Unit 2 (Chp. 8,9): Bonding & Molecular Geometry Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Unit 2 (Chp. 8,9): Bonding & Molecular Geometry John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

Energetics of Ionic Bonding IE & EA & ?? (3rd factor) +energy electrostatic (Coulombic) attraction between Na+ ion & Cl– ion. Forming attraction release or absorb energy? attract Coulomb’s law: F = k q1q2 d2 There must be a third piece to the puzzle. –energy F : force of attraction q : charge d : distance between

Lattice Energy: _________ charge, and ___________ size. increasing energy released when a mole gaseous ions attract to form a solid ionic compound. Lattice energy increases with: _________ charge, and ___________ size. increasing decreasing (q) (d) q , d stronger attraction E =  q1q2 d  energy

Covalent Bonding share electrons covalentbond

Electronegativity attraction Zeff , shield -ability of an atom to attract electrons when bonded (sharing e–’s) to another atom. attraction Zeff , shield most electronegative

Polarity of Covalent Bonds the more electronegative F has more e– density. d+ d– F F H F NONpolar (DEN ~ 0) Polar (DEN > 0.4) (shared equally) (shared UNequally) dipole: partial charges (d+, d–) on atoms that are sharing e–’s UNequally dipole moment,  : a measure of polarity with direction

Polar Covalent Bonds The greater the ∆EN, the more polar the bond. H I Br H Cl H F

H2O Molecular Polarity Overall dipole moment a molecule is polar ONLY IF: has polar bonds (DEN) arranged asymmetrically (do not cancel out) H2O Polar Overall dipole moment

Drawing Lewis Structures In 5 easy steps: PCl3 Count (val e–’s) Connect (bonds) Octet (outer dots) Octet (central dots) (if necessary) form Multiple Bonds to fill central octet

Drawing Lewis Structures Draw a Lewis structure for CO2 –1 +1 or Formal Charge = (VE’s) – (NBE’s) –(½BE’s) FC = (val e– ’s) – (dots) – (lines) NOTE: FC: charges IF all atoms shared e–’s equally (same electronegativity) IF 100% covalent.

Drawing Lewis Structures Use FC to select the best Lewis structure ONLY when directed… …the fewest & smallest charges. …a “–” on the most electronegative atom. –2 +1 Which of the Lewis structures for NCO– is the most significant in terms of formal charge? –1 –1

Resonance Multiple resonance structures are required if… …bonds can not be clearly identified between identical atoms due to delocalized e–’s.

Less Than Octet For BF3: A filled octet places a – on the B and a + on F. A double bond between B and F is worse than leaving B with only 6 electrons. Principle: Do NOT fill B’s octet if it gives a +FC on a more electronegative outer atom.

Expanded Octet phosphate ion (PO43–) OR The better structure minimizes the FC’s. Principle: ONLY expand octets for atoms in periods > 3 as needed when… directed to consider Formal Charge.

Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes repulsions among them.”

Molecular Geometry Must know basic shapes: AB2 linear AB3 trigonal planar AB4 tetrahedral AB5 trigonal bipyramidal AB6 octahedral AB4 AB3E AB2E2 all variations have 1 or more E’s (e– pairs)

Two Types of Orbital Overlap Sigma () bond overlap forms: head-to-head overlap single bonds Sigma (single)

Two Types of Orbital Overlap Pi () bond overlap forms: side-to-side overlap. double & triple bonds. Pi (p orbitals)

Hybridization (of orbitals) Consider carbon (C): promote hybridize …4 degenerate sp3 orbitals. ground state

Multiple Bonds C sp2 orbital forms  overlap with O sp2 orbital and with H s orbital. unhybridized p orbitals of C & O form  bond overlap.

= MO Theory & Bond Order total # of bonds Bond Order = # of bonding domains What are the bond orders of these molecules: 2 1.5 resonance delocalized p e–’s =