1. 1 Chapter 11 Chemical Reactions

2. 2 All chemical reactions l have two parts l Reactants - the substances you start with l Products- the substances you end up with l The reactants turn into the products. Reactants  Products

3. 3 In a chemical reaction l The way atoms are joined is changed l Atoms aren’t created or destroyed. l Can be described several ways l In a sentence –Copper reacts with chlorine to form copper (II) chloride. l In a word equation Copper + chlorine  copper (II) chloride

4. 4 Symbols used in equations l Table 11.1 l the arrow separates the reactants from the products l Read “reacts to form” l The plus sign = “and” l (s) after the formula -solid l (g) after the formula -gas l (l) after the formula -liquid

5. 5 Symbols used in equations l (aq) after the formula - dissolved in water, an aqueous solution.  used after a product indicates a gas (same as (g))  used after a product indicates a solid (same as (s))

6. 6 Symbols used in equations l indicates a reversible reaction (More later) l shows that heat is supplied to the reaction l is used to indicate a catalyst used in this case, platinum.

7. 7 What is a catalyst? l A substance that speeds up a reaction without being changed by the reaction. l Enzymes are biological or protein catalysts.

8. 8 Skeleton Equation l Uses formulas and symbols to describe a reaction l doesn’t indicate how many. l All chemical equations are sentences that describe reactions.

9. 9 Convert these to equations l Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form solid iron (II) chloride and hydrogen sulfide gas.

10. 10 Convert these to equations l Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

11. 11 The other way Fe(g) + O 2 (g)  Fe 2 O 3 (s)

12. 12 The other way Cu(s) + AgNO 3 (aq)  Ag(s) + Cu(NO 3 ) 2 (aq)

13. 13 Balancing Chemical Equations

14. 14 Balanced Equation l Atoms can’t be created or destroyed l All the atoms we start with we must end up with l A balanced equation has the same number of atoms of each element on both sides of the equation.

15. 15 C + O 2  CO 2 l This equation is already balanced l What if it isn’t already? C + O O  C O O

16. 16 C + O 2  CO l We need one more oxygen in the products. l Can’t change the formula, because it describes what actually happens +  O C O O C O CC

17. 17 l Must have started with two C 2 C + O 2  2 CO +  O CC O CC l Must be used to make another CO l But where did the other C come from? O O

18. 18 Rules for balancing  Write the correct formulas for all the reactants and products  Count the number of atoms of each type appearing on both sides  Balance the elements one at a time by adding coefficients (the numbers in front)  Check to make sure it is balanced.

19. 19 Never l Change a subscript to balance an equation. –If you change the formula you are describing a different reaction. –H 2 O is a different compound than H 2 O 2 l Never put a coefficient in the middle of a formula –2 NaCl is okay, Na2Cl is not.

20. 20 Example H 2 +H2OH2OO2O2  Make a table to keep track of where you are at RP H O 2 2 2 1 Need twice as much O in the product H 2 +H2OH2OO2O2  2 2 Changes the OAlso changes the H 4 Need twice as much H in the reactant 2 Recount 4 The equation is balanced, has the same number of each kind of atom on both sides

21. 21 Example H 2 +H2OH2OO2O2  RP H O 2 2 2 1 2 2 4 2 4 This is the answer Not this

22. 22 Examples CH 4 + O 2  CO 2 + H 2 O

23. 23 Examples AgNO 3 + Cu  Cu(NO 3 ) 2 + Ag

24. 24 Examples Al + N 2  Al 2 N 3

25. 25 Examples P + O 2  P 4 O 10

26. 26 Examples Na + H 2 O  H 2 + NaOH

27. 27 Techniques l If an atom appears more than once on a side, balance it last. l If you fix everything except one element, and it is even on one side and odd on the other, double the first number, then move on from there. l C 4 H 10 + O 2  CO 2 + H 2 O

28. 28 Types of Reactions Predicting the Products

29. 29 Types of Reactions l There are too many reactions to remember l Fall into categories. l We will learn 5 types. l Will be able to predict the products. l For some we will be able to predict whether they will happen at all. l Must recognize them by the reactants

30. 30 #1 Combination Reactions l Combine - put together l A + B  AB l 2 elements, or compounds combine to make 1 compound. Ca +O 2  CaO SO 3 + H 2 O  H 2 SO 4 l We can predict the products if they are two elements. Mg + N 2 

31. 31 Write and balance Ca + Cl 2 

32. 32 Write and balance Fe + O 2  iron (II) oxide

33. 33 Write and balance Al + O 2  l Remember that the first step is to write the formula l Then balance l Also called synthesis reaction

34. 34 Combining two compounds l If they tell you it is combination, you will make one product l Two compounds will make a polyatomic ion. l CO 2 + H 2 O → l H 2 O + Cl 2 O 7 →

35. 35 #2 Decomposition Reactions l decompose = fall apart l AB  A + B l one reactant falls apart into two or more elements or compounds. l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2

36. 36 #2 Decomposition Reactions l Can predict the products if it is a binary compound l Made up of only two elements l Falls apart into its elements lH2OlH2O

37. 37 #2 Decomposition Reactions l HgO

38. 38 #2 Decomposition Reactions l If the compound has more than two elements you must be given one of the products l The other product will be from the missing pieces l NiCO 3 NiO + H 2 CO 3 (aq)  CO 2 +

39. 39 #3 Single Replacement l One element replaces another l A + BC  AC + B l Reactants must be an element and a compound. l Products will be a different element and a different compound. Na + KCl  K + NaCl F 2 + LiCl  LiF + Cl 2

40. 40 Na + KCl  K + NaCl Na K Cl

41. 41 F 2 + 2 LiCl  2 LiF + Cl 2 F Li Cl F Li Cl Li

42. 42 #3 Single Replacement l Metals replace metals (and hydrogen) Al + CuSO 4  Zn + H 2 SO 4  l Think of water as HOH l Metals replace one of the H, combine with hydroxide. Na + HOH 

43. 43 #3 Single Replacement l We can tell whether a reaction will happen l Some are more active than other l More active replaces less active l There is a list on page 333

44. 44 #3 Single Replacement l There is a list on page 333 l Higher on the list replaces lower. l If the element by itself is higher, it happens, l if element by itself is lower, it doesn’t

45. 45 #3 Single Replacement l Note the * l H can be replaced in acids by everything higher l Only the first 4 (Li - Na) react with water.

46. 46 #3 Single Replacement Al + HCl 

47. 47 #3 Single Replacement Fe + CuSO 4 

48. 48 #3 Single Replacement Pb + KCl 

49. 49 #3 Single Replacement Al + H 2 O 

50. 50 #3 Single Replacement l What does it mean that Ag is on the bottom of the list?

51. 51 #3 Single Replacement l Nonmetals can replace other nonmetals l Limited to F 2, Cl 2, Br 2, I 2 l The order of activity is that on the table. l Higher replaces lower. F 2 + HCl  Br 2 + KCl 

52. 52 #4 Double Replacement l Two things replace each other. l AB + CD  AD + CB l Reactants must be two ionic compounds or acids. l Usually in aqueous solution NaOH + FeCl 3  l The positive ions change place. NaOH + FeCl 3  Fe 3+ OH - + Na + Cl - NaOH + FeCl 3  Fe(OH) 3 + NaCl

53. 53 3NaOH + FeCl 3  Fe(OH) 3 + 3NaCl Na + O-O- H+H+ O-O- H+H+ O-O- H+H+ Fe 3+ Cl -

54. 54 #4 Double Replacement l Will only happen if one of the products –doesn’t dissolve in water and forms a solid –or is a gas that bubbles out. –or is a covalent compound usually water. l Polyatomic ions don’t change from side to side

55. 55 Complete and balance l assume all of the reactions take place. CaCl 2 + NaOH  CuCl 2 + K 2 S  KOH + Fe(NO 3 ) 3 

56. 56 Complete and balance KOH + Fe(NO 3 ) 3  H 3 PO 4 + Ca(OH) 2 

57. 57 How to recognize which type l Look at the reactants l E for element l C for compound l E + E Combination l CDecomposition l E + CSingle replacement l C + CDouble replacement

58. 58 Last Type l Combustion l If O2 is a reactant, it can be classified as combustion l A compound composed of only C H and maybe O is reacted with oxygen l If the combustion is complete, the products will be CO 2 and H 2 O. l If the combustion is incomplete, the products will be CO and H 2 O. l or just C and H 2 O. l O 2 will always be the second reactant

59. 59 Examples l Complete combustion of C 4 H 10 l Incomplete combustion of C 4 H 10

60. 60 Examples l Complete combustion of C 6 H 12 O 6 l Incomplete combustion of C 2 H 6 O

61. 61 Ionic Compounds and acids l Fall apart into ions when they dissolve l That’s why they conduct electricity when dissolved. l So when we write them as (aq) they are really separated l NaCl(aq) is really Na + (aq) and Cl - (aq) l K 2 SO 4 (aq) is really K + (aq) and SO 4 2- (aq)

62. 62 Reactions in aqueous solutions l Many reactions happen in solution l Makes it so the ions separate so they can interact. l Solids, liquids, and gases are not separated, only aqueous l You should be able to write the “molecular equation”, “complete ionic equation”, and “net ionic equation” l Balance with by mass and charge

63. 63 Complete Ionic Equation l Every aqueous compound is written as separate ions l Solids, liquids and gases as whole compounds l (molecular) MgCl 2 (aq) + PbSO 4 (aq) → MgSO 4 (aq) + PbCl 2 (s) l Is really l (complete) Mg 2+ (aq) + Cl - (aq) + Pb 2+ (aq) + SO 4 (aq) → Mg 2+ (aq) + SO 4 (aq) + PbCl 2 (s)

64. 64 Write the complete ionic equation for l FeBr 3 (aq) + KOH(aq) → KBr (aq) + Fe(OH) 3 (s) Fe 3+ (aq)Br - (aq)K + (aq)OH - (aq) +++ → + Br - (aq)Fe(OH) 3 (s) + K + (aq)

65. 65 Use the provided molecular equation, write the complete ionic equation for l CaCl 2 (aq) + MgSO 4 (aq) → CaSO 4 (s) + MgCl 2 (aq)

66. 66 Given the molecular equation, write the complete ionic equation for l Ba(OH) 2 (aq) + H 2 SO 4 (aq) → BaSO 4 (s) + HOH(l)

67. 67 If given the complete ionic equation, write the net ionic equation l (complete) Fe 3+ (aq)+ Br - (aq) + K + (aq) + OH - (aq) →K + (aq) +Br - (aq) + Fe(OH) 3 (s) l K + and Br - don’t change. l They are spectator ions l Could be eliminated l (net) Fe 3+ (aq) +OH - (aq) →Fe(OH) 3 (s) l This is what really changes

68. 68 Net ionic equation l Shows only those particles that change (aq  s, l, or g) before and after. l Eliminate spectator ions l Needs to be balanced in terms of both mass and charge l Fe 3+ (aq) +OH - (aq) →Fe(OH) 3 (s) l Fe 3+ (aq) +3 OH - (aq) →Fe(OH) 3 (s)

69. 69 From the molecular equation, write the complete and net ionic equation l HCl (aq) + Ba(OH) 2 (aq) → BaCl 2 (s) + HOH (l)

70. 70 From the molecular equation, write the complete and net ionic equation l Al + FeSO 4 (aq) → Al 2 (SO 4 ) 3 (aq) + Fe

71. 71 Write the net ionic equation l Cl 2 (s) + NaI(aq) → NaCl(aq) + I 2 (s)

72. 72 Write the net ionic equation l K 2 CO 3 (aq) + MgI 2 (aq) → MgCO 3 (s) + KI(aq)

73. 73 Net ionic equations l Written for single and double replacement.

74. 74 Predicting precipitates l Solids formed from aqueous solution. l You can predict them if you know some general rules for solubility.

75. 75 These things are soluble 1.Salts with alkali metals and ammonium 2.Salts of nitrates and chlorates 3.Salts of sulfates except Ag +, Pb 2+, Hg 2 2+, Ba 2+, and Sr 2+ 4.Salts of chlorides except Ag +, Pb 2+, and Hg 2 2+

76. 76 These things are insoluble 5.Carbonates, phosphates, chromates, sulfides, and hydroxides l Unless they fall under rule # 1

77. 77 Is it soluble? l LiBr l Ba(NO 3 ) 2 l CaSO 4 l PbCl 2 l CaCO 3 l K 2 CO 3 l Cd(ClO 3 ) 2

78. 78 Is there a reaction? l For double replacement- has to make gas, solid or water. l Water from an acid- H + and a hydroxide- OH - makes HOH l Solids- from solubility rules l Exchange ions and see if something is insoluble

79. 79 Is there a reaction? l MgSO 4 + NaOH → l H 2 SO 4 + KOH → l K 3 PO 4 + FeF 3 →

80. 80 Oxidation-Reduction Reacions, or “RedOx” reactions l You are only required to RECOGNIZE an oxidation or reduction (to be able to identify which species is reduced or oxidized) l Oxidation number = charge l A species is oxidized when its charge INCREASES (more +) l A species is reduced when its charge DECREASES (more -)

81. 81 Oxidation-Reduction Reacions, or “RedOx” reactions l Oxidation number of a single atom (as in an atom not part of a compound) AND of diatomic molecules is 0 l O is ALWAYS -2 in a compound l Ignore any coefficients l Identify the oxidation numbers of the following: –Ag (s)……= –O 2 (g)…….= –AlBr…….Al = Br = –CO 3 2-…..C = O =

82. 82 Oxidation-Reduction Reacions, or “RedOx” reactions l Determine which species is oxidized and which is reduced in the following equation: l SO 2  S + O 2 –First, identify the oxidation number (charge) of each element on both sides of the equation –Remember, to be oxidized means charge increases –To be reduced means charge decreases

83. 83 Oxidation-Reduction Reacions, or “RedOx” reactions l Determine which species is oxidized and which is reduced in the following equation: l 2Cu + N 2  2CuN –First, identify the oxidation number (charge) of each element on both sides of the equation –Remember, to be oxidized means charge increases –To be reduced means charge decreases

84. 84 Chapter 7 Summary

85. 85 An equation l Describes a reaction l Must be balanced to follow the Law of Conservation of Mass l Can only be balanced by changing the coefficients. l Has special symbols to indicate state, and if catalyst or energy is required.

86. 86 Reactions l Come in 5 types. l Can tell what type they are by the reactants. l Single Replacement happens based on the activity series l Double Replacement happens if the product is a solid, water, or a gas.

87. 87 The Process 1. Determine the type by looking at the reactants. 2. Put the pieces next to each other based on type 3. Use charges to write the formulas –Elements get 2? 4. Use coefficients to balance the equation.