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AP Chemistry

AP Chemistry. Ch 4: Types of Chemical Reactions and Solution Stoichiometry. Three types of reactions in aqueous solutions:. 1) Precipitation Reactions 2) Acid—Base Reactions 3) Oxidation—Reduction Reactions. Properties of Water . List as many as you can….

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AP Chemistry

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  1. AP Chemistry Ch 4: Types of Chemical Reactions and Solution Stoichiometry

  2. Three types of reactions in aqueous solutions: 1) Precipitation Reactions 2) Acid—Base Reactions 3) Oxidation—Reduction Reactions

  3. Properties of Water List as many as you can…

  4. Electron distribution in molecules of H2 and H2O Figure 4.2

  5. Demo—Conductivity Apparatus: Predict if the following substances are: SE, WE, or NE (assume all are 1.0 M soln’s) SE NE WE SE ___ NaCl (aq) ___ C12H22O11 (aq) ___ C2H5OH ___ NaOH ___ NH4OH (aq) ___ HCl ___ CH3COOH ___ CaCl2 (acetic acid) WE SE WE SE

  6. Strong & Weak Electrolytes NaCl (s) HF (s) HF F+ Na+ Cl- HF HF Na+ H+ Cl- Na+ + Cl- only ions are present H+ + F- mostly HF present; only some ions STRONG ELECTROLYTES Ionize 100% NaCl WEAK ELECTROLYTES Only partially ionize HF

  7. Figure 4.1 The electrical conductivity of ionic solutions

  8. PROPERTIES OF WATER Electrolytes Substances that conduct electricity when dissolved in H2O Strong Electrolytes Substances that completely ionize when dissolved in H2O Equation: NaCl (s)  Na+ (aq) + Cl- (aq)

  9. The dissolution of an ionic compound Figure 4.3

  10. Equation: CH3COOH H+ + CH3COO- 5 % PROPERTIES OF WATER con’t Weak Electrolytes Substances that do not completely ionize when dissolved in water (weak acids & bases)

  11. PROPERTIES OF WATER con’t Nonelectrolytes Substances that do not ionize when dissolved in H2O Equation: C12H22O11 (s) C12H22O11 (aq)

  12. In Conclusion Strong Electrolytes are: 1) soluble salts—see solubility rules 2) strong acids—produce H+ ions HBr, HCl, HI, HNO3, HClO4, H2SO4 HiBer, HiCkel, HI, HoNO, HiCkle-O, H-2-S-O-4 3) strong bases—produce OH- ions LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2,Ba(OH)2

  13. Table 4.2 Selected Acids and Bases Acids Bases Strong Strong hydrochloric acid, HCl sodium hydroxide, NaOH hydrobromic acid, HBr potassium hydroxide, KOH hydroiodic acid, HI calcium hydroxide, Ca(OH)2 nitric acid, HNO3 strontium hydroxide, Sr(OH)2 perchloric acid, HClO4 barium hydroxide, Ba(OH)2 sulfuric acid, H2SO4 Weak Weak hydrofluoric acid, HF ammonia, NH3 phosphoric acid, H3PO4 acetic acid, CH3COOH (or HC2H3O2)

  14. Sample Problem 4.1 Determining Moles of Ions in Aqueous Ionic Solutions PROBLEM: How many moles of each ion are in the following solutions? (b) 78.5g of cesium bromide dissolved in water (d) 35mL of 0.84M zinc chloride

  15. Figure 4.5 A precipitation reaction and its equation ©

  16. Soluble Ionic Compounds Insoluble Ionic Compounds Table 4.1 Solubility Rules For Ionic Compounds in Water 1. All common compounds of Group 1A(1) ions (Li+, Na+, K+, etc.) and ammonium ion (NH4+) are soluble. 2. All common nitrates (NO3-), acetates (CH3COO- or C2H3O2-) and most perchlorates (ClO4-) are soluble. 3. All common chlorides (Cl-), bromides (Br-) and iodides (I-) are soluble, except those of Ag+, Pb2+, Cu+, and Hg22+. 1. All common metal hydroxides are insoluble, except those of Group 1A(1) and the larger members of Group 2A(2)(beginning with Ca2+). 2. All common carbonates (CO32-) and phosphates (PO43-) are insoluble, except those of Group 1A(1) and NH4+. 3. All common sulfides are insoluble except those of Group 1A(1), Group 2A(2) and NH4+. Solubility Song…

  17. Pb2+ I- Na+ PRECIPITATION REACTIONS NaI (aq) + Pb(NO3)2 (aq)  ?

  18. NaI (aq) + Pb(NO3)2 (aq)  ? Precipitation Reactions con’t Molecular Equation: 2 NaI (aq) + Pb(NO3)2 (aq) 2 NaNO3 (aq) + PbI2 (s) Complete Ionic Equation: 2Na+ (aq) + 2I- (aq) + Pb2+ (aq) + 2 NO3- (aq)  2 Na+(aq) + 2 NO3-(aq) + PbI2 (s) Net Ionic Equation: Pb2+ (aq) + 2 I- (aq)  PbI2 (s) Spectator Ions: Na+ and NO3-

  19. Stoichiometry of Precipitation Reactions: CuSO4 (aq) + 2 NaOH (aq)Cu(OH)2 (s) + Na2SO4 (aq) Calculate the mass of solid formed when 1.25 L of 0.0500 M CuSO4 and 2.00 L of 0.0250 M NaOH are mixed.

  20. QUIZ Write the molecular, complete ionic & net ionic equations for the reaction between solutions of barium nitrate with ammonium sulfate. Ba(NO3)2 (aq) + (NH4)2SO4 (aq)  BaSO4 (s) + 2 NH4NO3 (aq) Ba2+(aq) + 2 NO3-(aq) + 2NH4+(aq) + SO42-(aq) BaSO4(s) + 2NH4+(aq) + 2NO3-(aq) Ba2+ (aq) + SO42- (aq)  BaSO4 (s) Spectator ions: NH4+ and NO3-

  21. ACID BASE REACTIONS Neutralization Reactions Also known as… Acid: produces H+ ions proton (H+) donor produces OH- ions proton acceptor Base:

  22. Acids (con’t) dissociate to form H+ ions (protons) Hydrogen AtomHydrogen Ion(proton) e- +

  23. Hydronium Ion Acids and the Solvated Proton HCl + H2O  H+ combines with H2O to form H3O+ H+(proton) is attracted to lone e- pairs in H2O H3O+ + Cl- + - +

  24. Figure 4.4 The hydrated proton

  25. Acid Burns (H2SO4)

  26. Sample Problem 4.4 Writing Ionic Equations for Acid-Base Reactions PROBLEM: Write balanced molecular, total ionic, and net ionic equations for each of the following acid-base reactions and identify the spectator ions. (a) strontium hydroxide(aq) + perchloric acid(aq) (b) barium hydroxide(aq) + sulfuric acid(aq)

  27. What is a Titration? How is one performed? Titration One solution of known concentration is used to determine the concentration of another solution through a monitored reaction Stock solution / Standardized Solution The solution in which the concentration is known Equivalence point/stoichiometric point/end point Moles of H+ = Moles of OH-

  28. Titrations con’t Indicator substance added at the beginning of the titration that changes at color near the equivalence point

  29. Acidic—before neutralization Standardized solution (NaOH) Neutral— ”End Point” Basic— “overshot endpoint”

  30. Sample Problem 4.5 Finding the Concentration of Acid from an Acid-Base Titration PROBLEM: You perform an acid-base titration to standardize an HCl solution by placing 50.00mL of HCl in a flask with a few drops of indicator solution. You put 0.1524M NaOH into the buret, and the initial reading is 0.55mL. At the end point, the buret reading is 33.87mL. What is the concentration of the HCl solution?

  31. Figure 4.8 An aqueous strong acid-strong base reaction on the atomic scale.

  32. Demo: Conductivity Puzzle Demo: When Does 1 + 1 ≠ 2

  33. Proton Transfer Hydronium Ion H3O+ + Cl- + - + Brønsted Lowry Acid: donates a proton Brønsted Lowry Base: accepts a proton HCl + H2O 

  34. Gas Forming Reactions Acid + ionic carbonate  salt + CO2 + H2O HCl (aq) + K2CO3 (s)  KCl (aq) + H2CO3 (aq) H2CO3 (aq)  H2O (l) + CO2 (g) HCl (aq) + K2CO3 (s)  KCl (aq) + H2O (l) + CO2 (g)

  35. Reactions of Weak Acids Acetic Acid (aq) + sodium hydroxide (aq)  ? Molecular Equation: HC2H3O2 (aq) + NaOH (aq)  H2O (l) + NaC2H3O2 (aq) Total Ionic Equation: HC2H3O2 (aq) + Na+ (aq) + OH- (aq)  H2O(l) + Na+ (aq) + C2H3O2– (aq) Net Ionic Equation: HC2H3O2 (aq) + OH- (aq)  H2O(l) + C2H3O2– (aq) Spectator Ions: Na+

  36. Quiz 80.0 g of NaOH are added to 500 ml of water and used in a titration with a sample of 2.0 M sulfuric acid. What volume of NaOH was needed to neutralize the acid?

  37. OXIDATION-REDUCTION (REDOX) REACTIONS Demo—write the equation for the burning of Mg 2 Mg + O2 2 MgO

  38. LEO GER Reduction Gain of Electrons Oxidation state decreases Oxidation Loss of Electrons Oxidation state increases

  39. Redox con’t 2 Mg + O2 2 MgO Oxidation ½ Rxn Mg  Mg2+ + 2 e- Reduction ½ Rxn ½ O2 + 2e- O2- O2 Oxidizing Agent? Reducing Agent? Mg

  40. Redox Definitions 2 Mg + O2 2 MgO  Oxidizing Agent what does the oxidizing (the species that is reduced) Reducing Agent what does the reducing (the species that is oxidized)

  41. Oxidation Number (or oxidation state) A way to keep track of electrons in redox reactions Allow us to arbitrarily assign electrons to atoms in covalent compounds where electrons are shared—assigned to most electronegative atoms The charge an atom would have if electrons were transferred completely (and not shared) Rules for assigning oxidation states (table 4.3 pg. 160)

  42. General rules Rules for specific atoms or periodic table groups 1. For Group 1A(1): O.N. = +1 in all compounds 2. For Group 2A(2): O.N. = +2 in all compounds 3. For hydrogen: O.N. = +1 in combination with nonmetals 4. For fluorine: O.N. = -1 in combination with metals and boron 5. For oxygen: O.N. = -1 in peroxides O.N. = -2 in all other compounds(except with F) 6. For Group 7A(17): O.N. = -1 in combination with metals, nonmetals (except O), and other halogens lower in the group Table 4.3 Rules for Assigning an Oxidation Number (O.N.) 1. For an atom in its elemental form (Na, O2, Cl2, etc.): O.N. = 0 2. For a monoatomic ion: O.N. = ion charge 3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ion’s charge.

  43. Sample Problem 4.7 Determining the Oxidation Number of an Element PROBLEM: Determine the oxidation number (O.N.) of each element in these compounds: (a) zinc chloride (b) sulfur trioxide (c) nitric acid

  44. e- Y X Figure 4.12 A summary of terminology for oxidation-reduction (redox) reactions transfer or shift of electrons X loses electron(s) Y gains electron(s) X is oxidized Y is reduced X is the reducing agent Y is the oxidizing agent X increases its oxidation number Y decreases its oxidation number

  45. (a) 2Al(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g) (a) 2Al(s) + 3H2SO4(aq) Al2(SO4)3(aq) + 3H2(g) (b) PbO(s) + CO(g) Pb(s) + CO2(g) (b) PbO(s) + CO(g) Pb(s) + CO2(g) (c) 2H2(g) + O2(g) 2H2O(g) (c) 2H2(g) + O2(g) 2H2O(g) Sample Problem 4.7 Recognizing Oxidizing and Reducing Agents PROBLEM: Identify the oxidizing agent and reducing agent in each of the following: SOLUTION:

  46. Balancing Redox Rxns by the ON Method • Assign oxidation numbers to all elements in the reaction. • From the changes in oxidation numbers, identify the oxidized and reduced species. • Compute the number of electrons lost in the oxidation and gained in the reduction from the oxidation number changes (draw tie-lines between these atoms to show the changes). • Multiply one or both of these numbers by appropriate factors to make the electrons lost equal to the electrons gained, and use the factors as balancing coefficients. • Complete the balancing by inspection, adding states of matter.

  47. Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l) (a) Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l) (a) Cu(s) + HNO3(aq) Cu(NO3)2(aq) + NO2(g) + H2O(l) (b) PbS(s) + O2(g) PbO(s) + SO2(g) 0 +1 +5 -2 +2 +5 -2 +4 -2 +1 -2 Sample Problem 4.8 Balancing Redox Equations by the Oxidation Number Method PROBLEM: Use the oxidation number method to balance the following equations: SOLUTION: O.N. of Cu increases because it loses 2e-; it is oxidized and is the reducing agent. O.N. of N decreases because it gains1e-; it is reduced and is the oxidizing agent. loses2e- balance other ions gains 1e- x2 to balance e- balance unchanged polyatomic ions

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