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Chapter 18

Chapter 18. Oxidation-Reduction Reactions. 18.1 Electron Transfer Reactions. To learn about metal-nonmetal oxidation–reduction reactions To learn to assign oxidation states. Oxidation-Reduction Reactions.

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Chapter 18

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  1. Chapter 18 Oxidation-Reduction Reactions

  2. 18.1 Electron Transfer Reactions • To learn about metal-nonmetal oxidation–reduction reactions • To learn to assign oxidation states

  3. Oxidation-Reduction Reactions • Oxidation-Reduction (redox) reaction – a reaction in which one ore more electrons are transferred • Oxidation – loss of electrons • Reduction – gain of electrons

  4. Oxidation-Reduction Reactions • Which element is oxidized? • Which element is reduced?

  5. In the following reactions, identify which element is oxidized and which element is reduced: 2Mg(s) + O2(g)  2MgO(s) 2Al(s) + 3I2(s)  2AlI3(s)

  6. Oxidation States • Lets us keep track of electrons in oxidation-reduction reactions by assigning charges to the various atoms in a compound • Binary ionic compounds: oxidation state = the charge of the ion • NaCl • MgO

  7. An atom in a pure element has an oxidation number of 0 • Na • Cl2

  8. Oxidation states in covalent compounds – equal to the imaginary charges we determine by assuming that the most electronegative atom in a bond possesses both of the shared electrons • water

  9. The most electronegative elements are given oxidation states equal to the charge of their anion • F • O • N • Cl

  10. The sum of the oxidation states for an electrically neutral compound must be 0 • NO2

  11. Assign oxidation state to all atoms in the following: • CO2 • SF6 • NO3-

  12. 18.2 Balancing Oxidation-Reduction Reactions • To understand oxidation and reduction in terms of oxidation states • To learn to identify oxidizing and reducing agents • To learn to balance oxidation-reduction equations using half reactions

  13. Oxidation-Reduction Reactions Between Nonmetals • Oxidation- increase in oxidation state (loss of electrons) • Reduction – decrease in oxidation state (gain of electrons)

  14. 2Na(s) + Cl2(g)  NaCl • Na  oxidized • Na is also called the reducing agent (electron donor). • Cl2 reduced • Cl2 is also called the oxidizing agent (electron acceptor).

  15. CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) • C  oxidized • CH4 is the reducing agent. • O2 reduced • O2 is theoxidizing agent.

  16. Identify the atoms that are oxidized and those that are reduced and specify the oxidizing and reducing agents. 2Al(s) + 3I2(s)  2AlI3(s) PbO(s) + CO(g)  Pb(s) + CO2(g)

  17. Balancing redox reactions by the half reaction method • Half reaction – equation that have electrons as reactants or products • One half reaction represents a reduction process • Electrons are reactants • One half reaction represents an oxidation process • Electrons are products

  18. MnO4-(aq) + Fe2+(aq)  Fe3+(aq) + Mn2+(aq)

  19. Pb(s) + PbO2(s) + H+(aq)  Pb2+(aq) + H2O(l)

  20. 18.3 Electrochemistry and Its Applications • To understand the concept of electrochemistry • To learn to identify the components of an electrochemical (galvanic) cell • To learn about commonly used batteries • To understand corrosion and ways of preventing it • To understand electrolysis • To learn about the commercial preparation of aluminum

  21. Electrochemistry: An introduction • Electrochemistry – a study of the interchange of chemical and electrical energy • Two types of processes: • Production of an electric current from a chemical (redox) reaction • The use of an electrical current to produce a chemical change

  22. Making an electrochemical cell

  23. If electrons flow through the wire charge builds up. • Solutions must be connected to permit ions to flow to balance the charge.

  24. Electrochemistry: An Introduction • A salt bridge or porous disk connects the half cells and allows ions to flow, completing the circuit.

  25. Electrochemical battery (galvanic cell) – device powered by an oxidation-reduction reaction where the oxidizing agent is separated from the reducing agent so the electrons must travel through a wire from the reducing agent to the oxidizing agent

  26. Anode – electrode where the oxidation occurs • Cathode – electrode where the reduction occurs

  27. Electrolysis – a process where electrical energy is used to produce a chemical change • 2H2O(l)  2H2(g) + O2(g)

  28. Batteries • Lead Storage Battery • Anode reaction - oxidation • Pb + H2SO4  PbSO4 + 2H+ + 2e • Cathode reaction-reduction • PbO2 + H2SO4 + 2e + 2H+ PbSO4 + 2H2O

  29. Overall reaction • Pb + PbO2 + 2H2SO4 2PbSO4 + 2H2O

  30. Electrical Potential – the pressure on electrons to flow from one electrode to the other in a battery • Measured in volts

  31. Dry Cell Batteries – do not contain a liquid electrolyte • Acid version • Anode reaction - oxidation • Zn  Zn2+ + 2e • Cathode reaction – reduction • 2NH4+ + 2MnO2 + 2e Mn2O3 + 2NH3 + 2H2O

  32. Alkaline version • Anode reaction - oxidation • Zn + 2OHZnO + H2O + 2e • Cathode reaction – reduction • 2MnO2 + H2O + 2e Mn2O3 + 2OH

  33. Other types • Silver cell – Zn anode, Ag2O cathode • Mercury cell –Zn anode, HgO cathode • Nickel-cadmium –rechargeable –products turned back into reactants by the use of external source of current

  34. Cathodic protection of an underground pipe Corrosion • Corrosion is the oxidation of metals to form mainly oxides and sulfides. • Some metals, such as aluminum, protect themselves with their oxide coating. • Corrosion of iron can be prevented by coatings, by alloying and cathodic protection.

  35. Electrolysis • Electrolysis – a process involving forcing a current through a cell to produce a chemical change that would not otherwise occur

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