.g. octahedral)

there is facial andmeridian

F i l

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When three identical ligands occupy one face of an octahedron

any two identical ligands are adjacent or cis to each other

Facial

If these three ligands and the metal ion are in one plane

Meridian

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Co – Octahedral geometry

Example

cis-[CoCl2(NH3)4]+ trans-[CoCl2(NH3)4]+

fac-[CoCl3(NH3)3] mer-[CoCl3(NH3)3]

Are “stereoisomers” also possible?

An analogy to organic chirality.

molecules that have the same molecular formula and sequence

Stereoisomer

of bonded atoms (constitution), but which differ only in thethree-dimensional orientations of their atoms in space

Molecules which can rotate light.

Enantiomers

– non-superimposable

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non superimposable

mirror images

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Complex Stabilities

Generally in aqueous solution, for a given metal and

ligand, complexes where the metal oxidation state is +3

are more stable than +2

Generally the stabilities of complexes of the first row of Generally the stabilities of complexes of the first row of

transition metals vary in reverse of their cationic radii

MnII < FeII < CoII < NiII > CuII > ZnII

Hard and soft Lewis acid-base theory

• small atomic/ionic radius

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Hard acids and bases

tend to have:

• high oxidation state

• low polarizabilty

• high electronegativity

• hard bases - energy low-lying HOMO

• hard acids - energy high-lying LUMO

Chelate effect - is the additional stability of a complex

containing a chelating ligand, relative to that of a complex

containing monodentate ligands with the same type and number

of donors as in the chelate.

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[Cu(H2O)4(NH3)2]2+ + en [Cu(H2O)4(en)]2+ + 2 NH3

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Cu(H2O)4(NH3)2]2+ + en = [Cu(H2O)4(en)]2+ + 2 NH3

Mainly an entropy effect.

When ammonia molecule dissociates ‐ swept off in solution and

the probability of returning is remote.

When one amine group of en dissociates from complex ligand

retained by end still attached so the nitrogen atom cannot move

away – swings back and attach to metal again.

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away swings back and attach to metal again.

Therefore the complex has a smaller probability of dissociating.

Example

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Metal carbonyl

Compounds that have the metal bonded to the carbon

monoxide, giving a general formula of M(CO)n

M + CO M(CO)n

C OM ∏-orbitals in CO are very empty

Molecular orbital diagram (CO)

Therefore the bond order is:

4 – 1 = 3

Bond order: No. of e- pairs in the bonding orbital — No. of e- pairs in

the anti-bonding orbital

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Back-bonding (back donation)

Formation of ∏-bonding as a result of the overlap of metal d ∏-

orbitals and the ligand, CO, ∏* orbitals

Eff tEffects:

It enhances the bonding strength between the metal and the ligand.

The metal-ligand bond is shortened (M CO)

The becomes longer, weaker and the bond order decreases

Evidence and extent

C O

IR spectra – Vibration frequency

– The greater the extent of back bonding the lower the

stretching frequency (bond order decreases)

Free ≈ 2143 cm-1 M CO ≈ 1900 - 2125 cm-1

C O

C O

Effect of replacing the CO ligands

Non- ∏ accepting ligands (donor ligands)

Cr(CO) Cr(triens)(CO)3

TrienCr(CO)6 Cr(triens)(CO)3

2100 cm-1

2000 cm-1

1985 cm-1

1900 cm-1

1760 cm-1

Replacement of the 3 x (CO) groups with donor ligands (trien) increases ∏-

acidity of the remaining ligands (CO) so as to counter the accumulation of

the negative charge on the metal centre

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Effect of introducing a positive charge on metal complex

V(CO)6-

1 proton

1860 cm-1 2000 cm-1

V(CO)6 V(CO)6+

1 proton

2090 cm-1

Introducing a +ve charge on the metal inhibits shift of electrons from metal

to empty ∏*- orbital of the CO ligands

– This weakens ∏-bonding or decrease stretching frequencies of M-C

while the increases. (wave number or frequency increases)

1860 cm 1 2000 cm 1 2090 cm 1

C O

Thought

V(CO)- and Cr(CO) are isoelectronic yet

stretching frequencies of CO in V(CO)6

is lower than that of CO in Cr(CO)6 ?

The origin of colour - absorption

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The colour can change depending on a number of factorse.g.

Metal charge

Colours on coordination compounds

Ligand

Physical phenomenon

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Are there any simple theories to explain the colours in transition

metal complexes?

There is a simple electrostatic model used by chemists to

ti li th b d ltrationalize the observed results

This theory is called Crystal Field Theory

It is not a rigorous bonding theory but merely a simplistic

approach to understanding the possible origins of photo-

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pp g p g p

and electrochemical properties of the transition metal

complexes