Prof Gilbert LECTURE 24 CHEM 1211 11 4 10 Last time Chapter 8 Chemical Bonding Bond length bond strength energy and bond order the higher the order the shorter and stronger the bond Estimating Hrxn from bond energies Molecular Geometry Chapter 9 based on minimizing repulsions between pairs of valence shell electrons by keeping them as far apart as possible Steric numbers define electron pair geometries which in turn define molecular geometries shapes This time Chapter 9 Molecular Geometry continued Ex BeCl2 BF3 CH4 PF5 SF6 These electron pair geometries correspond to the same molecular geometries shapes when there are no lone non bonding pairs of electrons around the center atom If there are lone pairs then other geometries are possible For example if SN 4 Tetrahedral 4 single bonds trigonal pyramidal 3 bonds 1 lone pair Ex CH4 Ex NH3 Angular 2 bonds 2 lone pairs Ex H2O Bond angles 109 5 107 104 5 For SN 3 trigonal planar 3 bonds 1 bond angular 2 bonds 1 bond one lone pair Ex CH2O Ex O3 For SN 2 we get straight molecules 2 double bonds Ex CO2 For SN 5 For SN 6 For trigonal bipyramid any lone pairs will be in the base plane at equatorial positions where 2 other pairs are 120 away and 2 others are 90 away rather than in an axial position where 3 pairs are only 90 away Polar molecules and dipole moments The extent of separation of charge along polar bonds as in H F or within any polar molecule is expressed by the molecule s dipole moment Q r where Q is the partial charge that is separated by a distance r inside the molecule The value of in debye D can be determined experimentally and r is based on the distance between charge centers and the geometry of the molecule To calculate Q a unit conversion is required 1 D 2 24 10 30 coulomb meters The value of Q when divided by the charge on an electron 1 602 10 19 C yields the ionic character of a bond or polar substance Key concept To have a dipole moment a molecule must have 1 polar bonds and 2 they must be arranged asymmetrically so that their unequal sharing of electrons do not cancel out They do cancel out in CO2 CH4 SF6 But they don t in H2O or in HF Bond vibrations and global warming Covalent bonds vibrate stretch rotate and twist like springs When they vibrate they may create fluctuating electrical fields that interact with electromagnetic radiation i e infrared radiation Infrared inactive vibration Infrared active vibration The problem the left side vibrations can trap heat energy radiated by Earth s surface Valence Bond Theory Valence bond theory assumes that atomic orbitals mix hybridize among themselves to form orbitals with the shapes that give the observed electron and molecular geometries Hybridized orbitals atomic orbitals of different shapes and energies mix together Each has a major lobe that may overlap with a valence orbital on another atom to form a bond Ignore the minor lobe The carbon atom in CH4 that form 4 bonds to 4 H atoms Not every hybrid orbital forms a covalent bond filled hybrid orbitals remain lone pairs If a bond is needed a p orbital is left unhybridized so that it can overlap sideways with the corresponding p orbital on another atom such as the O in CH2O If a second bond is needed need to leave two p orbitals unhybridized as in acetylene H C C H Similarly in CO2 Expanded octets hybridization schemes need to include d orbitals as in SF6 Inquiry What are the hybridization schemes in HCN benzene SO4 2 and PO4 3 ions
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