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CHEM1612 - Pharmacy Week 8 : Complexes I

CHEM1612 - Pharmacy Week 8 : Complexes I. Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au. Unless otherwise stated, all images in this file have been reproduced from:

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CHEM1612 - Pharmacy Week 8 : Complexes I

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  1. CHEM1612 - Pharmacy Week 8: Complexes I Dr. SiegbertSchmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au

  2. Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille,Chemistry, John Wiley & Sons Australia, Ltd. 2008      ISBN: 9 78047081 0866

  3. Complexes • Blackman Chapter 13 and Sections 10.4, 11.8 • Biologically important metal-complexes • Complex ions • Kstab • Coordination compounds • Chelates • Geometry of complexes • Solubility and complexes • Nomenclature • Isomerism in complexes Co(EDTA)-

  4. [M(H2O)4]2+ • M2+(aq) (Hydrated M2+ ion) adduct M2+ H2O(l) Metal Ions as Lewis Acids • Whenever a metal ion enters water, a complex ion forms with water as the ligand. • Metal ions act as Lewis acid (accepts electron pair). • Water is the Lewis base (donates electron pair).

  5. Complex Ions • Definition: A central metal ion covalently bound to two or more anions or molecules, called ligands. • Neutral ligands e.g.: water, CO, NH3 • Ionic ligands e.g.: OH-, Cl-, CN- [Ni(H2O)6]2+, a typical complex ion. • Ni2+ is the central metal ion • Six H2O molecules are the ligands • overall 2+ charge. Blackman Figure 13.12

  6. Coordination Compounds • They consist of: • Complex ion (metal ion with attached ligands) • Counter ions (additional anions/cations needed for zero net charge) • Eg. [Co(NH3)6]Cl3 (s) [Co(NH3)6]3+(aq) + 3 Cl-(aq) Complex ion Counter ions In water coordination compounds behave like electrolytes: the complex ion exists as the cation and the 3 Cl- ions are separate. Note: the counter ion may also be a complex ion. e.g. [Co(H2O)6][CoCl4]3(s) [Co(H2O)6]3+(aq) + 3 [CoCl4]-(aq)

  7. Coordination compounds Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Ligands within the coordination sphere remain bound to the metal ion Coordination Compound Complex Ion Counter Ions

  8. Complex Ions e.g. Ag+(aq) + 2 NH3 Ag(NH3)2+(aq) • Ligands must have a lone pair to donate to the metal. • The ‘donation’ of the electron pair is sometimes referred to as a “dative” bond.

  9. Acidic solution Acidity of Aqueous Transition Metal Ions A small and multiply-charged metal ion acts as an acid in water, i.e. the hydrated metal ion transfers an H+ ion to water. 5 bound H2O molecules 1 bound OH- (overall charge reduced by 1) 6 bound H2O molecules Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

  10. ACID STRENGTH Metal Ion Hydrolysis Each hydrated metal ion that transfers a proton to water has a characteristic Ka value.

  11. Coordination number • The number of ligand atoms attached to the metal ion is called the coordination number. • varies from 2 to 8 and depends on the size, charge, and electron configuration of the metal ion. • Typical coordination numbers for some metal ions are: M+ Coord no. M2+Coord no. M3+Coord no. Cu+ 2,4 Mn2+ 4,6 Sc3+ 6 Ag+ 2 Fe2+ 6 Cr3+ 6 Au+ 2,4 Co2+ 4,6 Co3+ 6 Ni2+ 4,6 Au3+ 4 Cu2+ 4,6 Zn2+ 4,6

  12. Coordination Number and Geometry Remember Valence Shell Electron Pair Repulsion Theory (VSEPR)? : : : F : : : : : F : F S : F : : F : : : : F : : Blackman Chapter 5

  13. Coordination Number and Geometry

  14. Ligands • Ligands that can form 1 bond with the metal ion are called monodentate (denta – tooth) e.g. H2O, NH3, Cl- (a single donor atom). • Some ligands have more than one atom with lone pairs that can be bonded to the metal ion – these are called CHELATES (greek: claw) • Bidentate ligands can form 2 bonds e.g. ethylenediamine • Polydentate ligands – can form more than 2 bonds e.g. EDTA - (hexadentate, can form 6 bonds)

  15. Bidentate chelate ligands MX+(en) Ethylenediamine (en) has two N atoms that can form a bond with the metal ion, giving a five-memberedring. Blackman, Bottle, Schmid, Mocerino & Wille, Figure 13.10

  16. Hexadentate ligand: EDTA Ethylenediaminetetraacetate tetraanion (EDTA4-) EDTA forms very stable complexes with many metal ions. EDTA is used for treating heavy-metal poisoning, because it removes lead and other heavy metal ions from the blood and other bodily fluids. Co(III) N=blue O=red [Co(EDTA)]-

  17. Examples of ligands Table from Silberberg, “Chemistry”, McGraw Hill, 2006.

  18. Examples of ligands The charge of a complex ion is the charge of the metal ion plus the charge of its ligands: e.g. [Ni(H2O)6]2+ charge of complex ion is that of the Ni2+ ion. eg [NiCl4]2- Ni2+ ion coordinated to four chloride (Cl-) ions giving overall (2-) charge. [Fe(en)3]3+ [Fe(EDTA)]- [Fe(H2O)6]3+ monodentate ligands bidentate ligands hexadentate ligands

  19. NH3 3 more steps M(H2O)42+ 3NH3 M(H2O)3(NH3)2+ M(NH3)42+ Lewis bases: water and ammonia The stepwise exchange of NH3 for H2O in M(H2O)42+. Ammonia is a stronger Lewis base than water Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

  20. Equilibrium Constant Kstab Metal Ion + nLigand Complex • The larger Kstab, the more stable the complex, e.g. The complex formation equilibrium is characterised by a stability constant, Kstab(also called formation constant): Ag+(aq) + 2 NH3 Ag(NH3)2+(aq)

  21. Stepwise stability constant • Metal ions gain ligands one at a time. • Each step characterised by “stepwise stability constant” aka “stepwise formation constant”. Overall formation constant = Kstab = K1 x K2…x Kn • Example: Ag+(aq) + NH3(aq) Ag(NH3)+(aq) K1 = 2.1 · 103 Ag(NH3)+(aq) + NH3(aq) Ag(NH3)2+(aq) K2 = 8.2 · 103 Ag+(aq) + 2 NH3(aq) Ag(NH3)2+(aq) Kstab = Kstab = K1 x K2 = [Ag(NH3)2+] = 1.7 · 107 [Ag+] [NH3]2

  22. Demo: Nickel complexes Ni2+ forms three complexes with ethylenediamine: • Mix [Ni(H2O)6]2+ and en in ratio 3:1 → some [Ni(en)(H2O)4]2+and [Ni(H2O)6]2 Green blue-green • Mix [Ni(H2O)6]2+ and en in ratio 1:1 → mostly [Ni(en)(H2O)4]2+ light blue • Mix [Ni(H2O)6]2+ and en in ratio 1:3 → mostly [Ni(en)3]2+purple

  23. Biologically Important Complexes • Many biomolecules contain metal ions that act as Lewis acids. Give some examples of naturally occurring complexes. • Heme • Chlorophyll • Vitamin B12 • Enzyme Carbonic anhydrase

  24. Heme O2 bound to Fe2+ Heme is a square planar complex of Fe2+ and the tetradentate ring ligand porphyrin (bonds to 4 donor N atoms). Present in hemoglobin, which carries oxygen in blood, and myoglobin, which stores oxygen in muscle. Porphyrin ring Myoglobin protein Blackman Figure 13.37

  25. Chlorophyll • Chlorophyll is a photosynthetic pigment, that gives leaves the characteristic green colour. It is a complex of Mg2+ and a porphyrin ring system (four N atoms are the chelae). Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

  26. Vitamin B12 Dorothy Crowfoot Hodgkin The Nobel Prize in Chemistry 1964 Nobelprize.org Image download from Wikipedia

  27. Tetrahedral complex of Zn2+. Catalyses reaction between water and carbon dioxide during respiration. Coordinated to 3 N, fourth site left free to interact with molecule whose reaction is being catalysed (here with water). By withdrawing electron density, makes water acidic to lose proton and OH- attacks partial positive C of CO2 much more vigorously. Cd2+ is toxic because it competes with zinc for this spot. Carbonic anhydrase Figure downloaded from Wikipedia CO2(g) + 2H2O(l) H3O+(aq) + HCO3-(aq)

  28. Exercise 0.01 moles of AgNO3 are added to a 500 mL of a 1.00 M solution of KCN. Then enough water is added to make 1.00 L of solution. Calculate the equilibrium [Ag+] given Kstab [Ag(CN)2]– =1020 M–2. (careful with the direction of the equation represented by Kstab!) Ag+ + 2CN– [Ag(CN)2]– initial /M 0.01 0.500 0 change ~ -0.01 -0.02 0.01 equilibrium /M x 0.480 0.01

  29. Complex Formation and solubility • Metal complex formation can influence the solubility of a compound. e.g. AgCl(s) + 2 NH3 [Ag(NH3)2]+ + Cl- • This occurs in 2 stages: AgCl(s) Ag+ + Cl- (1) Ag+ + 2 NH3 [Ag(NH3)2]+ (2) • Complex formation removes the free Ag+ from solution and so drives the dissolution of AgCl forward.

  30. Complex ion formation affects solubility • Example: AgBr(s) Ag+(aq) + Br-(aq) • Calculate the solubility of AgBr in: a) water b) 1.0 M sodium thiosulfate (Na2S2O3) c) 1.0 M NH3 (Ksp (AgBr)= 5.0·10-13, Kstab ([Ag(S2O3)2]3- )= 4.7·1013; Kstab(Ag(NH3)2+)=1.7·107) a) Solubility of AgBr in water Ksp = [Ag+][Br-] AgBr(s) Ag+(aq) + Br-(aq) x x Ksp = x2 = 5.0·10-13x = 7.1 ·10-7 M

  31. [Ag(S2O3)23-][Br-] [S2O32-]2 b) Solubility of AgBr in sodium thiosulfate 1.0 M Na2S2O3 (1) AgBr(s) Ag+(aq) + Br-(aq) Koverall = Ksp x Kstab = = 5.0·10-13 x 4.7·1013 = 24 Ag+(aq) + 2S2O32-(aq) [Ag(S2O3)2]3-(aq) AgBr(s) + 2S2O32-(aq) [Ag(S2O3)2]3-(aq) + Br-(aq) (2) (1)+(2) 0 +x x 1.0 M -2x 1.0 -2x Initial Conc. Change Equilibrium Conc. 0 +x x Substitute: Koverall = x2/(1.0 - 2x)2 = 24 x = 0.45 Solubility of AgBr in thiosulfate is 0.45 M (c.f. in water 7.1 x 10-7 M)

  32. [Ag(NH3)2+][Br-] [NH3] c) Solubility of AgBr in ammonia 1.0 M NH3 (1) AgBr(s) Ag+(aq) + Br-(aq) Koverall = Ksp x Kstab = = 5.0·10-13 x 1.7·107 = 8.5·10-6 Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) AgBr(s) + 2NH3(aq) [AgNH3]+(aq) + Br-(aq) (2) (1)+(2) 0 +x x 1.0 M -2x 1.0 - 2x Initial Conc. Change Equilibrium Conc. 0 +x x Substitute: Koverall = x2/(1.0-2x)2 = 8.5·10-6 x = 2.9·10-3 M Solubility of AgBr in NH3 is 2.9·10-3 M (c.f. in thiosulfate 0.45 M)

  33. Start with a AgNO3 aqueous solution. Add sequentially : Ag+ + OH- AgOH(s) (brown) 2 AgOH(s) + HPO42- Ag3PO4(s) (yellow) Ag3PO4(s) + HNO3 3Ag+ + NO3- + HPO42- Ag+ + Cl- AgCl (s) (white) AgCl(s) + 2NH3 [Ag(NH3)2]+ + Cl- [Ag(NH3)2]+ + Br- AgBr (s)(green/white) AgBr(s) + 2S2O32- [Ag(S2O3)2]3- + Br- [Ag(S2O3)2]3- + I- AgI (s) (yellow) AgI(s) + 2CN- [Ag(CN)2]- + I- 2 Ag(CN)2- + S2- Ag2S + CN-(black) The One Pot Reaction + NaOH + Na2HPO4 + HNO3 + NaCl + NH3 + KBr + Na2S2O3 + KI + KCN + Na2S Ksp = 10-7.70 M2 Ksp = 10-16 M3 Ksp = 1.8 x 10-10 M2 Kstab = 1.7 x 107 M-2 Ksp = 5 x 10-13 M2 Kstab = 2.5 x 1013 M-2 Ksp = 8.3 x 10-17 M2 Kstab = 6.3 x 1019 M-2 Ksp = 8 x 10-51 M3

  34. Nomenclature Rules for nomenclature of coordination compounds: • Name cation, then anion, as separate words. Examples: [Pt(NH3)4Cl2](NO2)2 tetraamminedichloridoplatinum(IV) nitrite [Pt(NH3)4(NO2)2]Cl2 tetraamminedinitritoplatinum(IV) chloride • Name the ligands then the metal, all in same word. • Number of ligands as Greek prefixes (di-, tri-, tetra-, penta-, hexa-), except ligands that already have numerical prefixes which use Latin prefixes (bis, tris, tetrakis…) • e.g. bis(ethylenediamine) for (en)2

  35. Fluorido Chlorido Bromido Iodido Hydroxido Cyanido Nomenclature II • Oxidation state in Roman numeral in parentheses after name of metal • e.g. [Ag(NH3)2]NO3 diamminesilver(I) nitrate • Anionic ligands end in '-ido'; • Neutral ligands named as molecule, except those listed here: New IUPAC Nomenclature: all anions ending in – ‘ide’ become -‘ido’. (Please modify accordingly pp.518-519 of your book)

  36. Nomenclature of Ligands • Ligands named in alphabetical order (but prefixes do not affect the order) • e.g. [Co(NH3)5Cl]SO4 pentaamminechloridocobalt(III) sulfate • Anionic complexes end in ‘-ate’ • e.g. K3[CrCl6] potassium hexachloridochromate(III) • Some metals in anionic complexes use Latin -ate names: Not Ironate Not Copperate Not Leadate Not Silverate Not Goldate Not Tinnate

  37. Nomenclature - Exercises • [Co(H2O)6]CO3 hexaaquacobalt(II) carbonate • [Cu(NH3)4]SO4 tetraamminecopper(II) sulfate • (NH4)3[FeF6] ammonium hexafluoridoferrate(III) • K4[Mn(CN)6] potassium hexacyanidomanganate(II)

  38. Assigning oxidation numbers • Example 1: Find O.N. of Co in : [Co(NH3)5Cl]SO4pentaamminechloridocobalt(?) sulfate [Co(NH3)5Cl]2+ ammine is neutral, chloride is -1 O.N. -1 = +2 (sum of O.N.s = overall charge) O.N. = +3 • Example 2: Find O.N. of Mn in :K4[Mn(CN)6] potassium hexacyanidomanganate(?) [Mn(CN)6]4- (CN) is -1 overall O.N. + 6x(-1) = -4 (sum of O.N.s = overall charge) ON = +2

  39. About naming complexes • You won’t be asked to draw formulae of complicated biological complexes. • You should be able to use the naming rules to write formulae from names and names from formulae.

  40. Isomerism in Complexes Complexes can have several types of isomers: • Structural Isomers: different atom connectivities • Coordination sphere isomerism • Linkage isomerism • Stereoisomers: same atom connectivities but different arrangement of atoms in space • Geometric isomerism • Optical isomerism

  41. Coordination compounds Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Ligands within the coordination sphere remain bound to the metal ion Complex Ion Counter Ions Coordination Compound

  42. Coordination Isomers • Ligands and counter-ions exchange place: Example: • [Pt(NH3)4Cl2](NO2)2 tetraamminedichloridoplatinum(IV) nitrite • [Pt(NH3)4(NO2)2]Cl2 tetraamminedinitritoplatinum(IV) chloride • Two sets of ligands are reversed: [Cr(NH3)6][Co(CN)6] NH3 is a ligand for Cr3+ [Co(NH3)6][Cr(CN)6] NH3 is a ligand for Co3+ ligands counterions

  43. cyanato NCO:→ isocyanato OCN:→ cyanate ion Linkage isomers • Occur when a ligand has two alternative donor atoms. • Example 1: Thiocyanato NCS:→ Isothiocyanato SCN:→ Thiocyanate ion Pentaammineisothiocyanatocobalt(III) pentaamminethiocyanatocobalt (III)

  44. N O N O O O [Co(NH3)5(NO2)]Cl2 Pentaamminenitrocobalt(III) chloride [Co(NH3)5(ONO)]Cl2 Pentaamminenitritocobalt(III) chloride Linkage Isomers NO2- nitro O2N:→ nitrito ONO:→ • Example 2: Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

  45. Isomerism in Complexes Complexes can have several types of isomers: • Structural Isomers: different atom connectivities • Coordination sphere isomerism • Linkage isomerism • Stereoisomers: same atom connectivities but different arrangement of atoms in space • Geometric isomerism • Optical isomerism

  46. Stereoisomers: Geometric Isomers Square planar complex. Four coordinate: cis- and trans-[Pt(NH3)2Cl2] Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. cisplatin – highly effective anti-tumour agent No anti-tumour effect

  47. Stereoisomers: Geometric Isomers Octahedral complex. Six coordinate: cis- and trans- [Co(NH3)4Cl2]+ 2 Cl next to each other violet 2 Cl axial to each other green

  48. Stereoisomers: Optical Isomers • When a molecule is non-superimposable with its mirror image. • Example: four different substituents about tetrahedral centre. • Same physical properties, except direction in which they rotate the plane of polarized light. [NiClBrFI]2-

  49. cis-[Co(NH3)4Cl2]+ cis-[Co(en)2Cl2]+ + + Has no optical isomers Has optical isomers Stereoisomers: Optical isomers • Metal atoms with tetrahedral or octahedral geometries (but not square planar) may be chiral due to having different ligands. • For the octahedral case, several chiralities are possible, e.g. • Complex with four ligands of two types.

  50. [M(en)3]n+ complexes have optical isomers: Not superimposable 3+ 3+ Mirror plane Stereoisomers: Optical isomers Having three bidentate ligands of only one type - gives a propeller-type structure. www.pt-boat.com

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