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Electrochemistry

Electrochemistry. Electrochemical Cells/ Chemical Cells. Also called voltaic or galvanic cells A redox reaction produces electricity Occurs spontaneously. Electrochemical Cell. Half Cells. Each ½ of the redox reaction occurs in a separate container One for oxidation and one for reduction

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Electrochemistry

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  1. Electrochemistry

  2. Electrochemical Cells/Chemical Cells • Also called voltaic or galvanic cells • A redox reaction produces electricity • Occurs spontaneously

  3. Electrochemical Cell

  4. Half Cells • Each ½ of the redox reaction occurs in a separate container • One for oxidation and one for reduction • They are connected by a salt bridge • Salt Bridge: allows ions to flow between the two cells

  5. Electrodes • Metals which provide a surface for oxidation or reduction to occur • Solids • Oxidation Number = 0 • Anode • Cathode

  6. ANODE • Oxidation occurs at the anode • Negative electrode • CATHODE • Reduction occurs at the cathode • Positive electrode Red Cat – An Ox Reduction at the Cathode Oxidation at the Anode

  7. Flow of Electrons • The electrodes are connected by a wire • Electrons flow from the anode to the cathode through the wire

  8. Why does the cell produce electricity? • There is a difference of electric potential between the two electrodes • Electrons will flow between the two electrodes until equilibrium is reached • At equilibrium the cell’s voltage would be zero

  9. Zn Zn2+ + 2e- Electrons needed here for reduction Electrons released here by oxidation Zn + CuSO4 Cu + ZnSO4 Red Cat -reduction takes place…electrons are gained. An Ox -oxidation takes place…electrons are lost. Cu2+ + 2e - Cu0 e- e- e- e- e- - e- + e- e- e- e- e-

  10. Batteries • Use a redox reaction which produces electricity spontaneously • Batteries are recharged by reversing the reaction • Dry Cell (Acid or Alkaline), Lead Storage (Car), Rechargeable (Ni/Cd)

  11. Corrosion • Oxidation of a metal • Metal combines with element (usually oxygen) Example: 4Fe + O2 2Fe2O3 (rust)

  12. Prevention of Rust • Cover the metal – paint, oil, another (more reactive) metal • Cathodic Prevention • metal is placed in contact with a more reactive metal • That metal will be oxidized (acts as the anode), the original metal acts as the cathode • Alloys – mixture of metals • Brass, stainless steel (Fe + Cr), cast iron (C + Si)

  13. Electrolytic Cells • Also called electrolysis • An electric current is used to produce a chemical reaction • An electric current is used to force a non-spontaneous reaction to occur

  14. Oxidation occurs at the anode • Reduction occurs at the cathode • Electrons flow from anode to cathode • The cathode is the negative electrode • The anode is the positive electrode • This is opposite of the chemical cell because the external current causes the polarities to switch

  15. Electroplating • Object to be plated is the CATHODE, negative • Metal to be plated onto the object is the ANODE, positive • Solution must contain ions of the metal to be plated

  16. Silver Plating • Cathode = • Anode = • Solution = • What happens to the mass of each electrode during the reaction?

  17. Electrolysis of Water 2 H2O  2 H2 + O2 • The H+ is reduced at the (-) cathode, producing H2 (g), which is trapped in the tube • The O2 is oxidized at the (+) anode, yielding O2 (g), which is trapped in the tube

  18. Hydrogen Fuel Cells • Uses hydrogen gas as the fuel 2 H2 + O2  2 H2O

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