1. Transition Elements and Their Coordination Compounds

2. The Transition Elements and Their Coordination Compounds
1. Overview of the Transition Elements
2. Coordination Compounds
Properties
Naming
3. Crystal field theory on the Bonding and Properties of Complexes

3. The transition elements (d block) and inner transition elements (f block) in the periodic table.

4. Properties of the Transition Metals
Transition metals (d-block and f-block) comprise of majority of all metals
Many transition metal compounds are colored and paramagnetic, whereas most main-group ionic compounds are colorless and diamagnetic.
The properties of transition metal compounds are related to the electron configuration of the metal ion.

5. Appearance of pure transition metals: The Period 4 transition metals.

6. Electron Configurations of Transition Metals and their Ions
The d-block elements have the general condensed ground-state configuration [noble gas]ns2(n – 1)dx where n = 4 to 7 and x = 1 to 10.
Periods 6 and 7 elements include the f sublevel:
[noble gas]ns2(n – 2)f14(n – 1)dx where n = 6 or 7.
Transition metals form ions through the loss of the ns electrons before the (n – 1)d electrons.

7. Orbital Occupancy of the Period 4 Transition Metals
The number of unpaired electrons increases in the first half of the series and decreases in the second half, when pairing begins.

8. Trends in key atomic properties of Period 4 elements.

9. Why such Trends in the Properties of Transition Metals?
Atomic size decreases at first, then remains relatively constant.
- The 3d electrons fill inner orbitals, so they shield outer electrons very efficiently and the 4s electrons are not pulled closer by the increasing nuclear charge.
Electronegativity and ionization energies also increase relatively little across the transition metals of a particular period.

10. Trends in the Properties of Transition Metals
Within a group the trends also differ from those observed for main group elements.
Atomic size increases from Period 4 to 5, but not from Period 5 to 6.
- A Period 6 element has 32 more protons than its preceding Period 5 group member instead of only 18.
- The extra shrinkage from the increase in nuclear charge (called the lanthanide contraction) is roughly equal to the normal size increase due to adding an extra energy level.

11. Trends in Each Group of Transition Metals: Think about Cu  Ag  Au
Electronegativity increases within a group from Period 4 to 5, generally similar between Period 5 to 6.
Although atomic size increases slightly down the group, nuclear charge increases much more, leading to higher EN values.
Ionization energy values generally increase down a transition group
Density increases dramatically down a group: Atomic volumes change little while atomic masses increase significiantly.

12. Vertical trends in key properties within the transition elements.

13. Oxidation States of Transition Metals
Most transition metals have multiple oxidation states.
The highest oxidation state for elements in Groups 3B(3) through 7B(7) equals the group number.
- These states are seen when the elements combine with the highly electronegative oxygen or fluorine.
Elements in Groups 8B(8), 8B(9) and 8B(10) exhibit fewer oxidation states. The higher oxidation state is less common and never equal to the group number.
- The +2 oxidation state is common because the ns2 electrons are readily lost.

14. Aqueous oxoanions of transition elements.
Mn2+
MnO42−
MnO4− +2 +6 +7
The highest oxidation state for Mn equals its group number.
VO43−
Cr2O72−
MnO4− +5 +6 +7
Transition metal ions are often highly colored.

15. Oxidation States and d-Orbital Occupancy of the Period 4 Transition Metals*

16. Metallic Behavior of Transition Metals
The lower the oxidation state of the transition metal, the more metallic its behavior.
TiCl2 : melting point > 1000C; ionic
TiCl4 : melting point -24 C, covalent
Metal oxides become less basic (more acidic) as the oxidation state increases.
A metal atom in a positive oxidation state has a greater attraction for bonded electrons, and therefore a greater effective electronegativity than in the zero oxidation state. This effect increases as its oxidation state increases.

17. Color and Magnetic Behavior
Most main-group ionic compounds are colorless and diamagnetic because the metal ion has no unpaired electrons.
Many transition metal ionic compounds are highly colored and paramagnetic because the metal ion has one or more unpaired electrons.
Transition metal ions with a d0 or d10 configuration are also colorless and diamagnetic.

18. Colors of representative compounds of the Period 4 transition metals.
potassium ferricyanide
nickel(II) nitrate hexahydrate
zinc sulfate heptahydrate
titanium(IV) oxide
sodium chromate
scandium oxide
vanadyl sulfate dihydrate
manganese(II) chloride tetrahydrate
cobalt(II) chloride hexahydrate
copper(II) sulfate pentahydrate

19. Some Properties of Group 6B(6) Elements
Atomic Radius (pm)
IE1(kJ/mol)
E° (V) for M3+(aq)/M(s) Cr
128
653
-0.74 Mo
139
685
-0.20 W
770
-0.11
IE1 increases down the group, so reactivity decreases. This trend is opposite to that seen in main-group elements.

20. Lanthanides and Actinides
Lanthanides: aka Rare earth elements.
Their atomic properties vary little across the period, and chemical properties are also very similar.
Most have the ground-state electron configuration [Xe]6s24fx5d0.
Actinides: radioactive, and very similar physical and chemical properties.
The +3 oxidation state is common for both lanthanides and actinides.

21. Writing Electron Configurations of Transition Metal Atoms and Ions
PROBLEM:
Write condensed electron configurations for the following: (a) Zr; (b) V3+; (c) Mo3+. (Assume that elements in higher periods behave like those in Period 4.)
[Kr]5s24d2
[Ar]3d2
[Kr]4d3

22. Coordination Compounds
A coordination compound contains at least one complex ion (a central metal cation bonded to molecules and/or anions called ligands).
The complex ion is associated with counter ions of opposite charge.
Complex ion [Cr(NH3)6]3+ = central Cr3+ ion + 6 NH3 ligands.
The complex ion behaves like a polyatomic ion in solution.

23. Coordination Number
The coordination number: the number of ligand atoms bonded directly to the central metal ion.
Coordination number is specific for a given metal ion in a particular oxidation state and compound.
- [Cr(NH3)6]3+ has a coordination number of 6.
The most common coordination number in complex ions is 6, but 2 and 4 are often seen.
- [CuCl4]2+ has a coordination number of 4.
- [Ag(CN)2]+ has a coordination number of 2.

24. Components of a coordination compound.
[Co(NH3)6]Cl3 dissolves in water. The six ligands remain bound to the complex ion.
[Pt(NH3)4]Br2 has four NH3 ligands and two Br- counter ions.

25. Coordination Numbers and Shapes of Some Complex Ions
Examples 2
Linear
[CuCl2]-, [Ag(NH3)2]+, [AuCl2]- 4
Square planar
[Ni(CN)4]2-, [PdCl4]2-, [Pt(NH3)4]2+, [Cu(NH3)4]2+
Tetrahedral
[Cu(CN)4]3-, [Zn(NH3)4]2+, [CdCl4]2-. [MnCl4]2- 6
Octahedral
[Ti(H2O)6]3+, [V(CN)6]4-, [Cr(NH3)4Cl2]+, [Mn((H2O6]2+, [FeCl6]3-, [Co(en)3]3+
The geometry of a given complex ion depends both on the coordination number and the metal ion.
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26. Ligands
The ligands of a complex ion are molecules or anions with one or more donor atoms. Fluoride ion in AlF63-
Each donor atom donates a lone pair of electrons to the metal ion to form a covalent bond.
Ligands are classified in terms of their number of donor atoms, or “teeth”:
- Monodentate ligands bond through a single donor atom.
- Bidentate ligands have two donor atoms, each of which bonds to the metal ion.
- Polydentate ligands have more than two donor atoms.

27. Some Common Ligands in Coordination Compounds

28. Chelates
Bidentate and polydentate ligands give rise to rings in the complex ion. Ring structures are often very stable.
A complex ion containing this type of structure is called a chelate because the ligand seems to grab the metal ion like claws.
Chelation therapy:
EDTA has six donor atoms and forms very stable complexes with metal ions. Treat lead poisoning, calcium deposit.

29. Formulas of Coordination Compounds
A coordination compound may consist of
a complex cation with simple anionic counterions,
a complex anion with simple cationic counterions, or
a complex cation with complex anion as counterion.
When writing the formula for a coordination compound
the cation is written before the anion,
the charge of the cation(s) is/are balanced by the charge of the anion(s), and
neutral ligands are written before anionic ligands, and the formula of the whole complex ion is placed in square brackets.

30. Determining the Charge of the Metal Ion
The charge of the cation(s) is/are balanced by the charge of the anion(s).
K2[Co(NH3)2Cl4]
First, determine anion
NH3 ligands are 0, Cl- ligands are -1 each.
Co2+, cobalt(II) ion.

31. Naming Coordination Compounds
The cation is named before the anion.
In the complex ion, ligands in alphabetical order before the metal ion.
Anionic ligands drop the –ide and add –o after the root name. Cl-
A numerical prefix to indicate the number of ligands (di-, tri-, tetra-, penta-, hexa-). CuCl42-: tetrachloro
Prefixes do not affect the alphabetical order of ligand names.
Ligands that has a numerical prefix: Prefixes bis (2), tris (3), or tetrakis (4). Ethylenediamine (H2NCH2CH2NH2)
A Roman numeral is used to indicate the oxidation state for a metal that can have more than one state. Exceptions: Zn2+, Ag+, Al3+, etc.
complex anion: the metal ends with –ate. CuCl42-: tetrachlorocuperate(II)

32. Names of Some Neutral and Anionic Ligands
Formula
Aqua
H2O
Fluoro F-
Ammine
NH3
Chloro
Cl-
Carbonyl CO
Bromo
Br-
Nitrosyl NO
Iodo I-
Hydroxo
OH-
Cyano
CN-

33. Names of Some Metal Ions in Complex Anions: Not system name
Name in Anion
Iron
Ferrate
Copper
Cuprate
Lead
Plumbate
Silver
Argentate
Gold
Aurate
Tin
Stannate

34. Writing Names of Coordination Compounds
PROBLEM:
(a) What is the systematic name of Na3[AlF6]?
(b) What is the sytematic name of [Co(en)2Cl2]2SO4?
Naming rules:
Cation + anion
Prefix? + Ligand (alphabetical order) + Metal ion (Roman numeral?)
Anionic ligand: –ide  –o
bis (2), tris (3), or tetrakis (4) for ligands with prefixes in the name. E.g., Ethylenediamine.
Neutral ligands: ammine, aqua, carbonyl, nitrosyl
Complex anion ends with –ate.

35. Formula of Coordination Compounds
Cation + Anion
Brackets are used for complex ions: [Fe(CN)6]
Within the complex ion, the central metal ion is followed by ligands
Ligands are placed in the alphabetical order of their names. Example: [Co(NH3)4Br2]+
Suffix is used to indicate the number of ligands of a particular type
The charges of complex ions are to be balanced by the counter ions: [Co(NH3)4Br2]3[AlF6]

36. Writing Formulas of Coordination Compounds
PROBLEM:
tetraamminebromochloroplatinum(IV) chloride?
triaamminebromodichlorocobalt(IV) tetrachloroferrate(III)?
The cation + the anion.
Brackets for complex ions
Ligands before metal ion
Ligands in the alphabetical order of their names. Example: [Co(NH3)4Br2 ]+
Suffix is used to indicate the number of ligands
Balance the charge

37. Constitutional Isomers of Coordination Compounds
Constitutional isomers: Compounds with the same formula, but with the atoms connected differently.
Coordination isomers: the composition of the complex ion, but not the compound, is different.
- This can occur by the exchange of a ligand and a counter ion, or be by the exchange of ligands. [Co(en)2Cl2]NO2 vs. [Co(en)2Cl(NO2)]Cl
Linkage isomers: the composition of the complex ion is the same but the ligand donor atom is different.
- Some ligands can bind to the metal through either of two donor atoms. Nitrogen and Oxygen atoms in nitrite ion.

38. A pair of linkage (constitutional) isomers
The nitrite ion can bind either through the N atom or either one of the O atoms.

39. Ligands that have more than one donor atom

40. Stereoisomers of Coordination Compounds
Stereoisomers: compounds that have the same atomic connections but different spatial arrangements of their atoms.
Geometric or cis-trans isomers occur when atoms or groups can either be arranged on the same side or on opposite sides of the compound relative to the central metal ion.
Optical isomers (enantiomers) are non-superimposable mirror images of each other.

41. Geometric (cis-trans) isomerism.
The cis and trans isomers of [Pt(NH3)2Cl2].
The cis isomer (cisplatin) is an antitumor agent while the trans isomer has no antitumor effect.

42. Geometric (cis-trans) isomerism.
The cis and trans isomers of [Co(NH3)4Cl2]+.
Note the placement of the Cl- ligands (green spheres).

43. Optical isomerism in an octahedral complex ion:
Right hand vs. Left hand
Structure I and its mirror image, structure II, are optical isomers of cis-[Co(en)2Cl2]+.

44. Practice: Which one of the following structures is different from the other three?
wedged bond: closer to the viewer (“out of the screen”)
dashed bond: object farther from the viewer (“into the screen”)

45. Bonding in Complex Ions
When a complex ion is formed, each ligand donates an electron pair to the metal ion.
The ligand acts as a Lewis base, while the metal ion acts as a Lewis acid.
This type of bond is called a coordinate covalent bond since both shared e- originate from one atom in the pair.
In terms of valence bond theory, the filled orbital of the ligand overlaps with an empty orbital of the metal ion.
The VB model proposes that the geometry of the complex ion depends on the hybridization of the metal ion.

46. Review: Number of Effective Pairs (VSEPR groups), Molecular geometry, and the Hybrid Orbital Set Required

47. Hybrid orbitals and bonding in the octahedral [Cr(NH3)6]3+ ion.

48. Hybrid orbitals and bonding in the square planar [Ni(CN)4]2- ion.

49. Hybrid orbitals and bonding in the tetrahedral [Zn(OH)4]2- ion.

50. Colors of coordination complexes
Each color has a complementary color; the one opposite it on the artist’s wheel.
The color an object exhibits depends on the wavelengths of light that it absorbs.
An object will have a particular color because
it reflects light of that color, or
it absorbs light of the complementary color.

51. Wavelength: Relation Between Absorbed and Observed Colors
Absorbed Color
λ (nm)
Observed Color
Violet
400
Green-yellow
560
Blue
450
Yellow
600
Blue-green
490
Red
620
Yellow-green
570
410
580
Dark blue
430
Orange
650
Green
520

52. Crystal Field Theory
Crystal field theory explains color and magnetism in terms of the effect of the ligands on the energies of the d-orbitals of the metal ion.
The bonding of the ligands to the metal ion cause the energies of the metal ion d-orbitals to split.
Although the d-orbitals of the unbonded metal ion are equal in energy, they have different shapes, and therefore different interactions with the ligands.
The splitting of the d-orbitals depends on the relative orientation of the ligands and the structure of ligands

53. Higher energy orbitals
The five d-orbitals in an octahedral field of ligands.
Higher energy orbitals
Lower energy orbitals
The ligands approach along the x, y and z axes. Two of the orbitals point directly at the ligands, while the other three point between them.

54. Orbital Splitting Diagram: d-orbital splitting in an octahedral field of ligands
The d orbitals split into two groups.
The difference in energy between these groups is called the crystal field splitting energy (Δ).

55. The effect of ligands and splitting energy on orbital occupancy
Weak field ligands lead to a smaller splitting energy.
Electrons tend NOT share orbitals by filling higher energy orbitals.
Strong field ligands lead to a larger splitting energy. Electrons tend to share (pairing up) low energy orbitals.

56. The color of [Ti(H2O)6]3+ The hydrated Ti3+ ion is purple.
Green and yellow light (~500 nm) are absorbed while other wavelengths are transmitted. This gives a purple color.

57. Crystal field theory on the color of [Ti(H2O)6]3+
When the ion absorbs light, electrons can move from the lower t2g energy level to the higher eg level.
The difference in energy between the levels (Δ) determines the wavelengths of light absorbed.
The visible color is given by the combination of the wavelengths transmitted.

58. The Colors of Transition Metal Complexes
The color of a coordination compound is determined by the Δ of its complex ion.
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.

59. Effects of oxidation state and ligand on color.
[V(H2O)6]2+
[V(H2O)6]3+
[Cr(NH3)6]3+
[Cr(NH3)5Cl ]2+
Substitution of an NH3 ligand with a Cl- ligand affects the color of the complex ion.
A change in oxidation state causes a change in color.

60. Effect of Ligands on Splitting Energy: The spectrochemical series.
I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO
WEAKER FIELD
STRONGER FIELD
LARGER D
SMALLER D
LONGER 
SHORTER 
As Δ increases, shorter wavelengths (higher energies) of light must be absorbed to excite electrons.
For reference H2O is considered a weak-field ligand.

61. Magnetic Properties of Transition Metal Complexes
Magnetic properties are determined by the number of unpaired electrons in the d orbitals of the metal ion.
#Unpaired e- : 0 ~ 5 (low spin ~ high spin).
Hund’s rule states that e- occupy orbitals of equal energy one at a time. When all lower energy orbitals are half-filled:
- The next e- can enter a half-filled orbital and pair up by overcoming a repulsive pairing energy, (Epairing).
- The next e- can enter an empty, higher, energy orbital by overcoming Δ.
The number of unpaired e- will depend on the relative sizes of Epairing and Δ.

62. Spectrochemical series vs
Spectrochemical series vs. magnetic property of octahedral complex ions of Mn2+ : [Mn(H2O)6]2+ vs. [Mn(CN)6]4-
I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO

63. Orbital occupancy for high-spin and low-spin octahedral complexes of d4 through d7 metal ions.
high spin: weak-field ligand
low spin: strong-field ligand
high spin: weak-field ligand
low spin: strong-field ligand

64. Application of crystal field theory
Example:
Iron (II) forms a complex in hemoglobin. For each of the two octahedral complex ions [Fe(H2O)6]2+ and [Fe(CN)6]4,
A. draw an energy diagram showing orbital splitting
B. identify the ion as low spin or high spin.
C. Which complex ion has absorption at longer wavelength?
I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO

65. Splitting of d-orbital energies by a tetrahedral field of ligands.
The splitting of d-orbital energies is less in a tetrahedral than an octahedral complex, and the relative d-orbital energies are reversed. Only high-spin tetrahedral complexes are known because Δ is small.

66. Splitting of d-orbital energies by a square planar field of ligands.
Square planar complexes are low-spin and usually diamagnetic because the four pairs of d electrons fill the four lowest-energy orbitals.

67. Hemoglobin and the octahedral complex in heme.
Chemical Connections
Hemoglobin and the octahedral complex in heme.
Hemoglobin consists of four protein chains, each with a bound heme. In oxyhemoglobin (B), the octahedral complex in heme has an O2 molecule as the sixth ligand for iron(II).
(Illustration by Irving Geis. Rights owned by Howard Hughes Medical Institute. Not to be used without permission.)

68. Some Transition Metal Trace Elements in Humans

69. Chemical Connections
The tetrahedral Zn2+ complex in carbonic anhydrase.

70. Determining the Charge of the Metal Ion
The charge of the cation(s) is/are balanced by the charge of the anion(s).
K2[Co(NH3)2Cl4] contains a complex anion.
First, determine anion as [Co(NH3)2Cl4]2-.
Then, determine charge of the central metal ion, based on the charges of the ligands: NH3 ligands are 0, Cl- ligands are -1 each.
Charge of complex ion = charge of metal ion + total charge of ligands
= charge of metal ion + [(2 x 0) + (4 x -1)]
Charge of metal ion = (-2) – (-4) = +2 or 2+
The metal ion in this complex anion is Co2+, cobalt(II) ion.

71. Writing Names and Formulas of Coordination Compounds
PROBLEM:
(a) What is the systematic name of Na3[AlF6]?
(b) What is the sytematic name of [Co(en)2Cl2]NO3?
(c) What is the formula of tetraamminebromochloroplatinum(IV) chloride?
(d) What is the formula of hexaamminecobalt(III) tetrachloroferrate(III)?
sodium hexafluoroaluminate
dichlorobis(ethylenediamine)cobalt(III) nitrate
[Pt(NH3)4BrCl]Cl2
[Co(NH3)6][FeCl4]3

72. [Ti(CN)6]3- > [Ti(NH3)6]3+ > [Ti(H2O)6]3+
Practice:
PROBLEM:
Rank the ions [Ti(H2O)6]3+, [Ti(NH3)6]3+, and [Ti(CN)6]3- in terms of Δ and of the energy of visible light absorbed.
I- < Cl- < F- < OH- < H2O < SCN- < NH3 < en < NO2- < CN- < CO
[Ti(CN)6]3- > [Ti(NH3)6]3+ > [Ti(H2O)6]3+