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Chapter 18:

Chapter 18:. Oxidation-Reduction ( Redox ) Reactions and Electrochemistry. Students will learn about. What is redox reaction? Assigning oxidation states Oxidizing agent; Reducing agent Balancing redox reactions Half-reaction method Applications of electrochemistry electroplating

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Chapter 18:

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  1. Chapter 18: Oxidation-Reduction (Redox) Reactions and Electrochemistry

  2. Students will learn about.. • What is redox reaction? • Assigning oxidation states • Oxidizing agent; Reducing agent • Balancing redox reactions • Half-reaction method • Applications of electrochemistry • electroplating • batteries • corrosion • electrolysis

  3. Oxidation-Reduction (Redox) • Transferring electrons from atoms to atoms • Oxidation: losing electrons (Ex) Na  Na+ + e‒ *Electrons are shown as a product • Reduction: gaining electrons (Ex) Cl2 + 2e‒ 2Cl ‒ *Electrons are shown as a reactant

  4. Examples • Identify which element is oxidized and which is reduced • 2Mg + O2 2MgO • 2Al + 3I2  2AlI3

  5. Oxidation State (≈oxidation number) • the charge number assigned to each atomof a compound • These charge numbers are assigned according to the rules • The rules are similar to assigning the charge numbers based on the number of valence electrons • Not a permanent number

  6. Rules continue… 1) Oxidation state of an atom in an element = 0 (Ex)Na = 0; O2 = 0 2) Oxidation state of monatomic ion = charge of the ion (Ex) Na+1 = +1; O2- = -2 3) Oxygen = 2 (except in peroxides where O= 1 or when it is bonded to a more electronegative element ) (Ex) H2O: O = -2; H2O2 (peroxide): O = -1; OF2: O = +2 4) In covalent compounds, hydrogen = +1 (Ex) H2O: H = +1; CH4: H = +1

  7. Rules for Assigning Oxidation States 5) In binary compounds, element with the greater electronegativity = -1 (Ex) Fluorine = 1 in any compound *Other 7A nonmetals = -1 if not bonded with F (Ex)ClF: F = -1, Cl = +1; NaCl: Na = +1, Cl = -1 6) 1A metals = +1 7) 2A metals = +2 8) All other elements, follow the following rules: (a) Sum of oxidation states = 0 in compounds (b) Sum of oxidation states = charge of the ion in ions

  8. Examples Find the oxidation states for each of the elements in each of the following compounds: • K2Cr2O7 • CO32- • SO3 • SO2 • C2H6

  9. Revisit Oxidation & Reduction • Oxidation • losing electron(s) • Electron(s) written as a product • Electron donor • increasing the oxidation number • Reduction • gaining electron(s) • Electron(s) written as a reactant • Electron acceptor • decreasing the oxidation number

  10. Example Identify the atoms that are oxidized and those reduced. PbO + CO  Pb + CO2

  11. Oxidizing and Reducing Agent • Oxidizing agent • The element that makes another element oxidized • Itself is reduced • Reducing agent • The element that makes another element reduced • Itself is oxidized

  12. Examples (a) 2Na(s) + Cl2(g) 2NaCl(s) • Na  oxidized • Na = reducing agent (electron donor). • Cl2 reduced • Cl2= oxidizing agent (electron acceptor).

  13. (b) CH4(g) + 2O2(g) CO2(g) + 2H2O(g) • C  oxidized • CH4 = reducing agent. • O2 reduced • O2 = oxidizing agent (c) 2 CuCl CuCl2 + Cu • CuCl oxidized and reduced • CuCl = oxidizing agent and reducing agent

  14. Half-reaction • Oxidation half-reaction (Ex) Na  Na+ + e ‒ • Reduction half-reaction (Ex) Na+ + e ‒  Na

  15. Balancing Redox reactions • #of electrons lost by oxidation = # of electrons gained by reduction • The mass is conserved • Use H2O to balance the numbers of H and O atoms • If the redoxrxns occur in acidic solutions, H+ shows in the final equation • If the redoxrxns occur in basic solutions, OH ‒shows in the final equation • Total charge of all reactants = Total charge of all products

  16. Examples (a) Balance by the half-reaction method. Ce2+ + Co2+→ Ce3+ + Co

  17. Examples (b) The following reaction takes place in an acidic solution. Balance by the half-reaction method. Cr2O72-(aq) + SO32-(aq)  Cr3+(aq) + SO42-(aq)

  18. Applications of RedoxRxns • Electrochemistry • Batteries • Electrolysis • Corrosion

  19. Electrochemistry • Study of the interchange of chemical and electrical energy • Production of electric current • Using electric current to produce chemical change

  20. Electrochemical Cell • Apparatus in which electrochemical reactions take place (Ex) 8H+ + MnO4– + 5e– → Mn2+ + 4H2O Fe2+→ Fe3+ + e–

  21. Anode– electrode where oxidation occurs • Cathode – electrode where reduction occurs

  22. Lead Car Battery Pb + PbO2 + 2H2SO4 2PbSO4 + 2H2O

  23. Dry Cell Battery • Zn  Zn2+ + 2e ‒ 2NH4+ + 2MnO2 + 2e ‒  Mn2O3 + 2NH3 + H2O

  24. Corrosion • Fe(s)  Fe2+(aq) + 2e- • O2(g) + 2H2O(l) + 4e- 4OH-(aq) • Fe2+(aq) + 2OH-(aq)  Fe(OH)2(s) • Fe(OH)2(s) + O2(g)  2Fe2O3 •H2O(s) + 2H2O(l)

  25. Electrolysis

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