CH11_covalent_bond - Models for Covalent Bonding There are...

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Unformatted text preview: Models for Covalent Bonding There are two popular and successful models for describing covalent bonding. In the localized electron (LE) model a molecule is considered to be a collection of atoms which have localized electron pairs. Bonding is achieved whenever two atoms share the same electron pair. If a localized electron pair is not being shared it remains a lone pair on the atom. The other model takes an opposite point of View by considering a molecule to be a collection of atoms whose electrons are delocalized over the entire molecule. Bonding is achieved whenever the probability for the delocalized electron is concentrated between two atoms. In quantum mechan— ical formulations the LE model yields valence bond theory while the delocalized model yields molecular orbital theory. While the two models may appear to be diametrically opposed, in a quantum mechanical framework when an approximate treatment of both models is refined to become more "exact", the two models merge to give the same result. The localized electron model is relatively easy to use and intuitive — it accords with our perception of what a bond should be. It underlies many useful concepts that we have already employed: 1. Lewis structures — In 1902 G. N. Lewis (of Lewis acid/base fame) conceived of the octet rule (while lecturing to general chemistry students). He was the first to establish the concept of an electron pair bond. The localized picture is evident in a) octet rule b) bonding pairs c) lone pairs d) formal charges e) resonance structures 2. Valence Shell Electron Pair Repulsion Theory - VSEPR was originally proposed in 1940 by Sidgwick and Powell. In 1957 it was further developed by Nyholm and especially Gillespie. a) Thesis - a molecule adapts a shape which maximizes the distance between pairs of elec- trons in its valence shell. This minimizes the electron pair repulsion whose order is BP/BP < BP/LP < LP/LP. b) Uses 1) number of electron pairs (steric number, SN) -> electronic geometry 2) molecular geometry -> bond angles and their distortions 3) qualitative prediction of dipole moment and polarity 3. Hybridization — In the 1920’s Linus Pauling (Nobel Laureate in both chemistry and peace) showed exactly how one could take a linear combination of an s and px, py, and p Z orbitals and obtain four sp3 hybridized orbitals directed to the corners of a tetrahedron which explained the bonding in CH4. In 1931 he similarly explained the bonding in octahedral complexes by taking a linear combination of an s, three p’s (px, py, pl), and the dx2_y2 and dz2 orbitals (those either just above or just below the s and p orbitals). For ns and np, either ndeyz and ndzz or (n— 1)dx2_y2 and (n—1)dzz to form six sp3d2 or dzsp3 hybridized -2- orbitals directed to the six corners of an octahedron. The following hybridization schemes are intimately connected to valence bond theory where the designation beneath the steric number refers to the electronic geometry about a central atom (which is identical to the molecular geometry when there are no lone pairs on the central atom): SN=2 sp linear s + pz S — pz commonly d 10: Cu+, Ag+‘, Au+, Hg2+ SN=3 sp2 trigonal s — Vipx planar s — VE/pr +‘l6/2py s — «IE/2px — «16/2 py SN=4 sp3 tetrahedral S + px + Py + Pz ___ S+px_py_pz dx)’ d1xz dxz S—px—py'l'pz S " Px + py — pz At =4/9 A0 l dx2_y2 a; high spin complexes commonly d 10: Zn2+, Cd2+, Hg2+ SN=4 dsp2 square planar s an Py d x2_y2 d x2_y2 big $ _ gfl _dz2_ de dyz low spin complexes usually d8: Ir+, Ni2+, Pd2+, Pt“, Au3+ HCN No; [CuBr2]‘ [Hg (CH3 )2] so2 O3 CH2=CH2 [511(01'1)3l~ [Cu<CN)3]2' H20 H3O+ SiC14 [Be(H20)412+ [BF4l' [TiC14]‘ [FeC14]* [Ni(NH3)4]2+ [Zn(OH)4]2‘ [Co(CN)4]2‘ [Ni(CN)412' [PdCl412' [PtCl4]2‘ [Cll(1\11'13)4l2+ [AgF4l— [AuCl4]' N3 C02 [Ag(NH3 )2]+ [chu] Sbc1§ BF3 H2C=O [Pb(0H)3l_ [HgI3l_ H28 NH3 CH4 [SnCl4] [A1H41‘ [VC14] [CoBr4]2‘ [CuBr4]2_ [Cd(CN)4]2‘ [RC12(PH3)21_ [Pd(NH3)412+ [P'[(1\H‘I3)4l2+ [Cu(H20)4]2+ [Au(CN)4]_ BCHZ H—CEC-H [Au(CN)2l_ NO; 803 benzene SnClg NH: [Cr0412‘ [GaBr4]' [MnCl4]2" [C014]2' [Cu(CN)4]3“ [HgI412' [Pt(CN)412‘ SN=5 dsp3/sp3d trigonal S bipyramidal px, py, pz dzz d—zz dx2_y2 @ dyz d xz SN=6 d2sp3/sp3d2 octahedral S __ px, py’ pz d x2_y2 dzz deyZ 7 d 12 T A0 i dxy dyz dxz all d8, d9, d” are sp3d2 Sp3d2‘ d4 -— d7: high spin weak ligand field small A0 all d1, d2, d3 are dzsp3 dzsp3: d4 — d7: low spin strong ligand field large A0 -3- sp3d: RnClZ BI'F3 PB r5 959:: [TiCl5]3‘ [Mn(CN)5] [Fe(CO)5] [Co(CN)s]“‘ [Cuc1513' [CdC15]3' sp3 d 2 : XeF4 XeOF4 [11101613— [GeF6]2' [BiC16]” [Mn(H20)6]2+ [MnC16]4‘ [Fe(H20)612+ [F6F6]3_ [Co<H20)612+ [CoF6]3‘ [Ni(H20)6]2+ [Ptar— [Cu(NH3)6]2+ [Zn(H20)612+ dzsp3: [Sc(H20)6]3+ [Ti(H20)613+ [TiF6]3‘ [V(H20)6]3+ [V(CO)6]‘ [Cr(HzO)6]3+ [Cr(NH3>613+ [M0(CO)6] [Mn(CN)6]5' [Fe<CN>614' [Ru(CN)614‘ [Co(NH3>613+ [Co<N02)613' [RhBr6]3' [IfClsl3_ [Pt(NH3 )6]4+ Bra; SF4 SnClg [VFS] [Mn(co)s]‘ [Ru(co)51 [Ni(CN)5]3‘ [Hg01513' IF; SF6 [TlCl6]3‘ [SnBr6]2' [Cr(H20)612+ [Mn<H20)613+ [MnCl6]3‘ [Fe(H20)613+ [Fe(SCN)6]3‘ [C0(NH3 )6]2+ [Comf— [Ni(NH3)6]2+ [Cu(H20)6]2+ [Ti(H20)6]2+ [T113612— [V(H20)5]2+ [V(c0)6] [Cr(CN)6]4' [CrF613‘ [W(co)6] [Mn(CN)6]4‘ [Fe(CN)613" [Ru(NH3)6]2+ [Co(H20)6]3+ [Co(CN)6]4‘ [Rh(NH3 )6]3+ [Irc1612‘ [PtC16]2' £5 5'3 :11 e [Cr(co>512‘ [Tc<co)5]+ [0s(Co)5] C1F5 Bng [MD(NH3 )6]2+ [MnF614‘ [Fe01614‘ [TiCl6]2‘ [Zrc1614‘ [V(CN)613“ [V(N2)6] [Cr(CN)613‘ [Cr(co)6] [Mn<CN>6]3' [Fe(co>612+ [0s016]2“ [Co(CN)6]3' [Rh<CN)6]3' [Ir(NH3)613+ Transition elements in low oxidation states, even negative, can appear to have two more d elec— trons than expected due to promotion of their outer s electrons to the d subshell; typically found with such ligands as CO and CN‘. ...
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CH11_covalent_bond - Models for Covalent Bonding There are...

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