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Electrochemistry

Electrochemistry. AP Chapter 20. Electrochemistry. Electrochemistry relates electricity and chemical reactions. It involves oxidation-reduction reactions (aka – redox ) They are identified by the change in the oxidation state of one or more elements. Identifying Redox Reactions.

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Electrochemistry

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  1. Electrochemistry AP Chapter 20

  2. Electrochemistry • Electrochemistry relates electricity and chemical reactions. • It involves oxidation-reduction reactions (aka – redox) • They are identified by the change in the oxidation state of one or more elements.

  3. Identifying Redox Reactions

  4. Redox Reactions • In every redox reaction, one substance is oxidized and the other is reduced. • Substances that lose electrons from one side of the equation to the other are oxidized. • Substances that gain electrons are reduced.

  5. Redox • Loss of electrons – LEO • Gain of electrons - GER

  6. Using oxidation states to recognize a redox reaction.

  7. Oxidizing and Reducing Agents • A reducing agent is the agent or substance that causes another substance to be reduced. • An oxidizing agent is the agent or substance that causes another substance to be oxidized. • Generally, the substance that is reduced is the oxidizing agent and the substance that is oxidized is the reducing agent.

  8. Balancing Redox Reactions • These reactions are balanced by showing the transfer of electrons from one substance to another using half-reactions. • The electrons lost by one substance are gained by the other. • Electrons are never created or destroyed, just transferred.

  9. Half-Reactions • Take place simultaneously. • Sn2+(aq) + 2 Fe3+(aq) → Sn4+(aq) + 2Fe2+(aq) • LEO: Sn2+(aq) → Sn4+(aq) + 2 e- • GER: 2 Fe3+(aq) + 2 e-→ 2Fe2+(aq) • Oxidation half-reactions have the e- on the product side and reduction half-reactions have the e- on the reactant side.

  10. Voltaic Cells • The energy released in a spontaneous redox reaction can be used to perform electric work. • Voltaic cells are also called galvanic cells, where the transfer of electrons takes place through an external pathway (wire) rather than directly between reactants.

  11. Electron flow goes A → C LEO Anode GER Cathode Salt Bridge allows the flow of ions

  12. Cells Potential • The potential difference between the two electrodes in a voltaic cell provides the driving force that pushes electrons through the wire, or external circuit.

  13. EMF • This potential difference is called the electromotive (causing electron motion) force, or emf. • The emf of a cell, Ecell, is also called the cell potential. • Ecell is measured in volts, and is often referred to as cell voltage. • For any cell reaction that is spontaneous, (voltaic cell) the potential is positive.

  14. Standard Cell Potential • Under standard conditions, (i.e. 25°C), the emf is called the standard emf, or standard cell potential, E°cell.

  15. Standard Reduction Potentials • EMF depends on the particular cathode and anode half-cells involved. • The standard reduction potential, E°red,can be assigned for an individual half-reaction. • The potential is compared to the standard hydrogen electrode, (SHE), which has an E°red = 0.

  16. Standard Reduction Potential The standard hydrogen electrode (SHE) is used as a reference electrode.

  17. Standard Reduction Potentials • Whenever we assign an electrical potential to a half-reaction, we write the reaction as a reduction. • Standard reduction potentials are often called half-cell potentials (Appendix E.) • These can be combined to calculate emfs of a variety of voltaic cells.

  18. For each of the half-cells in a voltaic cell, the standard reduction potential provides a measure of the driving force for the reaction to happen. • The more positive the value of E°red, the greater the driving force for reduction under standard conditions.

  19. Standard cell potential of a voltaic cell. The cell potential measures the difference in the standard reduction potentials of the cathode and anode reactions: E°cell = E°red(cathode) - E°red(anode)

  20. Strengths of Oxidizing and Reducing Agents • The more positive the E°redvalue for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced, and therefore, to oxidize another substance.

  21. Substances that are strong oxidizing agents produce products that are weak reducing agents, and vice versa.

  22. Free Energy and Redox Reactions • E° = E°red(reduction process) - E°red(oxidation process) • A positive value of E indicates a spontaneous process, and a negative value of E indicates a nonspontaneous process.

  23. EMG and ΔG • The relationship between emf and the free-energy change is ΔG = -nFE nis the number of electrons transferred in the reaction F is Faraday’s constant (quantity of charge on 1 mol e-) E is the emf Both n and F are positive values: a positive value of E and a negative value of ΔG both indicate that a reaction is spontaneous.

  24. Faradays • 1 Faraday = 96,485 C/mol. • E > 0 is spontaneous • E < 0 is nonspontaneous

  25. Nernst Equation • The emf of a redox reaction varies with temperature and concentration of reactants and products. • The Nernst equation relates the emf under non-standard conditions to the standard emf and the reaction quotient, Q. • E = E° - (RT/nF)lnQ = E° - (0.0592/n) log Q

  26. Batteries • A battery is a self-contained electrochemical power source that contains one or more voltaic cells.

  27. Lead-Acid Battery • A 12-V lead-acid car battery consists of voltaic cells in series, each producing 2 V. • The standard cell potential can be applied: • E°cell = E°red (cathode) - E°red (anode) • = (+1.685 V) – (-0.356 V) = +2.041V

  28. When batteries are connected in series, the total emf is the sum of the individual emfs.

  29. Alkaline Batteries • These are the most common nonreachargable batteries. • The anode is powdered Zn in gel that is in contact with a concentrated solution of KOH.

  30. Corrosion • The corrosion of iron is electrochemical in nature, involving oxidation and reduction, while the metal conducts electricity.

  31. Preventing Corrosion • Covering the surface with paint prevents the surface from coming into contact with H2O and O2. • Galvanized iron involves coating the iron with a thin layer of zinc. • Protecting a metal from corrosion is called cathodic protection.

  32. Cathodic Protection

  33. Electrolysis • Electrolysis is conducted in an electrolytic cell. • This process requires an energy supply to provide the energy needed to force the electrons to go in the opposite direction. (It makes a nonspontaneous reaction happen.)

  34. Electroplating • Since electrolysis pumps electrons in the direction that doesn’t happen spontaneously, it is used for electroplating. • Electroplating uses electrolysis to deposit a thin layer of one metal onto the surface of another. (i.e. silverplating.)

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