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AP Chem - Unit 2

AP Chem - Unit 2

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AP Chem - Unit 2

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  1. AP Chem - Unit 2 Reactions and Solution Stoichiometry

  2. Types of Reactions • Combination (Synthesis) • Decomposition • Combustion • Single Replacement • Double Replacement (Metathesis)

  3. Combination/Synthesis Basic Form A + B  AB Rxns of elements with oxygen and sulfur C(s) + O(s)  CO2 (g) Rxns of metals with halogens 2Na(s) + Cl2 (g)  2NaCl Synthesis rxns with oxides CaO(s) + H2O (l)  Ca(OH)2 (s)

  4. Decomposition Basic Form AB  A + B Binary compounds 2H2O (l)  2H2 (g) + O2 (g) Metal hydroxides Cu(OH)2 (s)  CuO(s) + H2O (l)

  5. More Decomposition Examples Metal carbonates PbCO3 (s)  PbO (s) + CO2 (g) Metal chlorates 2KClO3 (s)  2KCl (s) + 3O2 (g) Oxyacids H2CO3 (l)  H2O (l) + CO2 (g)

  6. Combustion A substance combines with oxygen, releasing a large amount of energy in the form of light and heat. True or False: The products of every combustion reaction are CO2 and H2O False

  7. Combustion Examples Reactive elements combine with oxygen P4(s) + 5O2(g)  P4O10(s) (This is also a synthesis reaction) The burning of natural gas, wood, gasoline C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

  8. Single Replacement Basic Form A + BY  AY + B Metals by other metals 3Mg + 2FeCl3  2Fe + 3MgCl2 Hydrogen in water by another metal 2Na +2H2O  H2 + 2NaOH

  9. More Single Replacement Hydrogen in an acid by another metal 2Li + 2HCl  H2 + 2LiCl Halogens by more active halogens Cl2 + 2KI  I2 + 2KCl

  10. The Activity Series of the Metals • Lithium • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver • Platinum • Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Metals above hydrogen can replace hydrogen in acids. Metals from sodium upward can replace hydrogen in water

  11. The Activity Series of the Halogens • Fluorine • Chlorine • Bromine • Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F2(g)  2NaF(s) + Cl2(g) MgCl2(s) + Br2(g)  ??? No Reaction

  12. Double Replacement (Metathesis) Basic Form AX + BY  AY + BX One of the compounds formed is usually a precipitate, an insoluble gas that bubbles out of solution, or a molecular compound, usually water. How do you tell if a precipitate will form? Use your Solubility Rules!!!

  13. Mostly Soluble

  14. Mostly Insoluble

  15. The Solubility Rule Song!!!! • (Sing to Rhythm of 99 Bottles) •  Alkali metals and ammonium salts,Whatever they may be, • Can always be depended on for solubility • When asked about the nitrates • The answer is always clear, • They each and all are soluble, • Is all we want to hear. • Most every chloride's soluble • At least we've always read • Save silver, mercury one • And chloride of lead • Every single sulfate • Is soluble, 'Tis said • 'Cept barium, strontium, mercury one And calcium and lead. • - • Hydroxides in general • don't dissolve at all • But, barium, strontium and calcium • Are slightly soluble • *but don't forget * • Alkali metals and ammonium salts • Whatever they may be • Can always be depended on • For solubility  • The carbonates are insoluble, • It's lucky that it is so, • Or else, our marble buildings • Would melt away like snow. • *but, once again, don't forget * • Alkali metals and ammonium salts • Whatever they may be • Can always be depended on • For solubility

  16. Net Ionic Equations Written to show only the species that react or undergo change in aqueous solution. • Steps for writing net ionic equations • Write a balanced molecular equation • Rewrite the equation to show the ions that form in solution • Identify and cancel spectator ions Spectator Ions do NOT participate in the reaction

  17. Sample Problem • What compound precipitates when solutions of Fe2(SO4)3 and LiOH are mixed? Write the complete chemical reaction for the reactants above. Write the net ionic equation. List any spectator ions.

  18. Another Type of Metathesis • Acid/Base Reactions • Two Types • Neutralization • HCl + NaOH NaCl + H2O • Gas Formation • 2HCl +Na2S  H2S(g) + 2NaCl • HCl + NaHCO3  NaCl + H2CO3 • H2CO3  H2O + CO2(g)

  19. Oxidation Reduction Rxns • Electrons are transferred • Oxidation = Loses Electrons • Reduction = Gains Electrons OIL RIG LEO goes GER!! OR

  20. RedoxRxn Each sodium atom loses one electron atom gains one electron: Each chlorine atom gains one electron:

  21. LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

  22. OIL RIG: • Oxidation Is Loss of electrons Sodium is oxidized • Reduction Is Gain of electrons Chlorine is reduced

  23. Reducing Agents and Oxidizing Agents • The substance reduced is the oxidizingagent • The substance oxidized is the reducingagent Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

  24. Rules for Assigning Oxidation #s • Atoms in elemental form are always zero • Monatomic ions equal the charge on the ion • Oxygen is usually -2 (except peroxide is -1) • Hydrogen is +1 with NM and -1 with M • Fluorine is always -1 • Sum of the ox #s of all atoms in a neutral compound is zero • Sum of the ox #s in a polyatomic ion = the ion charge

  25. Sample Problem • Determine the oxidation #s for the following: • H2S • S8 • SCl2 • Na2SO3 • SO42-

  26. Oxidation of metals by acids and salts • Similar to single replacement rxns, sometimes called Displacement Rxns • General form • A + BX  AX + B • Examples • Zn(s) + 2HBr (aq)  ZnBr2(aq) + H2 (g) • Mn(s) + Pb(NO3)2(aq)  Mn(NO3)2 (aq)+ Pb(s)

  27. Rules for Balancing RedoxRxns • Assign oxidation #s • Determine which species are being oxidized or reduced • Divide the equation into 2 half-reactions • Balance each half reaction • Balance elements other than H and O • Balance O by adding water as needed • Balance H by adding H+ as needed • Balance charge by adding e- • Multiply half reactions by integers so # of e- are the same in each reaction • Add half reactions together, simplifying when you can • Check to see if equation is charge and mass balanced

  28. Sample Problem • Balance the following MnO4-(aq) + C2O42-(aq) Mn2+(aq) + CO2(aq)

  29. Other Rxns • Metal Oxide in Water • MO + H2O  MOH • BaO + H2O  Ba(OH)2 • Nonmetal Oxide in Water • NO + H2O  oxyacid • N2O3 + H2O  HNO2

  30. Solutions

  31. Solute A solute is the dissolved substance in a solution. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

  32. Definition of Electrolytes and Nonelectrolytes An electrolyte is: • A substance whose aqueous solution conducts an electric current. A nonelectrolyte is: • A substance whose aqueous solution does not conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes…

  33. Which of the following are electrolytes? • Pure water • Tap water • Sugar solution • Sodium chloride solution • Hydrochloric acid solution • Lactic acid solution • Ethyl alcohol solution • Pure, solid sodium chloride

  34. Answers… ELECTROLYTES: NONELECTROLYTES: • Tap water (weak) • NaCl solution • HCl solution • Lactate solution (weak) • Pure water • Sugar solution • Ethanol solution • Pure, solid NaCl

  35. Electrolytes & Net Ionic Equations Ionic CompoundsDissociate • NaCl(s)  Na+(aq) + Cl-(aq) • AgNO3(s)  Ag+(aq) + NO3-(aq) • MgCl2(s)  Mg2+(aq) + 2 Cl-(aq) • Na2SO4(s)  2 Na+(aq) + SO42-(aq) • AlCl3(s)  Al3+(aq) + 3 Cl-(aq) These compounds should be written as ions in net ionic equations.

  36. Strong acids such as HCl are completelyionized in solution. So they are written as ions in the net ionic equation Other examples of strong acids include: • Sulfuric acid, H2SO4 • Nitric acid, HNO3 • Hydriodic acid, HI • Perchloric acid, HClO4 Weak acids such as lactic acid usually ionize less than 5% of the time. So they are not written as ions in the net ionic equation.

  37. Molarity The concentration of a solution measured in moles of solute per liter of solution. mol= M L

  38. Sample Problem • How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaClsol’n?

  39. Serial Dilution It is not practical to keep solutions of many different concentrations on hand, so chemists prepare more dilute solutions from a more concentrated “stock” solution. MstockVstock = MdiluteVdilute

  40. Sample Problem • What volume of stock (11.6 M) hydrochloric acid is needed to prepare a 3.0 M - 250. ml solution?

  41. Sample Problem • How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?

  42. Sample Problem • A sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M silver nitrate, resulting in the formation of a precipitate. • Write the molecular equation for the reaction • What is the limiting reactant? • Calculate the theoretical yield, in grams, of the ppt that forms.

  43. Sample Problem • 45.7 mL of 0.500 H2SO4 is required to neutralize a 20.0 mL sample of NaOH solution. What is the concentration of the NaOH solution?