Covalent bonding uthere are many different approaches to describing covalent bonds which one you pick is often a compromise between ease of use and accuracy umolecular orbital theory is the best approach if good agreement with experiment is needed Lewis theory ulewis theory of bonding was one one the earliest to have any success ubased on the octet rule main group elements like to have eight electrons when they form compounds (except hydrogen) Bond order uin many cases Lewis structures can be used to calculate bond orders that correlate well with experimentally measured bond strengths and lengths :N:::N: triple bond O::O double bond 1
Formal charges ua molecule may have more than one plausible Lewis structure uthe best Lewis structure is the one that has the least charge separation Resonance structures umolecules and ions with more than one distinct but equivalent (by rotation or reflection etc.) Lewis structures can occur the real structure is usually an average of these different resonance forms Failures of the Lewis model ua number of molecules with odd numbers of electrons exist (no octet) e.g. NO uan atom may not have enough electrons to complete its octet without having ridiculous formal charges e.g. BF 3 ua central atom may clearly have more than 8 electrons e.g. SF 6 uo 2 is paramagnetic!! 2
Molecular Orbital theory uin principle, the electronic structure of molecules can be worked out in the same way as for atoms solve the schrodinger equation uthis gives molecular orbitals rather than atomic orbitals uhowever, it is difficult to solve the Schrodinger equation for molecular species LCAO approximation ugood approximations to the molecular orbitals can be obtained by taking linear combinations of atomic orbitals Rules for use of MOs uwhen two AOs mix to give MOs, two MOs will be produced ufor mixing AOs must have similar energies ueach orbital can have two electrons max ufill lowest energy orbitals first uif you have unpaired electrons they should be spin parallel (Hund s rule) ubond order is number of bonding pairs minus number of antibonding pairs 3
Period 1 diatomics Overlapping p-orbitals Oxygen and fluorine 4
Nitrogen Bond orders Experimental verification uuv-photoelectron spectroscopy can be used to verify the MO diagrams. Molecules are ionized using monochromatic light» N 2(g) + h ---> N 2 + (g) + e - The kinetic energy of the resulting photoelectrons is measured 5
PES spectrum for N 2 Heteronuclear diatomics umo diagrams for heteronuclear species are constructed in a similar fashion to those for homonuclear species However, the AO energies are different Molecular shapes uthe prediction of molecular shapes can be done in a number of ways MO theory is very effective. However, it requires complicated calculations VSEPR (Valence Shell Electron Pair Repulsion) theory is good for main group compounds and predictions can be made easily 6
VSEPR uvsepr does not have a solid theoretical foundation uit is based on the idea that pairs of valence electrons (either bonding or lone pairs) will try and avoid each other as much as possible molecule adopts a geometry that allows electron pairs to be as far apart as possible Electron pair count and geometry utwo pairs on central atom - linear uthree pairs on central atom - trigonal ufour pairs on central atom - tetrahedral ufive pairs on central atom - trigonal bipyramidal usix pairs on central atom - octahedral useven pairs - a number of possibilities Electron counting ufor molecules with single bonds to the central atom count all valence electrons of central atom plus one electron for each ligand atom BCl 3 3 electrons from boron and 1 from each chlorine so there are 3 pairs of electrons uwhen there are double bonds present count the four electrons associated with the double bond as a single pair 7
Linear geometry Trigonal geometry Tetrahedral geometry 8
TBP geometry Octahedral geometry Seven coordinate species uif 7, XeF 6, UF 7 2-, NbF 7 2-9
Species that violate VSEPR ulike nearly all truly useful sets of rules there are exceptions uspecies that are sterically crowded often do not obey VSEPR» XeF 6 obeys VSEPR (7 pairs)» TeCl 6 2- does not (7 pairs but is octahedral) Hybridization uthe bonding around atoms with different geometries is often pictures as consisting of overlapping hybrid orbitals a hybrid orbital is a mixture of AOs mixing in different ways gives different geometries Different types of hybridization 10
How useful are hybrid orbitals? uthe use of hybrid orbitals provides a picture of the bonding around an atom uhybridization arguments are not predictive, just descriptive umo theory is predictive but complicated to use An MO view of H 2 O umo theory predicts the lone pairs and bonds in H 2 O are not equivalent unlike the hybridization view Intermolecular forces uattractive interactions between molecules allow the formation of molecular liquids and solids uthere are a number of different types of intermolecular forces dispersion/london/van der Waals forces dipolar interactions hydrogen bonds 11
Dispersion forces uall molecular and atomic species are attracted to each other by dispersion forces ufluctuations in charge distribution polarize nearby atoms and molecules. This induced dipole interacts with the original uneven charge distribution The strength of dispersion forces uthe strength of the interaction depends on number of electrons in atom or molecule spatial extent of atom or molecule uincreasing electron count tends to increase the strength of the interaction more charge to move around uincreasing spatial extent increase the interaction can get better charge separation The BPs of MH 4 species 12
Interactions between polar molecules usome molecules have a permanent dipole moment due to a low symmetry charge distribution uto get low symmetry charge distribution you must have polar bonds uneven distribution of charge between two bonded atoms upolar bonds do not guarantee a polar molecule! Electronegativity usome atoms are better at attracting electrons than others the ability to attract electrons is called electronegativity uelectronegativity was quantified by Pauling Dipole-dipole versus dispersion umolecular dipoles can give a significant contribution to intermolecular forces CO (BP 82K) and N 2 (BP 77K) uhowever, dispersion forces are usually more important HCl (BP 188K) and HBr (BP 206 K) 13
The effect of hydrogen bonding Hydrogen bonding uvery polar molecules (HF, H 2 O, NH 3 ) often have anomalous properties uthe interactions between molecules are strong intermolecular distances drop down into the region where a covalent contribution is to be expected 14