PowerPoint PresentationSlide 2Slide 3Slide 4Slide 5Slide 63.4 Covalent Bonds and Lewis StructuresThe Lewis Model of Chemical BondingCovalent Bonding in H2Covalent Bonding in F2The Octet RuleSlide 12Slide 133.4 Double Bonds and Triple BondsSlide 15Slide 163.4 Formal ChargesNitric acidSlide 19Slide 20Slide 21Slide 22Slide 23Slide 24Slide 25Slide 26Slide 273.5 Drawing Lewis StructuresConstitutionTable 1.4 How to Write Lewis StructuresSlide 31Slide 32Slide 33Slide 34Slide 35Slide 36Slide 37Slide 38Slide 39Slide 40Slide 41Slide 42Slide 43Condensed structural formulasBond-line formulasSlide 463.5 Constitutional IsomersConstitutional isomersA Historical NoteExamples of constitutional isomers3.5 ResonanceResonanceSlide 53Slide 54Slide 55Slide 56Resonance Structures of Methyl NitriteSlide 58Why Write Resonance Structures?ExampleSlide 613.7 The Shapes of Some Simple MoleculesSlide 63Slide 64MethaneSlide 66Valence Shell Electron Pair RepulsionsWaterAmmoniaBoron TrifluorideMultiple BondsFormaldehyde: CH2=OFigure 1.12: Carbon DioxideSlide 74Slide 753.7: Polar Covalent Bonds and ElectronegativityElectronegativity is a measure of an element to attract electrons toward itself when bonded to another element.Slide 78Slide 79GeneralizationSlide 813.7 Molecular Dipole MomentsDipole MomentSlide 84Slide 85Slide 86Slide 87Molecular Dipole MomentsSlide 89Comparison of Dipole MomentsCarbon tetrachlorideDichloromethaneSections 3.3 & 3.4 Covalent Bonding and Lewis StructuresLearning goals:Writing valid Lewis structures for molecular substancesPredicting molecular geometry from Lewis structures (VSEPR theory)Understanding electronegativity and how this concept allows the distinction between polar bonds and non-polar bondsUsing Lewis structures to determine whether a molecule has a dipole moment or notUsing the octet rule to compute formal charges on atoms and multiple bonding between atomsSections 3.3 & 3.4 Covalent Bonding and Lewis Structures(1) Lewis “dot” (electron) structures of valence electrons for atoms(2) Use of Periodic Table to determine the number of “dots”(3) Use of Lewis structures to describe the electronic structures of atoms and molecules(4) Works best for covalent bonds and for elements in the first full row of the Periodic Table: H, He, Li, Be, B, C, N, O, F, Ne(5) Works with restrictions for second full row of the Periodic Table and beyond: Na, Mg, Al, Si, P, S, Cl, ArSome issues about Lewis Structures to be discussed:(1) Drawing “valid” Lewis structures which follow the “octet” rule (holds almost without exception for first full row)(2) Drawing structures with single, double and triple bonds(3) Dealing with isomers (same composition, different constitution)(4) Dealing with resonance structures (same constitution, different bonding between atoms)(5) Dealing with “formal” charges on atoms in Lewis structures(6) Dealing with violations of the octet rule:Molecules which possess an odd number of electronsMolecules which are electron deficientMolecules which are capable of making more than four covalent bondsLewis “dot-line” representations of atoms and molecules(1) Electrons of an atom are of two types: core electrons and valence electrons. Only the valence electrons are shown in Lewis dot-line structures.(2) The number of valence electrons is equal to the group number of the element for the representative elements.(3) For atoms the first four dots are displayed around the four “sides” of the symbol for the atom.(4) If there are more than four electrons, the dots are paired with those already present until an octet is achieved.(5) Ionic compounds are produced by complete transfer of an electron from one atom to another.(6) Covalent compounds are produced by sharing of one or more pairs of electrons by two atoms.The valence capacity of an atom is the atom’s ability to form bonds with other atoms. The more bonds the higher the valence.The valence of an atom is not fixed, but some atoms have typical valences which are most common:Carbon: valence of 4Nitrogen: valence of 3 (neutral molecules) or 4 (cations)Oxygen: valence of 2 (neutral molecules) or 3 (cations)Fluorine: valence of 1(neutral molecules) or 2 (cations)Covalent bonding and Lewis structures(1) Covalent bonds are formed from sharing of electrons by two atoms.(2) Molecules possess only covalent bonds.(3) The bedrock rule for writing Lewis structures for the first full row of the periodic table is the octet rule for C, N, O and F: C, N, O and F atoms are always surrounded by eight valence electrons.(4) For hydrogen atoms, the doublet rule is applied: H atoms are surrounded by two valence electrons.3.4Covalent Bonds and LewisStructures•In 1916 G. N. Lewis proposed that atomscombine in order to achieve a more stableelectron configuration.•Maximum stability results when an atomis isoelectronic with a noble gas.•An electron pair that is shared between two atoms constitutes a covalent bond.The Lewis Model of Chemical BondingThe Lewis Model of Chemical BondingCovalent Bonding in H2Covalent Bonding in H2HH..HH..Two hydrogen atoms, each with 1 electron,Two hydrogen atoms, each with 1 electron,can share those electrons in a covalent bond.can share those electrons in a covalent bond.HH ::HH•Sharing the electron pair gives each hydrogen an electron configuration analogous to helium.Covalent Bonding in F2Covalent Bonding in F2Two fluorine atoms, each with 7 valence electrons,Two fluorine atoms, each with 7 valence electrons,can share those electrons in a covalent bond.can share those electrons in a covalent bond.•Sharing the electron pair gives each fluorine an electron configuration analogous to neon.........FF..FF..::::........FF::FF::::................The Octet RuleThe Octet Rule•The octet rule is the most useful in cases involving covalent bonds to C, N, O, and F.FF::FF::::................In forming compounds, atoms gain, lose, or In forming compounds, atoms gain, lose, or share electrons to give a stable electron share electrons to give a stable electron configuration characterized by 8 valence configuration characterized by 8 valence electrons.electrons.ExampleExampleExampleExampleCC........FF::..........Combine carbon (4 valence electrons) andCombine carbon (4 valence electrons) andfour fluorines (7 valence electrons each)four fluorines (7 valence electrons each)to write a Lewis structure for CFto write a Lewis structure for CF44..::FF::........CC::FF::........::FF::........::FF::........The