Chapter 9 Bonding and Molecular Structure Orbital Hybridization and Molecular Orbitals 9 1 Orbitals and Theories of Chemical Bonding 2 common approaches to rationalizing bonding based on orbitals 1 Valence bond VB theory developed largely by Linus Pauling 2 Molecular Orbital MO theory developed by Robert Mulliken Valence Bond Theory Premise Bonding involves valence electrons Half filled atomic orbitals on bonding atoms overlap to form bonds Bonds are localized between atoms or as lone pairs Leads to prediction of molecular shape The idea that bonds are formed by overlap of atomic orbitals is the basis for valence theory Molecular Orbital Theory Bonding electrons reside in Molecular Orbitals that arise from the original atomic orbitals of the bonding atoms Bonding electrons are spread over the entire molecule delocalized Explains paramagnetism and electronic spectroscopy in molecules The only theory that describes bonding accurately in NO and O2 9 2 Valence Bond Theory The Orbital Overlap Model of Bonding Orbital overlap increases probability of finding the bonding electron in the region of space between the 2 nuclei The main points of the valence bond approach to bonding are Orbitals overlap to form a bond between two atoms Overlapping orbitals hold two electrons of opposite spin Usually one electron is supplied by each of the two bonded atoms The bonding electrons are localized with a higher probability of being found within a region of space between the bonding nuclei Both electrons are simultaneously associated with both nuclei Sigma bond A sigma bond is a bond in which electron density is greatest along the axis of the bond Describes a single bond Fig 9 2 p 403 Hybridization of Atomic Orbitals H H C H H Methane show orbital box diagram and hybridization Fig 9 4 p 406 Valence Bond Theory for Ammonia NH3 VB of ammonia Valence Bond Theory for Water H2O Hybrid Orbitals for Molecules and Ions with Trigonal Planar Electron Pair Geometries Examples BF3 O3 NO3 CO32 H H C H H e d e geo mol geo how many bonds do they need hybridization Fig 9 5 p 409 Hybrid Orbitals for Molecules and Ions with Linear Electron Pair Geometries Examples BeCl2 acetylene CHCH CO2 H H C H H Fig 9 6 p 409 Hybrid Orbitals for Molecules and Ions with Trigonal Bipyramidal or Octahedral ElectronPair Geometries d orbital Participation Examples trigonal bipyramidal PF5 AsF5 PCl5 Octahedral SF6 SeF6 SCl6 Trigonal Bipyramidal electronic arrangement sp3d AB5 AB4U AB3U2 and AB2U3 If lone pairs are incorporated into the trigonal bipyramidal structure there are three possible new shapes 1 One lone pair Seesaw shape 2 Two lone pairs T shape 3 Three lone pairs linear H H H C H H H H C H The lone pairs occupy equatorial positions because they are 120o from two bonding pairs and 90o from the other two bonding pairs Results in decreased repulsions compared to lone pair in axial position H H H C H 19 Octahedral electronic arrangement sp3d2 AB6 AB5U and AB4U2 If lone pairs are incorporated into the octahedral structure there are two possible new shapes 1 One lone pair square pyramidal 2 Two lone pairs square planar The lone pairs occupy axial positions because they are 90o from four bonding pairs Results in decreased repulsions compared to lone pairs in equatorial positions H H H C H H H C H H 20 Table 8 3a p 316 Table 8 3b p 316 Table 8 3c p 317 Table 8 4 p 321 Fig 9 3 p 405 Pi bond Used for double or triple bonds Above and below the sigma bond show ethene and acetylene Double bond 1 plus 1 Triple bond 1 plus 2 Multiple Bonds Double Bonds ethene Triple bonds Acetylene ethyne CO N2 Acetic acid Reminder Double bond 1 plus 1 Triple bond 1 plus 2 Cis trans isomerism A Consequence of Bonding Benzene A Special Case of Bonding 9 3 Molecular Orbital Theory MO theory assumes that pure atomic orbitals of the atoms in the molecules combine to produce orbitals that are spread out or delocalized over several atoms or even over an entire molecule Molecular Orbital Theory MO approaches bonding between atoms from a different approach than Valence Bond Theory In Valence Bond theory the atomic orbitals of a bonding atom mix or hybridize to form localized bonds that take on the electronic geometry predicted by VSEPR theory In MO theory the atomic orbitals are treated like waves that constructively or destructively add to form new Molecular Orbitals The electrons of the molecule are distributed over the entire molecule as a whole delocalized Molecular Orbital Theory has several advantages and differences over VSEPR VB theory MO does a good job of predicting electron pair spectra and paramagnetism where VSEPR and the VB theories don t MO theory like VB theory predicts the bond order of molecules however it does not need resonance structures to describe molecules The main drawback to our discussion of MO theory is that we are limited to talking about diatomic molecules molecules that have only two atoms bonded together or the theory gets very complex Molecular Orbitals for H2 Bonding molecular orbital 1s Antibonding molecular orbital 1s means antibonding Bonding molecular orbital is lower in energy than antibonding molecular orbital Fig 9 14 p 418 Principles of Molecular Orbital Theory First principle of molecular orbital theory is that the total number of molecular orbitals is always equal to the total number of atomic orbitals contributed by the atoms that have combined H2 Second principle of molecular orbital theory is that the bonding molecular orbital is lower in energy than the parent orbitals and the antibonding orbital is higher in energy He2 explains why it doesn t exist Third principle of molecular orbital theory is that the electrons of the molecule are assigned to orbitals of successively higher energy according to the Pauli exclusion principle and Hund s rule Li2 Be2 B2 N2 N2 MO Fig 9 17b p 422 Molecular Orbitals from Atomic p orbitals P block atom has two p orbitals in planes perpendicular to the bond connecting the two atoms so give two bonding molecular orbitals and two antibonding molecular orbitals O2 Bond Order Is the net number of bonding electron pairs linking a pair of atoms Bond order of e in bonding MOs of e in antibonding MOs H2 He2 Li2 He2 Bond length and bond strength depend on the bond order The more bonding electrons versus antibonding electrons shared between atoms the more tightly bound they are This leads to stronger shorter bonds O2 O2 O2As Bond Order increases bond length increases and bond strength decreases get