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Part 2.8: Coordination Chemistry 1. Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory.

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Presentation on theme: "Part 2.8: Coordination Chemistry 1. Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory."— Presentation transcript:

1 Part 2.8: Coordination Chemistry 1

2 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 2

3 Ancient times through Alchemy: – Descriptive chemistry, techniques, minerals (Cu compounds), glasses, glazes, gunpowder 17th Century – Mineral acids (HCl, HNO 3, H 2 SO 4 ), salts and their reactions, acid and bases – Quantitative work became important, molar mass, gases, volumes – 1869: The periodic table Late 1800s: Chemical Industry – Isolate, refine, purify metals and compounds 1896: Discovery of Radioactivity – Atomic structure, quantum mechanics, nuclear chemistry (through early 20th century) History of Inorganic Chemistry 3

4 Inorganic History Side Note Friedrich Wöhler (1828) Potassium CyananteAmmonium Sulfate Ammonium Cyanante Urea “I can no longer, so to speak, hold my chemical water and must tell you that I can make urea without needing a kidney.” Wöhler in a letter to Berzelius

5 20th Century – Coordination chemistry, organometallic chemistry – WWII & Military projects: Manhattan project, jet fuels (boron compounds) 1950s – Crystal field theory, ligand field theory, molecular orbital theory 1955 – Organometallic catalysis of organic reaction (polymerization of ethylene) History of Inorganic Chemistry 5

6 Metal Coordination Complexes Coordination complexes or coordination compounds- consists of a central atom, which is usually metallic, and a surrounding array of bound molecules or ions, that are in turn known as ligands or complexing agents. Stable in light and air. Accidentally discovered while trying to make a red dye (1705). Prussian blue Iron-hexacyanoferrate First synthetic blue dye. Known for centuries. 6

7 Prussian blue Iron-hexacyanoferrate Metal Coordination Complexes The Great Wave off Kanagawa Starry Night Structure of coordination complexes not understood until

8 Metal Coordination Complexes Ligands are ions or neutral molecules that bond to a central metal atom or ion. Denticity refers to the number of donor groups in a single ligand that bind to a central atom in a coordination complex. Ligand biting the metal. M = transition metal L = ligand Monodentate (one tooth) Bidentate (two teeth) Polydentate (many teeth) 8

9 Monodentate Ligands 9

10 Bidentate Ligands 10

11 Polydentate Ligands 11

12 EDTA ethylenediaminetetraacetate Ligands that bind to more than one site are called chelating agents. M = Mn(II), Cu(II), Fe(III), Pb (II) and Co(III) Added to foods to prevent catalytic oxidation In cleaning solutions (reduce water hardness) Chelation therapy for Hg and Pb poisoning Analytical titrations 12

13 Coordination Complex Isomers The same connectivities but different spatial arrangements. Different connectivities (same formula). 13

14 Coordination Isomers Same formula different bonding to the metal. [Cr(NH 3 ) 5 SO 4 ]Br and [Cr(NH 3 ) 5 Br]SO 4 Co + (NH 3 ) 5 + Cl + Br [Co(NH 3 ) 6 ] 3+ and [Cr(CN) 6 ] 3- ) [Cr(NH 3 ) 6 ] 3+ and [Co(CN) 6 ] 3- Cr + (NH 3 ) 5 + SO 4 + Br Co + Cr + (NH 3 ) 6 + (CN) 6 14

15 Linkage Isomers Composition of the complex is the same, but the point of attachment of the ligands differs. FormulaName NO 2 - nitrito (via O) NO 2 - nitro (via N) 15

16 Linkage Isomers The compounds have different properties and colors. Linear vs. bent nitrosyl N or S bond thiocyanate M-NCS M-SCN 16

17 Geometric Isomers In geometric isomers, the ligands have different spatial arrangements about the metal ion. Square planar complexes like [MX 2 Y 2 ]. Example: [Pt(NH 3 ) 2 Cl 2 ]. Octahedral complexes like [MX 4 Y 2 ]. Example: [Pt(NH 3 ) 4 Cl 2 ]. 17

18 Geometric Isomers Octahedral complexes with the formula [MX 3 Y 3 ] can be fac (facial) or mer (meridional). In geometric isomers, the ligands have different spatial arrangements about the metal ion. 18

19 Optical Isomers Optical isomers are compounds with non-superimposable mirror images (chiral molecules). C 1, C n, and D n also T, O, and I Chiral molecules lack an improper axis of rotation (S n ), a center of symmetry (i) or a mirror plane (σ)! Common for octahedral complexes with three bidentate ligands. 19

20 Optical Isomers Can be viewed like a propeller with three blades. 20

21 Optical Isomers Co(en) 2 Cl 2 Not Optically active Optically active 21

22 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 22

23 Organic Bonding Kekule proposes the correct structure of benzene Couper proposed that atoms joined to each other like modern-day Tinkertoys. Oxalic acidEthanol 23

24 Inorganic Complexes Late 1800s- Blomstrand and Jorgenson Co 3+, 4 x NH 3, 3 x Cl Their rules Charge on the metal ion determined the number of bonds - Co 3+ = 3 bonds Similar bonding concepts to organics NH 3 can form chains like -CH 2 - Only Cl - attached to an NH 3 could dissociate Did not explain isomers. 24

25 Inorganic Complexes Werner’s Theory Co 3+, 6 x NH 3, 3 x Cl His rules Metals interact with 6 ligands in octahedral geometry to form “complex ions” -Primary/inner coordination sphere: bound to metal -Secondary/outer coordination sphere: balance charge Blomstrand StructureWerner Structure 25

26 Werner’s Theory – Explains multiple complexes of the same sets of ligands in different numbers  [Co(NH 3 ) 6 ]Cl 3 [Co(NH 3 ) 5 Cl]Cl 2 [Co(NH 3 ) 4 Cl 2 ]Cl [Co(NH 3 ) 3 Cl 3 ]  Different numbers of ions are produced due to outer sphere dissociation – Explains multiple complexes with exact same formula = isomers Werner Complexes 26

27 Werner’s Other Contributions Werner Complexes » Coordination Number = Most first row transition elements prefer 6 ligands. Pt 2+ prefers 4 ligands. » CoA 4 B 2 only has two isomers. Not trigonal prismatic because trigonal antiprimatic because they would give 3 isomers. Octahedral because it only has two possible isomers. » PtA 2 B 2 only has two isomers so it must be square planar. Tetrahedral would have only 1 isomer. » Water completes the Inner Sphere coordination in aqueous solutions: NiCl 2 + H 2 O [Ni(H 2 O) 6 ]Cl 2 27

28 Werner’s Other Contributions Werner Complexes In 1914, Werner resolved hexol, into optical isomers, overthrowing the theory that only carbon compounds could possess chirality. 28

29 Werner Complexes Werner was awarded the Nobel Prize in 1913 (only inorg. up until 1973) 29

30 Coordination Complexes Shortcomings of Werner’s Theory – Does not explain the nature of bonding withing the coordination sphere. – Does not account for the preference between 4- and 6- coordination. – Does not account for square planar vs tetrahedral. Crystal Field TheoryLigand Field Theory 30

31 Crystal Field Theory Electrostatic approach to bonding. First Applied to ionic crystalline substances. Assumptions: 1)Metal ion at the center. 2)Ligands are treated as point charges. 3)Bonding occurs through M+ and L- electrostatic attraction. 4)Bonding is purely ionic. 5)M and L electrons repel each other. 6)d orbital degeneracy is broken as ligands approach. 31

32 Crystal Field Theory 32

33 Octahedral Splitting E M d-orbitals align along the octahedral axis will be affected the most. d z2 d x2-y2 d xy d yz d xz 33

34 d x2-y2 d z2 d xz d xy d yz Tetrahedral SplittingTetrahedral 34

35 Other Geometries 35

36 Other Geometries 36

37 Crystal Field Theory Merits of crystal field theory: 1)Can be used to predict the most favorable geometry for the complex. 2)Can account for why some complexes are tetrahedral and others square planar. 3)Usefull in interpreting magnetic properties. 4)The colors of many transition metal complexes can be rationalized. Limitations of crystal field theory: 1)Becomes less accurate as delocalization increases (more covalent character). 2)Point charge does not accurately represent complexes. 3)Does not account for pi bonding interactions. 4)Does not account for the relative strengths of the ligands. 37

38 Ligand Field Theory Application of molecular orbital theory to transition metal complexes. Ligands are not point charges. Takes into account  bonding. Can be used to explain spectrochemical series. Better than valence-bond model or crystal field theory at explaining experimental data. 38

39 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 39

40 Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry Outline Octahedral  bonding  bonding -Ligand Field Strength Square Planar  bonding  bonding Tetrahedral Organometallics 40

41 Octahedral  Only MOs 1.Assign a point group 2.Choose basis function 3.Apply operations -if the basis stays the same = +1 -if the basis is reversed = -1 -if it is a more complicated change = 0 4.Generate a reducible representation H s orbitals OhOh  Fs in-between H through H-M-H 41

42 Octahedral  Only MOs 1.Assign a point group 2.Choose basis function 3.Apply operations -if the basis stays the same = +1 -if the basis is reversed = -1 -if it is a more complicated change = 0 4.Generate a reducible representation 5.Reduce to irreducible representation 6.Combine orbitals by their symmetry 7.Fill MOs with e - 8.Generate SALCs of peripheral atoms 9.Draw peripheral atoms SALC with central atom orbital to generate bonding/antibonding MOs. H s orbitals OhOh 42

43 Octahedral  Only MOs 5.Reduce to irreducible representation  Hs  Hs : A 1g + T 1u + E g 43

44 Octahedral  Only MOs 5.Irreducible reps for M orbitals s d p 44

45 T 1u EgEg Octahedral  Only MOs 6.Combine the orbital's by their symmetry M 6 x H A 1g T 1u 4s 4p EgEg T 1u A 1g EgEg T 2g 3d E g,T 2g A 1g oo 45

46 Octahedral  Only MOs 6.Combine the orbital's by their symmetry M L EgEg EgEg EgEg T 2g 3d E g,T 2g oo 46

47 Octahedral  Only MOs 6.Combine the orbital's by their symmetry EgEg EgEg EgEg T2gT2g E g,T 2g oo EgEg EgEg EgEg T2gT2g oo Weak  donor Weak Lewis base Weaker bonding interaction Weak Field Smaller  o M L M L Stronger  donor Strong Lewis base Stronger bonding interaction Strong Field Larger  o 47

48 Octahedral  Only MOs 6.Combine the orbital's by their symmetry  o : I - < Br - < Cl - < F - Stronger Lewis base = Larger  o Smaller ligands = Larger  o 48 EgEg EgEg EgEg T2gT2g E g,T 2g oo EgEg EgEg EgEg T2gT2g oo M L M L

49 Octahedral  Only MOs 6.Combine the orbital's by their symmetry M 6 x H AgAg T 1u 4s 4p EgEg T 1u A 1g EgEg EgEg T 2g 3d E g,T 2g T 1u A 1g T 1u oo 49

50 4.Fill MOs with e - 5.Generate SALCs of peripheral atoms 6.Draw peripheral atoms SALC with central atom orbital to generate bonding/antibonding MOs. 50

51 Octahedral  Only MOs AgAg T 1u 4s 4p EgEg T 1u A 1g EgEg EgEg T 2g 3d E g,T 2g T 1u A 1g T 1u  Hs : A 1g + T 1u + E g s obitals p obitals  p : A 1g + T 1u + E g What about p orbitals? M L 51

52 Octahedral  +  MOs 52

53 1.Assign a point group 2.Choose basis function (  bonds) 3.Apply operations -if the basis stays the same = +1 -if the basis is reversed = -1 -if it is a more complicated change = 0 4.Generate a reducible representation 5.Reduce to irreducible representation  orbitals OhOh LL in-between L through L-M-L Octahedral  +  MOs  L  = T 1g + T 2g + T 1u + T 2u 53

54 6.Combine the orbital's by their symmetry M L AgAg T 1u 4s 4p EgEg T 1u A 1g EgEg EgEg T 2g 3d E g,T 2g T 1u A 1g T 1u Octahedral  +  MOs  orbitals  L  = T 1g + T 2g + T 1u + T 2u  orbitals T 2g T 1g T 1u T 2u 54

55 6.Combine the orbital's by their symmetry M L AgAg T 1u 4s 4p EgEg T 1u A 1g EgEg 3d E g,T 2g EgEg T 2g T 1u A 1g T 1u Octahedral  +  MOs  orbitals  orbitals T 2g T 1g T 1u T 2u 55

56 6.Combine the orbital's by their symmetry M-L  EgEg T 2g Octahedral  +  MOs T 2g filled  donor  base donates to M T 2g EgEg EgEg oo oo oo empty  acceptor  acid accepts from M 56

57 Strong Field Weak Field oo oo Ligand Field Strength Filled  donor  base Donates to M Empty  acceptor  acid Accepts from M t 2g egeg egeg Weak  donor Weak Lewis base Weaker bonding interaction Stronger  donor Strong Lewis base Stronger bonding interaction  bonding  bonding 57

58 Pure  donating ligands:    en > NH 3  donating ligands:   : H 2 O > F > RCO 2 > OH > Cl > Br > I  accepting ligands:   : CO, CN -, > phenanthroline > NO 2 - > NCS - Ligand Field Strength oo oo t 2g egeg egeg The Spectrochemical Series CO, CN - > phen > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > RCO 2 - > OH - > Cl - > Br - > I - Note:   increases with increasing formal charge on the metal ion   increases on going down the periodic table (larger metal) 58

59 Ligand Field Strength oo oo t 2g egeg egeg The Spectrochemical Series CO, CN - > phen > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > RCO 2 - > OH - > Cl - > Br - > I - Larger   Smaller   Why do we care? Predict/Tune/Understand the: 1.Photophysical properties of metal coordination complexes. 2.Magnetic properties of metal coordination complexes. 3.And others. 59

60 Increasing  The Spectrochemical Series CO, CN - > phen > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > RCO 2 - > OH - > Cl - > Br - > I - Larger   Smaller   Photophysical Properties 60

61 d1d1 d2d2 d3d3 d4d4 Strong fieldWeak fieldStrong fieldWeak field Magnetic Properties 61

62 Pairing Energy,  The pairing energy, , is made up of two parts. 1)Coulombic repulsion energy caused by having two electrons in same orbital. Destabilizing energy contribution of  c for each doubly occupied orbital. High Energy 2)Exchange stabilizing energy for each pair of electrons having the same spin and same energy. Stabilizing contribution of  e for each pair having same spin and same energy. Medium Energy Hund's Rules Less repulsion Less p + screening Low EnergyMedium Energy 62

63 Side note: Exchange Energy,  e Excitation Internal Conversion Fluorescence Non-radiative decay Intersystem Crossing Phosphorescence S0S0 S1S1 S2S2 E T1T1 Ground State (S 0 ) Singlet Excited State (S 1 ) Triplet Excited State (T 1 )  E ST ≈  e ≈ 2J e 63

64 Pairing Energy,  The pairing energy, , is made up of two parts. 1)Coulombic repulsion energy caused by having two electrons in same orbital. Destabilizing energy contribution of  c for each doubly occupied orbital. High Energy 2)Exchange stabilizing energy for each pair of electrons having the same spin and same energy. Stabilizing contribution of  e for each pair having same spin and same energy. Medium Energy Hund's Rules  = sum of all  c and  e interactions Less repulsion Less p + screening Low EnergyMedium Energy High Energy Low Energy 64

65 d4d4 Strong field = Low spin (2 unpaired) Weak field = High spin (4 unpaired)  <  o  >  o When the 4 th electron will either go into the higher energy e g orbital at an energy cost of  0 or be paired at an energy cost of , the pairing energy.  vs.  o oo oo 65

66 1 u.e.5 u.e. d5d5 0 u.e.4 u.e. d6d6 1 u.e.3 u.e. d7d7 2 u.e. d8d8 1 u.e. d9d9 0 u.e. d 10 Magnetic Properties 66

67 Magnetic Properties The Spectrochemical Series CO, CN - > phen > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > RCO 2 - > OH - > Cl - > Br - > I - Larger   Smaller   High Spin Low Spin Diamagnetic- all electrons paired. Paramagnetic- unpaired electrons. 67

68 Pure  donating ligands:    en > NH 3  donating ligands:   : H 2 O > F > RCO 2 > OH > Cl > Br > I  accepting ligands:   : CO, CN -, > phenanthroline > NO 2 - > NCS - Ligand Field Strength oo oo t 2g egeg egeg 68 The Spectrochemical Series CO, CN - > phen > NO 2 - > en > NH 3 > NCS - > H 2 O > F - > RCO 2 - > OH - > Cl - > Br - > I - Larger   Smaller  

69 Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry Outline Octahedral  bonding  bonding -Ligand Field Strength Square Planar  bonding  bonding Tetrahedral Organometallics 69

70 Square Planar 70

71 Square Planar MOs p orbitals of L D 4h Use a local coordinate system on each ligand with: y pointing in towards the metal. (p y =  bonding) z being perpendicular to the molecular plane. (p z =   bonding) x lying in the molecular plane. (p x =  || bonding) 1.Assign a point group 2.Choose basis function (p orbitals of L)  orbitals (p y )  orbitals (p x,z ) 71

72 Square Planar MOs p orbitals of L D 4h  orbitals (p y ) 1.Assign a point group 2.Choose basis function (p orbitals of L) 3.Apply operations -if the basis stays the same = +1 -if the basis is reversed = -1 -if it is a more complicated change = 0   (py) : A 1g + B 1g + E u 72

73 Square Planar MOs p orbitals of L D 4h  orbitals (p y ) 1.Assign a point group 2.Choose basis function (orbitals) 3.Apply operations 4.Generate a reducible representation 5.Reduce to irreducible representation 6.Combine orbitals by their symmetry   (py) : A 1g + B 1g + E u 73

74 74 Square Planar MOs 5.Irreducible reps for M orbitals s d p

75 Square Planar MOs 75

76  Bonding in Square Planar MOs  orbitals of L D 4h  orbitals (p x,z ) 76

77  Bonding in Square Planar MOs 77

78 Complete Square Planar MOs 78

79 Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry Outline Octahedral  bonding  bonding -Ligand Field Strength Square Planar  bonding  bonding Tetrahedral Organometallics 79

80  Only T d MOs 80    : A 1 + T 2

81  Only T d SALC 11 22 33 44 81   : A 1 + T 2

82  Only T d MOs 82

83 Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field Theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry Outline Octahedral  bonding  bonding -Ligand Field Strength Square Planar  bonding  bonding Tetrahedral Organometallics 83

84 Organometallic Chemistry Organometallic compound- a complex with direct metal-carbon bonds. Zeise’s salt- the first organometallic compound Isolated in 1825 (by William Zeise) Structure confirmed in

85  -bonding Ligands 85

86 History of Ferrocene Pauson and Kealy (1951 ) orange solid of "remarkable stability" FeCl 3 + Fulvalene Nature 1951, 168, G. Wilkinson, M. Rosenblum, M. C. Whiting, R. B. Woodward Journal of the American Chemical Society 1952, 74, 2125– Wilkinson and Fischer (1952) E. O. Fischer, W. Pfab Zeitschrift für Naturforschung B 1952, 7, 377–

87 The first sandwich complex. Fuel additives-anitknocking agents. Electrochemical standard. Some derivatives show anti-cancer activity. Small rotation barrier (~ 4 kJmol ‐1 ) and ground state structures of ferrocene can be D 5d or D 5h. Ferrocene D 5d D 5h D5D5 What about the bonding? 87

88  MOs of Cyclopentadienyl C5H5-C5H5- D 5h Decomposition/Reduction Formula 88

89  MOs of Cyclopentadienyl Generate SALC Energy increases as the # of nodes increases. 89

90  MOs of Ferrocene C5H5-C5H5- D 5h Fe(C 5 H 5 ) 2 D 5d 90

91 Decomposition/Reduction Formula  MOs of Ferrocene Fe(C 5 H 5 ) 2 D 5d 91

92 Generate SALC  MOs of Ferrocene From the equation Assemble 2 x C 5 H

93  MOs of Ferrocene 2 x 93

94  MOs of Ferrocene 94

95  MOs of Ferrocene D 5h D 5d A2”A2” E 1 ” E2”E2” E2”E2” E 2g E 2u E 2g E 2u E 1g E 1u E 1g E 1u A 2u A 1g 95

96  MOs of Ferrocene 96

97  MOs of Ferrocene 97

98  MOs of Ferrocene 98

99  MOs of Ferrocene 99

100  MOs of Ferrocene 100

101 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 101

102 Jahn-Teller Distortion Jahn-Teller theorem: “there cannot be unequal occupation of orbitals with identical energy” Molecules will distort to eliminate the degeneracy! E Distortion d3d3 1 u.e. d9d9 equal occupation unequal occupation 102

103 egeg t 2g E d xz dx2-y2dx2-y2 d yz d xy dz2dz2 Jahn-Teller Distortion [Cu(H 2 O) 6 ] Å 2.00 Å 103

104 Jahn-Teller Distortion 104

105 Jahn-Teller Distortion 105

106 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 106

107 Transition Metals in Biochemistry 107

108 Transport/storage proteins : Transferrin (Fe) Ferritin (Fe) Metallothionein (Zn) O 2 binding/transport:Myoglobin (Fe) Hemoglobin (Fe) Hemerythrin (Fe) Hemocyanin (Cu) Enzymes (catalysts) Hydrolases:Carbonic anhydrase (Zn) Carboxypeptidase (Zn) Oxido-Reductases: Alcohol dehydrogenase (Zn) Superoxide dismutase (Cu, Zn, Mn, Ni) Catalase, Peroxidase (Fe) Nitrogenase (Fe, Mo) Cytochrome oxidase (Fe, Cu) Hydrogenase (Fe, Ni) Isomerases: B 12 coenzymes (Co) Aconitase (Fe-S) Oxygenases: Cytochrome P450 (Fe) Nitric Oxide Synthases (Fe) Electron carriers:Cytochromes (Fe) Iron-sulfur (Fe) Blue copper proteins (Cu) Metals in Biochemistry Structural Skeletal roles via biomineralization Ca 2+, Mg 2+, P, O, C, Si, S, F as anions, e.g. PO 4 3 , CO 3 2 . Charge neutralization. Mg 2+, Ca 2+ to offset charge on DNA - phosphate anions Charge carriers: Na+, K +, Ca 2+ Transmembrane concentration gradients ("ion-pumps and channels") Trigger mechanisms in muscle contraction (Ca). Electrical impulses in nerves (Na, K) Heart rhythm (K). Hydrolytic Catalysts: Zn 2+, Mg 2+ Lewis acid/Lewis base Catalytic roles. Small labile metals. Redox Catalysts: Fe(II)/Fe(III)/Fe(IV), Cu(I)/Cu(II), Mn(II)/Mn(III)/(Mn(IV), Mo(IV)/Mo(V)/Mo(VI), Co(I)/Co(II)/Co(III) Transition metals with multiple oxidation states facilitate electron transfer - energy transfer. Biological ligands can stabilize metals in unusual oxidation states and fine tune redox potentials. Activators of small molecules. Transport and storage of O 2 (Fe, Cu) Fixation of nitrogen (Mo, Fe, V) Reduction of CO 2 (Ni, Fe) Organometallic Transformations. Cobalamins, B 12 coenzymes (Co), Aconitase (Fe-S) 108

109 Transition Metals in Biochemistry 109

110 Amino acid binding functionalities: -OH, -SH, -COOH, -NH, CONH 2 Biological Ligands 110

111 Biological Ligands 111

112 Bioinorganic Chemistry 112

113 Bioinorganic Examples Hemoglobin iron-containing oxygen-transport metalloprotein in the red blood cells of all vertebrates. hemoglobin in the blood carries oxygen from the respiratory organs (lungs or gills) to the rest of the body. 113

114 Bioinorganic Examples Nitrogenase Reduction of N 2 to 2NH 3 + H 2 Fe 7 MoS 9 cluster Mechanism not fully known. Mo sometimes replaced by V or Fe. Inhibited by CO. 114

115 Bioinorganic Examples Iron Sulfur Clusters Mediate electron transport. “Biological capacitors” Fe(II) and Fe(III) Found in a variety of metalloproteins, such as the ferredoxins, hydrogenases, nitrogenase, cytochrome c reductase and others. Ferredoxin 115

116 Metal Ions and Life 116

117 Not Enough Metal Ions 117

118 Argyria or argyrosis: a condition caused by inappropriate exposure to chemical compounds of the element silver. Excess Metal Ions Colloidal Silver Paul Karason- Used silver to “treat” dermatitis, acid reflux and other issues. Food and Drug Administration (FDA) doesn't approve of colloidal silver as a medical treatment! 118

119 To Much Ag 119

120 Outline Coordination Complexes – History – Ligands – Isomers Inorganic Bonding Crystal Field Theory Ligand Field theory Orbital Diagrams Ligand Field Jahn-Teller Distortion Bioinorganic Chemistry 120


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