1. Chemistry 481(01) Spring 2017 Instructor: Dr. Upali Siriwardane
Office: CTH 311 Phone
Office Hours:
M,W 8:00-9:00 & 11:00-12:00 am;
Tu,Th, F 9: :30 a.m.
April 4 , 2017: Test 1 (Chapters 1, 2, 3, 4)
April 27, 2017: Test 2 (Chapters  (6 & 7)
May 16, 2016: Test 3 (Chapters. 19 & 20)
May 17, Make Up: Comprehensive covering all Chapters

2. Chapter 7. An introduction to coordination compounds
The language of coordination chemistry
7.1 Representative ligands
7.2 Nomenclature
Constitution and geometry
7.3 Low coordination numbers
7.4 Intermediate coordination numbers
7.53Higher coordination numbers
7.6 Polymetallic complexes
Isomerism and chirality
7.7 Square-planar complexes
7.8 Tetrahedral complexes
7.9 Trigonal-bipyrmidal and square-pyramidal complexes
7.10 Octahedral complexes
7.11 Ligand chirality

3. Chapter 7. An introduction to coordination compounds
Thermodynamics of complex formation
7.12 Formation constants
7.13 Trends in successive formation constants
7.14 Chelate and macrocyclic effects
7.15 Steric effects and electron delocalization

4. Coordination compound
A compound formed from a Lewis acid and Lewis base.
A metal or metal ion acting Lewis acid (being an electron pair acceptor) and a atom or group of atoms with lone electron pairs Lewis base electron pair donor forms an adduct with dative or coordinative covalent bonds.
Ni(ClO4)2 (aq)+ 6NH3 → [Ni(NH3)6](ClO4)2 (aq)
The Lewis bases attached to the metal ion in such compounds are called ligands.

5. The coordination number (CN)
CN of a metal ion in a complex is defined as the number of ligand donor atoms to which the metal is directly bonded.
[Co(NH3)5Cl]2+
CN is 6, 1 chloride and 5 ammonia ligands each donating an electron pair.
For organometallic compounds. An alternative definition of CN would be the number of electron pairs arising from the ligand donor atoms to which the metal is directly bonded.

6. 1) What is a coordination compound?

8. Coordination sphere
Coordination sphere - the sphere around the central ion made up of the ligands directly attached to it. Primary and secondary coordination sphere.

9. Preparation of Complexes
[Fe(H2O)6] CN [Fe(CN)6] H2O
The figure at left shows cyanide ions (in the form of KCN), being added to an aq. solution of FeSO4.
Since water is a Lewis base, the Fe2+ ions were originally in the complex [Fe(H2O)6]2+
The CN- ions are driving out the H2O molecules in this substitution reaction that form the hexacyanoferrate(II) ion, [Fe(CN)6]4- .

10. Various Colors of d-Metal Complexes
The color of the complex depends on the identity of the ligands as well as of the metal..
Impressive changes of color often accompany substitution reactions.

13. Structures and symmetries
Six-coordinate complexes are almost all octahedral (a).
Four-coordinate complexes can be tetrahedral (b) or square planar (c).
(Square planar usually occurs with d8 electron configurations, such as in Pt2+ and Au3+.)

14. Representing Octahedral Shapes
Instead of a perspective drawing (a), we can represent octahedral complexes by a simplified drawing that emphasizes the geometry of the bonds (b).

15. Ligands
The Brønsted bases or Lewis base attached to the metal ion in such compounds are called ligands.
These may be
Simple ions such as Cl–, CN–
Small molecules such as H2O or NH3,
Larger molecules such as H2NCH2CH2NH2
N(CH2CH2NH2)3
Macromolecules, EDTA and biological molecules such as proteins.

16. Representative Ligands and Nomenclature Bidentate Ligands
Polydentate Ligands
Some ligands can simultaneously occupy more than one binding site.
Ethylenediamine (above) has a nitrogen lone pair at each end, making it bidentate. It is widely used and abbreviated “en”, as in [Co(en)3]3+.

18. Ethylenediaminetetraacetate Ion (EDTA)
EDTA4- is another example of a chelating agent. It is hexadentate.
This ligand forms complexes with many metal ions, including Pb2+, and is used to treat lead poisoning.
Unfortunately, it also removes Ca2+ and Fe2+ along with the lead.
Chelating agents are common in nature.

19. Porphyrins and phthalocyanins

20. Chelates
The metal ion in [Co(en)3]3+ lies at the center of the three ligands as though pinched by three molecular claws. It is an example of a chelate,
A complex containing one or more ligands that form a ring of atoms that includes the central metal atom.

21. Naming Transition Metal Complexes
Cation name first then anion name.
List first the ligands, then the central atom
The ligand names are made to end in -O if negative
Anion part of the complex ends in -ate
Eg. Cu(CN)64- is called the hexacyanocuprate(II) ion
The ligands are named in alphabetic order
Number of each kind of ligand by Greek prefix
The oxidation state of the central metal atom shown in parenthesis after metal name
Briding is shown with  ( -oxo)

22. Some Common Ligand Names

23. Names of Ligands (continued)

24. Coordination Sphere Nomenclature
Cationic coordination sphere
-ium ending
Anionic coordination sphere
-ate ending

25. Examples [Co(NH3)4Cl2]Cl: dichlorotetramminecobalt(III) chloride
[Pt(NH3)3Cl]2[PtCl4]: di(monochlorotriammineplatinum(II)) tetrachloroplatinate(II).
K3[Fe(ox)(ONO)4] :
potassium tetranitritooxalatoferrate(III)

26. Use bis and tris for di and tri for chelating ligands
[Co(en)3](NO3)2 :
tris(ethylenediamine)cobalt(II) nitrate
[Ir(H2O)2(en)2]Cl3
bis(ethylenediamine)diaquairidium(III) chloride
[Ni(en)3]3[MnO4] :
Tris(ethylenediamine)nickel(II) tetraoxomanganate(II)

27. Naming [Cu(NH3)4]SO4 tetraaminecopper(II) sulfate [Ti(H2O)6][CoCl6]
hexaaquatitanium(III) hexachlorocobaltate(III)
K3[Fe(CN)6]
potassium hexacyanoferrate(III)

28. 2) Give the formula of following coordination compounds a)Dichlorobis(ethylenediammine)nickle b) Potasium trichloro(ethylene)platinate(1-)

29. c) Tetrakis(pyridine)platinum(2+) tetrachloroplatinate(2-) d) Tetraamminebis(ethylenediamine) --hydroxo- -amidodicobalt(4+) chloride

30. 3) Give the names of following coordination compounds a) [Co(NH3)6]Cl3; b) trans-[Cr(NH3)4(NO2)2]+ ; c) K[Cu(CN)2] ; d) cis-[PtCl2(NH3)2] ; e) fac-[Co(NO2)3(NH3)3]Cl3

31. The Eta(h) System of Nomenclature
For for p bonded ligands number of atoms attached to the metal atom is shown by hn
(h5 -cyclopentadienyl) tricarbonyl manganese
tetracarbonyl (h3-allyl) manganese, Mn(C3H5)(CO)4

32. Isomers Both structural and stereoisomers are found.
The two ions shown below differ only in the positions of the Cl- ligand, but they are distinct species, with different physical and chemical properties.

33. 4)       What is the geometry and coordination number of compounds in the problem above?
a)       [Co(NH3)6]Cl3;
b)       trans-[Cr(NH3)4(NO2)2]+ ;
c)       K[Cu(CN)2] ;
d)       cis-[PtCl2(NH3)2] ;
e)       fac-[Co(NO2)3(NH3)3]Cl3

34. 5)  Draw the formula and find the BITE of following ligands.
a)       2,2'-bipyridine (bipy) ;
b)       terpy;
c)       cyclam;
d)       edta;

36. Ionization Isomers
These differ by the exchange of a ligand with an anion (or neutral molecule) outside the coordination sphere. [CoSO4(NH3)5]Br has the Br- as an accompanying anion (not a ligand) and [CoBr(NH3)5]SO4 has Br - as a ligand and SO42-as accompanying anion.

37. Ionization Isomers
The red-violet solution of [Co(NH3)5Br]SO4 (left) has no rxn w/ Ag+ ions, but forms a ppt. when Ba2+ ions are added.
The dark red solution of [CoSO4(NH3)5]Br (right) forms a ppt. w/ Ag+ ions, but does not react w/ Ba2+ ions.

38. Hydrate Isomers
These differ by an ex-change between an H2O molecule and another ligand in the coordination sphere.
The solid, CrCl3. 6H2O, may be any of three compounds.
[Cr(H2O)6]Cl3 (violet)
CrCl(H2O)5]Cl2.H2O (blue-green)
CrCl2 (H2O)4Cl.2H2O (green)
Primary and secondary coordination spheres

39. Other ligands capable or forming linkage isomers are
The triatomic ligand is the isothiocyanato, NCS-. In (b) it is the thiocyanato, SCN-.
Other ligands capable or forming linkage isomers are
NO2- vs. ONO -
CN - vs. NC - .
(a) NSC- ligand (the N is closest to the center); (b) SCN- ligand (S is closest the center)

40. Coordination Isomers
These occur when one or more ligands are exchanged between a cationic complex and an anionic complex.
An example is the pair [Cr(NH3)6][Fe(CN)6] and[Fe(NH3)6][Cr(CN)6].

41. Stereoisomers
Ionization, hydrate, linkage, and coordination isomers are all structural isomers.
In stereoisomers, the formulas are the same. The atoms have the same partners in the coordination sphere, but the arrangement of the ligands in space differs.
The cis- and trans- geometric isomers shown in next slide differ only in the way the ligands are arranged in space.
There can be geometric isomers for octahedral and square planar complexes, but not for tetrahedral complexes.

42. Square Planar Complexes Geometric Isomers
Properties of geometric isomers can vary greatly.
The cis- isomer below is pale orange-yellow, has a solubility of g/100 g water, and is used for chemotherapy treatment.
The trans- isomer is dark yellow, has a solu-bility of g/100 g water, and shows no hemotherapeutic effect.

43. 6)  Describe the geometrical isomerism in following compounds:
a)       [Co(NH3)4Cl2]+ ;
b)       [IrCl3(PPh3)3] ;
c)       [Cr(en)2Cl2] ;

44. cis and trans-PtCl2(NH3)2

45. Trans Effect & Influence

46. Preparation Geometrical Isomers

47. Note only four of the six ligands are different.
Optical Isomerism
The two complexes at left are mirror images. (The gray rectangle represents a mirror, through which we see somewhat darkly.)
No matter how the complexes are rotated, neither can be superimposed on the other.
Note only four of the six ligands are different.

48. Combined Stereoisomerisms
Both geometrical and optical isomerism can occur in the same complex, as below. The trans- isomer is green.
The two cis- isomers, which are optical isomers of each other, are violet.

50. Identifying Optical Isomerism
If a molecule or ion belong to a point group with a Sn axis is not optically active

52. Molecular Polarity and Chirality Polarity
Polarity:Only molecules belonging to the point groups Cn, Cnv and Cs are polar. The dipole moment lies along the symmetry axis formolecules belonging to the point groups Cn and Cnv.
Any of D groups, T, O and I groups will not be polar

53. Chirality Only molecules lacking a Sn axis can be chiral.
This includes mirror planes
and a center of inversion as
S2=s , S1=I and Dn groups.
Not Chiral: Dnh, Dnd,Td and Oh.

54. Optical Activity

55. Reactions of Metal Complexes
Formation constants
– the chelate effect
– Irving William Series
– Lability

56. 7) Pick the chiral compounds among the following:
a)       [Co(en)3]3+ ;
b)       cis-[Cr(en)2Cl2] ;
c)       c) trans-[Cr(en)2Cl2] ;

57. Formation of Coordination Complexes
typically coordination compounds are more labile or fluxional than other molecules X is leaving group and Y is entering group
MX + Y MY + X
One example is the competition of a ligand, L for a coordination site with a solvent molecule such as H2O
[Co(OH2)6]2+ + Cl [Co(OH2)5Cl]+ + H2O

58. Formation Constants
Consider formation as a series of formation equilibria:
Summarized as:

59. Values of Kn Typically: Kn>Kn+1
Expected statistically, fewer coordination sites
available to form MLn+1
eg sequential formation of Ni(NH3)n(OH2)6-n 2+

60. Breaking the Rules
Order is reversed when some electronic or chemical change drives formation
Fe(bipy)2(OH2)22+ + bipy Fe(bipy)32+
jump from a high spin to low spin complex
Fe(bipy)2(OH2)2 t2g4eg high spin
Fe(bipy) t2g low spin

63. Irving William Series
Values of log Kf for 2+ ions including transition metal species Lewis acidity (acceptance of e-) increases across the per. table, thus forming more and more stable complexes for the same ligand system
Kf series for transition metals:
Mn2+< Fe2+ < Co2+ < Ni2+ < Cu2+ >Zn2+

64. Irving William Series

65. Bonding and electronic structure
Bonding Theories of Transition Metal Complexes
Valance Bond Theory
Crystal Field Theory
Ligand Field Theory or Molecular Orbital Theory

66. Valance Bond Theory
”Outer orbital" (sp3d2) and ”Inner orbital" (d2sp3)
[CoF6]3- - Co3+ : d6
[Co(NH3)6]3+ - Co3+ : d6

67. Spectrochemical Series for Ligands
It is possible to arrange representative ligands in an order of increasing field strength called the spectrochemical series:
I¯ < Br¯ < -SCN¯ < Cl¯ < F¯ < OH¯ < C2O42¯ < H2O < -NCS¯ < py < NH3 < en < bipy < o-phen < NO2¯ < CN¯ < CO

68. 8) Use valence bond theory (VBT) to predict the electron configurations, the type of bonding (Inner and outer orbital) and number of unpaired electrons in following compounds:
a)       [Co(CN)6]3- ;
b)       [CoCl6]3-;
c)       [Fe(NH3)6]3+;

69. Crystal Field Theory In the electrical fields created by ligands
The orbitals are split into two groups: a set consisting of dxy, dxz, and dyz stabilized by 2/5Do, known by their symmetry
classification as the t2g set, and a set consisting of the dx2-y2 and dz2, known as the eg set, destabilized by 3/5Do where Do is the gap between the two sets.

70. Crystal Field Splitting of d Orbitals

71. Octahedral Crystal Field Splitting

72. 9)  What are the symmetry labels of s,p, and d orbitals in tetrahedral (NiCO)4) and square-planar ([PtCl4]2-) and octahedral (Cr(CO)6) compounds.

73. 10) Explain the effect of ligands on the d orbitals in octahedral, tetrahedral, trigonal-bipyramid and square-planar coordination compounds using Crystal Field Theory. Octahedral, Tetrahedral, Trigonal-bipyramid Square-planar

74. 11) [Ti(H2O)6]3+ shows a absorption at 20300 cm-1
11) [Ti(H2O)6]3+ shows a absorption at cm-1. Absorption values for similar coordination compounds of Ti3+ with different ligands are given below. Based on their absorption values arrange the following ligands in a Spectrochemical Series. Absorption(cm-1) Ligand H2O CN- PPh3 F- NH

75. Crystal Field Stabilization Energy
Crystal Field stabilization parameter Do

76. Crystal Field Stabilization Energy
d7 case.
Weak field case
The configurations would be written t2g5 eg2
5(-2/5Do) + 2(+3/5Do) = -4/5Do
Strong field case
The configurations would be written t2g6 eg1
6(-2/5Do) + 1(+3/5Do) = -9/5Do

77. CFSE & Paring Energy
[Fe(H2O)6]2+. Iron has a d6 configuration, the value of Do is 10,400 cm-1 and the pairing energy is 17600cm-1. (1 kJ mol-1 = cm-1.) We must compare the total of the CFSE and the pairing energy for the two possible configurations.

78. high spin (more stable)
CFSE = 4 x -2/5 x x 3/5 x = -4160cm-1 ( kJ mol-1)
Pairing energy (1 pair) = 1 x = cm-1 (50.32 kJ mol-1
Total = cm-1 (38.43 kJ mol-1)
low spin
CFSE = 6 x -2/5 x 10400= cm-1 ( kJ mol-1)
Pairing energy (3 pairs) = 3 x = (151.0 kJ mol-1)
Total = cm-1 (79.60 kJ mole-1)

79. Tetrahedral complexes
Splitting order or reversed. eg is now lower energy and t2g is hgher energy
Because a tetrahedral complex has fewer ligands, the magnitude of the splitting is smaller. The difference between the energies of the t2g and eg orbitals in a tetrahedral complex (t) is slightly less than half as large as the splitting in analogous octahedral complexes (o)
Dt = 4/9Do

80. Tetrahedral Ligand Arrangement
Dt = 4/9Do
Mostly forms high spin complxes

81. Octahedral Crystal Field Splitting

82. Square-planar Complexes-D4h

83. Generalizations about Crystal Field Splittings
The actual value of D depends on both the metal ion and the nature of the ligands:
The splitting increases with the metal ion oxidation state. For example, it roughtly doubles going from II to III.
The splitting increases by % per period down a group.
Tetrahedral splitting would be 4/9 of the octahedral value if the ligands and metal ion were the same.

84. Spectrochemical Series for Ligands
It is possible to arrange representative ligands in an order of increasing field strength called the spectrochemical series:
I¯ < Br¯ < -SCN¯ < Cl¯ < F¯ < OH¯ < C2O42¯ < H2O < -NCS¯ < py < NH3 < en < bipy < o-phen < NO2¯ < CN¯ < CO

85. Spectrochemical Series for Metals
It is possible to arrange the metals according to a spectrochemical series as well. The approximate order is
Mn2+ < Ni2+ < Co2+ < Fe 2+ < V2+ < Fe3+ < Co3+ < Mn3+ < Mo3 + < Rh3 + < Ru3 + < Pd4+ < Ir3+ < Pt 3+

86. Spectrum of [Ti(H2O)6]3+.
d1: t2g1eg0 –> t2g0eg1

87. Hydration Enthalpy.
M2+(g) + 6 H2O(l) = [M(O2H)6]2+(aq)

88. Irving-Williams Series

89. Ligand Field Splitting and Metals
the transition metal also impacts Do increases with increasing oxidation number
Do increases as you move down a group (i.e. with increasing principal quantum number n)

90. MO forML6 diagram Molecules0

91. Ligand Field Stabilization Energies
LFSE is a function of Do
weighted average of the splitting due to the
fact that they are split into groups of 3 (t2g)
and 2 (eg)

92. Weak Field vs. Strong Field
now that d orbitals are not degenerate how do we know what an electronic ground state for a d metal complex is? need to determine the relative energies of pairing vs. Do

93. Splitting vs. Pairing when you have more than 3 but fewer than 8 d
electrons you need to think about the relative merits
pairing vs. Do
• high-spin complex – one with maximum number of unpaired electrons
• low-spin complex – one with fewer unpaired electrons

94. Rules of Thumb for Splitting vs Pairing
depends on both the metal and the ligands
• high-spin complexes occur when o is small Do is small when:
• n is small (3 rather than 4 or 5)– high spin only really for 3d metals
• oxidation state is low– i.e. for oxidation state of zero or 2+
• ligands is low in spectrochemical series– eg halogens

95. Four Coordinate Complexes: Tetrahedral
Same approach but different set of orbitals with different ligand field
• Arrangement of tetrahedral field of point charges results in splitting of energy where dxy, dzx, dyz are repelled more by Td field of negative charges
• So the still have a split of the d orbitals into triply degenerate (t2) and double degenerate (e) pair but now e is lower energy and t2 is higher.

96. Tetrahedral Crystal Field Splitting

97. Ligand Field Splitting: Dt
describes the separation between reviouslydegenerate d orbitals
• Same idea as Do but Dt < 0.5 Do for comparable systems
• So …Almost Exclusively Weak Field

98. Electron configurations in octahedral fields
Weak field and strong fieled cases

99. Tetragonal Complexes Start with octahedral geometry and follow the
energy as you tetragonally distort the octahedron
Tetragonal distortion: extension along z and
compression on x and y
Orbitals with xy components increase in
energy, z components decrease in energy
Results in further breakdown of degeneracy
– t2g set of orbitals into dyz, dxz and dxy
– eg set of orbitals into dz2 and dx2-y2

100. Tetragonal Complexes

101. Square Planar Complexes
extreme form of tetragonal distortion
Ligand repulsion is completely removed from
z axis
Common for 4d8 and
5d8 complexes:
Rh(I), Ir(I)
Pt(II), Pd(II)

102. Jahn Teller Distortion
geometric distortion may occur in systems
based on their electronic degeneracy
This is called the Jahn Teller Effect:
If the ground electronic configuration of a
nonlinear complex is orbitally degenerate, the
complex will distort to remove the degeneracy
and lower its energy.

103. Jahn Teller Distortions
Orbital degeneracy: for octahedral geometry
these are:
– t2g3eg1 eg. Cr(II), Mn(III) High spin complexes
– t2g6eg1 eg. Co(II), Ni(II)
– t2g6eg3 eg. Cu(II)
basically, when the electron has a choice between one of the two degenerate eg orbitals, the geometry will distort to lower the energy of the orbital that is occupied.
Result is some form of tetragonal distortion

104. Ligand Field Theory
Crystal field theory: simple ionic model, does not accurately describe why the orbitals are raised or lowered in energy upon covalent bonding.
• LFT uses Molecular Orbital Theory to derive the ordering of orbitals within metal complexes
• Same as previous use of MO theory, build ligand group orbitals, combine them with metal atomic orbitals of matching symmetry to form MO’s

105. LFT for Octahedral Complexes
Consider metal orbitals and ligand group orbitals
Under Oh symmetry, metal atomic orbitals transform as:
Degeneracy Mulliken Label Atomic Orbital
eg dx2-y2, dz2
t2g dxy, dyz, dzx
t1u px, py, pz
a1g s

106. Sigma Bonding: Ligand Group Orbitals

107. Combinationsof Metal andLigand SALC’s

108. Molecular Orbital Energy Level Diagram: Oh

109. PI Bonding pi interactions alter the MOELD that results from
sigma bonding
• interactions occur between
frontier metal orbitals and the
pi orbitals of L
• two types depends on the ligand
–pi acid - back bonding accepts e- density from M
–pi base -additional e- density donation to the M
• type of bonding depends on relative energy level
of pi orbitals on the ligand and the metal orbitals

110. : PI Bases and the MOELD Oh
pi base ligands
contribute more
electron density to
the metal
• t2g is split to form a
bonding and
antibonding pair of
orbitals
Do is decreased
• halogens are good
pi donors

111. PI Acids and the MOELD: Oh
• pi acids accept electron
density back from the
metal
• t2g is split to form a
bonding and antibonding
pair of orbitals
• the occupied bonding
set of orbitals goes
down in energy so ..
• Do increases
• typical for phosphine
and carbonyl ligands

112. Magnetic Properties of Atoms
a) Diamagnetism?
Repelled by a magnetic field due to paired electrons. b)Paramagnetism?
attracted to magnetic field due to un-paired electrons.
c) Ferromagnetism?
attracted very strongly to magnetic field due to un-paired electrons.
d)Anti-ferromagnetic?
Complete cancelling of unpaired electrons in magnetic domains

113. Magnetic Suceptibility Vs Temperature

114. Types of magnetism

115. Magnetic Properties
A paramagnetic substance is characterised experimentally by its (molar) magnetic susceptibility, cm. This is measured by
suspending a sample of the compound under a sensitive balance between the poles of a powerful electro-magnet,

116. Number of Unparied Electrons
The magnetic moment of the substance is given by the Curie Law:
m = (cmT)½ (in units of Bohr magnetons)
The formula used to calculate the spin-only magnetic moment can be written in two forms
m= n(n+2) B.M.

117. Magnetic Properties of Atoms
Paramagnetism?
Ferromagnetism?
Diamagnetism?
Gouvy Balance

118. Octahedral Complexes

119. Tetrahedral Complexes