1. Prentice Hall © 2003Chapter 4 Chapter 4 Aqueous Reactions and Solution Stoichiometry CHEMISTRY The Central Science 9th Edition David P. White

2. Prentice Hall © 2003Chapter 4 Electrolytic Properties Aqueous solutions, solutions in water, have the potential to conduct electricity. The ability of the solution to conduct depends on the number of ions in solution. There are three types of solutions: Strong electrolytes, Weak electrolytes, and Nonelectrolytes. General Properties of Aqueous Solutions

3. Prentice Hall © 2003Chapter 4 Electrolytic Properties General Properties of Aqueous Solutions

4. Prentice Hall © 2003Chapter 4 Ionic Compounds in Water Ions dissociate in water. In solution, each ion is surrounded by water molecules. Transport of ions through solution causes flow of current. General Properties of Aqueous Solutions

5. Prentice Hall © 2003Chapter 4 Molecular Compounds in Water Molecular compounds in water (e.g., CH 3 OH): no ions are formed. If there are no ions in solution, there is nothing to transport electric charge. General Properties of Aqueous Solutions

6. Prentice Hall © 2003Chapter 4 Strong and Weak Electrolytes Strong electrolytes: completely dissociate in solution. For example: Weak electrolytes: produce a small concentration of ions when they dissolve. These ions exist in equilibrium with the unionized substance. For example: General Properties of Aqueous Solutions

7. Prentice Hall © 2003Chapter 4 When two solutions are mixed and a solid is formed, the solid is called a precipitate. Precipitation Reactions

9. Prentice Hall © 2003Chapter 4 Exchange (Metathesis) Reactions Metathesis reactions involve swapping ions in solution: AX + BY  AY + BX. Metathesis reactions will lead to a change in solution if one of three things occurs: –an insoluble solid is formed (precipitate), –weak or nonelectrolytes are formed, or –an insoluble gas is formed. Precipitation Reactions

10. Prentice Hall © 2003Chapter 4 Ionic Equations Ionic equation: used to highlight reaction between ions. Molecular equation: all species listed as molecules: HCl(aq) + NaOH(aq)  H 2 O(l) + NaCl(aq) Complete ionic equation: lists all ions: H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq)  H 2 O(l) + Na + (aq) + Cl - (aq) Net ionic equation: lists only unique ions: H + (aq) + OH - (aq)  H 2 O(l) Precipitation Reactions

11. Prentice Hall © 2003Chapter 4 Acids Dissociation = pre-formed ions in solid move apart in solution. Ionization = neutral substance forms ions in solution. Acid = substances that ionize to form H + in solution (e.g. HCl, HNO 3, CH 3 CO 2 H, lemon, lime, vitamin C). Acids with one acidic proton are called monoprotic (e.g., HCl). Acids with two acidic protons are called diprotic (e.g., H 2 SO 4 ). Acids with many acidic protons are called polyprotic. Acid-Base Reactions

12. Prentice Hall © 2003Chapter 4 Bases Bases = substances that react with the H + ions formed by acids (e.g. NH 3, Drano™, Milk of Magnesia™). Acid-Base Reactions

13. Prentice Hall © 2003Chapter 4 Strong and Weak Acids and Bases Strong acids and bases are strong electrolytes. –They are completely ionized in solution. Weak acids and bases are weak electrolytes. –They are partially ionized in solution. Acid-Base Reactions

14. Prentice Hall © 2003Chapter 4 Identifying Strong and Weak Electrolytes Water soluble and ionic = strong electrolyte (probably). Water soluble and not ionic, but is a strong acid (or base) = strong electrolyte. Water soluble and not ionic, and is a weak acid or base = weak electrolyte. Otherwise, the compound is probably a nonelectrolyte. Acid-Base Reactions

15. Prentice Hall © 2003Chapter 4 Identifying Strong and Weak Electrolytes Acid-Base Reactions

16. Prentice Hall © 2003Chapter 4 Neutralization Reactions and Salts Neutralization occurs when a solution of an acid and a base are mixed: HCl(aq) + NaOH(aq)  H 2 O(l) + NaCl(aq) Notice we form a salt (NaCl) and water. Salt = ionic compound whose cation comes from a base and anion comes from an acid. Neutralization between acid and metal hydroxide produces water and a salt. Acid-Base Reactions

17. Prentice Hall © 2003Chapter 4 Acid-Base Reactions with Gas Formation Sulfide and carbonate ions can react with H + in a similar way to OH . 2HCl(aq) + Na 2 S(aq)  H 2 S(g) + 2NaCl(aq) 2H + (aq) + S 2- (aq)  H 2 S(g) HCl(aq) + NaHCO 3 (aq)  NaCl(aq) + H 2 O(l) + CO 2 (g) Acid-Base Reactions

18. Prentice Hall © 2003Chapter 4 Oxidation and Reduction When a metal undergoes corrosion it loses electrons to form cations: Ca(s) +2H + (aq)  Ca 2+ (aq) + H 2 (g) Oxidized: atom, molecule, or ion becomes more positively charged. –Oxidation is the loss of electrons. Reduced: atom, molecule, or ion becomes less positively charged. –Reduction is the gain of electrons. Oxidation-Reduction Reactions

19. Prentice Hall © 2003Chapter 4 Oxidation and Reduction Oxidation-Reduction Reactions

20. Oxidation and Reduction Oxidation-Reduction Reactions

21. Prentice Hall © 2003Chapter 4 Oxidation Numbers Oxidation number for an ion: the charge on the ion. Oxidation numbers are assigned by a series of rules: 1.If the atom is in its elemental form, the oxidation number is zero. E.g., Cl 2, H 2, P 4. 2.For a monatomic ion, the charge on the ion is the oxidation state. Oxidation-Reduction Reactions

22. Prentice Hall © 2003Chapter 4 Oxidation Numbers 3.Nonmetal usually have negative oxidation numbers: a)Oxidation number of O is usually –2. The peroxide ion, O 2 2-, has oxygen with an oxidation number of –1. b)Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals. c)The oxidation number of F is –1. 4.The sum of the oxidation numbers for the atom is the charge on the molecule (zero for a neutral molecule). Oxidation-Reduction Reactions

23. Prentice Hall © 2003Chapter 4 Oxidation of Metals by Acids and Salts Metals are oxidized by acids to form salts: Mg(s) +2HCl(aq)  MgCl 2 (aq) + H 2 (g) During the reaction, 2H + (aq) is reduced to H 2 (g). Metals can also be oxidized by other salts: Fe(s) +Ni 2+ (aq)  Fe 2+ (aq) + Ni(s) Notice that the Fe is oxidized to Fe 2+ and the Ni 2+ is reduced to Ni. Oxidation-Reduction Reactions

24. Prentice Hall © 2003Chapter 4 Activity Series Some metals are easily oxidized whereas others are not. Activity series: a list of metals arranged in decreasing ease of oxidation. The higher the metal on the activity series, the more active that metal. Any metal can be oxidized by the ions of elements below it. Oxidation-Reduction Reactions

26. Prentice Hall © 2003Chapter 4 Molarity Solution = solute dissolved in solvent. Solute: present in smallest amount. Water as solvent = aqueous solutions. Change concentration by using different amounts of solute and solvent. Molarity: Moles of solute per liter of solution. If we know: molarity and liters of solution, we can calculate moles (and mass) of solute. Concentrations of Solutions

27. Prentice Hall © 2003Chapter 4 Molarity Concentrations of Solutions

28. Prentice Hall © 2003Chapter 4 Dilution We recognize that the number of moles are the same in dilute and concentrated solutions. So: M dilute V dilute = moles = M concentrated V concentrated Concentrations of Solutions

29. Prentice Hall © 2003Chapter 4 15.2 g of K 2 CO 3 dissolved in 120 ml solution what is M? What are the molar concentrations of each ion in solution?

30. Prentice Hall © 2003Chapter 4 See page 144 Note how a solution is made Needed.250 L of 1.00 M CuSO 4 How would you prepare this?

31. Prentice Hall © 2003Chapter 4 Dilution How many ml of 4.0 M HCl are needed to make 350. ml of.500 M HCl?.350 L.500 mole HCl 1 L 1000 ml 1 L 4.0 mole HCl 1L

32. Prentice Hall © 2003Chapter 4 There are two different types of units: –laboratory units (macroscopic units: measure in lab); –chemical units (microscopic units: relate to moles). Always convert the laboratory units into chemical units first. –Grams are converted to moles using molar mass. –Volume or molarity are converted into moles using M = mol/L. Use the stoichiometric coefficients to move between reactants and product. Solution Stoichiometry and Chemical Analysis

33. Prentice Hall © 2003Chapter 4 Solution Stoichiometry and Chemical Analysis

34. Prentice Hall © 2003Chapter 4 Titrations Suppose we know the molarity of a NaOH solution and we want to find the molarity of an HCl solution. We know: –molarity of NaOH, volume of HCl. What do we want? –Molarity of HCl. What do we do? –Take a known volume of the HCl solution, measure the mL of NaOH required to react completely with the HCl. Solution Stoichiometry and Chemical Analysis

35. Prentice Hall © 2003Chapter 4 Titrations What do we get? –Volume of NaOH. We know molarity of the NaOH, we can calculate moles of NaOH. Next step? –We also know HCl + NaOH  NaCl + H 2 O. Therefore, we know moles of HCl. Can we finish? –Knowing mole (HCl) and volume of HCl (20.0 mL above), we can calculate the molarity. Solution Stoichiometry and Chemical Analysis

36. Titrations

37. Prentice Hall © 2003Chapter 4 53.60 ml of.300 M NaOH is needed to neutralize a 25.00 ml sample of HCl. What is the concentration (M) of HCl?

38. Prentice Hall © 2003Chapter 4 How many grams of Ba(OH) 2 are needed to neutralize 20.0 ml of.25 M HCl?.0200 l.25 mole HCl 1 mole Ba(OH) 2 g 1 l 2 mole HCl 1

39. Prentice Hall © 2003Chapter 4 How many grams of Cl - are in a sample if 25.0 ml of.15 M Pb +2 is required to carry out the reaction? Pb +2 + 2Cl -  PbCl 2 (s).0250 l.15 mole Pb +2 2 mole Cl - 35.5 g 1 l 1 mole Pb +2 1 mole Cl -