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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 6: Chemical 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 --See examples in packet 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

  5. Objectives #1-3:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing (examples in lecture guide)

  6. 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”

  7. 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 + 2e-1 *What is the oxygen doing? O + 2e-1 --› O-2

  8. *Which element has been oxidized? Mg *Which element has been reduced? O

  9. 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: ∆ 2 Al(s) + 3 I2(s) --› 2 AlI3 (s) *What is the aluminum doing? Al --› Al+3 + 3e-1 *What is the iodine doing? I + e-1 --› I-1

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

  11. Objective #4 Assignment of Oxidation Numbers • In general, during REDOX reactions, • Metals tend to lose electrons and are oxidized • Nonmetalstend to gain electrons and are reduced

  12. 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 • Examples (see packet) • Demo Redox Reaction

  13. Objective #5 Balancing Redox Reactions *Writing Half-Reactions (charges and atoms must balance) (examples)

  14. 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 using hydrogen ions.

  15. Objective #5: Balancing Redox Reactions 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 atoms and chargesbalance. (examples)

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

  17. 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 reduction is said to be disproportionate.

  18. 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

  19. *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 oxidizing agent

  20. 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

  21. *______ 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

  22. Objective #9 Oxidation-Reduction Reactions *recall that oxidation-reduction reactions involve the transfer of electrons A. Synthesis Reactions *general formula: A + B --›AB *examples of types: Nonmetal + oxygen --› nonmetal oxide S + O2 --› SO3 N2 + O2 --› NO2

  23. 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 + S8 --› Rb2S Mg + S --› MgS

  24. *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

  25. 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

  26. BaCO3 --› BaO + CO2 Na2CO3 --› Na2O+CO2 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

  27. 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 + metal hydroxide Na + H2O --› H2 + NaOH Ca + H2O --› H2 + Ca(OH)2

  28. 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

  29. D. Combustion Reactions *examples: Element + element --› oxide Mg + O2 --› MgO Na + O2 --› Na2O Hydrocarbon + oxygen --› carbon dioxide + water CH4 + O2 --› CO2 + H2O C9H18 + O2 --› CO2 + H2O

  30. Objective #11 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 higher on activity series that the element in the compound it is trying to displace

  31. *it should be remembered however that an activity series should only be used as a general guide for predicting simple replacement reactions (see Table 3 on p.286) *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

  32. 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.

  33. Objective #10 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

  34. *Type II Formation of a Gas Ionic compound + ionic compound --› gas + aqueous solution + water NH4Cl + NaOH --› NH4OH + NaCl ^ NH3 + H2O Na2CO3 + HCl --› H2SO3 + NaCl ^ SO2 + H2O

  35. HCl + Na2CO3 --› H2CO3 + NaCl ^ CO2 + H2O *Type III Formation of Water (acid-base) Acid + Base --› water + salt NaOH + HCl --› H2O + NaCl Ca(OH)2 + HCl --› H2O + CaCl2

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

  37. 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)

  38. Part II Predicting Precipitation *use of the solubility table in lecture guide *examples:

  39. Objectives #12: Compounds in Aqueous Solutions Part III Writing Net Ionic Equations *net reaction vs. spectator ions (examples)

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