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Chapter 4

Chapter 4. You must become familiar with: 1.) the precipitation diagram (fig. 4.3) before you can write precipitation reactions. Chapter 4. 2.) The charges of the transition metal ions in solution (figure 4.2). 3.) Strong acids and bases (figure 4.1) if you are to write acid-base equations.

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Chapter 4

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  1. Chapter 4 You must become familiar with: 1.) the precipitation diagram (fig. 4.3) before you can write precipitation reactions.

  2. Chapter 4 • 2.) The charges of the transition metal ions in solution (figure 4.2). • 3.) Strong acids and bases (figure 4.1) if you are to write acid-base equations.

  3. Chapter 4 • Only molecular weak acids and bases are considered in this chapter; acidic and basic ions are covered in Chapter 13.

  4. Chapter 4 • The half-equation method is the one we will use for balancing redox equations. Advantage: trains you to break down reactions into reduction and oxidation; useful in Ch. 18

  5. Chapter 4 • Note: (even though it may be taught in Chem 1) it isn’t a good idea to construct net equations from ‘molecular equation’. It is a waste of time and isn’t what really happens in solution.

  6. Solutions • Dissociation: • Note that ionic compounds (at least to some extent) break up in water to produce cations and anions.

  7. Solutions • Ex: NaOH(s) • Ex: K2CrO4(s) • (Ionization Equations)

  8. Solutions • Solution Stoichiometry: • Molarity- mols of solute per liter of solution • Ex. 1: What volume of 12 M HCl must be taken to obtain 0.10 mol of HCl?

  9. Solutions • Ex 2: What mass of NaOH is contained in 125 ml of 6.00 M NaOH?

  10. Solutions • Ex 3: What are the molarities of the aluminum and sulfate ions in 0.100 M aluminum sulfate? (use dissociation equation)

  11. Precip Reactions • Solubility Rules (fig 4.3) - used in predicting the results of precipitation reactions.

  12. Precip Reactions • Ex. 1. Mix solutions of barium nitrate and sodium carbonate. What happens? What are the ions present? Possible precipitates?

  13. Precip Reactions • Ex. 1. Mix solutions of barium chloride and sodium hydroxide. What happens? What are the ions present? Possible precipitates?

  14. Acid/Base Reactions • Arrhenius acid and base definitions. • Strong vs weak acids and bases. • Ex. HCl vs HF, Calcium hydroxide vs ammonia. (draw the dissociation / ionization equations)

  15. Net Ionic Equations • Some ions in a neutralization reaction are considered spectator ions. That is, they are unchanged after the reaction.

  16. Net Ionic Equations • Chemists often write net ionic equations to show only those ions that actually take place in the reaction. Ex. Cons. Of mass lab

  17. Net Ionic Equations • The binary acids HCl, HBr, and HI are strong acids, the rest are weak. • Ternary acids with 2 or more oxygens versus hydrogens are usually strong.

  18. Net Ionic Equations • Organic acids are weak. • Polyprotic acids: second step always results in a weak acid (see next slide) • Group 1 and 2 bases are strong (usually).

  19. Polyprotic Acids • An acid that is capable of producing more than one mole of H+. • Dissociation equations for sulfuric acid and phosphoric acid.

  20. Net Ionic Equations • Molecules and weak acids and bases are not written in ionic form. • Soluble salts are written in ionic form, oxides and gases are written as molecules.

  21. Acid/Base Reactions • Strong acid and strong base (ex. Hydrochloric acid and sodium hydroxide) • Strong base and weak acid (ex. Hydrofluoric acid and calcium hydroxide)

  22. Acid/Base Reactions • Strong acid and weak base (ex. Hydrochloric acid and ammonia)

  23. Titrations • Suppose 22.0ml of 0.150 M HCl is reacted with 32.0ml of Ba(OH)2. What is the molarity of the barium hydroxide?

  24. Titrations It is found that 25.00ml of 0.0800 M Ca(OH)2 is required to react with 10.00ml of HCl. What is the molarity of the HCl? OH- + H+ --> H2O [HCl] = .400 M

  25. Solutions • Ex 4: What volume of 0.200 M copper (II) sulfate solution is required to react with 50.0ml of 0.100 M NaOH? (use net ionic)

  26. Redox Reactions • Oxidation numbers: “pseudocharge” assigned according to arbitrary rules; accounts for the movement of electrons.

  27. Redox Reactions • Rules: • 1. Elements in the uncombined form are assigned a number of zero.

  28. Redox Reactions • 2. Monatomic ions equal their charge number. • 3. Group one elements in a compound = +1, group 2 = +2, fluorine = -1, hydrogen normally = +1, oxygen is normally -2.

  29. Redox Reactions • 4. The sum of the oxidation numbers in a molecule = 0. In polyatomic ions, it equals the charge number.

  30. Redox Reactions • Ex. What is the oxidation number for sulfur in sulfuric acid? Chromium in the dichromate ion?

  31. Redox Reactions • Oxidation = an increase in oxidation number. • Reduction = a decrease in oxidation number. • Ex. Copper and zinc chloride (1/2 reactions)

  32. Redox Reactions • The oxidizing agent is reduced. The reducing agent is oxidized. No means yes. Yes means no.

  33. Redox Reactions • Balancing Redox Equations: • Ex. In solution the chlorate ion reacts with the iodide ion to produce the chloride ion and solid iodine.

  34. Redox Reactions • 1. Split into two half reactions (oxidation and reduction).

  35. Redox Reactions • 2. Balance half-equations separately in the following order • A. Balance atoms of element being oxidized or reduced.

  36. Redox Reactions • B. Balance oxidation numbers by adding electrons. • C. Balance charge by adding acid (H+) or base (OH-). (depends)

  37. Redox Reactions • D. Balance oxygen by adding water molecules. • E. Combine the half reactions so that the electrons cancel.

  38. Solutions • Balance Ex 5: I-(aq) + ClO3-(aq) --> I2(s) + Cl-(aq)

  39. Solutions • Ex 5: 6I-(aq) + ClO3-(aq) + 6H+(aq) --> 3 I2(s) + Cl-(aq) + 3 H2O

  40. Solutions • Suppose 22.0ml of 0.150 M potassium chlorate is required to react with a sample weighing 5.00g. What is the percent of the iodide ion in the sample?

  41. Solutions • Stoich Practice #1 Fe2+(aq) + Cr2O72-(aq) --> Fe2+(aq) + Cr3+(aq)

  42. Solutions • An iron ore sample with a mass of 0.9132g is dissolved in HCl(aq). The iron obtained is Fe2+. This solution is titrated with 28.72ml of 0.05051 M K2Cr2O7. What is the %Fe in the original ore?

  43. Solutions • Problem #2: • C2O42-(aq) + MnO4- --> Mn2+(aq) + CO2

  44. Solutions • 50.0ml of a saturated sodium oxalate solution requires 25.8ml of 0.02140 M KMnO4 in acid solution. What mass of sodium oxalate would be present in 1 L of this saturated solution?

  45. Solutions • Practice Problem #3: • As2O3(s) + MnO4-(aq) --> H3AsO4(aq) + Mn2+(aq)

  46. Solutions • A KMnO4 solution is standardized by titration with solid As2O3. A 0.1078g sample of As2O3 requires 22.15ml of KMnO4 to complete the titration. What is the molarity of the KMnO4 solution?

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