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Ch 19. d-Block Metals

Ch 19. d-Block Metals. D H vap (in kJ/mol) for Metals. T m Ba 725°C W 3410°C Au 1064°C. T m across TM. Quick Review of Redox Rxns. To balance a half-reaction: 1. Identify and balance redox atoms 2. Add e  as needed 3. Add H + or OH - to balance charge

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Ch 19. d-Block Metals

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  1. Ch 19. d-Block Metals

  2. DHvap (in kJ/mol) for Metals Tm Ba 725°C W 3410°C Au 1064°C

  3. Tm across TM

  4. Quick Review of Redox Rxns To balance a half-reaction: 1. Identify and balance redox atoms 2. Add eas needed 3. Add H+ or OH- to balance charge 4. Add H2O as needed Ex: Balance HMnO4  Mn2+ in acidic soln 5e + HMnO4 + 7H+  Mn2+ + 4 H2O Balance VO43  V2O3 in basic solution 4e + 2VO43 + 5H2O  V2O3 + 10 OHE = 1.37V at pH = 14

  5. Quick Review of Redox Rxns Nernst relation E = E - (0.059V / n) log Q What is E (VO43 / V2O3) at pH = 12 ? E = E - (0.059V / 4) log [OH]10 = E + (10) (0.059V / 4) ( pOH) = +1.37 V + (0.148) (2) = +1.66 V (E increases with decr pH because OH is produced)

  6. Latimer diagrams • Reverse direction, reverse sign • n E are additive, not E • Mn3+  Mn2+ Mn • E = (1) (1.5V) + 2(1.18V) / 3 = 0.28V • 3. E is independent of stoichiometry -1.18 1.5 Quick Review of Redox Rxns

  7. Quick Review of Redox Rxns e- + Fe3+ → Fe2+ E = 0.77 V e- + Fe(OH)3 + 3H+ → Fe2+ + 3 H2O E = E0 - 3(0.059) pH e- + Fe(OH)3→ Fe(OH)2 + OH- E = E0 - 0.059 pH

  8. TM redox trends TM Frost diagramsat pH=0 Electronegativity increases for TM going across the rows, therefore elements become more difficult to oxidize. A different way of stating this is that later TM elements are stronger oxidants at a given oxidation state. This is shown by the increasing upward slope for oxidation reactions in Frost diagrams.

  9. TM Pourbaix diagrams Pourbaix diagrams show increasing E for M/M2+ and M2+/M3+ equilibria

  10. Early vs late TMs 2 e + CoO2 Co2+ E = 1.66V 2 e + TiO2+  Ti 2+ E = - 0.14V Note that CoO2 is unstable in H2O because: 2 e + 4 H+ + CoO2 Co2+ + 2 H2OE = 1.66 2 H2O  O2 + 4 e + 4 H+ E = -1.23 2 CoO2 + 4 H+ 2 Co2+ + O2 + 2 H2O E = +0.43

  11. TM redox trends More valence e- going across the rows means higher oxidation states are possible, but later TM are too electronegative to be oxidized to their group number. 3 4 5 6 7 8 9 10 11 12 Sc Ti V Cr Mn Fe Co Ni Cu Zn +3 +4 +5 +6 +7 +6 +4 +2,3 +2,3 +2 Highest oxidation states accessible in aqueous solution

  12. TM redox trends Within a triad, 2nd and 3rd row TM are usually similar. Example: Group 6 = Cr, Mo, W triad Cr3+ is v. stable, unlike Mo3+ and W3+ Cr6+ is a strong oxidizer, unlike Mo6+ and W6+ Generally can get higher ox states for 2nd and 3rd row TMs Larger ions can have higher CN; CN = 6 is generally the max in 1st row TM complexes, but CN = 7-9 common for 2nd and 3rd row TM [Cr(CN)6]3 (Oh) vs [Mo(CN)8]3 (D4d square anti-prism)

  13. polyoxometallates Polyoxometallates • Metal atoms linked via shared ligands, usually corner or edge-shared Td or Oh • Common for groups 5 (V Nb Ta) and 6 (Cr Mo W) • pH dependence: • high pH Al(OH)4VO43MoO42no M-O-M • decr pH, • decr chg / vol • lower pH Al2O3 (s) V2O5(s) MoO3(s) extensive M-O-M

  14. Vanadates 2 H2VO4- + H+  H3V2O7- + H2O pKa ~ 4 metavanadate chains, (VO3) NaVO3

  15. Polyoxometallates decavanadate has edge-sharing Oh 6 MoO42 + 10 H+ Mo6O192 + 5 H2O M6O19n ;M = Nb,Ta (group 5); Mo,W (group 6) There are 6 edge-sharing Oh, each Oh has 1 unique O 1 4 shared O 4 x ½ 1 center O 1 x 1/6 total O / M 3 1/6 =M6O19

  16. Keggin structure [PMo12O40]3 Keggin structures Td site at cluster center, can also be As,Si,B,Te,Ti PO43- + 12 WO42- + 27 H+ H3PW12O40 + 12 H2O http://en.wikipedia.org/wiki/Keggin_structure (ref Fig below) X2M18O62n− Dawson structure

  17. Ferrodoxins

  18. Clusters (M-M bonding) “NaReCl4” is royal blue, diamagnetic [Mo2(CH3CO2)4] Mo-Mo = 2.11 Å 2 Mo(CO)6 + 4 CH3COOH  [Mo2(O2CCH3)4] + 4 H2 + 12 CO [Re2Cl8]2- D4h Re-Re = 2.24 Å < ClReRe = 104° Re(m) has Re-Re = 2.74 Å and Tm=3180°C ; Mo(m) Mo-Mo = 2.80Å

  19. M-M bonding interactions [M2X8]n common in groups 6-9 (Mo, W, Re, Ru, Rh)

  20. Electronic configurations Cluster ions config b.o. b.l. [Mo2(SO4)4]4Mo(II) d4 σ242 4 2.11 Å [Mo2(SO4)4]3Mo(II) d4 σ241 3.5 2.17 Å Mo(III) d3

  21. Electronic configurations Cluster ions config b.o. b.l. [Mo2(HPO4)4]2Mo(III) d3σ24 3 2.22 Å [Ru2Cl2(O2CCl)4] Ru(II) d6 Ru(III) d5 σ242*p*2 2.5 2.27 Å

  22. Electronic Configurations

  23. 3 Zr(s) + ZrCl4(g)  4 ZrCl (s) Larger Metal Clusters [Re3Cl12]3- ZrCl Zr-Zr bondlengths intrasheet 3.03 Å Intersheet 3.42 Å In Zr (m) 3.19 Å

  24. MoCl2 and [Mo6Cl14]2- HCl (aqu) [Mo6Cl14]2- MoCl2 4 of the 6 Cl bridge to other Mo6 clusters For each Mo6: 8 Cl capping faces 4 (½ Cl) bridging 2 Cl unique 12 Cl / Mo6 cluster Similar for M = Mo, W, Nb,Ta

  25. Groups 8-11 Noble metals : groups 8 – 11 except Fe, Co, Ni metallic forms can exist under environment conditions (see Pourbaix diagrams) Group 11 metals (Cu, Ag, Au) can even exist in strong acid, for example Au does not react with HCl (conc) Au (s)  [AuCl4]- (aq) + NO (g) [Au(CN)2]- (aq) NO3 oxidant, Cl forms stable complex 3 HCl / 1 HNO3 “aqua regia” O2 / CN

  26. Group 11 • +1 state = d10 no LFSE - usually CN = 2 linear (VSEPR) • often disproportionate • 2 Cu+  Cu (s) + Cu2+ E = +0.36 at pH = 0 • 1.2 at pH = 14 • - soft LA Kf I > Br > F • R3P > R3N • S2- > O2 • +3 state = d8 • usually D4h square planar (ex AuCl4-) • Ni(II) Cu(III) • Rh(I) Pd(II) Ag(III) • Ir(I) Pt(II) Au(III) AuCl sometimes Td AuF3

  27. Group 12 (Zn, Cd, Hg) Not noble metals; Zn, Cd are readily oxidized pH = 0 Fe/Fe2+ E0 = + 0.44V Cu/Cu2+ E0 =  0.34V Zn/Zn2+ E0 = +0.76V Why the aperiodic change from group 11 to 12 ? B–H approach: Cu Zn M (s)  M (g) + 338 +131 kJ/mol M (g)  M2+ (g) + 2 e +2704 +2639 M (s)  M2+ (g) + 2 e +3012 +2770 Zn(m) is used for anodic protection (sacrificial anode) www.boatzincs.com/shaft.html

  28. Group 12 Group 12 has d10s2 filled orbitals, much weaker M–M bonding, and lower IE MP Cu 1080°C Zn 420 C Cd 320 Hg - 39 Zn2+ common CN = 4 (6) Cd2+ common CN = 6 (4) Hg2+ common CN = 2 (linear) Hg2+ is stable in aqu solution HgCl – mercurous chloride (calomel) is [Hg2]2+ 2Cl Raman band at 171cm1 Hg–Hg stretch Diamagnetic (Hg+ would be d10s1) XRD bondlengths Hg (m) 300 pm Hg22+ 250-270 pm

  29. Hg32+ linear ion (catenation) (6-x) Hg + 3 AsF5 2 Hg3 x/2 AsF6 + AsF3 Superconductor Tc ~ 4 K Hg3NbF6 2D hex Hg plane SO2(l) Hg catenation Gray = Hg, white = F, black = Nb

  30. f-block elements Relatively constant electroneg across block (shielding keeps Z* = Z-σ nearly constant), so chemistry is very consistent across f-block Ions – have only f valence e Ce = [Xe]4f2 6s2 Ce3+ = [Xe]4f1 Ce4+ = [Xe] All Ln have 3+ as their most stable oxidation state Ce4+ is relatively stable (f) E0 (Ce4+/ Ce3+) = +1.76V strong oxidant Eu2+ “ “ “ (f7) E0 (Eu2+/ Eu3+) = + 0.35V mild reductant

  31. Actinide FrostDiagrams

  32. Pourbaix f-block

  33. Ligand interactions f-block metal – ligand interactions: Ligands have less influence on f orbitals f–f electronic transitions are sharp, relatively independent of ligand type, and long-lived (slow non-radiative energy transfer)  luminescence d–d transition forbidden (Laporte selection rules) Eu(III) 1 % gives bright orange-red luminescence Gd2O2S: Pr Gd(III) = f7 colorless (spin forbidden transitions) Pr(III) = f2 green

  34. Actinides actinides +3 oxidation state common, but high ox states also: Th4+ (f); U3+U6+ all common ArO22+ linear cation for U, Np, Pu, Am UO22+ uranyl cation (bright yellow) High CN common (8-10) ThO2 ThCl4 [UO2(NO3)2(OH2)4]

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