1. TYPES OF CHEMICAL REACTIONS Unit 6 AP CHEMISTRY

2. AIM: Precipitate Reactions DO NOW: 1. Take out the worksheet on predicting the reactions (from before midterm exams) Go to your note sheet that you printed out on Types of Chemical Reactions and Solution Stoichiometry Exercise 8 (pg. 8) predict the products of the three double replacement reactions ONLY predict the PRODUCTS

3. AIM: Precipitate Reactions Get your 3D- glasses! 3D Video – Physical vs. Chemical Changes Neutralization RedOx

4. CHEMICAL EQUATIONS chemical reaction--transforms elements and compounds into new substances balanced chemical equation--shows the relative amounts of reactants and products s, l, g, aq--solid, liquid, gas, aqueous solution NO ENERGY or TIME is given Antoine Lavoisier (1743-1794)--law of conservation of matter: matter can neither be created nor destroyed this means “balancing equations”

5. PRECIPITATION REACTIONS Precipitate is a driving force for a chemical reaction A precipitate is an insoluble solid that is formed when two aqueous solutions are mixed separate the precipitate (ppt) from solution by filtration in what is called a gravimetric analysis.

6. SOLUBILITY RULES!!!!!!

7. Practice w/ Solubility Rules Using the solubility rules, predict what will happen in the examples from your worksheet

8. Exercise 8 Using the solubility rules, predict what will happen when the following pairs of solutions are mixed. KNO3(aq) and BaC12(aq)

9. Exercise 8 Na2SO4(aq) and Pb(NO3)2(aq)

10. Exercise 8 KOH (aq) and Fe(NO3)3(aq)

11. DESCRIBING REACTIONS IN SOLUTIONS COMPLETE balanced equation—gives the overall reaction stoichiometry, but NOT the forms of the reactants & products as they exist in solution complete ionic equation—represents as IONS all reactants & products that are strong electrolytes net ionic equations—includes only those solution components undergoing a change. Spectator ions are NOT included spectator ions--not involved in the reaction process Λ started as an ion AND finished as an ion THERE IS ALWAYS A CONSERVATION OF CHARGE IN NET IONIC EQN’S.

12. EXERCISE 9 - For each of the following reactions, write the molecular equation, the complete ionic equation, and the net ionic equation. Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate plus aqueous potassium nitrate.

13. EXERCISE 9 Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to form a precipitate of iron(III) hydroxide and aqueous potassium nitrate.

14. ACID BASE REACTIONS DO NOW: 1.1. Write out the word equation for what happens when an acid and base react 2.2. Give examples of strong base/strong acid

15. ACID – BASE REACTIONS acids--any cmpd. that, on reaction with water, produces an ion called the hydronium ion, H3O+ [or H+], and an anion (Arrhenius definition) base--any cmpd. that provides a hydroxide, OH−,and a cation in water (Arrhenius definition) **ammonia, NH3 is an exception, so Bronsted-Lowry defined it as a proton acceptor!! neutralization—when moles acid = moles base each is neutralized [pH is not necessarily 7.0]. The products formed are a salt [ask yourself if it is soluble] and water

16. SOLUBILITY RULES These are strong electrolytes (100% ionized ) and written as ions 1. Strong Acids: HCl, HBr, HI, H2SO4, HNO3, HClO4, HClO3 Sulfuric acid (strong acid) can be written as H+ and SO42- or as H+ and HSO4-. If it says CONCENTRATED write it as a molecule***** 2. Strong Bases: Hydroxides of group IA and IIA(Ba, Sr, Ca are marginal Be and Mg are WEAK)

17. Rules for Acid – Base Reactions Know SOLUBILITY RULES!!!!! Ionize strong acids and bases Watch out for concentrated sulfuric acid

18. Example A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound as been added. 1.Write out reaction: H 2 SO 4 + Ba(OH) 2  BaSO 4 +H 2 O H + + SO 4 2- + Ba 2+ + OH -  BaSO 4 + H 2 O

19. Example Solutions of ammonia and hydrofluoric acid are mixed

20. Example Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide

21. Example A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound has been added

22. Example A solution of sodium hydroxide is added to a solution of sodium dihydrogen phosphate until the same number of moles of each compound has been added

23. Example Dilute nitric acid is added to crystals of pure calcium oxide

24. Example Equal volumes of 0.1 molar sulfuric acid and 0.1 molar potassium hydroxide are mixed

25. Example A solution of ammonia is added to a dilute solution of acetic acid

26. RedOx REACTIONS DO NOW: 1.1. What occurs in oxidation- reduction reactions? 2.2. What do the following mean? 3.OiL RiG 4.LEO says GER 5.3. Using the Rules for Assigning oxidation number on pg. 13 of Types of Chemical Reactions and Solution Stoichiometry do Exercise 16

27. Terms to Know OIL RIG – oxidation is loss, reduction is gain (of electrons) Oxidation – the loss of electrons, increase in charge Reduction – the gain of electrons, reduction of charge Oxidation number – the assigned charge on an atom Oxidizing agent (OA) – the species that is reduced and thus causes oxidation Reducing agent (RA) – the species that is oxidized and thus causes reduction

28. Rules for Assigning Oxidation Numbers

29. Exercise 17

30. Exercise 18

31. Balancing RedOx Reactions by Half Reaction Method Divide the equation into oxidation and reduction half reactions. [OILRIG] Balance all elements besides hydrogen and oxygen. Balance O’s by adding H2O’s to the appropriate side of each equation. Balance H’s by adding H+ Balance the charge by adding electrons. [OILRIG again] Multiply the half reactions to make electrons equal for both half-reactions. Cancel out any common terms and recombine the two half reactions. IF BASIC, neutralize any H+ by adding the SAME NUMBER of OH- to EACH side of the balanced equation. [This creates some waters that will cancel!] CHECK!!

32. Sample Problem Assign oxidation states to all atoms in the following equation, identify the oxidation and reduction half reactions, and the OA and RA. MnO 4− (aq) + Fe 2+ (aq)  Mn +2 (aq) + Fe 3+ (aq)

33. Sample Problem Balance the following equation using the half- reaction method. (acidic) MnO 4 − (aq) + I − (aq)  Mn +2 (aq) + I 2 (aq)

34. Sample Problem (basic) Ag(s) + CN − + O 2  Ag(CN) 2 − (aq)

35. Exercise 19

36. Exercise 20

37. Oxidation Reduction Reactions Redox reactions involve the transfer of electrons. The oxidation numbers of at least two elements must change. Single replacement, some combination and some decomposition reactions are redox reactions. To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and a reducing agent. When a problem mentions an acidic or basic solution, it is probably redox.

38. Oxidation Reduction Reactions

40. EXAMPLE Iron (III) ions are reduced by iodide ions

41. EXAMPLE Solution of tin (II) sulfate is added to a solution of iron (II) sulfate

42. EXAMPLE Metallic copper is heated with concentrated sulfuric acid

43. EXAMPLE Manganese (IV) is added to warm concentrated hydrobromic acid

44. EXAMPLE Chlorine gas is bubbled into cold dilute sodium hydroxide

45. EXAMPLE Solid iron (III) oxide is heated in excess carbon monoxide

46. EXAMPLE Hydrogen peroxide solution is added to acidified potassium iodide solution

47. EXAMPLE Potassium permanganate solution is added to concentrated hydrochloric acid

48. Decomposition Reactions Reactions where a compound breaks down into two or more elements or compounds. Heat, electrolysis, or a catalyst is usually necessary. A compound may break down to produce two elements. A compound may break down to produce an element and a compound. A compound may break down to produce two compounds.

49. Decomposition Reactions Metallic carbonates break down to yield metallic oxides and carbon dioxide, Metallic chlorates break down to yield metallic chlorides and oxygen. Hydrogen peroxide decomposes into water and oxygen. Ammonium carbonate decomposes into ammonia, water and carbon dioxide. Sulfurous acid decomposes into water and sulfur dioxide. Carbonic acid decomposes into water and carbon dioxide.

50. Example A solution of hydrogen peroxide is heated

51. Example Solid magnesium carbonate is heated

52. Example Solid ammonium carbonate is heated

53. Addition Reactions Two or more elements or compounds combine to form a single product. A group IA or IIA metal may combine with a nonmetal to make a salt. Two nonmetals may combine to form a molecular compound. The oxidation number of the less electronegative element is often variable depending upon conditions. Generally, a higher oxidation state of one nonmetal is obtained when reacting with an excess of the other nonmetal.

54. Addition Reactions When an element combines with a compound, you can usually sum up all of the elements on the product side. Two compounds combine to form a single product.

55. Addition Reactions A metallic oxide plus carbon dioxide yields a metallic carbonate. (Carbon keeps the same oxidation state) A metallic oxide plus sulfur dioxide yields a metallic sulfite. (Sulfur keeps the same oxidation state) A metallic oxide plus water yields a metallic hydroxide. A nonmetallic oxide plus water yields an acid.

56. Example The gases boron trifluoride and ammonia are mixed

57. Example A mixtures of solid calcium oxide and solid tetraphosphorus decaoxide is heated

58. Example Solid calcium oxide is exposed to a stream of carbon dioxide gas

59. Double Replacement Two compounds react to form two new compounds. No changes in oxidation numbers occur. All double replacement reactions must have a "driving force" that removes a pair of ions from solution.

60. Formation of a Precipitate A precipitate is an insoluble substance formed by the reaction of two aqueous substances. Two ions bond together so strongly that water can not pull them apart. You must know your solubility rules to write these net ionic equations Ex. Solutions of silver nitrate and lithium bromide are mixed. Ag + Br  AgBr

61. Formation of a Gas Gases may form directly in a double replacement reaction or can form from the decomposition of a product such as H2CO3 or H2SO3. Common gasses: CO2, SO2, SO3, H2S, NO2, NH3, O2, H2 Ex. Excess hydrochloric acid solution is added to a solution of potassium sulfite. H+ + SO3 2-  H20 + SO2

62. Formation of a Molecular Substance When a molecular substance such as water or acetic acid is formed, ions are removed from solution and the reaction "works". Ex. Dilute solutions of lithium hydroxide and hydrobromic acid are mixed. OH- + H+  H2O (HBr, HCI, and HI are strong acids) Ex. Gaseous hydrofluoric acid reacts with solid silicon dioxide. HF+SiO2  SiF4 +H2O

63. Example Hydrogen sulfide is bubbled through a solution of silver nitrate

64. Example An excess of sodium hydroxide solution is added to a solution of magnesium nitrate

65. Example Solutions of sodium iodide and lead nitrate are mixed

66. Single Replacement Reaction where one element displaces another in a compound. One element is oxidized and another is reduced. A + BC  B + AC

67. Single Replacement Active metals replace less active metals or hydrogen from their compounds in aqueous solution. Use an activity series or a reduction potential table to determine activity. The more easily oxidized metal replaces the less easily oxidized metal. The metal with the most negative reduction potential will be the most active. Ex. Magnesium turnings are added to a solution of iron(III) chloride. Mg + Fe 3+  Fe + Mg2+

68. Single Replacement Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Each halogen will displace less electronegative (heavier) halogens from their binary salts. Ex. Chlorine gas is bubbled into a solution of potassium iodide. Cl2 +I-  I2 +Cl-

69. Single Replacement Tricky redox reactions that appear to be ordinary single replacement reactions: Hydrogen reacts with a hot metallic oxide to produce the elemental metal and water. Ex. Hydrogen gas is passed over hot copper(II) oxide. H2+CuO  Cu2+ +H2O

70. Example Piece of aluminum metal is added to a solution of silver nitrate

71. Example Aluminum metal is added to a solution of copper (II) chloride

72. Example Hydrogen gas is passed over hot copper (II) oxide

73. Anhydrides Anhydride means "without water" Water is a reactant in each of these equations. Nonmetallic oxides (acidic anhydrides) plus water yield acids. Ex. Carbon dioxide is bubbled into water. CO2 + H2O  H2CO3

74. Anhydrides Metallic oxides (basic anhydrides) plus water yield bases. Ex. Solid sodium oxide is added to water. Na 2 O + H 2 O  Na + + OH -

75. Anhydrides Metallic hydrides (ionic hydrides) plus water yield metallic hydroxides and hydrogen gas. Ex. Solid sodium hydride is added to water. NaH + H 2 O  Na + + OH - + H 2

76. Anhydrides Phosphorus halides react with water to produce an acid of phosphorus (phosphorous acid or phosphoric acid) and a hydrohalic acid. The oxidation number of the phosphorus remains the same in both compounds. Phosphorus oxytrichloride reacts with water to make the same products. Ex. Phosphorus tribromide is added to water. PBr 3 + H 2 O  H 3 PO 3 + H + + Br - Group I&II nitrides react with water to produce the metallic hydroxide and ammonia. Amines react with water to produce alkylammonium ions and hydroxide ions Ex) Methylamine gas is bubbled into distilled water CH 3 NH 2 + H 2 O  CH 3 NH 3 + + OH -

77. Example Excess water is added to solid calcium hydride

78. Example Solid lithium hydride is added to water

79. Example Liquid phosphorus trichloride is poured into a large excess of water

80. Combustion Elements or compounds combine with oxygen. Hydrocarbons or alcohols combine with oxygen to form carbon dioxide and water. Ammonia combines with limited oxygen to produce NO and water and with excess oxygen to produce N02 and water.

81. Combustion Nonmetallic hydrides combine with oxygen to form oxides and water. Nonmetallic sulfides combine with oxygen to form oxides and sulfur dioxide. Ex. Carbon disulfide vapor is burned in excess oxygen. CS2 + O2  CO2 + SO2 Ex. Ethanol is burned completely 'in air. C2H5OH + O2  CO2 + H2O

82. Example Lithium metal is burned in air

83. Example The hydrocarbon hexane is burned in excess oxygen

84. Gaseous diborane B2H6 is burned in excess oxygen

85. Complex Ion Reactions Complex ion- the combination of a central metal ion and its ligands Ligand- group bonded to a metal ion Coordination compound- a neutral compound containing complex ions [Co(NH3)6]Cl3 (NH3 is the ligand, [Co(NH3)6]3+ is the complex ion)

86. Complex Ion Reactions

87. Example Concentrated (15M) ammonia solution is added in excess to a solution of copper (II) nitrate

88. Example Excess of nitric acid solution is added to a solution of tetraammine copper (II) sulfate

89. Example