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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