Chapter 23. Transition Metals and Coordination Chemistry
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 23.1 The Position of the Transition Metals 23.1 The Transition Metals
In the Periodic Table
Figure 23.4 Nonzero Oxidation States of the Period 23.1 The Transition Metals
4 Transition Metals
Figure 23.5 The Relative Orientation of Electron 23.1 The Transition Metals
Spins in Various Types of Compounds
Chlorophyll
Table 23.5 Some Common Ligands and Their 23.4 Nomenclature and Isomerism in Coordination
Names Chemistry
Figure 23.19 Forms of Isomerism in Coordination 23.4 Nomenclature and Isomerism in Coordination
Compounds Chemistry
Figure 23.22 Optical Isomerism 23.4 Nomenclature and Isomerism in Coordination
Chemistry
Figure 23.23 Using Polarized Light to Detect 23.4 Nomenclature and Isomerism in Coordination
Activities: Section:
Geometries of MLn Complexes 23.2 Transition Metal Complexes
Color Wheel 23.5 Color and Magnetism in Coordination
Chemistry
Chapter 23
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Animations: Section:
Isomerism 23.4 Nomenclature and Isomerism in Coordination
Chemistry
Chirality 23.4 Nomenclature and Isomerism in Coordination
Chemistry
Optical Activity 23.4 Nomenclature and Isomerism in Coordination
Chemistry
3-D Models: Section:
cis-tetraamminedichlorocobalt(III) 23.2 Transition Metal Complexes
trans-tetraamminedichlorocobalt(III) 23.2 Transition Metal Complexes
Other Resources
Further Readings: Section:
Trends in Ionization Energy of Transition-Metal 23.1 The Transition Metals
Elements
A Stability Ruler for Metal-Ion Complexes 23.2 Transition Metal Complexes
The Concept of Oxidation States in Metal Complexes 23.2 Transition Metal Complexes
Some Linguistic Detail on Chelation 23.3 Common Ligands in Coordination Chemistry
Selecting and Using Chelating Agents 23.3 Common Ligands in Coordination Chemistry
EDTA-Type Chelating Agents in Everyday 23.3 Common Ligands in Coordination Chemistry
Transition Metals and Coordination Chemistry
335
Iron Deficiency 23.3 Common Ligands in Coordination Chemistry
Introducing Stereochemistry to Non-science Majors 23.4 Nomenclature and Isomerism in Coordination
Chemistry
Live Demonstrations: Section:
The Copper Mirror 23.2 Transition Metal Complexes
Metals in Metal Salts: A Copper Mirror 23.2 Transition Metal Complexes
Demonstration
Cobalt Complexes: Changing Coordination 23.2 Transition Metal Complexes
Numbers
Changing Coordination Numbers: Nickel 23.2 Transition Metal Complexes
Complexes
Chapter 23
336
Chapter 23. Transition Metals and Coordination Chemistry
Common Student Misconceptions
• Students have difficulty naming coordination complexes.
Teaching Tips
• Students should be encouraged to review Chapters 2, 6, 7, 9, 11, 17 and 19 prior to covering this
chapter.
Lecture Outline
23.1 The Transition Metals
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• Transition metals occupy the d block of the periodic table.
• Most metals are found in nature in the form of solid inorganic compounds called minerals.
• Names of minerals are based on the location of their discovery, the person who discovered them,
or some characteristic of the mineral.
• Example: Iron can be separated from gangue in finely ground magnetite by using a magnet to
attract the iron.
• reduction (to obtain the free metal in the 0 oxidation state)
• purifying or refining (to obtain the pure metal)
• mixing with other metals (to form an alloy)
• Alloys are metallic materials composed of two or more elements.
Physical Properties
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• The physical properties of transition metals can be classified into two groups: atomic properties (e.g.,
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• The increase in size of the Cu and Zn triads is rationalized in terms of the completely filled d
orbital.
• In general, atomic size increases down a group.
• An important exception: Hf has almost the same radius as Zr (group 4B); we would expect Hf to be
larger than Zr.
Electron Configurations and Oxidation States
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• Even though the (n −1)d orbital is filled after the ns orbital, electrons are lost from the orbital with the
highest n first.
• That is, transition metals lose s electrons before the d electrons.
• Example: Fe: [Ar]3d64s2 Fe2+: [Ar]3d6.
Magnetism
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• Magnetism provides important bonding information.
• Electron spin generates a magnetic field with a magnetic moment.
• There are several types of magnetic behavior:
• diamagnetic (no atoms or ions with magnetic moments)
• ferromagnetic (coupled magnetic centers aligned in a common direction)
• Ferromagnetism is a special case of paramagnetism where the magnetic moments are
permanently aligned (e.g., Fe, Co, and Ni).
• Ferromagnetic oxides are used in magnetic recording tape (e.g., CrO2 and Fe3O4).
• Two additional types of magnetism involve ordered arrangements of unpaired electrons.
• Antiferromagnetism (the unpaired electrons on a given atom align so that their spins are
oriented in the opposite direction as the spins on neighboring atoms).
• Ferrimagnetism (has characteristics of both a ferromagnet and an antiferromagnet).
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• All magnetically ordered materials become paramagnetic when heated above a critical temperature.
• Curie temperature (Tc): critical temperature for ferromagnets and ferrimagnets.
• Néel temperature (Tn): critical temperature for antiferromagnets.
FORWARD REFERENCES
• Nickel used as a heterogeneous catalyst in hydrogenation of alkenes will be mentioned in
23.2 Transition-Metal Complexes
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• Metal complexes (or complexes) have a metal ion (which can have a 0 oxidation state) bonded to a
number of molecules or ions.
• If the complex has a net electrical charge, it is called a complex ion.
• Compounds that contain complexes are known as coordination compounds.
The Development of Coordination Chemistry: Werner’s Theory
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• Alfred Werner proposed:
• Metal ions exhibit a primary and secondary valence.
• Primary valence: The oxidation state of the metal.
• Secondary valence: The number of atoms directly bonded to the metal ion.
• This is the coordination number.
• The central metal and ligands bound to it are the coordination sphere of the complex.
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“A Stability Ruler for Metal Ion Complexes” from Further Readings
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“The Copper Mirror” from Live Demonstrations
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“Metals in Metal Salts: A Copper Mirror Demonstration” from Live Demonstrations
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• the chlorides are opposite each other.
The Metal-Ligand Bond
• The metal-ligand bond is an interaction between:
• a Lewis acid (the metal ion with its empty valence orbitals) and
Charges, Coordination Numbers, and Geometries
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• The charge on a complex ion equals the sum of the charge on the metal plus the charges on the
ligands.
• In a complex the donor atom is the atom bonded directly to the metal.
• The coordination number is the number of ligands attached to the metal.
• The most common coordination numbers are 4 and 6.
• Some metal ions have a constant coordination number (e.g., Cr3+ and Co3+ have coordination
numbers of 6).
23.3 Common Ligands in Coordination Chemistry
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• A donor atom is the ligand atom that binds to the central metal ion in a coordination complex.
• Monodentate ligands bind through one donor atom only.
• Therefore, they can occupy only one coordination site.
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“Cobalt Complexes: Changing Coordination Numbers” from Live Demonstrations
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Figure 23.9 from Transparency Pack
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“Geometries of MLn Complexes” Activity from Instructor’s Resource CD/DVD
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“Changing Coordination Numbers; Nickel Complexes” from Live Demonstrations
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• Because bidentate and polydentate ligands grasp the metal between two or more donor atoms and are
called chelating agents.
• Example: ethylenediamine (H2NCH2CH2NH2)
• The abbreviation for ethylenediamine is “en.”
• There are two nitrogen atoms that can act as ligands.
• The chelate effect refers to the larger formation constants for polydentate ligands as compared
with corresponding monodentate ligands.
• Chelating agents are sometimes referred to as sequestering agents.
• In medicine, sequestering agents are used to selectively remove toxic metal ions (e.g., Hg2+ and
Pb2+) while leaving biologically important metals.
• One very important chelating agent is ethylenediaminetetraacetate (EDTA4–).
Metals and Chelates in Living Systems
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• Ten of the twenty-nine elements required for human life are transition metals (V, Cr, Mn, Fe, Co, Cu,
Zn, Mo, Cd, and Ni).
• Many natural chelates coordinate to the porphine molecule.
• Porphine forms a tetradentate ligand with the loss of the two protons bound to its nitrogen atoms.
• A porphyrin is a metal complex derived from porphine.
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“EDTA-Type Chelating Agents in Everyday Consumer Products: Some Medicinal and Personal Care
Products” from Further Readings
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“Toxicity of Heavy Metals and Biological Defense: Principles and Applications in Bioinorganic
Chemistry, Part VII” from Further Readings
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“Heme (with bound O2)” 3–D Model from Instructor’s Resource CD/DVD
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• Two important porphyrins are heme (which contains Fe2+) and chlorophyll (which contains
Mg2+).
• Two important heme-containing molecules are myoglobin and hemoglobin.
• These proteins are important oxygen-binding proteins.
• A different metal complex is important in the process of photosynthesis.
• Photosynthesis is the conversion of CO2 and water to glucose and oxygen in plants in the
presence of light.
• The synthesis of one mole of sugar requires the absorption and utilization of 48 moles of photons.
• Chlorophylls are porphyrins that contain Mg(II). Photons of light are absorbed by chlorophyll-
containing pigments in plant leaves.
• Chlorophyll a is the most abundant chlorophyll.
• Plant photosynthesis sustains life on Earth.
FORWARD REFERENCES
• Photosynthesis will be covered in Chapter 24 (section 24.6).
• Tertiary structure of myoglobin will be mentioned in Chapter 24 (section 24.7).
23.4 Nomenclature and Isomerism in Coordination Chemistry
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• We can name complexes in a systematic manner using some simple nomenclature rules.
• For salts, the name of the cation is given before the name of the anion.
• Example: In [Co(NH3)5Cl]Cl2 we name [Co(NH3)5Cl]2+ before Cl–.
• Within a complex ion or molecule, the ligands are named (in alphabetical order) before the metal.
• Example: [Co(NH3)5Cl]2+ is pentaamminechlorocobalt(III).
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“Isomerism” Animation from Instructor’s Resource CD/DVD
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• Example [Co(en)3]Cl3 is tris(ethylenediamine)cobalt(III) chloride.
• If the complex is an anion, the name ends in -ate.
• For example, [CoCl4]2– is the tetrachlorocobaltate(II) ion.
Structural Isomerism
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• Two examples of structural isomerism in coordination chemistry are:
• linkage isomerism
• Linkage isomers: A ligand is capable of coordinating to a metal in two different ways.
• Example: Nitrite can coordinate via a nitrogen or an oxygen atom.
• If the nitrogen atom is the donor atom, the ligand is called nitro.
• [Cr(H2O)4Cl2]Cl.2H2O
Stereoisomerism
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• Stereoisomers have the same connectivity but different spatial arrangements of atoms.
• Two types of stereoisomerism are:
• geometric isomerism
• In geometric isomerism the arrangement of the atoms is different although the same bonds are
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Figure 23.19 from Transparency Pack
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“Pictorial Analogies VIII: Types of Formulas and Structural Isomers” from Further Readings
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“Introducing Stereochemistry to Non–science Majors” from Further Readings
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Figure 23.22 from Transparency Pack
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“Chirality” Animation from Instructor’s Resource CD/DVD
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• optical isomerism
• Optical isomers are nonsuperimposable mirror images.
• These are referred to as enantiomers.
• Complexes that exist as enantiomers are chiral.
• Chiral species are molecules or ions that cannot be superimposed on their mirror image.
• Most physical and chemical properties of enantiomers are identical.
• Therefore, enantiomers are very difficult to separate.
• Optical isomers are differentiated from each other by their interaction with plane-polarized
• Chiral molecules are said to be optically active because of their effect on light.
• Racemic mixtures contain equal amounts of l and d isomers.
• They have no overall effect on the plane of polarized light.
• The 2001 Nobel Prize in Chemistry was awarded to W. S. Knowles and K. B. Sharples of the United
States and R. Noyori of Japan for work on the catalysis of chiral reactions.
FORWARD REFERENCES
23.5 Color and Magnetism in Coordination Chemistry
Color
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• The color of a complex depends on the metal, the ligands present, and the oxidation state of the metal.
• For example, 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.
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• Thus, ions with completely empty (e.g., Al3+ or Ti4+) or completely filled (e.g., Zn2+) d subshells
are usually colorless.
Magnetism of Coordination Compounds
• Many transition-metal complexes are paramagnetic (i.e., they have unpaired electrons).
• Consider a d6 metal ion:
23.6 Crystal-Field Theory
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• 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.
• The dz2 and dx2-y2 orbitals lie on the same axes as negative charges.
• Therefore, there is a large, unfavorable interaction between the ligand and these orbitals.
• These orbitals form the degenerate high-energy pair of energy levels.
• The dxy, dyz, and dxz orbitals are oriented between the negative charges.
• Therefore, there is a smaller repulsion between ligands and these orbitals.
• These orbitals form the degenerate low-energy set of energy levels.
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• This transition is called a d–d transition because it involves exciting an electron from one set
of d orbitals to the other.
• A spectrochemical series is a listing of ligands in order of their ability to increase ∆:
Cl– < F– < H2O < NH3 < en < NO2– (N-bonded) < CN–
• Weak-field ligands lie on the low-∆ end of the spectrochemical series.
• Strong-field ligands lie on the high-∆ end of the spectrochemical series.
• Example: When the ligand coordinated to Cr3+ is changed from the weak-field ligand F– to the
strong-field ligand CN–, ∆ increases and the color of the complex changes from green (in
[CrF6]3+) to yellow (in [Cr(CN)6]3–).
Electron Configurations in Octahedral Complexes
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• Recall that when transition metals form cations, s electrons are lost first.
• Thus, Ti3+ is a d1 ion, V3+ is a d2 ion, and Cr3+ is a d3 ion.
Tetrahedral and Square Planar Complexes
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• By using the same arguments as for the octahedral case, we can derive the relative orbital energies for
d orbitals in a tetrahedral field.
• The splitting of the d orbitals is the opposite of that observed for an octahedral field.
• Because there are only four ligands, ∆ for a tetrahedral field is smaller than ∆ for an octahedral
field.
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Further Readings:
1. Paul S. Matsumoto, “Trends in Ionization Energy of Transition–Metal Elements,” J. Chem. Educ., Vol.
82, 2005, 1660−1661.
2. R. Bruce Martin, “A Stability Ruler for Metal Ion Complexes,” J. Chem. Educ., Vol. 64, 1987, 402.
7. Ei-Ichiro Ochiai, “Toxicity of Heavy Metals and Biological Defense: Principles and Applications in
Bioinorganic Chemistry, Part VII,” J. Chem. Educ., Vol. 72, 1995, 479–484.
8. Colin J. Rix, “The Biochemistry of Some Iron Porphyrin Complexes,” J. Chem. Educ., Vol. 59, 1982,
389–392.
9. N. M. Senozan and R. L. Hunt, “Hemoglobin: Its Occurrence, Structure, and Adaptation,” J. Chem.
Educ., Vol. 59, 1982, 173–178.
13. Nevin S. Scrimshaw, “Iron Deficiency,” Scientific American, October 1991, 46–52.
14. John J. Fortman, “Pictorial Analogies VIII: Types of Formulas and Structural Isomers,” J. Chem.
Educ., Vol. 70, 1993, 755.
15. Hannia Lujan-Upton, “Introducing Stereochemistry to Non–science Majors,” J. Chem. Educ., Vol. 78,
2001, 475–477.
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18. Robert E. Gawley, “Chirality Made Simple: A 1- and 2-Dimensional Introduction to
Stereochemistry,” J. Chem. Educ., Vol. 82, 2005, 1009–1012.
Live Demonstrations:
1. Lee R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “The Copper Mirror,” Chemical
Demonstrations, A Sourcebook for Teachers, Vol. 2 (Washington: American Chemical Society, 1988), pp.
187−188.
4. Lee R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Changing Coordination Numbers: Nickel
Complexes,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1987), pp. 73–74.
5. Lee. R. Summerlin,, Christie L. Borgford, and Julie B. Ealy, “Separating Metallic Iron from Cereal,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), p. 62. Iron is removed from fortified cereal in this simple demonstration.
6. Donald C. Bowman, “A Colorful Look at the Chelate Effect,” J. Chem. Educ., Vol. 83, 2006, 1158–
1160. A set of overhead projector demonstrations.
9. Lee. R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Green and Blue Copper Complexes,”
Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical
Society, 1988), pp.71−72. Three copper complexes are prepared in this demonstration.