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2P32 – Principles of Inorganic Chemistry
Dr. M. Pilkington
Lecture 2 - Introduction to Metal Complexes
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Metal Complexes: What are they?
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Why are so many transition metals six-coordinate?
Werner’s Coordination Theory.
Geometry of six-coordinate complexes, geometric
isomers?
Assignment 1 due in next Monday at 4.30pm, please write your name and
student ID on your work and staple all loose sheets together.
Transition Metals – located in the d-block of the periodic
table
alkali
metals
C, H, N, O, halogens
s-block
p-block
transition metals
d-block
d-block their d orbitals are filling
focus for coordination chem.
(f-block) inner transition metals
Rows are across are called Periods. Columns down are called Groups
Transition Elements begin in the 4th period
K
Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga
Ca
1
2
[Ar]4s [Ar]4s
d-block
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Ground State Electronic Configuration for the first
row Transition Metals
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Sc - Scandium
Ti – Titanium
V-Vanadium
Cr – Chromium
Mn – Manganes e
Fe – Iron
Co – Cobalt
Ni – Nickel
Cu – Copper
Zn – Zinc
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Ni
Pd
Pt
[Ar] 3d14s2
[Ar] 3d2 4s2
[Ar] 3d3 4s2
[Ar] 3d5 4s1
[Ar] 3d5 4s2
[Ar] 3d6 4s2
[Ar] 3d7 4s2
[Ar] 3d8 4s2
[Ar] 3d10 4s1
[Ar] 3d10 4s2
11
Cu
Ag
Au
12
Zn
Cd
Hg
Pd – Palladium
Pt – Platinum
Ag – Silver (Argentum)
Au – Gold (Aurum)
Cd – Cadmium
Hg – Mercury (hydrargyram) “liquid silver”
ƒ Most transition metals have several oxidation states.
ƒ Mn exists in 11 oxidation states -3 upto +7
ƒ Transition metals and their compounds are usually brightly
colored.
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“Metal Complexes” (coordination compounds)
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Origin of the name originates from the 1800’s, it was known
metal had “valencies” (oxidation numbers) that could be
satisfied by combination with elements having opposite
valencies.
Cr3+ valence of +3
O2- valence of -2
Cl- valence of -1
Examples of metal complexes CrCl3, Cr2O3
BUT CrCl3 reacts with ammonia (NH3) to form a new compound.
CrCl3 + 6NH3 → CrCl3.6NH3
Called a complex because nobody at the time
understood how they formed.
Alfred Werner – late 1800’s the father of coordination
chemistry.
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Studied in Switzerland at the University of Zurich.
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He lectured in both organic and inorganic chemistry.
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He developed the theory of coordination chemistry.
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He prepared and studied coordination compounds and
discovered optically active forms of 6-coordinate
octahedral complexes.
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His coordination chemistry extended through a whole
range of systematic inorganic chemistry and into
organic chemistry and he was awarded the Nobel
Prize in Chemistry in 1913.
Nobel Lecture
http://nobelprize.org/chemistry/laureates/1913/werner-lecture.html
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Werner studied the following metal complexes:
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CoCl3 forms four different compounds with NH3.
CoCl3.6NH3 AgNO3 (excess)
(1:6 mol ratio)
Yellow
CoCl3.5NH3
Lavender
AgNO3 (excess)
3AgCl (ppt)
2AgCl (ppt)
One Cl does not react
CoCl3.4NH3
Green
CoCl3.4NH3
Violet
AgNO3 (excess)
1AgCl (ppt)
AgNO3 (excess)
1AgCl (ppt)
Werner’s Conclusions:
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2.
3.
4.
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In this series of compounds, cobalt has a constant
coordination number of 6 (coordination number is the
number of groups that can bond directly to the metal).
As the NH3 molecules are removed they are replaced by
Cl- which acts as if it is covalently bonded to cobalt.
Chloride and Ammonia are now called ligands.
Ligands are a Lewis base/electron pair donors that can
bind to a metal ion.
A metal complex – metal ion combined with ligands.
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6.
Coordination complexes are neutral and counter ions are not
bonded to the central metal ion but balance the charge.
For example:
+3
0
-3
[Co(NH3)6]Cl3
counter ions
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The ligands directly coordinated to the metal are contained
within the square bracket.
Six NH3 bonded to Co.
3 chloride ions are not bonded to the Co these are counter ions,
they balance the charge (Co3+) they are “free” to react with
AgNO3 to give 3 moles of AgCl.
React with 3 moles of AgNO3
H2O
[Co(NH3)6]Cl3 →
Yellow
[Co(NH3)6]3+ + 3Cl-
React with 2 moles of AgNO3
[Co(NH3)5Cl]Cl2 rewrite as [Co(NH3)5Cl]2+ + 2ClLavender (now only two reactive Cl-’s).
React with 1 mole of AgNO3
[Co(NH3)4Cl2]Cl rewrite as [Co(NH3)4Cl2]+ + ClTwo isomers green and violet
Isomers - have the same formula but different structures, i.e.
different spatial arrangements
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What is the geometry of [Co(NH3)4Cl2]Cl ?
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Consider Hexagonal Planar
Cl
Cl
Cl
H3N
H3N
Co
Co
H3N
NH3
Cl
NH3
Co
Cl
H3N
NH3
H3N
NH3
NH3
1,2 isomer
ortho
1,3 isomer
meta
H3N
NH3
Cl
1,4 isomer
para
There are three possibilities so this does not fit with
the observations
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Consider Trigonal Prism
Cl
Cl
Co
Cl
Cl
or
Co
or
Co
Cl
Cl
There are three possibilities so this does not fit with
the observations
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Consider Octahedral
http://www.iumsc.indiana.edu/morphology/symmetry/octahedral.html
6 vertices, 8 sides
M
Metal ion in the center, ligands are on the vertices, all six
vertices are identical.
Geometry of [Co(NH3)4Cl2]Cl is Octahedral
Cl
Co
Cl
0
180
or
900
Cl
Co
Cl
trans isomer the two Cl
ligands are far apart (1800)
cis isomer - the two Cl
ligands are close to each other (900)
Many other examples of complexes of this coordination geometry
known.
This geometry reduces the steric crowding that is a problem in
other geometries and makes them unfavourable.
We accept Werner’s conclusions, today further evidence to
confirm his conclusions is provided by X-ray crystallography.
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Why is the coordination number 6 so common?
We have to consider the sizes (ionic radii) of the metal ions and
ligands.
Plot the ionic radii of transition metal ions, most of them are in
the range 75-90 pm which can accommodate 6 ligands and hence
favours 6-coordination.
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Consider the size of common ligands such as O as in H2O, N as
in NH3 or S. For example:
O2-
126 pm
126 pm
126 pm = radius of O (1 pm = 10-12 m)
Radius Ratio Rules
The structures of many crystals can be rationalised to a first approximation
by considering the relative sizes and numbers of ions present. The radius
ratio r+/r- can be used to make a first guess at the likely coordination
number and geometry around the cation using a set of simple rules:
Value of r+/r-
Predicted
Coordination
Number of Cation
Predicted Coordination
Geometry of
Cation
< 0.15
2
Linear
0.15-0.22
3
Trigonal Planar
0.22-0.41
4
Tetrahedral
0.41-0.73
6
Octahedral
> 0.73
8
Cubic
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To Summarize:
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For [M(H20)6]n+ in order for the metal ion to accommodate six ligands it
must have at least a 52 pm radius.
Six waters fit exactly around a metal ion with 52 pm radius (not
crowded).
At 92 pm’s the six waters have moved far enough apart that more
waters can fit.
If the metal ion is smaller than 52pm you will fit fewer H2O’s, if it is
larger then you have more space and can go to a larger coordination
number.
Hence transition metal ions display a number of coordination numbers
but 6 is very common.
N and O donors are the most abundant in biology (amino acids), S is also
present but it is allot larger than O and as a consequence Fe fits 6 O’s
but only 4 S’s. This is apparent in thiocluster compounds of Fe.
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