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Chapter 5 Molecular View of Reactions in Aqueous Solutions

Chapter 5 Molecular View of Reactions in Aqueous Solutions. Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop. Reactions in Solution. For reaction to occur Reactants needs to come into physical contact Happens best in gas or liquid phase Movement occurs Solution

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Chapter 5 Molecular View of Reactions in Aqueous Solutions

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  1. Chapter 5 Molecular View of Reactions in Aqueous Solutions Chemistry: The Molecular Nature of Matter, 6E Jespersen/Brady/Hyslop

  2. Reactions in Solution • For reaction to occur • Reactants needs to come into physical contact • Happens best in gas or liquid phase • Movement occurs Solution • Homogeneous mixture • 2 or more components mix freely • Molecules or ions completely intermingled • Contains at least 2 substances

  3. Definitions: Solvent • Medium that dissolves solutes • Component present in largest amount • Can be gas, liquid, or solid • Liquids most common • Aqueous solution—water is solvent Solute • Substance dissolved in solvent • Solution is named by solute • Can be gas—CO2 in soda • Liquid—Ethylene glycol in antifreeze • Solid—Sugar in syrup

  4. Iodine Molecules in Ethanol Crystal of solute placed in solvent Solute molecules dispersed throughout solvent

  5. Solutions • May be characterized using Concentration • Solute-to-solvent ratio • Percent Concentration or

  6. Relative Concentration Dilute solution • Small solute to solvent ratio Ex. Eyedrops Concentrated solution • Large solute to solvent ratio Ex. Pickle brine • Dilute solution contains less solute per unit volume than more concentrated solution

  7. Concentration Solubility • Temperature dependent Saturated solution • Solution in which no more solute can be dissolved at a given temperature Unsaturated solution • Solution containing less solute than maximum amount • Able to dissolve more solute

  8. Solubilities of Some Common Substances

  9. Concentrations Supersaturated Solutions • Contains more solute than required for saturation at a given temperature • Formed by careful cooling of saturated solutions • Unstable • Crystallize out when add seed crystal – results in formation of solid or precipitate (ppt.)

  10. Preciptates Precipitate • Solid product formed when reaction carried out in solutions and one product has low solubility • Insoluble product • Separates out of solution Precipitation reaction • Reaction that produces precipitate Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq) • 1 mol Pb(NO3)2 2 mol KI • 0.100 mol Pb(NO3)2 0.200 mol KI

  11. Electrolytes in Aqueous Solution • Ionic compounds conduct electricity • Molecular compounds don’t conduct electricity Why? Bright light No light Ions present Molecular CuSO4 & water Sugar & water

  12. Ionic Compounds (Salts) in Water • H2O molecules arrange themselves around ions & removethem from lattice. Dissociation • Break salts apart into ions when enter solution Separated ions • Hydrated • Conduct electricity • Note: Polyatomic ions remain intact • Ex. KIO3 K+ + IO3 NaCl(s) Na+(aq) + Cl–(aq)

  13. Molecular Compounds In Water • When molecules dissolve in water • Solute particles are surrounded by water • Molecules are not dissociated

  14. Electrical Conductivity Electrolyte • Solutes that yield electrically conducting solutions • Separate into ions when enter into solution Strong electrolyte • Electrolyte that dissociates 100% in water • Yields aqueous solution that conducts electricity • Good electrical conduction • Ionic compounds • Strong acids and bases Ex. NaBr, KNO3, HClO4, HCl

  15. Electrical Conductivity Weak electrolyte • Aqueous solution that weakly conducts electricity due to low ionization • Weak acids and bases Ex. Acetic acid (HC2H3O2), ammonia (NH3) Non-electrolyte • Aqueous solution that doesn’t conduct electricity • Molecules remain intact in solution Ex. Sugar, alcohol

  16. Your Turn How many ions form on the dissociation of Na3PO4? • 1 • 2 • 3 • 4 • 8

  17. Your Turn How many ions form on the dissociation of Al2(SO4)3? • 2 • 3 • 5 • 9 • 14

  18. Equations for Dissociation Reactions • Ionic compound dissolves to form hydrated ions • Hydrated = surrounded by water molecules • In chemical equations, hydrated ions are indicated by • Symbol (aq)after each ions • Ions are written separately KBr(s) K+(aq)+ Br(aq) Mg(HCO3)2(s) Mg2+(aq)+ 2HCO3(aq)

  19. Learning Check Write the equations that illustrate the dissociation of the following salts: • Na3PO4(aq)→ • Al2(SO4)3(aq)→ • CaCl2(aq)→ • Ca(MnO4)2(aq)→ 3 Na+(aq) + PO43(aq) 2 Al3+(aq) + 3 SO42(aq) Ca2+(aq)+ 2 Cl(aq) Ca2+(aq) + 2 MnO4(aq)

  20. Equations of Ionic Reactions • Consider the reaction of Pb(NO3)2 with KI PbI2(s) Pb2+ NO3– K+ I–

  21. Equations of Ionic Reactions • When two soluble ionic solutions are mixed, sometimes an insoluble solid forms. • Three types of equations used to describe • Molecular Equation • Substances listed as complete formulas • Ionic Equation • All soluble substances broken into ions • Net Ionic Equation • Only lists ions that actually take part in reaction

  22. Equations of Ionic Reactions 1. Molecular Equation • Complete formulas for all reactants and products • Formulas written with ions together • Does not indicate presence of ions • Gives identities of all compounds • Good for planning experiments Ex. Pb(NO3)2(aq)+ 2KI(aq) PbI2(s)+ 2KNO3(aq)

  23. Equations of Ionic Reactions 2. Ionic Equation • Emphasizes the reaction between ions • All strong electrolytes dissociate into ions • Used to visualize what is actually occurring in solution • Insoluble solids written together as they don’t dissociate to any appreciable extent Ex. Pb2+(aq)+ 2NO3(aq) + 2K+(aq)+ 2I(aq)  PbI2(s) + 2K+(aq)+ 2NO3(aq)

  24. Equations of Ionic Reactions Spectator Ions • Ions that don’t take part in reaction • They hang around and watch • K+ & NO3 in our example 3. Net Ionic Equation • Eliminate all spectator ions • Emphasizes the actual reaction • Focus on chemical change that occurs Ex. Pb2+(aq) + 2I(aq) PbI2(s)

  25. Net Ionic Equations • Many ways to make PbI2 • Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq) • Pb(C2H3O2)2(aq) + 2NH4I(aq)PbI2(s) + 2NH4C2H3O2(aq) • Different starting reagents • Same net ionic equation • Pb2+(aq) + 2I(aq) PbI2(s)

  26. Converting Molecular Equations to Ionic Equations Strong electrolytes exist as dissociated ions in solution Strategy • Identify strong electrolytes • Use subscript coefficients to determine total number of each type of ion • Separate ions in all strong electrolytes • Show states as recorded in molecular equations

  27. Learning Check: Convert Molecular to Ionic Equations: Write the correct ionic equation for each: Pb(NO3)2(aq)+ 2NH4IO3(aq)→Pb(IO3)2(s)+ 2NH4NO3(aq) 2NaCl(aq)+ Hg2(NO3)2 (aq)→ 2NaNO3(aq)+ Hg2Cl2(s) Pb2+(aq) + 2NO3–(aq) + 2NH4+(aq) + 2IO3–(aq)→ Pb(IO3)2(s) + 2NH4+(aq) + 2NO3–(aq) 2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq)→2Na+(aq) + 2NO3–(aq) + Hg2Cl2(s)

  28. Your Turn Consider the following reaction : Na2SO4(aq) + BaCl2(aq) → 2NaCl(aq) + BaSO4(s) Which is the correct ionic equation? • 2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq)+ 2Cl–(aq) + BaSO4(s) • 2Na+(aq) + SO42–(aq) + Ba2+(aq) + 2Cl–(aq) → 2Na+(aq) + 2Cl–(aq) + BaSO4(s) • 2Na+(aq) + SO42–(aq) + Ba2+(aq) + Cl22–(aq) → 2Na+(aq) + 2Cl–(aq) + Ba2+(s) + SO42–(s) • Ba2+(aq) + SO42–(aq) → BaSO4(s) • Ba2+(aq) + SO42–(aq) → Ba2+(s) + SO42–(s)

  29. Converting Ionic Equations to Net Ionic Equations Strategy • Identify spectator ions • Eliminate from both sides • Rewrite equation using only ions that actually react. • Show states as recorded in molecular and ionic equations

  30. Learning Check: Convert Ionic Equation to Net Ionic Equation Write the correct net ionic equation for each. Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2IO3–(aq)→Pb(IO3)2(s) + 2K+(aq) + 2NO3–(aq) 2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq)→ 2Na+(aq) + 2NO3–(aq) + Hg2Cl2(s) Pb2+(aq) + 2IO3–(aq)→ Pb(IO3)2(s) 2Cl–(aq)+ Hg22+(aq)→Hg2Cl2(s)

  31. Your Turn Consider the following molecular equation: (NH4)2SO4(aq)+ Ba(CH3CO2)2(aq)→ 2NH4CH3CO2(aq)+ BaSO4(s) Which is the correct net ionic equation? • Ba2+(aq) + SO42–(aq) → BaSO4(s) • 2NH4+(aq) + 2CH3CO2–(aq)→ 2NH4CH3CO2(s) • Ba2+(aq) + SO42–(aq) → BaSO4(aq) • 2NH4+(aq) + Ba2+(aq) + SO42–(aq)+2CH3CO2–(aq)→ 2NH4+(aq) + 2CH3CO2–(aq) + BaSO4(s) • 2NH4+(aq) + 2CH3CO2–(aq)→ 2NH4CH3CO2(aq)

  32. Criteria for Balancing Ionic and Net Ionic Equations Material Balance • There must be the same number of atoms of each kind on both sides of the arrow Electrical Balance • The net electrical charge on the left must equal the net electrical charge on the right • Charge does not have to be zero

  33. Learning Check: Balancing Equations for Mass & Charge Balance Molecular Eqn. for mass 2Na3PO4(aq) + 3Pb(NO3)2(aq) 6NaNO3(aq) + Pb3(PO4)2(s) • Can keep polyatomic ions together when counting BalanceIonic Eqn.for charge 6Na+(aq)+ 2PO43(aq)+ 3Pb2+(aq)+ 6NO3(aq) 6Na+(aq)+ 6NO3(aq)+ Pb3(PO4)2(s) • Charge must add up to zero on both sides. Net Ionic Eqn. Balanced for both mass & charge 3Pb2+(aq) + 2PO43(aq) Pb3(PO4)2(s)

  34. Acids & Bases as Electrolytes • Many common laboratory chemicals and household products Indicators • Dye molecules that change color in presence of acids or bases Acids • Turn blue litmus red • Lemon juice, vinegar, H2SO4 Bases • Turn red litmus blue • Drano (lye, NaOH), ammonia (NH3)

  35. Neutralization Reaction • Important reaction of acids and bases • Acid reacts with base to form water and salt (ionic compound). Acid + base  salt + H2O Ex. HCl(aq)+ NaOH(aq)  NaCl(aq)+ H2O HBr(aq)+ LiOH(aq)  LiBr(aq)+ H2O • 1:1 mole ratio of acid:base gives neutralsolution Ionization reactions • Ions form where none have been before • Reactions of acids or bases with water

  36. Arrhenius • Acid-base neutralization is H+(aq)+ OH–(aq) H2O • In solution, H+ attaches itself to H2O to form H3O+ or hydronium ion in water • H+ does not ever exist in aqueous solution • When H3O+ reacts, it releases H+ • H+ is active ingredient • Often use just H+ for simplicity

  37. Arrhenius Acid • Substance that reacts with water to produce the hydronium ion, H3O+ Acid + H2O  Anion + H3O+ HA + H2O  A– + H3O+ HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq) HCl(g) + H2O  Cl–(aq) + H3O+(aq)

  38. Acids Categorized by Number of H+s Monoprotic Acids • Furnish only one H+ HNO3(aq) + H2O  H3O+(aq) + NO3–(aq) HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq) Polyprotic acids • Furnish more than one H+ Diprotic acids — furnish two H+ H2SO3(aq) + H2O H3O+(aq) + HSO3–(aq) HSO3–(aq) + H2O  H3O+(aq) + SO32–(aq)

  39. Acids Catagorized by Number of H+s Polyprotic acids • Triprotic acids — furnish three H+ H3PO4 H2PO4– HPO42– PO43– • Stepwise equations H3PO4(aq) + H2O H3O+(aq) + H2PO4–(aq) H2PO4–(aq)+ H2O  H3O+(aq) + HPO42–(aq) HPO42–(aq)+ H2O  H3O+(aq) + PO43–(aq) Net: H3PO4(aq) + 3H2O  3H3O+(aq) + PO43–(aq) –H+ –H+ –H+

  40. Acidic Anhydrides Nonmetal Oxides • Act as Acids • React with water to form molecular acids that contain hydrogen SO3(g) + H2O  H2SO4(aq) sulfuric acid N2O5(g) + H2O  2HNO3(aq) nitric acid CO2(g) + H2O  H2CO3(aq) carbonic acid

  41. Arrhenius Bases • Ionic compounds that contain hydroxide ion, OH–, or oxide ion, O2–. or • Molecular compounds that react with water to give OH–. 1. Ionic compounds containing OH– or O2– a.Metal Hydroxides • Dissociate into metal & hydroxide ions NaOH(s) Na+(aq) + OH–(aq) Mg(OH)2(s) Mg2+(aq) + 2OH–(aq)

  42. Ionic Oxides b. Basic Anhydrides • Soluble metal oxides • Undergo ionization(hydrolysis)reactionto form hydroxide ions • Oxide reacts with water to form metal hydroxide CaO(s) + H2O  Ca(OH)2(aq) • Then metal hydroxide dissociates in water Ca(OH)2(aq)  Ca2+(aq) + 2OH–(aq) H2O 2OH– O2–

  43. Strong vs. Weak Electrolyte HCl(aq) NH3(aq) CH3COOH(aq)

  44. Strong Acids • Dissociate completely when dissolved in water Ex.HBr(g) + H2O  H3O+(aq) + Br–(aq) • Good electrical conduction • Any acid not on this list, assume weak

  45. Arrhenius Bases 2. Molecular Bases • Undergo ionization (hydrolysis)reactionto form hydroxide ions Base + H2O BaseH+(aq) + OH–(aq) B + H2O BH+(aq) + OH–(aq) NH3(aq) + H2O  NH4+(aq) + OH–(aq)  NH4+ NH3 OH– H2O

  46. Strong Bases • Bases that dissociate completely in water • Soluble metal hydroxides • KOH(aq) K+(aq) + OH–(aq) • Good electrical conductors • Behave as (aq) ionic compounds • Common strong bases are: • Group IA metal hydroxides • LiOH, NaOH, KOH, RbOH, CsOH • Group IIA metal hydroxides • Ca(OH)2, Sr(OH)2, Ba(OH)2

  47. Weak Acids • Any acid other than 7 strong acids • Only ionize partially (<100%) Organic acids HC2H3O2(aq)+ H2OH3O+(aq) + C2H3O2–(aq) Ex. HCO2H(aq) + H2O  H3O+(aq) + HCO2–(aq) Acetic Acid Molecule,HC2H3O2 Acetate ion, C2H3O2– Only this H comes off as H+

  48. Why is Acetic Acid Weak? H2O + C2H3O2–(aq) HC2H3O2(aq)+ H3O+(aq) H3O+(aq) + C2H3O2–(aq)HC2H3O2(aq)+ H2O

  49. Dynamic Equilibrium • 2 opposing reactions occurring at same rate • Also called Chemical equilibrium Equilibrium • Concentrations of substances present in solution do not change with time Dynamic • Both opposing reactions occur continuously • Represented by double arrow HC2H3O2(aq)+ H2O H3O+(aq) + C2H3O2–(aq) Forward reaction – Forms ions Reverse reaction – Removes ions

  50. Weak Bases • Molecular bases • Do not dissociate • Accept H+ from water inefficiently • Accept H+ from acids preferentially NH3(aq) + HCl(aq) NH4Cl(aq) Ex. NH3(aq) + H2O NH4+(aq) + OH(aq) Or for general base B(aq) + H2O  BH+(aq) + OH(aq)

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