2. 1 Chapter 8 “Chemical Reactions” Chemistry Judson High School Mr. Trotts

3. 2 Describing Chemical Reactions l OBJECTIVES: –Describe how to write a word equation.

4. 3 Describing Chemical Reactions l OBJECTIVES: –Describe how to write a skeleton equation.

5. 4 Describing Chemical Reactions l OBJECTIVES: –Describe the steps for writing a balanced chemical equation.

6. 5 All chemical reactions… l have two parts: –Reactants - the substances you start with –Products- the substances you end up with l The reactants turn into the products. Reactants  Products

7. 6 Reactants Products

8. 7 CHEMICAL REACTION: a reaction in which one or more substances are changed to new substances.

9. 8 REACTANTS: substances involved in a chemical reaction.

10. 9 PRODUCTS: new substances produced in a chemical reaction.

11. 10 The relationship between reactants and products can be written as: REACTANTS  PRODUCTS

12. 11 In a chemical reaction l Atoms aren’t created or destroyed. l A reaction can be described several ways: 1. In a sentence ( every item is a word) Copper reacts with chlorine to form copper (II) chloride. 2. In a word equation (some symbols used) Copper + chlorine  copper (II) chloride

13. 12 Symbols in equations l the arrow separates the reactants from the products –Read as “reacts to form” or yields l The plus sign = “and” l (s) after the formula = solid: AgCl (s) l (g) after the formula = gas: CO 2(g) l (l) after the formula = liquid: H 2 O (l)

14. 13 Symbols used in equations l (aq) after the formula = dissolved in water, an aqueous solution: NaCl (aq) is a salt water solution  used after a product indicates a gas has been produced: H 2 ↑  used after a product indicates a solid has been produced: PbI 2 ↓

15. 14 Symbols used in equations ■ indicates a reversible reaction (more later) ■ shows that heat is supplied to the reaction ■ is used to indicate a catalyst is supplied, in this case, platinum.

16. 15 What is a catalyst? l A substance that speeds up a reaction, without being changed or used up by the reaction. l Enzymes are biological or protein catalysts.

17. 16 3. The Skeleton Equation l Uses formulas and symbols to describe a reaction –but doesn’t indicate how many; this means they are NOT balanced l All chemical equations are a description that describe reactions.

18. 17 Write a skeleton equation for: 1. Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. 2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

19. 18 Now, read these: Fe(s) + O 2 (g)  Fe 2 O 3 (s) Cu(s) + AgNO 3 (aq)  Ag(s) + Cu(NO 3 ) 2 (aq) NO 2 (g) N 2 (g) + O 2 (g)

20. 19 4. Balanced Chemical Equations l Atoms can’t be created or destroyed in an ordinary reaction: –All the atoms we start with we must end up with l A balanced equation has the same number of each element on both sides of the equation.

21. 20 I. CONSERVATION OF MASS: In a chemical reaction, matter is not created or destroyed but is conserved.

22. 21 II. IN OTHER WORDS: The starting mass of the reactants equals the final mass of the products.

23. 22 Rules for balancing: 1)Assemble the correct formulas for all the reactants and products, use + and → 2)Count the number of atoms of each type appearing on both sides 3)Balance the elements one at a time by adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST! (I prefer to save O until the very last) 4)Check to make sure it is balanced.

24. 23 l Never 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 2NaCl is okay, but Na2Cl is not.

25. 24 WRITING EQUATIONS

26. 25 III. Chemical reactions can be described with words such as:

27. 26 solid lead (II) nitrate, dissolved in water, plus solid potassium iodide, dissolved in water, produces solid lead (II) iodide plus potassium nitrate, dissolved in water

28. 27 A. All of this information is important to letting a scientist know what the reactants and products are as well as their physical states.

29. 28 B. A shorthand method has been developed to describe chemical reactions.

30. 29 C. This method uses: 1. Chemical formulas (NaCl for sodium chloride)

31. 30 2. Coefficients (numbers to indicate how many molecules) 3. Symbols (for physical state, catalysts, direction of reaction, etc.)

32. 31 The reaction described above would look like this: Pb(NO 3 ) 2 (aq) + 2KI(aq)  PbI 2 (s) + 2KNO 3 (aq)

33. 32 IV. COEFFICIENTS: are the numbers placed to the left of the formulas for the reactants and products.

34. 33 A. The coefficients represent the number of units of each substance taking part in a reaction.

35. 34 B. In the reaction above, there are: 1. 2 units each of KNO 3 (aq) and KI(aq) 2. 1 unit each of Pb(NO 3 ) 2 (aq) and PbI 2 (s) Pb(NO 3 ) 2 (aq) + 2KI(aq)  PbI 2 (s) + 2KNO 3 (aq)

36. 35 V. Symbols are used to indicate what is happening in the reaction.

37. 36 A. The physical state of the reactants B. Things added to help the reaction take place

38. 37 C. An indicator to show which chemicals are the reactants and which are the products

39. 38 D. Some commonly used symbols are in the following table.

40. 39 SYMBOLS USED IN CHEMICAL EQUATIONS

41. 40 Reactants Products C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(g) 1 unit propane gas; 5 units oxygen gas 3 units carbon dioxide gas; 4 units water vapor

42. 41 VII. REACTION ENERGY -- Chemical reactions either absorb heat (endothermic) or release heat (exothermic).

43. 42 A. Endothermic reactions must have energy added (usually thermal, sometimes electrical, or light) for the reaction to take place.

44. 43 1. This is due to more energy being required to break bonds than to make new bonds.

45. 44 2. Some endothermic reactions cause the reaction container to feel cool to the touch.

46. 45 b. Exothermic reactions have some form of energy released by the reaction (usually it is thermal or light).

47. 46 1. This is due to less energy being required to break bonds than to form new bonds.

48. 47 2. Exothermic reactions cause the reaction container to feel warm or even hot to the touch.

49. 48 BALANCING EQUATIONS THERE ARE 5 STEPS TO BALANCING EQUATIONS.

50. 49 Step 1: Remember these two rules:

51. 50  Each capital letter begins a new element

52. 51  Coefficients (numbers that multiply the whole formula) can only be placed in front of the chemical formula

53. 52 Step 2: Determine each type of element present on both sides of the equation

54. 53  Write the equation.

55. 54  Make a table under the equation that lists each type of element.

56. 55  List the elements in the center of the table

57. 56 Example:

58. 57 Step 3: Determine the number of each type of element on both sides of the equation.

59. 58 Example:

60. 59 Step 4: Determine which elements are not balanced. In the example, the unbalanced element is oxygen.

61. 60 Step 5: to balance the number of atoms for the element, Put a coefficient in front of the formula with the lower number of atoms for that element to produce an equal number of atoms.

62. 61  Remember: The coefficient multiplies all of the elements in the chemical formula.

63. 62  Start balancing by trying to make the numbers of elements in all the other formulas equal with the most complicated chemical formula (the formula with the most elements and subscripts).

64. 63  Always balance Hydrogen and Oxygen last.

65. 64  Water can always be added to the equation

66. 65 Example: Place a coefficient of 2 in front of the MgO. Then multiply the number of atoms in the table by the coefficient.

67. 66

68. 67 Go to the other side of the equation and place a coefficient of 2 in front of the MgBr 2.

69. 68

70. 69 Place a coefficient of 2 in front of the Br 2.

71. 70

72. 71 The equation is now balanced with 2-Mg, 4-Br, and 2-O.

73. 72 OTHER HINTS

74. 73  The diatomic molecules (Br I N Cl H O F) are always written as two-atom molecules (Br 2, I 2, N 2, Cl 2, H 2, O 2, F 2 ).

75. 74  A coefficient can be changed when needed to get every element to balance.

76. 75  Many times there can be an even amount of atoms on one side and an odd amount of atoms on the other side. ] Find a common multiple.

77. 76 Example:

78. 77 1. Determine the common multiple between 2 and 3 (it is 6 ).

79. 78 2. Multiply the H 2 by 3 and the NH 3 by 2.

80. 79 The equation is now balanced with 2-N and 6-H.

81. 80  Sometimes hydrogen and oxygen can cause problems. Two or more sources of hydrogen or oxygen on one side of the equation can occur.

82. 81 Example: ZnS + O 2  ZnO + SO 2 The product side has oxygen from 2 sources.

83. 82 1. Make the odd source of oxygen even by multiplying ZnO by 2.

84. 83 2. Place a 2 in front of the ZnS.

85. 84 3. Place a 2 in front of the SO 2.

86. 85 4. Multiply the O 2 by 3 to balance the equation.

87. 86 The equation is now balanced with 2-Zn, 2-S, and 6-O on each side.

88. 87  A polyatomic ion that is not broken into parts by the reaction (it stays the same on both sides) is treated like an element.

89. 88 Example: The equation, H 3 PO 4 + Ca(OH) 2  Ca 3 (PO 4 ) 2 + H 2 O, has two polyatomic ions, phosphate (PO 4 ) and hydroxide (OH).

90. 89 The phosphate has remained intact. The hydroxide was broken up to form water.

91. 90 1. Make a table and list all of the elements.

92. 91 Note that the oxygen in the phosphate ion is not included when the number of oxygen atoms is listed.

93. 92 2. The most complicated formula is the Ca 3 (PO 4 ) 2. Make the other formulas balance with the Ca 3 (PO 4 ) 2.

94. 93 Start by balancing the calcium, multiply Ca(OH) 2 by 3.

95. 94 3. Multiply the H 3 PO 4 by 2. This leaves 12 H’s and 6 O’s. The ratio of H to O is 2:1. This means that the reaction made 6 molecules of water.

96. 95 4. Finish balancing the equation by multiplying the water by 6. The equation is now balanced with 2-PO 4, 3-Ca, 12-H and 6-O.

97. 96 Practice Balancing Examples 1. _AgNO 3 + _Cu  _Cu(NO 3 ) 2 + _Ag 2. _Mg + _N 2  _Mg 3 N 2 3. _P + _O 2  _P 4 O 10 4. _Na + _H 2 O  _H 2 + _NaOH 5. _CH 4 + _O 2  _CO 2 + _H 2 O

98. 97 Types of Chemical Reactions l OBJECTIVES: –Describe the five general types of reactions.

99. 98 Types of Chemical Reactions l OBJECTIVES: –Predict the products of the five general types of reactions.

100. 99 Types of Reactions l There are millions of reactions. l We can’t remember them all, but luckily they will fall into several categories. l We will learn 5 major types. l Will be able to predict the products. l For some, we will be able to predict whether or not they will happen at all. l We recognize them by their reactants

101. 100 #1 - Combination Reactions or Synthesis Reactions l Combine = put together l 2 substances combine to make one compound. Ca +O 2  CaO SO 3 + H 2 O  H 2 SO 4 l We can predict the products if the reactants are two elements. Mg + N 2 

102. 101 Complete and balance: Ca + Cl 2  Fe + O 2  (assume iron (II) oxide is product) Al + O 2  l Remember that the first step is to write the correct formulas – you can still change the subscripts at this point, but not later! l Then balance by using the coefficients only

103. 102 #1 - Combination l Note: a) Some nonmetal oxides react with water to produce an acid: SO 2 + H 2 O  H 2 SO 3 b) Some metallic oxides react with water to produce a base: CaO + H 2 O  Ca(OH) 2

104. 103 #2 - Decomposition Reactions l decompose = fall apart l one reactant breaks apart into two or more elements or compounds. l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2 l Note that energy (heat, sunlight, electricity, etc.) is usually required

105. 104 #2 - Decomposition Reactions l Can predict the products if it is a binary compound-Made up of only two elements –breaks apart into its elements: lH2OlH2O l HgO

106. 105 #2 - Decomposition Reactions l If the compound has more than two elements you must be given one of the products –The other product will be from the missing pieces l NiCO 3 CO 2 + ___ H 2 CO 3 (aq)  CO 2 + ___ heat

107. 106 #3 - Single Replacement l One element replaces another 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

108. 107 #3 Single Replacement l Metals replace other metals (and they can also replace hydrogen) K + AlN  Zn + HCl  l Think of water as: HOH –Metals replace one of the H, and then combine with the hydroxide. Na + HOH 

109. 108 #3 Single Replacement l We can even tell whether or not a single replacement reaction will happen: –Some chemicals are more “active” than others –More active replaces less active l There is a list on page 333 - called the Activity Series of Metals l Higher on the list replaces lower.

110. 109 The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1)Metals can replace other metals provided that they are above the metal that they are trying to replace. 2)Metals above hydrogen can replace hydrogen in acids. 3)Metals from sodium upward can replace hydrogen in water. Higher activity Lower activity

111. 110 #3 Single Replacement Practice: 6. Fe + CuSO 4  7. Pb + KCl  8. Al + HCl 

112. 111 #4 - Double Replacement l Two things replace each other. –Reactants must be two ionic compounds. –Usually in aqueous solution NaOH + FeCl 3  –The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 NaOH + FeCl 3  Fe(OH) 3 + NaCl

113. 112 The Activity Series of the Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl (s) + F 2(g)  2NaF (s) + Cl 2(g) MgCl 2(s) + Br 2(g)  ??? No Reaction ??? Higher Activity Lower Activity

114. 113 #4 - Double Replacement l Has certain “driving forces” –Will only happen if one of the products: a) doesn’t dissolve in water and forms a solid (a “precipitate”), or b) is a gas that bubbles out, or c) is a molecular compound (usually water).

115. 114 Complete and balance: assume all of the following reactions actually take place: 9. CaCl 2 + NaOH  10. CuCl 2 + K 2 S  11. KOH + Fe(NO 3 ) 3  12. (NH 4 ) 2 SO 4 + BaF 2 

116. 115 How to recognize which type l Look at the reactants: E + E =Combination (synthesis) C =Decomposition E + C =Single replacement C + C =Double replacement

117. 116 Practice Examples: reaction type and product 13. H 2 + O 2  14. H 2 O  15. Zn + H 2 SO 4  16. HgO  17. KBr +Cl 2  18. AgNO 3 + NaCl  19. Mg(OH) 2 + H 2 SO 3 

118. 117 #5 - Combustion l Means “add oxygen” l Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning” 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 (or possibly just C) and H 2 O.

119. 118 Combustion Examples: C 4 H 10 + O 2  (assume complete) C 4 H 10 + O 2  (incomplete) C 6 H 12 O 6 + O 2  (complete) C 8 H 8 +O 2  (incomplete)

120. 119 SUMMARY: an equation... l Describes a reaction l Must be balanced in order to follow the Law of Conservation of Mass l Can only be balanced by changing the coefficients. l Has special symbols to indicate physical state, if a catalyst or energy is required, etc.

121. 120 Reactions l Come in 5 major types. l We can tell what type they are by looking at the reactants. l Single Replacement happens based on the Activity Series l Double Replacement happens if the product is a precipitate (insoluble solid), water, or a gas.

122. 121 Reactions in Aqueous Solution l OBJECTIVES: –Describe the information found in a net ionic equation.

123. 122 Reactions in Aqueous Solution l OBJECTIVES: –Predict the formation of a precipitate in a double replacement reaction.

124. 123 Net Ionic Equations l Many reactions occur in water- that is, in aqueous solution l Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water l Now we are ready to write an ionic equation

125. 124 Net Ionic Equations l Example: –AgNO 3 + NaCl  AgCl + NaNO 3 1. this is the full equation 2. now write it as an ionic equation 3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation

126. 125 Net ionic equations Na + Al 3+ S 2– 2Ca 2+ PO 4 3– 3Cl –

127. 126 Review: forming ions l Ionic (i.e. salt) refers to +ve ion plus -ve ion l Usually this is a metal + non-metal or metal + polyatomic ion (e.g. NaCl, NaClO 3, Li 2 CO 3 ) l Polyatomic ions are listed on page 95 l (aq) means aqueous (dissolved in water) l For salts (aq) means the salt exists as ions l NaCl(aq) is the same as: Na + (aq) + Cl – (aq) l Acids form ions: HCl(aq) is H + (aq) + Cl – (aq), Bases form ions: NaOH(aq) is Na + + OH – Q - how is charge determined (+1, -1, +2, etc.)? A - via valences (periodic table or see pg. 95) l F, Cl gain one electron, thus forming F –, Cl – l Ca loses two electrons, thus forming Ca 2+

128. 127 l Charge can also be found via the compound l E.g. in NaNO 3 (aq) if you know Na forms Na +, then NO 3 must be NO 3 – (NaNO 3 is neutral) l By knowing the valence of one element you can often determine the other valences Q - Write the ions that form from Al 2 (SO 4 ) 3 (aq)? Step 1 - look at the formula: Al 2 (SO 4 ) 3 (aq) Step 2 - determine valences: Al 3 (SO 4 ) 2 Background: valences and formulas (Al is 3+ according to the periodic table) Step 3 - write ions: 2Al 3+ (aq) + 3SO 4 2– (aq) l Note that there are 2 aluminums because Al has a subscript of 2 in the original formula

129. 128 Practice with writing ions Q - Write ions for Na 2 CO 3 (aq) A - 2Na + (aq) + CO 3 2– (aq) (from the PT Na is 1+. There are 2, thus we have 2Na +. There is only one CO 3. It must have a 2- charge) l Notice that when ions form from molecules, charge can be separated, but the total charge (and number of each atom) stays constant. Q - Write ions for Ca 3 (PO 4 ) 2 (aq) & Cd(NO 3 ) 2 (aq) A - 3Ca 2+ (aq) + 2PO 4 3– (aq) A - Cd 2+ (aq) + 2NO 3 – (aq) Q - Write ions for Na 2 S(aq) and Mg 3 (BO 3 ) 2 (aq) A - 2Na + (aq) + S 2– (aq), 3Mg 2+ (aq)+ 2BO 3 3– (aq)

130. 129 Types of chemical equations Equations can be divided into 3 types 1) Molecular, 2) Ionic, 3) Net ionic l Here is a typical molecular equation: Cd(NO 3 ) 2 (aq) + Na 2 S(aq)  CdS(s) + 2NaNO 3 (aq) l We can write this as an ionic equation (all compounds that are (aq) are written as ions): Cd 2+ (aq) + 2NO 3 – (aq) + 2Na + (aq) + S 2– (aq)  CdS(s) + 2Na + (aq) + 2NO 3 – (aq) l To get the NET ionic equation we cancel out all terms that appear on both sides: Net: Cd 2+ (aq) + S 2– (aq)  CdS(s)

131. 130 Equations must be balanced l There are two conditions for molecular, ionic, and net ionic equations Materials balance Both sides of an equation should have the same number of each type of atom Electrical balance Both sides of a reaction should have the same net charge Q- When NaOH(aq) and MgCl 2 (aq) are mixed, _______(s) and NaCl(aq) are produced. Write balanced molecular, ionic & net ionic equations Mg(OH) 2

132. 131 NaOH(aq) + MgCl 2 (aq)  Mg(OH) 2 (s) + NaCl(aq) Next, balance the equation First write the skeleton equation 2 2 Ionic equation: 2Na + (aq) + 2OH - (aq) + Mg 2+ (aq) + 2Cl - (aq)  Mg(OH) 2 (s) + 2Na + (aq) + 2Cl - (aq) LiNO 3 Ca 3 (PO 4 ) 2 NaC 2 H 3 O 2

133. 132 Net ionic equation: 2OH - (aq) + Mg 2+ (aq)  Mg(OH) 2 (s) Write balanced ionic and net ionic equations: CuSO 4 (aq) + BaCl 2 (aq)  CuCl 2 (aq) + BaSO 4 (s) Fe(NO 3 ) 3 (aq) + LiOH(aq)  ______(aq) + Fe(OH) 3 (s) Na 3 PO 4 (aq) + CaCl 2 (aq)  _________(s) + NaCl(aq) Na 2 S(aq) + AgC 2 H 3 O 2 (aq)  ________(aq) + Ag 2 S(s)

134. 133 Cu 2+ (aq) + SO 4 2– (aq) + Ba 2+ (aq) + 2Cl – (aq)  Cu 2+ (aq) + 2Cl – (aq) + BaSO 4 (s) Net: SO 4 2– (aq) + Ba 2+ (aq)  BaSO 4 (s) Fe 3+ (aq) + 3NO 3 – (aq) + 3Li + (aq) + 3OH – (aq)  3Li + (aq) + 3NO 3 – (aq) + Fe(OH) 3 (s) Net: Fe 3+ (aq) + 3OH – (aq)  Fe(OH) 3 (s) For more lessons, visit www.chalkbored.com www.chalkbored.com

135. 134 2Na 3 PO 4 (aq) + 3CaCl 2 (aq)  Ca 3 (PO 4 ) 2 (s)+ 6NaCl(aq) 6Na + (aq) + 2PO 4 3– (aq) + 3Ca 2+ (aq) + 6Cl – (aq)  Ca 3 (PO 4 ) 2 (s)+ 6Na + (aq) + 6Cl – (aq) Net: 2PO 4 3– (aq) + 3Ca 2+ (aq)  Ca 3 (PO 4 ) 2 (s) 2Na + (aq) + S 2– (aq) + 2Ag + (aq) + 2C 2 H 3 O 2 – (aq)  2Na + (aq) + 2C 2 H 3 O 2 – (aq) + Ag 2 S(s) Net: S 2– (aq) + 2Ag + (aq)  Ag 2 S(s)

136. 135 Predicting the Precipitate l Insoluble salt = a precipitate l General solubility rules are found: a)Table 8.3, p. 227 b)Reference section - page 887 Table A.7 (back of textbook) gives combinations of cations and anions that are soluble or insoluble in water