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To learn about metal-nonmetal oxidation–reduction reactions To learn to assign oxidation states

Objectives. To learn about metal-nonmetal oxidation–reduction reactions To learn to assign oxidation states . A. Oxidation-Reduction Reactions . Oxidation-reduction reaction – a chemical reaction involving the transfer of electrons Oxidation – loss of electrons

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To learn about metal-nonmetal oxidation–reduction reactions To learn to assign oxidation states

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  1. Objectives • To learn about metal-nonmetal oxidation–reduction reactions • To learn to assign oxidation states

  2. A. Oxidation-Reduction Reactions • Oxidation-reduction reaction – a chemical reaction involving the transfer of electrons • Oxidation – loss of electrons • Reduction – gain of electrons

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

  4. B. Oxidation States • Oxidation states – allow us to keep track of electrons in oxidation-reduction reactions

  5. B. Oxidation States

  6. Objectives • 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

  7. A. Oxidation-Reduction Reactions Between Nonmetals • 2Na(s) + Cl2(g) 2NaCl(s) • Na  oxidized • Na is also called the reducing agent (electron donor). • Cl2 reduced • Cl2 is also called the oxidizing agent (electron acceptor).

  8. A. Oxidation-Reduction Reactions Between Nonmetals • CH4(g) + 2O2(g) CO2(g) + 2H2O(g) • C  oxidized • CH4 is the reducing agent. • O2 reduced • O2 is theoxidizing agent.

  9. B. Balancing Oxidation-Reduction Reactions by the Half-Reaction Method • Half reaction – equation which has electrons as products or reactants

  10. B. Balancing Oxidation-Reduction Reactions by the Half-Reaction Method

  11. Objectives • 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

  12. A. Electrochemistry: An Introduction • Electrochemistry – the study of the interchange of chemical and electrical energy • Two types of processes • Production of an electric current from a chemical reaction • The use of electric current to produce chemical change

  13. A. Electrochemistry: An Introduction • Making an electrochemical cell

  14. A. Electrochemistry: An Introduction • If electrons flow through the wire charge builds up. • Solutions must be connected to permit ions to flow to balance the charge.

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

  16. A. Electrochemistry: An Introduction • Electrochemical battery (galvanic cell) – device powered by an oxidation-reduction reaction where chemical energy is converted to electrical energy • Anode – electrode where oxidation occurs • Cathode – electrode where reduction occurs

  17. A. Electrochemistry: An Introduction • Electrolysis – process where electrical energy is used to produce a chemical change • Nonspontaneous

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

  19. B. Batteries • Overall reaction • Pb + PbO2 + 2H2SO4 2PbSO4 + 2H2O

  20. B. Batteries • Electric Potential – the “pressure” on electrons to flow from anode to cathode in a battery

  21. B. Batteries • 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

  22. B. Batteries • Dry Cell Batteries – do not contain a liquid electrolyte • Alkaline version • Anode reaction - oxidation • Zn + 2OH ZnO + H2O + 2e • Cathode reaction – reduction • 2MnO2 + H2O + 2e Mn2O3 + 2OH

  23. B. Batteries • Dry Cell Batteries – do not contain a liquid electrolyte • Other types • Silver cell – Zn anode, Ag2O cathode • Mercury cell –Zn anode, HgO cathode • Nickel-cadmium – rechargeable

  24. Cathodic protection of an underground pipe C. 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.

  25. D. 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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