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Chapter 4

Chapter 4. Aqueous solutions Types of reactions. mixture. homogeneous. heterogeneous. t wo phases oil/water milk. o ne phase Tab water. solution. Solute(s). solvent. Less abundant or other component(s) of mixture salts in tab water. more abundant component of mixture

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Chapter 4

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  1. Chapter 4 Aqueous solutions Types of reactions

  2. mixture homogeneous heterogeneous two phases oil/water milk one phase Tab water solution Solute(s) solvent • Less abundant or other component(s) of mixture • salts in tab water • more abundant • component of mixture • water in tab water Water always solvent even in 98% H2SO4

  3. solutes Non-electrolytes Electrolytes Solute undergoes Dissociation • Sugar / H2O strong weak complete dissociation partial dissociation • HAc / H2O • NaCl / H2O • HCl / H2O

  4. Hydration of Solid Solute • At edges, fewer oppositely charged ions around • H2O can come in; Ion-dipole forces; Remove ion • New ion at surface • Process continues until all ions in solution • Hydration of ions • Completely surrounded by solvent

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

  6. Electrical conductivity of electrolyte solutions • Ionic compounds • Strong acids and bases Ex.NaBr, KNO3, HClO4, HCl, KOH • Weak acids and bases Ex.Acetic acid (HC2H3O2), ammonia (NH3) Ex.Sugar, alcohol

  7. 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) Ca2+(aq)+ 2 Cl(aq)

  8. Solubility • Maximum amount of a substance that can be dissolved in a given amount of solvent at a given temperature. • Usually g/100 mL. Saturated solution: Solution in which no more solute can be dissolved at a given temperature Unsaturated solution: Solution containing less solute than max. amount; Able to dissolve more solute

  9. Solubilities of Some Common Substances

  10. “Like dissolves Like” • Ethanol (C2H5OH) dissolves in water: • polar↔ polar • Glucose (C6H12O6) and sucrose (C12H22O11) dissolve in water: polar↔ polar • Oil doesn’t dissolve in water: nonpolar↔ polar • Oil dissolves in benzene: nonpolar↔ nonpolar

  11. Salts are polar. insoluble soluble • AgCl • NaCl Water unable to separate Ag+ from Cl- Interaction very strong

  12. Relative Concentration Solute-to-solvent ratio 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

  13. Molarity quantitatively • abbreviated M • 1 M = 1 mol solute / 1 liter solution

  14. Preparing Solution of Known Molarity a b c d e • Weigh solid and transfer to volumetric flask • Add part of the water • Dissolve solute completely • Add water to reach etched line • Stopper flask and invert to mix thoroughly

  15. Concentration of each type of ions in 0.50 M Co(NO3)2(aq)? 1 mol 1 mol 2 mol In 1.00 L 0.50 mol 0.50 mol 1.00 mol Molarity 0.50 M 0.50 M 1.00 M Concentration of each type of ions in 0.50 M Fe(ClO4)3(aq)? Molarity 0.50 M 0.50 M 1.50 M

  16. Moles of Cl- 1.75 L of 1.0×10-3 M ZnCl2(aq)? 1 2 1.75×10-3mol ?

  17. Practice • How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution? • What would the concentration be if you used 27g of CaCl2 to make 500. mL of solution? What is the concentration of each ion? • Describe how to make 1.00 L of a 0.200 M K2CrO4solution. • Describe how to make 250. mL of an 2.0 M copper (II) sulfate dihydrate solution. • Calculate the concentration of a solution made by dissolving 45.6 g of Fe2(SO4)3 to 475 mL. What is the concentration of each ion?

  18. Describe how to make 1.00 L of a 0.200 M K2CrO4solution.

  19. No solid K2CrO4 available in the lab . But 2.00 M K2CrO4 solution is available .

  20. Dilution

  21. Prepare 150 mL of 0.100 M H2SO4 from 16.0 M solution. • What volume of a 1.7 M solution is needed to make 250 mL of a 0.50 M solution? • 18.5 mL of 2.3 M HCl is added to 250 mL of water. What is the concentration of the solution? • You have a 4.0 M stock solution. Describe how to make 1.0 L of a 0.75 M solution.

  22. Types of Chemical Reactions Reduction-Oxidation Metathesis Double Replacement Electron transfer AB + CD  AD + CB Acid-Base Reaction precipitation reaction • a solid is formed from solution precipitate Formation of a weak electrolyte Formation of a gas

  23. Precipitation reactions

  24. Molecular equation Ionic equation Net ionic equation: describes what really happens. Spectator ions: A reaction takes place if it has a net ionic equation

  25. Solubility Rules • All nitrates and acetates are soluble • Salts of alkali metals ions and NH4+ ions are soluble. • Chlorides, bromides and iodides (salts of Cl-, Br- and I-) are soluble except those of Ag+, Pb2+, and Hg22+. • Most sulfates are soluble, except those of Pb2+, Ba2+, Hg2+, and Ca2+. • Most hydroxides are slightly soluble (insoluble) except those of alkali metals (Ba(OH)2, Sr(OH)2 and Ca(OH)2 are marginally soluble). • Sulfides (S2-), carbonates (CO32-), chromates (CrO42-) and phosphates (PO43-), are insoluble except those of alkali metals and NH4+.

  26. Does the following mixing process involve a chemical reaction?

  27. Precipitation reactions • NaOH(aq) + FeCl3(aq)® ?? • NaOH(aq) + FeCl3(aq)® NaCl + Fe(OH)3 • NaOH(aq) + FeCl3(aq)®NaCl(aq) + Fe(OH)3(s) • Na+(aq)+OH-(aq) + Fe3+ (aq) + Cl-(aq) ® Na+(aq) + Cl-(aq) +Fe(OH)3(s) • OH-(aq) + Fe3+ (aq)® Fe(OH)3(s)

  28. Precipitation reactions • BaCl2(aq) + KNO3(aq)® ?? • BaCl2(aq) + KNO3(aq)® KCl + Ba(NO3)2 • BaCl2(aq) + KNO3(aq)® KCl(aq) + Ba(NO3)2(aq) • Ba2+(aq)+2 Cl-(aq) + K+ (aq) + NO3-(aq) ® K+(aq) + Cl-(aq) +Ba2+(aq)+ 2 NO3-(aq) • No net ionic equation • No reaction

  29. Practice iron (III) sulfate and potassium sulfide Lead (II) nitrate and sulfuric acid. solutions of NaOH and NiCl2 are mixed.

  30. 1 1 0.15 mol ?

  31. 1.25 L of 0.0500 M Pb(NO3)2 mixed with 2.0 L of 0.0250 M Na2SO4. Calculate the mass of precipitate. 1 1 0.0625 mol ? 1 1 0.0500 mol ?

  32. Stoichiometry of Precipitation • What mass of solid is formed when 100.00 mL of 0.100 M Barium chloride is mixed with 100.00 mL of 0.100 M sodium hydroxide? • What volume of 0.204 M HCl is needed to precipitate the silver from 50.0 ml of 0.0500 M silver nitrate solution ? • 25 mL 0.67 M of H2SO4 is added to 35 mL of 0.40 M CaCl2 . What mass CaSO4 Is formed?

  33. Arrhenius Acid  Cl–(aq) + H3O+(aq) HCl(g) + H2O Bronsted-Lowry Acid: H+ donor Substance thatreacts with water to produce the hydronium ion, H3O+ Acid + H2O  Anion + H3O+ HA + H2O  A– + H3O+ HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq)

  34. Arrhenius Bases Bronsted Base: H+ acceptor • Substance that reacts with water to give OH–. • Metal Hydroxides NaOH(s)  Na+(aq) + OH–(aq) Mg(OH)2(s)  Mg2+(aq) + 2OH–(aq) • Basic Anhydrides CaO(s) + H2O  Ca(OH)2(aq) Ca(OH)2(aq) Ca2+(aq) + 2OH–(aq) c. Molecular bases: NH3(aq)+H2O  NH4+(aq)+ OH-(aq)

  35. Acid-Base Reactions Ionic equation Net ionic equation: Weak electrolyte: H2O + H2O  H3O+(aq)+ OH-(aq) Any strong acid + strong base

  36. Formation of Weak electrolyte: HAc + H2O  H3O+(aq)+ Ac-(aq) • Acid - Base Reactions are often called neutralizationreaction Because the acid neutralizes the base.

  37. Volume of 0.100 M HCl needed to neutralize 25.0 mL of 0.350 M NaOH ? 1 1 ? 8.75×10-3

  38. 28.0 mL of 0.250 M HNO3 mixed with 53.0 mL of 0.320 M KOH; Amount of water formed Concentrations of H+ and OH- at the end of rct 1 1 7.0 mmol ?  1 1 17.0 mmol ?

  39. HNO3 is Limiting reactant: reacts completely No HNO3 left HNO3→ H+ + NO3- No H+ at the end of reaction How much remains from KOH? KOH → K+ + OH- 10 → 10 mmol

  40. Volumetric analysis: Titration Controlled addition of 1 reactant to another until rxn is complete. Acid-Base Titration: Very common type of titration Ex. Analysis of citric acid in orange juice by neutralization with NaOH

  41. An indicator is needed: organic substance that changes color according to solution acidity • Where the indicator changes color is the endpoint. • Endpoint must be very close to the equivalence point. Acid (Base) added equivalent to base (acid) present Phenolphthalein Acidic Basic

  42. Standardization of NaOH solution • Know the exact concentration! • Its weight is inaccurate . • NaOH is hygroscopic and it absorbs CO2. • Cannot be used to prepare solutions with exactly known M. • Not a primary standard. • KHP is a primary standard: high purity, no weighing problems, • Potassium hydrogen phthalate: KHC8H4O4. • Monoprotic! 41.2 mL of NaOH solution is needed to react exactly with 1.300 g of KHP (M=204.22 g/mol). MNaOH=?

  43. practice • 75 mL of 0.25M HCl is mixed with 225 mL of 0.055 M Ba(OH)2 . What is the concentration of the excess H+ or OH-? • A 50.00 mL sample of aqueous Ca(OH)2 requires 34.66 mL of 0.0980 M Nitric acid for neutralization. What is [Ca(OH)2 ]?

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