Prof Gilbert LECTURE 23 CHEM 1211 11 3 10 Last time Chapter 8 Chemical Bonding Polar bonds and electronegativity Polar covalent vs ionic bonds Resonance Formal charges close to zero negative values on the more electronegative elements Exceptions to the octet rule 1 odd electron molecules 2 low Z elements 3 expanded octets This time Chapter 8 Chemical Bonding continued More on expanded octets Many compounds and polyatomic ions with atoms from the 3rd period onwards can accommodate more than an octet because the d orbitals are low enough in energy to participate in bonding Consider SF6 SO4 2 PO4 3 Bond length bond strength and bond order Bond Order the of shared pairs of electrons between atoms For example a double bond has a bond order of 2 Bond order is related to bond length which we can measure Consider these O O bonds Note inverse relationship between bond order and bond length This trend applies to bonds between all the elements Bond Strength average bond energy the average change in enthalpy H that is required to break a mole of bonds of a given type in the gas phase Bond energies vary depending on their molecular environment e g the bond energy of C O bonds is 743 kJ mol but to break the C O bonds in CO2 requires 799 kJ mol The trends in bond strength and bond length are illustrated for C C and C N bonds in the table Calculation Estimating Hrxn from bond energies BE Add up the energy required to breaks the bonds that hold together the molecules of the reactants with the energy released when the bonds in the products form from free atoms NOTE The signs of the H values for breaking bonds are but the signs of the H values for forming bonds are i e H Consider the combustion of methane Hrxn 4 H C H 2 H O O 2 H C O 4 H O H 808 kJ Molecular Geometry Chapter 9 Basic idea each pair of valence electrons wants to be as far from the other pairs as possible this is the Valence Shell Electron Pair Repulsion VSEPR model There are different electron group geometries for atoms with 2 3 4 5 and 6 independent pairs of valence electrons bonds lone pairs on the center atom of the molecules This sum is called the steric number SN of the central atom If all the pairs form single bonds then electron pair geometries molecular geometries Ex BeCl2 BF3 CH4 PF5 SF6 Additional shapes are possible if there are lone pairs of electrons on the central atom For SN 4 Tetrahedral 4 single bonds trigonal pyramidal 3 bonds 1 lone pair Angular 2 bonds 2 lone pairs Ex CH4 Ex NH3 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 permanent dipole 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
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