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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 Transition metals

Main-group elements

Lanthanides

Actinides

1. Main-group elements

2. Transition metals

3. Lanthanides

4. Actinides

Transition metals vs. Main-group metals

There is some controversy about the classification of the elements

i.e. Zinc (Zn), Cadmium (Cd) and Mercury (Hg)

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 complexes

• form stable compounds with neutral molecules

• forms one or more stable ions which have incompletely filled d orbitals

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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: [Ar] 4s2 3d7

Co2+: [Ar] 3d7

Co3+: [Ar] 3d6

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

before they are removed from valence d orbitals when transition

metals are ionized.

Co has 27 electrons

[Ar] has 18 electrons

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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.

22 23 24 25 26 27 28 29

Ti V Cr Mn Fe Co Ni Cu

Cross off the first 2,

d2 d3 d4 d5 d6 d7 d8 d9

EXAMPLES

Elements Outer e- configuration

Sc 4s23d1

V 4s23d3

Cr 4s13d5

Fe 4s23d6

Ni 4s23d8

Cu 4s13d10

Zn 4s23d10

Half-filled or

Filled subshell

Pauli exclusion principle

Hand's rule

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Evaluating the oxidation state

[CoCl(NO2)(NH3)4]+

x = +3

x - 2 = +1

+1

Co3+

Neutralzero charge

Exercise

Question 1

(a) For [Co(Br)(NO2)(H2O)4]+, write

(i) the coordination number for Co _________ (1)

(ii) the oxidation state for Co ____________ (1)

(b) Write the outer electron configuration of

(i) Co ______________________ (1)

(ii) Mn2+ ________________________ (1)

School of ChemistryUNIVERSITY OF KWAZULU-NATAL, HOWARD COLLEGE

April 2008 TestCHEM261: APPLIED INORGANIC CHEMISTRY FOR CHEMICAL

ENGINEERSMarks : 40 Time : 45 minutesNAME : ____________STUDENT NO. : __________

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Why do these elements exhibit a variety of oxidation states?

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

The most prevalent oxidation numbers are shown in green.

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

Co +1 +2 +3 +4 +5

Ni +1 +2 +3 +4

Cu +1 +2 +3

Zn +2

Oxidation states and their relative stabilities

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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.

The stability of higher oxidation states decreases in moving from Sc

to Zn. Increase in effective nuclear charge across (L→R)

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

oxidation states of Ni, Cu and Zn are unknown.

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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.

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

M = [Ar]4s23dx

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

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

Example

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

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

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

Cr2O72− are strong oxidizer

Chromium

• Exists in solution as +2 or +3 state

• Elemental iron reacts with non-oxidizing acids

to form Fe2+, which oxidizes in air 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

solution

Iron

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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

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 two sites, a tridentate three sites etc.

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Ligands

molecular formula

Lewis base/ligand

Lewis acid

donor atom

coordination 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

ox2- Oxalato

EDTA4- Ethylenediamine-tetraacetato

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Chelating ligands bond to metal

Coordination numbers and geometries

Five or six atoms rings are common

(i.e. including metal)

forms rings – chelate rings

Linear

Square planar Tetrahedral Octahedral

The basic protocol in coordination nomenclature is to name the

ligands attached to the metal as prefixes before the metal name.

Nomenclature of Coordination Compounds

Some common ligands and their names are listed below.

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

)

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 the name of the ligand itself has such a prefix,

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

[Co(NH2CH2CH2NH2)2Cl2]+

Example

Dichlorobis(ethylenediammine)cobalt(III)

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

Exercise 1

Name the following coordination complexes:

(i) Cr(NH3)Cl3

(ii) Pt(en)Cl2

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

Exercise 2

Give the structures of the following coordination complexes:

(i) Tris(acetylacetanato)iron(III)

(ii) Hexabromoplatinate(2-)

(iii) Potassium diamminetetrabromocobaltate(III)

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Isomers

Primarily in coordination numbers 4 and 6.

Different arrangement of ligands in space and also can be the

ligands themselves.

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.

Different compounds with similar formula

[Co(NH3)3 (NO2)3 ]° ( n = 1)

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

[Co(NH3)6 ]3+ [Co(NO2)6 ] 3− ( n = 2)

e.g.

Types

[MXx Bb ]n

Hydration isomers exist for crystals of complexes containing water

molecules

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

[Cr(H2O)5 Cl]Cl2·H2O

[Cr(H2O)6 ]Cl3

e.g. CrCl3·6H2O

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

Hydration isomers

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Coordination isomers

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

Linkage isomers

e.g. Nitro and nitritoN or O coordination

possible

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

of ligands can vary, giving rise to isomers.

(a) [Co(NO2)(NH3)5]2+

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

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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.

Pt

Cl

Cl

NH3

NH3

cis-[PtCl2(NH3)2]

Pt

Cl

NH3

H3N

Cl

trans-[PtCl2(NH3)2]

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For hexacoordinate systems

GreenPurple

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For M(X)3(Y)3 systems there is

facial and meridional

Cis-[CoCl2(NH3)4]+ Trans-[CoCl2(NH3)4]+

Co – Octahedral geometryExample

Cis/Trans Vs. Fac/Mer

Cis-

[CoCl2(NH3)4]+

Trans-

[CoCl2(NH3)4]+

Fac-

[CoCl3(NH3)3]+

Mer-

[CoCl3(NH3)3]+

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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

transition metals vary in reverse of their cationic radii

MnII < FeII < CoII < NiII > CuII > ZnII

Properties of hard

acids and bases:

• small atomic/ionic radius

• high oxidation state

• low polarizabilty

• high electronegativity

• hard bases - energy low-lying HOMO

• hard acids - energy high-lying LUMO

80 76 74 69 71 74 pm

Hard and soft Lewisacid-base theory

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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.

[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

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.

Therefore, the chelate complex has a smaller probability of

dissociating. Thus, more stable

Mainly an entropy effect.

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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)

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

the anti-bonding orbital

What is the bond order ?

Back-bonding (back donation)

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

orbitals and the ligand, CO, ∏* orbitals.

Effects:

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

Infra red (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

C O

Effect of replacing the CO ligands

Non- ∏ accepting ligands (donor ligands)

Cr(CO)63NH3

2100 cm-1

2000 cm-1

1985 cm-1

1900 cm-1

1760 cm-1

Cr(NH3)3(CO)3

Replacement of the 3 x (CO) groups with donor ligands, 3 x (NH3)3

increases ∏-acidity of the remaining ligands (CO) so as to counter the

accumulation of the negative charge on the metal centre.

Metal-carbon (M─CO) bond enhanced while carbon-oxygen (C≡O) bond is

weakened, hence, lower wavenumbers on IR spectra.

Effect of introducing a positive charge on metal complex

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 C≡O increases. (wave number or frequency increases)

V(CO)6-

1 proton

1860 cm-1 2000 cm-1

V(CO)6 V(CO)6+

1 proton

2090 cm-1

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 ? Why?

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The origin of colour - absorption

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

Metal charge

Ligand

Colours on coordination compounds

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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

rationalize 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-

and electrochemical properties of the transition metal

complexes