Chapter 9Chemical Bonding§9.2 Types of Chemical Bonds§9.3 Lewis Dot Structures for Atoms§9.4 How to Quantify Bond Strengths§9.4 Lattice Energy§9.4 Born-Haber CycleSlide 8§9.4 Coulomb’s LawSlide 10Slide 11§9.5 Covalent Bonds§9.6 Polar Covalent Bonds§9.6 Ionic-Covalent Bond Continuum§9.6 Electronegativity§9.6 Electronegativity Trends§9.6 Ionic vs Covalent CompoundsSlide 18Slide 19Slide 20§9.6 Dipole MomentSlide 22Slide 23§9.7 Drawing Lewis Structures§9.7 Lewis Dot Structures§9.7 Drawing Lewis Dot Structures§9.8 Resonance§9.8 Benzene Resonance§9.8 Formal ChargeSlide 30§9.8 Formal Charge (FC)§9.9 Octet Rule Exceptions§9.9 Octet Rule Exceptions§9.9 Nitrogen Dioxide§9.10 Bond Energies§9.10 Average Bond Energies§9.10 DHrxn from Bond EnergiesSlide 38§9.10 Average Bond LengthsSlide 40NI3 Demo§9.8 Resonance - Formal ChargeSlide 43Slide 44CHEMICAL BONDING I: LEWIS THEORYChapter 9Chemical BondingChapter 9 and 10 each cover a different theory on the chemical bond.Both models are far from perfect, but where one fails the other succeeds.§9.2 Types of Chemical BondsCovalent bond – sharing of electrons between two non-metal nucleiIonic bond – transfer of electrons between a metal and a non-metal atomMetallic bonding – why most metals are solid§9.3 Lewis Dot Structures for AtomsLDS represent valence electrons as dots around the atom. Here are the LDS for the second row elements: Li· Be: ·B: ·C: ·N: :O: :F: :Ne:H follows the duet rule: it forms stable compounds when it shares 2 electrons. Most main-group atoms follow the octet rule: they tend to have 8 electrons (bonds + lone pairs) in their valence shell.Transition metal behavior is addressed in Chapter 24... . .:::..§9.4 How to Quantify Bond StrengthsIonic bonds – lattice energyCovalent bonds – bond energies§9.4 Lattice EnergyLattice energy – the H released (always exothermic) when ions in the gas phase form one mole of an ionic solid:M+(g) + X-(g) → 1 MX(s)LE is a measure of ionic bond strength.§9.4 Born-Haber CycleBorn-Haber Cycle – calculates lattice energy (H°lattice) using , ionization energy and electron affinity. Consider the example of LiF:Li (s) → Li (g) ½ F2 (g) → F (g) Li (g) → Li+ (g) + e-F (g) + e- → F- (g)Li+ (g) + F- (g) → LiF (s)Li (s) + ½ F2 (g) → LiF (s)161 kJ½(154 kJ) = 77 kJ 520 kJ-328 kJH°lattice-617 kJ161 kJ + 77 kJ + 520 kJ - 328 kJ + H°lattice = -617 kJH°lattice = -1047 kJNotice the magnitude of the lattice energy dominates the multistep process, ionic solids only exist because H°lattice exceeds the energetically unfavorable electron transfer.§9.4 Lattice EnergyLi+(g) + F-(g) → LiF(s)§9.4 Coulomb’s LawThe energy of any 2 ions as a function of charge and distance.Applies to ions of opposite and like charges.Q = charge on each ion (include sign)r = distance between nuclei§9.4 Lattice EnergyLattice energy as a variation on Coulomb’s law: ∴ lattice energy is more exothermic when the ions have higher charges.k = a proportionality constant§9.4 Lattice EnergyOf the following pairs, which has the more exothermic lattice energy?a) NaCl KClb) LiF LiClc) Mg(OH)2MgOd) Fe(OH)2Fe(OH)3e) NaCl Na2Of) MgO BaSQ = charge on each ionr = distance between nucleik = a proportionality constantNa+ smaller than K+F- smaller than Cl-O2- larger charge than HO-Fe3+ larger charge than Fe2+O2- larger charge than Cl-Both ions smaller§9.5 Covalent BondsSingle, double and triple bonds are represented by one, two or three dashes between atoms. Each dash represents an electron pair.oI2oO2oN2§9.6 Polar Covalent BondsFor some covalent bonds, electrons are not equally shared. Such polar covalent bonds can be represented by:1) Partial charges ()2) Dipole; arrow points to more electronegative atom3) Electrostatic potential map, electron-rich and electron-poor regions§9.6 Ionic-Covalent Bond Continuumcovalent bondpolar covalent bondionic bondnuclei§9.6 ElectronegativityElectronegativity (EN) – the ability of an atom to attract shared electrons to itself in a chemical bond.EN can be used to predict bond type:EN Bond Type ExampleSmall, < 0.4 Covalent Cl2Medium, 0.4 – 2.0Polar CovalentHClLarge, > 2.0 Ionic NaCl§9.6 Electronegativity TrendsFour observations:1) EN generally increases left to right across a period (row).2) EN generally increases going up a group (column).3) F is the most EN.4) Fr is the least EN.EN is inversely related to atomic size.§9.6 Ionic vs Covalent CompoundsIonic CompoundsLarge ENCation-anion arraysSolids at STPHigh melting/boiling pointsConducting when liquid17Covalent Compounds Small ENDiscrete moleculesSolid/liquid/gas at STPLow melting/boiling pointsInsulating when liquid192.4 °C§9.6 Ionic vs Covalent Compounds2.11.81.51.20.80.50EN192.4 °C§9.6 Ionic vs Covalent Compounds801714192-69-94-121-102melting point (°C)192.4 °C§9.6 Ionic vs Covalent Compounds10210110-50 000conductivity at mp (S/m)§9.6 Dipole MomentA measure of positive and negative charge separation in a molecule.Dipole moment is the vector sum of individual bond dipoles.MoleculeEN Dipole MomentCl20 0ClF 1.0 0.88HF 1.9 1.82LiF 3.0 6.33§9.6 Dipole MomentEN is calculated, dipole moment is measured.A DM exists for species with permanent uneven charge distributions. DM is measured as flow of current between oppositely charged electrical plates containing a sample:§9.6 Dipole MomentA molecule with polar bonds may or may not have a dipole moment. The individual bond dipoles may cancel or reinforce each other.§9.7 Drawing Lewis Structures1) Count valence electrons from all atoms.2) Connect atoms using single (2-electron) bonds.3) Leftover electrons go to multiple bonds and lone pairs to follow duet (H) and octet rules.§9.7 Lewis Dot StructuresDraw the Lewis dot structure for H2CODraw the Lewis dot structure for HNO2H + C + O = 2(1) + 4 + 6 = 12 e-∴ 6 bonds and lone pairsH + N + O = 1 + 5 + 6 = 12 e-∴ 6 bonds and lone pairsH HCO2468::1012:H N O24681012:::::1 2 3 4 3 2 1 0§9.7 Drawing Lewis Dot StructuresHow many covalent bonds form in neutral compounds of the second-row elements? Li· ·Be· ·B· ·C · ·N: :O: :F: