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Reactions in Aqueous Solution

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  1. Reactions in Aqueous Solution Chapter 5 Chapter 5

  2. Compounds in Aqueous Solution HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) This is a reaction in which reactants are in solution Solution – homogeneous mixture composed of two parts: solute – the medium which is dissolved solvent – the medium which dissolves the solute. Chapter 5

  3. Compounds in Aqueous Solution Compounds in Water Some compounds conduct electricity when dissolved in water – electrolytes Those compounds which do not conduct electricity when dissolved in water are called – nonelectrolytes Chapter 5

  4. Compounds in Aqueous Solution Ionic Compounds in Water (Electrolytes) The conductivity of the solution is due to the formation of ions when the compound dissolves in water These ions are not the result of a chemical reaction, they are the result of a dissociation of the molecule into ions that compose the solid. Chapter 5

  5. Compounds in Aqueous Solution Ionic Compounds in Water Chapter 5

  6. Compounds in Aqueous Solution Molecular Compounds in Water(Nonelectrolytes) In this case no ions are formed, the molecules just disperse throughout the solvent. Chapter 5

  7. Compounds in Aqueous Solution Molecular Compounds in Water(Nonelectrolytes) There are exceptions to this, some molecules are strongly attracted to water and will react with it. Chapter 5

  8. Compounds in Aqueous Solution Strong and Weak Electrolytes Strong electrolytes – A substance which completely ionizes in water. For example: Chapter 5

  9. Compounds in Aqueous Solution Strong and Weak Electrolytes Weak electrolyte: A substance which partially ionizes when dissolved in water. For example: Chapter 5

  10. Compounds in Aqueous Solution Strong and Weak Electrolytes Notice that the arrow in this reaction has two heads, this indicates that two opposing reactions are occurring simultaneously. Chapter 5

  11. Compounds in Aqueous Solution Strong and Weak Electrolytes Since both reactions occur at the same time, this is called a chemical equilibrium. Chapter 5

  12. Precipitation Reaction Chapter 5

  13. Precipitation Reaction • A reaction which forms a solid (precipitate) • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • AgCl is classified as an insoluble substance Chapter 5

  14. Precipitation Reaction • Net Ionic Equation • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • AgNO3 and NaNO3 are electrolytes in solution so they actually occur as free ions. • Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq) Chapter 5

  15. Precipitation Reaction • Net Ionic Equation • Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq) • Notice that NO3-(aq) and Na+(aq) occur in both the left and right side of the equation. • These are called spectator ions. Chapter 5

  16. Precipitation Reaction • Net Ionic Equation • Ag+(aq) + Cl-(aq) AgCl(s) • With the spectator ions removed, the resulting equation shows only the ions involved in the reaction remain. • This is a net ionic equation. Chapter 5

  17. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds • Most nitrates (NO3-) and acetates (CH3CO2-) are soluble in water. • All chlorides are soluble except: Hg+, Ag+, Pb2+, Cu+ • All sulfates are soluble except: Sr2+, Ba2+, Pb2+ • Carbonates (CO32-), Phosphates (PO43-), Borates (BO33-),Arsenates (AsO43-), and Arsenites (AsO33-) are insoluble. • Hydroxides (OH-) of group Ia and Ba2+ and Sr2+ are soluble. • Most sulfides (S2-) are insoluble. Chapter 5

  18. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds Predict the solubility of the following compounds: PbSO4 AgCH3CO2 (NH4)3PO4 KClO4 Chapter 5

  19. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds Predict the solubility of the following compounds: PbSO4 Insoluble AgCH3CO2 (NH4)3PO4 KClO4 Chapter 5

  20. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds Predict the solubility of the following compounds: PbSO4 Insoluble AgCH3CO2 Soluble (NH4)3PO4 KClO4 Chapter 5

  21. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds Predict the solubility of the following compounds: PbSO4 Insoluble AgCH3CO2 Soluble (NH4)3PO4 Soluble KClO4 Chapter 5

  22. Compounds in Aqueous Solution Solubility Guidelines for Ionic Compounds Predict the solubility of the following compounds: PbSO4 Insoluble AgCH3CO2 Soluble (NH4)3PO4 Soluble KClO4 Soluble Chapter 5

  23. Acids and Bases Acid - substance which ionizes to form hydrogen cations (H+) in solution Examples: Hydrochloric Acid HCl Nitric Acid HNO3 Acetic Acid CH3CO2H Sulfuric Acid H2SO4 Sulfuric acid can provide two H+’s - Diprotic acid, The other acids can provide only one H+ - Monoprotic acid. Chapter 5

  24. Acids and Bases Diprotic acid H2SO4 H+ + HSO4- HSO4-Û H+ + SO42- Chapter 5

  25. Acids and Bases Base - substance which reacts with H+ ions. Examples: ammonia NH3 sodium hydroxide NaOH Chapter 5

  26. Acids and Bases • Acid-Base Reaction • H+ + OH- H2O • It is clear that the metal hydroxides (NaOH for example) provide OH- by disassociation. • Bases like ammonia make OH- by reacting with water (ionization) • NH3 + H2O Û NH4+ + OH- Chapter 5

  27. Acids and Bases Strong and Weak Acids and Bases Strong acids and bases are strong electrolytes. Weak acids and bases are weak electrolytes. Chapter 5

  28. Acids and Bases • Strong Acids • The strength of acids and bases are concerned with the ionization (or dissociation) of the substance, not its chemical reactivity. • Example: • Hydrofluoric acid (HF) is a weak acid, but it is very chemically reactive. • this substance can’t be stored in glass bottles because it reacts with glass (silicon dioxide). Chapter 5

  29. Acids and Bases Common Strong Acids and Bases Common Strong Acids Hydrochloric Acid HCl Hydrobromic Acid HBr Hydroiodic acid HI Nitric Acid HNO3 Perchloric Acid HClO4 Sulfuric Acid H2SO4 Chapter 5

  30. Acids and Bases Common Strong Acids and Bases Common Strong Bases Lithium Hydroxide LiOH Sodium Hydroxide NaOH Potassium Hydroxide KOH Chapter 5

  31. Acids, Bases, and Salts • Neutralization Reaction • Reaction between an acid and a base. • HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) • “The neutralization between acid and metal hydroxide produces water and a salt” • Salt – an ionic compound whose cation comes from a base and anion from an acid. Chapter 5

  32. Acids, Bases, and Salts Neutralization Reaction HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) - despite the appearance of the equation, the reaction actually takes place between the ions. Chapter 5

  33. Acids, Bases, and Salts Neutralization Reaction HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Total Ionic Equation H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O(l) + Na+(aq) + Cl-(aq) Chapter 5

  34. Acids, Bases, and Salts Neutralization Reaction HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Total Ionic Equation H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O(l) + Na+(aq) + Cl-(aq) Chapter 5

  35. Acids, Bases, and Salts Neutralization Reaction HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq) Total Ionic Equation H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O(l) + Na+(aq) + Cl-(aq) Net Ionic Equation H+(aq) + OH-(aq)  H2O(l) Chapter 5

  36. Gas-Forming Reactions Metal Carbonates and Acid • 2 HCl(aq) + Na2CO3(aq)  2 NaCl(aq) + H2O(l) + CO2(g) • Metal carbonates (or bicarbonates) always form a salt, water and carbon dioxide Chapter 5

  37. Gas-Forming Reactions Metal Sulfide and Acid • 2 HCl(aq) + Na2S(s)  H2S(g) + 2 NaCl(aq) • Metal sulfides form a salt and hydrogen sulfide. Chapter 5

  38. Gas-Forming Reactions Metal Sulfite and Acid • 2 HCl(aq) + Na2SO3(s)  SO2(g) + 2 NaCl(aq) + H2O(l) • Metal sulfites form a salt, sulfur dioxide and water. Chapter 5

  39. Gas-Forming Reactions Ammonium Salt and Strong Base • NH4Cl(s) + NaOH(aq)  NH3(g) + NaCl(aq) + H2O(l) • This reaction forms ammonia, salt and water Chapter 5

  40. Oxidation-Reduction Reactions • Reaction where electrons are exchanged. 2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g) Na(s)  Na+(aq) + 1 e- oxidation – loss of electrons 2 H+(g) + 2 e-  H2(g) reduction – gain of electrons Chapter 5

  41. Oxidation-Reduction Reactions • Reaction where electrons are exchanged. 2 Na(s) + 2 H2O(l)  2 NaOH(aq) + H2(g) • An alternate approach is to describe how one reagent effects another. • Reducing Agent, a substance that causes another substance to be reduced. Na(s)  Na+(aq) + 1 e- • Oxidizing Agent, a substance that causes another substnace to be oxidized. 2 H+(g) + 2 e-  H2(g) Chapter 5

  42. Oxidation-Reduction Reactions • Oxidation Numbers • Each atom of a pure element has an oxidation number of zero(0). • For monatomic ions, the oxidation number equals the charge on the ion. • Fluorine always has an oxidation state of -1 in compounds. • Cl, Br, and I always have oxidation numbers of -1, except when combined with oxygen or fluorine. • The oxidation number of H is +1 and O is -2 in most compounds. • The sum of the oxidation numbers must equal the charge on the molecule or ion. Chapter 5

  43. Oxidation-Reduction Reactions Oxidation Numbers Examples PCl5 P 1( ) = Cl 5( ) = ________ 0 Chapter 5

  44. Oxidation-Reduction Reactions Oxidation Numbers Example PCl5 P 1( ? ) = ? Cl 5(-1) = __-5____ 0 Chapter 5

  45. Oxidation-Reduction Reactions Oxidation Numbers Example PCl5 P 1(+5) = +5 Cl 5(-1) = __-5____ 0 Chapter 5

  46. Oxidation-Reduction Reactions Oxidation Numbers Example CO32- C 1( ) = O 3( ) = _____ -2 Chapter 5

  47. Oxidation-Reduction Reactions Oxidation Numbers Example CO32- C 1(?) = ? O 3(-2) = __-6__ -2 Chapter 5

  48. Oxidation-Reduction Reactions Oxidation Numbers Example CO32- C 1(+4) = +4 O 3(-2) = __-6__ -2 Chapter 5

  49. Oxidation-Reduction Reactions Oxidation Numbers Example K2CrO4 K 2( ) = O 4( ) = Cr 1( ) = ________ 0 Chapter 5

  50. Oxidation-Reduction Reactions Oxidation Numbers Example K2CrO4 K 2(+1) = +2 O 4(-2) = -8 Cr 1( ? ) = ___?____ 0 Chapter 5