2. IUPAC Nomenclature of co ordination compounds
• The cation is named first in both positively and negatively charged
coordination entities.
• The ligands are named in alphabetic order before the name of central
atom/ion.
• Names of anionic ligands end in-o, those neutral and cationic ligands are
same except aqua for H2O, ammine for NH3, carbonyl for CO, and nitosyl for
NO. while writing the formula of the coordination entity, these are enclosed
in brackets.
• Oxidation number of the metal is indicated by Roman numeral in brackets.
• If the complex is a cation, the metal is named the same as the element. If the
complex is an anion, the name of the metal ends with the suffix -ate.
13. FORMULAS OF MONONUCLEAR COORDINATION
ENTITIES:
The following rules are applied while writing the formulas:
• Central atom is listed first.
• Ligands are then listed in alphabetical order. The placement of a ligand in the list
does not depend on its charge.
• Polydentate ligands are also listed alphabetically. In case of abbreviated ligand, the
first letter of the abbreviation is used to determine the position of the ligand in the
alphabetical order.
• The formula for the entire coordination entity, whether charged or not, is enclosed
in square brackets. When ligands are polyatomic, their formulas are enclosed in
parentheses. Ligand abbreviations are also enclosed in parentheses.
14. • There should be no space between the ligands and the metal within a
coordination sphere.
• When the formula of a charged coordination entity is to be written without
that of the counter ion, the charge is indicated outside the square brackets
as a right superscript with the number before the sign. For example,
[Co(CN)6]3-, [Cr(H2O)6]3+, etc.
• The charge of the cation(s) is balanced by the charge of the anion(s).
15. PROPERTIES OF COORDINATION COMPOUNDS
• The coordination compounds formed by the transition elements are coloured due to the
presence of unpaired electrons that absorb light in their electronic transitions.
• For example, the complexes containing Iron(II) can exhibit green and pale green
colours, but the coordination compounds containing iron(III) have a brown or yellowish-
brown colour.
• When the coordination centre is a metal, the corresponding coordination complexes have
a magnetic nature due to the presence of unpaired electrons.
• Coordination compounds exhibit a variety of chemical reactivity. They can be a part of
inner-sphere electron transfer reactions as well as outer-sphere electron transfers.
• Complex compounds with certain ligands have the ability to aid in the transformation of
molecules in a catalytic or a stoichiometric manner.
16. TYPES OF COORDINATION COMPLEXES
based on whether complex ion is a cation/anion
• 1.Cationic complexes: In this co-ordination sphere is a cation.
Example: [Co(NH3)6]Cl3
• 2.Anionic complexes: In this co-ordination sphere is Anion.
Example: K4[Fe(CH)6]
• 3.Neutral Complexes: In this co-ordination sphere is neither cation
or anion. Example: [Ni(CO)4]
17. WERNER’S EXPERIMENT
• Werner conducted an experiment by mixing AgNO3(silver nitrate) solution
with CoCl3·6NH3, all three chloride ions got converted to AgCl (silver
chloride).
• However, when AgNO3 was mixed with CoCl3·5NH3, two moles of AgCl were
formed. Further, on mixing CoCl3·4NH3 with AgNO3, one mole of AgCl was
formed.
• Based on this observation, Werner’s theory was postulated explaining the
structure of coordination compounds.
18.
19. WERNER’S THEORY
• POSTULATES OF WERNER’S THEORY:
1.The central metal atom in the coordination compound exhibits two types of valences,
namely, primary and secondary linkages or valences.
2.Primary linkages are ionizable and are satisfied by the negative ions.
3. Secondary linkages are non-ionizable. These are satisfied by negative ions or neutral
molecules. Also, the secondary valence is fixed for any metal and is equal to its
coordination number.
20. 4. The ions bounded by the secondary linkages to the metal exhibit
characteristic spatial arrangements corresponding to different
coordination numbers.
24. Bonding in coordination compounds:
• It primarily the work of Linus Pauling
The postulates of valence bond theory:
• The central metal atom/ion makes available a number of vacant orbitals
equal to its coordination number. These vacant orbitals form covalent bonds
with the ligand orbitals.
VALENCE BOND THEORY (VB THEORY):
25. • A coordinate covalent bond is formed by the overlap of a vacant metal
orbital and filled ligand orbitals. This complete overlap leads to the
formation of a metal ligand, σ (sigma) bond.
• A strong covalent bond is formed only when the orbitals overlap to the
maximum extent. This maximum overlapping is possible only when the
metal vacant orbitals undergo a process called ‘hybridisation’.
• A hybridised orbital has a better directional characteristics than an
unhybridized one.
26.
27.
28. If it is a strong field ligand then the electrons in the metal atom get paired up but in weak field ligands they will
not pair up.
29.
30.
31.
32.
33.
34.
35. The following table gives the coordination number,
orbital hybridisation and geometry
Coordination
number
Types of
hybridization
Geometry
2 Sp Linear
4 sp3 Tetrahedral
4 dsp2 square planar
6 d2sp3 Octahedral
6 sp3d2 Octahedral
48. MAGNETIC MOMENT
A species having at least one unpaired electron, is said to be
paramagnetic.
• It is attracted by an external field. The paramagnetic moment is given
by the following spin-only formula.
• BM
• μs = spin-only magnetic moment , n=number of unpaired electrons
49. For metal ions with upto three electrons in the d orbitals, like Ti3+ (d1); V3+ (d2);
Cr3+ (d3); two vacant d orbitals are available for octahedral hybridisation with 4s
and 4p orbitals. The magnetic behaviour of these free ions and their coordination
entities is similar.
When more than three 3d electrons are present, the required pair of 3d orbitals for
octahedral hybridisation is not directly available (as a consequence of Hund’s rule).
Thus, for d4 (Cr2+, Mn3+), d5 (Mn2+, Fe3+), d6 (Fe2+, Co3+) cases, a vacant pair of d
orbitals results only by pairing of 3d electrons which leaves two, one and zero
unpaired electrons, respectively.
50. Inner orbital complexes with d2sp3 hybridisation
Outer orbital complexes with sp3d2 hybridisation
51.
52.
53.
54. LIMITATIONS OF VALENCE BOND THEORY:
• It involves a number of assumptions.
• It does not give a quantitative interpretation of magnetic data.
• It does not explain the colour exhibited by coordination compounds.
• It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of
coordination compounds.
• It does not make exact predictions regarding the tetrahedral and square planar structures of 4-
coordinate complexes.
• It does not distinguish between weak and strong ligands.
55. CRYSTAL FIELD THEORY (CFT)
• Main postulates of crystal field theory are
• The crystal field theory (CFT) is an electrostatic model which considers the metal-
ligand bond to be ionic arising purely from electrostatic interactions between the
metal ion and the ligand.
• Ligands are treated as point charges in case of anions or point dipoles in case of
neutral molecules.
56. • The five d orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they are degenerate.
• This degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the
metal atom/ion.
• However, when this negative field is due to ligands (either anions or the negative ends of dipolar
molecules like NH3 and H2O) in a complex, it becomes asymmetrical, and the degeneracy of the d
orbitals is lifted. It results in splitting of the d orbitals.
• The pattern of splitting depends upon the nature of the crystal field. Let us explain this splitting in
different crystal fields.
57.
58.
59.
60. In an octahedral coordination entity with six ligands surrounding the metal atom/ion,
there will be repulsion between the electrons in metal d orbitals and the electrons (or
negative charges) of the ligands. Such a repulsion is more when the metal d orbital is
directed towards the ligand than when it is away from the ligand.
CRYSTAL FIELD SPLITTING IN
OCTAHEDRAL COMPLEXES
61. Thus, the dx2 −y2 and dz2 orbitals which point towards the axes along the direction
of the ligand will experience more repulsion and will be raised in energy; and the
dxy, dyz and dxz orbitals which are directed between the axes will be lowered in
energy relative to the average energy in the spherical crystal field.
62.
63. Thus, the degeneracy of the d orbitals has been removed due to ligand
electron-metal electron repulsions in the octahedral complex to yield three
orbitals of lower energy, t2g set and two orbitals of higher energy, eg set.
This splitting of the degenerate levels due to the presence of ligands in a
definite geometry is termed as crystal field splitting and the energy
separation is denoted by ∆o (the subscript o is for octahedral).
64. Thus, the energy of the two eg orbitals will increase by (3/5) ∆o and
that of the three t2g will decrease by (2/5)∆o.
65. • The crystal field splitting, ∆o, depends upon the field produced by the ligand
and charge on the metal ion.
• Some ligands are able to produce strong fields(Pairing occurs) in which
case, the splitting will be large whereas others produce weak fields(electrons
will not pair) and consequently result in a small splitting of d orbitals.
High spin complex low spin complex
66.
67. • In general, ligands can be arranged in a series in the order of increasing field strength
as given below:
Such a series is termed as spectrochemical series.
It is an experimentally determined series based on the absorption of light by complexes
with different ligands.
68. 1)Weak field ligands
2)No pairing
3)Crystal field is less.
4)High spin complex
1)Strong field ligands
2)pairing occurs
3)Crystal field is more.
4)Low spin complex
69. For d4 ions, two possible patterns of electron distribution arise:
(i) the fourth electron could either enter the t2g level and pair with an existing electron,
or
(ii) it could avoid paying the price of the pairing energy by occupying the eg level.
Which of these possibilities occurs, depends on the relative magnitude of the crystal
field splitting, ∆o and the pairing energy, P (P represents the energy required for electron
pairing in a single orbital).
The two options are:
70. SPECTROCHEMICAL SERIES.
The arrangement of ligands in order of their increasing CFSE values is
known as spectrochemical series. The ligands with small CFSE values are
called weak field ligands, whereas those with large value of CFSE are called
strong field ligands.
71. CRYSTAL FIELD SPLITTING IN
TETRAHEDRAL COMPLEXES
• In tetrahedral coordination entity formation, the d orbital splitting is
inverted and is smaller as compared to the octahedral field splitting.
• For the same metal, the same ligands and metal-ligand distances, it
can be shown that ∆t = (4/9) ∆0.
• Consequently, the orbital splitting energies are not sufficiently large
for forcing pairing and, therefore, low spin configurations are rarely
observed.
• The ‘g’ subscript is used for the octahedral and square planar
complexes which have centre of symmetry.
• Since tetrahedral complexes lack symmetry, ‘g’ subscript is not used
with energy levels.
72.
73. 1) In tetrahedral coordination entity formation, the d orbital splitting is inverted and is smaller
as compared to the octahedral field splitting.
74. 2) The orbital splitting energies are not sufficiently large for forcing pairing and, therefore,
low spin configurations are rarely observed.
75. • The crystal field modal is successful in explaining the formation, structures,
colour and magnetic properties of coordination compounds to a large extent.
• However, from the assumptions that the ligands are point charges, it follows
that anionic ligands should exert the greatest splitting effect.
• The anionic ligands actually are found at the low end of the spectrochemical
series. Further, it does not take into account the covalent character of
bonding between the ligand and the central atom.
• These are some of the weaknesses of CFT, which are explained by ligand
field theory (LFT) and molecular orbital theory which are beyond the scope
of the present study.
Limitationsof CrystalFieldTheory
76. Colour in co ordination compounds
• The colour of the complex is complementary to that which is
absorbed.
• If green is absorbed by the complex, it appears red.
77. • In CFT, the colour in the co ordination compounds can be readily
explained.
• In [Ti(H2O)6]3+ Ti3+ is a 3d1 system and is in the t2g level.
• If light corresponding to energy of blue green region is absorbed by
the complex, it would excite electron from t2g level to eg
level.(t2g
1eg
0→t2g
0eg
1)
78. • The homoleptic carbonyls (compounds containing carbonyl ligands only) are formed
by most of the transition metals.
• These carbonyls have simple, well defined structures. Tetracarbonylnickel(0) is
tetrahedral, pentacarbonyliron(0) is trigonalbipyramidal while hexacarbonyl
chromium(0) is octahedral.
• Decacarbonyldimanganese(0) is made up of two square pyramidal Mn(CO)5 units
joined by a Mn – Mn bond.
• Octacarbonyldicobalt(0) has a Co – Co bond bridged by two CO groups.
Bonding in Metal Carbonyls
79.
80. • The metal-carbon bond in metal carbonyls possesses both σ and π character.
The M–C σ bond is formed by the donation of lone pair of electrons on the
carbonyl carbon into a vacant orbital of the metal.
• The M–C π bond is formed by the donation of a pair of electrons from a
filled d orbital of metal into the vacant antibonding π* orbital of carbon
monoxide.
• The metal to ligand bonding creates a synergic effect which strengthens the
bond between CO and the metal
