Transition Metals and Coordination Compounds:
1- electronic configurations and #of unpaired electrons for transition metals
and their ions
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Most transition metals have unfilled d orbitals, which creates a large number
of other electrons that can be removed.
 Transition metals generally lose the s electron(s) to form +1 and +2 ions, but
they can also lose some (or all) of the d electrons to form other oxidation
states as well.
 When electrons are removed from transition metals, we find the 4s
depopulates first.
 The # of d electrons each metal has depends on the column they r in; for
example: Co- 9th column 9-2 s electrons= 7 d electrons, when it becomes an
ion, it loses the 2 electrons first then the d electrons; for example: Co= [Ar]
4s^2 3d^7, Co 2+ = [Ar] 4s^2 3d^7, Co 3+ = [Ar] 3d^6
 Exceptions: Cr [Ar] 3d54s1 (5 ½ filled d orbital and s orbital is more stable),
Cu- [Ar]3d104s1 (fully filled d orbital is more stable)
2- oxidation states of transition metals in their coordination compounds and
complex ions
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One of the key features of transition metal chemistry is the wide range of
oxidation states (oxidation numbers). Due to the relatively low reactivity of
unpaired d electrons, these metals typically form several oxidation states and
therefore can have several oxidation numbers.
Iron has two common oxidation states (+2 and +3) in, for example, Fe2+ and
Fe3+
Why are there different compounds with different oxidation states: look at
overall energy: amt of energy needed to ionize the metal (take off electrons)
but also the amt of energy released when compound is formed- lattice energy
The more energy released, the more stable the compound.
The more highly charged the ion, the more electrons you have to remove and
the more ionization energy you will have to provide.
In the Fe case- the extra ionization energy is compensated more or less by
the extra lattice energy evolved when the 3+ compound is made.
The net effect of all this is that the overall enthalpy change isn't vastly
different whether you make, say, FeCl2 or FeCl3. That means that it isn't too
difficult to convert between the two compounds.
In complex ions or molecules, the oxidation numbers of these atoms can be
calculated if we assume that the oxidation numbers of the other atoms in the
species are fixed.
3- structure of coordination compounds and complex ions
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A coordination complex is the product of a Lewis acid-base reaction in which
neutral molecules or anions (called ligands) bond to a central metal atom (or
ion) by coordinate covalent bonds.
A complex ion is comprised of two important parts: the central atom and its
surrounding ligands. The central atom can be any metallic ion (usually a
transition metal). The ligands are any combination of anions that can donate
an electron pair, effectively meaning they are all Lewis bases. When
combined they form coordinate covalent bonds.
These shapes are defined by orbital overlap between ligand and metal
orbitals and ligand-ligand repulsions, which tend to lead to certain regular
geometries.
4- coordination number and geometry of complex ions
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1.
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Two common forms are the square planar, in which four ligands are
arranged at the corners of a hypothetical square around the central metal
atom, and the octahedral, in which six ligands are arranged, four in a plane
and one each above and below the plane.
The coordination number of a particular complex is determined by the
relative sizes of the metal atom and the ligands, by spatial (steric) constraints
governing the shapes (conformations) of polydentate ligands, and by
electronic factors, most notably the electronic configuration of the metal
ion. Although coordination numbers from 1 to 16 are known, those below 3
and above 8 are rare.
CN=3 trigonal planar, trigonal pyramidal
CN=4 square planar, tetrahedral
CN=6 octahedral
nomenclature for coordination compounds and complex ions
Ligands are named first in alphabetical order.
The name of the metal comes next.
The oxidation state of the metal follows, noted by a Roman numeral in
parentheses (II, IV)
Ligands that act as anions which end in "-ide" are replaced with an ending "o"
Anions ending with "-ite" and "-ate" are replaced with endings "-ito" and "ato" respectively
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Coordination_Chemistry
/Basics_of_Coordination_Chemistry/Nomenclature_of_Coordination_Complex
es
name of cation first, anion last
prefixes denoting the number of a particular ligand are ignored when
alphabetizing
2. structural isomers and stereoisomer’s and common geometries for transition
metal coordination compounds
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if a ligand can bind to the metal with one or another atom as the donor atom
linkage isomers are formed- ex NO2- linkage isomerism
if isomers differ in what ligands are bonded to the metal and what is outside
the coordination sphere- coordination-sphere isomers
stereoisomers- geometric isomers- cis and trans & optical
isomers=enantiomers- mirror images of each other; diasteromers- cannot be
superimposed on each other
TOPIC AREA:
1. determining the # elec in the d orbitals
see above
2. valence bond theory
mixing of atomic orbitals to form the same number of hybrid orbitals
explains chemical bonding and shapes
hybridization- combining of atomic orbitals- ex sp, sp2, sp3, sp3d- 1 s, 3 p, 1 d
s- 1 orbital
p-3 orbitals
d-5 orbitals
f-7 orbitals
each orbital can have 2 electrons occupied
ex: mixing one s and all three p orbitals produces a set of four equivalent sp3 hybrid
atomic orbitals- tetrahedron ex CH4
d2sp3- two 3d, one 3s, and three 3p atomic orbitals are mixed, central atom bonds
to six other atoms, octahedron shape
3. crystal field theory
strong field ligands- low spin complexes
weak field ligands- high spin complexes
 ability of ligands to cause a large splitting of the energy between the orbitals
is essentially independent of the metal ion and the SPECTROCHEMICAL
SERIES is a list of ligands ranked in order of their ability to cause large
orbital separations.
 When metal ions that have between 4 and 7 electrons in the d orbitals form
octahedral compounds, two possible electron distributions can occur. These
are referred to as either weak field - strong field or high spin - low spin
configurations.
 Tetrahedral- upper layers (t2): dxy, dyz, dxz, lower levels (eg)dx2-y2, dz2
 Splitting energy in tetrahedral: ∆ t = 4/9 ∆ o
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All tetrahedral complexes are high spin
the electrons in the d-orbitals and those in the ligand repel each other due to
repulsion between like charges. Thus the d-electrons closer to the ligands
will have a higher energy than those further away which results in the dorbitals splitting in energy. This splitting is affected by the following factors:
nature of ligands, arrangement of ligands, metal’s ox state, and nature of
metal ion
know octahedral split:
know square planar split:
4. colors absorbed or reflected by a complex ion
interactions between electrons on a ligand and the orbitals on the metal cause
differences in energies between orbitals in the complex
some ligands, such as F, tend to make the gap between the orbitals higher, and some
like cyano CN groups make the gap smaller
when the splitting energy is increased=more energy
the larger the gap, the shorter the wavelength of light absorbed by electrons
jumping from a lower energy orbital to a higher one, wavelength of light observed in
the complex is longer- closer to red
as energy gap gets smaller- light absorbed is of longer wavelength ad the shorter
wavelength light is reflected
5. energy, frequency, and wavelength
see #4