1. Chapter 9 Covalent Bonding

2. This chapter is hard You must do your homework and study every day You must know your polyatomics and be able to write chemical formulas You must learn what I tell you to learn

3. Covalent Bonding In ionic bonding, we talked about the transferring of electrons. Covalent bonding involves the sharing of pairs of electrons. Atoms that bond covalently are called molecules. Covalent bonds occur nonmetal to nonmetal. It can involve a metalloid. In covalent bonding all atoms involved contribute.

4. Covalent Bonding and the Periodic Table Because of the number of valence e- in each group, nonmetals in certain groups will form a certain number of bonds. Group 17 will share a pair of e-. – Ex. F 2 Group 16 will share 2 pair of e-. – Ex. O 2 Group 15 will share 3 pair of e-. – Ex. NH 3

5. Types of Bonds Single bond – Also called a sigma bond( ) – Sharing of a single pair of electrons – In a sigma bond, the s orbital holding the shared e- will be centered between two atoms and overlap end to end.

6. Double bond – 2 shared pair of e- – 1 sigma bond and 1 pi bond Pi bond (π) – The orbitals containing the shared electrons overlap above and below the plane of the sigma bond. – Pi bonds cannot be alone; They must accompany a sigma bond.

7. Triple bond – 3 shared pair of e- – Contains 2 pi bonds and one sigma bond.

9. Bonding examples – Draw and model C 2 H 4 C 2 H 2

10. Bond Length and Strength The shorter the length the stronger the bond. Triple bonds are the shortest and have the greatest strength. Double is smaller than single. Single is the longest and weakest.

11. Bond Association Energy The total energy of a chemical reaction is determined by bonds formed and bonds broken. Determining type of reaction – Endothermic vs Exothermic Formula: energy of bonds broken – energy of bonds formed = net energy of products. If the value is positive it is an endothermic reaction. If the value is negative it is an exothermic reaction.

12. Naming Binary Molecular Compounds Use prefixes listed on page 248 – You must learn these – Name root-ide – Add prefixes Do not add mono- to the first word Do not use double vowels – No i-i, o-o, a-a

13. Examples CO 2 – Carbon dioxide CO – Carbon monoxide CCl 4 – Carbon tetrachloride As 2 O 3 – Diarsenic trioxide Pg. 249 – H 2 O – Water – NH 3 - Ammonia

14. Acids and Bases Acids – Arrhenius acid – H+ – Bronsted-Lowry acid – H+ donor Bases – Arrhenius base – OH- – Bronsted-Lowry base – H+ acceptor

15. Naming Acids Acids always contain H and they are in an aqueous solution – They are different as a gas There are 2 types – Binary acids that do not contain oxygen H and a monoatomic anion in aqueous solution Exception – Polyatomic ions without O – Oxyacids H and an O

16. Naming Binary Acids Hydro- root of anion –ic acid – Exception You always use the full name for sulfur – HBr (aq) – hydrobromic acid – HCl (aq) – hydrochloric acid – H 2 S (aq) – hydrosulfuric acid – Exception: Polyatomic without O HCN (aq) – hydrocyanic acid

17. Naming Oxyacids H and an oxyanion DO NOT USE HYDRO Polyatomic anions end in –ate or –ite. This indicates the number of O atoms in the polyatomic. You will use this for naming oxyacids. If the polyatomic ends in –ate you will name the acid ending with –ic. If the polyatomic ends in –ite you will name the acid ending with –ous. Example – HNO 3 – nitric acid – HNO 2 – nitrous acid

18. Acids that contain Halogens Polyatomics that contain a halogen have 4 names. Example: perchlorate, chlorate, chlorite, hypochlorite. When naming these you will apply the same rules of –ate to –ic and –ite to –ous AND indicate per- and hypo-. Example – HClO 4 – perchloric acid – HClO 3 – chloric acid – HClO 2 – chlorous acid – HClO – hypochlorous acid

19. Name These Acids HIO 2 HBr HBrO 3 H 2 SO 3 H 2 CO 3 HF HIO 4 Iodous acid Hydrobromic acid Bromic acid Sulfurous acid Carbonic acid Hydrofluoric acid Periodic acid

20. Naming Hydroxides Cation and hydroxide – Most are metals; ammonium is not. – Examples NaOH – Sodium hydroxide NH 4 OH – Ammonium hydroxide Mg(OH) 2 – Magnesium hydroxide

21. Groups Pg 250 – 18-22

23. Naming Covalent Compounds Practice Worksheets – Covalent compounds and Acids

24. Pg 272 88-98 Pg 274 128

26. Molecular Structures Fact you must know – Lone or nonbonding pairs of electrons require more space than bonding pairs. Rules for depicting structural formula Example: NH 3 1)Predict atom location – Terminal atoms – comes off of the central atom(s) H is always terminal – Central atoms – usually located closest to the middle of the periodic table. (Only nonmetals – We are talking covalent bonding) Everything will stem off of it

27. 2.Get the sum of all valence e- – NH 3 5e- + 3(1e-) = 8e- Divide this number by two in order to get number of pairs of electrons NH 3 has 4 pair 3.Use one pair to bond each terminal atom to the central atom H N H H

28. 4.Equally distribute remaining pairs to the terminals. Then, add any remaining pairs to the central atom. – Make sure each atom has 4 pair. Why? – What about H?

29. .. H N H H

30. 5.If the central does not have 4 pairs of e- then make multiple bonds. Example: CO 2

32. Polyatomic Lewis Structures Follow the same steps given for structural formula. The only difference is you must indicate charge by using [ ] - or [ ] +. Examples – ClO 4 - – NH 4 +

33. Resonance Structures What is resonance? – Guitar string – Something vibrating back and forth very quickly – Resonance structures occur when there is more than one Lewis structure for a molecule or compound. Example: NO 3 - and SO 2 In NO 3 -, one bond resonates between 3 places. You must draw each resonating structure.

35. Exceptions to the Octet Rule Some atoms have odd numbers of electrons – Example: NO 2 has 17 e- Boron tends to form 3 bonds when it is the central atom. – Example: BH 3 Expanded Octets – Some atoms form more than 4 bonds on the central atom. Almost always involves halogens PCl 5

36. Positions to know when looking at shape Cl P Cl P Cl Equatorial Position Cl Axial Position

37. You must learn page 260. All of it Group Work – Pg 255 #30-34

39. Pg. 273 – 99-104

40. Naming Sheet – No book – No notes

41. Molecular Shape Linear – Usually formed by atoms with multiple bonds Example: CO 2 Trigonal Planer – Flat – All angles are equal (120) Example: BH 3 Tetrahedral – There are four equal bonding sites – All angles are equal (109.5) Example: CH 4

42. The next four are deformation of the tetrahedral. Trigonal Pyramidal – Involves a lone pair. Lone pairs take up more space than bonded pairs. – The angles between the terminal atoms decreases from the angle of the tetrahedral (107.3). – Example: NH 3

43. Bent – 2 lone pairs – The angle decreases more (104.5) – Example: H 2 O Trigonal Bypyramidal – What does bi mean 2 pyramids – Only occurs with elements in period 3 or greater – Must have d orbitals – Has 2 different angles: 90 and 120

44. Octahedran – All bond angles are equal (90) – Forms 6 bonds and 8 (octa-) faces – Example: SF 6

45. Hybridization Hybrid orbitals – Mixing of orbitals to form new orbitals Example: C – Show with electron configuration and orbital notation – Carbon tends to form 4 bonds – The outer most energy level is 2 – Forms an sp 3 hybrid

46. Bond Polarity Bonds have oppositely charged ends Related to electronegativity in a chemical bond – The higher the electronegativity value the more strongly the atom pulls electrons to itself in a chemical bond. Pg. 263 In a bond between F and H which atom will pull more on an electron? – F – 3.98 and H – 2.20

47. Because F has a greater pull on the electron it creates polarity. How to represent polarity. – H F or – Is the symbol for indicating a partial charge

48. Compare the electronegativity values of the bonded atoms one bond at a time. Rules for determining polarity – If the difference is less than or equal to 0.4 then the bond is nonpolar. – If the difference is between 0.5 and 1.7 then the bond is polar. – If the bond is greater than 1.7 then the bond is ionic. The only true covalent bond exist between bonded identical atoms. In order for a molecule to be polar you must have positive and negative ends. How does symmetry of a molecule relate to polarity?

49. Polarity Example Problem H 2 O

50. Homework Pg 273 105-114