CHEM 340
TRANSITION METAL CHEMISTRY
Lect-12
Preliminaries/Background
Just as element names/symbols were memorized for FRO8 atoms
(Li, Be, B, C, N, O, F, Ne), now you are asked to memorize the FROTM
atoms (Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn). These elements provide most of the examples
for (1) main group chemistry and (2) now TM chemistry, so we need to be very familiar with
their names, symbols, and ground state and ion electron configurations, and term states.
2
n
2
FROTM have vse configurations of 4 s 3d . The 4 s electrons are first lost, so all
TM elements have a common + 2 ox.st. (In the FROTM only copper has a + 1ox.st., and in
later rows silver and gold.) Many TM atoms also display + 3 and higher ox.sts., when 3d
electrons are removed.
Angular Distribution Diagrams of d atomic orbitals Reacquaint yourselves with names,
shapes, locations and trigonometric signs of d atomic orbitals. (d orbitals are gerade)
d xy four lobes in the XY plane and between X and Y axes
d xz
d yz
dx − y
dz
2
2
four lobes in the XZ plane and between X and Z axes
four lobes in the YZ plane and between Y and Z axes
2
four lobes in the XY plane and are centered on X and Y axes
seems to be different, but is really a combination of x 2 - z 2 plus y2 - z 2
2
2
=
2
(2 z - x - y ). Main ( + ) lobes centered on Z axis, ( - ) ring in XY plane.
TERM STATE SYMBOLS
Note the pattern obtained for term state symbol
values upon incrementing the number of d electrons:
dn
term symbol
d1
d2
d3
d4
d5
d6
d7
d8
d9
d 10
2
3
4
5
6
5
4
3
2
1
D
F
F
D
S
D
F
F
D
S
S , D, F terms are met. S terms are always associated with
5
10
filled and half-filled subshells, i.e., d and d cases. Each of the remaining eight
cases has one of two terms, D or F . Note that several FROTM ions can have
n
4
identical term states and d configurations, i.e., Cr(II) and Mn(III) for d , Fe(II) and
6
Co(III) for d . Work in TM chemistry is greatly facilitated by seamless transitions from
d n configuration notation to actual TM atom/ion symbol and charge.
Observe that only
Magnetic Susceptibilities (Moments)
Paramagnetic substances (with unpaired electrons)
are attracted to a magnetic field, whereas diamagnetic substances (all electrons paired)
are repelled. Presence of unpaired electrons is common for many TM atoms/ions so
they and their compounds are often paramagnetic because of this feature.
The magnetic character of a substance is determined from two mass measurements,
first in the absence of a magnetic field (i.e., a "normal" weighing procedure), and
secondly in the presence of a magnetic field. (Refer to text, Figure 10-1, page 314.)
Masses of paramagnetic substances are greater in a magnetic field.
Masses of diamagnetic substa nces are less in a magnetic field.
An expression developed to calculate the magnetic character
( µ) involves the
( S ) and total orbital angular momentum
= g S ( S + 1) + ( L( L + 1) / 4 (Bohr Magnetons)
properties of total spin angular momentum
( L) as shown:
µ( S + L )
where g is the gyromagnetic ratio (which has a value very close to 2), and
S = ∑ ms and L = ∑ ml
8
Apply this expression to the
Ni(II) = d
case of Ni(II) ion, and compare
to the experimental magnetic
moment of 2.83 Bohr Magnetons…
S =2/2
ml =
2
1
0
-1
-2
µ( S + L) = 4.5 B.M.
L =3
This does not compare well with the experimental value. A modification of this
expression suggests that for lighter metals (like FROTM) the contribution of the spin
term is more important than the orbital term. This leads to a spin-only formula:
µ( spin only ) = 2 S ( S + 1)
or equivalently
n = number of unpaired electrons. For Ni(II) µ( so )
µ( so ) = n( n + 2)
where
= 2.83 B.M. which agrees well with
the experimental value. Accordingly, we will use the
µ( so )
formula for FROTM.
Magnetic moment determinations are important because they afford a means of (1)
COUNTING unpaired electrons by using experimental µ( so ) value and calculating n
from the above expression, and (2) determining the oxidation of the TM.
Some Examples of Science "in Action"
First, accumulate sets of information, and
then analyze for trends and/or anomalies.
Consider several examples:
a.
Ionization Energy IE vs. at.no.
yields a graph with a maxima at
each noble gas. Note that
mini reversals are met after
2
3
configurations of ns and np ,
This information is analyzed
and intepreted to show very stable electron configurations at (2), (2)+(2+6),
(2)+(2+6)+(2+6), and (2)+(2+6)+(2+6+10)+(2+6) electrons.
b.
Radii for TM free atoms (decreasing size is almost linear), but
Radii for TM IONS (shows double peaked curve) minimum at Fe(II) and Co(III). Why?
c.
Lattice energies (double peaked curve) minimum at Mn(II). Why?
d.
Enthalpy of hydration(double peaked curve) with minimum at Mn(II) and Fe(III).
maximum at Cr(III) & Co(III). Why?
The questions are posed. Let's begin looking for answers (and finding understanding).
Descriptive Ligand Field Theory
CHEMICAL BONDS
(1) The "chemical glue" binding two atoms together has its
origin in vs electrons shared by the atoms, and (2) remember that electrons reside in orbitals.
Accordingly, knowledge of vs orbitals informs where electrons are to be found, and that's
where bonds are expected to form. Accordingly, knowledge of TM vs orbital
shapes, locations, and names/labels are critical to developments that follow.
Six-coordinate Complexes in an Oh Field
C.N. = 6 and pt. gp. symmetry = Oh
d orbitals in Oh pt. gp. symmetry (a char.table look-up)
d z and d x − y transform as the eg two-dimensional irreducible representation.
d xy , d xz , d yz transform as the t2 g three-dimensional irreducible representation
transformation properties of
2
2
2
Consider six ligands positioned at Oh sites as
they approach a TM atom/ion to form an
Oh coordination complex, as shown to the right:
Now superimpose on this diagram the five
d orbitals. Note that orientation of the
dz
2
and
dx − y
2
2
pair allows for greater
L1
+z
L2
TM
L6
+x
L3
+y
L5
L4
interaction with ligand orbitals , then does
the d xy , d xz , d yz set, which has lobes oriented between incoming ligands. Based on
TM-ligand e-e repulsions, TM electrons in the former set will be perturbed more by ligand
electrons and this leads to a differentiation (splitting) of energies for d orbitals in the complex.
(1)
Sigma-bonded Oh TM Complexes
Transformation properties of sigma-bonding p-type atomic orbitals for six ligands…
Using methods met and used in Chapter 5, ligand σ -bonding orbitals transform as
Γ( σ bonds ) = a1 g + eg + t1u
the eg symmetry orbitals, i.e.,
d z and d x − y of the TM.
The TM d xy , d xz , d yz are
2
2
The only TM - ligand bonding interaction involves
eg
2
t2g
non-bonding in this case. The MO
diagram for a sigma bonded Oh TM
complex is shown as Figure 10-5 on
page 321. Isolation of TM d orbitals from this diagram shows the above pattern. The
energy difference between
eg and t2 g
TM orbitals in Oh complexes is labeled
∆o
and is represented by the arrow in the above diagram. The MAGNITUDE of the
∆o
energy difference
varies, depending on (a) inherent bonding strength of the
ligand, and (b) charge of the TM ion.
Energy is conserved in the splitting, so the net destabilization of the
orbitals) is offset by the stabilization of the
t2 g
set (three orbitals).
eg
set (two
Or,
3

 −2

2 x (destabiliz ) + 3 x ( stabiliz) = 0
5

 5

Applying factors of (3 / 5) and ( -2 / 5) to electrons in eg and t2 g sets respectively,
gives the Ligand Field Stabilization Energy.
LFSE can be used to show which
n
particular d configurations will be more (or less) favored for Oh complex formation.
Refer to text Table 10-5 on page 325, and the following development…
d 1 [as in Ti(III)]
LFSE = (3/5)[ 0 ] + (-2/5)[ 1 ] = -2/5
d 2 [as in Ti(II),
LFSE = (3/5)[ 0 ] + (-2/5)[ 2 ] = -4/5
V(III) ]
d 3 [as in V(II),
LFSE = (3/5)[ 0 ] + (-2/5)[ 3 ] = -6/5
Cr(III) ]
Two situations arise for d 4 through
∆o
d 7 depending on the MAGNITUDE of ∆o .
When energy of
is small, relative to energy required for electron pairing, then a
WEAK FIELD (with respect to ligand bonding strength) or HIGH SPIN (with respect to
number of unpaired electrons) condition prevails and so the maximum number of
unpaired electrons will be present.
When energy of
is large, relative to energy required for electron pairing, then a
∆o
STRONG FIELD or LOW SPIN condition prevails and the
t2 g set is filled before the eg .
Consider the following diagrams and LFSE's…
WEAK FIELD
STRONG FIELD
d 4 [as in Cr(II),
Mn(III) ]
LFSE = (3/5)[ 1 ] + (-2/5)[ 4 ] = -3/5
LFSE = (3/5)[ 0 ] + (-2/5)[ 4 ] = -8/5
∆o (weak field) < ∆o (strong field)
4
Notice that the d weak field case results in FOUR unpaired electrons, but the strong field
case has only TWO unpaired electrons. So measurements of magnetic moments can be
used to identify and distinguish between strong field and weak field ligands and complexes.
5
6
7
Convince yourself that the following results would be obtained for d , d , d cases…
[ WEAK FIELD ] LFSE
n
[ STRONG FIELD ]LFSE
5
d
d6
d7
n
0
5
-10/5
1
-2/5
4
-12/5
0
-4/5
3
-9/5
1
d vse, only weak field cases are possible so they would be as indicated below.
d8
-6/5 2
and
d9
-3/5
1
6
5
LFSE are largest for d in a strong field, and zero for d in a weak field. This information can
For 8 and 9
be used to account for some observations in the accumulated data presented at the beginning
d 6 strong
5
Mn(II) and Fe(III) have lowest hydration energies, and they both are d weak field
of the hour: Co(III) has the largest hydration energy , a nd its aqua complex is a
field case.
cases having LFSE = 0.
Magnetic moments are very informative in identifying ox.sts. and field strengths of complexes.
Consider some information about iron complexes. Octahedral complexes of Fe(II), a
can be strong field/low spin with n =0, or weak field/high spin with n =4. Octahedral
5
complexes of Fe(III), a d case, can be strong field/low spin with
with n =5. Magnetic moments for the four situations would be
diamagnetic,
µ( so ) = 4.9,
µ( so ) = 1.8,
and
(2)
d 6 case,
n =1, or weak field/high spin
µ( so ) = 5.9 B.M. respectively.
Pi-bonded Oh TM Complexes
Carbon Monoxide, a pi-acceptor ligand.
Consider the mo diagram for CO, a ten
electron system.
2p
2p
The HOMO is associated more with
carbon, as is the empty pi* LUMO .
The sigma bond from CO to the TM
involves the HOMO, and is why the
bond is TM - C instead of TM - O.
2s
C
2s
O
Energy of the empty pi* LUMO is
similar to that of the TM d vse, and more
importantly, they are symmetry matched, so
bonding between them is possible. In this instance however, the TM furnishes
electrons to the empty pi* LUMO, in a manner known as "back-bonding".
The combination of a sigma bond (donation from C in CO to the TM) and pi-bond (backbonding from TM to empty pi* LUMO in CO) makes for exceptionally strong net
bonding. This is the reason for the toxicity of CO; it binds so strongly to heme iron in
blood that oxygen can no longer be processed. Cyanide anion is isoelectronic with CO
and is also an excellent pi-acceptor ligand. Pi-acceptor ligands are strong field ligands.
Compare CO and CN 1 - to chloride anion, Cl 1 - . Chloride anion has a complete octet
of vse; all of its vs are filled so it cannot accept (pi) electrons from TM. However, it can
donate pi electrons to the TM. But the TM in a complex is already in enough trouble
w/r to electrons and doesn't need any more. Consequently, pi-donor ligands like
chloride anion are weak field ligands; they present too much electron density to the TM.
(3)
Relative Bonding Strengths of Ligands - The Spectrochemical Series
When the bonding effects ( i.e.,
∆o values)
of a series of ligands are compared for
∆o
similar TM's, and then arranged in order of magnitudes of
, a qualitative sequence
of ligands results that is called the Spectrochemical Series. A simple example follows:
STRONG FIELD
CO CN1 - o-phen
pi-acceptor ligands
MODERATE FIELD
NO 2
1-
en
NH 3
H2 O
F
1-
WEAK FIELD
RCO 2
sigma donor ligands
1-
OH
1-
Cl
1-
Br
1-
I
1-
pi-donor ligands
∆o
The effect of larger
values can be demonstrated by the color of coordination
complexes. Consider two complexes of Ni(II), the hexaaqua and hexaammine
8
complexes. As will become apparent soon, Ni(II), a d system, displays a three-band
absorption spectrum. The hexaaqua complex is green and the hexaammine blue.
These colors are observed b/c the energies they represent are not absorbed, as
shown in the following diagram of the visible range of energies…
Visible region of the
electromagnetic
spectrum
[ Ni ( H 2 O ) 6 ] 2 +
complex, absorbs
energies (in black)
and appears green
in color.
[ Ni ( NH 3 ) 6 ] 2 +
complex, absorbs
energies (in black)
and appears blue
in color.
In both spectra the color is due to the three-band absorption system of the nickel(II) ion
However, ammonia is a stronger field ligand than water so the 3 bands absorb at higher
energies. This shifts the transparent "window" between the center and right absorption
bands to higher energies and imparts a blue color to the ammine complex.
Water is classified as a ligand of moderate/weak donating ability, so it would be
anticipated that aqua complexes are WEAK FIELD / HIGH SPIN cases. Indeed that is
found to be the case , except for one FROTM ion. The one TM cation forming a strong
d 6 system. Recall that d 6 strong field
6
Note that complexes of d strong field are diamagnetic.
field complex with water as ligand is Co(III), a
case has the highest LFSE .
Four-coordinate Complexes in a Square-Planar Field
C.N. = 4 and pt. gp. symmetry =
D4 h
d orbitals in D4 h pt. gp. symmetry (a char.table look-up)
d z transforms as the a1 g irreducible representation
d x − y transforms as the b1 g irreducible representation.
d xy transforms as the b2 g irreducible representation, and
d xz , d yz transform as the eg two-dimensional irreducible representation
transformation properties of
2
2
2
+z
Consider four ligands in a square-planar
arrangement approaching on the ± x and
± y axes.
Clearly the
dx − y
2
2
L4
would be
most directly in-line with the ligands and
has the highest energy. Remaining d
orbitals would be affected in the follo wing
order d xy then d z2 and lastly the d xz ,
+x
d yz
L3
TM
+y
L1
pair. This leads to the splitting pattern
as shown to the right. The numbers indicate the
ordering of entering electrons into this pattern.
Ni(II), a
L2
b1g
d 8 system, forms square-planar complexes.
Introducing 8 electrons into the sq-planar pattern
results in a diamagnetic condition. This feature
allows identification and differentiation of sq-planar
Ni(II) complexes from Oh Ni(II) complexes because
the Oh compounds will be paramagnetic with
µ( so ) =2.83 B.M.
1,3
7,8
b2g
5,6
a1g
2,4
eg
All available information is used when characterizing coordination compounds, including
chemical analysis, magnetic moments, conductances, and (from the next chapter)
absorption spectra. So while it may be possible to differentiate sq-planar Ni(II) with four
donor atoms, from Oh Ni(II) with six donor atoms by chemical analysis alone, all
information available is collected and applied to the problem. Differences in magnetic
moments for sq-planar and Oh Ni(II) offers additional reliable information for making
the determination.
Four-coordinate Complexes in a Tetrahedral Field
C.N. = 4 and pt. gp. symmetry =
Td
d orbitals in Td pt. gp. symmetry (a char.table look-up)
d z and d x − y transform as the e two-dimensional irreducible representation.
d xy , d xz , d yz transform as the t2 three-dimensional irreducible representation
transformation properties of
2
2
2
Note in Td symmetry there are no
inversion center.
g
or
u subscripts.
In a Td symmetry field, the splitting pattern
for the five d orbitals is as shown to the right.
Note that it is the inverse of the Oh splitting
pattern.
This is b/c in Td symmetry there is no
t2
e
The energy difference between the two levels is labeled ∆t and is represented by the arrow
in the diagram. A Td field is less than an Oh field because there are only four donor atoms
instead of six. It can be shown that
4
∆t ≈ ∆o .
9
Consequently, Td complexes are
always WEAK FIELD / HIGH SPIN cases.
Ni(II) forms Td as well as Oh and sq-planar complexes, and even five -coordinate complexes.
Its coordination number and geometry depend very much on the ligand itself. One of the many
fun parts of coordination chemistry is met in the design and subsequent preparation of specific
ligand types to cause formation of a complex with a particular geometrical arrangement.
Chemical analysis and magnetic moment determinations of nickel(II) complexes often allows
differentiation between coordination numbers of six, five and four, and in the latter case,
between tetrahedral and square-planar geometries.