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Chemistry I Honors—Unit 6 Chemical Equations, Reactions, & Redox

Chemistry I Honors—Unit 6 Chemical Equations, Reactions, & Redox. Objectives #1-3: 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 6Chemical Equations, Reactions, & Redox

  2. Objectives #1-3: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 #1-3:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing II. B. Writing Unbalanced Equations “Liquid hydrogen peroxide decomposes to form water vapor and pxygen 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 thethe presence 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. 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. Objectives #1-3:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing _____H2O2 _____ H2O + ______O2 ___CaC2 + ___H2O  ___C2H2 + __Ca(OH)2 ___C2H2 + ____O2  ____CO2 + _____H2O ____Pb(NO3)2 + ___NaI  ___NaNO3 + ___PbI2

  8. Objective #4: Assignment of Oxidation Numbers Part I: Oxidation vs. Reduction • Oxidation is the loss of electrons; during this process the charge of a species increases • Reduction is the gain of electrons; during this process the charge of a species decreases • “OIL RIG” or “LEO the lion goes GER”

  9. Objective #4: 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

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

  11. Objective #4: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

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

  13. Objective #4:Assignment of Oxidation Numbers • In general, during REDOX reactions, • Metalstend to lose electrons and are oxidized • Nonmetals tend to gainelectrons and are reduced

  14. Objective #4: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)!!

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

  16. Demo Redox Reaction: NaI(s) + H2SO4 (l) + MnO2 (s) I2 (g) + MnSO4(aq) + Na2SO4 (aq)+ H2O(l)

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

  25. Objective #6-8:Oxidizing and Reducing Agents • Examples—see packet

  26. Objective #6-8: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 #6-8:Oxidizing and Reducing Agents • Oxidizing and Reducing Ability • Example Demo: Cu + AgNO3 Cu(NO3)2 + Ag • assignment of oxidation numbers: 0 +1 +5 -2 +2 +5 -2 0 Cu + AgNO3 Cu(NO3)2 + Ag • Cu has been oxidized and therefore Cu is the reducing agent • Ag has been reduced and therefore AgNO3 is the oxidizing agent

  28. The more easily a species can lose electrons, the greater its ability to be a reducing agent and cause another species to gain electrons. • A species that loses electrons readily is unlikely to gain electrons and be reduced; such a species would not cause another species to lose electrons readily and therefore would act as a poor oxidizingagent

  29. Objective #6-8: Oxidizing and Reducing Agents • Example: Na + FeCl3 NaCl + Fe • assignment of oxidation numbers: 0 +3 -1 +1 -1 0 Na + FeCl3 NaCl + Fe • _____ is oxidized Na • _____ is reduced Fe

  30. *______ is the reducing agent and therefore would act as a ______ oxidizing agent Na, poor *______ is the oxidizing agent and therefore would act as a _______ reducing agent FeCl3, poor

  31. Obj. 9 & 10—Types of ChemRxns A. Synthesis Reactions • General formula: A + B  AB B. Decomposition Reactions • General formula: AB  A + B C. Single-Displacement Reactions • General formula: A + BC  AC + B D. Double Replacement Reactions • General Formula: AB + CD  AD + CB E.Combustion Reactions • General Formula: Hydrocarbon + O2 CO2 + H2O

  32. Objective #9: Oxidation-Reduction Reactions • Recall that oxidation-reduction reactions involve the transfer of electrons A. Synthesis Reactions • General formula: A + B AB • Examples: Nonmetal + oxygen nonmetal oxide S + O2 SO3 N2 + O2 NO2

  33. Metal + oxygen metal oxide Rb + O2 Rb2O Mg + O2MgO Nonmetal + sulfur nonmetal sulfide C + S  CS2 S + O2 SO3 (additional info needed) Metal + sulfur metal sulfide Rb + S Rb2S Mg + S MgS

  34. Metal + halogen  metal halide Na + Cl2 NaCl Ca + I2 CaI2 Metal oxide + water metal hydroxide (base) Na2O + H2O NaOH MgO + H2O  Mg(OH)2 Nonmetal oxide + water acid SO3 + H2O  H2SO4(add. info. needed) SO2 + H2O  H2SO3

  35. B. Decomposition Reactions General formula: AB  A + B Examples: Decomposition of binary compounds 2 elements H2O  H2 + O2 NaCl Na + Cl2 Decomposition of metal carbonates  carbon dioxide + metal oxide BaCO3BaO + CO2 Na2CO3 Na2O+CO2

  36. Decomposition of metal hydroxides  water + metal oxide NaOH H2O + Na2O Ca(OH)2H2O + CaO Decomposition of metal chlorates oxygen + metal chloride KClO3KCl + O2 Ca(ClO3)2 CaCl2 + O2 Decompostion of acids water + nonmetal oxide H2SO4  H2O + SO2

  37. C. Single-Displacement Reactions • General formula: A + BC  AC + B Examples: High metal + compound low metal + compound Fe + CuSO4 Cu + FeSO4 Cu + AgNO3 Ag + Cu(NO3)2 Active metal + water hydrogen + (low electronegativity) metal hydroxide Na + H2O  H2 + NaOH Ca + H2O  H2 + Ca(OH)2

  38. Metal + acid hydrogen + salt Zn + HCl ZnCl2 + H2 Mg + H3PO4 H2 + Mg3(PO4)2 High halogen + compound low halogen + compound F2 + NaCl Cl2 + NaF Br2 + NaI I2 + NaBr

  39. Objective #11: Using an 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 (SEE RXN. PACKET!!) • 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.

  40. Remember, however, that an activity series should only be used as a general guide for predicting simple replacement reactions (see Table 3 on p.286) • Examples: • Predict if the following reactions will occur: Zn + H2O --› (assume Zn is +2 if rx. occurs) No Rx. Sn + O2 --› (assume Sn is +4 if rx. occurs) Rx. Occurs; SnO2 will form

  41. Cd + Pb(NO3)2(assume Cd has a +2 charge if rx. occurs) Rx. occurs ; Cd(NO3)2 + Pb Cu + HCl(assume Cu has a charge of +2 if rx. occurs) No Rx.

  42. D. Double Replacement Reactions • General Formula: AB + CD  AD + CB Type I: Formation of a Precipitate (precipitation) Ionic compound + ionic compound  aqueous solution + precipitate Pb(NO3)2 + NaI NaNO3 + PbI2(s) Na2S + Pb(NO3)2PbS(s) + NaNO3

  43. Type II: Formation of a Gas Ionic compound + ionic compound  gas + aqueous solution + water NH4Cl + NaOH NH4OH + NaCl NH3 + H2O Na2SO3 + HCl H2SO3 + NaCl SO2 + H2O

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

  45. E. Combustion Reactions Examples: Element + oxygen oxide Mg + O2MgO Na + O2Na2O Hydrocarbon + oxygen carbon dioxide + water CH4 + O2 CO2 + H2O C9H18 + O2 CO2 + H2O

  46. Practice in Predicting the Products of Chemical Reactions (see example in lecture guide)

  47. Objectives #12: 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) + 2Cl-1(aq) Al(NO3)3(aq) Al+3(aq) + 3NO3-1(aq)

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

  49. Objective #12: Compounds in Aqueous Solutions Part III: Writing Net Ionic Equations • Net Reaction vs. Spectator Ions Examples:

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