Chapter 8: Advanced Theories of Covalent Bonding Notes PDF

Chapter 8 Advanced Theories of Covalent Bonding 8.1 Valence Bond Theory Covalent bonds form When electrons are shared between two atoms...

Chapter 8 Advanced Theories of Covalent Bonding 8.1 Valence Bond Theory Covalent bonds form When electrons are shared between two atoms · C.: 4 bonds Electrons in overlapping orbitals must have opposite spins - Number of bonds formed by an atom is determined by the #of unpaired electrons ↳ : 2 bonds These facts explain: the bonding in diatomic molecules with only single bonds the lack of bonding experienced by the noble gases. Orbital Overlap orientation of orbitals: want greatest amt of overlap type of orbitals: creates different type of bond nucleus b b S - Ss - p P P - P - P (head on Sigma (s) and Pi Bonds (p) (sideways) Single bond I Sigma : bond (H- ) : Double bond : I sigma and 1 pibond (:0 = 0: ) Triple bond i : sigma and z pi bonds (iNEN :) Sigma (s) Bonds single bonds e- density is between the 2 nuclei S-s, s-p, or hybrid Orbital overlap P-p orbital head-to-head overlap z regions Pi (p) Bonds - 2 & 3 bonding double or triple bond - I bond e-density above and below the 2 nuclei P-p Orbital side-to-side overlap Example: Count s and p bonds 2 Th 9 0 Limitations of the Valence Bond Theory Valence bond theory must be modified to explain the covalent bonds formed in other molecules. Ex. NH3 1) ↓ IU Ire [s22p3 2 ? N Is : Is 3H : 1s  8.2 Hybrid Atomic Orbitals mixing of two or more nonequivalent atomic orbitals (ex. s and p) to form newly shaped orbitals # of hybrid orbitals is equal to # of pure atomic orbitals used in the hybridization process Named from the atomic orbitals that are combined iSp , Sp2 Sp3 , , spid , Sp3d.... ↓ ↳1s sp Hybrid Orbitals Is + 1p + 3p + 2d · 180 ; linear geometry - · two p Orbital remain one s Orbital + one p Orbital · central atom surrounded by two regions of electron Orbitals density ↳ two sp hybrid sp2 Hybrid Orbitals occurs when Central atom Surrounded by · three regions of electron density 120 ; trigonal · planar geometry · one p Orbital remains orbitals ones orbital +two p ↳ three sp2 hybrid Orbitals sp3 Hybrid Orbitals surrounded by four regions · central atom · 109 5 °; tetrahedral. geometry · zeroP Orbitals remain Y eX C , Orbital H H one p orbital+ three p - - I ↳ 4 Sp3 hybrid Orbital H One + three p + sp3d Hybrid Orbitals sp3d2 Hybrid Orbitals Twod -6 sp3d2 one so three p + one d trigonal bipyramid ↳ Fue spod Orbitals · 90' octahedral  single one region bonds - Hybridization Summary double of e-density triple What is the hybridization of the selected atom? -3 + Sp P 0 T > P - = , - & [ ↑ 4 + sp3 3 - Sp7 Ogoss e & Oesp3 & Se sp3d 8 go 8.3 Multiple Bonds Single bonds (sigma, s) are formed by the direct overlap of: 1. two hybrid Orbitals 2. porbitals (head-to-head) -- 3. s Orbital S 4. Combo of above 2nd & 3rd bonds (pi, p) are formed by side-by-side overlap of: 1. two regular pl atomic orbitals 8 sorbital of + top ↓ right % 2π = 90 bottom bond : (0) Csp2-Osp2 making identify orbitals sp3 bond + it is Sp3 ↳ o always & Spe ! ↳ (π)p bond < P P-P - O bond 7 q -Osp3-Hs < E O bond bond So 3 Ssp3-Csp3 3 Csp3-Nsp 109 5. : - 109 5.. Delocalized Orbitals the electrons in the # bonds are not located in One set of p orbitals, but rather delocalized throughout the molecule O - Orbital s can be - C Overlapping C anywhere Paramagnetism: phenomenon in which a material is not magnetic itself but is attracted to a magnetic field; it occurs when there are unpaired electrons present Diamagnetism: phenomenon in which a material is not magnetic itself but is repelled by a magnetic field; it occurs when there are only paired electrons present 8.4 Molecular Orbital Theory model that describes the behavior of electrons delocalized throughout a molecule in terms of the combination of atomic wave functions Comparison of Bonding Theories Valence Bond Theory Molecular Orbital Theory considers bonds as localized between considers electrons delocalized throughout the entire molecule one pair of atoms creates bonds from overlap of atomic orbitals (s, p, d…) and hybrid combines atomic orbitals to form molecular orbitals (σ, σ*, π, π*) orbitals (sp, sp2, sp3…) creates bonding and antibonding interactions based on which orbitals forms σ or π bonds are filled predicts molecular shape based on the number of regions of electron predicts the arrangement of electrons in molecules density needs multiple structures to describe resonance Combining Atomic Orbitals into Molecular Orbitals - & constructive and destructive interference - - - o - & - bonding and antibonding molecular orbitals constructive destructive * anti bonding antibonding high energy = , : bonding: low energy sigma & sigma*, for s and px  pi & pi*, for py and pz C density anti bonding : on outside of bond Molecular Orbital Energy Diagrams combined atomic Orbitals # of molecular orbitals = # of Mo always more stable (lower E) than anti bonding bonding Mo each Orbital Spin fill 2e of opposite fill low-energy Orbitals first (outbau principle) Mos of the same energy hunds when electrons to use rune adding energy A Molecular Orbital Energy Diagrams for H2 – N2 Molecular Orbital Energy Diagrams for O2 – Ne2 fipped valence - - e ↑ *>4 - T H ↑ H at om atom molecule MO sequence MO sequence * * * * * (s1s)(s 1s)(s2s)(s 2s) (p2py p2pz)(s2px)(p 2py p 2pz)(s 2px) (s1s)(s 1s)(s2s)(s 2s)(s2px)(p2py p2pz)(p*2py p*2pz)(s*2px) * * Nile- Superscript sum = e# 0 : 2e T : 4e- Bond Order ( # ote- Bond order = ↓ ( - in antibonding MOS Higher the bond order, stronger the bond Bond order = 0, no bond exists ↳ no negatives exsts Examples He2 > - 42 - N2+ > - 13e 1 & ↑ ↑ 1 ↓N - 1V X 1 -v X 1 ↑ Pl & & v N M N + He He N2 + Hez 50 % (0 s)(0 *, s)(02s))0 zs)(2py))πzpz))02px) * 0. -4 , O unpairede => diamagnetic I unpairede = paramagnetic BO : (2-2) = 0 iDNE Bo : (9 4) - = 25. 5 B2- I1 > - es O2 & - 1 14 & IV ↑v i 1 1V 1 1 ↑L 11 - 12 ↑L X -V -V 11 - O - O B B- (0 (0 *. ](02(02s(Thapy) I unpaired -> Paramagnetic Bo : (7-4) = 1 3. Aktiv Ch 8. 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