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Voltaic Cells

Voltaic Cells. Chapter 20. Voltaic Cells. In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released. Also known as electrochemical cells. Voltaic Cells. We can use that energy to do work if we make the electrons flow through an external device.

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Voltaic Cells

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  1. Voltaic Cells Chapter 20

  2. Voltaic Cells • In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released. • Also known as electrochemical cells

  3. Voltaic Cells • We can use that energy to do work if we make the electrons flow through an external device. • We call such a setup a voltaic cell.

  4. Voltaic Cells • A typical cell looks like this. • The oxidation occurs at the anode. • The reduction occurs at the cathode.

  5. Voltaic Cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop.

  6. Voltaic Cells • Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. • Cations move toward the cathode. • Anions move toward the anode (since oxidation is at anode the positive charge attracts the negative anions)

  7. Voltaic Cells • In the cell, then, electrons leave the anode and flow through the wire to the cathode. • As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

  8. Voltaic Cells • As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. • The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

  9. Electromotive Force (emf) • Water only spontaneously flows one way in a waterfall. • Likewise, electrons only spontaneously flow one way in a redox reaction—from higher to lower potential energy.

  10. Electromotive Force (emf) • The potential difference between the anode and cathode in a cell is called the electromotive force (emf). • It is also called the cell potential, and is designated Ecell. • The cell potential is determined by reduction potentials for the half reactions which is the tendency of a substance to gain electrons. • the more positive the reduction potentials, the more likely the reaction will be spontaneous

  11. J C 1 V = 1 Cell Potential Cell potential is measured in volts (V).

  12. Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated. These can be reduced to calculate the standard cell potentials of any redox reaction

  13. Standard Hydrogen Electrode • Their values are referenced to a standard hydrogen electrode (SHE). • By definition, the reduction potential for hydrogen is 0 V: 2 H+(aq, 1M) + 2 e− H2(g, 1 atm)

  14. = Ered (cathode) − Ered (anode)   Ecell  Standard Cell Potentials The cell potential at standard conditions can be found through this equation: Because cell potential is based on the potential energy per unit of charge, it is an intensive property.

  15. Ered = −0.76 V  Ered = +0.34 V  Cell Potentials • For the oxidation in this cell, • For the reduction,

  16. = (anode) (cathode) − Ered Ered   Ecell  Cell Potentials = +0.34 V − (−0.76 V) = +1.10 V

  17. Free Energy G for a redox reaction can be found by using the equation G = −nFE where n is the number of moles of electrons transferred, and F is a constant, the Faraday. 1 F = 96,485 C/mol = 96,485 J/V-mol

  18. Free Energy Under standard conditions, G = −nFE Therefore, if the cell potential is positive the reaction is spontaneous because G will be negative

  19. Real systems • Many real systems do not operate at standard conditions so the E cell potential must account for the effect of concentration for the products and reactants (If E cell has equal concentrations of products and reactants the cell potential = 0 because the system is at equilibrium) • Think of the cell potential as a driving force towards equilibrium so that at equilibrium K=Q and E cell = 0 • The greater the magnitude of the cell potential, the farther the reaction is from equilibrium • In concentration cells, the direction of spontaneous electron flow can be determined by considering the direction needed to reach equilibrium

  20. Ecell  • For such a cell, would be 0, but Q would not. Concentration Cells • a cell could be created that has the same substance at both electrodes. • Therefore, as long as the concentrations are different, E will not be 0.

  21. Faraday’s Law • Faraday’s First Law of electrolysis: the mass of the deposited material in a redox reaction is directly proportional to the quantity of electrical charge (which is the current determined by the number of electrons) that passes through the electrolyte. The longer a current is applied the greater the number of electrons and mass that can be deposited. • The larger the charge of the ionic species, the greater the mass deposited as well

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