Zumdahl’s Chapter 20

1. Zumdahl’s Chapter 20 Transition Metals

2. e– configuration Oxidation #s & IP Coordination Compounds Coordination # Ligands Nomenclature Isomerism Structural Isomerism Stereoisomerism Bonding in Complex Ions Crystal Field Theory Octahedral Tetrahedral Chapter Contents

3. Electronic Configurations • d – block transition metals • ns2 (n–1)d X where n = 4,5,6,7 • Potential for high spin (Hund’s Rule) • Ions lose s electrons first. • f – block transition elements • ns2 (n–1)d0,1 (n–2)f X where n = 6,7 • Lanthanides & Actinides are even more similar than members of d – block.

4. Oxidation States • Often lose e– to Rare Gas configuration. • But beyond Mn, transition metal ions do not achieve that high. • Because the 8th IP is prohibitively expensive!

5. Coordination Compounds • Often complex ions (both cat– and an–) • But neutrals possible if ligands exactly balance metal ion’s charge. • Often highly colored • Since MO energy separations match visible light photon energies,  absorb visible light. • Often paramagnetic • Duhh! These are transition metals, no? • Dative bonded by e– donating ligands.

6. But to only one of many solvent water molecules. Here’s Gd bonding to a ligand called DOTA 6 ways … Coordination Number • The number of ligand bonds • Usually 6(octahedral) but as few as 2(linear) and as many as 8(prismatic or antiprismatic cube). For a bizarre 7 coordination.

7. 6-coordinated metals like cobalt sepulchrate : C12H24N8Co2+ Or the one we used in lab, MgEDTA2– C10H12O8N2Mg2– Sane Coordination Numbers

8. ethylene diamine halides Ligands • From Latin ligare, “to bind” • Must be a Lewis base (e– donor) • Could, as does EDTA, have several Lewis base functionalities: polydentate! • If monodentate, should be small enough to permit others to bind. • Relative bonding strengths: • X– < OH– < H2O < NH3 < en < NO2– < CN–

9. Naming Anionic Names • Anions that electrically balance cationic coordination complexes can also be present as ligands in that complex! • So they need different names that identify when they’re being used as ligands:

10. Naming Neutral Names • But ligands needn’t be anions; many neutral molecules are Lewis bases. • And they too get new names appearing as ligands in coordination complexes:

11. Name That Complex, Oedipus • [ Cr Br2 (en)2 ] Br • Anion, bromide, is named last (no surprise) • chromium(III) is named next-to-last • Ligands named 1st in alphabetical order: • Number of a ligands is shown as Greek prefix: • dibromo … • Unless it already uses “di” then use “bis” • Dibromobis(ethylenediammine) … • Dibromobis(ethylenediammine)chromium(III) bromide

12. Charge Overrun • Since ligands are often anions, their charge may swamp the transition metal, leaving the complex ion negative! • Na2 [ PbI4 ] (from Harris p. 123) • Sodium tetraiodoplumbate(II) • While lead(II) is the source, the Latin root is used for the complex with “ate” denoting anion. • Li [ AgCl2 ], lithium dichloroargentate

13. Isomeric Complications • dichlorobis(diethylsulfide)platinate(II) would appear to be the name of the square planar species above, but • The square planar configuration can have another isomer where the Cl ligands are on opposite sides of the platinum, so it’s really • cis-dichlorobis(diethylsulfide)platinate(II) • and this is not the only way isomers arise!

14. Complex Isomerization Simplified • Stereoisomers preserve bonds • Geometric (cis-trans) isomers • Optical (non-superimposable mirrors) • Structural isomers preserve only atoms • Coordination isomers swap ligands for anions to the complex. • Linkage isomers swap lone pairs on the ligand as the bonding site.

15. Coordination Isomers • Unique to coordination complexes • [ Pb (en)2 Cl2 ] Br2 • bis(ethylenediammine)dichlorolead(IV) bromide • Only 1 of 3 possible coordination isomers • The other 2 are • [ Pb Br (en)2 Cl ] Br Cl • bromobis(ethylenediammine)chlorolead(IV) • bromide chloride • [ Pb Br2 (en)2 ] Cl2 • dibromobis(ethylenediammine)lead(IV) chloride

16. Optical Isomers • We need to compare the mirror image of a sample complex to see if it can be superimposed on the original. These views of cobalt sepulchrate and its Mirror image demonstrate non-superimposition. They are optical isomers.

17. Colorful Complexes • Colors we see everywhere are due, for the most part, to electronic transitions. • Most electronic transitions, however, occur at energies well in excess of visible h. • d-electrons transitions ought not to be visible at all, since they are degenerate. • But, in a complex, that degeneracy is broken! Transition energies aren’t then 0.

18. Breaking Degeneracy • 5 d orbitals in a tetrahedral charge field split as a doublet (E) and a triplet (T).

19. Symmetry Tells Not All • While the symmetry tables assure us that there are now 2 energy levels for d orbitals instead of 1, we don’t know the energies themselves. • That depends upon the field established by the ligands and the proximity of the d s. • See Zumdahl’s Fig. 20.26 for a visual argument why dxy,dxz,dyz are lower energy.

20. Other Ligand Symmetries • Octahedral, Oh, (6-coordinate, Fig. 20.20) • Eg symmetic species for (2z2–x2–y2, x2–y2) • T2g symmetric species for (xy, xz, yz) • Square Planar, D4h(Fig. 20.27a) • A1g symmetric species for z2 • B1g symmetric species for x2–y2 • B2g symmetric species for xy • Eg symmetric species for (xz, yz)

21. Consequences • Degeneracies work in Hund’s favor to separate e– pairs and maximize spin. • With high enough energy separations, , Aufbau (lowest level) wins instead. • High field case,  large, e– pairs in lower energy states. • Low field case,  small, e– unpaired as much as feasible.

22. Symmetry and  • tetrahedral = (4/9) octahedral(same ligands) • As a consequence of symmetry. • If some ligand was 9/4 as strong as the weakest to give octahedral strong field, then strong field (low-spin) tetrahedral might exist. But none does. • Field strengths of ligands vary as: • X– < OH– < H2O < NH3 < en < NO2– < CN–