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Chemistry I Honors—Unit 6: Chemical Equations / Reactions/Redox

Chemistry I Honors—Unit 6: Chemical Equations / Reactions/Redox. Objectives #1-2: Introduction to Chemical Reactions, Reaction Interpretation, and Balancing. Chemical Equations Describe chemical reactions Starting substances are called reactants Ending substances are called products

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Chemistry I Honors—Unit 6: Chemical Equations / Reactions/Redox

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  1. Chemistry I Honors—Unit 6: Chemical Equations/Reactions/Redox

  2. Objectives #1-2:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing • Chemical Equations • Describe chemical reactions • Starting substances are called reactants • Ending substances are called products • All chemical reactions must follow the Law of Conservation of Matter by being balanced

  3. II. Interpreting Chemical Equations A. Symbols

  4. Objectives #1-2:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing II. B. Writing Unbalanced Equations “Liquid hydrogen peroxide decomposes to form water vapor and oxygen gas in the presence of the catalyst manganese (IV) oxide.” “Solid calcium carbide (CaC2) reacts with water to form ethyne gas and aqueous calcium hydroxide.”

  5. “Ethyne gas (C2H2) reacts with oxygen in thepresence of a flame to produce carbon dioxide gas and water vapor.” • “Aqueous solutions of lead (II) nitrate and sodium iodine react to form lead (II) iodide and aqueous sodium nitrate.”

  6. Objectives #1-2:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing III. Balancing Chemical Equations • Basic Procedures: • Be sure all formulasare correct before attempting to balance • Never balance by changing subscripts • Use coefficientsto balance • Typeand number of atoms on each side of reaction must balance • Coefficients used must be in the lowest ratio possible

  7. Examples: _____H2O2 _____ H2O + ______O2 ___CaC2 + ___H2O  ___C2H2 + __Ca(OH)2

  8. ___C2H2 + ____O2  ____CO2 + _____H2O ____Pb(NO3)2 + ___NaI  ___NaNO3 + ___PbI2

  9. Objective #3:Assignment of Oxidation Numbers Part I: Oxidation vs. Reduction • Oxidation is the lossof electrons; during this process the charge of a species increases • Reduction is the gain of electrons; during this process the chargeof a species decreases • “OIL RIG”: oxidation is loss, reduction is gain • “LEO the lion goes GER”

  10. Objective #3: Assignment of Oxidation Numbers • Example I: Solid magnesium is reacted with oxygen gas in the air to produce solid magnesium oxide • Equation: 0 0 +2 -2 • Mg (s) + O2 (g) 2 MgO(s) *What is the magnesium doing? Mg  Mg+2 + 2 e-1 *What is the oxygen doing? O + 2e-1  O-2

  11. Which element has been oxidized? Mg • Which element has been reduced? O

  12. Objective #3:Assignment of Oxidation Numbers • Example II: Water is added to produce sufficient heat to react solid forms of aluminum and iodine. The resulting reaction produces solid aluminum iodide. • Equation: 0 0∆ +3-1 2 Al(s) + 3 I2(s)2 AlI3 (s) *What is the aluminum doing? Al  Al +3 + 3 e -1 *What is the iodine doing? I + e -1  I -1

  13. *Which element has been oxidized? Al *Which element has been reduced? I

  14. Objective #3 :Assignment of Oxidation Numbers • In general, during REDOX reactions, • Metals tend to lose electrons and are oxidized • Nonmetals tend to gain electrons and are reduced

  15. Objective #3:Assignment of Oxidation Numbers Part II: Utilization of Oxidation Number Rules • See text p.232-233 • The “Big 4”: Group I elements are +1 Group II elements are +2 H is usually +1 O is usually -2 • Remember: 1)Elements are always neutral (zero)! • 2) The total of the oxidation • numbers in a compound must be • neutral (zero)!!

  16. Oxidation Number Examples: He NaCl Na2Cr2O7 Ca(ClO3)2 OF Mg3(PO4)2 CrO4-2

  17. Objective #4:Balancing Redox Reactions • Writing Half-Reactions (charges and atoms must balance to in order to be conserved! ) • Examples: • K  K+1 + _____ (__________) • S + _______  S-2 (__________) • Mg  Mg+2 + _______ (__________) • _____F-1  ______+ F2 (__________)

  18. Objective #4: Balancing Redox Reactions • Key Steps: 1.Write half-reactionsfor the oxidation and reduction sections of the reaction. 2. Balance all elements except hydrogenand oxygen. 3. Balance oxygen by using water. 4. Balance hydrogen by usinghydrogen ions.

  19. 5. Balance charge by adding electrons to the side that is deficientin electrons. 6. Equalize electrons lost and gained by multiplyingeach half-reaction by an appropriate factor. 7.Addtogether half-reactions and cancellike species. 8. Check that atomsand chargesbalance.

  20. Example #1: • MnO4-1 + Fe -2 Fe +3 + Mn-2

  21. Example #2: • Cr2O7-2 + Cl-1 Cr+3 + Cl2

  22. Example #3: • Ce+4 + H3AsO3 Ce+3 + H3AsO4

  23. Example #4: I2 + OCl-1 IO3-1 + Cl -1

  24. Objective #5: Oxidizing & Reducing Agents • Examples: • Copper + silver nitrate  silver + copper (II) nitrate • Element Oxidized:_______ “The Box:” • Element Reduced: _______ O: OA: • Oxidizing Agent:________ • Reducing Agent: ________ R: RA:

  25. Objective #5:Oxidizing and Reducing Agents • Examples—see packet

  26. Objective #5:Oxidizing and Reducing Agents • Summary: • The charge of the element oxidized goes up • The charge of the element reduced goes down • The item oxidized is the reducingagent • The item reduced is the oxidizingagent • A species that is the source of BOTH oxidation and reductionis said to be disproportionate.

  27. Objective #6: Oxidation-Reduction Reactions • Recall that oxidation-reduction reactions involve the transfer of electrons A. Synthesis Reactions • General Equation: A + B  AB • Examples: • Mg + O2  • Na + F2  • Na2O + H2O  • SO2 + H2O 

  28. B. Decomposition Reactions • General Equation: AB  A + B • Examples: • NaCl • Na2CO3  • KClO3 

  29. C. Single-Displacement Reactions • General Equation: A + BC  AC + B • Examples: • Fe + CuCl2 • Na + H2O  • Zn + HCl  • F2 + NaCl 

  30. D. Combustion Reactions • Examples: Element + oxygen  _____ oxide Hg + O2  Na + O2  Hydrocarbon + oxygen water and carbon dioxide CH4 + O2  C9H18 + O2 

  31. Objective #7: Activity Series • An activity series is a vertical listing of elements in terms of their chemical reactivity; elements that are more reactive are listed at the top and less reactive elements are listed near the bottom. • A reactive element can readily transfer its valence electrons to another element. • In general, for a single replacement reaction to go to completion, the lone element in the reaction must be higheron activity series that the element in the compound it is trying to displace.

  32. It should be remembered, however, that an activity series should is used as a general guide for predicting single replacement reactions. It can also be used to predict the reactions of metals with oxygen, if that information is included on the chart. (See Table 3 on p.286) • Examples: Predict if the following rxns will occur and determine the products: • Zn + H2O (assume Zn is +2 if rx. occurs) No Rxn. • Sn + O2 (assume Sn is +4 if rx. occurs) Rxn. Occurs --SnO2

  33. Cd + Pb(NO3)2 (assume Cd has a +2 charge if rxn. occurs) Rxn. Occurs  Cd(NO3)2 + Pb • Cu + HCl (assume Cu has a charge of +2 if rxn. occurs) No Rxn.

  34. Objective #8: Double Replacement Reactions • General Equation: AB + CD  AD + CB • Type I: Formation of a Precipitate (precipitation) Ionic compound + ionic compound  aqueous solution + precipitate* Examples: Pb(NO3)2 (aq)+ NaI(aq)NaNO3 (aq)+ PbI2 (s) Na2S (aq)+ Pb(NO3)2(aq)PbS(s) + NaNO3 (aq) * Check the solubility chart to determine the ppt.

  35. Type II: Formation of a Gas Ionic compound + ionic compound  gas + aqueous solution + water NH4Cl + NaOH NH4OH + NaCl(aq) NH3 (g)+ H2O (l) Na2SO3 + HCl H2SO3 + NaCl(aq) SO2 (g)+ H2O (l)

  36. Type III: Formation of Water (acid-base) Acid + Base  water + salt* NaOH(aq)+ HCl(aq)H2O (l)+ NaCl(aq) Ca(OH)2 (aq)+ HCl(aq)H2O(l) + CaCl2 (aq) *SALT= an ionic compound that does NOT contain H+ or OH-

  37. Practice: Predicting the Products of Chemical Reactions Use the General Lab Fact Sheet to help predict the products of the reactions listed in the packet…

  38. Objective #9: Compounds in Aqueous Solutions Part I: Dissociation of Ionic Compounds • Dissociation process: The separation of ions that occurs when an ionic compound is dissolved in water. • Examples: CaCl2(aq) Ca+2(aq) + 2 Cl-1(aq) Al(NO3)3(aq) Al+3(aq) + 3 NO3-1(aq)

  39. Part II: Predicting Precipitation • Use of the solubility table in lecture guide • Examples:

  40. Part III:Writing Net Ionic Equations • Net Reaction vs. Spectator Ions Net Reaction shows the formation of the solid precipitate from aqueous ions. Spectator ions do not change from the aqueous state during the reaction. • Examples:

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