Chemical Bonding I: Lewis TheoryChemical BondingBondsMolecular compounds boil at low temperatures because only weak intermolecular forces must be disrupted.Slide 5Ionic BondsSlide 7Slide 8Lattice energy dependence on atomic sizesLattice energy dependence on ionic chargesFormation of ionic compoundsEnergetics of NaCl FormationSlide 13Slide 14Slide 15Slide 16Slide 17Slide 18magnesium bromideMg + Cl2  MgCl2 DH = -642 kJ/mol Mg + Br2  MgBr2 DH = -561 kJ/mol Why are they different?Slide 21Slide 22Covalent BondingRepresenting Atoms, Ions, and Molecules as Lewis Electron Dot StructuresSlide 25Polar Covalent Bonds and ElectronegativitySlide 27Bond PolarityElectronegativitySlide 30Lewis Electron Dot StructuresWriting Lewis Dot StructuresSlide 33Slide 34Formal ChargeExpanded OctetsLewis Structures of ionsResonanceBond length - the optimum distance between nuclei in a covalent bond.Bond dissociation energy – (Bond Strength)Slide 41DefinitionsChemical Bonding I: Lewis TheoryChapter 9Chemical BondingAtoms gain, lose, or share electrons in order to achieve a full outer shell electron configuration.BondsIonic Bonds•Composed of ions that have gained or lost electrons to achieve a full outer shell•Electrostatic attractive forces•Crystalline solids – no discrete molecules - formula units•Identified by empirical formulas •Metal + non-metal Covalent Bonds•Composed of atoms that are sharing electrons to achieve a full outer shell•Shared electron bonds•Discrete molecules, forms gases, liquids, and solids•Identified by molecular formulas •Non-metal + non-metalMolecular compounds boil at low temperatures because only weak intermolecular forces must be disrupted.Ionic compounds boil at high temperatures because strong electrostatic bonding forces must be disrupted.Ionic Bonds•Ions are held together by ionic attraction where the force of attraction is governed by Coulomb’s Law.• • •Makes sense•Large Z  strong attraction  larger E•Large d  charge felt less  smaller EdZkZFddZkZF21221E and Z=ion charged = distance between nucleiLattice EnergiesLattice energy dependence on atomic sizes LiCl 834 kJ/molNaCl 787 kJ/molLCl 701 kJ/molCsCl 657 kJ.molLattice energy dependence on ionic chargesNaF 910 kJ/molCaO 3414 kJ/molFormation of ionic compounds•Get elements as atoms (generally requires energy)•Form ions (anions are energetically favorable, cations are unfavorable)•Bring ions together (favorable)•Condense to solid phase (favorable)Na(s) Cl2(g)Na(g)Na+(g)Cl(g)Cl-(g)NaCl(s)vaporizeionize(IE)dissociateionize(EA)lattice energy++Energetics of NaCl Formation•Na(s) Na(g) +107.3 kJ/mol•Na(g)  Na+ + 1e+495.8 kJ/mol•1/2 Cl2(g)  Cl (g) +122 kJ/mol•Cl(g)  Cl(g) 348.6 kJ/mol•Na+(g) + Cl(g)  NaCl(s) 787 kJ/mol•====================================•Na(s) + 1/2 Cl2(g)  NaCl(s) 411 kJ/mol•Determine the energy of formation of MgBr2 from the elements.magnesium bromide•Mg(s)  Mg(g) +147.7 kJ/mol•Mg(g)  Mg+ + 1e+737.7 kJ/mol•Mg+(g)  Mg2+(g) + 1e+1450.7 kJ/mol•Br2(g)  2 Br (g) +193 kJ/mol•2Br(g)  2Br(g) 2(325 kJ/mol)• = 650 kJ/mol•Mg2+(g) + 2Br(g)  MgBr2(s) 2440 kJ/mol•====================================•Mg(s) + Br2(g)  MgBr2(s)  561 kJ/molMg + Cl2  MgCl2H = -642 kJ/molMg + Br2  MgBr2H = -561 kJ/molWhy are they different?•Calculate the energy released in kJ/mol in the reaction• •Na(s) + 1/2 I2(s)  NaI(s)• •The energy of vaporization of Na(s) is 107 kJ/mol. The sum of the•enthalpies of dissociation and vaporization of I2(s) is 214 kJ/mol, and the lattice energy of NaI is 704 kJ/mol.•Calculate the energy released in kJ/mol when LiH is formed in the reaction•Li(s) + ½ H2(g)  LiH(s)• •Heat of vaporization, Li 161 kJ/mol•Dissociation energy, H2436 kJ/mol•Lattice energy, LiH -917 kJ/mol•Ionization energy, Li 520 kJ/mol•Electron affinity, H -73 kJ/mol• •Answer: -91 kJ/mole net changeCovalent Bonding•Shared electron bonds•Due to overlap of atomic orbitals –(Valence Bond Theory)•Allows each atom to fill valence shell with electronsRepresenting Atoms, Ions, and Molecules as Lewis Electron Dot Structures•Use dots to represent valence electronsPolar Covalent Bonds and Electronegativity•Polar bonds are bonds where the electron density is not shared equally between the two bonded atoms.•In polar bonds there is a positive and a negative end to the bond.Bond PolarityNaCl Cl-ClHClElectronegativity•The ability of an atom in a bond to attract electrons toward itself.-Electron greed-Note that electronegativity increases up and to the right as do the ionization energy and the electron affinityLewis Electron Dot Structures•Bonding electrons pairs – electron pairs involved in bonds•Lone electron pairs – electron pairs that do not participate in bonding•Bond order = number of bondsWriting Lewis Dot Structures•Decide which atoms are bonded together - draw a skeleton structure•Count the total number of valence electrons available.•Find the number of electrons needed to give an octet around all atoms -- (remember H needs 2, all else need 8).Writing Lewis Dot Structures•Determine number of electrons short.•Number of bonds needed = number of electrons short/2.•Distribute bonds -- (1st hook atoms together and then add double bonds where appropriate).•Calculate number of electrons used in bonds.Writing Lewis Dot Structures•Calculate electrons remaining.•Distribute remaining electrons to give all atoms an octet.•Done!!Formal Charge•The result of a method of electron bookkeeping that tells whether an atom in a molecule has gained or lost electrons compared to an isolated atom.•Formal charge = # valence electrons – (# bonds + # electrons as lone pairs)Expanded Octets•Elements beyond neon have available d orbitals that may be used to accept additional electrons if necessary.-If you the number of bonds necessary to hook all atoms together is greater than the number needed to give all an octet then put in necessary bonds and distribute extra electrons on atoms that have available d orbitals in which to expand.Lewis Structures of ions •for anions add the extra electrons to the number available•for cations subtract the lost electrons from the number availableResonance•In some Lewis structures, the multiple bonds