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William A. Goddard, III, wag@kaist.ac.kr

Lecture 26, December 3, 2009. Nature of the Chemical Bond with applications to catalysis, materials science, nanotechnology, surface science, bioinorganic chemistry, and energy. Course number: KAIST EEWS 80.502 Room E11-101 Hours: 0900-1030 Tuesday and Thursday.

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William A. Goddard, III, wag@kaist.ac.kr

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  1. Lecture 26, December 3, 2009 Nature of the Chemical Bond with applications to catalysis, materials science, nanotechnology, surface science, bioinorganic chemistry, and energy Course number: KAIST EEWS 80.502 Room E11-101 Hours: 0900-1030 Tuesday and Thursday William A. Goddard, III, wag@kaist.ac.kr WCU Professor at EEWS-KAIST and Charles and Mary Ferkel Professor of Chemistry, Materials Science, and Applied Physics, California Institute of Technology Senior Assistant: Dr. Hyungjun Kim: linus16@kaist.ac.kr Manager of Center for Materials Simulation and Design (CMSD) Teaching Assistant: Ms. Ga In Lee: leeandgain@kaist.ac.kr Special assistant: Tod Pascal:tpascal@wag.caltech.edu EEWS-90.502-Goddard-L15

  2. Schedule changes Dec. 3, Thursday, 9am, L26, as scheduled Dec. 7-10 wag in Pasadena; no lectures, Dec. 14, Monday, 2pm, L27, additional lecture, room 101 Dec. 15, Final exam 9am-noon, room 101 EEWS-90.502-Goddard-L15

  3. Last time EEWS-90.502-Goddard-L15

  4. Bonding in metallic solids Mosty of the systems discussed so far in this course have been covalent, with the number of bonds related to the number of valence electrons. Thus we have discussed the bonding of molecules such as CH4, benzene, O2, and Ozone.. The solids such as diamond, silicon, GaAs, are generally insulators or semiconductors We have also considered covalent bonds to metals such as FeH+, (PH3)2Pt(CH3)2, (bpym)Pt(Cl)(CH3), The Grubbs Ru catalysts We have also discussed the bonding in ionic materials such as (NaCl)n, NaCl crystal, and BaTiO3, where the atoms are best modeled as ions with the bonding dominated by electrostatics Next we consider the bonding in bulk metals, such as iron, Pt, Li, etc. where there is little connection between the number of bonds and the number of valence electrons. EEWS-90.502-Goddard-L15

  5. Bringing atoms together to form the solid As we bring atoms together to form the solid, the levels broaden into energy bands, which may overlap . Thus for Cu we obtain Energy Fermi energy (HOMO and LUMO Thus we can obtain systems with no band gap. Density states EEWS-90.502-Goddard-L15

  6. Metals vs inulators EEWS-90.502-Goddard-L15

  7. conductivity EEWS-90.502-Goddard-L15

  8. The elements leading to metallic binding There is not yet a conceptual description for metals of a quality comparable to that for non-metals. However there are some trends, as will be described EEWS-90.502-Goddard-L15

  9. Body centered cubic (bcc), A2 A2 EEWS-90.502-Goddard-L15

  10. Face-centered cubic (fcc), A1 EEWS-90.502-Goddard-L15

  11. Alternative view of fcc EEWS-90.502-Goddard-L15

  12. Closest packing layer EEWS-90.502-Goddard-L15

  13. Stacking of 2 closest packed layers EEWS-90.502-Goddard-L15

  14. Hexagaonal closest packed (hcp) structure, A3 EEWS-90.502-Goddard-L15

  15. Cubic closest packing EEWS-90.502-Goddard-L15

  16. Double hcp The hexagonal lanthanides mostly exhibit a packing of closest packed layers in the sequence ABAC ABAC ABAC This is called the double hcp structure EEWS-90.502-Goddard-L15

  17. Structures of elemental metals some correlation of structure with number of valence electrons EEWS-90.502-Goddard-L15

  18. Binding in metals Li has the bcc structure with 8 nearest neighbor atoms, but there is only one valence electron per atom. Similarly fcc and hcp have 12 nearest neighbor atoms, but Al has only three valence electrons per atom. Clearly the bonding is very different than covalent One model (Pauling) resonating valence bonds Problem is energetics: Li2 bond energy = 24 kcal/mol 12 kcal/mol per valence electron Cohesive energy of Li (energy to atomize the crystal is 37.7 kcal/mol per valence electron. Too much to explain with resonance New paradigm: Interstitial electron model (IEM). Each valence electron localizes in a tetrahedron between four Li nuclei. Bonding like in Li2+, which is 33.7 kcal/mol per valence electron EEWS-90.502-Goddard-L15

  19. GVB orbitals of ring M10 molecules Get 10 valence electrons each localized in a bond midpoint Calculations treated all 11 valence electrons of Cu, Ag, Au using effective core potential. All electrons for H and Li R=2 a0 EEWS-90.502-Goddard-L15

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  27. New EEWS-90.502-Goddard-L15

  28. Hypervalent compounds It was quite a surprize to most chemists in 1962 when Neil bartlett reported the formation of a compound involving Xe-F bonds. But this was quickly folllowed by the synthesis of XeF4 (from Xe and F2 at high temperature and XeF2 in 1962 and later XeF6.Indeed Pauling had predicted in 1933 that XeF6 would be stable, but noone tried to make it. Later compounds such as ClF3 and ClF5 were synthesized These compounds violate simple octet rules and are call hypervalent EEWS-90.502-Goddard-L15

  29. Noble gas dimers Recall from L17 that there is no chemical bonding in He2, Ne2 etc This is explained in VB theory as due to repulsive Pauli repulsion from the overlap of doubly occupied orbitals It is explained in MO theory as due to filled bonding and antibonding orbitals (sg)2(su)2 EEWS-90.502-Goddard-L15

  30. Noble gas dimer positive ions - On the other hand the positive ions are strongly bound (L17) This is explained in MO theory as due to one less antibonding electron than bonding, leading to a three electron bond for He2+ of 2.5 eV, the same strength as the one electron bond of H2+ (sg)2(su)1 The VB explanation is a little less straightforward. Here we consider that there are two equivalent VB structures neither of which leads to much bonding, but superimposing them leads to resonance stabilization Using (sg) = L+R and (su)=L-R Leads to (with negative sign EEWS-90.502-Goddard-L15

  31. Re-examine the bonding of HeH - He+ H- He H IP=+24.6 eV EA = 0.7 eV Why not describe HeH as (sg)2(su)1 where (sg) = L+R and (su)=L-R Would this lead to bonding? The answer is no, as easily seen with the VB form where the right structure is 23.9 eV above the left. Thus the energy for the (sg)2(su)1 state would be +12.0 – 2.5 = 9.5 eV unbound at R=∞ Adding in ionic stabilization lowers the energy by 14.4/2.0 = 7.2 eV (too big because of shielding) , still unbound by 2.3 eV EEWS-90.502-Goddard-L15

  32. Examine the bonding of XeF Consider the energy to form the charge transfer complex Xe Xe+ The energy to form Xe+ F- can be estimated from Using IP(Xe)=12.13eV, EA(F)=3.40eV, and R(IF)=1.98 A, we get E(Xe+ F-)=1.45eV Thus there is no covalent bond for XeF, which has a weak bond of ~ 0.1 eV and a long bond EEWS-90.502-Goddard-L15

  33. Examine the bonding in XeF2 Xe+ We saw that the energy to form Xe+F-, now consider, the impact of putting a 2nd F on the back side of the Xe+ Since Xe+ has a singly occupied pz orbital pointing directly at this 2nd F, we can now form a bond to it? How strong would the bond be? Probably the same as for IF, which is 2.88 eV. Thus we expect F--Xe+F- to have a bond strength of ~2.88 – 1.45 = 1.43 eV! Of course for FXeF we can also form an equivalent bond for F-Xe+--F. Thus we get a resonance We will denote this 3 center – 4 electron charge transfer bond as FXeF EEWS-90.502-Goddard-L15

  34. Stability of XeF2 Ignoring resonance we predict that XeF2 is stable by 1.43 eV. In fact the experimental bond energy is 2.69 eV suggesting that the resonance energy is ~ 1.3 eV. The XeF2 molecule is stable by 2.7 eV with respect to Xe + F2 But to assess where someone could make and store XeF2, say in a bottle, we have to consider other modes of decomposition. The most likely might be that light or surfaces might generate F atoms, which could then decompose XeF2 by the chain reaction XeF2 + F  {XeF + F2}  Xe + F2 + F Since the bond energy of F2 is 1.6 eV, this reaction is endothermic by 2.7-1.6 = 1.1 eV, suggesting the XeF2 is relatively stable. Indeed it is used with F2 to synthesize XeF4 and XeF6. EEWS-90.502-Goddard-L15

  35. XeF4 Putting 2 additional F to overlap the Xe py pair leads to the square planar structure, which allows 3 center – 4 electron charge transfer bonds in both the x and y directions, leading to a square planar structure The VB analysis would indicate that the stability for XeF4 relative to XeF2 should be ~ 2.7 eV, but maybe a bit weaker due to the increased IP of the Xe due to the first hypervalent bond and because of some possible F---F steric interactions. There is a report that the bond energy is 6 eV, which seems too high, compared to our estimate of 5.4 eV. EEWS-90.502-Goddard-L15

  36. XeF6 Since XeF4 still has a pz pair, we can form a third hypervalent bond in this direction to obtain an octahedral XeF6 molecule. Here we expect a stability a little less than 8.1 eV. Pauling in 1933 suggested that XeF6 would be stabile, 30 years in advance of the experiments. He also suggested that XeF8 is stable. However this prediction is wrong EEWS-90.502-Goddard-L15

  37. Estimated stability of other Nobel gas fluorides (eV) 1.3 1.3 1.3 1.3 1.3 1.3 -5.3 -0.1 1.0 2.7 3.9 -2.9 Using the same method as for XeF2, we can estimate the binding energies for the other Noble metals. Here we see that KrF2 is predicted to be stable by 0.7 eV, which makes it susceptible to decomposition by F radicals RnF2 is quite stable, by 3.6 eV, but I do not know if it has been observed EEWS-90.502-Goddard-L15

  38. XeCl2 Since EA(Cl)=3.615 eV and R(XeCl+)=2.32A and De(XeCl+)=2.15eV, can estimate that XeCl2 is stable by 1.14 eV with respect to Xe + Cl2. However since the bond energy of Cl2 is 2.48 eV, the energetics of the chain dempostion process are exothermic by 1.34 eV, suggesting at most a small barrier Thus XeCl2 would be difficult to observe EEWS-90.502-Goddard-L15

  39. Halogen Fluorides, ClFn The IP of ClF is 12.66 eV which compares well to the IP of 12.13 for Xe. This suggests that the px and py pairs of Cl could be used to form hypervalent bonds leading to ClF3 and ClF5. Indeed these estimates suggest that ClF3 and ClF5 are stable. Indeed the experiment energy for ClF3  ClF +F2 is 2.6 eV, quite similar to XeF2. EEWS-90.502-Goddard-L15

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  49. Origin of reactivity in the hypervalent reagent o-iodoxybenzoic acid (IBX) Julius Su and William A. Goddard III Materials and Process Simulation Center, California Institute of Technology EEWS-90.502-Goddard-L15

  50. Hypervalent iodine assumes many metallic personalities Oxidations CrO3/H2SO4 Radical cyclizations SnBu3Cl Electrophilic alkene activation HgCl2 CC bond formation Pd(OAc)2 Can we understand iodine as we understand metals? Martin, J. C. organo-nonmetallic chemistry– Science 1983 221(4610):509-514 EEWS-90.502-Goddard-L15

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