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Chapter 4 Aqueous Reactions and Solution Stoichiometry

Chapter 4 Aqueous Reactions and Solution Stoichiometry

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Chapter 4 Aqueous Reactions and Solution Stoichiometry

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  1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 4Aqueous Reactions and Solution Stoichiometry

  2. SAMPLE EXERCISE 4.6 Identifying Strong, Weak, and Nonelectrolytes Classify each of the following dissolved substances as a strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl2 , HNO3, C2H5OH (ethanol), HCHO2 (formic acid), KOH. Solution Solve: Two compounds fit the criteria for ionic compounds: CaCl2 and KOH. As Table 2.3 tells us that all ionic compounds are strong electrolytes, that is how we classify these two substances. The three remaining compounds are molecular. Two, HNO3 and HCHO2 , are acids. Nitric acid, HNO3 is a common strong acid, as shown in Table 4.2, and therefore is a strong electrolyte. Because most acids are weak acids, our best guess would be that HCHO2 is a weak acid (weak electrolyte). This is correct. The remaining molecular compound, C2H5OH is neither an acid nor a base, so it is a nonelectrolyte. Comment: Although C2H5OH has an OH group, it is not a metal hydroxide; thus, it is not a base. Rather, it is a member of a class of organic compounds that have C––OH bonds, which are known as alcohols. (Section 2.9)

  3. SAMPLE EXERCISE 4.6 continued PRACTICE EXERCISE Consider solutions in which 0.1 mol of each of the following compounds is dissolved in 1 L of water: Ca(NO3)2 (calcium nitrate), C6H12O6 (glucose), NaC2H3O2 (sodium acetate), and HC2H3O2 (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. Answers: C6H12O6 (nonelectrolyte) < HC2H3O2 (weak electrolyte, existing mainly in the form of molecules with few ions) < NaC2H3O2 (strong electrolyte that provides two ions, and (strong electrolyte that provides three ions, Ca2+ and 2 NO3–

  4. These metathesis reactions do not give the product expected. The expected product decomposes to give a gaseous product (CO2 or SO2). CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l) NaHCO3 (aq) + HBr (aq) NaBr (aq)+ CO2 (g) + H2O (l) SrSO3 (s) + 2 HI(aq) SrI2 (aq) + SO2 (g) + H2O (l) Gas-Forming Reactions

  5. 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: Na2S (aq) + H2SO4 (aq) Na2SO4 (aq) + H2S (g) Gas-Forming Reactions

  6. Summary of gas rxns • H2S is a gas If you produce: • H2CO3then it breaks down to H2O + CO2(g) • H2SO3then it breaks down to H2O + SO2(g) • NH4OH then it breaks down to H2O + NH3(g)

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

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

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

  10. Rules for assigning oxidation numbers A metal or a nonmetal in the free state has an oxidation number of zero. Mg S8 Cl2 A monoatomic ion has an oxidation number equal to its ionic charge K+ = +1 Mg2+=+2 P3-=-3 A hydrogen atom is usually +1 HCl H3PO4

  11. Rules for assigning oxidation numbers An oxygen atom is usually -2 MgO Al2O3 For a molecular compound (2 nonmetals bonded together), the total compound charge must equal zero CCl4 P2O5

  12. Rules for assigning oxidation numbers For an ionic compound (a metal with a nonmetal) the sum of the oxidation numbers must equal zero BaCl2 Fe2O3 Ir2O7 For a polyatomic ion, the sum of the oxidation numbers for each of the atoms in the compound is equal to the ionic charge on the polyatomic ion. NO3- SO42- PO43-

  13. SAMPLE EXERCISE 4.8 Determining Oxidation Numbers Determine the oxidation number of sulfur in each of the following: (a) H2S, (b) S8 , (c) SCl2, (d) Na2SO3, (e) SO42–.

  14. PRACTICE EXERCISE What is the oxidation state of the boldfaced element in each of the following: (a) P2O5 , (b) NaH, (c) Cr2O7–2, (d) SnBr4, (e) BaO2? Answers:(a) +5, (b) –1, (c) +6, (d) +4, (e) –1

  15. BrO3- C2O42- F2 CaH2 H2SiO4 SO42-

  16. Oxidation Reduction Rxns Also known as “redox” These types of reactions have a transfer of electrons LEO the lion says GER Lose Electrons Oxidized Gain Electrons Reduced

  17. Oxidation Reduction Rxns Fe(s) + S(s)  FeS(s)

  18. Oxidation Reduction Rxns Fe2O3(l) + 2Al(l) 2Fe(l) + Al2O3(l)

  19. Oxidation Reduction Rxns I2(s) + SO32-(aq) + H2O(l) 2I-(aq) + SO42-(aq) + 2H+(aq)

  20. Displacement Reactions • In displacement reactions (aka single replacement rxns), ions oxidize an element. • The ions, then, are reduced. • ALL single replacement rxns are redox!!!!

  21. Displacement Reactions In this reaction, silver ions oxidize copper metal. Cu (s) + 2 Ag+ (aq) Cu2+ (aq) + 2 Ag (s)

  22. Displacement Reactions The reverse reaction, however, does not occur. Cu2+ (aq) + 2 Ag (s)  Cu (s) + 2 Ag+ (aq) x

  23. Activity Series

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

  25. Acids There are only seven strong acids: • Hydrochloric (HCl) • Hydrobromic (HBr) • Hydroiodic (HI) • Nitric (HNO3) • Sulfuric (H2SO4) • Chloric (HClO3) • Perchloric (HClO4)

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

  27. Bases (know your ABCS!) The strong bases are the soluble salts of hydroxide ion: • Alkali metals • Barium • Calcium • Strontium

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

  29. Generally, when solutions of an acid and a base are combined, the products are a salt and water. HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) Neutralization Reactions

  30. When a strong acid reacts with a strong base, the net ionic equation is… HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) H+ (aq)+ Cl- (aq)+ Na+ (aq) + OH-(aq) Na+ (aq)+ Cl- (aq)+ H2O (l) Neutralization Reactions

  31. When a strong acid reacts with a strong base, the net ionic equation is… HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) H+ (aq)+ Cl- (aq)+ Na+ (aq) + OH-(aq) Na+ (aq)+ Cl- (aq)+ H2O (l) H+ (aq)+ Cl- (aq)+ Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l) Neutralization Reactions

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

  33. 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 (HC2H3O2) and barium hydroxide [Ba(OH)2]. (b) Write the net ionic equation for this reaction.

  34. PRACTICE EXERCISE (a) Write a balanced molecular equation for the reaction of carbonic acid (H2CO3) and potassium hydroxide (KOH). (b) Write the net ionic equation for this reaction. (H2CO3 is a weak acid and therefore a weak electrolyte, whereas KOH, a strong base, and K2CO3, an ionic compound, are strong electrolytes.)

  35. 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 Molarity (M) = volume of solution in liters Molarity

  36. SAMPLE EXERCISE 4.11 Calculating Molarity Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate (Na2SO4) in enough water to form 125 mL of solution. Thus, the molarity is

  37. SAMPLE EXERCISE 4.11continued PRACTICE EXERCISE Calculate the molarity of a solution made by dissolving 5.00 g of glucose (C6H12O6) in sufficient water to form exactly 100. mL of solution. Answer: 0.278 M

  38. 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? Solve: Calcium nitrate is composed of calcium ions (Ca2+) and nitrate ions (NO3–) so its chemical formula is Ca(NO3)2. Because there are two NO3– ions for each Ca2+ ion in the compound, each mole of Ca(NO3)2 that dissolves dissociates into 1 mol of Ca2+ and 2 mol of NO3–. Thus, a solution that is 0.025 M in Ca(NO3)2 is 0.025 M in Ca2+ and 2  0.025 M = 0.050 M in NO3–. PRACTICE EXERCISE What is the molar concentration of K+ ions in a 0.015 M solution of potassium carbonate? Answer: 0.030 M K+

  39. PRACTICE EXERCISE (a) How many grams of Na2SO4 are there in 15 mL of 0.50 M Na2SO4? (b) How many milliliters of 0.50 M Na2SO4 solution are needed to provide 0.038 mol of this salt? Answers: (a) 1.1 g, (b) 76 mL

  40. Mixing a Solution

  41. Dilution M1 x V1 = M2 x V2

  42. Using Molarities inStoichiometric Calculations

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

  44. Titration

  45. Solve: The product of the molar concentration of a solution and its volume in liters gives the number of moles of solute: Because this is an acid-base neutralization reaction, HNO3 and Ca(OH)2 react to form and the salt containing Ca2+ and NO3 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 HNO3?

  46. PRACTICE EXERCISE (a) How many grams of NaOH are needed to neutralize 20.0 mL of 0.150 M H2SO4 solution? (b) How many liters of 0.500 M HCl(aq) are needed to react completely with 0.100 mol of Pb(NO3)2(aq), forming a precipitate of PbCl2(s)? Answers: (a) 0.240 g, (b) 0.400 L

  47. 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? From the balanced equation we see that Using this information and the molar mass of Cl, we have

  48. SAMPLE EXERCISE 4.16 continued (b) Plan: To calculate the percentage of Cl– in the sample, we compare the number of grams of Cl– in the sample, 7.17  10–2 g, with the original mass of the sample, 10.0 g. Solve: Comment: Chloride ion is one of the most common ions in water and sewage. Ocean water contains 1.92% Cl–. Whether water containing Cl– tastes salty depends on the other ions present. If the only accompanying ions are Na+, a salty taste may be detected with as little as 0.03% Cl–.

  49. PRACTICE EXERCISE A sample of an iron ore is dissolved in acid, and the iron is converted to Fe2+. The sample is then titrated with 47.20 mL of 0.02240 M MnO4– solution. The oxidation-reduction reaction that occurs during titration is as follows: (a) How many moles of MnO4– were added to the solution? (b) How many moles of Fe2+ 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? Answers: (a) 1.057  10–3 mol MnO4–(b) 5.286  10–3 mol Fe2+(c) 0.2952 g, (d) 33.21%

  50. SAMPLE 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. (b) AgNO3 is the limiting reactant