Chapter Nine Chemical Reactions in Aqueous Solutions

1. Chapter NineChemical Reactions in Aqueous Solutions

2. Section 9.1General Properties of Aqueous Solutions

3. Review • A solution is a homogeneous mixture • Gas example: air • Liquid liquid: salt water • Solid example: brass • Solute: substance being dissolved • Typically lesser in quantity • Solvent: substance doing the dissolving • Typically greater in quantity

4. Types of Solutes: Electrolytes vs. Nonelectrolytes • Electrolyte: substance that when dissolved in water conducts electricity • Sodium Chloride (or table salt) • Has ions in solution (dissociation) • Nonelectrolyte: substance that when dissolved in water does NOT conduct electricity • Sucrose (or sugar) • Does NOT have ions in solution, but molecules

5. Electrolytes vs. Nonelectrolytes

6. Strong vs. Weak Electrolytes • All water-soluble ionic compounds will dissociate completely • Therefore, they are strong electrolytes (i.e. substances that completely dissociate) • There are only 7 molecular compounds that are also considered strong electrolytes • HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

7. Strong vs. Weak Electrolytes • Most molecular compounds are weak electrolytes OR nonelectrolytes • Weak electrolytes produce some ions upon dissolving but exist mostly of molecules that aren’t ionized • Acids are electrolytes (they produce H+ ions) • HCl(g)  H+(aq) + Cl-(aq) • Bases are electrolytes (they produce OH- ions) • NH3(g)  NH4+(aq) + OH-(aq)

8. Strong vs. Weak Electrolytes • For acids/bases that are WEAK, the reaction goes in both directions simultaneously • HC2H3O2(l)  H+(aq) + C2H3O2-(aq) • “” reaction occurs in both directions • Dynamic Chemical Equilibrium • A + B2  AB2

9. Strong Electrolyte, Weak Electrolyte, or Nonelectrolyte???

10. Classify if Strong Electrolyte, Weak Electrolyte, or Nonelectrolyte • Sucrose (C12H22O11) • Fructose (C6H12O6) • Sodium Citrate (Na3C6H5O7) • Potassium Citrate (K3C6H5O7) • Ascorbic Acid (H2C6H6O6)

11. Section 9.2Precipitation Reactions

12. Precipitation Reactions • Reaction where a “precipitate” forms

13. Precipitation Reactions

14. Solubility • Maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature

15. Solubility Rules for Ionic Compounds

16. Molecular, Ionic, & Net Ionic Equations • Pb(NO3)2(aq) + NaI(aq)  • Ionic Equation: Shows equation with ions dissociated • Net Ionic Equation: Shows only what’s involved in the reaction • Removes “Spectator Ions”

17. Group Quiz #1 • For the following reaction, correctly predict the products to write the balanced molecular equation. Then write the ionic equation and the net ionic equation. • Aqueous solutions of Lead Acetate and Calcium Chloride

18. Section 9.3Acid-Base Reactions

19. Acid-Base Models • Arrhenius Model: • Acids produce H+ ions • Bases produce OH- ions • Bronsted Model: • Acids are H+ donors (or proton donors) • Bases are H+ acceptors (or proton acceptors)

20. More about Acids and Bases

21. Acid-Base Neutralization • Reaction between an acid and base • Produce water (most of the time) and a salt (ionic compound)

22. Section 9.4Oxidation-Reduction Reactions

23. Oxidation-Reduction Reactions • A.K.A. “Redox” Reactions • Chemical Reaction where electrons are being transferred from one reactant to another.

24. Example Redox Reaction • Consider Zn(s) + CuCl2(aq)  ZnCl2(aq) + Cu(s)

25. Example Redox Reaction

26. Some definitions • Oxidation is loss of electrons • Reduction is gain of electrons • “OIL RIG” • Oxidizing Agent: species that causes oxidation • Takes the electrons • Reducing Agent: species that causes reduction • Gives the electrons

27. Oxidation Numbers • A.K.A. Oxidation State (or charge) • Help us determine what elements were oxidized and reduced • In order to determine an element’s oxidation number, you must follow the guidelines on the next two slides:

28. Guidelines with Oxidation Numbers

29. Guidelines with Oxidation Numbers

30. Determining Oxidation Numbers • What is the oxidation number of each atom in the following: • SO2 • NaH • CO32- • H2SO4

31. What is Oxidized and What is Reduced? • 2Fe + 6HBr  3H2 + 2FeBr3 • N2 + 3H2  2NH3 • 2KClO3  2KCl + 3O2

32. Group Quiz #2 • What is the oxidation number for chlorine in the compound HClO4? • What species is the reducing agent in the following equation? • Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) • Does the following equation represent a redox reaction? Why? • 2Mg(s) + O2(g)  2MgO(s)

33. Section 9.5Concentrations of Solutions

34. Concentration • Measure of amount of solute dissolved in a certain amount of solvent or solution • More solute: • Concentrated • Less solute: • Diluted

35. Molarity (One type of concentration) • Molarity = moles of solute/ L of solution • A.K.A. molar concentration • Represented by “M” ex: 1.5 M • If you have exactly 1 L of 1.5 M glucose, it contains 1.5 moles of glucose

36. Example • Suppose you wanted to make a 0.150 M solution of KMnO4 using a 25o.00 mL volumetric flask. How would you do this?

37. Preparing molar solutions

38. Group Quiz #3 • You need to make 500. mL of a 0.650 M solution of Sodium Hydroxide (NaOH). What mass of NaOH do you need to use? • What is the molar concentration (M) of a solution prepared by dissolving 58.50 g of Copper Chloride (CuCl2) in water to yield a 1.50 L solution?

39. Dilution • Preparing less concentrated solutions • Typically done by adding water to concentrated solution • Dilution formula: McVc = MdVd • C = concentrated • D = diluted

40. Dilution Examples • What volume in mL of a 1.20 M HCl solution must be diluted in order to prepare 1.00 L of 0.0150 M HCl? • How much water was added?

41. Number of Ions in Solution • Recall: Soluble Ionic Compounds dissociate completely (all ionize) • If you have 0.500 M of KMnO4, then there is 0.500 M of K+ and 0.500 M of MnO4- (1:1 ratio between ions) • [ ] are usually used to show concentration • [KMnO4] = 0.500 M, [K+] = 0.500 M, [MnO4-] = 0.500M

42. Number of Ions in Solution • If you have soluble ionic compounds with ratios other than 1:1 for ions, use subscripts to determine ion concentration • Ex: Na2SO4 • [Na2SO4] = 0.35 M, • [Na+] = 0.70 M, • [SO42-] = 0.35 M Suppose you had a 1.55 L solution of this ionic compound. How many moles of each ion do you have? How many individual ions do you have?

43. Section 9.6Aqueous Reactions and Chemical Analysis

44. Gravimetric Analysis • Analytical technique based on mass • Uses percent composition • Ex: A 0.8633-g sample of an ionic compound containing chloride ions and unknown metal cations is dissolved in water and treated with excess AgNO3. If 1.5615 g of AgCl precipitate, what is the percent by mass of Cl in the original compound?

45. Acid-Base Titrations • Process where • Solution of known concentration (standard solution) is added gradually to • Another solution of unknownconcentrationtill • The reaction is complete • Equivalence point: # of moles of H+ ions equals # of moles of OH- ions • End point: Color change in solution (visually indicates the equivalence point)

46. Acid-Base Titrations

47. Examples • What volume of a 0.203 M NaOH solution is needed to neutralize 25.0 mL of a 0.188 M H2SO4 solution? • If it takes 26.79 mL of 0.560 M HCl solution to neutralize 85.70 mL of Ba(OH)2, what is the molarity of the base?

48. One more example • What is the molar mass of a diprotic acid if 30.5 mL of 0.1112 M NaOH is required to neutralize a 0.1365-g sample?

49. Group Quiz #4 • How many milliliters of a 1.89 M H2SO4 solution are needed to neutralize 91.9 mL of a 0.336 M KOH solution? • Explain the difference between an endpoint and an equivalence point.