1. Chapter 4: Chemical Reactions in Aqueous Solutions Precipitation reactions Acid-base Reactions Reduction-Oxidation Reactions

2. Copyright © Cengage Learning. All rights reserved. 4 | 2 Contents and Concepts Ions in Aqueous Solution Explore how molecular and ionic substances behave when they dissolve in water to form solutions. Ionic Theory of Solutions and Solubility Rules Molecular and Ionic Equations

3. Copyright © Cengage Learning. All rights reserved. 4 | 3 Types of Chemical Reactions Investigate several important types of reactions that typically occur in aqueous solution: precipitation reactions, acid–base reactions, and oxidation– reduction reactions. 3. Precipitation Reactions 4. Acid–Base Reactions 5. Oxidation–Reduction Reactions 6. Balancing Simple Oxidation–Reduction Equations

4. Copyright © Cengage Learning. All rights reserved. 4 | 4 Working with Solutions Now that we have looked at how substances behave in solution, it is time to quantitatively describe these solutions using concentration. 7. Molar Concentration 8. Diluting Solutions Quantitative Analysis Using chemical reactions in aqueous solution, determine the amount of substance or species present in materials. 9. Gravimetric Analysis 10. Volumetric Analysis

5. 5 A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount SolutionSolventSolute Soft drink (l) Air (g) Soft Solder (s) H2OH2O N2N2 Pb Sugar, CO 2 O 2, Ar, CH 4 Sn

6. 6 An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. nonelectrolyte weak electrolyte strong electrolyte

7. Copyright © Cengage Learning. All rights reserved. 4 | 7 A strong electrolyte dissolves to produce ions. The ions, as moving charges, complete the circuit. When a light bulb is attached to the circuit, it shines.

8. Copyright © Cengage Learning. All rights reserved. 4 | 8 A strong electrolyte is an electrolyte that exists in solution almost entirely as ions.

9. 9 Strong Electrolyte – 100% dissociation NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq)

10. 10 Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.   H2OH2O

11. 11 Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) H2OH2O

12. 12 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb 2+ + 2NO 3 - + 2Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 - Na + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) precipitate Pb 2+ + 2I - PbI 2 (s)

13. 13 Precipitation of Lead Iodide PbI 2 Pb 2+ + 2I - PbI 2 (s)

14. 14 Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.

15. 15 Examples of Insoluble Compounds CdS PbSNi(OH) 2 Al(OH) 3

16. Copyright © Cengage Learning. All rights reserved. 4 | 16 Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. K 3 PO 4 + CaCl 2  1.Determine the product formulas: K + and Cl − make KCl Ca 2+ and PO 4 3− make Ca 3 (PO 4 ) 2 2.Determine whether the products are soluble: KCl is soluble Ca 3 (PO 4 ) 2 is insoluble

17. Copyright © Cengage Learning. All rights reserved. 4 | 17 Molecular Equation (Balance the reaction and include state symbols) 2K 3 PO 4 (aq) + 3CaCl 2 (aq)  6KCl(aq) + Ca 3 (PO 4 ) 2 (s) Ionic Equation 6K + (aq) + 2PO 4 3− (aq) + 3Ca 2+ (aq) + 6Cl − (aq)  6K + (aq) + 6Cl − (aq) + Ca 3 (PO 4 ) 2 (s) Net Ionic Equation 2PO 4 3− (aq) + 3Ca 2+ (aq)  Ca 3 (PO 4 ) 2 (s)

18. Copyright © Cengage Learning. All rights reserved. 4 | 18 Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. NaOH + MgCl 2  1.Determine the product formulas: Na + and Cl − make NaCl Mg 2+ and OH − make Mg(OH) 2 2.Determine whether the products are soluble: NaCl is soluble Mg(OH) 2 is insoluble

19. Copyright © Cengage Learning. All rights reserved. 4 | 19 Molecular Equation (Balance the reaction and include state symbols) 2NaOH(aq) + MgCl 2 (aq)  2NaCl(aq) + Mg(OH) 2 (s) Ionic Equation 2Na + (aq) + 2OH − (aq) + Mg 2+ (aq) + 2Cl − (aq)  2Na + (aq) + 2Cl − (aq) + Mg(OH) 2 (s) Net Ionic Equation 2OH − (aq) + Mg 2+ (aq)  Mg(OH) 2 (s)

20. Acid-Base Reaction

21. 21 Properties of Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Cause color changes in plant dyes. 2HCl (aq) + Mg (s) MgCl 2 (aq) + H 2 (g) 2HCl (aq) + CaCO 3 (s) CaCl 2 (aq) + CO 2 (g) + H 2 O (l) Aqueous acid solutions conduct electricity.

22. 22 Have a bitter taste. Feel slippery. Many soaps contain bases. Properties of Bases Cause color changes in plant dyes. Aqueous base solutions conduct electricity. Examples:

23. Copyright © Cengage Learning. All rights reserved. 4 | 23 Household Acids and Bases

24. Copyright © Cengage Learning. All rights reserved. 4 | 24 Arrhenius Acid A substance that produces hydrogen ions, H +, when it dissolves in water. Arrhenius Base A substance that produces hydroxide ions, OH −, when it dissolves in water.

25. 25 Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water

26. Copyright © Cengage Learning. All rights reserved. 4 | 26 Brønsted–Lowry Acid The species (molecule or ion) that donates a proton to another species in a proton−transfer reaction. Brønsted–Lowry Base The species (molecule or ion) that accepts a proton from another species in a proton−transfer reaction.

27. 27 A Brønsted acid is a proton donor A Brønsted base is a proton acceptor acidbaseacidbase A Brønsted acid must contain at least one ionizable proton!

28. Copyright © Cengage Learning. All rights reserved. 4 | 28 Acid−Base Indicator A dye used to distinguish between an acidic and basic solution by means of the color changes it undergoes in these solutions. The sample beakers show a red cabbage indicator in beakers varying in acidity from highly acidic (left) to highly basic (right).

29. Copyright © Cengage Learning. All rights reserved. 4 | 29 Strong Acid An acid that ionizes completely in water. It is present entirely as ions; it is a strong electrolyte. Common strong acids: HNO 3 H 2 SO 4 HClO 4 HClHBrHI

30. Copyright © Cengage Learning. All rights reserved. 4 | 30 Weak Acid An acid that only partly ionizes in water. It is present primarily as molecules and partly as ions; it is a weak electrolyte.

31. 31 Monoprotic acids HCl H + + Cl - HNO 3 H + + NO 3 - CH 3 COOH H + + CH 3 COO - Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids H 2 SO 4 H + + HSO 4 - HSO 4 - H + + SO 4 2- Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids H 3 PO 4 H + + H 2 PO 4 - H 2 PO 4 - H + + HPO 4 2- HPO 4 2- H + + PO 4 3- Weak electrolyte, weak acid

32. 32

33. Copyright © Cengage Learning. All rights reserved. 4 | 33 Strong Base A base that ionizes completely in water. It is present entirely as ions; it is a strong electrolyte. Common strong bases: LiOHNaOH KOH Ca(OH) 2 Sr(OH) 2 Ba(OH) 2

34. Copyright © Cengage Learning. All rights reserved. 4 | 34 Weak Base A base that is only partly ionized in water. It is present primarily as molecules and partly as ions; it is a weak electrolyte. These are often nitrogen bases such as NH 3 : NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH − (aq) If a base is not strong, it is weak.

35. 35 Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH 3 COO -, (c) H 2 PO 4 - HI (aq) H + (aq) + I - (aq)Brønsted acid CH 3 COO - (aq) + H + (aq) CH 3 COOH (aq) Brønsted base H 2 PO 4 - (aq) H + (aq) + HPO 4 2- (aq) H 2 PO 4 - (aq) + H + (aq) H 3 PO 4 (aq) Brønsted acid Brønsted base

36. Copyright © Cengage Learning. All rights reserved. 4 | 36 Neutralization Reaction A reaction of an acid and a base that results in an ionic compound (a salt) and possibly water.

37. 37 Neutralization Reaction acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O

38. 38 Neutralization Reaction Involving a Weak Electrolyte weak acid + base salt + water HCN (aq) + NaOH (aq) NaCN (aq) + H 2 O HCN + Na + + OH - Na + + CN - + H 2 O HCN + OH - CN - + H 2 O

39. 39 Neutralization Reaction Producing a Gas acid + base salt + water + CO 2 2HCl (aq) + Na 2 CO 3 (aq) 2NaCl (aq) + H 2 O +CO 2 2H + + 2Cl - + 2Na + + CO 3 2- 2Na + + 2Cl - + H 2 O + CO 2 2H + + CO 3 2- H 2 O + CO 2

40. Copyright © Cengage Learning. All rights reserved. 4 | 40 Write the molecular, ionic, and net ionic equations for the neutralization of sulfurous acid, H 2 SO 3, by potassium hydroxide, KOH.

41. Copyright © Cengage Learning. All rights reserved. 4 | 41 Molecular Equation (Balance the reaction and include state symbols) H 2 SO 3 (aq) + 2KOH(aq)  2H 2 O(l) + K 2 SO 3 (aq) Ionic Equation H 2 SO 3 (aq) + 2K + (aq) + 2OH − (aq)  2H 2 O(l) + 2K + (aq) + SO 3 2− (aq) Net Ionic Equation H 2 SO 3 (aq) + 2OH − (aq)  2H 2 O(l) + SO 3 2− (aq)

42. Copyright © Cengage Learning. All rights reserved. 4 | 42 Acid−base reactions with gas−formation Sulfides, carbonates, sulfites react with acid to form a gas. Na 2 S(aq) + 2HCl(aq)  2NaCl(aq) + H 2 S(g) Na 2 CO 3 (aq) + 2HCl(aq)  2NaCl(aq) + H 2 O(l) + CO 2 (g) Na 2 SO 3 (aq) + 2HCl(aq)  2NaCl(aq) + H 2 O(l) + SO 2 (g)

43. Copyright © Cengage Learning. All rights reserved. 4 | 43 Write the molecular, ionic, and net ionic equations for the reaction of copper(II) carbonate with hydrochloric acid.

44. Copyright © Cengage Learning. All rights reserved. 4 | 44 Molecular Equation (Balance the reaction and include state symbols) CuCO 3 (s) + 2HCl(aq)  CuCl 2 (aq) + H 2 O(l) + CO 2 (g) Ionic Equation CuCO 3 (s) + 2H + (aq) + 2Cl − (aq)  Cu 2+ (aq) + 2Cl − (aq) + H 2 O(l) + CO 2 (g) Net Ionic Equation CuCO 3 (s) + 2H + (aq)  Cu 2+ (aq) + H 2 O(l) + CO 2 (g)

45. 45 Oxidation-Reduction Reactions (electron transfer reactions) 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO

46. Copyright © Cengage Learning. All rights reserved. 4 | 46 Half−reaction One of two parts of an oxidation–reduction reaction, one part of which involves a loss of electrons (or increase in oxidation number) and the other part of which involves a gain of electrons (or decrease in oxidation number).

47. Copyright © Cengage Learning. All rights reserved. 4 | 47 Oxidation The half−reaction in which there is a loss of electrons by a species (or an increase in oxidation number). Reduction The half−reaction in which there is a gain of electrons by a species (or a decrease in oxidation number).

48. Copyright © Cengage Learning. All rights reserved. 4 | 48 Oxidizing Agent A species that oxidizes another species; it is itself reduced. Reducing Agent A species that reduces another species; it is itself oxidized.

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50. Copyright © Cengage Learning. All rights reserved. 4 | 50 Common Oxidation–Reduction Reactions Combination reaction Decomposition reaction Displacement reaction Combustion reaction

51. Copyright © Cengage Learning. All rights reserved. 4 | 51 For example:2Na(s) + Cl 2 (g)  2NaCl(s)

52. Copyright © Cengage Learning. All rights reserved. 4 | 52 Decomposition Reaction A reaction in which a single compound reacts to give two or more substances. For example: 2HgO(s)  2Hg(l) + O 2 (g)

53. Copyright © Cengage Learning. All rights reserved. 4 | 53 Displacement Reaction A reaction in which an element reacts with a compound, displacing another element from it. For example: Zn(s) + 2HCl(aq)  H 2 (g) + ZnCl 2 (aq)

54. Copyright © Cengage Learning. All rights reserved. 4 | 54 Combustion Reaction A reaction in which a substance reacts with oxygen, usually with the rapid release of heat to produce a flame. For example: 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s)

55. 55 Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Zn is oxidized Zn Zn 2+ + 2e - Cu 2+ is reduced Cu 2+ + 2e - Cu Zn is the reducing agent Cu 2+ is the oxidizing agent Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Cu Cu 2+ + 2e - Ag + + 1e - Ag Ag + is reduced Ag + is the oxidizing agent

56. 56 Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. 1.Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1. 4.4

57. 57 4.The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 5.Group IA metals are +1, IIA metals are +2 and fluorine is always –1. HCO 3 - O = – 2H = +1 3x( – 2) + 1 + ? = – 1 C = +4 What are the oxidation numbers of all the elements in HCO 3 - ? 7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O 2 -, is –½.

58. 58 The Oxidation Numbers of Elements in their Compounds

59. Balancing Simple Oxidation−Reduction Reactions: Half−Reaction Method First, identify what is oxidized and what is reduced by determining oxidation numbers.

60. Copyright © Cengage Learning. All rights reserved. For the reaction Zn(s) + Ag + (aq)  Zn 2+ (aq) + Ag(s) 0+1+20 Zn is oxidized from 0 to +2. Ag + is reduced from +1 to 0.

61. Copyright © Cengage Learning. All rights reserved. Next, write the unbalanced half−reactions. Zn(s)  Zn 2+ (aq)(oxidation) Ag + (aq)  Ag(s)(reduction) Now, balance the charge in each half reaction by adding electrons. Zn(s)  Zn 2+ (aq) + 2e − (oxidation) e − + Ag + (aq)  Ag(s)(reduction)

62. Copyright © Cengage Learning. All rights reserved. Since the electrons lost in oxidation are the same as those gained in reduction, we need each half−reaction to have the same number of electrons. To do this, multiply each half−reaction by a factor so that when the half−reactions are added, the electrons cancel. Zn(s)  Zn 2+ (aq) + 2e − (oxidation) 2e − + 2Ag + (aq)  2Ag(s)(reduction)

63. Copyright © Cengage Learning. All rights reserved. 4 | 63 Lastly, add the two half−reactions together. Zn(s) + 2Ag + (aq)  Zn 2+ (aq) + 2Ag(s)

64. Copyright © Cengage Learning. All rights reserved. 4 | 64 Balance the following oxidation−reduction reaction: FeI 3 (aq) + Mg(s)  Fe(s) + MgI 2 (aq) The oxidation numbers are given below the reaction. FeI 3 (aq) + Mg(s)  Fe(s) + MgI 2 (s) +3−100+2−1 Now, write the half−reactions. Since Iodide is a spectator ion it is omitted at this point. Mg(s)  Mg 2+ (aq)(oxidation) Fe 3+ (aq)  Fe(s)(reduction)

65. Copyright © Cengage Learning. All rights reserved. 4 | 65 Balancing the half−reactions: Mg(s)  Mg 2+ (aq) + 2e − (oxidation) Fe 3+ (aq) + 3e −  Fe(s)(reduction) Multiply the oxidation half−reaction by 3 and the reduction half−reaction by 2. 3Mg(s)  3Mg 2+ (aq) + 6e − (oxidation) 2Fe 3+ (aq) + 6e −  2Fe(s)(reduction)

66. Copyright © Cengage Learning. All rights reserved. 4 | 66 Add the half−reactions together. 2Fe 3+ (aq) + 3Mg(s)  2Fe(s) + 3Mg 2+ (aq) Now, return the spectator ion, I −. 2FeI 3 (aq) + 3Mg(s)  2Fe(s) + 3MgI 2 (aq)

67. 67 Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Precipitation Acid-Base Redox (H 2 Displacement) Redox (Combination) Classify each of the following reactions.