January 20th 2017 Chapter 11 HW Chapter 11 and 12 Due Feb 3rd iClicker SLW 212 meets next week Recall Molecular Geometry A summary of common molecular shapes with two to six electron groups 11 1 Valence Bond VB Theory and Orbital Hybridization The basic principle of VB theory A covalent bond forms when the orbitals of two atoms overlap and a pair of electrons of opposite spins occupy the overlap region Extent of orbital overlap depends on orbital shape and direction Figure 11 1 Orbital overlap and spin pairing in H2 Orbital overlap and spin pairing in F2 11 1 Valence Bond VB Theory and Orbital Hybridization What about bonding when atoms are not identical How can we account for molecular shapes The orbitals that form when bonding occurs are different from the atomic orbitals in the isolated atoms Atomic orbitals mix or hybridize when bonding occurs to form hybrid orbitals and type of hybrid orbitals formed and type of atomic orbitals mixed s p d f Formation and orientation hybrid orbitals Example BeCl2 Example BeCl2 Be Z 4 He 2s2 2Cl Z 17 Ne 3s23p5 BeCl2 sp hybridized Cl Be Hybridized Be sp Cl 2Cl BeCl2 2p 3s 3p Covalent bonds form from overlap of sp and p orbitals Formation and orientation hybrid orbitals F Example BF3 sp2 hybridized B F F Hybridized B atom Isolated B atom Covalent bonds form from overlap of sp2 and p orbitals Formation and orientation hybrid orbitals Example CH4 H Covalent bonds form from overlap of sp3 and s orbitals H sp3 hybridized C H H Formation and orientation hybrid orbitals Example PCl5 Cl Cl P sp3d hybridized Cl Cl Cl Covalent bond forms from overlap of sp3d and p orbitals Formation and orientation hybrid orbitals F Example SF6 F sp3d2 hybridized F F S F Covalent bond forms from overlap of sp3d2 and p orbitals F Formation and orientation hybrid orbitals What is the hybridization of the central atom in the following molecule A sp B ClF3 F sp2 Cl C sp3 D sp3d E sp3d2 F F 5 electron groups attached to central atom 3 bonding pairs 2 lone pairs Hybrid orbitals are formed by mixing one sorbital three p orbitals and one d orbital to create five sp3d hybrid orbitals 11 2 The Mode of Orbital Overlap and the Types of Covalent Bonds Orbitals can overlap by two modes to form bonds end to end or side to side which gives rise to two different types of covalent bonds sigma bonds and pi bonds respectively Consider the molecular shapes predicted by VSEPR theory for ethane C2H6 ethylene C2H4 and acetylene C2H2 H H H C C H H H H H C H H C H C C H End to End Overlap and Sigma Bonding Ethane tetrahedral H H H C C H H Sigma Bond greatest electron density along bond axis free rotation about bond All single bonds are bonds tetrahedral H both C are sp3 hybridized s bond formed by s sp3 overlap End to end sp3 sp3 overlap to form a s bond Side to Side Overlap and Pi Bonding Trigonal planar Ethylene H H C H unhybridized 2p orbitals C H sp2 sp2 overlap in one position s Pi Bond two regions lobes of electron density above and below sigma bond restricted rotation Any double bond consists of one bond and one bond Side to Side Overlap and Pi Bonding Acetylene linear H C sp sp overlap in one position s linear C H Each C is sp hybridized and has two unhybridized p orbitals Triple bond consists of 1 bond and 2 bonds no rotation Orbital Overlap in Single and Multiple bonds Orbitals can overlap by two modes to form bonds 1 End to end Sigma Bond Free rotation electron density along bond axis all single bonds are sigma bonds 2 Side to side Pi Bond Restricted rotation electron density above and below sigma bond 11 3 Molecular Orbital MO Theory and Electron Delocalization VSEPR describes molecular shape VB theory describes orbital overlap Molecular orbital MO theory explains magnetic and spectral properties and electron delocalization MO model is a quantum mechanical treatment for molecules where molecules have molecular orbitals MOs of given energies and shapes that are occupied by the molecule s electrons MO theory pictures molecules as a collection of nuclei with orbitals delocalized over the whole molecule and occupied by electrons Formation of Molecular Orbitals The combination of orbitals to form bonds is viewed as the combination of wave functions Atomic wave functions AOs combine to form molecular wave functions MOs Addition of AOs forms a bonding MO which has a region of high electron density between the nuclei Subtraction of AOs forms an antibonding MO which has a node or region of zero electron density between the nuclei Shape and Energy of H2 MOs Example H2 The bonding MO is always lower in energy and the antibonding MO is always higher in energy than the AO s combined to form them Figure 11 14 Both the bonding and antibonding MO s of H2 are sigma s MOs because they are cylindrically symmetric about the imaginary axis running between the two nuclei bonding s and antibonding s Electrons in MOs MOs are filled in order of increasing energy aufbau principle A MO can hold a maximum of 2 electrons with opposite spins exclusion principle Orbitals of equal energy are half filled with parallel spins before filled Hund s rule MO diagram for H2 MO bond order e in bonding MO e in antibonding MO Example Bond order for H2 2 0 1 Bond order Bond order 0 molecule more stable than separate atoms forms Bond order 0 molecule stable as separate atoms does not form MO diagram for He2 MO diagram for He2 1s 2 1s 2 1s 2 1s 1 Bond order 2 2 0 Bond order 2 1