Download The electronic spectra of the complex [Cr(NH3)

Advanced Inorganic Chemistry
By Dr. Mahmou N.
Electronic Spectra
of Coordinated Compounds
The way of how to analyze the electronic spectra of complexes improve our
understanding of their bonding. First of all we shall start with the electronic spectra of
atoms.
Electronic Spectra of Atoms: Electronic configuration for elements give us an idea about
the number of electrons in each orbital , but it doesn't give us the arrangements of
electrons in atoms. For configuration 2p2, as an example, the two electrons might occupy
any of p orbitals with different orientations which means that we may have several
different states of total orbital angular momentum, and each on corresponds to an
occupation of orbitals with different values of ml and ms and this case called microstates.
For the notation s, p, d, …. for orbitals with l = 0,1,2,…. , the total orbital angular
momentum of an atomic term is donated by the upper case equivalent :
L = 0 1 2 3 4 …….
S P D F G ( then alphabetical omitting J )
The total spin quantum number (S ) of an atom is normally reported as the
multiplicity of the term, the value of 2S +1 :
S = 0 ½ 1 1½ 2 ……
2S +1 = 1 2 3 4 5 ……
The process of combining electron angular momenta by summing first the spin, then
the orbital momenta and finally the two resultants is called Russell – Saunders Coupling
The multiplicity is written as a left superscript on the letter representing the value of
L, and the entire label of a term is called a term symbol . Thus, the term symbol 3P
donates a term with L = 1 and S =1 an is called a triplet term.
Microstates that correspond to different relative spatial distributions of
electrons have different energies. For 2p2, as an example, there are 15 possible
microstates configurations, while for 3d2 there are 45 possible microstates ?, so we
can conclude that the terms of a 3d2 configuration are 1G, 3F, 1D, 3P and 1S which
account for all 45 permitted states. The Pauli principle restricts the microstates that
can occur in configuration, and consequently it affects the terms that can occur.
The energies of the terms: For a given configuration, the values of S and L can be
calculated and it is possible to predict the ground term by using Hund`s rules .
There it was expressed as , the lowest energy configuration is achieved if the
electrons are parallel, so we can write the following two statements ;
1- For a given configuration, the term with the greatest multiplicity lies lowest
in energy.
2- For a term of given multiplicity, the greatest the value of L, the lower the
energy.
These two rules give us that, in 3d2 the configuration 3F is lower in energy
than 3P, so for Ti+2 the ground term is 3F. Hund`s rules are reasonably reliable for
predicting which term has lower energy( the ground state), but are not very reliable
for predicting the order in which terms of higher energy lie. Thus, for Ti+2 the rules
predict the order
3F < 3P < 1G < 1D < 1S
but the observed order is
3F < 1D < 3P < 1G < 1S
The spin multiplicity rule is fairly reliable for predicting the ordering of n
terms, but the greatest L rule is reliable only for predicting the ground terms,
there is generally little correlation of L with the order of the higher terms. The
procedure for predicting the ground term of an atom or ion may be summarized
as follows:
1- Identify the microstate that has the highest value of multiplicity.
2- Identify the highest permitted value of L for that multiplicity.
Electronic Spectra of Complexes: The region for the electronic spectra of the
complexes usually applied at the range 200 – 800nm. For transition metal
complexes we can identify a maximum of four types of transitions as follows:
1- Absorption bands attributed to π – π* and n – π* transitions for the
organic part of the complexes( ligands), located at the ultraviolet region.
2- Absorption bands attributed to charge transfer transition ( metal to ligand
or ligand to metal), located between last part of ultraviolet and first part of
visible region.
3- Absorption bands attributed to d – d transition of the metal ion located at
the visible region.
4- Absorption bands attributed to the counter ion ( if presents), located at the
same area of the charge transfer absorption bands.
Absorption bands for 1 and 2 seems to be very strong while it is medium for
4 and weak for 3, why?. Any complex should had two absorption bands ( 2 and 3)
while some complexes could had the other two or one of them, why? .
Let us take the simplest possible case with a d1 configuration, [Ti(H2O)6]+3 the
d electron will occupy a t2g orbital. On irradiation with light of frequency c, equal
to ∆o/h, where h is Plank`s constant. The d electron captured the quantum of
radiation and excited from t2g to the eg orbital. The absorption bands found to
be in the visible spectrum of complex and is responsible for the violet color. Three
features of this absorption band are of importance, the position, the intensity and
the breadth. Fig. 1
The position of the band is related to the splitting of the d orbitals, the
spectrum tell us that ∆o in the complex is 20.000cm-1. Since there are 83.7cm-1 per
kilojoules, that means the splitting energy is about 250kJ per mole.
The intensity of the absorption band is extremely weak compared to the other
bands because the transition involved in this band is in the same orbital (d – d )and
according to the Laporta`s rule it is forbidden ( in tetrahedral complexes the
absorption band is greater with the factor 10* , why?).
The absorption band seems to be a broad band due to the vibration of the
complex.
Complexes with a d9 configuration ( CuII) will have a same spectrum as d1
although it is not simple as Ti( III) complex, several nearly superposed bands, due
to the distortion of the octahedral complex ( Jahn – Teller effect) .
Fig. 1 The visible absorption spectrum of [Ti(H2O)6]+3
Ti(III) is a d1 complex and exhibits ONE absorption in its electronic
spectrum due to transition of the electron from the t2g orbitals to the eg orbitals.
The energy of the absorption corresponds to DO.
Br–
Cl–
(H2N)2C=O
NCS–
F–
H2O
CN–
11,400
13,000
17,550
18,400
18,900 20,100 22,300
For complexes having more than one but less than nine electrons, we must
employ an energy level diagram based upon the Russell – Saunders states of the
relevant dn configuration in the free ( uncomplexed ) ion. It can be shown that
just as the set of five d orbitals is split apart by the electrostatic field of
surrounding ligands to give two or more sets of lower degeneracy, so also are the
various Russell – Saunders states of a dn configuration. The number and types of
the components into which an octahedral field will split a state of given L is the
same regardless of the dn configuration from which it arises ( Table 1 ).
Although the states into which a given free ion state is split are the same in
number and type in both octahedral and tetrahedral field, the pattern of energies
is inversed in one case relative to the other, why?.
The electronic spectra of the complex [Cr(NH3)6]+3 in aqueous solution
showed two types of absorption bands, Fig. 1,The bands at the lowest energy
attributed to d – d transition, while the very intense band attribute to charge
transfer transition . A closer analysis shows that there is also a third transition
hidden under the very intense CT band. The ground configuration for this
complex is t32g so the excited state is t22geg1 t32g ; because there are three t2g
orbitals and two eg orbitals, there are in fact six possible transitions because any
of the three t2g electrons can migrate to either of the two eg orbitals. In the
absence of interelectron repulsions, all six transitions occur at the same energy.
However, because there are interelectron repulsions, the transition energies
depend on which orbitals are specially involved in the transition.
Table – 1 Russell – Saunders states in Oct. & Tet. Electronic field
Fig. 2 The spectrum of the d3 complex [Cr(NH3)6]+3
It may be helpful to see qualitatively from the viewpoint of simple MOT why
there are two bands. The dz2 dxy transition promotes an electron from the xy
plane into the already electron rich z direction ( it is electron rich because both dyz
and dxz are occupied). However, the dz2 dxz transition merely relocates an electron
that is already largely concentrated along the z axis , Fig. 3. The repulsion between
electrons in these two cases are not the same, and as a result the two eg t2g
transitions lie at different energies. All the other transitions resemble one or other
of these two cases, and three transitions fall into one group and the other three fall
into the second group.
The spectrum of octahedral and tetrahedral complexes of Mn(II) in Mn(H2O)6]+2
and [MnBr4]-2 showed for the first one, a very weak bands compared to other
octahedral complexes and a large number of bands and last a great variation in the
width of the bands, with one extremely narrow indeed( Fig. 4 ). The correspond to
the 6S ground state of the free ion, which is not split by the ligand field. All the
excited states have different spin multiplicity from the ground state, and transition
to them are spin forbidden. Because of weak spin orbit interaction, such transition
are not totally absent, but they are very weak.
Tetrahedral complexes of Mn(II) are yellow – green, and the color is more
intense than that of the octahedral complexes. As shown in Fig. 4 the molar
absorbance is 100 times than that of the octahedral due probably to the fact that in
tetrahedral environment there is a mixing between p orbitals of the ligands and d
orbital of the metal ion, which is facilitated by overlap of metal d orbitals with
ligand orbitals in the tetrahedral complexes.
Fig. 3 The shift in electron density
a) relocate of
electron density toward the ligands on z axis b) less
relocation
Fig. 4 The d – d absorption spectra of the Mn(H2O)6]+2 ion
(solid) and the [MnBr4]-2 ion ( dashed curve)