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6/4/2011
1
APPLIED INORGANIC CHEMISTRY FOR CHEMICAL ENGINEERS
Transition Metal Chemistry
CHEM261HC/SS1/01
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+
22 23 24 25 26 27 28 29
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.
22 23 24 25 26 27 28 29
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
CHEM261HC/SS1/13
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
CHEM261HC/SS1/16
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.
CHEM261HC/SS1/17
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
CHEM261HC/SS1/18
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.
CHEM261HC/SS1/19
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For hexacoordinate systems
Purple Green
CHEM261HC/SS1/20
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
CHEM261HC/SS1/22
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
CHEM261HC/SS1/23
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.
CHEM261HC/SS1/24
[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.
CHEM261HC/SS1/25
away swings back and attach to metal again.
Therefore the complex has a smaller probability of dissociating.
Example
CHEM261HC/SS1/26
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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
CHEM261HC/SS1/27
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The colour can change depending on a number of factorse.g.
Metal charge
Colours on coordination compounds
Ligand
Physical phenomenon
CHEM261HC/SS1/29
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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-
CHEM261HC/SS1/30
pp g p g p
and electrochemical properties of the transition metal
complexes