1. 1 Recap – Last Lecture Cl 2 Covalent Bond HClPolar Covalent Bond NaClIonic Bond Na + Cl - Cl

2. 2 Lewis Structures Lewis Structures provide a way to predict stable bonding arrangements in covalent molecules. Based on maximum number of electrons allowed in the valence shell: –H 2 electrons –C, N, O, F 8 electrons –Period 3 & below: up to 18 electrons but usually 8, 10 or 12

3. 3 Rules for Lewis Structures 1.Arrange the atoms. –Place the least electronegative atom (not H) in the centre. 2.Count the total number of valence electrons. –Remember to add or subtract e- for anions and cations. 3.Allocate two electrons between each pair of atoms which are assumed to be covalently bonded. 4.Use remaining valence electrons to form lone pairs. –Start with the surrounding atoms (centre atom last). 5.Check number of electrons on the central atom. –If too few, move lone pairs from the (least electronegative) surrounding atoms into the bonding region (make double bonds).

4. 4 Period 2 – CH 4 1.C is at the centre 2.Total number of valence electrons = 4 (C) + 4 × 1 (H) = 8 3.Four C-H bonds require 4 × 2 electrons: 4.Electrons remaining = 8 (valence) – 8 (bonding) = 0 5.Carbon has 8 electrons: stable

5. 5 Period 3 – SF 4 1.S is at the centre 2.Total number of valence electrons 6 (S) + 4 × 7 (F) = 34 3.Four S-F bonds require 8 electrons: 4.Electrons remaining = 34 (valence) – 8 (bonding) = 26 Place 3 lone pairs on each fluorine (24 electrons) Place remaining electrons on sulfur (2 electrons) 5.Fluorine has 8 electrons: stable Sulfur has 10 electrons: stable

6. 6 Polyatomic ions – NH 4 + 1.N is at the centre 2.Total number of valence electrons 5 (N) + 4 × 1 (H) -1 (positive charge) = 8 3.Four N-H bonds require 4 × 2 electrons: 4.Electrons remaining = 8 (valence) – 8 (bonding) = 0 5.Nitrogen has 8 electrons: stable

7. 7 Molecules with multiple bonds O 2 Total number of valence electrons: 2 x 6 = 12 Each oxygen has 8 electrons: stable N 2 Total number of valence electrons: 2 x 5 = 10 Each nitrogen has 8 electrons: stable

8. 8 Ozone O 3 Total number of valence electrons: 3 x 6 = 18 Small rings are highly ‘strained’ and do not normally form. Try a different arrangement of oxygen atoms Another valid Lewis structure would be:

9. 9 Resonance When two or more Lewis structures are possible for a molecule it is said to exhibit resonance. This is indicated by double headed arrows:  This does not mean the molecule flips back and forth between the possible structures – the true structure is an average of the possible resonance structures. Bonds are neither single nor double – they are intermediate between the two. They have intermediate length and strength.

10. 10 Resonance Benzene: C 6 H 6 Benzene is best described as being a resonance hybrid of these two bonding arrangements.

11. 11 Resonance Benzene: C 6 H 6 Benzene is also about 140 kJ mol -1 more stable than predicted for 1,3,5-cyclohexatriene (the structure with localised double bonds). Bond LengthBond Strength C–C154 pm356 kJ mol -1 C=C133 pm636 kJ mol -1 benzene139 pm518 kJ mol -1

12. 12 Resonance – polyatomic ions Nitrate, NO 3 - Total number of valence electrons: 5 + (3 x 6) + 1 = 24

13. 13 NO 3 - Nitrate ion All N-O bond lengths the same and intermediate between N-O and N=O (bond order = 1 1/3) Negative charge distributed evenly across the three oxygen atoms (1/3 on each oxygen)

14. 14 Resonance If resonance occurs… The actual structure is an average of all possible Lewis structures. The actual structure is a single structure. The actual structure is more stable than predicted from the Lewis structure. There is no easy way of drawing the actual structure!

15. 15 Predicting the best structure Draw the Lewis structure, including lone pairs and multiple bonds. Check for the presence of ‘equivalent’ resonance structures. Check number of bonds on C, N, O & F as a prompt for ‘non-equivalent’ resonance structures. Choose best structure and check again for resonance structures.

16. 16 Example: POCl 3 Total number of valence electrons: 5 + 6 + (3 x 7) = 32 Oxygen has a valency of 2 so, when uncharged, forms two bonds.

17. 17 Example: POCl 3 P now has 10 electrons, but this is OK as valence electrons are in the third shell. O has the two bonds expected from its valency. These are resonance structures (only position of the electrons have moved) but not ‘equivalent’. Right hand structure is a better representation of the real structure.

18. 18 Applications Streptonigrin is an anti-biotic and anti-tumor agent. http://www.sigmaaldrich.com/content/dam/sigma- aldrich/structure0/131/mfcd00063401.eps/_jcr_content/r enditions/mfcd00063401-medium.png It contains flat benzene like rings. These rings can slot between adjacent base pairs in DNA, rather like putting a letter in a letterbox, and disrupt the replication of DNA in tumor cells.

19. By the end of this lecture, you should: −be able to draw out plausible Lewis structures for simple polyatomic molecules. −assign bond orders based on sharing of electron pairs and resonance structures. −identify the structure that most closely resembles the actual structure given non-identical resonance structures. −be able to complete the worksheet (if you haven’t already done so…). 19 Learning Outcomes:

20. 20 Questions to complete for next lecture: 1.Draw the Lewis structures of NH 2 -, NH 3 and NH 4 + and indicate how the number of bonds relates to the valency of nitrogen. 2.Draw the Lewis structure of PCl 3 and PCl 5 and indicate the number of electrons associated with the phosphorus atom. 3.Draw the ‘best’ Lewis structure of SO 2 and SO 3 and give a reason for your choice. 4.Draw the ‘best’ Lewis structure of HPO 3 2- and give a reason for your choice (note H is bonded to the P).