Introduction to coordination chemistry

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Introduction to coordination chemistry
Office hours Paula Diaconescu Tu 430-530 pm
(MSB 1515) Th 12-1 pm (MSB 3515) Erin
Broderick We 6 7 pm (Geology 4607) Kevin
Miller Th 6 7 pm
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Outline

• The chelate effect
• Hard and soft acids and bases
• Geometries for TM complexes
• Isomerism
• Structural
• Stereoisomerism

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The chelate effect

• Comparison between the reaction of a chelating
  ligand and a metal ion with the corresponding
  reaction involving comparable monodentate
  ligands 2,2'-bipyridine with pyridine or
  1,2-diaminoethane (ethylenediamineen) with
  ammonia.
• Reaction of NH3 and en with Cd2
• ?H values for the formation steps are almost
  identical.
• The ?S term changes from negative (unfavorable)
  to positive (favorable). Note as well that there
  is a dramatic increase in the size of the ?S
  term for adding two compared to adding four
  monodentate ligands. (-5 to -35 JK-1mol-1).

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The chelate effect explanation

• One explanation is to count the number of species
  on the left and right hand side of the equations
  above for chelating ligands there is an increase
  in disorder.
• An alternative view comes from trying to
  understand how the reactions might proceed (watch
  movie). To form a complex with 6 monodentates
  requires 6 separate favorable collisions between
  the metal ion and the ligand molecules. To form
  the tris-bidentate metal complex requires an
  initial collision for the first ligand to attach
  by one arm but the other arm is always nearby and
  only requires a rotation of the other end to
  enable the ligand to form the chelate ring.
• Also, when one considers dissociation steps,
  then when a monodentate group is displaced, it is
  lost into the bulk of the solution. On the other
  hand, if one end of a bidentate group is
  displaced the other arm is still attached and it
  is only a matter of the arm rotating around and
  it can be reattached again.

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Size of the rings

• 5-member ring is more stable than 6- or 7-member
  rings for bis-amino complexes in which the two N
  atoms are linked by an aliphatic chain.
• 6-member rings can become the most stable,
  depending on the ligand.

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The hard/soft acid/base principle
Irving-Williams stability series (1953) for a
given ligand, the stability increases in the
following order Ba2 lt Sr2 lt Ca2 lt
Mg2 lt Mn2 lt Fe2 lt Co2 lt Ni2 lt Cu2 lt Zn2
Certain ligands formed their most stable
complexes with metal ions like Al3, Ti4 and
Co3 while others formed stable complexes with
Ag, Hg2 and Pt2.
1958 Ahrland et al.
Type A metal cations Alkali metal cations Li
to Cs Alkaline earth metal cations Be2 to
Ba2 Lighter transition metal cations in higher
oxidation states Ti4, Cr3, Fe3, Co3 The
proton, H Type B metal cations Heavier
transition metal cations in lower oxidation
states Cu, Ag, Cd2, Hg, Ni2, Pd2, Pt2.
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Type A/B ligands
Ligands were classified as type A or type B
depending upon whether they formed more stable
complexes with type A or type B metals

Type A metals prefer to bind to type A ligands.

and
Type B metals prefer to bind to type B ligands.
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Pearson's hard/soft Lewis acid/base principle
Ralph Pearson (1960s) Lewis acid Lewis
base -gt complex Pearson classified Lewis acids
and Lewis bases as hard, borderline, or soft.
According to Pearson's hard soft Lewis acid
base (HSAB) principle Hard Lewis acids
prefer to bind to hard Lewis bases and
Soft Lewis acids prefer to bind to soft
Lewis bases
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Pearsons acids
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Pearsons bases
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Characteristics of hard/soft acids/bases

• Hard Lewis bases Small, highly solvated,
  electronegative atomic centers 3.0-4.0
  Species are weakly polarizable Difficult to
  oxidize High energy HOMO
• Soft Lewis bases Large atoms of
  intermediate electronegativity 2.5-3.0 Easy
  to polarize and oxidize Low energy HOMOs but
  large magnitude HOMO coefficients

• Hard Lewis acids Atomic centers of small
  ionic radius High positive charge Species
  do not contain electron pairs in their valence
  shells Low electron affinity Likely to be
  strongly solvated High energy LUMO
• Soft Lewis acids Large radius Low or
  partial (d) positive charge Electron pairs in
  their valence shells Easy to polarize and
  oxidize Low energy LUMOs, but large magnitude
  LUMO coefficients

Note it is not necessary for species to possess
all properties.
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Shape of TM complexes

• Lone pairs exert less influence on geometry in TM
  cxs than in main-group compounds (VSEPR model).
• Rule the shape of the complex is dictated by the
  number of ligands. In general, the geometry of a
  complex is that in which the ligands are set as
  far apart as possible from each other.

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Common coordination geometries
6-Coordinate Octahedral (90 180 angles)
4-Coordinate Square Planar or Tetrahedral
(90 180) (109)
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Other coordination geometries
5-Coordinate Trigonal Bipyramidal or Square
Pyramidal (90 120) (100 90)
3-Coordinate Trigonal planar (120)
http//www.d.umn.edu/pkiprof/ChemWebV2/index2.htm
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Isomerism
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Structural isomerism

• Linkage isomerism occurs with ambidentate
  ligands. These ligands are capable of
  coordinating in more than one way. The best known
  cases involve the monodentate ligands SCN- / NCS-
  and NO2- / ONO-.
• Coordination isomerism where compounds
  containing complex anionic and cationic parts can
  be thought of as occurring by interchange of some
  ligands from the cationic part to the anionic
  part.
• Ionization isomerism where the isomers can be
  thought of as occurring because of the formation
  of different ions in solution.

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Solvate isomerism

• Note that both anions are necessary to balance
  the charge of the complex and that they differ in
  that one ion is directly attached to the central
  metal but the other is not. A very similar type
  of isomerism results from replacement of a
  coordinated group by a solvent molecule (solvate
  isomerism). In the case of water, this is called
  hydrate isomerism.
• Hydrate isomerism the best known example of this
  occurs for chromium chloride "CrCl36H2O" which
  may contain 4, 5, or 6 coordinated water
  molecules.
• These isomers have very different chemical
  properties and on reaction with AgNO3 to test for
  Cl- ions, would find 1, 2, and 3 Cl- ions in
  solution, respectively.

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Stereoisomerism

• Stereoisomers have the same atoms, same sets of
  bonds, but differ in the relative orientation of
  these bonds. Ignoring special cases involving
  esoteric ligands, then
• Geometric isomers are possible for both square
  planar and octahedral complexes, but not
  tetrahedral.
• Optical isomers are possible for both tetrahedral
  and octahedral complexes, but not square planar.
• The earliest examples of stereoisomerism involve
  complexes of Co3. In 1889, Jorgensen observed
  purple and green salts of CoCl2(en)2, which
  Werner later correctly identified as the cis- and
  trans- geometric isomers.
• In 1911, the first resolution of optical isomers
  was reported by Werner and King for the complexes
  cis-CoX(NH3)(en)22, where XCl- or Br-.

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Geometrical isomers
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Optical isomers

• Optical isomers are related as non-superimposable
  mirror images and differ in the direction with
  which they rotate plane-polarized light. These
  isomers are referred to as enantiomers or
  enantiomorphs of each other and their
  non-superimposable structures are described as
  being asymmetric.
• The two isomers have identical chemical
  properties and just denoting their absolute
  configuration does NOT give any information
  regarding the direction in which they rotate
  plane-polarized light. This can ONLY be
  determined from measurement and then the isomers
  are further distinguished by using the prefixes
  laevo ((-) or l) and dextro (() or d) depending
  on whether they rotate left or right. The use of
  l- and d- is not recommended.

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Optical isomers

• To add to the confusion, when measured at the
  sodium D line (589nm), the tris(1,2-diaminoethane)
  M(III) complexes (M Rh(III) and Co(III)) with
  IDENTICAL absolute configuration, rotate plane
  polarized light in OPPOSITE directions!
• The left-handed (?)-Co(en)33 isomer gives a
  rotation to the right and therefore corresponds
  to the () isomer.
• Note that, although it is predicted that
  tetrahedral complexes with 4 different ligands
  should be able to give rise to optical isomers
  (compare to C chemistry), in general they are too
  labile and cannot be isolated.

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Assignment

• Lecture 1 book 1, chapters 1 and 2.1 2.5
• Lecture 2 book 1, chapters 2.6 2.8, 3