Chapter 24

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Chapter 24

  1. 1. CHEM 160 General Chemistry II Lecture Presentation Coordination Chemistry Chapter 24
  2. 2. JEWEL
  3. 3. Why Study Descriptive Chemistry of Transition Metals Transition metals are found in nature  Rocks and minerals contain transition metals  The color of many gemstones is due to the presence of transition metal ions  Rubies are red due to Cr  Sapphires are blue due to presence of Fe and Ti  Many biomolecules contain transition metals that are involved in the functions of these biomolecules  Vitamin B12 contains Co  Hemoglobin, myoglobin, and cytochrome C contain Fe
  4. 4. Why Study Descriptive Chemistry of Transition Metals Transition metals and their compounds have many useful applications  Fe is used to make steel and stainless steel  Ti is used to make lightweight alloys  Transition metal compounds are used as pigments  TiO2 = white  PbCrO4 = yellow  Fe4[Fe(CN)6]3 (prussian blue)= blue  Transition metal compounds are used in many industrial processes
  5. 5. Why Study Descriptive Chemistry of Transition Metals To understand the uses and applications of transition metals and their compounds, we need to understand their chemistry. Our focus will be on the 4th period transition elements.
  6. 6. Periodic Tabled block transition elements f block transition elements
  7. 7. Transition Metals General Properties Have typical metallic properties Not as reactive as Grp. IA, IIA metals Have high MP’s, high BP’s, high density, and are hard and strong Have 1 or 2 s electrons in valence shell Differ in # d electrons in n-1 energy level Exhibit multiple oxidation states
  8. 8. d-Block Transition Elements VIIIBIIIB IVB VB VIB VIIB IB IIBSc Ti V Cr Mn Fe Co Ni Cu Zn Y Zr Nb Mo Tc Ru Rh Pd Ag CdLa Hf Ta W Re Os Ir Pt Au Hg Most have partially occupied d subshells in common oxidation states
  9. 9. Electronic ConfigurationsElement Configuration Sc [Ar]3d14s2 Ti [Ar]3d24s2 V [Ar]3d34s2 Cr [Ar]3d54s1 Mn [Ar]3d54s2 [Ar] = 1s22s22p63s23p6
  10. 10. Electronic ConfigurationsElement Configuration Fe [Ar] 3d64s2 Co [Ar] 3d74s2 Ni [Ar] 3d84s2 Cu [Ar]3d104s1 Zn [Ar]3d104s2 [Ar] = 1s22s22p63s23p6
  11. 11. Transition Metals Characteristics due to d electrons: Exhibit multiple oxidation states Compounds typically have color Exhibit interesting magnetic properties  paramagnetism  ferromagnetism
  12. 12. Oxidation States of Transition ElementsSc Ti V Cr Mn Fe Co Ni Cu Zn +1 +1 +2 +2 +2 +2 +2 +2 +2 +2 +2+3 +3 +3 +3 +3 +3 +3 +3 +3 +4 +4 +4 +4 +4 +4 +5 +5 +5 +5 +6 +6 +6 +7
  13. 13. Oxidation States of Transition Elements Sc Ti V Cr Mn Fe Co Ni Cu Zn +1 +1 +2 +2 +2 +2 +2 +2 +2 +2 +2 +3 +3 +3 +3 +3 +3 +3 +3 +3 +4 +4 +4 +4 +4 +4 +5 +5 +5 +5 +6 +6 +6 +7 3/7/01 Ch. 24 11loss of ns e-s loss of ns and (n-1)d e-s
  14. 14. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+
  15. 15. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+
  16. 16. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 valence ns e-’s removed first
  17. 17. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 [Ar]3d6 valence ns e-’s removed first
  18. 18. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+
  19. 19. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+ Fe – 3e- → Fe3+
  20. 20. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+ Fe – 3e- → Fe3+ [Ar]3d64s2 valence ns e-’s removed first, then n-1 d e-’s
  21. 21. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+ Fe – 3e- → Fe3+ [Ar]3d64s2 [Ar]3d5 valence ns e-’s removed first, then n-1 d e-’s
  22. 22. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+
  23. 23. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+ Co – 3e- → Co3+
  24. 24. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+ Co – 3e- → Co3+ [Ar]3d74s2 valence ns e-’s removed first, then n-1 d e-’s
  25. 25. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+ Co – 3e- → Co3+ [Ar]3d74s2 [Ar]3d6 valence ns e-’s removed first, then n-1 d e-’s
  26. 26. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+
  27. 27. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+ Mn – 4e- → Mn4+
  28. 28. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+ Mn – 4e- → Mn4+ [Ar]3d54s2 valence ns e-’s removed first, then n-1 d e-’s
  29. 29. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+ Mn – 4e- → Mn4+ [Ar]3d54s2 [Ar]3d3 valence ns e-’s removed first, then n-1 d e-’s
  30. 30. Coordination Chemistry Transition metals act as Lewis acids  Form complexes/complex ions Fe3+(aq) + 6CN-(aq) → Fe(CN)63-(aq) Lewis acid Lewis base Complex ion Ni2+(aq) + 6NH3(aq) → Ni(NH3)62+(aq) Lewis acid Lewis base Complex ion Complex contains central metal ion bonded to one or more molecules or anions Lewis acid = metal = center of coordination Lewis base = ligand = molecules/ions covalently bonded to metal in complex
  31. 31. Coordination Chemistry Transition metals act as Lewis acids  Form complexes/complex ions Fe3+(aq) + 6CN-(aq) → [Fe(CN)6]3-(aq) Lewis acid Lewis base Complex ion Ni2+(aq) + 6NH3(aq) → [Ni(NH3)6]2+(aq) Lewis acid Lewis base Complex ion Complex with a net charge = complex ion Complexes have distinct properties
  32. 32. Coordination Chemistry Coordination compound Compound that contains 1 or more complexes Example  [Co(NH3)6]Cl3  [Cu(NH3)4][PtCl4]  [Pt(NH3)2Cl2]
  33. 33. Coordination Chemistry Coordination sphere Metal and ligands bound to it Coordination number number of donor atoms bonded to the central metal atom or ion in the complex  Most common = 4, 6  Determined by ligands  Larger ligands and those that transfer substantial negative charge to metal favor lower coordination numbers
  34. 34. Coordination ChemistryComplex charge = sum of charges on the metal and the ligands [Fe(CN)6]3-
  35. 35. Coordination ChemistryComplex charge = sum of charges on the metal and the ligands [Fe(CN)6]3- +3 6(-1)
  36. 36. Coordination Chemistry Neutral charge of coordination compound = sum ofcharges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 neutral compound
  37. 37. Coordination Chemistry Neutral charge of coordination compound = sum ofcharges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 +2 6(0) 2(-1)
  38. 38. Coordination Chemistry Ligands classified according to the number of donor atoms Examples  monodentate = 1  bidentate = 2  tetradentate = 4  hexadentate = 6  polydentate = 2 or more donor atoms
  39. 39. Coordination Chemistry Ligands classified according to the number of donor atoms Examples  monodentate = 1  bidentate = 2 chelating agents  tetradentate = 4  hexadentate = 6  polydentate = 2 or more donor atoms
  40. 40. Ligands Monodentate  Examples:  H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides), CO, O2- Example Complexes  [Co(NH3)6]3+  [Fe(SCN)6]3-
  41. 41. Ligands Bidentate Examples  oxalate ion = C2O42-  ethylenediamine (en) = NH2CH2CH2NH2  ortho-phenanthroline (o-phen) Example Complexes  [Co(en)3]3+  [Cr(C2O4)3]3-  [Fe(NH3)4(o-phen)]3+
  42. 42. Ligands oxalate ion ethylenediamineO O 2- CH2 CH2 C C H2N NH2O O * ** * ortho-phenanthroline CH * N CH * N C CHDonor Atoms HC C C HC C CH CH CH
  43. 43. Ligandsoxalate ion ethylenediamine H C C O M N M
  44. 44. Ligands
  45. 45. Ligands Hexadentate  ethylenediaminetetraacetate (EDTA) = (O2CCH2)2N(CH2)2N(CH2CO2)24- Example Complexes  [Fe(EDTA)]-1  [Co(EDTA)]-1
  46. 46. Ligands O EDTA O*O C CH2 CH2 C O* * N * CH2 CH2 N*O C CH2 CH2 C O* O O Donor Atoms
  47. 47. Ligands EDTA O HC N M
  48. 48. LigandsEDTA
  49. 49. Common Geometries of ComplexesCoordination Number Geometry 2 Linear
  50. 50. Common Geometries of ComplexesCoordination Number Geometry 2 Linear Example: [Ag(NH3)2]+
  51. 51. Common Geometries of ComplexesCoordination Number Geometry 4 tetrahedral (most common) square planar (characteristic of metal ions with 8 d e-’s)
  52. 52. Common Geometries of ComplexesCoordination Number Geometry 4 tetrahedral Examples: [Zn(NH3)4]2+, [FeCl4]- square planar Example: [Ni(CN)4]2-
  53. 53. Common Geometries of ComplexesCoordination Number Geometry 6 octahedral
  54. 54. Common Geometries of ComplexesCoordination Number Geometry 6 Examples: [Co(CN)6]3-, [Fe(en)3]3+ octahedral
  55. 55. Porphine, an importantchelating agent found in nature N NH NH N
  56. 56. Metalloporphyrin N 2+ N Fe N N
  57. 57. Myoglobin, a protein that stores O2 in cells
  58. 58. Coordination Environment of Fe2+ inOxymyoglobin and Oxyhemoglobin
  59. 59. FG24_014.JPGFerrichrome (Involved in Fe transport in bacteria)
  60. 60. Nomenclature of Coordination Compounds: IUPAC Rules The cation is named before the anion When naming a complex: Ligands are named first  alphabetical order Metal atom/ion is named last  oxidation state given in Roman numerals follows in parentheses Use no spaces in complex name
  61. 61. Nomenclature: IUPAC Rules The names of anionic ligands end with the suffix -o -ide suffix changed to -o -ite suffix changed to -ito -ate suffix changed to -ato
  62. 62. Nomenclature: IUPAC Rules Ligand Name bromide, Br- bromo chloride, Cl- chloro cyanide, CN- cyano hydroxide, OH- hydroxo oxide, O2- oxo fluoride, F- fluoro
  63. 63. Nomenclature: IUPAC Rules Ligand Namecarbonate, CO32- carbonato oxalate, C2O42- oxalato sulfate, SO42- sulfatothiocyanate, SCN- thiocyanatothiosulfate, S2O32- thiosulfato Sulfite, SO32- sulfito
  64. 64. Nomenclature: IUPAC Rules Neutral ligands are referred to by the usual name for the molecule Example  ethylenediamine Exceptions  water, H2O = aqua  ammonia, NH3 = ammine  carbon monoxide, CO = carbonyl
  65. 65. Nomenclature: IUPAC Rules Greek prefixes are used to indicate the number of each type of ligand when more than one is present in the complex  di-, 2; tri-, 3; tetra-, 4; penta-, 5; hexa-, 6 If the ligand name already contains a Greek prefix, use alternate prefixes:  bis-, 2; tris-, 3; tetrakis-,4; pentakis-, 5; hexakis-, 6  The name of the ligand is placed in parentheses
  66. 66. Nomenclature: IUPAC Rules If a complex is an anion, its name ends with the -ate appended to name of the metal
  67. 67. Nomenclature: IUPAC RulesTransition Name if in Cationic Name if in Anionic Complex Metal Complex Sc Scandium Scandate Ti titanium titanate V vanadium vanadate Cr chromium chromate Mn manganese manganate Fe iron ferrate Co cobalt cobaltate Ni nickel nickelate Cu Copper cuprate Zn Zinc zincate
  68. 68. Isomerism Isomers compounds that have the same composition but a different arrangement of atoms Major Types structural isomers stereoisomers
  69. 69. Structural Isomers Structural Isomers isomers that have different bonds
  70. 70. Structural Isomers Coordination-sphere isomers differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere
  71. 71. Coordination-Sphere Isomers Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl
  72. 72. Coordination-Sphere Isomers Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl Consider ionization in water [Co(NH3)5Cl]Br → [Co(NH3)5Cl]+ + Br- [Co(NH3)5Br]Cl → [Co(NH3)5Br]+ + Cl-
  73. 73. Coordination-Sphere Isomers Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl Consider precipitation[Co(NH3)5Cl]Br(aq) + AgNO3(aq) → [Co(NH3)5Cl]NO3(aq) + AgBr(s)[Co(NH3)5Br]Cl(aq) + AgNO3(aq) → [Co(NH3)5Br]NO3(aq) + AgCl(aq)
  74. 74. Structural Isomers Linkage isomers differ in the atom of a ligand bonded to the metal in the complex
  75. 75. Linkage Isomers Example [Co(NH3)5(ONO)]2+ vs. [Co(NH3)5(NO2)]2+
  76. 76. Linkage Isomers
  77. 77. Linkage Isomers Example [Co(NH3)5(SCN)]2+ vs. [Co(NH3)5(NCS)]2+  Co-SCN vs. Co-NCS
  78. 78. Stereoisomers Stereoisomers Isomers that have the same bonds, but different spatial arrangements
  79. 79. Stereoisomers Geometric isomers Differ in the spatial arrangements of the ligands
  80. 80. Geometric Isomerscis isomer trans isomer Pt(NH3)2Cl2
  81. 81. Geometric Isomerscis isomer trans isomer [Co(H2O)4Cl2]+
  82. 82. Stereoisomers Geometric isomers Differ in the spatial arrangements of the ligands Have different chemical/physical properties  different colors, melting points, polarities, solubilities, reactivities, etc.
  83. 83. Stereoisomers Optical isomers isomers that are nonsuperimposable mirror images  said to be “chiral” (handed)  referred to as enantiomers A substance is “chiral” if it does not have a “plane of symmetry”
  84. 84. Example 1 mirror plane cis-[Co(en)2Cl2]+
  85. 85. Example 1 rotate mirror image 180° 180 °
  86. 86. Example 1 nonsuperimposable cis-[Co(en)2Cl2]+
  87. 87. Example 1 enantiomers cis-[Co(en)2Cl2]+
  88. 88. Example 2 mirror plane trans-[Co(en)2Cl2]+
  89. 89. Example 2 rotate mirror image 180° 180 ° trans-[Co(en)2Cl2]+
  90. 90. Example 2 Superimposable-not enantiomers trans-[Co(en)2Cl2]+
  91. 91. Properties of Optical Isomers Enantiomers possess many identical properties  solubility, melting point, boiling point, color, chemical reactivity (with nonchiral reagents) different in:  interactions with plane polarized light
  92. 92. Optical Isomers polarizing filter plane polarized light light unpolarizedsource light (random vibrations) (vibrates in one plane)
  93. 93. Optical Isomerspolarizing filter plane polarized optically active sample light in solution rotated polarized light
  94. 94. Optical Isomers polarizing filter plane polarized optically active sample light in solutionDextrorotatory (d) = rightrotationLevorotatory (l) = left rotationRacemic mixture = equal rotated polarizedamounts of two enantiomers; no lightnet rotation
  95. 95. Properties of Optical Isomers Enantiomers possess many identical properties  solubility, melting point, boiling point, color, chemical reactivity (with nonchiral reagents) different in:  interactions with plane polarized light  reactivity with “chiral” reagents Example d-C4H4O62-(aq) + d,l-[Co(en)3]Cl3(aq) → d-[Co(en)3](d-C4H4O62- )Cl(s) + l-[Co(en)3]Cl3(aq) +2Cl-(aq)
  96. 96. Properties of Transition Metal Complexes Properties of transition metal complexes: usually have color  dependent upon ligand(s) and metal ion many are paramagnetic  due to unpaired d electrons  degree of paramagnetism dependent on ligand(s)  [Fe(CN)6]3- has 1 unpaired d electron  [FeF6]3- has 5 unpaired d electrons
  97. 97. Crystal Field Theory Crystal Field Theory Model for bonding in transition metal complexes  Accounts for observed properties of transition metal complexes Focuses on d-orbitals Ligands = point negative charges Assumes ionic bonding  electrostatic interactions
  98. 98. Y Z d orbitals X X Y dx2-y2 Z dz2 ZX X Y dxy dxz dyz
  99. 99. Crystal Field Theory Electrostatic Interactions (+) metal ion attracted to (-) ligands (anion or dipole)  provides stability lone pair e-’s on ligands repulsed by e-’s in metal d orbitals  interaction called crystal field  influences d orbital energies  not all d orbitals influenced the same way
  100. 100. Crystal Field TheoryOctahedral Crystal Field - (-) Ligands attracted to (+) metal ion; provides stability - - + - - d orbital e-’s repulsed by (–) ligands; increases d orbital - potential energy ligands approach along x, y, z axes
  101. 101. Crystal Field Theory Lobes directed at ligandsgreater electrostatic repulsion = higher potential energy
  102. 102. Crystal Field Theory Lobes directed between ligandsless electrostatic repulsion = lower potential energy
  103. 103. Crystal Field Theory octahedral crystal field dz2 dx2- y2 _ _ d orbital energy levels _ _ _E dxy dxz dyz isolated metal ion metal ion in octahedral _____ complex d-orbitals
  104. 104. Crystal Field Splitting Energy dz2 dx2- y2 Determined by metal∆ ion and ligand dxy dxz dyz
  105. 105. metal ion in octahedraloctahedral crystal field complex dz2 dx2- y2d orbital energy levels _ _ ∆ _ _ _E dxy dxz dyz isolated _____ metal ion Metal ion and the nature of the ligand determines ∆ d-orbitals
  106. 106. Properties of Transition Metal Complexes Properties of transition metal complexes: usually have color  dependent upon ligand(s) and metal ion many are paramagnetic  due to unpaired d electrons  degree of paramagnetism dependent on ligand(s)  [Fe(CN)6]3- has 1 unpaired d electron  [FeF6]3- has 5 unpaired d electrons
  107. 107. Crystal Field Theory Crystal Field Theory Can be used to account for  Colors of transition metal complexes  A complex must have partially filled d subshell on metal to exhibit color  A complex with 0 or 10 d e-s is colorless  Magnetic properties of transition metal complexes  Many are paramagnetic  # of unpaired electrons depends on the ligand
  108. 108. Colors of Transition Metal Complexes Compounds/complexes that have color: absorb specific wavelengths of visible light (400 –700 nm)  wavelengths not absorbed are transmitted
  109. 109. Visible Spectrum wavelength, nm (Each wavelength corresponds to a different color)400 nm 700 nm higher lower energy energy White = all the colors (wavelengths)
  110. 110. Visible Spectrum
  111. 111. Colors of Transition Metal Complexes Compounds/complexes that have color: absorb specific wavelengths of visible light (400 –700 nm)  wavelengths not absorbed are transmitted  color observed = complementary color of color absorbed
  112. 112. absorbed observedcolor color
  113. 113. Colors of Transition Metal Complexes Absorption of UV-visible radiation by atom, ion, or molecule: Occurs only if radiation has the energy needed to raise an e- from its ground state to an excited state  i.e., from lower to higher energy orbital  light energy absorbed = energy difference between the ground state and excited state  “electron jumping”
  114. 114. Colors of Transition Metal Complexes green lightwhite red light observedlight absorbed For transition metal Absorption raises ancomplexes, ∆ corresponds to electron from the lower d energies of visible light. subshell to the higher d subshell.
  115. 115. Colors of Transition Metal Complexes Different complexes exhibit different colors because: color of light absorbed depends on ∆  larger ∆ = higher energy light absorbed  Shorter wavelengths  smaller ∆ = lower energy light absorbed  Longer wavelengths magnitude of ∆ depends on:  ligand(s)  metal
  116. 116. Colors of Transition Metal Complexeswhite green light red lightlight observed absorbed (lower energy light) [M(H2O)6]3+
  117. 117. Colors of Transition Metal Complexeswhite blue light orange lightlight absorbed observed (higher energy light) [M(en)3]3+
  118. 118. Colors of Transition Metal ComplexesSpectrochemical SeriesSmallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
  119. 119. Properties of Transition Metal Complexes Properties of transition metal complexes: usually have color  dependent upon ligand(s) and metal ion many are paramagnetic  due to unpaired d electrons  degree of paramagnetism dependent on ligand(s)  [Fe(CN)6]3- has 1 unpaired d electron  [FeF6]3- has 5 unpaired d electrons
  120. 120. Electronic Configurations of Transition Metal Complexes Expected orbital filling tendencies for e -’s: occupy a set of equal energy orbitals one at a time with spins parallel (Hund’s rule)  minimizes repulsions occupy lowest energy vacant orbitals first These are not always followed by transition metal complexes.
  121. 121. Electronic Configurations of Transition Metal Complexes d orbital occupancy depends on ∆ and pairing energy, P e-’s assume the electron configuration with the lowest possible energy cost If ∆ > P (∆ large; strong field ligand)  e-’s pair up in lower energy d subshell first If ∆ < P (∆ small; weak field ligand)  e-’s spread out among all d orbitals before any pair up
  122. 122. d-orbital energy level diagrams octahedral complex d1
  123. 123. d-orbital energy level diagrams octahedral complex d2
  124. 124. d-orbital energy level diagrams octahedral complex d3
  125. 125. d-orbital energy level diagrams octahedral complex d4high spin low spin ∆<P ∆>P
  126. 126. d-orbital energy level diagrams octahedral complex d5high spin low spin ∆<P ∆>P
  127. 127. d-orbital energy level diagrams octahedral complex d6high spin low spin ∆<P ∆>P
  128. 128. d-orbital energy level diagrams octahedral complex d7high spin low spin ∆<P ∆>P
  129. 129. d-orbital energy level diagrams octahedral complex d8
  130. 130. d-orbital energy level diagrams octahedral complex d9
  131. 131. d-orbital energy level diagrams octahedral complex d10
  132. 132. Electronic Configurations of Transition Metal Complexes Determining d-orbital energy level diagrams: determine oxidation # of the metal determine # of d e-’s determine if ligand is weak field or strong field draw energy level diagram
  133. 133. Colors of Transition Metal ComplexesSpectrochemical SeriesSmallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
  134. 134. d-orbital energy level diagrams tetrahedral complex
  135. 135. d-orbital energy level metal ion in tetrahedral complexdiagram dxy dxz dyz _ _ _ ∆ _ _E dz2 dx2- y2 isolated metal ion only high spin _____ d-orbitals
  136. 136. d-orbital energy level diagrams square planar complex
  137. 137. d-orbital energy level metal ion in squarediagram planar complex __ dx2- y2 __ dxy __ dz2E __ __ isolated dxz dyz metal ion _____ only low spin d-orbitals
  138. 138. Myoglobin, a protein that stores O2 in cells
  139. 139. Porphine, an importantchelating agent found in nature N NH NH N
  140. 140. Metalloporphyrin N 2+ N Fe N N
  141. 141. Coordination Environment of Fe2+ inOxymyoglobin and Oxyhemoglobin
  142. 142. Arterial Blood Strong field O2 NN large ∆ FeN N N NH Bright red due to globin absorption of greenish (protein) light
  143. 143. Venous Blood Weak field OH2N N FeN N small ∆ N NH Bluish color due to globin absorption of orangish (protein) light
  144. 144. End of Presentation

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