Studies in structural chemistry
Abstract
The molar refractions of the rare earth nitrates,
the ionic refractions and the polarizabilities of the
trivalent rare earth ions decrease from lanthanum to
ytterbium with a break at gadolinium, as a consequence of
lanthanide contraction. The observed polarizabilities
of the rare earth ions show that these are hard Lewis
acids in Pearson's classification. Accordingly, these
ions will form complexes with oxygen- or nitrogen?containing
donors.
The infrared spectra of the rare earth perchlorates
indicate that the perchlorate ion is ionic and non?coordinating.
However, the band of the perchlorate ion is
found to be split. This has been explained to be due to
the asymmetry of the ligand field of the water molecules
coordinated to the metal ion.
Correlation of the donor properties of the ligands studied:
Antipyrine (Ap), dimethylsulphoxide (DMSO) and
pyridine N?oxide (PyO) complexes as well as the hydrated
perchlorates behave as strong electrolytes. The number
of ligand molecules attached to the metal atoms differ
as is evident from Table 9.1.
Table 9.1. Number of monodentate ligands attached to the
metal Complex No. of ligandsH?OApLa³?5.5Ce³?7.0Pr³?7.0Nd³?6.5Sm³?4.5Eu³?8.0Y³?7.0
The infrared spectra of hydrated perchlorates and
antipyrine, pyridine N?oxide and dimethylsulphoxide complexes
show that the N–O, C=O, N?O and S=O stretching
vibrations respectively shift to lower frequencies compared
with the ligands (Table 9.2). From this it is
inferred that the ligands coordinate to the metal through
oxygen in all the cases. The results in Table 9.2 show
that the shift for all the rare earth ions for a particular
ligand is constant, indicating the close similarity among
rare earths.
Table 9.2 Lowering of the stretching frequencies on
bonding of N–O, C=O, N?O and S=O in the
complexes Frequency in cm?¹Ligand aloneH–O (H?O): 3652, C=O (Ap): 1658, N?O (PyO): 1265, S=O (DMSO): 1045La³?3425 (227)Ce³?3425 (227)Pr³?3380 (272)Nd³?3420 (232)Sm³?3410 (242)Eu³?3405 (247)Y³?3400 (252)
(Figures in brackets show the total shift in cm?¹.)
The shift in the stretching frequency can be taken
as a rough measure of the metal–oxygen bond strength,
which in turn is a measure of the stability of the complex.
The shifts of 240 cm?¹ for the hydrated perchlorate,
58 cm?¹ for antipyrine, 41 cm?¹ for pyridine N?oxide and
49 cm?¹ for dimethylsulphoxide complexes show that the
bond strength between the metal ion and the ligand or the
donating capacity of the ligands decreases in the order:
H?O > Ap > DMSO > PyO.
From a consideration of an electrostatic model
(S.E. Livingston, Quart. Rev., 19, 386 (1965)), the coordinating
ability of a unidentate ligand will depend not only
on the electronegativity but also on the total dipole
moment (?) of the ligand. In the ligands studied, the
electronegativities of the elements bonded to oxygen are in
the order: H < S < C < N. The order of donor strength
should be the reverse of this, i.e., H?O > S > C > N. However,
infrared studies show that the order of donor strengths
is H?O > Ap > DMSO > PyO. The deviation of dimethyl?
sulphoxide from the simple electronegativity order is
probably due to the fact that the accepted electronegativity
scale can be applied only to oxo?bonds with first?row
elements, and does not relate to compounds containing
other group elements (like S, P, As), since empty d?orbitals
are available on these elements for ??bond formation.
It can be expected that the donor strengths of the
ligands studied will be determined by the electron density
on the oxygen atom. The charge on the oxygen atom can be
calculated by Sanderson's method. Such a calculation
(Table 9.3) shows that the donor strength of the ligands should be in the same order as derived from infrared data,
i.e., Ap > DMSO > PyO.
Table 9.3 Ligands: Dipole moment and charge on oxygen
Ligand Dipole moment (Debye)Charge on oxygenWater1.87–0.2479Antipyrine5.50–0.3035Dimethylsulphoxide3.96–0.2975Pyridine N?oxide4.24–0.2887
Rare earth trichloroacetates are non?electrolytes
in dimethylformamide. Infrared spectra of these compounds
suggest that the trichloroacetate ion coordinates to the
metal ion as a bidentate ligand. The water molecules are
weakly coordinated to the metal. The separation of the
two COO stretching frequencies shows that the metal–oxygen
bond strength increases in the order Y > Gd > Sm > Pr.
Calculation of partial charges by Sanderson’s method:
According to Sanderson
(E.T. Sanderson, "Chemical Periodicity", Reinhold, 1960),
the partial charge on a combined atom of a compound is
the ratio of the electronegativity change undergone by the
atom during bond formation to the electronegativity change
that it would undergo in acquiring a unit charge.
The partial charge on an atom of element E is given by:
q=2.08(S?Sm)q = 2.08 \left( S - S_m \right)q=2.08(S?Sm?)
where S is the electronegativity (stability ratio) of the atom
and S? is the electronegativity of the molecule (or complex ion).
The stability ratio represents the relative
electronic density of an atom compared to that of an
isoelectronic (hypothetical) inert atom:
S=DiDwhereD=Z4?r3S = \frac{D_i}{D}
\quad\text{where}\quad
D = \frac{Z}{4\pi r^3}S=DDi??whereD=4?r3Z?
where Z is atomic number, V is volume, and r is the
non?polar covalent radius. The molecular stability ratio
is the geometric mean of the stability ratios of all atoms
before combination.
The stability ratios for the atoms of
dimethylsulphoxide are:
C = 3.79
S = 4.11
O = 5.21
H = 3.55
Sm(for DMSO)=(3.792×3.552×4.11×5.21)1/6=3.797S_m (\text{for DMSO}) = (3.79^2 \times 3.55^2 \times 4.11 \times 5.21)^{1/6} = 3.797Sm?(for DMSO)=(3.792×3.552×4.11×5.21)1/6=3.797
qO=2.08(3.797?5.210)=?0.2975q_O = 2.08 (3.797 - 5.210) = -0.2975qO?=2.08(3.797?5.210)=?0.2975