Transcript Chapter1
Orbital Hybridization and Molecular Orbitals Chapter 9 Historical Models G.N.Lewis and I. Langmuir - laid out foundations for molecular structures - Ionic species were formed by electron transfer -Covalent molecules arise from electron sharing cannot predict molecular geometries? Valence shell electron pair repulsion theory (VSEPR) - predicts molecular shapes based on valence electrons - uses the Lewis dot structures and electron repulsions to construct models how do orbitals allow these geometries? Valence bond theory (VB) - a molecule arises from interaction (overlap) of complete atoms - binding through localized overlap of half filled valence-shell atomic orbitals retaining their original character (s, p, d, f) could not explain experimentally determined molecular properties? Molecular orbital theory (MO)- a more complete theory - a molecule is formed by the overlap of atomic orbitals to form molecular orbitals - electrons are then distributed into MOs -A molecule is a collection of nuclei with the orbitals delocalized (spread out) over the entire molecule Valence Bond Theory Sigma Bond Formation Two s orbitals overlap Two p orbitals overlap Valence Bond (VB) Theory Covalent bonds are formed by the overlap of atomic orbitals -Atomic orbitals on the central atom can mix and exchange their character with other atomic orbitals in a molecule -This process is called hybridization -Hybrid Orbitals: create geometries with the same shapes as seen in VSEPR theory Valence Bond Theory Linus Pauling -valence electrons are localized between atoms (or as lone pairs) -half-filled atomic orbitals overlap to form bonds Valence Bond (VB) Theory Regions of High Electron Density (“things” attached) 2 3 Electronic Geometry Hybridization Linear Trigonal planar sp sp2 4 5 Tetrahedral Trigonal bipyramidal sp3 sp3d 6 Octahedral sp3d2 Molecular Shapes and Bonding In the next sections we will use the following terminology: A = central atom B = bonding pairs around central atom U = lone pairs around central atom Linear Electronic Geometry: AB2 Some examples of molecules with this geometry are: BeCl2, BeBr2, BeI2, HgCl2, CdCl2 -All of these examples are linear, nonpolar molecules -Important exceptions occur when the two substituents are not the same BeClBr or BeIBr will be linear and polar Linear Electronic Geometry: AB2 1-p 1-p Trigonal Planar Electronic Geometry: AB3 Some examples of molecules with this geometry are: BF3, BCl3 -All of these examples are trigonal planar, nonpolar molecules -Important exceptions occur when the three substituents are not the same BF2Cl or BCI2Br will be trigonal planar and polar! Trigonal Planar Electronic Geometry: AB3 Trigonal Planar How to account for 3 bonds 120o apart using a spherical s orbital and p orbitals that are 90o apart? -Pauling said to modify VB approach with ORBITAL HYBRIDIZATION -mix bonding orbitals to form a new set of orbitals -HYBRID ORBITALS : give the maximum overlap in the correct geometry Valance Bond Theory 2p 2s hydridize orbs. 2 rearrange electrons three sp hybrid orbitals unused p orbital Valence Bond Theory The three hybrid orbitals are made from 1-s orbital and 2-p orbitals 3-sp2 hybrids Now we have 3, half-filled HYBRID orbitals that can be used to form B-F sigma bonds Valance Bond Theory Bonding in BF3 •• •• F •• •••• F •• B Boron configuration ••• F• 1s •• 2s 2p planar triangle angle = 120o Tetrahedral Electronic Geometry: AB4 Some examples of molecules with this geometry are: CH4, CF4, CCl4, SiH4, SiF4 -All of these examples are tetrahedral, nonpolar molecules -Important exceptions occur when the four substituents are not the same CF3Cl or CH2CI2 will be tetrahedral and polar Tetrahedral Electronic Geometry: AB4 What’s the electron configuration of carbon? How many free electrons does it have? How many bonds can it form? Tetrahedral Bonding How do we account for 4 C—H sigma bonds 109o apart? -Need to use 4 atomic orbitals -s, px, py, and pz to form 4 new hybrid orbitals -hybrid orbitals point in the correct orientation 109o Tetrahedral Electronic Geometry: AB4 Tetrahedral Electronic Geometry: AB4 sp3 Tetrahedral Electronic Geometry: AB3U Some examples of molecules with this geometry are: NH3, NF3, PH3, PCl3, AsH3 -These molecules are examples of central atoms with lone pairs of electrons -Thus, the electronic and molecular geometries are different -All three substituents are the same but molecule is polar -NH3 and NF3 are trigonal pyramidal, polar molecules Steps in predicting the hybrid orbitals used by an atom in bonding: 1. Draw the Lewis structure 2. Determine the electron pair geometry using the VSEPR model 3. Specify the hybrid orbitals needed to accommodate the electron pairs in the geometric arrangement NH3 1. Lewis structure 2. VSEPR indicates tetrahedral geometry with one non-bonding pair of electrons (structure itself will be trigonal pyramidal) 3. Tetrahedral arrangement indicates four equivalent electron orbitals Tetrahedral Electronic Geometry: AB2U2 Some examples of molecules with this geometry are: H2O, OF2, H2S -These molecules are examples of central atoms with two lone pairs of electrons -Thus, the electronic and molecular geometries are different -Both substituents are the same but molecule is polar. -Molecules are angular, bent, or V-shaped and polar Orbital Hybridization Bonds Shape (“Things”) Hybrid orbital Remaining orbitals 2 linear sp 2 p’s 3 trigonal planar sp2 1p 4 tetrahedral sp3 none Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 Some examples of molecules with this geometry are: PF5, AsF5, PCl5, etc. -These molecules are examples of central atoms with five bonding pairs of electrons -The electronic and molecular geometries are the same -Molecules are trigonal bipyramidal and nonpolar when all five substituents are the same -If the five substituents are not the same polar molecules can result, AsF4Cl is an example. Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 Some examples of molecules with this geometry are: SF6, SeF6, SCl6, etc. -These molecules are examples of central atoms with six bonding pairs of electrons -Molecules are octahedral and nonpolar when all six substituents are the same -If the six substituents are not the same polar molecules can result, SF5Cl is an example. Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 Molecular Geometry Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 Molecular Geometry Polarity Carbon-Carbon Double Bonds C atom has four valence electrons -Three electrons from each C atom are in sp2 hybrids -One electron in each C atom remains in an unhybridized p orbital 2s 2p three sp2 hybrids 2p C Double Bonds The single 2p orbital is perpendicular to the trigonal planar sp2 lobes (containing 3 e-) -The fourth electron is in the p orbital Side view of sp2 hybrid with p orbital included Double Bonds An sp2 hybridized C atom has this shape -Remember there will be one electron in each of the three lobes Top view of an sp2 hybrid Double Bonds The portion of the double bond formed from the head-on overlap of the sp2 hybrids is designated as a s bond p p Triple Bonds Ethyne or acetylene, C2H2, is the simplest triple bond containing organic compound -Compound must have a triple bond to obey octet rule Triple Bonds Lewis Dot Formula VSEPR suggests regions of high electron density are 180o apart Carbon-Carbon Triple Bonds Carbon has 4 electrons -Two of the electrons are in sp hybrids -Two electrons remain in unhybridized p orbitals C [He] 2s 2p 2-sp hybrids H-C=C-H 2p Triple Bonds A s bond results from the head-on overlap of two sp hybrid orbitals py pz Triple Bonds The unhybridized p orbitals form two p bonds -Note that a triple bond consists of one s and two p bonds - py - pz Summary of Geometries Molecular Orbital (MO) Theory VSEPR and VB theory are good to explain the molecular shape -BUT they did not explain the magnetic or spectral properties of molecules -Molecular orbital theory does! MOs are derived from the addition and subtraction of atomic orbitals represented as wavefunctions to form molecular orbitals There are two possible combinations: -Adding two atomic orbitals forms a bonding MO -Subtracting two atomic orbitals forms an antibonding MO Note: The number of atomic orbitals contributed equals the number of molecular orbitals generated Molecular Orbitals Consider H2 Molecular Orbital -Bonding MO is lower in energy than the unbound atomic orbitals -Antibonding MO is higher in energy Molecular Orbital Diagram Electrons fill according to Hund’s rule Molecular Orbital Theory Bond Order (B.O.) B.O. = 1/2 (B - AB) B = bonding electrons AB = antibonding electrons Molecular Orbital Theory Example: He2 MO Diagram for He2+ and H2- AO of He s*1s AO of He+ s1s MO of He2+ He2+ bond order = ?? Energy s*1s AO of AO of H Hs1s MO of H2 - H2- bond order = ?? MO Diagram for He2+ and H2- s*1s 1s 1s MO of He2+ 1s 1s s1s s1s AO of He Energy s*1s AO of He+ He2+ bond order = 1/2 AO of H- MO of H2 - AO of H H2- bond order = 1/2 Orbital Interaction for Li2 Molecule Li atom - 1s22s1 s*2s Li 2s Li 2s s2s Bond order for Li2? s*1s Li 1s Li 1s s1s Be2? Li2+? p Orbital Bonding p Orbital Bonding O2 molecule is an example with sigma and pi bonds forming between atoms -MO theory predicts that oxygen will be paramagnetic ½ ( 8 - 4) = 2