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Chapter 24

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This document discusses coordination chemistry and transition metals. It begins by explaining why transition metals are important to study, as they are found in nature and have many applications. It then discusses the electronic configurations of transition metals and how they can exist in multiple oxidation states. The focus is on coordination complexes formed when transition metals act as Lewis acids and bond to other ligands. Different types of ligands are described along with common coordination geometries. Rules for naming coordination compounds according to IUPAC nomenclature are also provided.

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CHEM 160 General Chemistry II Lecture Presentation Coordination Chemistry Chapter 24
JEWEL
Chapter 24
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
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
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.
Periodic Table d block transition elements f block transition elements
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
d-Block Transition Elements VIIIB IIIB IVB VB VIB VIIB IB IIB Sc Ti V Cr Mn Fe Co Ni Cu Zn Y Zr Nb Mo Tc Ru Rh Pd Ag Cd La Hf Ta W Re Os Ir Pt Au Hg Most have partially occupied d subshells in common oxidation states
Electronic Configurations Element Configuration Sc [Ar]3d14s2 Ti [Ar]3d24s2 V [Ar]3d34s2 Cr [Ar]3d54s1 Mn [Ar]3d54s2 [Ar] = 1s22s22p63s23p6
Electronic Configurations Element Configuration Fe [Ar] 3d64s2 Co [Ar] 3d74s2 Ni [Ar] 3d84s2 Cu [Ar]3d104s1 Zn [Ar]3d104s2 [Ar] = 1s22s22p63s23p6
Transition Metals  Characteristics due to d electrons: Exhibit multiple oxidation states Compounds typically have color Exhibit interesting magnetic properties  paramagnetism  ferromagnetism
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
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 11 loss of ns e-s loss of ns and (n-1)d e-s
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 valence ns e-’s removed first
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 [Ar]3d6 valence ns e-’s removed first
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+
Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+ Fe – 3e- → Fe3+
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
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
Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+
Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+ Co – 3e- → Co3+
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
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
Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+
Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+ Mn – 4e- → Mn4+
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
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
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
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
Coordination Chemistry  Coordination compound Compound that contains 1 or more complexes Example  [Co(NH3)6]Cl3  [Cu(NH3)4][PtCl4]  [Pt(NH3)2Cl2]
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
Coordination Chemistry Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3-
Coordination Chemistry Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3- +3 6(-1)
Coordination Chemistry Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 neutral compound
Coordination Chemistry Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 +2 6(0) 2(-1)
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
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
Ligands  Monodentate  Examples:  H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides), CO, O2- Example Complexes  [Co(NH3)6]3+  [Fe(SCN)6]3-
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+
Ligands oxalate ion ethylenediamine O O 2- CH2 CH2 C C H2N NH2 O O * * * * ortho-phenanthroline CH * N CH * N C CH Donor Atoms HC C C HC C CH CH CH
Ligands oxalate ion ethylenediamine H C C O M N M
Ligands
Ligands  Hexadentate  ethylenediaminetetraacetate (EDTA) = (O2CCH2)2N(CH2)2N(CH2CO2)24- Example Complexes  [Fe(EDTA)]-1  [Co(EDTA)]-1
Ligands O EDTA O *O C CH2 CH2 C O* * N * CH2 CH2 N *O C CH2 CH2 C O* O O Donor Atoms
Ligands EDTA O H C N M
Ligands EDTA
Common Geometries of Complexes Coordination Number Geometry 2 Linear
Common Geometries of Complexes Coordination Number Geometry 2 Linear Example: [Ag(NH3)2]+
Common Geometries of Complexes Coordination Number Geometry 4 tetrahedral (most common) square planar (characteristic of metal ions with 8 d e-’s)
Common Geometries of Complexes Coordination Number Geometry 4 tetrahedral Examples: [Zn(NH3)4]2+, [FeCl4]- square planar Example: [Ni(CN)4]2-
Common Geometries of Complexes Coordination Number Geometry 6 octahedral
Common Geometries of Complexes Coordination Number Geometry 6 Examples: [Co(CN)6]3-, [Fe(en)3]3+ octahedral
Porphine, an important chelating agent found in nature N NH NH N
Metalloporphyrin N 2+ N Fe N N
Myoglobin, a protein that stores O2 in cells
Coordination Environment of Fe2+ in Oxymyoglobin and Oxyhemoglobin
FG24_014.JPG Ferrichrome (Involved in Fe transport in bacteria)
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
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
Nomenclature: IUPAC Rules Ligand Name bromide, Br- bromo chloride, Cl- chloro cyanide, CN- cyano hydroxide, OH- hydroxo oxide, O2- oxo fluoride, F- fluoro
Nomenclature: IUPAC Rules Ligand Name carbonate, CO32- carbonato oxalate, C2O42- oxalato sulfate, SO42- sulfato thiocyanate, SCN- thiocyanato thiosulfate, S2O32- thiosulfato Sulfite, SO32- sulfito
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
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
Nomenclature: IUPAC Rules  If a complex is an anion, its name ends with the -ate appended to name of the metal
Nomenclature: IUPAC Rules Transition 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
Isomerism  Isomers compounds that have the same composition but a different arrangement of atoms  Major Types structural isomers stereoisomers
Structural Isomers  Structural Isomers isomers that have different bonds
Structural Isomers  Coordination-sphere isomers differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere
Coordination-Sphere Isomers  Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl
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-
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)
Structural Isomers  Linkage isomers differ in the atom of a ligand bonded to the metal in the complex
Linkage Isomers  Example [Co(NH3)5(ONO)]2+ vs. [Co(NH3)5(NO2)]2+
Linkage Isomers
Linkage Isomers  Example [Co(NH3)5(SCN)]2+ vs. [Co(NH3)5(NCS)]2+  Co-SCN vs. Co-NCS
Stereoisomers  Stereoisomers Isomers that have the same bonds, but different spatial arrangements
Stereoisomers  Geometric isomers Differ in the spatial arrangements of the ligands
Geometric Isomers cis isomer trans isomer Pt(NH3)2Cl2
Geometric Isomers cis isomer trans isomer [Co(H2O)4Cl2]+
Stereoisomers  Geometric isomers Differ in the spatial arrangements of the ligands Have different chemical/physical properties  different colors, melting points, polarities, solubilities, reactivities, etc.
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”
Chapter 24
Example 1 mirror plane cis-[Co(en)2Cl2]+
Example 1 rotate mirror image 180° 180 °
Example 1 nonsuperimposable cis-[Co(en)2Cl2]+
Example 1 enantiomers cis-[Co(en)2Cl2]+
Example 2 mirror plane trans-[Co(en)2Cl2]+
Example 2 rotate mirror image 180° 180 ° trans-[Co(en)2Cl2]+
Example 2 Superimposable-not enantiomers trans-[Co(en)2Cl2]+
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
Optical Isomers polarizing filter plane polarized light light unpolarized source light (random vibrations) (vibrates in one plane)
Optical Isomers polarizing filter plane polarized optically active sample light in solution rotated polarized light
Optical Isomers polarizing filter plane polarized optically active sample light in solution Dextrorotatory (d) = right rotation Levorotatory (l) = left rotation Racemic mixture = equal rotated polarized amounts of two enantiomers; no light net rotation
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)
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
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
Y Z d orbitals X X Y dx2-y2 Z dz2 Z X X Y dxy dxz dyz
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
Crystal Field Theory Octahedral 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
Chapter 24
Crystal Field Theory Lobes directed at ligands greater electrostatic repulsion = higher potential energy
Crystal Field Theory Lobes directed between ligands less electrostatic repulsion = lower potential energy
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
Crystal Field Splitting Energy dz2 dx2- y2 Determined by metal ∆ ion and ligand dxy dxz dyz
metal ion in octahedral octahedral crystal field complex dz2 dx2- y2 d orbital energy levels _ _ ∆ _ _ _ E dxy dxz dyz isolated _____ metal ion Metal ion and the nature of the ligand determines ∆ d-orbitals
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
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
Colors of Transition Metal Complexes  Compounds/complexes that have color: absorb specific wavelengths of visible light (400 –700 nm)  wavelengths not absorbed are transmitted
Visible Spectrum wavelength, nm (Each wavelength corresponds to a different color) 400 nm 700 nm higher lower energy energy White = all the colors (wavelengths)
Visible Spectrum
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
absorbed observed color color
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”
Colors of Transition Metal Complexes green light white red light observed light absorbed For transition metal Absorption raises an complexes, ∆ corresponds to electron from the lower d energies of visible light. subshell to the higher d subshell.
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
Colors of Transition Metal Complexes white green light red light light observed absorbed (lower energy light) [M(H2O)6]3+
Colors of Transition Metal Complexes white blue light orange light light absorbed observed (higher energy light) [M(en)3]3+
Chapter 24
Colors of Transition Metal Complexes Spectrochemical Series Smallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
Chapter 24
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
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.
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
d-orbital energy level diagrams octahedral complex d1
d-orbital energy level diagrams octahedral complex d2
d-orbital energy level diagrams octahedral complex d3
d-orbital energy level diagrams octahedral complex d4 high spin low spin ∆<P ∆>P
d-orbital energy level diagrams octahedral complex d5 high spin low spin ∆<P ∆>P
d-orbital energy level diagrams octahedral complex d6 high spin low spin ∆<P ∆>P
d-orbital energy level diagrams octahedral complex d7 high spin low spin ∆<P ∆>P
d-orbital energy level diagrams octahedral complex d8
d-orbital energy level diagrams octahedral complex d9
d-orbital energy level diagrams octahedral complex d10
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
Colors of Transition Metal Complexes Spectrochemical Series Smallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
d-orbital energy level diagrams tetrahedral complex
d-orbital energy level metal ion in tetrahedral complex diagram dxy dxz dyz _ _ _ ∆ _ _ E dz2 dx2- y2 isolated metal ion only high spin _____ d-orbitals
d-orbital energy level diagrams square planar complex
d-orbital energy level metal ion in square diagram planar complex __ dx2- y2 __ dxy __ dz2 E __ __ isolated dxz dyz metal ion _____ only low spin d-orbitals
Myoglobin, a protein that stores O2 in cells
Porphine, an important chelating agent found in nature N NH NH N
Metalloporphyrin N 2+ N Fe N N
Coordination Environment of Fe2+ in Oxymyoglobin and Oxyhemoglobin
Arterial Blood Strong field O2 N N large ∆ Fe N N N NH Bright red due to globin absorption of greenish (protein) light
Venous Blood Weak field OH2 N N Fe N N small ∆ N NH Bluish color due to globin absorption of orangish (protein) light
Chapter 24
End of Presentation

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Chapter 24

  • 1. CHEM 160 General Chemistry II Lecture Presentation Coordination Chemistry Chapter 24
  • 4. 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
  • 5. 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
  • 6. 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.
  • 7. Periodic Table d block transition elements f block transition elements
  • 8. 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
  • 9. d-Block Transition Elements VIIIB IIIB IVB VB VIB VIIB IB IIB Sc Ti V Cr Mn Fe Co Ni Cu Zn Y Zr Nb Mo Tc Ru Rh Pd Ag Cd La Hf Ta W Re Os Ir Pt Au Hg Most have partially occupied d subshells in common oxidation states
  • 10. Electronic Configurations Element Configuration Sc [Ar]3d14s2 Ti [Ar]3d24s2 V [Ar]3d34s2 Cr [Ar]3d54s1 Mn [Ar]3d54s2 [Ar] = 1s22s22p63s23p6
  • 11. Electronic Configurations Element Configuration Fe [Ar] 3d64s2 Co [Ar] 3d74s2 Ni [Ar] 3d84s2 Cu [Ar]3d104s1 Zn [Ar]3d104s2 [Ar] = 1s22s22p63s23p6
  • 12. Transition Metals  Characteristics due to d electrons: Exhibit multiple oxidation states Compounds typically have color Exhibit interesting magnetic properties  paramagnetism  ferromagnetism
  • 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
  • 14. 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 11 loss of ns e-s loss of ns and (n-1)d e-s
  • 15. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+
  • 16. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+
  • 17. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 valence ns e-’s removed first
  • 18. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe2+ Fe – 2e- → Fe2+ [Ar]3d64s2 [Ar]3d6 valence ns e-’s removed first
  • 19. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+
  • 20. Electronic Configurations of Transition Metal Ions  Electronic configuration of Fe3+ Fe – 3e- → Fe3+
  • 21. 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
  • 22. 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
  • 23. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+
  • 24. Electronic Configurations of Transition Metal Ions  Electronic configuration of Co3+ Co – 3e- → Co3+
  • 25. 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
  • 26. 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
  • 27. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+
  • 28. Electronic Configurations of Transition Metal Ions  Electronic configuration of Mn4+ Mn – 4e- → Mn4+
  • 29. 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
  • 30. 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
  • 31. 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
  • 32. 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
  • 33. Coordination Chemistry  Coordination compound Compound that contains 1 or more complexes Example  [Co(NH3)6]Cl3  [Cu(NH3)4][PtCl4]  [Pt(NH3)2Cl2]
  • 34. 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
  • 35. Coordination Chemistry Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3-
  • 36. Coordination Chemistry Complex charge = sum of charges on the metal and the ligands [Fe(CN)6]3- +3 6(-1)
  • 37. Coordination Chemistry Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 neutral compound
  • 38. Coordination Chemistry Neutral charge of coordination compound = sum of charges on metal, ligands, and counterbalancing ions [Co(NH3)6]Cl2 +2 6(0) 2(-1)
  • 39. 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
  • 40. 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
  • 41. Ligands  Monodentate  Examples:  H2O, CN-, NH3, NO2-, SCN-, OH-, X- (halides), CO, O2- Example Complexes  [Co(NH3)6]3+  [Fe(SCN)6]3-
  • 42. 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+
  • 43. Ligands oxalate ion ethylenediamine O O 2- CH2 CH2 C C H2N NH2 O O * * * * ortho-phenanthroline CH * N CH * N C CH Donor Atoms HC C C HC C CH CH CH
  • 44. Ligands oxalate ion ethylenediamine H C C O M N M
  • 46. Ligands  Hexadentate  ethylenediaminetetraacetate (EDTA) = (O2CCH2)2N(CH2)2N(CH2CO2)24- Example Complexes  [Fe(EDTA)]-1  [Co(EDTA)]-1
  • 47. Ligands O EDTA O *O C CH2 CH2 C O* * N * CH2 CH2 N *O C CH2 CH2 C O* O O Donor Atoms
  • 48. Ligands EDTA O H C N M
  • 50. Common Geometries of Complexes Coordination Number Geometry 2 Linear
  • 51. Common Geometries of Complexes Coordination Number Geometry 2 Linear Example: [Ag(NH3)2]+
  • 52. Common Geometries of Complexes Coordination Number Geometry 4 tetrahedral (most common) square planar (characteristic of metal ions with 8 d e-’s)
  • 53. Common Geometries of Complexes Coordination Number Geometry 4 tetrahedral Examples: [Zn(NH3)4]2+, [FeCl4]- square planar Example: [Ni(CN)4]2-
  • 54. Common Geometries of Complexes Coordination Number Geometry 6 octahedral
  • 55. Common Geometries of Complexes Coordination Number Geometry 6 Examples: [Co(CN)6]3-, [Fe(en)3]3+ octahedral
  • 56. Porphine, an important chelating agent found in nature N NH NH N
  • 57. Metalloporphyrin N 2+ N Fe N N
  • 58. Myoglobin, a protein that stores O2 in cells
  • 59. Coordination Environment of Fe2+ in Oxymyoglobin and Oxyhemoglobin
  • 60. FG24_014.JPG Ferrichrome (Involved in Fe transport in bacteria)
  • 61. 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
  • 62. 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
  • 63. Nomenclature: IUPAC Rules Ligand Name bromide, Br- bromo chloride, Cl- chloro cyanide, CN- cyano hydroxide, OH- hydroxo oxide, O2- oxo fluoride, F- fluoro
  • 64. Nomenclature: IUPAC Rules Ligand Name carbonate, CO32- carbonato oxalate, C2O42- oxalato sulfate, SO42- sulfato thiocyanate, SCN- thiocyanato thiosulfate, S2O32- thiosulfato Sulfite, SO32- sulfito
  • 65. 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
  • 66. 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
  • 67. Nomenclature: IUPAC Rules  If a complex is an anion, its name ends with the -ate appended to name of the metal
  • 68. Nomenclature: IUPAC Rules Transition 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
  • 69. Isomerism  Isomers compounds that have the same composition but a different arrangement of atoms  Major Types structural isomers stereoisomers
  • 70. Structural Isomers  Structural Isomers isomers that have different bonds
  • 71. Structural Isomers  Coordination-sphere isomers differ in a ligand bonded to the metal in the complex, as opposed to being outside the coordination-sphere
  • 72. Coordination-Sphere Isomers  Example [Co(NH3)5Cl]Br vs. [Co(NH3)5Br]Cl
  • 73. 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-
  • 74. 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)
  • 75. Structural Isomers  Linkage isomers differ in the atom of a ligand bonded to the metal in the complex
  • 76. Linkage Isomers  Example [Co(NH3)5(ONO)]2+ vs. [Co(NH3)5(NO2)]2+
  • 78. Linkage Isomers  Example [Co(NH3)5(SCN)]2+ vs. [Co(NH3)5(NCS)]2+  Co-SCN vs. Co-NCS
  • 79. Stereoisomers  Stereoisomers Isomers that have the same bonds, but different spatial arrangements
  • 80. Stereoisomers  Geometric isomers Differ in the spatial arrangements of the ligands
  • 81. Geometric Isomers cis isomer trans isomer Pt(NH3)2Cl2
  • 82. Geometric Isomers cis isomer trans isomer [Co(H2O)4Cl2]+
  • 83. Stereoisomers  Geometric isomers Differ in the spatial arrangements of the ligands Have different chemical/physical properties  different colors, melting points, polarities, solubilities, reactivities, etc.
  • 84. 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”
  • 86. Example 1 mirror plane cis-[Co(en)2Cl2]+
  • 87. Example 1 rotate mirror image 180° 180 °
  • 88. Example 1 nonsuperimposable cis-[Co(en)2Cl2]+
  • 89. Example 1 enantiomers cis-[Co(en)2Cl2]+
  • 90. Example 2 mirror plane trans-[Co(en)2Cl2]+
  • 91. Example 2 rotate mirror image 180° 180 ° trans-[Co(en)2Cl2]+
  • 92. Example 2 Superimposable-not enantiomers trans-[Co(en)2Cl2]+
  • 93. 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
  • 94. Optical Isomers polarizing filter plane polarized light light unpolarized source light (random vibrations) (vibrates in one plane)
  • 95. Optical Isomers polarizing filter plane polarized optically active sample light in solution rotated polarized light
  • 96. Optical Isomers polarizing filter plane polarized optically active sample light in solution Dextrorotatory (d) = right rotation Levorotatory (l) = left rotation Racemic mixture = equal rotated polarized amounts of two enantiomers; no light net rotation
  • 97. 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)
  • 98. 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
  • 99. 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
  • 100. Y Z d orbitals X X Y dx2-y2 Z dz2 Z X X Y dxy dxz dyz
  • 101. 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
  • 102. Crystal Field Theory Octahedral 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
  • 104. Crystal Field Theory Lobes directed at ligands greater electrostatic repulsion = higher potential energy
  • 105. Crystal Field Theory Lobes directed between ligands less electrostatic repulsion = lower potential energy
  • 106. 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
  • 107. Crystal Field Splitting Energy dz2 dx2- y2 Determined by metal ∆ ion and ligand dxy dxz dyz
  • 108. metal ion in octahedral octahedral crystal field complex dz2 dx2- y2 d orbital energy levels _ _ ∆ _ _ _ E dxy dxz dyz isolated _____ metal ion Metal ion and the nature of the ligand determines ∆ d-orbitals
  • 109. 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
  • 110. 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
  • 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
  • 112. Visible Spectrum wavelength, nm (Each wavelength corresponds to a different color) 400 nm 700 nm higher lower energy energy White = all the colors (wavelengths)
  • 114. 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
  • 115. absorbed observed color color
  • 116. 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”
  • 117. Colors of Transition Metal Complexes green light white red light observed light absorbed For transition metal Absorption raises an complexes, ∆ corresponds to electron from the lower d energies of visible light. subshell to the higher d subshell.
  • 118. 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
  • 119. Colors of Transition Metal Complexes white green light red light light observed absorbed (lower energy light) [M(H2O)6]3+
  • 120. Colors of Transition Metal Complexes white blue light orange light light absorbed observed (higher energy light) [M(en)3]3+
  • 122. Colors of Transition Metal Complexes Spectrochemical Series Smallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
  • 124. 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
  • 125. 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.
  • 126. 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
  • 127. d-orbital energy level diagrams octahedral complex d1
  • 128. d-orbital energy level diagrams octahedral complex d2
  • 129. d-orbital energy level diagrams octahedral complex d3
  • 130. d-orbital energy level diagrams octahedral complex d4 high spin low spin ∆<P ∆>P
  • 131. d-orbital energy level diagrams octahedral complex d5 high spin low spin ∆<P ∆>P
  • 132. d-orbital energy level diagrams octahedral complex d6 high spin low spin ∆<P ∆>P
  • 133. d-orbital energy level diagrams octahedral complex d7 high spin low spin ∆<P ∆>P
  • 134. d-orbital energy level diagrams octahedral complex d8
  • 135. d-orbital energy level diagrams octahedral complex d9
  • 136. d-orbital energy level diagrams octahedral complex d10
  • 137. 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
  • 138. Colors of Transition Metal Complexes Spectrochemical Series Smallest ∆ Largest ∆ ∆ increases I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- strong field weak field
  • 139. d-orbital energy level diagrams tetrahedral complex
  • 140. d-orbital energy level metal ion in tetrahedral complex diagram dxy dxz dyz _ _ _ ∆ _ _ E dz2 dx2- y2 isolated metal ion only high spin _____ d-orbitals
  • 141. d-orbital energy level diagrams square planar complex
  • 142. d-orbital energy level metal ion in square diagram planar complex __ dx2- y2 __ dxy __ dz2 E __ __ isolated dxz dyz metal ion _____ only low spin d-orbitals
  • 143. Myoglobin, a protein that stores O2 in cells
  • 144. Porphine, an important chelating agent found in nature N NH NH N
  • 145. Metalloporphyrin N 2+ N Fe N N
  • 146. Coordination Environment of Fe2+ in Oxymyoglobin and Oxyhemoglobin
  • 147. Arterial Blood Strong field O2 N N large ∆ Fe N N N NH Bright red due to globin absorption of greenish (protein) light
  • 148. Venous Blood Weak field OH2 N N Fe N N small ∆ N NH Bluish color due to globin absorption of orangish (protein) light