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William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 15 Complex Ions.

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Presentation on theme: "William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut Chapter 15 Complex Ions."— Presentation transcript:

1 William L Masterton Cecile N. Hurley Edward J. Neth University of Connecticut 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 Cu 2+ (aq) + 4NH 3 (aq) Cu(NH 3 ) 4 2+ (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(NH 3 ) 4 2+ 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(NH 3 ) 4 2+, 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(NH 3 ) 4 ]F 2 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(NH 3 ) 4 2+ The central metal ion is Cu 2+ The ligands are the NH 3 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(NO 3 ) 2 In water, the ion exists as Zn(H 2 O) 4 2+

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(H 2 O) 4 2+ Zn(H 2 O) 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 Pt 2+

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, C 2 O 4 2- Ethylenediamine, H 2 N-CH 2 CH 2 -NH 2 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 CuCl 2 - Ag(NH 3 ) 2 + Au(CN) 2 -

27 Coordination Number 4 Two geometries exist Tetrahedral Zn(NH 3 ) 4 2+ CoCl 4 2- Square planar Characteristic of Cu 2+ and metal ions with d8 configurations (Pt 2+, Ni 2+ )

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(NH 3 ) 2 Cl 2 Can have two different isomers: cis and trans Complexes of the form Ma 2 b 2 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(NH 3 ) 4 Cl 2 or Ma 4 b 2

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 Hunds Rule Recall that Hunds 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]3d 6 Orbital diagram: [Ar] () ( ) ( ) ( ) ( ) Fe 2+ 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 d x 2 -y 2 and d z 2 A low energy triplet, the d xy, d xz and d yz

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, Δ o The crystal field splitting energy is given the symbol Δ o The magnitude of Δ o will determine the way in which the electrons fill the d orbitals in the metal ion of the complex If Δ o is large, electrons will pair in the lower energy orbitals before occupying the higher energy ones If Δ o 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 Δ o is large, a low spin complex results Electrons fill the lower three orbitals before occupying the upper two When Δ o is small, a high spin complex results Electrons distribute themselves to all five orbitals by Hunds 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 Δ o is determined by the nature of the ligand(s) Strong field-ligands produce low spin complexes Example: CN -, NH 3 Weak field-ligands produce high spin complexes Example: H 2 O, Cl -

54 Example 15.5

55 Color Most transition metal complexes are brightly colored Exception: those with empty d sublevels (e.g., Sc 3+ ; those with full d sublevels (e.g., Zn 2+ ) 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 Ti 3+ in a complex There is only one d electron Because there is only one possible electronic transition, it is possible to calculate Δ o 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 Δ o for Titanium(III) By knowing the energy of the light absorbed, it is possible to calculate the value of Δ o For ions with more than one d electron, calculating the energy Δ o is more difficult

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

60 Light Absorption and Color

61 The Spectrochemical Series of Ligands CN - > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > Cl - Strong fieldweak 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 Cu 2+ (aq) + 4NH 3 (aq) Cu(NH 3 ) 4 2+ (aq) We can write the expression for K f (the formation constant for the ion: In many cases, K f 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 (Contd)

67 Practical Example Most cleaning compounds containing ammonia state that the product is not to be used on objects made of silver Consider the K f for Ag(NH 3 ) 2 + : 1.7 X 10 7 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


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

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