Molecular Orbitals and Bonding: An In-depth Analysis - Prof. John G. Georgiadis, Study notes of Mechanical Engineering

Download Molecular Orbitals and Bonding: An In-depth Analysis - Prof. John G. Georgiadis and more Study notes Mechanical Engineering in PDF only on Docsity! ME 498 –The Molecules of Life Carbon Bonds and Coordination Chemistry Lecture 3 23 January 2007 Outline ►Need to understand the richness of carbon bonding and the coordination ability of trace elements molecules of importance to life ►Ionic and Covalent bonds Hybrization ►p-orbitals (1st row) non-metals and multiple carbon bonds ►Stereochemistry Bond Stability ►d-orbitals (1st row) metals coordination chemistry Electrostatic Potential Maps ( ) red < orange < _ yellow < green <_ blue LiH Ha HF • Carbon shares electrons with other carbon atoms as well as with several different kinds of atoms – it is extremely “sociable” •Need to know how to distribute the electrons COVALENT BONDS Sharing e- according to Molecular Orbitals The 1s- Atomic Orbital of H (solution of Schrodinger eq.) Out-of-phase overlap antibonding MO In-phase overlap bonding MO Orbital Hybridization -(anti) bond Sigma bond (σ) is formed by end-on overlap of two p orbitals A σ bond is stronger than a π bond Orbital Hybridization σ bond Pi bond (π) is formed by sideways overlap of two parallel p orbitals sp3 hybridization accounts for shape of ammonia (trigonal pyramid) water (V-shaped) hydrogen fluoride (linear) • Orbitals not involved in bonding to H are still hybrized and carry lone electron pairs Other sp3 single bonds • Bond angles vary –less symmetry than methane lone-pair electrons are in an sp? orbital \ bong a 04.5° H,O water H Bonding in Water bond is formed by the overlap of an sp? orbital of oxygen with the s orbital of hydrogen ball-and-stick model of water electrostatic potential map for water Bonding in the Methyl Anion lone-pair electrons are in an sp3 orbital bond formed by sp3-s overlap ICH3 methyl anion ball-and-stick model of the methyl anion electrostatic potential map for the methyl anion Structures of Amines Now, N..., N..,,,, CHy— \ H CHy \ CH3 CH” \ CH; methylamine dimethylamine trimethylamine a primary amine a secondary amine a tertiary amine electrostatic potential maps for methylamine dimethylamine trimethylamine Hybrid Orbitals of Ethane sp? single bonds S- oD o* antibonding molecular orbital | sp? atomic sp? atomic orbital orbital o bonding molecular orbital Energy Conformations of Alkanes: Rotation about Carbon–Carbon Bonds The chair conformation of cyclohexane is free of strain very stable cyclic molecule Conformations of Cyclohexane and Energies 12.1 kcal/m 50.6 kJ/m 5.3 kcal/m 6.8 kcal/m 22 kJ/m 28 kJ/m • The bond angle in the sp2 carbon is 120° An sp2-Hybridized Carbon • The sp2 carbon is the trigonal planar carbon Bonding in the Methyl Cation empty p orbital bond formed by sp?-s overlap y H——Ct o *CH angled side view top view methyl cation ball-and-stick models of the methyl cation electrostatic potential map for the methyl cation Polymenisation of alkenes Pp orbitals overlap to form a7 bond Pp orbitals overlap to form a7 bond Pp orbitals overlap to form a7 bond To polymerize the alkene isomers (double bonds) break the π bond between the two sp2 carbons Benzene is unusually stable. Why ? • A planar molecule • Has six identical carbon–carbon bonds • Each π electron is shared by all six carbons • The π electrons are delocalized Delocalized Electrons contribute to stability of molecule CH3 NH2 CH3 CH CH2 localized electrons localized electrons delocalized electronsCH3C O O δ- δ- Drawing Resonance Contributors + ee +7 2 CH;CH, —N <> (CH,CH,—N ‘oO No resonance contributor resonance contributor resonance hybrid π electrons cannot delocalize in nonplanar molecules Benzene Resonance Hybrid Resonance contributors are imaginary, but the resonance hybrid is real Benzene Six m molecular orbitals % % ts Lees eae ogee eemeeesetgtereeeee energy of the p st sth atomic orbitals the energy levels Energy Ye p atomic orbitals of benzene ta Ys Energy _— antibonding rn, 45 wy bonding Bonding in Ethyne: sp triple bond • Bond angle of the sp carbon: 180° • A triple bond consists of one σ bond and two π bonds Table 1.6 Hydrogen-Halogen Bond Lengths and Bond Strengths Bond length Bond strength Hydrogen halide (A) kcal/mol kJ/mol 0.917 136 S71 1.2746 103 432 1.4145 87 366 1.6090 71 298 σ and π bonds - Summary • A π bond is weaker than a σ bond • The greater the electron density in the region of orbital overlap, the stronger is the bond • The more s character, the shorter and stronger is the bond • The more s character, the larger is the bond angle Dipole–Dipole Interaction Dipole–dipole interactions are stronger than van der Waals force but weaker than ionic or covalent bonds A hydrogen bond is a special kind of dipole–dipole interaction The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule Molecular Dipole Moment alcohols share the structure of water ►Low boiling point ethanol=78 oC ethane= -88 oC ►High solubility in water Ethanol infinite ethane 4.7 ml(gas) /100 ml water d- block p - block f- block Elements Essential to Biological Species s- block ELECTRONEGATIVITY p-block diatomic molecules Molecular Orbitals for B2, C2 , N2 ,O2 ,F2 ► MO-8: the highest energy antibonding combination of two pz orbitals, forming a σ* type ► MO-6 and MO-7: the antibonding interaction of two px and two py orbitals, forming two π* type MOs ► MO-4 and MO-5: the bonding interaction of two px and two py orbitals, forming two π type MOs ► MO-3: the bonding interaction of two pz atomic orbitals, forming a σ type MO ► MO-2: the antibonding combination of two s atomic orbitals, forming a σ* type MO ► MO-1: the lowest energy bonding interaction of two s atomic orbitals, forming a σ type MO Fill MOs with Valence Electrons: O2 ► e-: bonding (4 pairs), antibonding (1 pair, and 2 unpaired) ► O2 is a triplet (magn): ► O2 is paramagnetic & an electrophile ► B structure: 2 “half π-bonds” ► LUMO: MO-8 ► HOMO: MO-6 and MO-7 ► Consistent with Hund’s rule ► Two degenerate π* -type MOs (MO-6 and MO-7) split the 2 remaining e-, each with the same spin ► Fill the 5 lowest-in-energy orbitals MO-1 to MO-5 (weak s-p mixing) ► 12 valence electrons (6 from each O atom) Fill MOs with Valence Electrons: CO ► e-: bonding (4 pairs), antibonding (1 pair) ► Resonance Lewis structures ► C-O bond length is only 1.11Å triple bond ► For comparison: C-O single bond ~ 1.43 Å C-O double bond ~ 1.23 Å ► Fill the 5 lowest-in-energy orbitals MO-1 to MO-5 (strong s-p mixing) ► O’s atomic orbitals are slightly smaller than C’s ► 10 valence electrons (4 from C and 6 from O) O Molecular Orbitals of CO σ MO built from 3 p AOs. HOMO of CO. Its largest lobe is on C. CO makes bonds to metal atoms through the C (not O), in many organometallic complexes, where the empty orbital on the metal (LUMO) accepts the two electrons from the HOMO of CO.O http://courses.chem.psu.edu/Chem38/mol-gallery/gallery.html IONIC Crystal Field Theory ► If ions are “point charges”, Electrostatic Potential is proportional to q1 * q2/r ► Fine for large cations of low charge: K+, Na+ ► For transition metal cations becoming coordination centers, d-orbitals are not spherically symmetric tetrahedral, square planar, octahedral orientations http://wwwchem.uwimona.edu.jm:1104/courses/CFT_Orbs.html d orbitals of future coordination center Compound symmetry Crystal Field) d orb/ta/ Splitting Tce @ orbitals Metal ion in spherical aes dhe eryelal field . . Melal iam in tetrahedral crystal field Crystal Field d orbital Splitting Energy Diagrams for 3 stereo-chemistries Transition Metals Ligands (coordination chemistry) ► 5 d-orbitals are degenerate (same energy) when there are no ligands ► When a ligand approaches the metal ion (eg, octahedral Fe 3+), the d-orbitals will “split” in energy ► The size of the gap Δ between the two sets of orbitals depends on several factors, including the ligands (Ligand Field Theory) ► Spectrochemical list of ligands (small Δ to large Δ) I- < Br - < S2- < Cl- < NO3- < N3- < F- < OH- < H2O< CH3CN < py (pyridine) < NH3 < NO2- < CN- < CO [Fe(NO2)6]3- strong field ligand Low spin [Fe Br6]3- weak field ligand Heme centers and Transition Metals ► Heme centers are important to a very important homologous series of life- sustaining proteins Hemoglobins b- or c-cytochromes Fe-S proteins blue-Cu proteins ► Redox pathways in electron-transfer proteins: reactivities correlate with unpaired-spin electrons Fe coordination in Hemoglobin