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Chapter 20 The Transition Metals

Chapter 20 The Transition Metals. Survey of Transition Metals General Properties Transition metals have similar reactivity across the period Representative Elements: filling valence shells as you go across a period Transition Metals: filling core subshells (d or f) as you go across a period

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Chapter 20 The Transition Metals

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  1. Chapter 20 The Transition Metals • Survey of Transition Metals • General Properties • Transition metals have similar reactivity across the period • Representative Elements: filling valence shells as you go across a period • Transition Metals: filling core subshells (d or f) as you go across a period • d-block elements (Sc—Zn) • f-block elements (Ce—Lu) • Core electrons don’t effect reactivity as much as valence electrons

  2. All Transition elements are metals: luster, conductors, malleable, ductile • Transition metals do vary in their properties • Melting points: W = 3400 oC, Hg = 25 oC • Hardness: Titanium is hard, Gold is soft • Chemical Reactivity: Iron is very reactive (rust), Platinum is very inert • Formation of Ionic compounds with nonmetals • More than one oxidation number is possible: FeCl2 vs. FeCl3 • Different than Main Group metals: NaCl, CaCl2, AlCl3 • Formation of Complex Ions • Transition metal cation at the center • Multiple Lewis Bases (Ligands) donating lone pairs • Compounds are generally colored • Absorption of light by d-electrons • Most main group compounds are colorless (or black) • Many transition metal complexes are paramagnetic: unpaired d-electrons

  3. Electron Configuration of First Row Transition Metals (Sc—Zn) • Expected configuration is usual: Sc = [Ar]4s23d1, Ti = [Ar]4s23d2, V = [Ar]4s23d3 • Chromium has an unexpected configuration. Why? Cr = [Ar]4s13d5 • 3d and 4s subshells have almost same energies for first row metals • Cr occupies essentially degenerate orbitals equally: 4s__ 3d__ __ __ __ __ • Copper is the only other exception having a single 4s electron: Cu = [Ar]4s13d10 • For cations, 3d energy < 4s energy • 4s electrons are lost first because remaining electrons are more stable Cr2+ = d4 4s__ 3d__ __ __ __ __ • Mn2+ Mn0 = [Ar]4s23d5 Mn2+ = [Ar]3d5 • Ti3+ Ti0 = [Ar]4s23d2 Ti3+ = [Ar]3d1 • Oxidation state and ionization energy (Table 21.2) • For the first five metals: maximum oxidation state = loss of all 4s and 3d electrons Cr0 -----> Cr6+ (loss of 4s13d5) Ti0 -----> Ti4+ (loss of 4s23d2) • Last five metals: usually +2 or +3 • 3d energy lowers as the nuclear charge increases • Harder to remove more electrons 3) Ionization Energy increases across the period: M0 -----> Mn+ + n e-

  4. First row transition metal Ionization Energies Removal of first d-electron Removal of first s-electron

  5. D. Standard Reduction Potentials • Generally listed with metal as reducing agent (metal is oxidized) • This is the reverse of how reduction potentials are usually written • 2H+ + 2e- -----> H2 E = 0.0 V becomes • H2 -----> 2H+ + 2e- E = 0.0 V • “Oxidation Potentials” as written would have the most easily oxidized metal with the most positive potential • All but copper can reduce H+ to H2 • Zn + 2H+ -----> Zn2+ + H2 E0 = +0.76 • 5) Reducing Ability decreases from left to right (Zn and Cr are exceptions)

  6. 4d and 5d Transition Metals • Atomic Radii • Decreases from left to right: more +/- attraction in same shell • Increases as you go from 3d to 4d to 5d: filling larger shell • 4d and 5d metals are very similar in size = Lanthanide Contraction • Fifth period fills up 4f electrons before beginning to fill 5d subshell • Each 4f electron added decreases the radius, so the atoms are smaller than expected by the time we get to 5d • By the time we get to La, size is back down to near that of Y • Other Properties • Generally less reactive than first row metals (poor orbital size match) • Important for specialized materials • Range from rare to very rare in abundance on Earth’s crust

  7. Chemistry of the First Row Transition Metals • Scandium (Sc) 1) Rare element found mostly as Sc3+ in ionic compounds like Sc2O3 2) Chemistry is most similar to Lanthanides = colorless, no d-electrons, diamagnetic 3) Sc0 is used in some high intensity lamps • Titanium (Ti) • Common (0.6% of crust), strong, low density metal; structural uses (jet engines) • Fairly unreactive, so used for pipes, pumps, reaction vessels • Usually found in Ti4+ state, no d-electrons, colorless/white • TiO2 (titanium(IV) oxide) used as white pigment in paper, paint, etc… • Ti3+ is stable: Ti(H2O)63+ complex ion is purple (1 d-electron) • Ti2+ is not very stable, but TiO and TiCl2 do exist • Vanadium (V) • Common (0.02%) metal most often used as alloy with Fe or Ti • Usually found in V5+ state, colorless, no d-electrons • V2O5 is an industrial catalyst used in the manufacture of H2SO4 • V4+, V3+, V2+ can all exist; usually strong reducing agents (form V5+)

  8. Chromium (Cr) • Rare, but important industrially • Chromite Ore reduced to Ferrochrome = mixture of Fe and Cr used in steelmaking FeCr2O4 + 4C -----> Fe + 2Cr + 4CO • Pure Cr metal used to plate steel: shiny surface due to “invisible” Cr oxide • Cr2+, Cr3+, and Cr6+ oxidation states are common (Table 20.5) • Cr2+ is a powerful reducing agent (used to remove O2) • Cr6+ is a powerful oxidizing agent: Dichromate = Cr2O72- -----> Cr3+ Cr2O72- + 14H+ + 6e- -----> 2Cr3+ + 7H2O Eo = 1.33 V • Dichromate equilibrates in basic water to chromate = CrO42- • CrO3 in H2SO4 very strong oxidizer used to clean glassware • Manganese (Mn) • Abundant (0.1%) but no sources in the United States • With iron, it makes an especially hard steel (crush rock, bank vaults) • Nodules of almost pure Mn made on sea floor by marine organisms • Mn2+ up to Mn7+ are all possible, Mn2+ and Mn7+ are most common • Permanganate = MnO4- deep purple strong oxidizing agent MnO4- + 8H+ + 5e- -----> Mn2+ + 4H2O Eo = 1.5 V

  9. Iron (Fe) • Most abundant transition metal (4.7%) • Solid form rapidly reacts with O2 to form rust = Fe2O3 • Usually found as Fe2+, Fe3+ oxidation states • Makes steel in combination with other elements • Important biologically (later) • Cobalt (Co) • Rare, found in ores = smaltite = CoAs2 • Hard metal with a blue/white appearance • Mostly used in alloys with Fe, Cu, W • Co2+ and Co3+ are most common oxidation states • Historically important in the development of “Coordination Chemistry” • Nickel (Ni) • Abundant in ores with As, Sb, and S • Corrosion resistant, white metal used to plate other metals • Usually in the Ni2+ oxidation state, such as in the green compound Ni(H2O)62+ • Well characterized in many coordination compounds

  10. Copper (Cu) • Abundant in ores with S, As, Cl, and CO32- • High electrical conductance, low corrosion make it suitable for wires and pipes • Corrodes in air to green “patina”: Cu + 2H2O + SO2 + O2 -----> Cu3(OH)4SO4 • Usually found as Cu+ or Cu2+ as in blue Cu(H2O)62+ • Biologically necessary, but toxic in large amounts • Zinc (Zn) • Abundant in ore = sphalerite = ZnS • Good reducing agent: Zn -----> Zn2+ + 2e- E = 0.76 V • Colorless, full d-orbitals • Zn2+ is the only oxidation state commonly encountered • Coordination Compounds • Definitions • Complex Ion = metal cation with Lewis Bases covalently bonded to it • Ligands = Lewis bases covalently bonded to a metal ion • Counter Ions = anions present in the formula to balance the charge of complex ion • [Co(NH3)5Cl]Cl2 = Co(NH3)5Cl2+ (complex ion) + 2 Cl- (counter ions) • Solid = packed cations/anion; Solution = dissociated like NaCl

  11. Alfred Werner developed theory of coordination chemistry in 1890’s • Primary Valence = total positive charge = number of anions total • Secondary Valence = Lewis Acid ability = how many Lewis Bases bound • [Co(NH3)5Cl]Cl2 = CoCl3• 5 NH3 • Primary Valence: +3 is satisfied by 3 Cl- total • Secondary Valence: 6 Lewis bases bound to Co (5 NH3, 1 Cl) • Primary Valence = Oxidation State • Secondary Valence = Coordination Number • Coordination Number • Metals can bind between 2 and 8 ligands: Coordination Number = CN = 2-8 2) Size, charge, d-electron count decide how many ligands bind 3) Most prevalent CN = 6 > 4 > 2 4) A given metal ion can have any CN, depending on what the ligands are 5) Common coordination numbers: Table 21.12

  12. Ligands • Coordinate Covalent Bond = Lewis base donates both electrons (lone pair) • Monodentate = one-toothed = only able to donate one lone pair to the metal ion H2O, OH-, NH3, CN-, SCN-, Cl- • Bidentate = 2 bonds from the same ligand • Chelate = claw • Ethylenediamine, oxalate • Polydentate • Diethylenetriamine: H2NCH2CH2NHCH2CH2NH2 • EDTA = ethylenediaminetetraacetate • Very stable complexes • Used to scavenge toxic metals

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