Download d-Block metal chemistry: general considerations

Dr. Said El-Kurdi
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Inorganic Chemistry B
Chapter 20
d-Block metal chemistry:
general considerations
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 The term ‘transition elements (metals)’ is also widely
used. However, the group 12 metals (Zn, Cd and Hg) are
not always classified as transition metals.
 The elements in the f-block are sometimes called inner
transition elements.
 Each group of d-block metals consists of three members
and is called a triad.
 Metals of the second and third rows are sometimes
called the heavier d-block metals.
 Ru, Os, Rh, Ir, Pd and Pt are collectively known as the
platinum-group metals.
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Ground state electronic configurations
the ground state of chromium is
rather than
M2+ and M3+ ions of the first row d-block metals all have
electronic configurations of the general form [Ar]3dn, and so
the comparative chemistry of these metals is largely
concerned with the consequences of the successive filling of
the 3d orbitals.
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Physical properties
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The metallic radii (rmetal) for 12-coordination (Table 6.2 and
Figure 20.1) are much smaller that those of the s-block metals
of comparable atomic number.
Figure 20.1 also illustrates that values of rmetal:
 show little variation across a given row of the d-block;
 are greater for second and third row metals than for first
row metals:
 are similar for the second and third row metals in a given
triad. (lanthanoid contraction)
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Metals of the d-block are (with the exception of the group
12 metals) much harder and less volatile than those of the
s-block.
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Metals in the second and third rows generally possess
higher enthalpies of atomization than the corresponding
elements in the first row.
This is a substantial factor in accounting for the far greater
occurrence of metal–metal bonding in compounds of the
heavier d-block metals compared with their first row
congeners
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 The first ionization energies (IE1) of the d-block metals in a
given period are higher than those of the preceding s-block
metals.
 Across each of the periods K to Kr, Rb to Xe, and Cs to Rn, the
variation in values of IE1 is small across the d-block and far
greater among the s- and p-block elements.
 Within each period, the overall trend for the d-block metals
is for the ionization energies to increase, but many small
variations occur.
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 All 3d metals have values of IE1 and IE2 larger than those
of calcium, and all except zinc have higher values of
aHo these factors make the metals less reactive than
calcium.
 In the formation of species containing M2+ ions, all the
3d metals are thermodynamically less reactive than
calcium, and this is consistent with the standard
reduction potentials listed in Table 20.1
 Formation of a coherent surface film of metal oxide
often renders a metal less reactive than expected
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The reactivity of the metals
In general, the metals are moderately reactive and combine
to give binary compounds when heated with dioxygen,
sulfur or the halogens
product stoichiometry depending, in part, on the available
oxidation states
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Characteristic properties: a general perspective
The colors of d-block metal compounds are a characteristic
feature of species with ground state electronic
configurations other than d0 and d10.
[Cr(OH2)6]2+ is sky blue,
[Mn(OH2)6]2+ very pale pink,
MnO4 intense purple
salts of Sc(III) (d0) or Zn(II) (d10) are colorless.
The fact that many of the observed colors are of low
intensity is consistent with the color originating from
electronic ‘d–d’ transitions.
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The pale colors indicate that the probability of a transition
occurring is low.
The intense colors of species such as MnO4 have a different
origin, namely charge transfer absorptions or emissions.
The latter are not subject to selection rule 20.4 and are
always more intense than electronic transitions between
different d orbitals.
electronic spectra
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Paramagnetism
The occurrence of paramagnetic compounds of d-block
metals is common and arises from the presence of
unpaired electrons. This phenomenon can be investigated
using electron paramagnetic resonance (EPR) spectroscopy
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Complex formation
d-Block metal ions readily form complexes, with complex
formation often being accompanied by a change in color
and sometimes a change in the intensity of color.
 A central metal atom bonded to a group of molecules or
ions is a metal complex.
 If it’s charged, it’s a complex ion.
 Compounds containing complexes are coordination
compounds.
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Lewis acids and bases
A Lewis base is a molecule or ion that donates a lone pair
of electrons to make a bond
Examples:
NH3
OH2
-
-
Cl
F
Electrons in the highest occupied orbital (HOMO) of a
molecule or anion are the best Lewis bases
A Lewis acid is a molecule of ion that accepts a lone
pair of electrons to make a bond
Examples:
+
H
3+
Co
2+
Co
n+
M
Molecules or ions with a low lying unoccupied orbital
(LUMO) of a molecule or cation are the best Lewis acids
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 The molecules or ions coordinating to the metal are the
ligands.
 They are usually anions or polar molecules
 They must have lone pairs to interact with metal
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Alfred Werner: the father of the
structure of coordination
complexes
The Nobel Prize in Chemistry 1913
"in recognition of his work on the
linkage of atoms in molecules by
Alfred Werner
Switzerland
University of Zurich
Zurich, Switzerland
b. 1866
(in Mulhouse, then Germany)
d. 1919
which he has thrown new light on
earlier investigations and opened
up new fields of research especially
in inorganic chemistry"
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Same metal, same ligands, different number of ions
when dissolved
How did Werner deduce the structure of coordination
complexes?
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Werner suggested in 1893 that metal ions have
primary and secondary valences.
 Primary valence equal the metal’s oxidation number
 Secondary valence is the number of atoms directly
bonded to the metal (coordination number)
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Variable oxidation states
The occurrence of variable oxidation states and, often, the
interconversion between them, is a characteristic of most
d-block metals. Exceptions are in groups 3 and 12
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Electroneutrality principle
Pauling’s electroneutrality principle is an approximate method
of estimating the charge distribution in molecules and
complex ions.
The distribution of charge in a molecule or ion is such that
the charge on any single atom is within the range +1 to 1
(ideally close to zero)
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(a) a conventional diagram showing the
donation of lone pairs of electrons from
ligands to metal ion
the charge distribution that results from a
100% covalent model of the bonding
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the charge distribution that results from a
100% ionic model of the bonding
the approximate charge distribution that results
from applying the electroneutrality principle.
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Coordination numbers and geometries
most examples in this section involve mononuclear complexes
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Coordination environments are often described in terms of
regular geometries such as those in Table 20.4, in practice
they are often distorted
Detailed discussion of a particular geometry usually involves
bond lengths and angles determined in the solid state and
these may be affected by crystal packing forces
Small energy difference may also lead to the observation of
different structures in the solid state
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sterically demanding ligands favor low coordination numbers
at metal centers;
 high coordination numbers are most likely to be attained
with small ligands and large metal ions;
 the size of a metal ion decreases as the formal charge
increases, e.g. r(Fe3+) < r(Fe2+);
 low coordination numbers will be favored by metals in
high oxidation states with -bonding ligands.
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The Kepert model
VSEPR model in predicting the shapes of molecular species of
the p-block elements
we might reasonably expect the structures of the complex
ions
to vary as the electronic configuration of the metal ion
changes. However, each of these species has an octahedral
arrangement of ligands
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The VSEPR model is not applicable to d-block metal
complexes.
Kepert model, in which the metal lies at the center of a
sphere and the ligands are free to move over the surface
of the sphere.
Kepert ignores non-bonding electrons
Independent of the ground state electronic configuration of
the metal center
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 One of the most important classes of structure for
which the Kepert model does not predict the correct
answer is that of the square planar complex, and here
electronic effects are usually the controlling factor, as
we will discuss in Section 21.3.
 Another factor that may lead to a breakdown of the
Kepert model is the inherent constraint of a ligand. For
example:
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Chelate Effect
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Chelate Effect
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 The oxidation +2 state is common for almost all the
transition metals. Suggest an explanation.
 No compounds are known in which scandium is in the +2
oxidation state. Suggest an explanation.
 How many electrons are in the valence d orbitals in these
transition-metal ions? (a) Co3+ , (b) Cu+, (c) Cd2+ , (d) Os3+.
 Why can the NH3 molecule serve as a ligand but the BH3
molecule cannot?
 Would you expect ligands that are positively charged to be
common?
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Some, but not all, of these ligand arrangements are in accord
with the Kepert model.
For example, the coordination sphere in [Cu(CN)3]2 is
predicted by the Kepert model to be trigonal planar.
Indeed, this is what is found
experimentally.
The other option in Table 20.4 is trigonal pyramidal, but this
does not minimize interligand repulsions.
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the four nitrogen donor atoms of a porphyrin ligand are
confined to a square planar array
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tripodal ligands such as 20.3 have
limited flexibility which means that
the donor atoms are not
necessarily free to adopt the
positions predicted by Kepert;
macrocyclic ligands are less lexible
than open chain ligands.
polyether 18-crown-6
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Coordination numbers in the solid state
 molecular formula can be misleading in terms of
coordination number
For example in CdI2, each Cd center
is octahedrally sited, and molecular
halides or pseudohalides (e.g. [CN])
may contain MXM bridges and exist
as oligomers, e.g. -PdCl2 is
polymeric
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 when the bonding mode of a ligand can be described in
more than one way. This often happens in organometallic
chemistry, for example with cyclopentadienyl ligands
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Coordination number 2
Examples of coordination number 2 are uncommon, being
generally restricted to Cu(I), Ag(I), Au(I) and Hg(II), all d10
ions.
Examples include [CuCl2], [Ag(NH3)2]2+, [Au(CN)2],
(R3P)AuCl, [Au(PR3)2]+ (R = alkyl or aryl) and Hg(CN)2, in each
of which the metal center is in a linear environment.
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3-coordinate
 Bulky amido ligands, are often associated with low
coordination numbers.
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Coordination number 3
3-Coordinate complexes are not common. Usually,
trigonal planar structures are observed, and examples
involving d10 metal centers include:
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Coordination number 4
4-Coordinate complexes are extremely common, with a
tetrahedral arrangement of donor atoms being the most
frequently observed.
 The tetrahedron is sometimes ‘flattened’, distortions being
attributed to steric or crystal packing effects or, in some
cases, electronic effects.
 Tetrahedral complexes for d3 ions are rare
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Tetrahedral complexes for d4 ions have been stabilized only
with bulky amido ligands
for M = Hf or Zr
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Square planar complexes are rarer than tetrahedral, and are
often associated with d8 configurations where electronic
factors strongly favor a square planar arrangement
the steric demands of
the ligands distort the
structure from the square
planar structure expected
for this d8 metal centre
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Coordination number 5
 The limiting structures for 5-coordination are the trigonal
bipyramid and square-based pyramid.
 The energy differences between trigonal bipyramidal and
square-based pyramidal structures are often small
trigonal bipyramidal
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square-based pyramidal
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Coordination number 6
For many years after Werner’s proof from stereochemical
studies that many 6-coordinate complexes of chromium and
cobalt had octahedral structures
The regular or nearly regular octahedral coordination sphere
is found for all electronic configurations from d0 to d10,
low-spin and high-spin complexes
Jahn–Teller distortion
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there is a small group of d0 or d1 metal complexes in which
the metal center is in a trigonal prismatic or distorted
trigonal prismatic environment.
The octahedron and trigonal prism are closely related, and
can be described in terms of two triangles which are
staggered (20.9) or eclipsed (20.10).
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 contain regular trigonal prismatic (D3h) metal centres
 the coordination
environment is
distorted trigonal
prismatic (C3v)
The common feature of the ligands in these complexes is
that they are -donors, with no -donating or -accepting
properties
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For a regular trigonal prism, angle
 in 20.13 is 0o
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Coordination number 7
High coordination numbers (7) are observed most frequently
for ions of the early second and third row d-block metals and
for the lanthanoids and actinoids, i.e. rcation must be relatively
large
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capped trigonal prismatic [ZrF7]3
capped octahedral [TaCl4(PMe3)3]
pentagonal bipyramidal
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capped octahedral
Coordination number 8
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 Specifying the counter-ion is important since the energy
difference between 8-coordinate structures tends to be
small with the result that the preference between two
structures may be altered by crystal packing forces in two
different salts.
which possess square antiprismatic or dodecahedral structures
depending on the cation.
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Isomerism in d-block metal complexes
Stereoisomers possess the same connectivity of atoms, but
differ in the spatial arrangement of atoms or groups.
If the stereoisomers are not mirror images of one another,
they are called diastereoisomers.
Stereoisomers that are mirror images of one another are
called enantiomers.
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Self-study exercises P627
All the answers can be found by reading Section 2.9.
Structural isomerism: ionization isomers
Ionization isomers result from the interchange of an anionic
ligand within the first coordination sphere with an anion
outside the coordination sphere.
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Structural isomerism: hydration isomers
Hydration isomers result from the interchange of H2O and
another ligand between the first coordination sphere and the
ligands outside it.
When this is dissolved in
water, the chloride ions in the
complex are slowly replaced
by water to give blue-green
blue-green
violet
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Structural isomerism: coordination isomerism
Coordination isomers are possible only for salts in which both
cation and anion are complex ions; the isomers arise from
interchange of ligands between the two metal centers.
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Structural isomerism: linkage isomerism
Linkage isomers may arise when one or more of the ligands
can coordinate to the metal ion in more than one way, e.g. in
[SCN] , both the N and S atoms are potential donor
sites. Such a ligand is ambidentate.
N-bonded ligand, the correspondin vibrational
wavenumbers are 1310 and 1430 cm1.
For the O-bonded ligand, characteristic
absorption bands at 1065 and 1470 cm1 are
observed,
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The DMSO ligand (dimethylsulfoxide) can coordinate to
metal ions through either the S- or O-donor atom.
These modes can be distinguished by using IR spectroscopy:
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An example of the interconversion of linkage isomers
involving the DMSO ligand is
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 Indicate the coordination number of the metal and the
oxidation number of the metal as well as the number
and type of each donor atom of the ligands for each of
the following complexes:
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Square planar species
In a square planar species such as [PtCl4]2  , the four Cl
atoms are equivalent. Similarly, in [PtCl3(PMe3)] , there is only
one possible arrangement of the groups around the square
planar Pt(II) center.
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The introduction of two PMe3 groups to give [PtCl2(PMe3)2]
leads to the possibility of two stereoisomers
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Trigonal bipyramidal species
axial position
equatorial position
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Octahedral species
In EX2Y4 the X groups may be mutually cis or trans
[SnF4Me2]2
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If an octahedral species has the general formula EX3Y3, then
the X groups (and also the Y groups) may be arranged so as
to define one face of the octahedron or may lie in a plane
that also contains the central atom E
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 Indicate the likely coordination number of the metal in
each of the following complexes:
 what would you predict for the magnitude of the
equilibrium constant? Explain your answer.
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 Is the following ligand a chelating one? Explain.
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Distinguishing between cis- and trans-isomers of a square
planar complex or between mer- and fac-isomers of an
octahedral complex is most unambiguously confirmed by
structural determinations using single-crystal X-ray
diffraction.
Vibrational spectroscopy may also be of assistance.
The selection rule for an IR active vibration is that it must
lead to a change in molecular dipole moment
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[Pt(NH)2Cl2]
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Stereoisomerism: enantiomers
A pair of enantiomers consists of two molecular species
which are non-superposable mirror images of each other
The occurrence of enantiomers (optical isomerism) is
concerned with chirality
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[Cr(acac)3], an octahedral tris-chelate complex
A molecule is chiral if it is non-superposable on its mirror image
Enantiomers are distinguished by using the labels  and 
Chiral molecules rotate the plane of polarized light. This
property is known as optical activity. Enantiomers rotate the
light to equal extents, but in opposite directions, the
dextrorotatory (d) enantiomer to the right and the
laevorotatory (l) enantiomer to the left.
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A mixture of equal amounts of two enantiomers is called a
racemate.
The rotation, , may be measured in an instrument called a
polarimeter (Figure 20.14). In practice, the amount of
rotation depends upon the wavelength of the light,
temperature and the concentration of compound present in
solution.
The specific rotation, [],
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(+) and () prefixes: the specific rotation of enantiomers is
equal and opposite, and a useful means of distinguishing
between enantiomers is to denote the sign of []D. Thus, if
two enantiomers of a compound A have []D values of +12o
and 12o, they are labeled (+)-A and ()-A.
d and l prefixes: sometimes (+) and () are denoted by
dextro- and laevo- (derived from the Latin for right and left)
and these refer to right- and left-handed rotation of the
plane of polarized light respectively; dextro and laevo are
generally abbreviated to d and l.
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The +/ or d/l notation is not a direct descriptor of the
absolute configuration of an enantiomer (the arrangement
of the substituents or ligands) for which the following
prefixes are used.
R and S prefixes: the convention for labeling chiral carbon
atoms (tetrahedral with four different groups attached)
uses sequence rules
This notation is used for chiral organic ligands, and also for
tetrahedral complexes.
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 and  prefixes: enantiomers of octahedral complexes
containing three equivalent bidentate ligands (tris-chelate
complexes) are among those that are distinguished using 
(delta) and  (lambda) prefixes.
The octahedron is viewed down a 3-fold axis, and the
chelates then define either a right- or a left-handed helix.
The enantiomer with right-handedness is labeled  , and
that with left-handedness is .
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The complexes [Cr(acac)3]


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[Co(en)2Cl2]+


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Nomenclature
Ligands are frequently named using older trivial names
rather than the International Union of Pure and Applied
Chemistry (IUPAC) names
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Common Name
hydrido
fluoro
IUPAC Name
hydrido
fluoro
Formula
H
F
chloro
bromo
iodo
nitrido
azido
chloro
bromo
iodo
Cl
Br
I
N3
N 3
oxo
cano
thiocyano
isothiocyano
oxido
cano
thiocyanato-S(S-bonded)
isothiocyanato-N(N-bonded)
hydroxo
hydroxo
aqua
carbonyl
aqua
carbonyl
thiocarbonyl
nitrosyl
nitro
thiocarbonyl
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nitrosyl
nitrito -N (N-bonded)
CS
nitrito
nitrito- O (O-bonded)
methylisocyanide
phosphine
pyridine
ammine
methylisocyanide
phosphane
pyridine (abbrev. py)
ammine
methylamine
methylamine
ONO
CH3NC
PR3
C5H5N
NH3
MeNH2
nitrido
azido
O2
CN
SCN
NCS
OH
H2O
CO
NO+
NO2
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nitrido
azido
nitrido
N
azido
N 3
oxo
cano
oxido
cano
thiocyano
isothiocyano
hydroxo
thiocyanato-S(S-bonded)
isothiocyanato-N(N-bonded)
hydroxo
O2
CN
SCN
aqua
carbonyl
aqua
carbonyl
thiocarbonyl
nitrosyl
nitro
thiocarbonyl
nitrosyl
nitrito -N (N-bonded)
CS
nitrito
nitrito- O (O-bonded)
methylisocyanide
methylisocyanide
ONO
CH3NC
phosphine
phosphane
PR3
pyridine
ammine
pyridine (abbrev. py)
ammine
methylamine
amido
methylamine
azanido
C5H5N
NH3
MeNH2
imido
azanediido
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NCS
OH
H2O
CO
NO+
NO2
NH2
NH2
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Nomenclature Rules
1. The cation comes first, followed by the anion.
Examples: diamminesilver(I) chloride, [Ag(NH3)2]Cl
potassium hexacyanoferrate(III), K3[Fe(CN)6]
2. The inner coordination sphere is enclosed in square
brackets. Although the metal is provided first within the
brackets, the ligands within the coordination sphere are
written before the metal in the formula name.
Examples: tetraamminecopper(II) sulfate, [Cu(NH3)4]SO4
hexaamminecobalt(III) chloride, [Co(NH3)6]Cl3
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3. The number of ligands of each kind is indicated by prefixes
(Table 3). In simple cases, the prefixes in the second column
are used. If the ligand name already includes these prefixes
or is complicated, it is set off in parentheses, and prefixes in
the third column (ending in –kis) are used.
2
di
bis
3
4
5
6
7
8
9
10
tri
tetra
penta
hexa
hepta
octa
nona
deca
tris
tetrakis
pentakis
hexakis
heptakis
octakis
nonakis
decakis
dichlorobis(ethylenediamine)cobalt(III),
[Co(NH2CH2CH2NH2)2Cl2]+
tris(2,2-bipyridine)iron(II),
[Fe(C10H8N2)3]2+
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4. Ligands are generally written in alphabetical order according to
the ligand name, not the prefix.
Examples: tetraamminedichlorocobalt(III), [Co(NH3)4Cl2]+
amminebromochloromethylamineplatinum(II), Pt(NH3)BrCl(CH3NH2)
5. Anionic ligands are given an o suffix.
Neutral ligands retain their usual name.
Coordinated water is called aqua and
coordinated ammonia is called ammine .
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6. Two systems exist for designating charge or oxidation number:
a. The Stock system puts the calculated oxidation number of
the metal as a Roman numeral in parentheses after the metal
name.
b. The Ewing-Bassett system puts the charge on the
coordination sphere in parentheses after the name of the
metal.
In either case, if the charge is negative, the suffix -ate is added
to the name.
tetraammineplatinum(II) or tetraammineplatinum(2+), [Pt(NH3)4]2+
tetrachloroplatinate(II) or tetrachloroplatinate(2–), [PtCl4]2
hexachloroplatinate(IV) or hexachloroplatinate(2–), [PtCl6]2
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7. Prefixes designate adjacent (cis -) and opposite (trans -) geometric
locations.
cis - and trans -diamminedichloroplatinum(II), [PtCl2(NH3)2]
8. Bridging ligands between two metal ions (Figures 1) have the prefix m-.
tris(tetrammine-m-dihydroxocobalt)cobalt(6+), [Co(Co(NH3)4(OH)2)3]6+
m-amido-m-hydroxobis(tetramminecobalt)(4+),
[(NH3)4Co(OH)(NH2)Co(NH3)4]4+
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9. When the complex is negatively charged, the names for these metals are
derived from the sources of their symbols:
iron (Fe) ferrate
lead (Pb) plumbate
silver (Ag) argentate
tin(Sn) stannate
gold (Au) aurate
copper (Cu) cuprate
tetrachloroferrate(III) or tetrachloroferrate(1–), [FeCl4]
Examples:
dicyanoaurate(I) or dicyanoaurate(1–), [Au(CN)2]
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Name these coordination complexes:
a. Cr(NH3)3Cl3
b. Pt(en)Cl2
c. [Pt(ox)2]2
d. [Cr(H2O)5Br]2+
e. [Cu(NH2CH2CH2NH2)Cl4]2
f. [Fe(OH)4]
a. Triamminetrichlorochromium(III)
b. Dichloroethylenediamineplatinum(II)
c. Bis(oxalato)platinate(II) or bis(oxalato)platinate(2-)
d. Pentaaquabromochromium(III) or pentaaquabromochromium(2+)
e. Tetrachloroethylenediaminecuprate(II) or
tetrachloroethylenediaminecuprate(2-)
f. Tetrahydroxoferrate(III) or tetrahydroxoferrate(1-)
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Give the structures of these coordination complexes:
a. Tris(acetylacetonato)iron(III)
b. Hexabromoplatinate(2–)
c. Potassium diamminetetrabromocobaltate(III)
d. Tris(ethylenediamine)copper(II) sulfate
e. Hexacarbonylmanganese(I) perchlorate
f. Ammonium tetrachlororuthenate(1–)
a
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b
c
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d
f
e
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