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Chapter 22: Metal Complexes

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  1. Chapter 22: Metal Complexes Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

  2. Transition Metal Complexes Ligands • Neutral molecules or ions that have lone pair of electrons that can be used to form bond to metal • Lewis base = electron pair donor Metal • Lewis Acid = electron pair acceptor • Can accept more than one Ligand (Lewis base) M—L bond • Coordinate covalent bond • Lewis acid base adduct formation

  3. Transition Metal Complexes Coordinate Covalent Bond • Both electrons in shared pair come from same atom Coordination Complexes • Central metal atom surrounded by set of ligands • Complex ion: [Co(NH3)6]3+ , [PtCl4]2 • Coordination compound: Ni(CO)4

  4. Common Ligands 1. Monodentate M :L • 1 donor atom • 1 lone pair • 1 bond to metal Anions F–, Cl–, Br–, I–, O2–, S2–, NO2–, C–, OH–, SCN–, S2O32– Molecules H2O, NH3, CO

  5. Common Ligands 2. Chelate or Polydentate Ligands • Have two or more atoms on one molecule with lone pairs • Each of which can simultaneously form 2 e– bonds to Mn+ • Usually 5 or 6-membered rings with M • Sometimes form 4-membered rings • Must be nonlinear molecules

  6. Bidentate Ligands • Two possible sites of attachment • Two Lewis base sites that can attach to Mn+ • Part of larger class of chelate Ethylenediamine (en) 2,2’-bipyridine (bpy) Oxalate (ox2) 1,10-phenanthroline (phen)

  7. Important Polydentate Ligands EDTA4– Porphyrin Corrin ring

  8. Structure of EDTA Ligand and Complex • H4EDTA common polydentate ligand • EDTA4– used to complex metal ions • 6 donor atoms • Wraps around metal ion • Forms very stable complexes

  9. Complex Ions of Nickel (II) • Left = Green • Ni2+ with • Six water ligands • [Ni(H2O)6]2+ • Right = Blue • Ni2+ with • One water and five NH3 ligands • [Ni(NH3)5(H2O)]2+

  10. Formulas of Complex Ions • The symbol for the metal ion is always given first, followed by ligands. • When more than one kind of ligand is present, • Anionic ligands arefirst (in alphabetical order) • Neutral ligands are next(in alphabetical order) • Charge on complex is algebraic sum of charge on metal ion and charges on ligands • The formula is placed inside of square brackets with the charge of the complex as a superscript outside the brackets, if it is not zero.

  11. Ex. Co2+ with 6 H2O and 2 Cl– • Ligands • The six H2O molecules are bound to the Co2+ ion • Counterions • The two Cl– ions are there only to balance the charge • Electrical neutrality

  12. Learning Check 1. What is the formula for the complex made by Cu2+ and four ammonia (NH3) molecules? Decide if the complex could be isolated as a chloride salt or a potassium salt. Write the formula of the appropriate salt. • Cu2+ has +2 charge • NH3 is neutral • So overall charge of ion is +2 + 4(0) = +2 • [Cu(NH3)4]2+ • Need two Cl– to make neutral complex • [Cu(NH3)4]Cl2

  13. Learning Check 2. What is the formula for the complex made by Ag+ and two cyanide (CN–) ions? Decide if the complex could be isolated as a chloride salt or a potassium salt. Write the formula of the appropriate salt. • CN– is negative one charge • So overall charge of ion is +1 + 2(–1) = –1 • [Ag(CN)2]– • Need +1 to make neutral complex • So add one K+ ion • K[Ag(CN)2]

  14. Your Turn! What is the formula for the complex made by Cr2+ and six ammonia molecules? Add chloride or potassium ions to form the appropriate salt. • K2[Cr(NH3)6] • [Cr(NH3)6] • [Cr(NH3)6]Cl2 • [Cr(NH3)6]Cl • K[Cr(NH3)6]

  15. Chelate Effect • Extra stability that results when chelate ligands bind to metal ion Ni2+(aq) + 6NH3(aq)[Ni(NH3)6]2+(aq)Kform = 2.0 × 108 Ni2+(aq) + 3en(aq)[Ni(en)3]2+(aq)Kform = 1.4 × 1017 • en is bidentate ligand (NH2CH2CH2NH2) • NH3 is monodentate ligand • [Ni(en)3]2+is 2 × 109 times more stable than [Ni(NH3)6]2+

  16. Why this Extra Stability? • Entropy effect • Look at the reverse process where metal loses the ligands. [Ni(NH3)6]2+(aq) + 6H2O Ni2+(aq) + 6NH3(aq) Kinst = 5.0 × 10–9 [Ni(en)3]2+(aq) + 6H2O Ni2+(aq) + 3en(aq)Kinst = 2.4 × 10–18 • With NH3 ligands, same number of particles on each side • With chelate ligand, more molecules as reactants

  17. Metal Complex Nomenclature IUPAC Rules for naming Coordination Compounds • Cation named first, then anion • Names of anionic ligands always end in suffix –o • Ligands whose names end in –ide have suffix changed to –o

  18. Metal Complex Nomenclature Rules for naming Coordination Compounds • Ligands whose names end in –iteor –atebecome –itoand –ato respectively

  19. Some Common Ligands

  20. Metal Complex Nomenclature • Neutral ligands given same name as used for molecule except: • H2O aqua NH3 ammine • When there is more than one of a given ligand, specify number of ligands by prefixes.

  21. Metal Complex Nomenclature • Ordering of ligands • Informula, symbol of metal first followed by ligands Ligands order • Anionic ligands first in alphabetical order • Neutral ligands next also in alphabetical order • In the name of the complex, • Ligands named first in alphabetical order without regard to charge • Metal named last

  22. Nomenclature • If the complex ion has a negative charge, the suffix –ate is added to the name of the metal.

  23. Latin Names of Metals • Sometimes the Latin name of the metal is used. • Oxidation state of the metal is designated by Roman numeral in parentheses

  24. Learning Check: Name the Following: • [Ag(CN)2]– • dicyanoargentate(I) ion • [Zn(OH)4]2– • tetrahydroxozincate(II) ion • [Co(NH3)6]3+ • hexamminecobalt(III) ion • [Mn(en)3]Cl2 • tris(ethylenediamine)manganese(II) chloride

  25. Your Turn! What is the correct name for Ni(Br)(CN)(NH3)2? • nickel(II)cyanobromodiammine • diamminebromocyanonickel(II) • amminebromocyanonickel • bromocyanodiamminenickel(II) • bromocyanoammine nickel(II)

  26. Learning Check Predict the formula from the following names: • tetracyanocuprate(I) ion • [Cu(CN)4]3– • triamminethiocyanoplatinum(III) ion • [PtSCN(NH3)3]2+ • diamminetetraaquacopper(II) ion • [Cu(NH3)2(H2O)4]2+ • potassium hexacyanoferrate(III) • K3[Fe(CN)6]

  27. Your Turn! Predict the formula of the coordination complex triamminedichlorocyanocobalt(III) • [CoCl2(CN)(NH3)3] • [Co(NH3)3Cl2(CN)] • [Cl2(CN)(NH3)3Co] • [(NH3)3Cl2(CN)Co] • [CoCl(CN)(NH3)]

  28. Coordination Number (CN) • Number of bonds formed by metal ions to ligands in complex ions • Varies from 2 to 8 • Depends on: • Size of central atom • Steric interactions of ligands • Electrostatic interactions e.g.[Co(NH3)6]3+ CN = 6 [PtCl4]2 CN = 4 Ni(CO)4 CN = 4 CN 4 and 6 most common

  29. Some Common Coordination Numbers (CN) of Metal Ions

  30. Structures CN = 2 ML2 Linear CN = 6 ML6 Octahedral

  31. Structures CN = 4 ML4 Tetrahedral CN = 4 ML4 Square Planar

  32. Your Turn! What is the coordination number of cobalt in [CoCl2(en)2]+? • 2 • 3 • 4 • 5 • 6

  33. Isomers • Existence of two or more compounds with same chemical formula and different physical properties • Consider CrCl3·6H2O • Can isolate three compounds with this formula • Each with own characteristic color and distinct physical properties • [Cr(H2O)6]Cl3purple • [Cr(H2O)5Cl]Cl2·H2O blue-green • [Cr(H2O)4Cl2]Cl·2H2O green

  34. Coordination Isomers • Results from interchange of anionic ligand in first coordination sphere with anion outside coordination sphere Ex. • [Co(NH3)5Br](SO4) violet • [Co(NH3)5(SO4)]Br red • Easily distinguished by tests for counterion • SO42–(aq) + Ba2+(aq) BaSO4(s) • Br –(aq) + Ag+(aq) AgBr(s)

  35. Stereoisomerism • Difference among isomers that arises from various possible orientations of atoms in space • Same atoms attached, but in different order in space • Two major types • Geometric isomerism • Chirality or handedness

  36. 1. Geometric Isomerism CN = 4 Square planar • trans- and cis- isomers • Occurs only with ML2X2 • trans- isomer • Opposite each other • cis-isomer • Next to each other trans- cis-

  37. Isomerism, Geometric, CN = 6 1. Geometric (Structural) • CN = 6 Octahedral • ML4X2 • trans- and cis- isomers trans- cis-

  38. Isomerism, Geometric, CN = 6 • Octahedral geometry • Also get with twobidentate ligands (symbol:L-L) • M(L-L)2X2 cis- trans-

  39. Your Turn! Which of the following coordination complexes can have cis- and trans- isomers? • [Fe(H2O)6]3+ • [CuCl4]2– • [NiBr(NH3)5]+ • [Mn(NH3)4Cl2] • [Pt(en)2]2+

  40. Chirality • More subtle form of structural isomerism • Differ only in “handedness” • Right glove doesn’t fit left hand • Mirror-image object is different from original object

  41. Superimposable • If you can place mirror image on top of object and get same 3-D one-to-one coincidence • For molecules • Each atom in one molecule with equivalent atom in other molecule Chiral • Object and its mirror image are NOT superimposable Enantiomers • Two non-superimposable isomers e.g. [Co(en)3]2+

  42. Isomerism, Chirality Chiral or Optical Isomers • Most important occurs in octahedral geometry • [M(LL)3]n+ d- isomer l- isomer

  43. Geometric Isomers Not Necessarily Optical Isomers M(LL)2X2 • Two bidentate ligands and two monodentate ligands • cis- isomer has two optical isomers • Chiral with two enantiomers mirror d-cis l- cis

  44. Geometric Isomers Not Necessarily Optical Isomers M(LL)2X2 • Two bidentate ligands and two monodentate ligands • Trans-isomer has no optical isomers • Not chiral mirror trans

  45. Unpolarized and Polarized Light • Light possesses electric and magnetic components that behave like vectors • In unpolarized light, electromagnetic oscillations of photons oriented at random angles perpendicular to axis of light propagation

  46. How to Determine Chirality • Place solution in polarimeter and pass plane polarized light through it • Enantiomers rotate plane-polarized light in opposite directions

  47. Your Turn! Which two structures below represent a pair of optical isomers? A. D. B. E. C. B and C

  48. Crystal Field Theory • Localized electron model doesn’t work • No information about how energies of d orbitals are affected by ligands when they form • Transition metal complexes are usually colored • Different ligands often give different colors

  49. Crystal Field Theory 2. Magnetic properties of transition metal complexes often affected by what ligands are attached to metal • Because transition metals have incomplete d subshells, complexes often paramagnetic • But for given metal, number of unpaired spins varies e.g. [Fe(H2O)6]2+ four of its six 3d electrons are unpaired; [Fe(CN)6]4– has no unpaired spins • Any theory that attempts to explain bonding in transition metal complexes must account for color and magnetic properties

  50. Crystal Field Theory • Crystal field theory • Simplest model • Purely electrostatic (ionic) model • Ignores covalent bonding interactions with transition metals • Assumes ligand lone pair = point negative charge • Repels electrons in d orbital on transition metals • Allows us to understand and correlate all those properties that arise from presence of partly filled d subshells