Reading, etc. We will start Chapter 3 today - somematerial comes from Chapter 1. Material related to Chapter 4 beginsnext week – Read all of this is materialon symmetry.. Later half of Chapter 3 - coming latenext week. This is a lot of new stuff formost of you - try to read ahead. e-mail to [email protected] 1.3Electron Configurations,Periodic Properties, & thePeriodic TableFriday, Sept. 3CHEM 462T. HughbanksElectronegativity Pauling: “the power of an atom in a moleculeto attract the electrons to itself.” Mulliken Electronegativity: directly related toIE and EA: = (1/2)[IE + EA] Based on your knowledge of IE and EAvariations, how would you expectelectronegativity to vary in the periodic table?Pauling Electronegativity Assume we know the homonuclear bond dissociationenergies, DAA and DBB. Pauling reasoned that if therewas no electronegativity difference between A and B,then the “ideal” bond energy, DAB, would be the mean:DAB(calc.) = (1/2)(DAA + DBB) Note: in 1937, Pauling switched to using the geometricmean: DAB(calc.) = [DAA•DBB]1/2 Most experimental heteronuclear bond strengths,DAB(exp.), are larger than DAB(calc.), which Paulingthought was due to a stabilizing ionic component tothe bond.Pauling Electronegativity The electronegativity difference between Aand B was defined using the difference in theexperimental and calculated heteronuclearbond energies:A - B = [DAB(exp.) – DAB(calc.)]1/2or AB = [DAB]1/2 (units of DAB = eV)or AB = .102[DAB]1/2 (units = kJ/mol) Arbitrarily, Pauling chose theelectronegativity of Fluorine to be 4.0, fromwhich all other values are obtained bydifferences.ExamplesBond kJ/molH–H 434F–F 158H–F 535H–Cl 404Cl–Cl 242H–Br 339Br–Br 193H–I 272I–I 151DAB = DAB(exp.) – DAB(calc.)DAB(calc.) = [DAA•DBB]1/2DHF = 535 - 262 = 273HF = .102[273]1/2 = 1.69DHCl = 404 - 324 = 80HCl = .102[80]1/2 = 0.91DHBr = 339 - 289 = 50HBr = .102[50]1/2 = 0.72DHI = 272 - 256 = 16HI = .102[16]1/2 = 0.41Electronegativity Values are approximate, scale arbitrary Highest electronegativity: F,  = 4.0 Lowest electronegativity: Cs,  = 0.7 Generally, electronegativity increasesas you move up or to the right in theperiodic table.Electronegativity Difference In a purely covalent bond, the 2 atomsare identical: H2, N2, etc.– same electronegativity even sharing In an ionic bond, one atom has highelectronegativity, one low: NaCl(Na) = 0.9, (Cl) = 3.0  = 2.1Chlorine pulls an electron away fromsodium, forming ions.Polar Bonds For most bonds,  is moderate, notzero This gives an intermediate case:electrons are shared, but not equally. CO: (C) = 2.5, (O) = 3.5   = 1.0 Bond is “polar covalent.”Lewis Structures Easy, useful way of representingvalence electrons in a molecule(compared to the real physics). One electron = one dot; examples: One pair of shared electrons = one line Two pairs = two lines, etc.H......F....C..Writing Lewis Structures There are systematic methods fordoing this, which I will use followloosely. Various texts put different emphasis onthe “octet rule” for identifying ‘stable”Lewis structures. The octet rule isused to a greater extent whenconsidering molecules involving first-row atoms. For a simple scheme, Lewis structures(including Pauling’s “resonance” ideas)are powerful - but they are still just amodelLewis StructuresSystematic method1. Treat ions separately.2. Count the valence e-’s.3. Set up the bonding framework, usingtwo e-’s per bond4. 3 pairs of nonbonding e-’s on eachouter atom, except H (assumingenough e-’s)5. Remaining e-’s to inner atomsLewis Structures, cont.Systematic method6. Find formal charge on each atom.7. Minimize formal charges by shifting e-’sto make double and triple bonds.(a) 2nd row atom  4 occupied valenceorbitals (8e-’s  “octet rule”)(b) other atoms  formal charge to zero.Formal Charges A useful “accounting device,” not thereal charge on the atoms (because e–sin bonds not equally shared).FC = (# valence e-’s in free atom) - (# valence e-’s assigned in structure) Sum of FC’s = zero for a neutralmolecule, or total charge on an ion. Minimize FC’s to get “best” structure.Examples - no octet rule violationsCH4 (methane)C2H6 (ethane)CCl4 (carbon tetrachloride)Br2, O2, N2 (bromine, oxygen, nitrogen)H2O, NH3 (water, ammonia)C2H4, C3H6 (ethene, propene)HCOOH (formic acid)(NH2)2CO (urea)Lewis Structures & Resonance In many cases, no single Lewisstructure adequately represents thedistribution of electrons in a molecule.In such cases, we represent theelectron distribution as a combination ofLewis structures. Real molecule does NOT “bounce”between the different resonancestructures!Resonance — Examples equivalent resonance structures:O3(ozone), NO2-, NH4NO3, CaCO3, C6H6(benzene) resonance structures are inequivalent, butat least two are important:N2O (nitrous oxide), NCO- (cyanate),CH3CONHCH3(N-methylacetamide)— anexample of an amide bond.Bond Lengths - an experimental testN ON+-N ON+-1.129 1.188 (Å)NN1.094 ÅN OO+N2NO2+O =N = 1.150 Å (both)N2OAdding hybridization; COCl2 vs SOCl2Existence of double bondimplies a "unhybridized" porbital engaged in  bonding,therefore... C atom geometrymust be trigonal-planarClCClOsp2ClSClOsp3In the structure obeying octetrule, there’s no S=O (double)bond, implying sp3hybridization at S andpyramidal geometry (tetrahedralincluding the lone-pair) aroundS.More on SOCl2Lone-pairs on Cl are understood -and omitted here.A common depiction, with an S=Obond. This makes some inorganicchemists happy because it reducesthe formal charges but it confusesstudents who don't understand thatthe double bond involves use of 3dorbitals on S (which is questionableanyway).ClSClOsp3ClSClOOctet ‘Violations’ Electron “deficient” molecules (e.g., BF3) Even more serious: B2H6 - completefailure of classical structure theory“There are no electron deficient molecules,only theory deficient chemists.” – K. Wade “Hypervalence” - e.g., PCl5, SF4 Though not necessary, non-octetstructures are often drawn - e.g., SO3,SO42-Some Properties of Bonds; BondDissociation EnthalpiesA–B (g)  A(g) + B(g) H > 0H is the bond dissociation enthalpyCH4 (g)  C(g) +