1. 1 YAY! More Stoichiometry!

2. Solutions:  Homogeneous mixtures of two or more pure substances.  The solvent is present in greatest abundance.  All other substances are solutes. 2

3.  Guess what else you must memorize! Table 4.1 3

4. Classify the following ionic compounds as soluble or insoluble in water: (a) sodium carbonate (Na 2 CO 3 ), (b) lead sulfate (PbSO 4 ). Classify the following compounds as soluble or insoluble in water: (a) cobalt(II) hydroxide, (b) barium nitrate, (c) ammonium phosphate. 4

5. Dissociation  When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.  This process is called dissociation. 5

6. Electrolytes  Substances that dissociate into ions when dissolved in water.  A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so. 6

7. Electrolytes and Nonelectrolytes Soluble ionic compounds tend to be electrolytes. 7

8. Electrolytes and Nonelectrolytes Molecular compounds tend to be nonelectrolytes, except for acids and bases. 8

9. Electrolytes  A strong electrolyte dissociates completely when dissolved in water.  A weak electrolyte only dissociates partially when dissolved in water. 9

10. Strong Electrolytes Are… 10 Strong acids Strong bases Soluble ionic salts

11. Strong Electrolytes Are… 11 Strong acids Strong bases Soluble ionic salts

12. Strong Electrolytes Are… Strong acids Strong bases Soluble ionic salts 12

13. Identifying Strong, Weak, and Nonelectrolytes Classify each of the following dissolved substances as a strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl 2, HNO 3, C 2 H 5 OH (ethanol), HCHO 2 (formic acid), KOH. 13

14. PRACTICE EXERCISE Consider solutions in which 0.1 mol of each of the following compounds is dissolved in 1 L of water: Ca(NO 3 ) 2 (calcium nitrate), C 6 H 12 O 6 (glucose), NaC 2 H 3 O 2 (sodium acetate), and HC 2 H 3 O 2 (acetic acid). Rank the solutions in order of increasing electrical conductivity, based on the fact that the greater the number of ions in solution, the greater the conductivity. 14

15. Precipitation Reactions When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed. 15

16.  16 Metathesis (Exchange) Reactions  Metathesis comes from a Greek word that means “to transpose”  Formerly called double replacement reactions AgNO 3 ( aq ) + KCl ( aq )  AgCl ( s ) + KNO 3 ( aq )

17.  17 Metathesis (Exchange) Reactions  It appears the ions in the reactant compounds exchange, or transpose, ions AgNO 3 ( aq ) + KCl ( aq )  AgCl ( s ) + KNO 3 ( aq )

18.  18 Metathesis (Exchange) Reactions AgNO 3 ( aq ) + KCl ( aq )  AgCl ( s ) + KNO 3 ( aq )

19. SAMPLE EXERCISE 4.3 Predicting a Metathesis Reaction (a) Predict the identity of the precipitate that forms when solutions of BaCl 2 and K 2 SO 4 are mixed. (b) Write the balanced chemical equation for the reaction. PRACTICE EXERCISE (a) What compound precipitates when solutions of Fe 2 (SO 4 ) 3 and LiOH are mixed? (b) Write a balanced equation for the reaction. (c) Will a precipitate form when solutions of Ba(NO 3 ) 2 and KOH are mixed? 19

20.   It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution).  If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction. 20 Solution Chemistry

21.  21 Molecular Equation The molecular equation lists the reactants and products in their molecular form. AgNO 3 ( aq ) + KCl ( aq )  AgCl ( s ) + KNO 3 ( aq )

22.  22 Ionic Equation  In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions.  This more accurately reflects the species that are found in the reaction mixture. Ag + ( aq ) + NO 3 - ( aq ) + K + ( aq ) + Cl - ( aq )  AgCl ( s ) + K + ( aq ) + NO 3 - ( aq )

23.  23 Net Ionic Equation  To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. Ag + (aq) + NO 3 - ( aq ) + K + (aq) + Cl - ( aq )  AgCl ( s ) + K + (aq) + NO 3 - ( aq )

24.  24 Net Ionic Equation  To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right.  The only things left in the equation are those things that change (i.e., react) during the course of the reaction.  Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions. Ag + (aq) + NO 3 - ( aq ) + K + (aq) + Cl - ( aq )  AgCl ( s ) + K + (aq) + NO 3 - ( aq )

25.  25 Net Ionic Equation  The only things left in the equation are those things that change (i.e., react) during the course of the reaction. Ag + (aq) + Cl - ( aq )  AgCl ( s )

26.  1.Write a balanced molecular equation. 2.Dissociate all strong electrolytes. 3.Cross out anything that remains unchanged from the left side to the right side of the equation. 4.Write the net ionic equation with the species that remain. 26 Steps to Writing Net Ionic Equations

27. SAMPLE EXERCISE 4.4 Writing a Net Ionic Equation Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed. 27

28. PRACTICE EXERCISE Write the net ionic equation for the precipitation reaction that occurs when aqueous solutions of silver nitrate and potassium phosphate are mixed. 28

29.  Ye olde Brønsted–Lowry definition of acids & bases… acid is defined as anything that releases H 1+ ions base is defined as anything that accepts H 1+ ions. The Brønsted-Lowry concept is based on the transfer of a proton from one substance to another. i.e. an acid is a proton donor; a base is a proton acceptor. 29

30. Acids:  Substances that increase the concentration of H + when dissolved in water (Arrhenius).  Proton donors (Brønsted–Lowry). 30

31. There are only seven strong acids: Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Nitric (HNO 3 ) Sulfuric (H 2 SO 4 ) Chloric (HClO 3 ) Perchloric (HClO 4 ) 31

32. Solution Analyze: We are asked to rank three acids from strongest to weakest, based on schematic drawings of their solutions. Plan: We can examine the drawings to determine the relative numbers of uncharged molecular species present. The strongest acid is the one with the most H + ions and fewest undissociated acid molecules in solution. The weakest is the one with the largest number of undissociated molecules. Solve: The order is HY > HZ > HX. HY is a strong acid because it is totally ionized (no HY molecules in solution), whereas both HX and HZ are weak acids, whose solutions consist of a mixture of molecules and ions. Because HZ contains more H + ions and fewer molecules than HX, it is a stronger acid. SAMPLE EXERCISE 4.5 Comparing Acid Strengths The following diagrams represent aqueous solutions of three acids (HX, HY, and HZ) with water molecules omitted for clarity. Rank them from strongest to weakest. 32

33. PRACTICE EXERCISE Imagine a diagram showing ten Na + ions and ten OH – ions. If this solution were mixed with the one pictured on the previous page for HY, what would the diagram look like that represents the solution after any possible reaction? (H + ions will react with OH – ions to form H 2 O. ) SAMPLE EXERCISE 4.5 continued 33

34. Bases:  Substances that increase the concentration of OH − when dissolved in water (Arrhenius).  Proton acceptors (Brønsted–Lowry). 34

35. The strong bases are the soluble salts of hydroxide ion: Alkali metals Calcium Strontium Barium 35

36. Acid-Base Reactions In an acid-base reaction, the acid donates a proton (H + ) to the base. 36

37.  37 Neutralization Reactions Generally, when solutions of an acid and a base are combined, the products are a salt and water. HCl ( aq ) + NaOH ( aq )  NaCl ( aq ) + H 2 O ( l )

38.  38 Neutralization Reactions When a strong acid reacts with a strong base, the net ionic equation is… HCl ( aq ) + NaOH ( aq )  NaCl ( aq ) + H 2 O ( l ) H + ( aq ) + Cl - ( aq ) + Na + ( aq ) + OH - ( aq )  Na + ( aq ) + Cl - ( aq ) + H 2 O ( l )

39.  39 Neutralization Reactions When a strong acid reacts with a strong base, the net ionic equation is… HCl ( aq ) + NaOH ( aq )  NaCl ( aq ) + H 2 O ( l ) H + ( aq ) + Cl - ( aq ) + Na + ( aq ) + OH - ( aq )  Na + ( aq ) + Cl - ( aq ) + H 2 O ( l ) H + ( aq ) + Cl - ( aq) + Na + ( aq ) + OH - ( aq )  Na + ( aq ) + Cl - ( aq ) + H 2 O ( l )

40. Neutralization Reactions Observe the reaction between Milk of Magnesia, Mg(OH) 2, and HCl. 40

41. SAMPLE EXERCISE 4.7 Writing Chemical Equations for a Neutralization Reaction (a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (HC 2 H 3 O 2 ) and barium hydroxide [Ba(OH) 2 ]. (b) Write the net ionic equation for this reaction. 41

42. PRACTICE EXERCISE (a) Write a balanced molecular equation for the reaction of carbonic acid (H 2 CO 3 ) and potassium hydroxide (KOH). (b) Write the net ionic equation for this reaction. 42

43.  43 Gas-Forming Reactions  These metathesis reactions do not give the product expected.  The expected product decomposes to give a gaseous product (CO 2 or SO 2 ). CaCO 3 ( s ) + HCl ( aq )  CaCl 2 ( aq ) + CO 2 ( g ) + H 2 O ( l ) NaHCO 3 ( aq ) + HBr ( aq )  NaBr ( aq ) + CO 2 ( g ) + H 2 O ( l ) SrSO 3 ( s ) + 2 HI ( aq )  SrI 2 ( aq ) + SO 2 ( g ) + H 2 O ( l )

44.  44 Gas-Forming Reactions  This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular!  Just as in the previous examples, a gas is formed as a product of this reaction: Na 2 S ( aq ) + H 2 SO 4 ( aq )  Na 2 SO 4 ( aq ) + H 2 S ( g )

45. Oxidation-Reduction Reactions  An oxidation occurs when an atom or ion loses electrons.  A reduction occurs when an atom or ion gains electrons. 45

46. Oxidation-Reduction Reactions One cannot occur without the other. 46

47.  To determine if an oxidation-reduction reaction has occurred, we assign an oxidation number to each element in a neutral compound or charged entity. 47 Oxidation Numbers

48.   Elements in their elemental form have an oxidation number of 0.  The oxidation number of a monatomic ion is the same as its charge. 48 Rules for Assigning Oxidation Numbers

49.   Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.  Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.  Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal. 49 Rules continued

50.   Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.  Fluorine always has an oxidation number of −1.  The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions. 50 Rules continued

51.   The sum of the oxidation numbers in a neutral compound is 0.  The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. 51 Oxidation Numbers

52.  52 Oxidation Numbers examples Determine the oxidation number of sulfur in each of the following: H 2 S S 8 SCl 2 Na 2 SO 3 SO 4 2–

53. PRACTICE EXERCISE What is the oxidation state of the boldfaced element in each of the following: (a) P 2 O 5, (b) NaH, (c) Cr 2 O 7 –2, (d) SnBr 4, (e) BaO 2 ? 53

54. Displacement Reactions previously called single replacement  In displacement reactions, ions oxidize an element.  The ions, then, are reduced. 54 Remind me to set up copper-silver demonstration

55. Displacement Reactions In this reaction, silver ions oxidize copper metal. Cu ( s ) + 2 Ag + ( aq )  Cu 2+ ( aq ) + 2 Ag ( s ) 55

56. Displacement Reactions The reverse reaction, however, does not occur. Cu 2+ ( aq ) + 2 Ag ( s )  Cu ( s ) + 2 Ag + ( aq ) 56 x

57.  57 Activity Series Guess what else you kinda need to know off the top of your head…

58. SAMPLE EXERCISE 4.9 Writing Molecular and Net Ionic Equations for Oxidation-Reduction Reactions Write the balanced molecular and net ionic equations for the reaction of aluminium with hydrobromic acid. 58

59. PRACTICE EXERCISE (a) Write the balanced molecular and net ionic equations for the reaction between magnesium and cobalt(II) sulfate. (b) What is oxidized and what is reduced in the reaction? 59

60. SAMPLE EXERCISE 4.10 Determining When an Oxidation-Reduction Reaction Can Occur Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction. 60

61. 61 PRACTICE EXERCISE Which of the following metals will be oxidized by Pb(NO 3 ) 2 : Zn, Cu, Fe?

62.  62 Molarity  Two solutions can contain the same compounds but be quite different because the proportions of those compounds are different.  Molarity is one way to measure the concentration of a solution. moles of solute volume of solution in liters Molarity (M) =

63. SAMPLE EXERCISE 4.11 Calculating Molarity Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate (Na 2 SO 4 ) in enough water to form 125 mL of solution. 63

64. PRACTICE EXERCISE Calculate the molarity of a solution made by dissolving 5.00 g of glucose (C 6 H 12 O 6 ) in sufficient water to form exactly 100 mL of solution. 64

65. SAMPLE EXERCISE 4.12 Calculating Molar Concentrations of Ions What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate? 65

66. 66 Another PRACTICE EXERCISE What is the molar concentration of K + ions in a 0.015 M solution of potassium carbonate?

67. SAMPLE EXERCISE 4.13 Using Molarity to Calculate Grams of Solute How many grams of Na 2 SO 4 are required to make 0.350 L of 0.500 M Na 2 SO 4 ? 67

68. PRACTICE EXERCISE (a) How many grams of Na 2 SO 4 are there in 15 mL of 0.50 M Na 2 SO 4 ? (b) How many milliliters of 0.50 M Na 2 SO 4 solution are needed to provide 0.038 mol of this salt? 68

69.  69 Mixing a Solution

70.  70 Dilution

71. SAMPLE EXERCISE 4.14 Preparing a Solution by Dilution How many milliliters of 3.0 M H 2 SO 4 are needed to make 450 mL of 0.10 M H 2 SO 4 ? 71

72. Another PRACTICE EXERCISE (a) What volume of 2.50 M lead(II)nitrate solution contains 0.0500 mol of Pb 2+ ? (b) How many milliliters of 5.0 M K 2 Cr 2 O 7 solution must be diluted to prepare 250 mL of 0.10 M solution? (c) If 10.0 mL of a 10.0 M stock solution of NaOH is diluted to 250 mL, what is the concentration of the resulting stock solution? 72

73.  73 Using Molarities in Stoichiometric Calculations A Molar Road Map

74. SAMPLE EXERCISE 4.15 Using Mass Relations in a Neutralization Reaction How many grams of Ca(OH) 2 are needed to neutralize 25.0 mL of 0.100 M HNO 3 ? 74

75. Another PRACTICE EXERCISE (a) How many grams of NaOH are needed to neutralize 20.0 mL of 0.150 M H 2 SO 4 solution? (b) How many liters of 0.500 M HCl(aq) are needed to react completely with 0.100 mol of Pb(NO 3 ) 2 (aq), forming a precipitate of PbCl 2 (s)? 75

76. Titration The analytical technique in which one can calculate the concentration of a solute in a solution. 76

77.  77 Titration

78. SAMPLE EXERCISE 4.16 Determining the Quantity of Solute by Titration The quantity of Cl – in a municipal water supply is determined by titrating the sample with Ag +. The reaction taking place during the titration is The end point in this type of titration is marked by a change in color of a special type of indicator. (a) How many grams of chloride ion are in a sample of the water if 20.2 mL of 0.100 M Ag + is needed to react with all the chloride in the sample? (b) If the sample has a mass of 10.0 g, what percent Cl – does it contain? 78

79. (a) How many moles of MnO 4 – were added to the solution? (b) How many moles of Fe + were in the sample? (c) How many grams of iron were in the sample? (d) If the sample had a mass of 0.8890 g, what is the percentage of iron in the sample? Another PRACTICE EXERCISE A sample of an iron ore is dissolved in acid, and the iron is converted to Fe +. The sample is then titrated with 47.20 mL of 0.02240 M MnO 4 – solution. The oxidation-reduction reaction that occurs during titration is as follows: 79

80. SAMPLE EXERCISE 4.17 Determining Solution Concentration via an Acid-Base Titration One commercial method used to peel potatoes is to soak them in a solution of NaOH for a short time, remove them from the NaOH, and spray off the peel. The concentration of NaOH is normally in the range of 3 to 6 M. The NaOH is analyzed periodically. In one such analysis, 45.7 mL of 0.500 M H 2 SO 4 is required to neutralize a 20.0-mL sample of NaOH solution. What is the concentration of the NaOH solution? 80

81. Another PRACTICE EXERCISE What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H 2 SO 4 ? 81

82. INTEGRATIVE EXERCISE Putting Concepts Together Note: Integrative exercises require skills from earlier chapters as well as ones from the present chapter. A sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M silver nitrate, resulting in the formation of a precipitate. (a) Write the molecular equation for the reaction. (b) What is the limiting reactant in the reaction? (c) Calculate the theoretical yield, in grams, of the precipitate that forms. 82

83. Table 4.1 BACK 83