1. 1 Adventures of Oxygen Clip

2. 1. Compare & contrast ionic and covalent bonds in terms of electron position. 2. Predict formulas for stable binary ionic compounds based on balance of charges. 3. Use IUPAC nomenclature for transition between chemical names and chemical formulas of - binary ionic compounds - binary covalent compounds 4. Apply the Law of Conservation of Matter by balancing the following types of chemical equations: Synthesis Decomposition Single Replacement Double ReplacementGOALS 2 34 16 3. Determine the Types of ions formed by representative elements

3. Why do Atoms Form Compounds? Stability. What makes an atom stable? Full outer energy level. –E–Eight. They can either…… –1–1) Gain electrons –2–2) Lose electrons –3–3) Share electrons 3

4. A Chemical Bond holds atoms together in a compound. Two basic typesTwo basic types: 1. Ionic 2. Covalent 4

5. Ionic BondingT r a n s f e r o f e l e c t r o n s f r o m o n e a t o m t o a n o t h e r a t o m. O c c u r s b e t w e e n m e t a l s & n o n m e t a l s. Remember: Atoms need a full outer energy level to be stable. EIGHT! 5 C a l l e d c o m p o u n d s.

6. O c c u r s b e t w e e n m e t a l s a n d n o n m e t a l s. M e t a l s a r e e l e c t r o n d o n o r s. S O, t h e y b e c o m e P O S I T I V E N o n - m e t a l s a r e e l e c t r o n a c c e p t e r s. S O, t h e y b e c o m e N E G A T I V E. 6 OPPOSITS ATTRACT!

7. When Atoms gain or lose electrons, they are called Ions. Cation Anion 3P 7

8. Metals lose electrons to become stable. Nonmetals gain electrons to become stable. 8

9. Atoms can gain or lose electrons Ionization: requires energy Why do atoms lose and gain electrons? To become more stable. Stability=full outer energy level 9

11. OPPOSITS ATTRACT! 11

12. Properties of Ionic Compounds Crystalline solids at room temperature. Arranged in repeating three- dimensional patterns Have high melting points Can conduct electricity when melted or dissolved in water 12

13. Ionic Bonding CLIP 13

14. Covalent Bonding T h e s h a r i n g o f e l e c t r o n s b e t w e e n a t o m s. O c c u r s b e t w e e n n o n m e t a l s a n d n o n m e t a l s. C a l l e d M o l e c u l e s. 14

15. 15

16. Hydrogen and Fluorine Hydrogen and Chlorine 16

17. 17 Single, Double, Triple 2 e- 4e- 6e-

18. Clip 18

19. Unequal Sharing δ+δ+ Called Polar δ_δ_ 19 Polar molecules happen when one atom has a greater positive charge

20. Properties of Covalent Molecules Many are gases or liquids at room temperature Composed of two nonmetals. Have low melting and boiling points 20

21. Ionic and Covalent Bonding Review ClipClip

22. 1.CO 2 2.NaCl 3.H 2 O 4.MgCl 2 5.NO 2 6.Li 2 S 7.NaF 9.BeO 10.HCl 11.NaF 12.KCl 13.H 2 O 2 14.N 2 15.Cl 2 clip 21 Covalent or Ionic? (write the formula, then write “C” or “I”

23. Goals revisited

24. Writing chemical formulas is a shorthand way of indicating what a substance is made of. These formulas also let you know how many atoms of each type are found in a molecule. The chemical formula for water is H 2 O. Carbon Dioxide is CO 2. Why does oxygen combine in different ratios, in different compounds? The chemical formula for table salt is NaCl. Calcium Chloride is CaCl 2. Why does chlorine combine in different ratios, in different compounds? 22

25. The simplest compounds are ones with only two elements These are called binary KI, CO, H 2 O, NaCl 23

26. +1 +2 -2 -3+3 +4 -4 0 Oxidation numbers Tell you how many electrons an atom must gain, lose or share to become stable. 2424

27. We can predict the ratio of atoms in ionic compounds based on their oxidation numbers K Cl +1 KCl Tells you how many electrons an atom must gain, lose or share to become stable. 1 valence electron 7 valence electron All compounds are neutral That means the overall charge is ZERO! 25

28. Subscripts show the number of atoms of that kind in the compound Na Br +1 NaBr Ca Br +2 CaBr 2 To make it ZERO, you need 1 Ca & 2 Br. 26

29. Now You Try writing Binary Ionic formulas 1.K + Br 2.Mg + Cl 3.Ca + I 4.K + O 5.K + I 6.Sr + Br 7.Na + O 8.Ga + Br 9.Mg + O 10.Al + P 27

30. Some elements have more than one oxidation number (Chart p588) Fe O +3 -2 Fe 2 O 3 Fe O +2-2 FeO We call these elements- Multivalent Elements 28

31. Multivalent Practice 1.Fe +2 + O 2.Fe +3 + O 3.Cu +2 + F 4.Cr +3 + O 29

32. Cations:ammonium, NH 4 + Anions: nitrate, NO 3 - sulfate, SO 4 2- hydroxide, OH - phosphate, PO 4 3- carbonate, CO 3 2- chlorate, ClO 3 - permanganate, MnO 4 - chromate, CrO 4 2- Polyatomic Ions Groups of Covalently Bonded atoms that stay together. 30

33. Try these…… p591 1. Na + SO 4 2. Mg + PO 4 3. Ca + CO 3 4. Na + OH 5. Mg + OH 6. NH 4 + OH 7. K + PO 4 8. NH 4 + NO 3 9. H + SO 4 10. Ca + SO 4 11. K + NO 3 12. Na + PO 4 31 Mixed Practice

35. Naming Binary Compounds and Molecules Steps: –I–If it is Binary- 1.Decide if it is an ionic or covalent bond. –M–Metal- nonmetal….. »I»Ionic –N–Nonmetal- nonmetal…. »C»Covalent 32 Example: NaCl

36. If ionic ……. 2.C heck to see if any elements are multivalent or polyatomic. 3.I f all single valent, write the name of the positive ion first. 4.W rite the root of the negative ion and add –ide. Examples: 1.NaCl 2.K 2 O 3.AlCl 3 4.BaF 2 5.KI 6.Li 2 O 33

37. If ionic ……. 5.Check to see if any elements are multivalent. 6.If multivalent ions, determine the oxidation number of the element. 7.Use Roman numerals in parentheses after the name of the element. 8.Write the root of the negative ion and add –ide. Examples: 1.FeO 2.Fe 2 O 3 3.CuO 4.Cu 2 O 5.PbCl 4 6.PbI 2 34

38. If it contains a polyatomic ion... 2.Write the name of the positive ion. 3.Write the name of the polyatomic ion. Examples: 1.NaCO 3 2.KNO 3 3.NaC 2 H 3 O 2 Example: KOH Potassium Hydroxide CaCO 3 Calcium Carbonate 35

39. If Covalent... 2.Use Greek prefix to indicate how many atoms of each element are in the molecule 3.Add -ide to the more electronegative element Greek Prefixes 1- mono- 2- di- 3- tri- 4- tetra- 5- penta- 6- hexa- 7- hepta- 8- octa- Example: NO Nitrogen Monoxide PCl 3 Phosphorous trichloride 36

40. Name the following: Mixed Practice 1.KBr 2.HCl 3.MgO 4.CaCl 2 5.H 2 O 6.NO 2 7.CuSO 4 8.CaSO 4 9.NH 4 OH 10.CaCO 3 11.Cu(ClO 3 ) 2 12.Cr 2 O 3 13. FeO 14. LiBr 15 MgCl 37

41. 38 Goals revisited

42. Chemical Reactions A chemical reaction is a change in which one or more substances are converted into new substances. –R–R–R–Rearrangement of bonds in compounds and molecules. Chemical Equations make it possible to see clearly what is happening during a chemical reaction 39

43. Chemical equations are a shorthand way to show chemical reactions. ReactantsProducts H 2 + O 2 H 2 O 40

44. Conservation of Mass T h e m a s s o f t h e p r o d u c t s a l w a y s e q u a l s t h e m a s s o f t h e r e a c t a n t s 41

45. H 2 + O 2 H2OH2O 2 Hydrogen atoms 2 Oxygen atoms 2 Hydrogen atoms & one Oxygen atom Does this meet the Conservation of Mass Law? Must Balance the Equation to show Conservation of Mass. 42

46. 2 H 2 + O 2 H2OH2O 2 4122 2 4 Balanced!! Can add coefficients to Balance equations. Steps: 1. Count Atoms on both sides 2. If not Balanced, add coefficients to balance. 3. Recount atoms after adding each coefficient. 2 4. Keep adding coefficients until balanced. 43

47. 44

48. 45 Chemical Reactions You start with one or more compounds and turn it into different compounds. Vapors of hydrogen chloride in a beaker and ammonia in a test tube meet to form a cloud of a new substance, ammonium chloride.

49. 1.Synthesis 2.Decomposition 3.Single Replacement (Single Displacement) 4.Double Replacement (Double Displacement) 46

50. Synthesis “to make” A + B AB Cu + O CuO 2H 2 + O 2 2H 2 O

51. Decomposition “to breakdown” 2H 2 O 2H 2 + O 2 AB A + B NaOH Na + OH Animation

52. Single Replacement When one element replaces another element in a compound A + BC AC + B Cu+AgNO 3 Cu(NO 3 ) 2 + 2Ag The more reactive metal will always replace the less reactive metal. (p749)

53. Clip Single Replacement

54. Double Replacement Positive Ion of One compound replaces the positive ion of another compound and a Precipitate is formed. AB + CD AD + CB Ba(NO 3 ) 2 +KSO 4 2KNO 3 + BaSO 4 Review Clip

55. Animation

56. Review Clip

58. Chemical Reactions and Energy All chemical reactions release or absorb energy. –H–Heat, light, sound Chemical reactions are the making and breaking or bonds. 51

59. 1. Exergonic Chemical reactions that releases energy are called exergonic. –G–Glow sticks If heat is released, it is called exothermic.

60. 2. Endergonic Chemical reactions that require energy are called endergonic. Ex: Cold Packs If heat is absorbed, it is called endothermic

61. Catalysts and Inhibitors Some reactions proceed slowly. They can be sped up by a catalysts. –C–Catalysts are not used up in the reaction. –E–EX: enzymes (biological catalysts) Some reactions proceed too fast. They can be slowed down by inhibitors. –E–EX: Preservatives in food

62. 1. Compare & contrast ionic and covalent bonds in terms of electron position. 2. Predict formulas for stable binary ionic compounds based on balance of charges. 3. Use IUPAC nomenclature for transition between chemical names and chemical formulas of binary ionic compounds binary covalent compounds 4. Apply the Law of Conservation of Matter by balancing the following types of chemical equations: Synthesis Decomposition Single Replacement Double Replacement GOALS Revisited…..