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Chapter 5

Chapter 5. Reactions in Aqueous Solutions. K + (aq) + MnO 4 - (aq). Reactions in Aqueous Solution. Many reactions involve ionic compounds, especially reactions in water — aqueous solutions. KMnO 4 in water. An Ionic Compound, CuCl 2 , in Water. CCR, page 177. Aqueous Solutions.

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Chapter 5

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  1. Chapter 5 Reactions in Aqueous Solutions Dr. S. M. Condren

  2. K+(aq) + MnO4-(aq) Reactions in Aqueous Solution Many reactions involve ionic compounds, especially reactions in water — aqueous solutions. KMnO4 in water Dr. S. M. Condren

  3. An Ionic Compound, CuCl2, in Water CCR, page 177 Dr. S. M. Condren

  4. Aqueous Solutions How do we know ions are present in aqueous solutions? The solutions conduct electricity! They are called ELECTROLYTES HCl, CuCl2, and NaCl are strong electrolytes. They dissociate completely (or nearly so) into ions. Dr. S. M. Condren

  5. Aqueous Solutions HCl, CuCl2, and NaCl are strong electrolytes.They dissociate completely (or nearly so) into ions. Dr. S. M. Condren

  6. Aqueous Solutions Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq) ---> CH3CO2-(aq) + H+(aq) Dr. S. M. Condren

  7. Aqueous Solutions Acetic acid ionizes only to a small extent, so it is a weak electrolyte. CH3CO2H(aq) ---> CH3CO2-(aq) + H+(aq) Dr. S. M. Condren

  8. Aqueous Solutions Some compounds dissolve in water but do not conduct electricity. They are called nonelectrolytes. Examples include: sugar ethanol ethylene glycol Dr. S. M. Condren

  9. Electrolytes • Conduct electricity in solution due to the presence of ions • Strong electrolyte – completely ionized in solution • Weak electrolyte – partially ionized in solution • Non-electrolyte – nonionic solution Dr. S. M. Condren

  10. Solubility Rules 1. All nitrates (NO3-1) are soluble. 2. All compounds of Group IA metals and the ammonium ion, NH4+, are soluble. 3. All chlorides are soluble except: AgCl, Hg2Cl2 and PbCl2. 4. All sulfates are soluble except: PbSO4, BaSO4, and SrSO4. Dr. S. M. Condren

  11. Solubility Rules 5. All hydroxides (OH-1)and sulfides (S-2)are insoluble except those of the Group IA metals and the ammonium ion. 6. All carbonates (CO3-2) and phosphates (PO4-3) are insoluble except those of the Group IA metals and the ammonium ion. Dr. S. M. Condren

  12. Net Ionic Equation Balanced Chemical Equation Pb(NO3)2 + Na2SO4 ---> PbSO4 + 2NaNO3 Total Ionic Equation Pb+2 + 2NO3-1 + 2Na+1 + SO4-2 2Na+ + 2NO3-1 + PbSO4 Net Ionic Equation Pb+2 + SO4-2 PbSO4 Dr. S. M. Condren

  13. A solution of Ba(NO3)2 is added to a solution of Na2SO4 to make a precipitate. From a table of solubility rules, the product is barium sulfate, sodium nitrate Dr. S. M. Condren

  14. Types of Reactions • synthesis reactions or combination reactions • decomposition reactions • precipitation reactions • neutralization reactions • acid • base • oxidation-reduction reaction Dr. S. M. Condren

  15. Synthesis or Combination Reactions Formation of a compound from simpler compounds or elements. 2Na(s) + Cl2(g) 2 NaCl(s) Dr. S. M. Condren

  16. Decomposition Reactions Separation into constituents by chemical reaction. catalysis 2 H2O2 (aq) 2H2O(l) + O2(g) Dr. S. M. Condren

  17. Precipitation Reactions The process of separating a substance from a solution as a solid. AgNO3 + NaCl ---> AgCl + NaNO3 precipitate Total Ionic Equation Ag+1(aq) + NO3(aq)-1 + Na(aq)+1 + Cl(aq)-1 AgCl(s) + NO3(aq)-1 + Na(aq)+1 Net Ionic Equation Ag+1(aq) + Cl(aq)-1 AgCl(s) Dr. S. M. Condren

  18. Neutralization Reactions acid + base ---> “salt” + water Balanced Chemical Equation HCl + NaOH ---> NaCl + H2O Total Ionic Equation H+1 + Cl-1 + Na+1 + OH-1 Na+ + Cl-1 + H2O Net Ionic Equation H+1 + OH-1 H2O Dr. S. M. Condren

  19. Neutralization Reactions acid + base ---> “salt” + water Balanced Chemical Equation H2SO4 + 2KOH ---> K2SO4 + 2H2O Total Ionic Equation 2H+ + SO4-2 + 2K+ + 2OH-1 2H2O + SO4-2 + 2K+ Net Ionic Equation 2H+1 + 2OH-1 2H2O Dr. S. M. Condren

  20. Neutralization Reactions • acid • base • salt Household acids and Bases Dr. S. M. Condren

  21. Neutralization Reactions • acid • Any of a large class of sour-tasting substances whose aqueous solutions are capable of turning blue litmus indicators red, of reacting with and dissolving certain metals to form salts, and of reacting with bases or alkalis to form salts. • Substance that donates H+ ions to solution Dr. S. M. Condren

  22. HNO3 ACIDS An acid -------> H+ in water Some strongacids are HCl hydrochloric H2SO4 sulfuric HClO4 perchloric HNO3 nitric Dr. S. M. Condren

  23. Acetic acid Weak Acids WEAK ACIDS = weak electrolytes CH3CO2H acetic acid H2CO3 carbonic acid H3PO4 phosphoric acid HF hydrofluoric acid Dr. S. M. Condren

  24. ACIDS Nonmetal oxides can be acids CO2(aq) + H2O(liq) ---> H2CO3(aq) SO3(aq) + H2O(liq) ---> H2SO4(aq) and can come from burning coal and oil. Dr. S. M. Condren

  25. Neutralization Reactions • base • Any of a large class of compounds, including the hydroxides and oxides of metals, having a bitter taste, a slippery solution, the ability to turn litmus blue, and the ability to react with acids to form salts. • Substance that donates a OH-1 ion to solution Dr. S. M. Condren

  26. BASES Base ---> OH- in water NaOH(aq) Na+(aq) + OH-(aq) NaOH is a strong base Dr. S. M. Condren

  27. Ammonia, NH3An Important Base Dr. S. M. Condren

  28. Neutralization Reactions • salt • The term salt is also applied to substances produced by the reaction of an acid with a base, known as a neutralization reaction. • Salts are characterized by ionic bonds, relatively high melting points, electrical conductivity when melted or when in solution, and a crystalline structure when in the solid state. Dr. S. M. Condren

  29. pH Scale @ -2 -----> @ +16 pH = - log [H3O+] [H3O+] = -antilog pH = 10-pH For log problems, only decimal places are significant, and all decimal places count Dr. S. M. Condren

  30. Example: What is the pH of a 1.0 M solution of HCl? pH = - log [H3O+] [H3O+] = MHCl = 1.0 M • because it is a strong electrolyte pH =-log(1.0) = 0.00 2 significant figures Dr. S. M. Condren

  31. Example: What is the pH of a 1.0 M solution of HCl? pH = - log [H3O+] [H3O+] = MHCl = 1.0 M • because it is a strong electrolyte pH =-log(1.0) = 0.00 2 SF Dr. S. M. Condren

  32. Example: What is the pH of a 0.010 M solution of HCl? pH = - log [H3O+] [H3O+] = MHCl = 0.010 M • because it is a strong electrolyte pH =-log(0.010) = 2.00 2 SF Dr. S. M. Condren

  33. Example: What is the pH of a 2.0 M solution of HCl? pH = - log [H3O+] [H3O+] = MHCl = 2.0 M • because it is a strong electrolyte pH =-log(2.0) = -0.30 2 SF Dr. S. M. Condren

  34. Example: What is the [H3O+] of a solution that has a pH = 2.30? [H3O+] = -antilog pH = 10-pH [H3O+] = 10-pH = 10-2.30 = 5.0x10-3M decimal point pacement 2 SF 2 SF Dr. S. M. Condren

  35. pH Scale Dr. S. M. Condren

  36. Indicators Dr. S. M. Condren

  37. Oxidation-Reduction Reaction Oxidation - loss of electrons Reduction - gain of electrons Redox reaction oxidizing agent - substance that causes oxidation reducing agent - substance that cause reduction Dr. S. M. Condren

  38. Identify the oxidizing agent in the reaction: 2Al(s) + 6 H+ ==> 2 Al3+(aq) + 3 H2(g) Al, H+, Al3+, H2 Identify the oxidizing agent in the reaction: Dr. S. M. Condren

  39. Redox Reactions CuSO4(aq) + Zn(s) Cu(s) + ZnSO4(aq) 4 Al(s) + 3 O2(g) 2 Al2O3(s) 2HgO(s) 2 Hg(l) + O2(g) 2 Al(s) + 3 Br2 (l) Al2Br6 (s) Fe2O3(s) + 2 Al(s) 2 Fe(l) + Al2O3(s) Dr. S. M. Condren

  40. Solution Solutions, in chemistry, homogeneous mixtures of two or more substances. Dr. S. M. Condren

  41. Solute The substance that is present in smallest quantity is said to be dissolved and is called the solute. The solute can be either a gas, a liquid, or a solid. Dr. S. M. Condren

  42. Solvent The substance present in largest quantity usually is called the solvent. The solvent can be either a liquid or a solid. Dr. S. M. Condren

  43. Preparing a Solution Dr. S. M. Condren

  44. Preparing a Solution by Dilution Dr. S. M. Condren

  45. Molarity The number of moles of solute per liter of solution. molarity => M moles of solute M = -------------------- liter of solution units => molar = moles/liter = M Dr. S. M. Condren

  46. Example: A sample of NaNO3 weighing 0.38g is placed in a 50.0-mL volumetric flask. The flask is then filled with water to the mark on the neck, dissolving the solid. What is the molarity of the resulting solution? (0.38gNaNO3) M = (50.0-mL soln) = 0.089 mol/L = 0.089 molar = 0.089 M Dr. S. M. Condren

  47. Stoichiometric Roadmap Grams of A Moles of A Moles of B Grams of B g B x (mol A/g A) Divide by the molar mass mol B x (g B/mol B) Multiply by the molar mass mol A x (mol B/mol A) Multiply by the stoichiometric factor Volume solution A Volume solution B Vol A x (mol A/L A) Multiply by molarity Mol B x (L B/mol B) Diviide by molarity Dr. S. M. Condren

  48. Example: What mass of Na2CO3, in grams, is required for complete reaction with 50.0mL of 0.125 M HNO3? Na2CO3(aq) + 2 HNO3(aq)  2 NaNO`3(aq) + CO2(g) + H2O(l) #g Na2CO3 = (50.0 mL) = 0.331 g Na2CO3 Dr. S. M. Condren

  49. EXAMPLE: Lye, which is sodium hydroxide, can be neutralized by sulfuric acid. How many milliliters of 0.200 M H2SO4 are needed to react completely with 25.0 mL of 0.400 M NaOH? 2 NaOH(aq) + H2SO4(aq) -----> Na2SO4(aq) + 2 H2O (25.0 mL NaOH) #mL H2SO4 = ---------------------- (1 L) -------------- (1000 mL) (0.400 mol NaOH) ------------------------- (1 L NaOH) (1 mol H2SO4) ------------------- (2 mol NaOH) (1000 mL H2SO4) ------------------------ (0.200 mol H2SO4) = 25.0 mL H2SO4 Dr. S. M. Condren

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