galvanic cell

Examples of galvanic cell in the following topics:

• Free Energy and Cell Potential

  • In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, the change in Gibbs free energy must be negative.
  • The basis for an electrochemical cell, such as the galvanic cell, is always a redox reaction that can be broken down into two half-reactions: oxidation occurs at the anode, where there is a loss of electrons, and reduction occurs at the cathode, where there is a gain of electrons.
  • In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, the change in Gibbs free energy must be negative.
  • If E°cell > 0, then the process is spontaneous (galvanic cell)
  • A demonstration electrochemical cell setup resembling the Daniell cell.

• Electrochemical Cell Notation

  • Cell notation is shorthand that expresses a certain reaction in an electrochemical cell.
  • Cell notations are a shorthand description of voltaic or galvanic (spontaneous) cells.
  • The anode half-cell is described first; the cathode half-cell follows.
  • Using these rules, the notation for the cell we put together is:
  • A typical arrangement of half-cells linked to form a galvanic cell.

• Standard Reduction Potentials

  • In this galvanic cell, zinc reduces copper cations.

• Equilibrium Constant and Cell Potential

  • It can also be used to determine the total voltage, or electromotive force, for a full electrochemical cell.
  • The standard cell potential for the reaction is then +0.34 V - (-0.76 V) = +1.10 V.
  • The cell equilibrium constant, K, can be derived from the Nernst equation:
  • Schematic of a galvanic cell for the reaction between Zn and Cu.
  • Calculate the equilibrium constant, K, for a galvanic cell using the Nernst equation

• The Nernst Equation

  • In electrochemistry, the Nernst equation can be used to determine the reduction potential of an electrochemical cell.
  • In electrochemistry, the Nernst equation can be used, in conjunction with other information, to determine the reduction potential of a half-cell in an electrochemical cell.
  • Find the cell potential of a galvanic cell based on the following reduction half-reactions where [Ni2+] = 0.030 M and [Pb2+] = 0.300 M.
  • First, find the electromotive force for the standard cell, which assumes concentrations of 1 M.
  • The added half-reactions with the adjusted E0 cell are:

• Electrolytic Cells

  • Electrolysis uses electrical energy to induce a chemical reaction, which then takes place in an electrolytic cell.
  • Electrolysis can sometimes be thought of as running a non-spontaneous galvanic cell.
  • A cell used in elementary chemical experiments to produce gas as a reaction product and to measure its volume.

• Concentration of Cells

  • The standard potential of an electrochemical cell requires standard conditions for all of the reactants.
  • When reactant concentrations differ from standard conditions, the cell potential will deviate from the standard potential.
  • The change in Gibbs free energy for an electrochemical cell can be related to the cell potential.
  • Under standard conditions, the output of this pair of half-cells is well known.
  • Discuss the implications of the Nernst equation on the electrochemical potential of a cell

• Voltaic Cells

  • A voltaic cell is a device that produces an electric current from energy released by a spontaneous redox reaction in two half-cells.
  • This kind of cell includes the galvanic, or voltaic, cell, named after Luigi Galvani and Alessandro Volta.
  • In a typical voltaic cell, the redox pair is copper and zinc, represented in the following half-cell reactions:
  • Each half-cell is connected by a salt bridge, which allows for the free transport of ionic species between the two cells.
  • The cell consists of two half-cells connected via a salt bridge or permeable membrane.

• Predicting the Products of Electrolysis

  • Electrolysis is very important commercially as a stage in the separation of elements from naturally occurring sources, such as ores, using an electrolytic cell.
  • You may have noticed that this is the opposite of a galvanic cell, where the anode is negative and the cathode is positive.

• Preventing Corrosion

  • The corrosion-prone iron alloy steel is commonly coated with zinc, a more active metal, in a process known as galvanizing.
  • Galvanic sacrificial anode attached to the hull of a ship; here, the sacrificial anode shows corrosion but the metal it is attached to does not.
  • The galvanic anode continues to corrode, consuming the anode material until eventually it must be replaced, but the cathodic material is protected.
  • Protecting iron alloys with a coating of a more active metal through the process of galvanizing prevents the alloys from corroding.