IIIIII 8.4 and 8.5 Ch. 8 – Molecular Structure

A. VSEPR Theory n Valence Shell Electron Pair Repulsion Theory n Electron pairs orient themselves in order to minimize repulsive forces.  As far apart from each other as possible!

A. VSEPR Theory n Types of e - Pairs  Bonding pairs - form bonds  Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!

A. VSEPR Theory n Lone pairs reduce the bond angle between atoms. Bond Angle

4. Example n BF 3 vs. SO 2 B S F F F 120° O O <120°

n Draw the Lewis Diagram. n Tally up e - pairs on central atom.  * double/triple bonds = ONE pair n Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles! B. Determining Molecular Shape

C. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR 180° BeH 2

3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 C. Common Molecular Shapes

3 total 2 bond 1 lone BENT <120° SO 2

4 total 4 bond 0 lone TETRAHEDRAL 109.5° CH 4 C. Common Molecular Shapes

4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° NH 3 C. Common Molecular Shapes

4 total 2 bond 2 lone BENT 104.5° H2OH2O C. Common Molecular Shapes

5 total 5 bond 0 lone TRIGONAL BIPYRAMIDAL 120°/90° PCl 5 C. Common Molecular Shapes

6 total 6 bond 0 lone OCTAHEDRAL 90° SF 6 C. Common Molecular Shapes

D. Summary n See Word Document Section 6-5 Outline

n PF 3 4 total 3 bond 1 lone TRIGONAL PYRAMIDAL 107° F P F F Examples

n CO 2 O C O 2 total 2 bond 0 lone LINEAR 180° Examples

C. Bond Polarity zMost bonds are a blend of ionic and covalent characteristics. zBond type determined by difference in electronegativity

C. Bond Polarity zElectronegativity yAttraction an atom has for a shared pair of electrons. yhigher electroneg. atom   - ylower electroneg. atom   +

C. Bond Polarity zElectronegativity Trend yIncreases up and to the right.

zNonpolar Covalent Bond ye - are shared equally ysymmetrical e - density yusually identical atoms C. Bond Polarity

++ -- zPolar Covalent Bond ye - are shared unequally yasymmetrical e - density yresults in partial charges (dipole)

EN Diff:BOND TYPE < 0.4Nonpolar cov. 0.4 – 1.7Polar covalent > 1.7Ionic C. Bond Polarity

Bond Polarity Examples n Cl – 3.16 = 0 nonpolar covalent n HCl 3.16 – 2.20 = 0.96 polar covalent n NaCl 3.16 – 0.93 = 2.23 ionic **Note: When more than two atoms are involved, must examine geometry and dipoles to determine polarity

G. Intermolecular Forces n intermolecular forces: (def) forces of attraction between molecules n vary in strength but generally weaker than bonds between atoms in molecules, ions in ionic compounds, or metals atoms in solid metals

G. Intermolecular Forces n Dipole-dipole forces: a. dipole: (def): created by equal but opposite charges that are separated by a short distance; direction is from dipole’s positive to negative end c. dipole-dipole force: force of attraction between polar molecules: negative region of one polar molecule attracts positive region in neighboring molecules

G. Intermolecular forces b. H – Cl d. Additive dipoles create overall molecular dipole H 2 ONH 3 CCl 4 CO 2

G. Intermolecular Forces 3. Hydrogen bonding a. (def) Occurs when a hydrogen atom is bonded to a highly electronegative atom and is attracted to an unshared pair of e- in a nearby molecule b. Represented by dotted lines connecting the hydrogen-bonded H to the unshared electron pair of the electronegative atom to which it is attracted (pic)

G. Intermolecular Forces 4. London Dispersion Forces a. (def) intermolecular attractions resulting from the constant motion of e- and the creation of instantaneous dipoles b. present in all atoms and molecules but the only intermolecular force acting among noble gas atoms, nonpolar molecules, and slightly polar molecules

From this point on… z FYI information.

E. Hybridization 1. hybridization: (def) the mixing of two or more atomic orbitals of similar energies on the same atom in order to produce new orbitals of equal energies 2. hybrid orbitals: (def) orbitals of equal energy produced by combination of 2 or more orbitals in the same atom

E. Hybridization 3. Example: CH 4 C e- configuration: 1s 2 2s 2 2p 2 Energy ___ 1s ___ 1s ___ 2s __ __ __ 2p __ __ sp 3 BEFOREAFTER

E. Hybridization Atomic Orbitals Type of Hybridization # of hybrid orbitals Geometry s,psp2 linear s,p,psp 2 3 trigonal planar s,p,p,psp 3 4 tetrahedral