1. Bonding and Molecular Structure

2. Bonds  A chemical bond forms when two atoms rearrange their valence electrons which causes an attraction between the atoms.  An ionic bond forms when electrons are transferred, creating positive and negative ions.  A covalent bond forms when electrons are shared.

3. Valence and Core Electrons  You should know what valance electrons are and be able to make an educated guess what core electrons are.  Valence electrons-  Core electrons-

4. Valence Electrons for Transition Metals  For transition metals, the ns and (n-1)d orbitals are valence electrons.  For example, iron would have 8 valence electrons. Two from 4s and six from 3d.  How many valence electrons do the following elements have?  Co  Ca  Cr

5. Drawing Lewis Structures  A Lewis structure depicts an element with its valence electrons.  To draw them, write the symbol of the element and place a dot around the symbol for each valence electron the element has.  Example: Sulfur  Practice: CalciumSiliconKrypton

6. Lewis Structures for Compounds  Starting with the structures for elements, we can rearrange electrons to form compounds. Atoms will end up with a full valance shell. 1. Determine arrangement of atoms. The central atom has the lowest electron affinity. 2. Determine total number of valence electrons. Charges change this number! 3. Form single bonds between all bonded atoms. 4. Add lone pairs and multiple bonds until all atoms have a full shell  Example: Carbon DioxideChlorate ion

7. Practice  Draw Lewis structures for the following:  Nitronium (NO 2 + )carbonate (CO 3 2- )

8. Assignment  Page 355 Exercise 8.1  Page 395 2, 8

9. Predicting Lewis Structures  Lewis structures have patterns.  Carbon typically forms four bonds.  Nitrogen typically forms three bonds and has one lone pair.  Oxygen typically forms two bonds and has two lone pairs.  Fluorine typically forms one bond and has three lone pairs.

10. Practice  Draw Lewis structures for CCl 4 and NF 3 using predictions.

11. Isoelectronic Species  Molecules and ions that have the same number of valence electrons and the same Lewis structures are isoelectronic.  Isoelectronic species will have the same number of atoms and the same number of valence electrons.  Example: NO +, N 2, CN -

12. Formal Charges  The formal charge is the charge on an atom in a molecule or ion.  The sum of the formal charges gives the overall charge of the molecule or ion.  Example:

13. Practice  Find the formal charge of each atom in chlorate, ClO 3 -, and phosphate, PO 4 3-.

14. Assignment  Page 357 Exercise 8.2  Page 358 Exercise 8.4  Page 360 Exercise 8.5  Page 396 14, 16

15. Resonance  Sometimes a single Lewis structure doesn’t show the actual electronic structure.  Different possible structures for a molecule are called resonance structures.

16. Resonance Hybrid  When you have a mix of single and double bonds, you don’t really have single or double bonds, the bonds are all equal and somewhere between single and double bonds.  The two resonance structures above show the extremes, but the truth lies in- between.  The bottom shows how the true structure is represented. The electrons forming the double bonds are spread out amongst all the carbon atoms

17. Practice  Draw the resonance structures for nitrite, NO 2 -. Are the N-O bonds single, double, or intermediate?

18. Assignment  Page 363 Exercise 8.6  Page 396 10

19. Exceptions to the Octet Rule  Although the majority of elements want eight valence electrons, there are some exceptions.  Boron is often stable with only six valence electrons  Elements in period 3 and on often form compounds with more than eight valence electrons.

20. Practice  Draw the Lewis structure for ClF 4 -

21. Question  How do you arrange three objects so they are as far apart from each other as possible?  What about four objects?

22. Molecular Shapes  VSEPR (valence shell electron pair repulsion) theory allows us to predict the shape of different molecules.  The basis is that bonds and lone pairs want to be as far away from each other as possible, similar to a sibling that is getting on your nerves.  To the right are the shapes for molecules with no lone pairs.

23. Lone Pairs  Lone pairs also take up space and affect the shape of the molecule. We will need the Lewis structure to correctly predict the shape.  Lone pairs take up more space in fact, which moves the bonds closer.

24. To the right are the possible shapes for VSEPR. There are similar charts in your book on pages 371 and 372. Example: Find the shape of H 3 O + Practice: Find the shapes of SiCl 4, ClF 2 +, ICl 4 -, and NO 3 -.

25. Assignment  Page 396 18, 20, 22, 24

26. Polar Bonds  In covalent bonds, often the electrons aren’t shared equally.  This makes a polar covalent bond.  The atom that has the electrons closer to it has a partial negative charge and the other is partially positive.

27. Electronegativity  Linus Pauling, an American chemist, created the idea of electronegativity.  It is a value that tells us how well an atom attracts electrons.  A bond is polar when the two atoms have a large difference in electronegativity values.  Practice: Which bond is more polar? Label which is + and -.  B-F or B-Cl  Si-O or P-P

28. Electronegativity and Formal Charges  Sometimes formal charges are illogical when we consider electronegativity.  In BF 4 -, the formal charge of each fluorine is 0 and the boron is -1.  This doesn’t make sense because fluorine has a higher electronegativity. The fluorine would be expected to be negative.  What happens is the negative charge is spread across the molecule.  Also, it is preferred to have charges of zero when possible, this explains why some resonance structures are preferred.

29. Practice  We discussed how boron is often stable with only three bonds. Why doesn’t it just form a double bond to complete an octet? Draw Lewis structures for BF 3 and find formal charges for each atom.

30. Dipole Moment  We talked about how bonds can be polar. As a result, whole molecules can be polar as well.  The dipole moment tells us how polar a molecule is.  If we know the polarity of each bond and the shape of the molecule, we can determine if it is a polar molecule.

31. Practice  Which of the following molecules are polar? If they are polar, where is the dipole? CCl 4, SF 4

32. Assignment  Page 377 Exercise 8.12  Page 386 Exercise 8.14  Page 386 Exercise 8.15  Page 397 28, 30, 38

33. Bond Order  Bond order is the number of pairs of electrons shared between two atoms.

34. Bond Length  Bond Length is the distance between the nuclei of two bond atoms.  It is determined by the size of the atoms and the order of the bond.  We can use the trend in atomic size to predict the bond length.  Practice: Which bond is longer? C-F, C-O, or C-N?  Multiple bond decrease the bond length.

35. Bond Dissociation Enthalpy  There is an enthalpy change when a bond is broken.  The enthalpy change is always positive when breaking bonds. It is an endothermic process. “It takes energy to break bonds.”  Question: Which bond would take more energy to break? C-C, C=C or C ≡ C?  These enthalpies can be used to find the enthalpy change for a reaction.  Δ r H = Σ ΔH(bonds broken) - Σ ΔH(bonds formed)

36. Example  Calculate the enthalpy change for acetone reacting with hydrogen to form isopropanol.  Which bonds are broken? Which are formed? + H 2 →

37. Assignment  Page 388 Exercise 8.16  Page 391 Exercise 8.17  Page 398 42, 44, 48, 50, 52, 58, 64, 70