Download Ch 12- Transition Metales and coordination compound

Transition Metales and
Coordination Compound
coordination
bond
[ Cu
ionic bond
( NH3 )4 ]2+
Complex ion
（inner sphere）
SO42Counter ion
（outer sphere）
coordination compounds
★ coordination molecules have no outer sphere
but inner sphere
The molecules or ions that surround a
metal ion in a complex are called
ligands.
Ligands may be anions or polar
molecules. Some common ligands are
water, ammonia, hydroxide, cyanide,
and chloride.
Each of these has at least one unshared pair of
electrons that can be used to coordinate (form a
coordinate covalent bond to the metal ion). Metal
ions can act as Lewis acids, accepting electrons
from ligands, which act as Lewis bases. The atom
that is bound directly to the metal ion is called
the donor atom. It's the atom that donates a lone
pair of electrons to the metal ion.
[ Cu（NH3）4 ] 2+ [ Fe(H2O) 6]3+ [ SiF6 ] 2Cu(Ⅱ)
Fe(Ⅲ)
NH3
H2O
F
O
F
N
Si(Ⅳ)
-
：
cations or atoms offering
（central atom） empty orbitals
：
（Ligand）
the molecules or ions that
surround the metal ion in a
complex
： the atoms of ligand bound
（donor atom） directly to the metal
★ Ligand is from the Latin
word ligare, meaning “to bind”
 Coordination sphere
– Metal and ligands bound to it
 Coordination number
– number of donor atoms bonded to the central
metal atom or ion in the complex
• Most common = 4, 6
• Determined by ligands
– Larger ligands and those that transfer
substantial negative charge to metal
favor lower coordination numbers
 Ligands
– classified according to the number of donor
atoms
– Examples
• monodentate = 1
• bidentate = 2
• tetradentate = 4
• hexadentate = 6
• polydentate = 2 or more donor atoms
Coordination Compounds
 Common ligands:
– water
– ammonia
H-O-H
NH3
– chloride ion
Cl
-
– cyanide ion
C
N
-
（monodentate ligand )：
— the ligands possess a single donor
atom and are able to occupy only one
site in a coordination sphere
• Monodentate
– Examples:
• H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides),
CO, O2-
– Example Complexes
• [Co(NH3)6]3+
• [Fe(SCN)6]3-
（bidentate ligand )：
— the ligands possess two donor atoms
– Examples
• oxalate ion = C2O42• ethylenediamine (en) = NH2CH2CH2NH2
• ortho-phenanthroline (o-phen)
– Example Complexes
• [Co(en)3]3+
• [Cr(C2O4)3]3• [Fe(NH3)4(o-phen)]3+
H
H
H
N - CH2- CH2 - N
H
Ethylenediamine (en)
oxalate ion
Orthophenanthroline
carbonate ion
bipyridine
Structures of other
bidentate ligands.
The coordinating
atoms are shown
in blue.
H
H
Ethylenediamine (en)
H
N - CH2- CH2 - N
H
（hexadentate ligand )：
— the ligands possess six donor atoms
ethylenediaminetetraacetate (EDTA) =
(O2CCH2)2N(CH2)2N(CH2CO2)24– Example Complexes
• [Fe(EDTA)]-1
• [Co(EDTA)]-1
Coordination Compounds
 Uses for EDTA
– Used in food products to complex metals that
promote decomposition reactions
– Used to remove toxic heavy metals from the
body
– Used to complex metal ions that inhibit or
promote polymerization reactions
EDTA
Porphine, an important
chelating agent found in
nature
Metalloporphyrin
N
N
2+
NH
NH
N
N
Fe
N
N
Coordination Compounds
• Porphyrins are naturally occurring
complexes derived from porphin.
– Heme (Fe2+)
– Chlorophyll (Mg2+)
Hymoglobin, a protein that stores O2 in cells
Coordination Environment of Fe2+ in Oxymyoglobin,
and Oxyhemoglobin
Ferrichrome
(Involved in Fe transport in bacteria)
Complex charge
☛ The charge on a central metal ion can be
determined in much the same way that
oxidation numbers are determined in
molecular compounds. It is important to
recognize common molecules, to remember
that they are neutral, and to know the charges
on both monoatomic and polyatomic ions. The
sum of charges on the metal ion and the
ligands must equal the net charge on the
complex.
Complex charge = sum of charges on the metal
and the ligands
[Fe(CN)6]3+3
6(-1)
Neutral charge of coordination compound = sum of
charges on metal, ligands, and counterbalancing ions
neutral compound
[Co(NH3)6]Cl2
+2
6(0)
2(-1)
☛ Coordination numbers of 2, 4, and 6 are
common. The geometry of a complex is
determined in part by its coordination number.
(Size of the ligands can also impact the
geometry.)
Coordination
Geometry
Number
2
Linear
4
Tetrahedral(more common)
or square planar
6
Octahedral
Structures of geometries,
(a) [Zn(NH3)4]2+
(b) [Pt(NH3)4]2+,
tetrahedral
square-planar
These are the two common geometries for complexes in
which the metal ion has a coordination number of 4.
Attentions:
 In mono-core coordination compounds, each
donor atom can only bind directly to one central
atom to form coordination bond.
e.g：H2O、 F Though some ligands have two or more donor
atoms, they are still monodentate ligand because
only one donor atom can form bond with central
atom.
e.g：SCN-、 NCS-、 NO2- 、ONO-
1. Simple complex
— The one formed by central metal atom
and monodentate ligand.
e.g:
[Cu(NH3)4]SO4、K[Ag(CN)2]、 K2[PtCl4]、etc.
2. (Chelate compound)
— the ringed coordination compounds
formed by metal atom and polydentate
Polydentate ligands tend to coordinate
more readily and form more stable
complexes than monodentate ligands.
This phenomenon is known as the
Chelate effect .
(chelating agents)
— polydentate ligands are
also known as chelating
agents which are
usually organic ligands
possessing donor atoms
like N、P、O、S.
Cd2+
The [CoEDTA]– ion, showing how
the Ethylene diamine tetraacetate
ion is able to wrap around a metal
ion, occupying six positions in the
coordination sphere.
Nomenclature of CC (IUPAC Rules)
 The cation is named before the anion
 When naming a complex:
– Ligands are named first
• alphabetical order
– Metal atom/ion is named last
• oxidation state given in Roman numerals follows
in parentheses
– Use no spaces in complex name
[Co(NH3)5Cl]2+
Pentaamminechlorocobalt (III)
Nomenclature of CC (IUPAC Rules)
 The names of anionic
ligands end with the
suffix -o
– -ide suffix changed to -o
– -ite suffix changed to -ito
– -ate suffix changed to -ato
Ligand
Name
bromide, Br-
bromo
chloride, Cl-
chloro
cyanide, CN-
cyano
hydroxide,
OH-
hydroxo
oxide, O2-
oxo
fluoride, F-
fluoro
Nomenclature of CC (IUPAC Rules)
Ligand
Name
carbonate, CO32-
carbonato
oxalate, C2O42-
oxalato
sulfate, SO42-
sulfato
thiocyanate, SCN-
thiocyanato
thiosulfate, S2O32-
thiosulfato
Sulfite, SO32-
sulfito
Nomenclature of CC (IUPAC Rules)
 Neutral ligands are referred to by the usual
name for the molecule
– Example
• ethylenediamine
– Exceptions
• water, H2O = aqua
• ammonia, NH3 = ammine
• carbon monoxide, CO = carbonyl
Nomenclature of CC (IUPAC Rules)
 Greek prefixes are used to indicate the number of each
type of ligand when more than one is present in the
complex
– di-, 2; tri-, 3; tetra-, 4; penta-, 5; hexa-, 6
 If the ligand name already contains a Greek prefix, use
alternate prefixes:
– bis-, 2; tris-, 3; tetrakis-,4; pentakis-, 5; hexakis-, 6
– The name of the ligand is placed in parentheses
e.g:
[Co(en)3]Cl3－tris(ethylenediamine)cobalt(III) chloride
Nomenclature of CC (IUPAC Rules)
 If a complex is an anion, its name ends
with the -ate
– appended to name of the metal
e.g:
K4[Fe(CN)6]
potassium hexacyanoferrate(II)
[CoCl4]2-
tetrachlorocobaltate(II) ion
Nomenclature of CC (IUPAC Rules)
Transition
Metal
Name if in Cationic
Complex
Name if in Anionic Complex
Sc
Scandium
Scandate
Ti
titanium
titanate
V
vanadium
vanadate
Cr
chromium
chromate
Mn
manganese
manganate
Fe
iron
ferrate
Co
cobalt
cobaltate
Ni
nickel
nickelate
Cu
Copper
cuprate
Zn
Zinc
zincate
★ Some names of coordination compounds
[Co(NH3)4Cl2]Cl:
tetraamminedichlorocobalt(Ⅲ) chloride
[Pt(NH3)2Cl2]:
diamminedichloroplatinum(Ⅱ)
K[Ni(C2O4)2]:
potassium bisoxalatonicketate(Ⅱ)
oxalic acid
[Cr (NH3)5H2O](NO3)3:
pentaammineaquachromium (Ⅲ) nitrate
Questions
1. Which name is correct for the coordination
compound [Fe(en)2Cl2]Cl?
A dichlorobisethylenediamineiron(III) chloride
B diethylenediaminechlorineiron(I) chloride
C iron(III) bisethylenediaminedichloro chloride
D iron(III) trichloridebisethylenediamine
Classes of isomers
Dr.Monther F.Salem
42
✽ Isomerism of CC
 isomer ：Two or more compounds with the same
elemental composition but different
arrangements of atoms.
— Major Types
 structural isomers
 stereoisomers
Structural Isomers
 structural isomers ：
different bonds.
— isomers that have
1. Coordination-sphere isomers
– differ in a ligand bonded to the metal in the
complex, as opposed to being outside the
coordination-sphere
• Example
[Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl
[Co(NH3)5Cl]Br
vs.
[Co(NH3)5Br]Cl
— Consider ionization in water
[Co(NH3)5Cl]Br  [Co(NH3)5Cl]+ + Br[Co(NH3)5Br]Cl  [Co(NH3)5Br]+ + Cl-
— Consider precipitation
[Co(NH3)5Cl]Br(aq) + AgNO3(aq)  [Co(NH3)5Cl]NO3(aq) + AgBr(s)
[Co(NH3)5Br]Cl(aq) + AgNO3(aq)  [Co(NH3)5Br]NO3(aq) + AgCl(aq)
Linkage Isomers
2. Linkage isomers
— differ in the atom of a ligand bonded
to the metal in the complex
Example:
– [Co(NH3)5(ONO)]2+
vs.
[Co(NH3)5(NO2)]2+
Example:
– [Co(NH3)5(SCN)]2+
vs.
[Co(NH3)5(NCS)]
Co-SCN vs. Co-NCS
Stereoisomers
 stereoisomers
— isomers with the same bonds but different
spatial arrangements of those bonds.
1. Geometric isomers
－Differ in the spatial arrangements of the
ligands
－Have different chemical/physical
properties different colors, melting
points, polarities, solubilities, reactivities,
etc.
Geometric Isomers
cis isomer,
trans isomer，
Pt(NH3)2Cl2
[PtCl2(NH3)2] (cis-platin),
[PtCl2(NH3)2] (trans-platin),
Geometric Isomers
cis isomer
trans isomer
[Co(H2O)4Cl2]+
2. Optical isomers
– isomers that are nonsuperimposable
mirror images
• said to be “chiral” (handed)
• referred to as enantiomers
– A substance is “chiral” if it does not
have a “plane of symmetry”
Stereoisomers
• Other stereoisomers, called optical isomers or
enantiomers, are mirror images of each other.
• Just as a right hand will not fit into a left glove, two
enantiomers cannot be superimposed on each other.
Enantiomers
A molecule or ion that exists as a pair of
enantiomers is said to be chiral.
Example 1
mirror plane
cis-[Co(en)2Cl2]+
Example 1
rotate mirror image 180°
180 °
Example 1
Nonsuperimposable, enantiomers
cis-[Co(en)2Cl2]+
Example 2
mirror plane
trans-[Co(en)2Cl2]+
Example 2
rotate mirror image 180°
180 °
trans-[Co(en)2Cl2]+
Example 2
Superimposable-not enantiomers
trans-[Co(en)2Cl2]+
(a) Trans isomer of Co(en)2Cl2+ and its mirror image
a
are identical(superimposable)
(b) cis isomer of Co(en)2Cl2+
No Optical activity
Does have Optical activity
Properties of Optical Isomers
 Enantiomers
– possess many identical properties
• solubility, melting point, boiling point, color,
chemical reactivity (with nonchiral reagents)
– different in:
• interactions with plane polarized light
• reactivity with “chiral” reagents
Example
Optical Isomers
polarizing
filter
light
source
plane
polarized
light
unpolarized
light
(random
vibrations)
(vibrates in one plane)
Optical Isomers
(+) dextrorotatory (right rotation )
(-) levorotatory (left rotation )
polarizing filter
plane
polarized
light
optically active sample
in solution
rotated polarized
light
polarizing filter
plane
polarized
light
optically active sample
in solution
Dextrorotatory (d) = right rotation
Levorotatory (l) = left rotation
Racemic mixture = equal amounts of
two enantiomers; no net rotation
rotated polarized
light
Geometry of NH3
NH3
NH3
2+
Zn
NH3
NH3
2+
Geometry of [Zn(NH3)4] 2+
2+ [Ar]18
Zn
30
4s
4p
3d10
sp3
[Zn(NH3)4] 2+
[Ar]18
3d10
Tetrahedron
4 个sp3
Crystal-Field Theory
When ligands coordinate to a metal ion, they increase the
energies of the metal‘s d orbitals by ligand d-orbital
repulsion. Because of their spatial arrangement, the d
orbitals do not all experience the same increase in energy.
If we consider the formation of an octahedral complex in
which the ligands approach the metal ion along the x, y,
and z axes, we see that only orbitals that lie along the axes
(the dz2 and dx2-y2) are approached directly.
(红宝石)
Crystal-Field Theory
 Crystal field theory describes bonding in
transition metal complexes.
 The formation of a complex is a Lewis acid-base
reaction.
 Both electrons in the bond come from the ligand
and are donated into an empty, hybridized
orbital on the metal.
 Charge is donated from the ligand to the metal.
 Assumption in crystal field theory: the
interaction between ligand and metal is
electrostatic.
 The more directly the ligand attacks the metal
orbital, the higher the energy of the d orbital.
Crystal Field Theory
Octahedral Crystal Field
-
(-) Ligands attracted to (+)
metal ion; provides stability
-
+
d orbital e-’s repulsed by (–)
ligands; increases d orbital
potential energy
-
-
ligands approach along x, y, z axes
greater electrostatic repulsion = higher potential energy
less electrostatic repulsion = lower potential energy
Splitting of d orbital of metal atoms
 The complex metal ion has a
lower energy than the
separated metal and ligands.
 In an octahedral field, the five
d orbitals do not have the same
energy: three degenerate
orbitals are higher energy than
two degenerate orbitals.
 The energy gap between them
is called , the crystal field
splitting energy.
Splitting of d-orbital Energies by a
Tetrahedral Field and a Square Planar
Field of Ligands
Factors affecting the magnitude of ∆
1.
effect of ligands
2. Effect of metal atoms
[Co(H2O)6]3+ △o=18600 cm－1
[Co(H2O)6]2+ △o= 9300 cm－1
[Co(NH3)6]3+ △o= 23000 cm－1
[Co(NH3)6]2+ △o= 10100 cm－1
[Co(NH3)6]3+ △o=23000 cm－1
[Rh(NH3)6]3+ △o= 33900 cm－1
[Ir(NH3)6]3+
△o= 40000 cm－1
Electronic Configurations of
Transition Metal Complexes
• Expected orbital filling tendencies for e-’s:
– occupy a set of equal energy orbitals one at a time
with spins parallel (Hund’s rule)
• minimizes repulsions
– occupy lowest energy vacant orbitals first
• These are not always followed by transition
metal complexes.

d orbital occupancy depends on  and
pairing energy, P
– e-’s assume the electron configuration with the lowest
possible energy cost
– If 
> P ( large; strong field ligand)
• e-’s pair up in lower energy d subshell first
– If 
< P ( small; weak field ligand)
• e-’s spread out among all d orbitals before any pair
up
d-orbital energy level diagrams
octahedral complex
d1
d-orbital energy level diagrams
octahedral complex
d2
d-orbital energy level diagrams
octahedral complex
d3
d-orbital energy level diagrams
octahedral complex
d4
high spin
<P
low spin
>P
d-orbital energy level diagrams
octahedral complex
d5
high spin
<P
low spin
>P
d-orbital energy level diagrams
octahedral complex
d6
high spin
<P
low spin
>P
d-orbital energy level diagrams
octahedral complex
d7
high spin
<P
low spin
>P
d-orbital energy level diagrams
octahedral complex
d8
d-orbital energy level diagrams
octahedral complex
d9
d-orbital energy level diagrams
octahedral complex
d10
Electronic Configurations of
Transition Metal Complexes
 Determining d-orbital energy level diagrams:
– determine oxidation # of the metal
– determine # of d e-’s
– determine if ligand is weak field or strong field
– draw energy level diagram
High-spin and low-spin complex
ions of Mn2+
Orbital Occupancy
for High- and LowSpin Complexes of
d4 Through d7 Metal
Ions
High spin
weak-field
ligand
d4
d5
d6
d7
Low spin
strong-field
ligand
Application of crystal field theory
Crystal field theory can offer a good explanation
of many properties of transition metal, such as
stability、
spectrum、
and magnetism.
Properties of Transition Metal
Complexes
 Properties of transition metal complexes:
– usually have color
• dependent upon ligand(s) and metal ion
– many are paramagnetic
• due to unpaired d electrons
• degree of paramagnetism dependent on
ligand(s)
– [Fe(CN)6]3- has 1 unpaired d electron
– [FeF6]3- has 5 unpaired d electrons
 Crystal Field Theory
– Can be used to account for
• Colors of transition metal complexes
– A complex must have partially filled d
subshell on metal to exhibit color
– A complex with 0 or 10 d e-s is colorless
• Magnetic properties of transition metal
complexes
– Many are paramagnetic
– # of unpaired electrons depends on the
ligand
1. Account for the color of complexs

Compounds/complexes that have color:
– absorb specific wavelengths of visible light
(400 –700 nm)
• wavelengths not absorbed are transmitted
• color observed = complementary color of
color absorbed
Visible Spectrum
wavelength, nm
(Each wavelength corresponds to a different color)
400 nm
700 nm
higher energy
lower energy
White = all the colors (wavelengths)
absorbed
color
observed
color
Relation Between Absorbed and
Observed Colors
Absorbed
Color
 (nm)
Violet
Blue
Blue-green
Yellow-green
Yellow
Orange
Red
400
450
490
570
580
600
650
Observed
Color
Green-yellow
Yellow
Red
Violet
Dark blue
Blue
Green
 (nm)
560
600
620
410
430
450
520
Color & Visible Spectra
Color& Spectra
 The plot of absorbance versus
wavelength is the absorption
spectrum.
—For example, the absorption
spectrum for [Ti(H2O)6]3+ has
a maximum absorption occurs
at 510 nm (green and yellow).
—So, the complex transmits all
light except green and yellow.
—Therefore, the complex
appears to be purple.
 Absorption of UV-visible radiation by
atom, ion, or molecule:
– Occurs only if radiation has the energy needed
to raise an e- from its ground state to an
excited state
• i.e., from lower to higher energy orbital
• light energy absorbed = energy difference
between the ground state and excited state
• “electron jumping”
white
light
red light
absorbed
For transition metal
complexes,  corresponds to
energies of visible light.
green light
observed
Absorption raises an
electron from the lower d
subshell to the higher d
subshell.
 Different complexes exhibit different colors
because:
– color of light absorbed depends on 
• larger  = higher energy light (Shorter wavelengths)
absorbed
• smaller  = lower energy light(Longer wavelengths)
absorbed
– magnitude of  depends on:
• ligand(s)
• metal
white
light
red light
absorbed
(lower
energy
light)
[M(H2O)6]3+
green light
observed
Colors of Transition Metal Complexes
white
light
blue light
absorbed
(higher
energy
light)
[M(en)3]3+
orange light
observed
Colour of transition metal complexes
Ruby
Corundum
Al2O3 with Cr3+ impurities
Sapphire
Corundum
octahedral metal centre
coordination number 6
Al2O3 with Fe2+ and Ti4+ impurities
Emerald
Beryl
AlSiO3 containing Be with Cr3+ impurities
The wavelength of light absorbed depends on the size of the
energy gap, ∆ , between lower- and higher-energy d orbitals.
The size of ∆ depends in part on the identity of the ligands
coordinating to the metal ion. Figure below shows a series of
four chromium(III) complexes, each with a different ∆ value.
∆1
3+
24Cr
∆2
∆3
∆4
violet
yellow
yellow
violet
yellow
violet
3d3
green
red
The energy gap between the three lower-energy d
orbitals and the two higher-energy d orbitals is
such that visible light causes this electron
promotion. Thus, a particular wavelength of
visible light is absorbed, giving a coordination
complex its characteristic color. The gap between
lower- and upper-level d orbitals in the
[Ti(H2O)6]3+ ion is of the magnitude that light of
510 nm wavelength promotes an electron.
Absorption of this wavelength of light (510 nm is
a green wavelength) gives the ion its
characteristic red color.
22
Ti
3+
22Ti
[Ar]3d24s2
[Ar]3d 4s
1
Ti: titanium
0
[Ti(H2O)6]3
• Color of a complex depends on: (i) the metal
and (ii) its oxidation state.
• Pale blue [Cu(H2O)6]2+ can be converted into
dark blue [Cu(NH3)6]2+ by adding NH3(aq).
• A partially filled d orbital is usually required
for a complex to be colored.
• So, d 0 metal ions are usually colorless.
Exceptions: MnO4- and CrO42-.
• Colored compounds absorb visible light.
• The color our eye perceives is the sum of the
light not absorbed by the complex.