1. Chapter 15 Complex Ions

2. Outline 1. Composition of complex ions 2. Geometry of complex ions
3. Electronic structure of complex ions
4. Formation constants of complex ions

3. Compounds of Transition Metals
Components
Transition metal
Associated ions and molecules
Counter-ion

4. Components of Complex Ions and Compounds of Them
Composition of complex ions
Geometry of complex ions
Electronic structure of central metal atom or ion
Equilibrium constant for the formation of the complex ion

5. Composition of Complex Ions
Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq)
When ammonia is added to a solution of aqueous copper(II) ion, the color changes from pale blue to deep blue
The species that accounts for the color is a complex made up of copper(II) ion and four associated ammonia molecules

6. The Nature of the Complex Ion, Cu(NH3)42+
Each ammonia molecule has a lone pair on the nitrogen
A bond forms between the ammonia and the copper(II) ion
Both electrons come from ammonia
Result is a coordinate covalent bond

7. Terminology The species is referred to as a complex ion
Central metal ion
Small molecules or ions surround it
These are called ligands
The number of atoms bonded to the metal is the coordination number
This may or may not be equal to the number of ligands, as we will see

8. Complex Ion or Compound?
The species, Cu(NH3)42+, is clearly a complex ion; because of the charge, it cannot be a compound
Complex ions can associate with other ions to form compounds:
[Cu(NH3)4]F2
Two fluorides are needed to balance the positive charge on the complex
The square brackets indicate the complex ion, i.e., everything inside is part of the complex ion

9. Cu(NH3)42+ The central metal ion is Cu2+
The ligands are the NH3 molecules
The coordination number is 4

10. Figure 15.1

11. The Nature of the Metal Ion
Complex ions form from
Transition metals
Main-group metals: Al, Sn, Pb
Complex ions exist in aqueous solution
Consider Zn(NO3)2
In water, the ion exists as Zn(H2O)42+

12. Lewis Acid-Base Principles
Lewis bases donate lone pairs
Lewis acids accept lone pairs

13. The Nature of the Complex Ion
Consider that hydrated metal cations are acidic
Brønsted
Zn(H2O)42+ ⇌ Zn(H2O)3(OH)+ (aq) + H+ (aq)
Lewis
The metal can accept a pair of electrons and is therefore a Lewis acid
Consider that the ligand possesses at least one lone pair
This makes the ligand a Lewis base
The complex ion can be described as the product of a Lewis acid-base reaction

14. Lewis and Brønsted Acids and Bases
Lewis bases are also Brønsted bases
Can accept an electron pair (or pairs)
Lewis acids need not be Brønsted acids
The Lewis model broadens the definition of an acid

15. Charges of Complexes
Charge of complex = oxidation number of central metal + charges of ligands
Platinum(II)

16. Table 15.1 – Complexes of Pt2+

17. Example 15.1

18. Chelating Agents Ligands
Any molecule or anion with at least one unshared pair of electrons
Some ligands have more than one pair of unshared electrons
These ligands can coordinate using multiple pairs of electrons
Ligands that can form more than one bond to a central metal ion are called chelates

19. Two Common Chelating Ligands
Oxalate, C2O42-
Ethylenediamine, H2N-CH2CH2-NH2
Both of these ligands coordinate in two places per ligand species
Bidentate chelating ligand

20. Figure 15.2

21. Geometry of Complex Ions
Review of Chapter 7
Recall that the geometry of a molecule can be determined by the way in which the central atom coordinates with terminal atoms
For a complex ion, the geometry is determined by the shape taken by the complex, as determined by the coordination number and nature of the metal ion and ligands

22. Coordination Number The most common coordination number is 6
Geometry is octahedral
Two other coordination numbers are common
2; Geometry is linear
4; Geometry may be square planar or tetrahedral

23. Table 15.2

24. Figure 15.2

25. Example 15.2

26. Coordination Number 2 2-coordinate complexes are linear CuCl2-
Ag(NH3)2+
Au(CN)2-

27. Coordination Number 4 Two geometries exist Tetrahedral Square planar
Zn(NH3)42+
CoCl42-
Square planar
Characteristic of Cu2+ and metal ions with d8 configurations (Pt2+, Ni2+)

28. Figure 15.3

29. Coordination Number 6 6-coordinate complexes are octahedral
The six ligands are equidistant from the central metal
The octahedron can be considered a derivative of a square plane, with two ligands added, one above and one below the plane

30. Example 15.3

31. Isomerism
Two or more species with the same formula, but different chemical and physical properties are called isomers
Complex ions can show several different types of isomerism
Only type to be considered here is geometric isomerism
Only the spatial orientation of ligands differs between geometric isomers

32. Isomerism in Square Planar Complexes
Pt(NH3)2Cl2
Can have two different isomers: cis and trans
Complexes of the form Ma2b2 will show cis-trans isomerism
a and b are different ligands

33. Meaning of cis and trans
Cis positions are 90° apart
Trans positions are 180° apart

34. Figure 15.4

35. Isomerism in Octahedral Complexes
Co(NH3)4Cl2 or Ma4b2

36. Chelates and Isomers
In general, chelating ligands can bridge only cis positions
The bridge is not long enough to stretch across a trans position
Chelates, due to their binding in two locations, generally produce more stable complexes
Partially a result of the geometry (ring size)
Partially as a result of the nature of the bonds

37. Figure 15.6

38. Figure 15.7

39. Physical Properties of Isomers
Note in the last figure that the color of the two complexes differ
The cis isomer is reddish-purple
The trans isomer is green
Geometric isomerism can lead to great differences in the physical and chemical properties of compounds containing complex ions

40. Electronic Structure of Complex Ions
Crystal Field Model explains
Color of transition metal complexes
Magnetic properties of transition metal complexes
Considers the nature of the
Metal
Ligands

41. Transition Metal Cations
In a simple transition metal cation
There are no outer s electrons
Electrons are distributed among the five d orbitals by Hund’s Rule
Recall that Hund’s rule results in the maximum number of unpaired spins
Magnetic properties depend on distribution of electrons
Diamagnetism: no unpaired electrons
Paramagnetism: unpaired electrons

42. Iron(II) ion Iron(II) ion has 26-2 or 24 electrons
Shorthand notation: [Ar]3d6
Orbital diagram:
[Ar] (↑↓) (↑ ) (↑ ) (↑ ) (↑ )
Fe2+ is paramagnetic; there are four unpaired electrons

43. Figure 15.8 – Colors of Transition Metal Compounds

44. Example 15.4

45. d-orbitals Recall that there are five d orbitals
In uncomplexed metal cations, these orbitals have the same energy

46. Figure 15.9

47. Octahedral Complexes We can collect the d orbitals into two groups:
A high energy pair, the dx2-y2 and dz2
A low energy triplet, the dxy, dxz and dyz

48. Why the Split into Two Groups?
In the absence of any ligands, the d orbitals have the same energy
Note that the two orbitals whose energy is considered high lie on the xyz axes
Electron density in metal interacts with the electron density of ligands brought toward the metal on the axes
Note that the two orbitals whose energy is considered lower lie between the xyz axes
There is less interaction between the metal electron density and that of the ligand

49. The Splitting Energy, Δo
The crystal field splitting energy is given the symbol Δo
The magnitude of Δo will determine the way in which the electrons fill the d orbitals in the metal ion of the complex
If Δo is large, electrons will pair in the lower energy orbitals before occupying the higher energy ones
If Δo is small, electrons will distribute themselves in all five orbitals, pairing only in cases where there are more than five electrons

50. Figure 15.10

51. Figure 15.11

52. Electronic Arrangements in Complexes
When Δo is large, a low spin complex results
Electrons fill the lower three orbitals before occupying the upper two
When Δo is small, a high spin complex results
Electrons distribute themselves to all five orbitals by Hund’s rule

53. Notes on High and Low Spin
For a given cation, a high spin complex will always have a larger number of unpaired electrons than a low spin complex
The value of Δo is determined by the nature of the ligand(s)
Strong field-ligands produce low spin complexes
Example: CN-, NH3
Weak field-ligands produce high spin complexes
Example: H2O, Cl-

54. Example 15.5

55. Color Most transition metal complexes are brightly colored
Exception: those with empty d sublevels (e.g., Sc3+; those with full d sublevels (e.g., Zn2+)
The splitting of the d sublevel results in an energy difference that corresponds to the visible region of the electromagnetic spectrum
Visible light is absorbed in the transition of an electron in the ion
Some of the wavelengths of white light are removed
Complex appears colored

56. Figure 15.12

57. Titanium(III) Consider Ti3+ in a complex There is only one d electron
Because there is only one possible electronic transition, it is possible to calculate Δo for this ion
The ion absorbs at 510 nm (green)
The complex appears as the color complement of green (i.e., red-violet or purple)

58. Calculating Δo for Titanium(III)
By knowing the energy of the light absorbed, it is possible to calculate the value of Δo
For ions with more than one d electron, calculating the energy Δo is more difficult

59. Δo and Wavelength
The smaller Δo is, the longer the wavelength of absorption
Small Δo stem from weak field ligands
Large Δo stem from strong field ligands

60. Light Absorption and Color

61. The Spectrochemical Series of Ligands
CN- > NO2- > en > NH3 > NCS- > H2O > F- > Cl-
Strong field weak field

62. Equilibrium and Complex Ions
We can consider the formation of a complex ion in aqueous solution as an equilibrium between
The reactants
Bare metal ion and ligands
The product
The complex ion
As with acid-base and gas phase equilibria, an expression for K can be written

63. Formation Constants Consider Cu2+ (aq) + 4NH3 (aq) ⇌ Cu(NH3)42+ (aq)
We can write the expression for Kf (the formation constant for the ion:
In many cases, Kf is a very large number, indicating that the complex which forms is very stable

64. Table 15.4

65. Example 15.6

66. Example 15.6 (Cont’d)

67. Practical Example
Most cleaning compounds containing ammonia state that the product is not to be used on objects made of silver
Consider the Kf for Ag(NH3)2+: 1.7 X 107
The complex of silver with ammonia is a very stable one
Using an ammonia-based cleaner on silver will partially dissolve the silver

68. Chelates in Nature Chelating agents are abundant in nature
Iron is complexed with the chelate heme
EDTA is used to complex metal ions
Treat heavy metal poisoning
Mask metal ions that promote food spoilage

69. Heme

70. EDTA

71. Key Concepts
1. Relate the composition of a complex ion to its charge, coordination number, and the oxidation number of the central metal
2. Sketch the geometry of a complex ion and identify geometric isomers
3. Give the electron configuration and the orbital diagram for the transition metal at the center of the complex
4. Derive the orbital diagram (high or low spin) for the complex
5. Relate the formation equilibrium constant to the concentrations of metal ion, ligands and complex ion