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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 --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 #1-2:Introduction to Chemical Reactions, Reaction Interpretation, and Balancing (examples in lecture guide)

  6. Objective #3: 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” (oxidation is loss, reduction is gain)

  7. Objective #3: Assignment of Oxidation Numbers (video clip) • 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) 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 #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 Al(s) + I2(s) --› 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 #3 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 #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 • Examples (see packet) • Demo Redox Reaction

  13. Demo Reaction

  14. Objective #4 Balancing Redox Reactions *Writing Half-Reactions (charges and atoms must balance) (examples)

  15. 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 using hydrogen ions.

  16. Objective #4: 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)

  17. Redox Examples

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

  19. 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

  20. Objective #6 Oxidation-Reduction Reactions *recall that oxidation-reduction reactions involve the transfer of electrons A. Synthesis Reactions *general formula: A + B --›AB *examples:

  21. B. Decomposition Reactions *general formula: AB --› A + B *examples:

  22. C. Single-Displacement Reactions *general formula: A + BC --› AC + B *examples:

  23. D. Combustion Reactions *examples: Element + oxygen -- oxide Hydrocarbon + oxygen -- water and carbon dioxide

  24. Objective #8 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

  25. *it should be remembered however that an activity series should only be used as a general guide for predicting singlereplacement 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

  26. 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.

  27. Objective #7 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

  28. *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

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

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

  31. Objectives #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) + 2Cl-1(aq) Al(NO3)3(aq) --› Al+3(aq) + 3NO3-1(aq)

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

  33. Objectives #9: Compounds in Aqueous Solutions Part III Writing Net Ionic Equations *net reaction vs. spectator ions (examples)

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