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Understanding Electron Transfer and Electrochemistry: Principles and Applications

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This resource covers essential concepts of electron transfer reactions, oxidation-reduction (redox) processes, and electrochemistry. Learn to recognize oxidation states, balance redox reactions using the half-reaction method, and identify oxidizing and reducing agents. The significance of electrochemical cells, types of batteries, and corrosion prevention techniques are also explored. Gain insights into the conversion of chemical energy to electrical energy and the role of electrolysis in driving nonspontaneous reactions, crucial for various commercial applications.

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Understanding Electron Transfer and Electrochemistry: Principles and Applications

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  1. Unit 18 Section 1 – Electron Transfer Reactions Section 2 – Balancing Oxidation Reduction Reactions Section 3 – Electrochemistry and its Applications

  2. Objectives Section 1 – Electron Transfer Reactions • To learn about metal-nonmetal oxidation–reduction reactions • To learn to assign oxidation states

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

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

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

  6. Oxidation States

  7. Objectives Section 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

  8. 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).

  9. 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.

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

  11. Balancing Oxidation-Reduction Reactions by the Half-Reaction Method

  12. Objectives Section 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

  13. 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

  14. Electrochemistry: An Introduction • Making an electrochemical cell

  15. 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.

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

  17. 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

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

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

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

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

  22. 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

  23. 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

  24. 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

  25. 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.

  26. 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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