2. AP Chapter 7 Covalent Bonding

3. Lewis Structures A Lewis structure shows the distribution of outer (valence) electrons in an atom, molecule, or polyatomic ion. Unshared electrons are shown as dots, bonds are shown as straight lines

5. Transparency #1 In H 2 O and HF, as in most molecules and polyatomic ions, nonmetal atoms except H are surrounded by eight electrons, an octet. In this sense, each atom has a noble gas structure Lewis structures are written following a stepwise procedure

6. A) Rules for writing Lewis Structures (single bonds) 1.Count valence electrons available. Number of valence electrons contributed by a nonmetal atoms is equal to the last digit of its group number in the periodic table (one for H) Add electrons to take into account negative charge

7. OCl - ion: 6 + 7 + 1 = 14 valence e - CH 3 OH molecule: 4 + 4(1) + 6 = 14 valence e - SO 3 -2 ion: 6 + 3(6) + 2 = 26 valence e -

8. 2. Draw skeleton structure, using single bonds Transparency #2 (next slide) Note that carbon almost always forms four bonds. Central atom is written first in formula, terminal atoms are most often H, O, or a halogen

10. 3. Deduct two electrons for each single bond in the skeleton OCl - ion: 14 – 2 = 12 valence e - left CH 3 OH molecule: 14 – 10 = 4 valence e - left SO 3 -2 ion: 26 – 6 = 20 valence e- left

11. 4. Distribute these electrons to give each atom a noble gas structure, if possible Transparency # 3

13. Examples of central atom POCl 3 P is central atom SO 4 -2 S is central atom SO 3 -2 S is central atom PO 4 -3 P is central atom SCl 2 S is central atom

14. Draw Lewis Dot Structure Ethane C 2 H 6 Check answer on page 167

15. B. Too few electrons; form multiple bonds Structure of NO 3 - ion? Number of valence electrons = ? 5 + 18 + 1 = 24 Go to transparency #4

16. Exceptions to the octet When we must exceed the octet, extra electrons go on central atom. (expanded octet) ClF 3 XeO 3 ICl 4 - BeCl 2

17. Consider XeF 4, 36 valence electrons. Octet structure uses 32 electrons. Put extra e - around Xe Transparency 5 In a few molecules, there are less than eight electrons around the central atom Transparency 5

18. Resonance Sometimes there is more than one valid structure for an molecule or ion. NO 3 - Transparency 6 True structure is a hybrid of those three forms Use double arrows to indicate it is the “average” of the structures. It doesn’t switch between them.

19. Note that….. Resonance forms are obtained by moving electrons, not atoms. Resonance can be expected when is possible to draw more that one structure that follows the octet rule

20. VSEPR Lewis structures tell us how the atoms are connected to each other. They don’t tell us anything about shape. The shape of a molecule can greatly affect its properties. Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

21. VSEPR Molecules take a shape that puts electron pairs as far away from each other as possible. Have to draw the Lewis structure to determine electron pairs. bonding nonbonding lone pair Lone pair take more space. Multiple bonds count as one pair.

23. Octahedral

24. Atomic Orbitals - Hybridization Valence bond theory modified In molecules, the orbitals occupied by electron pairs are seldom “pure” s or p orbitals Instead, they are “hybrid” orbitals, formed by combining s, p, d orbitals

25. Formation of hybrid orbital s orbital + p orbital  two sp hybrid orbitals –Ex. Be in BeF 2 s orbital + two p orbitals  three sp 2 hybrid orbitals –Ex. B in BF 3 s orbital + three p orbitals  four sp 3 hybrid orbitals –Ex. C in CH4

26. Hybridization with 5 or 6 electron pairs –sp 3 d, sp 3 d 2

27. I will put samples on board, then you check answers

28. Multiple bonds The extra electron pairs in a multiple bond are not located in hybrid orbitals (One pair in a double bond, two pairs in a triple bond) Pg. 186 Example 7.10

29. Sigma and Pi bonds Sigma bond – single “lobe” All single bonds are sigma bonds (σ) Unhybridized electron pairs in multiple bonds have different shape, pi bond One of the electron pairs in a multiple bond is a sigma bond; the others are pi bonds Pg. 187 Example 7.11