1. Ch. 6 & 7 - Covalent Bonding Molecular or Covalent Compounds (p. 164 – 17 5, 211 – 213) Sharing Valence Electrons Nonmetals Only

2. Molecules Water molecule H 2 O Oxygen molecule, O 2 Sucrose molecule, C 12 H 22 O 11

3. Molecular Nomenclature Prefix System (binary compounds) 1.Less e - neg atom comes first. 2.Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 3.Change the ending of the second element to -ide.

4. PREFIXmono-di-tri-tetra-penta-hexa-hepta-octa-nona-deca-NUMBER12345678910 Molecular Nomenclature

5. Name These N2ON2ON2ON2O NO 2 Cl 2 O 7 CBr 4 CO 2 BaCl 2

6. Write formulas for these diphosphorus pentoxide tetraiodine nonoxide sulfur hexaflouride nitrogen trioxide Carbon tetrahydride phosphorus trifluoride aluminum chloride h.w. p. 235 # 10,11

7. The Terminology The smallest entity of a covalent compound is called a ___________. A bond is formed when one pair of electrons is shared. Some elements,( nitrogen, carbon, and oxygen) can share more than one pair of electrons and form bonds.

8. The Terminology A bond is produced when two pairs of electrons are shared. A bond is produced when three pairs of electrons are shared. Covalent bonding can also occur between two atoms of the same nonmetal. This forms a molecule.

9. N OF Cl Br I H Diatomic Elements There are seven diatomic elements: H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2

10. Lewis Structures Octet Rule –Most atoms form bonds in order to obtain 8 valence e - –Full energy level stability ~ Noble Gases Ne

11. Lewis Structures Electron dot notations that represent compounds. Pairs of dots represent shared electrons between bonded atoms And unshared pairs or lone pairs that belong exclusively to one atom

12. Drawing Lewis Structures Determine the type and number of atoms in the molecule. Determine the type and number of atoms in the molecule. Write the electron dot notation for each atom. Write the electron dot notation for each atom. Determine the total number of valence electrons in the atoms

13. Drawing Lewis Structures Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is central - otherwise the least electronegative element is central. (hydrogen is never central) Then connect the atoms by electron pair bonds.

14. Drawing Lewis Structures Add lone pairs so that hydrogen has a duet and each other nonmetal has an octet Add lone pairs so that hydrogen has a duet and each other nonmetal has an octet Count the electrons to be sure that the number of valence electrons used is equal to the number available.

15. Drawing Lewis Structures single bonds CF 4 1 C × 4e - = 4e - 4 F × 7e - = 28e - 32e - 32e - F F C F F Each wants an octet: 1 C × 8e - = 8e - 4 F × 8e - = 32e - 40e - 40e - -32e - -32e - must share 8e -

16. Practice Single Bonds CH 3 I H 2 F 2 H 2 O HCl NH 3 H 2 S CH 4 H.W. p. 197 # 39

17. Drawing Lewis Structures multiple bonds Multiple bonds between nitrogen, carbon, and oxygen are possible

18. Drawing Lewis Diagrams multiple bonds CO 2 1 C × 4e - = 4e - 2 O × 6e - = 12e - 16e - 16e - O C O Each wants an octet: 1 C × 8e - = 8e - 2 O × 8e - = 16e - 24e - 24e - -16e - -16e - must share 8e -

19. Practice Multiple Bonds C 2 H 2 CH 2 O HCN CH 4 CCl 4 h.w. p. 197 #41

20. Polyatomic Ions ( are held together by covalent bonds) To find total # of valence e - : –Add 1e - for each negative charge. –Subtract 1e - for each positive charge. Place brackets around the ion and label the charge.

21. Polyatomic Ions ClO 4 - 1 Cl × 7e - = 7e - 4 O × 6e - = 24e - 31e 31e O O Cl O O + 1e - + 1e - 32e - 32e - Each wants an octet: 1 Cl × 8e - = 8e - 4 O × 8e - = 32e - 40e - 40e - -32e - -32e - must share 8e -

22. NH 4 + 1 N × 5e - = 5e - 4 H × 1e - = 4e - 9e - 9e - H H N H H - 1e - - 1e - 8e - 8e - 4 H × 2e - = 8e - 1 N × 8e - = 8e - 16e - 16e - - 8e - - 8e - must share 8e - Polyatomic Ions

23. Practice Polyatomic Ions PhosphateSulfateNitrateChlorateChloriteHypochlorite h.w. p. 197 #42

24. Resonance Structures Molecules that can’t be correctly represented by a single Lewis diagram. Actual structure is an average of all the possibilities. Show possible structures separated by a double-headed arrow.

25. Acids Acids –Compounds that form H + in water. –Formulas usually begin with ‘H’. Examples: –HCl – hydrochloric acid –HNO 3 – nitric acid –H 2 SO 4 – sulfuric acid

26. Naming Binary acids Includes hydrogen and another nonmetal use the prefix hydro- and change -ide to - ic acid HCl - hydrogen ion and chloride ion hydrochloric acid H 2 S hydrogen ion and sulfide ion hydrosulfuric acid

27. Naming Oxyacids Includes hydrogen and a polyatomic ion containing oxygen Do not include hydro- in the name. If the polyatomic ion ends in -ate, change it to -ic acid HNO 3 - Hydrogen and nitrate ions Nitric acid If the polyatomic ion ends in -ite, change it to -ous acid HNO 2 Hydrogen and nitrite ions Nitrous acid

28. Acid Nomenclature oxyacidsbinary acids

29. Name these acids HF H3PH3PH3PH3P H 2 SO 4 H 2 SO 3 HCN H 2 CrO 4

30. Writing Formulas Binary Acids hydro- nonmetal root-ic acid Will include hydrogen and a nonmetal Criss-cross to make the charges cancel out. (the charge on the other element becomes the subscript for hydrogen) Ex. Hydrochloric acid Hydrogen and chlorideHCl

31. Writing Formulas Oxyacids If acid end in –ic, then polyatomic ion ends in –ate. If acid end in –ous, then polyatomic ion ends in –ite. Criss-cross to make the charges cancel out. (the charge on the polyatomic ion becomes the subscript for hydrogen) Ex. Iodic acid hydrogen and iodateHIO 3 hydrogen and iodateHIO 3

32. Write formulas for these hydrobromic acid hydrofluoric acid phosphoric acid bromic acid Hypochlorous acid hydroiodic acid acetic acid carbonic acid phosphorous acid Oxalic acid h.w. p. 235 #14,15

33. Bell Ringer p. 235 # 12