1. 1

2. 2 Chemists use chemical equations to describe reactions they observe in the laboratory or in nature. Chemical equations provide us with the means to 1.summarize the reaction 2.display the substances that are reacting 3.show the products 4.indicate the amounts of all component substances in a reaction.

3. 3 The Chemical Equation

4. 4 Chemical reactions always involve change. Atoms, molecules or ions rearrange to form new substances. The substances entering the reaction are called reactants. The substances formed in the reaction are called products. During reactions chemical bonds are broken and new bonds are formed.

5. 5 A chemical equation uses the chemical symbols and formulas of the reactants and products and other symbolic terms to represent a chemical reaction. A chemical equation is a shorthand expression for a chemical change or reaction.

6. 6 Balancing Equations WB ___W 8 + ___B 12 ___WB ReactantsReactants ProductsProducts Making Hot dogs: How many packages wieners & buns to buy so none is left over. 3224 http://www.youtube.com/watch?v=oYIHLUxzRr8

7. 7 Al + Fe 2 O 3  Fe + Al 2 O 3 reactantsproducts Al + Fe 2 O 3  Fe + Al 2 O 3 Chemical Equation iron oxygen bonds break aluminum oxygen bonds form

8. 8 Coefficients (whole numbers) are placed in front of substances to balance the equation and to indicate the number of units (atoms, molecules, moles, or ions) of each substance that is reacting.

9. 9 Al + Fe 2 O 3  Fe + Al 2 O 3 coefficient 2 2

10. 10 Conditions required to carry out the reaction may be placed above or below the arrow.

11. 11 Al + Fe 2 O 3  Fe + Al 2 O 3 coefficient 2 2   heat

12. 12 The physical state of a substance is indicated by symbols such as (l) for liquid.

13. 13 2Al(s) + Fe 2 O 3 (s)  2Fe(l) + Al 2 O 3 (s) All atoms present in the reactant must also be present in the products. In a chemical reaction atoms are neither created nor destroyed. (s)(s) (l)(l)(s)(s)(s)(s)

14. 14 Symbols Used in Chemical Reactions

15. 15 placed between substances + symbol plus meaning location

16. 16 placed between substances symbol resonance meaning location

17. 17  symbol yields meaning between reactants and products location

18. 18 symbol equilibrium meaning between reactants and products location

19. 19 (s)(s) symbol solid meaning after formula location

20. 20 (l)(l) symbol liquid meaning location after formula

21. 21 (g)(g) symbol gas meaning location after formula

22. 22 (aq) symbol aqueous meaning after formula location

23. 23  symbol heat meaning written above  location

24. 24 h symbol light energy meaning written above  location

25. 25  symbol gas formation meaning after formula location

26. 26  symbol precipitate formation meaning after formula location

27. 27 Writing and Balancing Equations

28. 28 To balance an equation, adjust the number of atoms of each element so that they are the same on each side of the equation. Golden Rule of Balancing: Balance with coefficients only! Do not mess with the subscripts!!!

29. 29 Steps for Balancing Equations

30. 30 Step 1 Identify the reaction. Write a description or word equation for the reaction. Mercury(II) oxide decomposes to form mercury and oxygen. mercury(II) oxide → mercury + oxygen

31. 31 HgO  Hg + O 2 –The formulas of the reactants and products must be correct. –The reactants are written to the left of the arrow and the products to the right of the arrow. Step 2 Write the unbalanced (skeleton) equation. The formulas of the reactants and products can never be changed.

32. 32 Step 3a Balance the equation. –Count and compare the number of atoms of each element on both sides of the equation. –Determine the elements that require balancing.

33. 33 2HgO  2Hg + O 2  THE EQUATION IS BALANCED 

34. 34 sulfuric acid + sodium hydroxide → sodium sulfate + water Balance the Equation

35. 35 H 2 SO 4 (aq) + NaOH(aq) → Na 2 SO 4 (aq) + H 2 O(l) 2 Balance the Equation

36. 36 H 2 SO 4 (aq) + NaOH(aq) → Na 2 SO 4 (aq) + H 2 O(l) 22  THE EQUATION IS BALANCED 

37. 37 butane + oxygen → carbon dioxide + water Balance the Equation

38. 38 C 4 H 10 (g) + O 2 (g) → CO 2 (g) + H 2 O(l) 4 Balance the Equation

39. 39 C 4 H 10 (g) + O 2 (g) → CO 2 (g) + H 2 O(l) 45

40. 40 C 4 H 10 (g) + O 2 (g) → CO 2 (g) + H 2 O(l) 4 5 13 2 Remove fractions by multiplying by the lowest common multiple

41. 41 C 4 H 10 (g) + O 2 (g) → CO 2 (g) + H 2 O(l) 105 28  THE EQUATION IS BALANCED  13

42. 42 What Information Does an Equation Tell Us?

43. 43 The meaning of a formula is context dependent. The formula H 2 O can mean: 1.2 H and 1 O atom 2.1 molecule of water 3.1 mol of water 4.6.02 x 10 23 molecules of water 5.18.02 g of water

44. 44 In an equation formulas can represent units of individual chemical entities or moles. H2H2 +Cl 2 2HCl→ 1 molecule H 2 1 molecule Cl 2 2 molecules HCl 1 mol H 2 1 mol Cl 2 2 mol HCl

45. 45 Formulas Number of molecules Number of atoms Number of moles Mole weights

46. 46 Types of Chemical Equations

47. 47 Combination (synthesis) Decomposition (analysis) Single-Displacement (substitution) Double-Displacement (metathesis) Combustion (oxidation)

48. 48 Combination Reactions

49. 49 A + B  AB Two reactants combine to form one product.

50. 50Examples

51. 51 2Ca(s) + O 2 (g)  2CaO(s) Metal + Oxygen → Metal Oxide 4Al(s) + 3O 2 (g)  2Al 2 O 3 (s)

52. 52 S(s) + O 2 (g)  SO 2 (g) Nonmetal + Oxygen → Nonmetal Oxide N 2 (g) + O 2 (g)  2NO(g)

53. 53 2K(s) + F 2 (g)  2KF(s) Metal + Nonmetal → Salt 2Al(s) + 3Cl 2 (g)  2AlCl 3 (s)

54. 54 Na 2 O(s) + H 2 O(l)  2NaOH(aq) Metal Oxide + Water → Metal Hydroxide CaO(s) + 2H 2 O(l)  2Ca(OH) 2 (aq)

55. 55 SO 3 (g) + H 2 O(l)  H 2 SO 4 (aq) Nonmetal Oxide + H 2 O(l) → Oxy-acid N 2 O 5 (g) + H 2 O(l)  2HNO 3 (aq)

56. 56 Decomposition Reactions

57. 57 AB  A + B A single substance breaks down to give two or more different substances.

58. 58ExamplesExamples

59. 59 2Ag 2 O(s)  4Ag(s) + O 2 (g) Metal Oxide → Metal + Oxygen Metal Oxide → Metal Oxide + Oxygen 2PbO 2 (s)  2PbO(s) + O 2 (g)

60. 60 Carbonate → CO 2 (g) CaCO 3 (s)  CaO(s) + CO 2 (g) 2NaHCO 3 (s)  Na 2 CO 3 (s) + H 2 O(g) + CO 2 (g) Hydrogen carbonate → CO 2 (g)

61. 61 Miscellaneous Reactions 2KClO 3 (s)  2KCl(s) + O 2 (g) 2NaNO 3 (s)  2NaNO 2 (s) + O 2 (g) 2H 2 O 2 (l)  2H 2 O(l) + O 2 (g)

62. 62 Single Displacement Reactions

63. 63 A + BC  AC + B One element reacts with a compound to replace one of the elements of that compound.

64. 64 Mg(s) + HCl(aq)  H 2 (g) + MgCl 2 (aq) 2Al(s) + 3H 2 SO 4 (aq)  3H 2 (g) + Al 2 (SO 4 ) 3 (aq) salt Metal + Acid → Hydrogen + Salt salt

65. 65 Na(s) + 2H 2 O(l)  H 2 (g) + NaOH(aq) Ca(s) + 2H 2 O(l)  H 2 (g) + Ca(OH) 2 (aq) Metal + Water → Hydrogen + Metal Hydroxide metal hydroxide

66. 66 Metal + Water → Hydrogen + Metal Oxide metal oxide Fe(s) + 4H 2 O(g)  4H 2 (g) + Fe 3 O 4 (s)

67. 67 The Activity Series

68. 68 Metals K Ca Na Mg Al Zn Fe Ni Sn Pb H Cu Ag Hg An atom of an element in the activity series will displace an atom of an element below it from one of its compounds. Sodium (Na) will displace an atom below it from one of its compounds. increasing activity

69. 69 Examples Metal Activity Series

70. 70 Mg(s) + PbS(s)  MgS(s) + Pb(s) Metal Higher in Activity Series Displacing Metal Below It Magnesium is above lead in the activity series. Metals Mg Al Zn Fe Ni Sn Pb

71. 71 Ag(s) + CuCl 2 (s)  no reaction Metal Lower in Activity Cannot Displace Metal Above It Metals Pb H Cu Ag Hg Silver is below copper in the activity series.

72. 72 Example Halogen Activity Series

73. 73 Cl 2 (g) + CaBr 2 (s)  CaCl 2 (aq) + Br 2 (aq) Halogen Higher in Activity Series Displaces Halogen Below It Halogens F 2 Cl 2 Br 2 I 2 Chlorine is above bromine in the activity series.

74. 74 Double Displacement Reactions

75. 75 AB + CD  AD + CB Two compounds exchange partners with each other to produce two different compounds. The reaction can be thought of as an exchange of positive and negative groups. A displaces C and combines with D B displaces D and combines with C

76. 76 The Following Accompany Double Displacement Reactions formation of a precipitate release of gas bubbles release of heat formation of water

77. 77Examples

78. 78 Acid Base Neutralization HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l) H 2 SO 4 (aq) + 2NaOH(aq)  Na 2 SO 4 (aq) + 2H 2 O(l) acid + base → salt + water

79. 79 Formation of an Insoluble Precipitate AgNO 3 (aq) + NaCl(aq)  AgCl(s) + NaNO 3 (aq) Pb(NO 3 ) 2 (aq) + 2KI(aq)  PbI 2 (s) + 2KNO 3 (aq) ↓ ↓

80. 80 Metal Oxide + Acid CuO(s) + 2HNO 3 (aq)  Cu(NO 3 ) 2 (aq) + H 2 O(l) CaO(s) + 2HCl(aq)  CaCl 2 (s) + H 2 O(l) metal oxide + acid → salt + water

81. 81 Formation of a Gas H 2 SO 4 (aq) + 2NaCN(aq)  Na 2 SO 4 (aq) + 2HCN(g) NH 4 Cl(aq) + NaOH(aq)  NaCl(aq) + NH 4 OH(aq) NH 4 OH(aq)  NH 3 (g) + H 2 O(l) indirect gas formation ↑ ↑

82. 82 Combustion Reactions

83. 83 Hydrocarbon reacts with oxygen to produce carbon dioxide and water

84. 84Examples

85. 85 Combustion of Methane CH 4 (g) + O 2 (g)  CO 2 (g) + H 2 O (l) Balance by first balancing C, then H, and finally O

86. 86 1 CH 4 (g) + O 2 (g)  1 CO 2 (g) + H 2 O (l) 1 CH 4 (g) + O 2 (g)  1 CO 2 (g) + 2 H 2 O (l) 1 CH 4 (g) + 2 O 2 (g)  1 CO 2 (g) + 2 H 2 O (l)

87. 87 Combustion of Ethane C 2 H 6 (g) + O 2 (g)  CO 2 (g) + H 2 O (l) Balance by first balancing C, then H, and finally O. Use fractions if needed to balance, then change to whole numbers.

88. 88 1 C 2 H 6 (g) + O 2 (g)  2 CO 2 (g) + H 2 O (l) 1 C 2 H 6 (g) + O 2 (g)  2 CO 2 (g) + 3 H 2 O (l) 2 C 2 H 6 (g) +7 O 2 (g)  4 CO 2 (g) + 6 H 2 O (l)

89. 89 Heat in Chemical Reactions

90. 90 Energy changes always accompany chemical reactions. One reason why reactions occur is that the product attains a lower energy state than the reactants. When this occurs, energy is released to the surroundings.

91. 91 Energy changes always accompany chemical reactions. One reason why reactions occur is that the product attains a lower energy state than the reactants. When this occurs, energy is released to the surroundings.

92. 92 H 2 (g) + Cl 2 (g) → 2HCl(g) + 185 kJ (exothermic) N 2 (g) + O 2 (g) + 185 kJ → 2NO(g) (exothermic) Exothermic reactions liberate heat. Endothermic reactions absorb heat. The amounts of substances are expressed in moles. 1 mol 2 mol 1 mol 2 mol

93. 93 For life on Earth the sun is the major provider of energy. The energy for plant photosynthesis is derived from the sun. glucose 6CO 2 + 6H 2 O + 2519 kJ → C 6 H 12 O 6 + 6O 2

94. 94 Energy of Activation

95. 95 A certain amount of energy is always required for a reaction to occur. The energy required to start a reaction is called the energy of activation.

96. 96 This reaction will not occur unless activation energy is supplied. The activation energy can take the form of a spark or a flame. 6CH 4 + 2O 2 → CO 2 + 2 H 2 O + 890 kJ

97. 97 8.1 8.2

98. 98