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Chapter #7

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Chapter #7

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  1. Chapter #7 Reactions in Aqueous Media

  2. Reaction Driving Forces Five Driving Forces Favor Chemical Change • Formation of a solid • Formation of water • Transfer of electrons • Formation of a gas • Formation of a weak electrolyte

  3. Precipitation Precipitation is the formation of a solid when two solutions are combined.

  4. Types of Aqueous Solutions Solutions are homogeneous mixtures of a solute and a solvent. • The solute is the solution component in the smallest amount while the solvent is the larger component of a solution. • Solutes whose solutions conduct electricity are called electrolytes • Solutes whose solutions do not conduct electricity are called nonelectrolytes • Electrolytes are solutes that form ions when they dissolve. Ionic solutes or acids usually form solutions that conduct electricity.

  5. Strong electrolyte Weak electrolyte Nonelectrolyte Solution Conductivity

  6. Solution Formation Water is one of the best solvents known. It is able to dissolve ionic solutes, such as sodium chloride, to produce solutions that conduct electricity. Molecules, containing a positive and negative regions, are called polar. Water is an example of a polar molecule and can dissolve ionic solutes by the positive region of water attracting to the negative ion of an ionic solute thus separating the crystal lattice in to a solution of solvated ions.

  7. Sodium Chloride Crystal Lattice

  8. H2O K+ The Solvation Process If the attractive forcebetween the surface ion and the solvent is greater than the forces between the ion and the solid then the ion will enter the solution phase. The ion that has left the solid and becomes completedsurrounded by water molecules. It has become solvated or hydrated.

  9. Solvated Ions The process continues as new water moleculesapproach the crystal until the crystal has been fully dissolved. - - Note the different orientation of water molecules around the oppositely charged ions. + - + + - + - + +

  10. Solubility Rules How do we determine if an ionic solute will dissolve in water?

  11. Solubility Rules How do we determine if an ionic solute will dissolve in water? Use the solubility rules

  12. Writing Ionic Equations Ionic compounds that are water soluble are indicated with and (aq) after the symbol, and if insoluble then (s) goes after the symbol. Consider the following equation. AgNO3(aq) + NaCl (aq) What are the products?

  13. Writing Ionic Equations Ionic compounds that are water soluble are indicated with and (aq) after the symbol, and if insoluble then (s) goes after the symbol. Consider the following equation. AgNO3(aq) + NaCl (aq) NaNO3 (aq) + AgCl (s) Called a formula equation Ag+ (aq) + NO3-(aq) + Na+(aq)+ Cl-(aq) Na+(aq) + NO3-(aq) + AgCl (s) Called an ionic equation

  14. Writing Ionic Equations Ionic compounds that are water soluble are indicated with and (aq) after the symbol, and if insoluble then (s) goes after the symbol. Consider the following equation. AgNO3(aq) + NaCl (aq) NaNO3 (aq) + AgCl (s) Called a formula equation Ag+ (aq) + NO3-(aq) + Na+(aq)+ Cl-(aq) Na+(aq) + NO3-(aq) + AgCl (s) Called an ionic equation Spectator Ions

  15. Writing Ionic Equations Ionic compounds that are water soluble are indicated with and (aq) after the symbol, and if insoluble then (s) goes after the symbol. Consider the following equation. AgNO3(aq) + NaCl (aq) NaNO3 (aq) + AgCl (s) Called a formula equation Ag+ (aq) + NO3-(aq) + Na+(aq)+ Cl-(aq) Na+(aq) + NO3-(aq) + AgCl (s) Called an ionic equation Spectator Ions Ag+ (aq) + AgCl (s) Cl-(aq) Called the net ionic equation

  16. Writing Ionic Equations Ionic compounds that are water soluble are indicated with and (aq) after the symbol, and if insoluble then (s) goes after the symbol. Consider the following equation. AgNO3(aq) + NaCl (aq) NaNO3 (aq) + AgCl (s) Called a formula equation Ag+ (aq) + NO3-(aq) + Na+(aq)+ Cl-(aq) Na+(aq) + NO3-(aq) + AgCl (s) Called an ionic equation Spectator Ions Ag+ (aq) + AgCl (s) Cl-(aq) Called the net ionic equation This reaction goes to completion because the solid silver chloride is formed. It is called a precipitate.

  17. Ionic Equations It is possible for all of the reactants and products to be water soluble and thus produce all spectator ions. If this is the case then all of the ions cancel out and there is no net ionic equation. When this occurs we say that there is No Reaction, and give the label NR. This makes sense, since in order for reactions to go to completion a solid, water, gas, or electron transfer must occur in order for a reaction to go to completion.

  18. Acids as Electrolytes Strong acids and bases ionize 100%! Memorized Strong acids and bases: Bases Acids HCl (aq) HI (aq) HBr (aq) HNO3 H2SO4 HClO4 Hydroxides of group I and II metals, except Be and Mg

  19. Types of Chemical Reactions REDOX reactionswhere the oxidation number changes from reactants to products. Oxidation is when the oxidation number increases, by losing of electrons. Reductionis when the oxidation numberdecreases by gaining electrons. Consider the following equation: H2 + O2 H2O What are the oxidation numbers of hydrogen and oxygen?

  20. Types of Chemical Reactions REDOX reactionswhere the oxidation number changes from reactants to products. Oxidation is when the oxidation number increases, by losing of electrons. Reductionis when the oxidation numberdecreases by gaining electrons. Consider the following equation: H2 + O2 H2O What are the oxidation numbers of hydrogen and oxygen? 0 0

  21. REDOX REACTIONS 0 0 2(1+) 2- = 0 H2 + O2 H2O How about hydrogen and oxygen in water?

  22. REDOX REACTIONS 0 0 2(1+) 2- = 0 H2 + O2 H2O How about hydrogen and oxygen in water? Oxidation is caused by the oxygen molecule, so it is referred to as the oxidizing agent (OA) Reduction is caused by the hydrogen molecule, so it is referred to as the reducing agent (RA) reduced oxidized

  23. REDOX REACTIONS Note: • All of the previously discussed reactions from Ch#6 are REDOX except the double replacement reactions. • The number of electrons lost is equal to the number of electrons gained in a reaction. Why? • Most elements have variable oxidation numbers, except for hydrogen, oxygen, and the memorized polyatomic ions.

  24. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H 1+ and oxygen is 2- 1+ 4(2-)=0

  25. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0

  26. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0 5+

  27. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0 5+ 3+

  28. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0 5+ 3+ 1+

  29. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0 5+ 3+ 1+ 0

  30. REDOX REACTIONS Oxidation numbers for a compound must add up to equal zero, while the oxidation numbers for a polyatomic ion must up to equal the charge of that ion. Consider the following chlorine compounds HClO4, HClO3, HClO2, HClO, Cl2, HCl What is the oxidation number of chlorine in each of these compounds, assuming H is 1+ and oxygen is 2- 1+ 7+ 4(2-)=0 5+ 3+ 1+ 0 1-

  31. REDOX REACTIONS 3(2-)=2- How about sulfur in SO3 2-

  32. REDOX REACTIONS 3(2-)=2- How about sulfur in SO3 2-

  33. REDOX REACTIONS 4+ 3(2-)=2- How about sulfur in SO3 2- How about carbon in C6H12O6 12(1+) +6(2-)=0

  34. REDOX REACTIONS 4+ 3(2-)=2- How about sulfur in SO3 2- How about carbon in C6H12O6 0 + 12(1+) +6(2-)=0

  35. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O Balance remaining hydrogen atoms by adding H+ Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  36. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O Balance remaining hydrogen atoms by adding H+ Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ ? 3(2-)=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  37. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O Balance remaining hydrogen atoms by adding H+ Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ 5+ 3(2-)=0 2(?)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  38. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  39. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. ? + 2(1-)=0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  40. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2+ + 2(1-)=0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  41. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2(1-)=0 ? + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  42. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2(1-)=0 2 + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O

  43. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2+ 2(1-)=0 2 + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → Cu(NO3)2 + NO + H2O oxidized reduced

  44. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2+ 2(1-)=0 2 + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + Cu2O → 2 Cu(NO3)2 + NO + H2O oxidized reduced

  45. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2+ 2(1-)=0 2 + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • HNO3 + 3 Cu2O → 3 (2) Cu(NO3)2 + NO + H2O Oxidized3( -2e) Reduced 2(+3)e

  46. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 2+ 2(1-)=0 2 + 2- =0 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 • 2HNO3 + 3 Cu2O → 3 (2) Cu(NO3)2 + 2NO + H2O Oxidized3( -2e) Reduced 2(+3)e

  47. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 2+ 2(1-)=0 2 + 2- =0 • 2HNO3 + 3 Cu2O → 3 (2) Cu(NO3)2 + 2NO + H2O Oxidized3( -2e) Reduced 2(+3)e

  48. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 2+ 2(1-)=0 2 + 2- =0 • 14HNO3 + 3 Cu2O → 3 (2) Cu(NO3)2 + 2NO + 7 H2O Oxidized3( -2e) Reduced 2(+3)e

  49. Balancing Redox Reactions I. Oxidation Number Method a. Assign oxidation numbers to each element b. Determine the elements oxidized and reduced c. Balance the atoms that are oxidized and reduced d. Balance the electrons lost or gained, to conform to the Law of Conservation of Matter, by placing coefficients in front of the formulas containing the atoms oxidized and reduced to both sides of the equation. e. The remaining atoms are balanced by inspection f. Balance oxygen, or hydrogen by adding H2O g. Balance remaining hydrogen atoms by adding H+ h. Simplify i. For basic reactions add the same number of OH- ions to both sides of the equation as there are H+ ions. j. Combine H+ and OH- ions to make water k. Simplify again if necessary. 1+ 5+ 3(2-)=0 2(1+)+ 2-=0 2+ 2(1-)=0 2 + 2- =0 • 14 HNO3 + 3Cu2O → 6 Cu(NO3)2 + 2 NO + 7 H2O Oxidized3( -2e) Reduced 2(+3)e

  50. OX # BALANCING EXAMPLE MnO4 - + Cl- → Mn2+ + Cl2