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Chapter 24. Transition Metals & Coordination Compounds. Gemstones. The colors of rubies and emeralds are both due to the presence of Cr +3 ions – the difference lies in the crystal hosting the ion In rubies, some Al +3 ions in the Al 2 O 3 are replaced by Cr +3 ions.

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Chapter 24

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    1. Chapter 24 Transition Metals & Coordination Compounds

    2. Gemstones • The colors of rubies and emeralds are both due to the presence of Cr+3 ions – the difference lies in the crystal hosting the ion • In rubies, some Al+3 ions in the Al2O3 are replaced by Cr+3 ions. • In emeralds, some Al+3 ions in the Be3Al2(SiO3)6 are replaced by Cr+3 ions.

    3. Electron Configuration • For 1st & 2nd transition series = ns2 (n−1)dx • Fe = [Ar]4s23d6; Zr= [Kr]5s24d2 • For 3rd transition series =ns2 (n−2)f14 (n−1)dx • Re = [Xe] 6s2 4f14 5d5 • Some individuals deviate from the general pattern by “promoting” one or more s electrons into the underlying d to complete the subshell • Form ions by losing the ns electrons first, then the (n – 1)d

    4. Lewis Acids & Bases • Section 15.11 • G.N. Lewis – noticed that acid-base chemistry always involves an electron pair. • BH3 + NH3 H3B:NH3 • Acid = electron pair acceptor • Base = electron pair donor • Greatly expands what we view as an “acid” • LEP #1

    5. Complexes • An ion like [Ag(NH3)2]+1, are called complex ions as well as coordination compounds. • The molecules or ions that bond to the metal are known as ligands. • The coordination sphere is the metal and the total number of ligands bonded to it. • The complex is a neutral charge salt, which may contain additional cations or anions not bonded to the metal.

    6. Complexes • [Cu(NH3)4] SO4 • The complex ion charge is _____. • The charge of Copper is _____. • Coordination number is the number of lone pairs donated to the metal. • Coordination numbers of 2, 4, and 6 are most common.

    7. Complexes

    8. Molecular Geometry

    9. Chelates • Ligands are sometimes referred to as chelates (Greek = claw). • Most are monodentate (one “toothed”) like NH3, Cl-, CN-, etc. • A few are bidentate (two “toothed”) like ethylenediamine and the oxalate ion. • A few are polydentate like EDTA.

    10. Chelates • The formation of complexes favors the products as seen in Chapter 17. • Ni+2(aq) + 6 NH3(aq) Ni(NH3)6+2 ; Kf = 4 E8 • Ni+2(aq) + 3 en(aq)  Ni(en)3+2(aq) ; Kf = 2 E18 • The larger K for the bidentate ligand is known as the chelating effect. • Uses of EDTA and the EDTA challenge.

    11. Metals in Living Systems • Nine metals important to life – V, Cr, Mn, Fe, Co, Ni, Cu, Zn, and Mo – owe their roles to their ability to form complexes with ligands. • The role of iron in hemeglobin is a perfect example. • In hemeglobin, the iron is bonded to four N atoms in a molecule called porphoryn. • The fifth site is bonded to the protein (globin). • This leaves one position empty in the octahedral geometry.

    12. Metals in Living Systems Porphine molecule

    13. Nomenclature • Complexes are named using a systematic method. • Rules: • Cation named first, then anion • Name of the complex is always one word, name of ligands come first and in alphabetical order • Name of ligands include prefixes if more than one Anionic ligands get an –o suffix • Name of metal also includes oxidation number in ( ). If complex is an anion, metal name ends in –ate. Ex) Vandium = Vanadate, Ferrum = Ferrate Note: Some metals use old Latin names!

    14. Nomenclature LEP #2, #3

    15. Isomers • Isomers are compounds with the same formula but either atoms are in a different order (structural) or atoms are in a different spatial arrangement (stereoisomers).

    16. Structural Isomers • A linkage isomer occurs when a ligand can bond through a different atom. • NO2- can bond through the N (NO2-) or the O (ONO-). • Another one is SCN-.

    17. Structural Isomers • A coordination sphere isomer occurs when the ligands bonded to the metal are exchanged for ones outside of the coordination sphere. • For example, the formula CrCl36H2O has several forms. • [Cr(H2O)6] Cl3 is purple • [Cr(H2O)5Cl] Cl2H2O is green

    18. Stereoisomers • A geometric isomer occurs when the spatial orientation of a complex can be changed. These are referred to as cis-trans isomers. • Example is the square planar geometry of PtCl2(NH3)2.

    19. Stereoisomers • Can also produce cis-trans for octahedral complexes if general formula is: MX4Y2. • Example is Co(NH3)4Cl2+.

    20. Stereoisomers • A second type of geometric isomerism can occur if the general formula is MX3Y3 called fac-mer (short for facial and meridian). • An example is Co(NH3)3Cl3.

    21. Stereoisomers • An optical isomer occurs when the mirror image of the complex is non-superimposable. • The pair of isomers are called enantiomers.

    22. Stereoisomers • In complexes, the only way to get optical isomerism is with a 6-coordinate system and two or three bidentate ligands. • Most of the chemical and physical properties of any enantiomer pair are identical. • However, towards other optically active molecules only one might react.

    23. Stereoisomers • If the two mirror image complex ions can be separated, then they can be tested with plane polarized light.

    24. Color • Some ions are highly colored. • Cu+2= blue • Ni+2 = green • Co+2 = pink • Some ions are not colored. • Zn+2 • Ba+2 • Al+3

    25. Color • Color depends on two factors: • _______________ • _______________ • Compounds must absorb some visible light to have a color.

    26. Color • A compound’s color can be due to: • either it absorbs all wavelengths but that color • OR, it absorbs one color exclusively • For the second choice, the color is then the complimentary color. Spectrum for Ti(H2O)6+3

    27. Color Wheel • The color wheel shows the complimentary colors. • Those that are opposite are complimentary.

    28. Spectrum of Ni+2

    29. Electron Configurations • In period 4, the d orbitals start with Sc. • Sc: [Ar] 4s2 3d1 • Orbital diagram – shows how each of the d orbitals are filled. • Example) Fe: [Ar] 4s2 3d6 • Will see many metal ions, so that means you have to remove some of the electrons. • Co+3

    30. Magnetism • Unpaired electrons = paramagnetic • Paired electrons = diamagnetic • Zn(Cl4)-2 = diamagnetic • CoF6-3 = paramagnetic • Co(CN)6-3 = diamagnetic • ???

    31. Crystal Field Theory (CFT) • As the ligand donates its electron pair to form the bond, it interacts with the metal’s d orbitals. • Not all the d orbitals are affected in the same way. • This splits the d orbitals into different levels.

    32. d orbitals

    33. d orbitals

    34. CFT

    35. CFT

    36. CFT

    37. High and Low Spin • Normally, electrons fill the d orbitals one at a time WITH parallel spins. • Octahedral complexes • Small D = fill each level first before pairing • Large D = fill the lower level completely before moving to upper level • only matters for d4 to d7 configurations

    38. High and Low Spin • [CoF6]-3 • [Co(CN)6]-3 • They are different!

    39. Spectrochemical Series • Ranks the ligands from weak to strong field. Cl- < F- < H2O < NH3 < en < NO2- < CN-increasing D

    40. Tetrahedral and Square Planar • D is always small for tetrahedral complexes so these are always high spin. • D is always large for square planar complexes so these are always low spin.

    41. Tetrahedral and Square Planar • Ni+2 can be either • d8 • [NiCl4]-2 • [Ni(CN4)]-2

    42. Measuring Delta • D = hc / l • Remember, though, if a compound is red, then it absorbs green. • Use wavelength in green part of the spectrum!