1. CHAPTER 4 REACTIONS IN AQUEOUS SOLUTIONS CHE 106 4.1 General Properties of Aqueous Solutions

2. Solution Composition Terminology Review Solution: Solvent: Solute: Concentration: 4.1 General Properties of Aqueous Solutions

3. Molarity 4.1 General Properties of Aqueous Solutions Molarity is defined as: Example: Calculate the molarity of a solution made by dissolving 0.0575 mol NH 4 Cl in 400 mL of water.

4. Molarity 4.1 General Properties of Aqueous Solutions Example: Calculate the molarity of a solution made by dissolving 9.93g of sodium sulfate in enough water to form 650. mL. (MW of Na 2 SO 4 = 142). Example: A. Calculate the number of moles of HNO 3 in 2.0L of a 0.200 M HNO 3 solution. B. What volume of a 0.30M HNO 3 solution is required to supply 2.0 mol of HNO 3 ?

5. Molarity 4.1 General Properties of Aqueous Solutions Pure acetic acid is a liquid with a density of 1.049 g/mL at 25 0 C. Calculate the molarity of a solution of acetic acid prepared by dissolving 10.00mL of acetic acid at 25 o C in enough water to make a 100.0 mL solution. Acetic acid = C 2 H 4 O 2, MW 60.1 amu.

6. Molarity and Dilution 4.1 General Properties of Aqueous Solutions Definition of dilution: Key Idea: Therefore the formula for dilutions:

7. Molarity and Dilution 4.1 General Properties of Aqueous Solutions An experiment requires 150 mL of a 2M HCl solution. All that is available in the laboratory is a 10.0M HCl stock solution. How would you prepare the required solution?

8. Molarity and Dilution 4.1 General Properties of Aqueous Solutions How many mL of a 0.250M HCl solution would be required to do each of the following: a) prepare 50.0mL of a 0.100 M HCl sol’n. b) react with the OH- in 15.0 mL of a 0.200 M Ba(OH) 2 sol’n. c) dissolve 0.500 g CaCO3 according to the rxn: CaCO 3 + 2H +  Ca +2 + H 2 O + CO 2

9. Electrolytes 4.1 General Properties of Aqueous Solutions Electrolytes vs. Non-electrolytes: Strong vs. Weak Electrolytes:

10. Electrolytes 4.1 General Properties of Aqueous Solutions

12. Electrolytes 4.1 General Properties of Aqueous Solutions Strong vs. Weak Electrolytes and Equilibrium

13. Electrolytes 4.1 General Properties of Aqueous Solutions

14. Precipitation Reactions 4.2 Precipitation Reactions Precipitation reactions and solubility rules: Precipitation Reactions:

15. Precipitation Reactions 4.2 Precipitation Reactions

16. Precipitation Reactions 4.2 Precipitation Reactions SOLUBLE - All Nitrates (NO 3 -1 ), Acetates (C 2 H 3 O 2 -1 ) and Ammoniums (NH 4 + ) - All Group 1 metals, fluorides - All Chlorides, Bromides, Iodides: Except Ag, Hg, Pb - All Sulfates: Except Ca, Sr, Ba, Pb, Hg and Ag. INSOLUBLE - all sulfides (S -2 ) except those with Grp 1 and 2 metals and ammonium - all carbonates (CO 3 -2 ) except those of Grp 1 metals and ammonium - all phosphates (PO 4 -3 ) except those with Grp 1 metals and ammonium - all hydroxies (OH -1 ) except those with Grp 1 metals and Ba, Sr, Ca

17. Precipitation Reactions 4.2 Precipitation Reactions Use your solubility guidelines to predict whether each of these are soluble or insoluble in water: a. AgI b. Na 2 CO 3 c. BaCl 2 d. Al(OH) 3 e. Zn(CH 3 COO) 2

18. Precipitation Reactions 4.2 Precipitation Reactions Metathesis Reactions: Ion Exchange Format: These reactions are driven to completion by the removal of ions from solution. This is accomplished in 3 possible ways: 1. 2. 3.

19. Precipitation Reactions 4.2 Precipitation Reactions Driving Force 1: Formation of an insoluble precipitate Predict the precipitate that forms when aqueous solutions of barium chloride and potassium sulfate are mixed. Write the balanced equation for this process.

20. Precipitation Reactions 4.2 Precipitation Reactions Driving Force 2: Formation of a non electrolytes or weak electrolyte examples – often acid base reactions fall into this type of reaction. a. HCl(aq) + NaC 2 H 3 O 2 (aq)  HC 2 H 3 O 2 (aq) + NaCl(aq) b. HCl(aq) + NaOH(aq) + NaCl (aq) + H 2 O(l) Weak electrolytes

21. Precipitation Reactions 4.2 Precipitation Reactions Driving Force 3: The formation of a gas that is insoluble in H 2 O. The gases that are often formed from aqueous solutions are hydrogen sulfide and carbon dioxide. Reaction from lab! 1. HCl(aq) + NaHCO 3 (aq)  NaCl(s) + H2CO 3 (aq) 2. H 2 CO 3 (aq)  H 2 O(l) + CO 2 (g)

22. Ionic Equations 4.2 Precipitation Reactions Molecular Equation: Complete Ionic Equation: Net Ionic Equation

23. Ionic Equations 4.2 Precipitation Reactions To write the complete ionic equation: 1. Start with the balanced molecular equation 2. Break all soluble strong electrolytes (compounds with aq) into their ions 3. Bring down any compound that has an (s), (l) or (g) as is. Make sure to: 1. - Include the correct formula and charge of each ion. 2. - Indicate the correct number of each ion 3. - Write (aq) after each ion.

24. Ionic Equations 4.2 Precipitation Reactions Example: Write the molecular equation and complete ionic equation for the metathesis reactions between sodium phosphate and calcium chloride. Molecular: Complete Ionic:

25. Ionic Equations 4.2 Precipitation Reactions The net ionic equation simplifies the complete ionic equation and eliminates any spectator ions. To recognize spectator ions: look for ions that are present on both sides of the equation. To convert the complete ionic equation into a net ionic equation: - Cross out the spectator ions that are present - Write the “leftovers” as the net ionic equation.

26. Ionic Equations 4.2 Precipitation Reactions Complete ionic equation: 6 Na +1 (aq) + 2 PO 4 -3 (aq) + 3 Ca +2 (aq) + 6 Cl -1 (aq)  6 Na +1 (aq) + 6 Cl - 1 (aq) + Ca 3 (PO 4 ) 2 (s) Net Ionic:

27. Ionic Equations 4.2 Precipitation Reactions Examples: a. 3(NH 4 ) 2 CO 3 (aq) + 2 Al(NO 3 ) 3 (aq)  6 NH 4 NO 3 (aq) + Al 2 (CO 3 ) 3 (s) b. Pb(NO 3 ) 2 (aq) + Na 2 SO 4 (aq)  c. CuBr 2 (aq) + NaOH (aq) 

28. Ionic Equations 4.1 General Properties of Aqueous Solutions d. NaOH + H 2 SO 4 e. Mg + HCl f. Zn + CuSO 4 g. Na 2 CO 3 + HNO 3  NaNO 3 + H 2 O(l) + CO 2 (g)

29. Ionic Equations 4.2 Precipitation Reactions d. Separate samples of a solution of an unknown ionic compound are treated with dilute AgNO 3, Pb(NO 3 ) 2 and BaCl 2. Precipitates form in all three cases. Which of the following could be the anion of the unknown salt: Br -, CO 3 - 2, NO 3 -. e. Three solutions are mixed together to form a single solution. One contains 0.2 mol Pb(CH 3 COO) 2, the second 0.1 mol Na 2 S and the third contains 0.1 mol CaCl 2. Write the net ionic equations for the precipitation reactions that occur. What are the spectator ions in the solution?

30. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Acids are considered proton donors: Mono, Di, Triprotic acids: H 3 PO 3 (aq)  H 2 PO 3 -1 (aq)  HPO 3 -2 (aq)  Ionization of Diprotic acid example: H 2 SO 4 

31. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Bases can be defined two ways: 1. 2.

32. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions As acids and bases are added, the concentrations of [H+] and [OH-] change in response… When strong acids are added: When strong bases are added:

33. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Strength of Acids and Bases: Strong: Weak: Strong Acids: HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 3 and HClO 4. Strong Bases: Group 1 and 2 metal hydroxides.

34. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Summary of Electrolytic Properties: Strong Electrolytes: Weak Electrolytes: Nonelectrolytes:

35. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Neutralization Reactions: Basic Format: Molecular Equation: Ionic Equation: Net Ionic Equation:

36. Acid Base Chemistry Introduction 4.3 Acids, Bases and Neutralization Reactions Complete and balance each of the following: 1. Barium hydroxide reacts with nitric acid: 2. Phosphoric acid reacts with lithium hydroxide: 3. Acetic acid is neutralized with sodium hydroxide 4. Ammonia is added to hydrochloric acid:

37. Acid Base Chemistry Introduction 4.1 General Properties of Aqueous Solutions Because there are other compounds that are capable of accepting a Hydrogen ion, some neutralization reactions result in products other than a salt and water. Two ions that often will accept a H+ ion: Example 1: Example 2:

38. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions Example: 4 Na(s) + O 2 (g)  2 Na 2 O (s) Oxidation: Reduction: Corrosion:

39. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions Assigning Oxidation Number Rules: 1. Oxidation number of an element = 0. 2. Oxidation number of a monoatomic ion = charge 3. For binary compounds: the element with the greater EN is assigned the negative oxidation number. 4. Sum of oxidation numbers must equal the charge of the overall molecule.

40. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions 5. Assign oxidation numbers in the following order: Group 1 metals = +1 Group 2 metals = +2 Aluminum = +3 Fluorine = -1 Hydrogen = +1 Oxygen = -2 Halides = -1 Assign the top oxidation number starting from elements from the right.

41. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions Determine the oxidation numbers of each element in the following compounds: HCCl 3 NH 4 +1 HNO 2 SCl 2 ClF 3

42. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions In addition to reacting with gases in the atmosphere, metals also react with acids or other salts in a displacement reaction, often in an aqueous solution. Basic Format: A + BX  AX + B Example: Zn (s) + 2 HCl (aq)  ZnCl 2 (aq) + H 2 (g) Net Ionic: Example: Li + Zn(NO 3 ) 2  LiNO 3 + Zn Net Ionic:

43. Oxidation and Reduction 4.1 General Properties of Aqueous Solutions How do we predict when oxidation and reduction reactions will occur? Activity Series: Using the activity series to predict reactions:

44. 4.1 General Properties of Aqueous Solutions Alkali Metals: More Easily Oxidized Noble Metals: More Difficult Oxidation Activity Series In order to react displace the H+ ion in an acid, the metal must be above Hydrogen.

45. 4.1 General Properties of Aqueous Solutions If we compare iron to copper: Can iron oxidize copper? Can copper oxidize iron? Does copper react with H+ ions? Does nickel react with H+ ions? Is Nickel oxidized by copper? Activity Series

46. 4.1 General Properties of Aqueous Solutions Which of the following metals will be oxidized by ZnSO 4 ? Lithium Silver Chromium Nickel Hydrogen Magnesium Activity Series

47. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions We can use stoichiometry principles and apply it to reactions within an aqueous environment to predict the amount of products or reactants necessary. Example: How much Ba(OH) 2 is necessary to neutralize 15 mL of 0.25M HCl?

48. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions To experimentally determine the concentration of an unknown solution, we perform titrations. Standard Solution: Titration: Equivalence Point:

49. Solution Stoichiometry Chapt. 4.5 Grams of Compound A Grams of Compound B Moles of Compound A Moles of Compound B Volume or Molarity of Compound A of Compound A Volume or Molarity of Compound B of Compound B use MW use MW use M=mol/V use M=mol/V use bal. eqn.

50. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions Titration Formula: M (strd) V (strd) = M (unk) V (unk) Example: What volume of 0.827 M KOH solution is required to completely neutralize 35.00 mL of 0.737 H 2 SO 4 solution?

51. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions Example: A 17.5 mL sample of acetic acid (CH 3 CO 2 H) required 29.6 mL of 0.250M NaOH to neutralize it. What was the molarity of the acetic acid solution? Example: Acetic acid (HC2H3O2) reacts with sodium hydroxide in a neutralization reaction. If 8.30 mL of an acetic acid solution requires 42.4 mL of 0.1050 M NaOH to reach the equivalence point, how many grams of acetic acid are in a 500 mL sample?

52. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions Example: In the laboratory, 6.67 g of Sr(NO 3 ) 2 was dissolved to make 750 mL of solution. A 100 mL sample of this solution was titrated with 0.0460 M Na 2 CrO 4 solution. What volume of the Na 2 CrO 4 solution is required to precipitate all of the Sr +2 as SrCrO 4 ? (MW Sr(NO 3 ) 2 = 211.6 Sr(NO 3 ) 2 (aq) + Na 2 CrO 4 (aq)  SrCrO 4 (s) + NaNO 3 (aq)

53. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions Example: The arsenic in a 1.22 gram sample of pesticide was chemically converted to AsO 4 3-. It was then titrated with Ag +1 to form Ag 3 AsO 4 as a precipitate. If it took 25.0 mL of a 0.102 M Ag +1 solution to reach the equivalence point, what is the percentage arsenic is the pesticide? Equation:

54. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions A 4.36 g sample of an unknown alkali metal hydroxide is dissolved in 100.0 mL of water. An acid base indicator is added and the resulting solution is titrated with 2.50 M HCl. The indicator changes color signaling that the equivalence point has been reached after 17.0 mL of the hydrochloric acid solution has been added. a) What is the molar mass of the metal hydroxide? b) What is the identify of the metal cation: Li, Na, K, Rb, or Ca?

55. Solution Stoichiometry 4.1 General Properties of Aqueous Solutions A 1.248 sample of limestone rock is pulverized and treated with 30.00 mL of 1.035 M HCl. The excess acid then requires 11.56 of 1.010 M NaOH for titration. Calculate the percent by mass of calcium carbonate in the rock, assuming that it is the only substance reacting with the HCl solution.