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Chapter 23
The Transition Elements and
Their Coordination Compounds
23-1
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The Transition Elements and Their Coordination Compounds
23.1 Properties of the Transition Elements
23.2 The Inner Transition Elements
23.3 Highlights of Selected Transition Metals
23.4 Coordination Compounds
23.5 Theoretical Basis for the Bonding and Properties of Complexes
23-2
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Figure 23.1
23-3
The transition elements (d block) and inner transition
elements (f block) in the periodic table.
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General Properties:
• The transition metals show great similarities
within a given period and a group. Their
Chemistry does not change as number of
valence electron change.
• They are metals good conductors of heat and
electricity: Ex: Ag.
• More than one oxidation state.
23-4
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General Properties:
• Complex ions: Species where transition metal
ion is surrounded by a certain number of
ligands.
• Ligand: Molecules or ions that behave as
Lewis bases.
• Paramagnetic-unpaired electrons.
23-5
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23-6
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Electronic Configurations:
•
•
•
•
•
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Cr:
Cu:
Mo3+
Ag+
The energy of 3d orbitals in transition metal
ions is less than 4s orbitals.
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Atomic and Physical Properties of
Transition Elements:
1. Atomic size decreases from left to right across
the period , but then remain fairly constant.
2. Transition elements exhibit a small change in
electronegativity. The values are intermediate.
3. First Ionization energies increase little.
4. Lanthanide contraction- 4f orbitals are filled,
increases overall charge on nucleus and not the
size.
23-8
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Figure 23.3
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Horizontal trends in key atomic properties of the
Period 4 elements.
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Atomic and Physical Properties of
Transition Elements:
5.Nuclear charge increases down a group.
6. Heavier transition metals exhibit more
covalent character.
7. First I.E increases down a transitional group.
8. Densities increase as atomic mass increases.
23-10
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Figure 23.4
Vertical trends in key properties within the transition elements.
2nd and 3rd element nearly same size Electronegativity increases down a group.
1st IE highest at bottom of trans group. Densities increase as mass increases
23-11
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Chemical Properties
1. They have multiple oxidation states.
2. +2 oxi state is most common as ns2 electrons are
readily lost.
3. Ionic bonding occurs in lower O.S and covalent in
higher O.S.
4. Electrons in a partially filled d sublevels can absorb
visible wavelengths and hence their compounds are
colored.
5. They are paramagnetic (unpaired d electrons)
6. IE1 increases down a group.
23-12
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23-13
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Sample Problem 23.2
PROBLEM:
PLAN:
Finding the Number of Unpaired Electrons
The alloy SmCo5 forms a permanent magent because both
samarium and cobalt have unpaired electrons. How many
unpaired electrons are in the Sm atom (Z = 62)?
Write the condensed configuration of Sm and, using Hund’s
rule and the aufbau principle, place electrons into a partial
orbital diagram.
SOLUTION: Sm is the eighth element after Xe. Two electrons go into the 6s
sublevel and the remaining six electrons into the 4f (which fills
before the 5d).
Sm is [Xe]6s24f6
6s
4f
There are 6 unpaired e- in Sm.
23-14
5d
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Coordination compounds:
• They contain atleast one Complex ion , bonded to
ligands and associated with other counter ions.
• Coordination compound: Complex transitional metal
ion attached to ligands.
• Two types of valance: 1. secondary Valence- Ability
of metal ion to bind to a Lewis base(ligands)Coordination number.
• 2. Primary valence-Ability of metal ion to form ionic
bonds with oppositely charged ions.-Oxidation state.
23-15
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Coordination compounds:
• Complex ions: Species where transition metal
ion is surrounded by a certain number of
ligands.
• Ligand: Molecules or ions that behave as
Lewis bases.
23-16
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Structures of Complex Ions:
Coordination Numbers, Geometries, and Ligands
•Coordination Number - the number of ligand atoms that are bonded
directly to the central metal ion. The coordination number is specific for
a given metal ion in a particular oxidation state and compound.
•Geometry - the geometry (shape) of a complex ion depends on the
coordination number and nature of the metal ion.
•Donor atoms per ligand - molecules and/or anions with one or more
donor atoms that each donate a lone pair of electrons to the metal ion to
form a covalent bond.
23-17
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The coordination number
•
•
•
•
•
Varies from 2-8.
6 ligands-octahedral arrangement.
4-Tetrahedral/Square planar.
2- Linear.
Most common coord. Number is 6.
23-18
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Figure 23.9 Components of a coordination compound.
models
6 ligands-octahedral
23-19
wedge diagrams
chemical formulas
4 ligands-square planar
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23-20
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Ligands:
• Neutral molecule or ion having a lone pair of
electron that can be used to form a bond to
metal ion.
• Metal (Lewis acid-e pair acceptor) _______
Nonmetal ( Lewis base- e pair donor)
• Monodentate/ Unidentate ligand- Ligand
forms 1 bond.CN-, H2O, NH3.
23-21
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Ligands
• Chelating ligands/ Chelates: Ligands have
more than one atom with a lone pair of
electrons that can be used to bond a metal ion.
• Bidentate ligand- can form 2 bonds.Ex:
ethylenediamine(en), oxalate
• Polydentate ligands- can form more than 2
bonds.Ex: EDTA- 6 bonds.
23-22
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23-23
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23-24
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Formulas and Names of Coordination Compounds
Rules for writing formulas:
1. The cation is written before the anion.
2. The charge of the cation(s) is balanced by the charge
of the anion(s).
3. In the complex ion, neutral ligands are written before
anionic ligands, and the formula for the whole ion is
placed in brackets.
23-25
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Formulas and Names of Coordination Compounds
Rules for naming complexes:
continued
1. The cation is named before the anion.
2. Within the complex ion, the ligands are named, in alphabetical
order, before the metal ion.
3. Neutral ligands generally have the molecule name, but there
are a few exceptions. Anionic ligands drop the -ide and add
-o after the root name.
4. A numerical prefix indicates the number of ligands of a
particular type.
5. The oxidation state of the central metal ion is given by a
Roman numeral (in parentheses).
6. If the complex ion is an anion we drop the ending of the metal
name and add -ate.
23-26
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Nomenclature:
•
•
•
•
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Name cation before anion.
In naming a complex ion, name ligands
before the metal ion.
In naming ligands, add o to the root name of
anion(chloro), Use the full name for a neutral
ligand .
Exceptions to # 3: aqua, ammine,
methylamine, carbonyl, nitrosyl.
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Nomenculature
• . Use prefix mono, di, tri, tetra , penta and hexa for
simple ligands.
• 6. Use prefix bis,tris,tetrakis for complicated ligands
that alreadt have bi,tri.
• 7. Oxidation state for metal in Roman numerals in ()
• 8. When more than one type of ligand are present
name alphabetically.
• 9. If complex ion has negative charge add the suffix ate to the name of the metal.(Latin name)
• Iron
copper
lead
silver
gold
tin
23-28
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Sample Problem 23.3
PROBLEM:
Writing Names and Formulas of Coordination
Compounds
(a) What is the systematic name of Na3[AlF6]?
(b) What is the systematic name of [Co(en)2Cl2]NO3?
(c) What is the formula of tetraaminebromochloroplatinum(IV)
chloride?
(d) What is the formula of hexaaminecobalt(III) tetrachloroferrate(III)?
PLAN:
Use the rules presented -
SOLUTION:
and
.
(a) The complex ion is
3-.
[AlF6] Six
(hexa-) fluorines (fluoro-) are the ligands - hexafluoro
Aluminum is the central metal atom - aluminate
Aluminum has only the +3 ion so we don’t need Roman
numerals.
sodium hexafluoroaluminate
23-29
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Sample Problem 23.3
Writing Names and Formulas of Coordination
Compounds
continued
(b) There are two ligands, chlorine and ethylenediamine dichloro, bis(ethylenediamine)
The complex is the cation and we have to use Roman numerals for
the cobalt oxidation state since it has more than one - (III)
The anion, nitrate, is named last.
dichlorobis(ethylenediamine)cobalt(III) nitrate
(c)
Pt4+
ClCl
tetraaminebromochloroplatinum(IV) chloride
4 NH3
Br-
[Pt(NH3)4BrCl]Cl2
(d)
6 NH3
Co3+
4 Cl-
Fe3+
hexaaminecobalt(III) tetrachloro-ferrate(III)
[Co(NH3)6][Cl4Fe]3
23-30
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23-31
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Figure 23.10
Important types of isomerism in coordination compounds.
ISOMERS
Same chemical formula, but different properties
Constitutional (structural) isomers
Stereoisomers
Atoms connected differently
Different spatial arrangement
Coordination
isomers
Linkage
isomers
Ligand and
counter-ion
exchange
Different donor
atom
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Geometric (cistrans) isomers
(diastereomers)
Different
arrangement
around metal ion
Optical isomers
(enantiomers)
Nonsuperimposable
mirror images
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Isomerism:
• Same formula but different properties.
• Structural isomerism: Isomers contain same
atoms but different bonds.
– Coordination isomerism: Composition of complex
ion varies.
• Linkage isomerism: Point of attachment of
atleast one of the ligands differs
23-33
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Isomerism
•
Stereoisomers: All bonds are same but
different spatial arrangements.
–
–
Geometrical isomerism –cis-trans-Atoms or
group of atoms can assume different positions
around a rigid ring or bond.
Optical Isomerism: Have opposite effects on
plane polarized light.
•
•
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Chiral: Objects that have nonsuperimposable mirror
images.
Enantiomers: Isomers that are nonsuperimposable
mirror images of each other.
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Linkage isomers
23-35
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Figure 23.11
23-36
Geometric (cis-trans) isomerism.
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Figure 23.12
Optical isomerism in an
octahedral complex ion.
23-37
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Sample Problem 23.4
PROBLEM:
PLAN:
Determining the Type of Stereoisomerism
Draw all stereoisomers for each of the following and state the type
of isomerism:
(a) [Pt(NH3)2Br2]
(b) [Cr(en)3]3+ (en = H2NCH2CH2NH2)
Determine the geometry around each metal ion and the nature of
the ligands. Place the ligands in as many different positions as
possible. Look for cis-trans and optical isomers.
SOLUTION: (a) Pt(II) forms a square planar complex and there are two pair
of monodentate ligands - NH3 and Br.
Br
NH3
H3N
Pt
H3N
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Pt
Br
trans
Br
H3N
Br
cis
These are geometric isomers;
they are not optical isomers
since they are superimposable
on their mirror images.
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Sample Problem 23.4
Determining the Type of Stereoisomerism
(b) Ethylenediamine is a bidentate ligand. Cr3+ is
hexacoordinated and will form an octahedral geometry.
continued
Since all of the ligands are identical, there will be no geometric isomerism
possible.
3+
3+
N
N
N
N
N
N
Cr
Cr
N
The mirror images are
nonsuperimposable
and are therefore
optical isomers.
N
N
N
N
N
rotate
3+
N
N
N
Cr
N
N
N
23-39
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Figure 23.13
Hybrid orbitals and bonding in the octahedral [Cr(NH3)6]3+ ion.
23-40
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Figure 23.14
Hybrid orbitals and bonding in the square planar [Ni(CN)4]2- ion.
23-41
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Figure 23.15
Hybrid orbitals and bonding in the tetrahedral [Zn(OH)4]2- ion.
23-42
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The Crystal Field Model
• It focuses on the energies of d orbitals.
• Metal –ligand bond is ionic.
• Ligands are negative point charges.
23-43
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Figure 23.16
23-44
An artist’s wheel.
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23-45
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Octahedral complexes:
• Dz2 and dx2-y2 orbitals have lobes that point
directly at the ligands.
• Dxy, dyz,dxy point their lobes between
charges.
• Electrons fill the d orbitals farthest from the
ligands to minimize repulsion.
• Dxy, dyz,dxy ( t2g set) are at lower energy in
octahedral complex first.
• Dz2 and dx2-y2 (eg set) is at higher energy.
23-46
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Octahedral complexes:
• Splitting of 3d orbital energies explains color
and magnetism.
• Strong field case- splitting produced by ligands
is very large, electrons will pair in lower t2g
orbitals. Diamagnetic (all electrons are paired)
• Weak field case-splitting produced by ligands
is small, electrons will occupy all 5 orbitals
before pairing occurs. Paramagnetic( unpaired
electrons)
23-47
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Figure 23.17
23-48
The five d-orbitals in an octahedral field of ligands.
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Problems:
1.Fe (CN)63- has one unpaired electron . Does
the CN- ligand produce a strong or weak
field?
2.Predict the number of unpaired electrons in the
complex ion [Cr(CN)6] 4 –
23-49
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Figure 23.22
The spectrochemical series.
•For a given ligand, the color depends on the oxidation state of the metal ion.
•For a given metal ion, the color depends on the ligand.
I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO
WEAKER FIELD
23-50
STRONGER FIELD
SMALLER D
LARGER D
LONGER 
SHORTER 
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Color of octahedral compounds:
• Absorbed color is different than observed
color.
• Transition metals absorb colors in the visible
region.
• ∆E=hc/ , ∆E= energy spacing, =walength
needed to move an electron from t2g to eg.
• Color of solution changes as the ligand
changes.
23-51
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In tetrahedral complexes:
• None of the 3d orbitals point at the ligands.
• Tetrahedral splitting is 4/9 times that of
octahedral.
• Dxy, dyz,dxy are closer to pint charges than
Dz2 and dx2-y2.
• Weak field case always applies in tetrahedral
complexes.
23-52
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Figure 23.18
Splitting of d-orbital energies by an octahedral field of
ligands.
D is the splitting energy
23-53
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Figure 23.19
23-54
The effect of ligand on splitting energy.
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Figure 23.20
23-55
The color of [Ti(H2O)6]3+.
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Figure 23.21
Effects of the metal oxidation state and of ligand identity on color.
[V(H2O)6]3+
[V(H2O)6]2+
[Cr(NH3)6]3+
23-56
[Cr(NH3)5Cl ]2+
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Problem:
Give the crystal field diagram for tetrahedral
complex ion CoCl 4 2-
23-57
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Sample Problem 23.5
PROBLEM:
Ranking Crystal Field Splitting Energies for
Complex Ions of a Given Metal
Rank the ions [Ti(H2O)6]3+, [Ti(NH3)6]3+, and [Ti(CN)6]3- in terms of
the relative value of D and of the energy of visible light absorbed.
PLAN: The oxidation state of Ti is 3+ in all of the complexes so we are
looking at the crystal field strength of the ligands. The stronger the
ligand the greater the splitting and the higher the energy of the light
absorbed.
SOLUTION:
The field strength according to
is CN- > NH3 > H2O. So the
relative values of D and energy of light absorbed will be
[Ti(CN)6]3- > [Ti(NH3)6]3+ > [Ti(H2O)6]3+
23-58
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Figure 23.23
23-59
High-spin and low-spin complex ions of Mn2+.
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Figure 23.24 Orbital occupancy for high- and low-spin complexes
of d4 through d7 metal ions.
high spin:
weak-field
ligand
23-60
low spin:
strong-field
ligand
high spin:
weak-field
ligand
low spin:
strong-field
ligand
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Sample Problem 23.6
PROBLEM:
PLAN:
Identifying Complex Ions as High Spin or Low Spin
Iron (II) forms an essential complex in hemoglobin. For each of the
two octahedral complex ions [Fe(H2O)6]2+ and [Fe(CN)6]4-, draw an
orbital splitting diagram, predict the number of unpaired electrons,
and identify the ion as low or high spin.
The electron configuration of Fe2+ gives us information that the
iron has 6d electrons. The two ligands have field strengths shown
in
.
potential energy
Draw the orbital box diagrams, splitting the d orbitals into eg and
t2g. Add the electrons noting that a weak-field ligand gives the
maximum number of unpaired electrons and a high-spin complex
and vice-versa.
[Fe(CN)6]42+
[Fe(H2O)6]
SOLUTION:
4 unpaired e-eg
(high spin)
eg
no unpaired e-(low spin)
t2g
t2g
23-61
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Figure 23.25
Splitting of d-orbital energies by a tetrahedral field
and a square planar field of ligands.
tetrahedral
square planar
23-62
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Figure B23.1
23-63
Hemoglobin and the octahedral complex in heme.
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23-64
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Figure B23.2
23-65
The tetrahedral Zn2+ complex in carbonic anhydrase.