Honors Chemistry Chapter 9: Chemical Bonding I 9.1 Lewis Dot Diagrams • Symbol surrounded by dots for valence e• Separate dots as much as possible 9.2 Ionic Bonding • Electrostatic force holding ions together • Ions formed by electron transfer • Low Eion loses e-, high e.a. gains e- • • • Na + Cl Na+ ClMg + S Mg2+ S2Li + S 2 Li+ S2- 9.3 Lattice Energy • Energy that holds ionic compounds together in a crystal lattice • Transfer of e- requires energy (Eion) and releases energy (e.a.) • In general, the cation requires more energy than the anion releases, which makes bond formation unstable • Lattice energy releases additional energy, making bond formation stable 9.4 The Covalent Bond • Covalent bond = shared pair of electrons • F + F F–F • Shared pair – shared electrons, bond • Lone pair – electrons not involved in bond 9.4 Lewis Structures • Representation of covalent compounds using dots for e- and lines for bonds • Octet rule • atoms bond in such a way as to gain 8 e- in valence shell • Exceptions – H and He • Multiple Bonds • Double bond – share 2 pairs; eg, O2 • Triple bond – share 3 pairs; eg, N2 9.4 Bonding Summary 9.4 Ionic / Covalent Properties • Intermolecular attractive forces • Ionic – strong, covalent – weak • Consider phase, density, solubility, conductivity • Ionic Covalent • • • • Solid High density Usually soluble Good conductor Liquid or gas Low density Often insoluble Poor conductor 9.5 Electronegativity • Element’s relative attraction for shared e- 9.5 Pauling Electronegativities 9.5 Electronegativity and Atomic Radius 9.5 Bond Character • Degree of sharing of the bonded e• Depends on the difference in electronegativity • Small electronegativity difference • Equal sharing of bonded e• True covalent bond (nonpolar covalent) • Moderate electronegativity difference • Unequal sharing of bonded e• Polar covalent bond • Large electronegativity difference • Ionic bond 9.5 Bond Character 3.0 DEN 2.0 0.0 9.6 Writing Lewis Structures • Draw a reasonable skeletal structure for the compound • Count the total valence electrons available • Draw single bonds between all atoms and use remaining electrons to fulfill the octet rule • If there are not enough electrons to fulfill the octet rule, form double or triple bonds • Draw Lewis structures for NH3, O3, CO32- 9.7 Formal Charges • Difference between the number of e- an atom has in a Lewis structure and the number of e- it has as a free element • Assigning formal charges • Atom counts all its nonbonding e-’s • Atom counts 1 e- from each bond • Total formal charge must add up to the total charge of the molecule / ion 9.7 Formal Charges • • • • The most plausible structures have: The fewest formal charges Formal charges of smallest magnitude Negative formal charges on the most electronegative elements • No adjacent charges of the same type 9.8 Resonance • Some molecules have more than one plausible Lewis structure O O O O O O • Resonance –use of both structures to represent a molecule • Reality is that bonds are delocalized O O O 9.8 Resonance • • • • • • Draw resonance structures for N2O HSCN NO3CO32Extreme resonance – C6H6 9.9 Exceptions to the Octet Rule • • • • • • • Incomplete Octets Not enough electrons to make an octet Usually occurs with Groups IIA and IIIA Examples: BeH2, BF3 Consider resonance in BF3 BF3 + NH3 F3BNH3 Coordinate covalent bond • One atom donates both shared electrons 9.9 Exceptions to the Octet Rule • • • • • • Odd Electron Molecules Often called radicals Examples: NO, NO2 Single e- goes on element with lower EN Very reactive Tend to form dimers 9.9 Exceptions to the Octet Rule • • • • • Expanded Octets More than 8 e- on central atom Requires the atom to have a d orbital Can’t happen with periods 1 and 2 Examples: SF6, XeF4, ClO4-