1. Molecular shapes Balls and sticks

2. Learning objectives  Apply VSEPR to predict electronic geometry and shapes of simple molecules  Distinguish between polar and nonpolar bonds in molecules  Predict polarity of simple molecules from bond polarity and molecular shape

3. Roadmap to polarity  Establish skeleton of molecule  Determine Lewis dot structure using S = N – A  Determine electronic geometry using VSEPR  Identify molecular geometry from molecular  Count number of polar bonds  Perform polarity analysis

4. Valence shell electron pair repulsion  Lewis dot structure provides 2D sketch of the distribution of the valence electrons among bonds between atoms and lone pairs; it provides no information about molecular shape  First approach to this problem is to consider repulsion between groups of electrons (charge clouds)

5. Electron groups (clouds) minimize potential energy  Valence shell electron pair repulsion (VSEPR)  Identify all groups of charge: non- bonding pairs or bonds (multiples count as one)  Bonded atoms – single, double or triple count as 1  Lone pairs count as 1  Distribute them about central atom to minimize potential energy (maximum separation)

6. Choices are limited  Groups of charge range from 2 – 6  Only one electronic geometry in each case  More than one molecular shape follows from electronic geometry depending on number of lone pairs  One surprise: the lone pairs occupy more space than the bonded atoms (with very few exceptions)  Manifested in bond angles (examples follow)  Molecular shape selection (particularly in trigonal bipyramid)

7. Total number of groups dictates electronic geometry  Octet rule:  Two – linear  Three – trigonal planar  Four – tetrahedral  Additional possibilities (expand octet):  Five – trigonal bipyramidal  Six - octahedral

8. Stage 3: Molecular shape:  What you get from electronic geometry considering atoms only  Same tetrahedral electronic geometry – different molecular shape

9. Two groups: linear  Except for BeH 2, all cases with two groups involve multiple bonds

10. Three groups: trigonal planar  Two possibilities for central atoms with complete octets:  Trigonal planar (H 2 CO)  Bent (SO 2 )  BCl 3 provides example of trigonal planar with three single bonds  B is satisfied with 6 electrons

11. Four groups: tetrahedral  Three possibilities:  No lone pairs (CH 4 ) - tetrahedral  One lone pair (NH 3 ) – trigonal pyramid  Two lone pairs (H 2 O) – bent  Note: H-N-H angle 107°H-N-H angle 107° H-O-H angle 104.5°H-O-H angle 104.5° Tetrahedral angle 109.5°Tetrahedral angle 109.5°

12. Representations of the tetrahedron

13. Groups of charge Lone electron pairs Electronic geometry Molecular shape 20Linear 30Trigonal planar 31 Bent 40Tetrahedral 41 Trigonal pyramid 42TetrahedralBent

14. Important properties related to polarity  Solubility: polar molecules dissolve in polar solvents; nonpolar molecules dissolve in nonpolar solvents  Oil (nonpolar) and water (polar) don’t mix  Ammonia (polar) dissolves in water  Melting and boiling points  Polar substances have high intermolecular forces:  Melting and boiling points are much higher than with nonpolar substances (H 2 O is a liquid, CO 2 is a gas)

15. Roadmap to polarity  Establish skeleton of molecule  Determine Lewis dot structure using S = N – A  Determine electronic geometry using VSEPR  Identify molecular geometry from molecular  Count number of polar bonds  Perform polarity analysis

16. Polar bonds and polar molecules  Not all molecules containing polar bonds will themselves be polar.  Need to examine the molecular shape  Ask the question:  Do the individual bond polarities cancel out?  If so, non polar. If not, polar.

17. Consider some examples  In CO 2 (linear molecule) the two polar bonds oppose each other exactly  In chemical tug-o-war there is stalemate

18. The most important polar molecule  In BF 3 the three bonds cancel out – tug of war stalemate  In H 2 O (bent) the polar bonds do not directly oppose – no stalemate  Lone pair also adds some component  Overall net polarity  Consequence of polarity: H 2 O is a liquid, CO 2 is a gas

19. Symmetry and polarity  If the molecule “looks” symmetrical it will be nonpolar  If the molecule “looks” non-symmetrical it will be polar

20. Rules of thumb for evaluation of polarity  Presence of one lone pair of electrons  Only one polar bond  Always polar molecules  Two or more polar bonds  Do polar bonds perfectly oppose?  If no, polar molecule

21. Two bonds  Equal bonds oppose (linear)  Nonpolar (CO 2 )  Unequal bonds oppose (linear)  Polar (HCN)  Equal bonds do not oppose (bent)  Polar (H 2 O)

22. Three bonds  Equal bonds oppose in trigonal planar arrangement  Nonpolar  Unequal bonds in trigonal planar arrangement  Polar