1. The shapes of things Molecular shape determines properties Bonding determines shape

2. Learning objectives  Write Lewis dot structures for simple molecules  Predict shape of simple molecules  Predict polarity of simple molecules

3. Covalent molecular compounds  Covalent compounds are usually molecular  Bonds between atoms are covalent  Interactions between molecules are very weak  Atoms in a covalent molecule don’t stack like marbles  Bonds have specific directions

4. Molecules have specific shapes  Shape will depend on The number of atoms bonded to the central atom The number of atoms bonded to the central atom The number of lone pairs around the central atom The number of lone pairs around the central atom  Distinguish between Electronic geometry (molecular geometry) Electronic geometry (molecular geometry) Consider atoms and lone pairs Consider atoms and lone pairs Molecular shape Molecular shape Consider atoms only Consider atoms only

5. Lewis dot structures: doing the dots  Molecular structure in simplest terms: arrange valence electrons as dots in a 2-dimensional figure Only valence electrons are shown Only valence electrons are shown Electrons are either in: Electrons are either in: bonds bonds lone pairs (stable molecules do not contain unpaired electrons – with very few exceptions) lone pairs (stable molecules do not contain unpaired electrons – with very few exceptions) Octet rule is guiding principle: each atom has 8 dots round it (H has 2 dots) Octet rule is guiding principle: each atom has 8 dots round it (H has 2 dots)

6. Lewis dot structures made easy: the S = N –A machine  Start with the skeleton of the molecule  Least electronegative element is the central atom  S = N - A N = total number of electrons required to fill octet for each atom in the molecule (8 for each element, except 2 for H and 6 for B) N = total number of electrons required to fill octet for each atom in the molecule (8 for each element, except 2 for H and 6 for B) A = total number of valence electrons A = total number of valence electrons S = total number of electrons in bonds S = total number of electrons in bonds  We are given N and A; we need to find S

7. Applying the rules  Calculate N for the molecule  Calculate A (all the dots) include charges for ions (add one for each –ve charge and subtract one for each +ve charge) include charges for ions (add one for each –ve charge and subtract one for each +ve charge)  Determine S (no of dots in bonds) (S = N – A) (S = N – A)  Satisfy all octets and create number of bonds dictated by S (may be multiple bonds)  NF 3  N = 8(N) + 3 x 8(F) = 32  A = 5(N) + 3 x 7(F) = 26  S = 32 – 26 = 6 N FF F NFF F

8. Two tests for dot structures  Is the number of dots in the molecule equal to the number of valence electrons?  Are all the octets satisfied?  If both yes structure is valid  If either no then back to the drawing board

9. Electronic geometry  Identify central atom. Many molecules have more than one. Central atom has more than one atom bonded to it Central atom has more than one atom bonded to it

10. Methanol has two central atoms  O is one central atom – bonded to H and C  C is another central atom – bonded to O, H, H and H  Consider geometry around each one separately

11. Counting regions of charge  Count only atoms and lone pairs immediately bonded to central atom  Count the regions of electrons Bonds – single, double or triple count as 1 Bonds – single, double or triple count as 1 Lone pairs count as 1 Lone pairs count as 1  Number will be between 2 and 4 for molecules that obey octet rule

12. Counting groups  OF 2 two bonds, two lone pairs Total groups = 4 Total groups = 4  CF 4 four bonds, no lone pairs Total groups = 4 Total groups = 4

13. Double or triple bonds count as one  CO 2 has two groups  HCN has two groups

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

15. Summary of possible molecular shapes

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? Do the individual bond polarities cancel out? If so, non polar. If not, polar. 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 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 Lone pair also adds some component Overall net polarity 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 Always polar molecules  Two or more polar bonds  Do polar bonds perfectly oppose? If no, polar molecule If no, polar molecule