half-cell

Examples of half-cell in the following topics:

• Electrochemical Cell Notation

  • Cell notation is shorthand that expresses a certain reaction in an electrochemical cell.
  • The anode half-cell is described first; the cathode half-cell follows.
  • Within a given half-cell, the reactants are specified first and the products last.
  • A double vertical line ( || ) represents a salt bridge or porous membrane separating the individual half-cells.
  • A typical arrangement of half-cells linked to form a galvanic 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 redox reaction consists of two half-reactions.
  • 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 if a Metal Will Dissolve in Acid

  • Each half-cell is associated with a potential difference whose magnitude depends on the nature of the particular electrode reaction and on the concentrations of the dissolved species.
  • In order to express them in a uniform way, we follow the rule that half-cell potentials are always defined for the reduction direction.
  • Therefore, the half-cell potential for the Zn/Zn2+ electrode always refers to the reduction reaction:
  • For this reason, the potential difference contributed by the left half-cell has the opposite sign to its conventional reduction half-cell potential.
  • Set up the oxidation and reduction half-reactions with their cell potential:

• Concentration of Cells

  • The standard potential of an electrochemical cell requires standard conditions for all of the reactants.
  • The change in Gibbs free energy for an electrochemical cell can be related to the cell potential.
  • The Nernst equation can be used to calculate the output voltage changes in a pair of half-cells under non-standard conditions.
  • 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

• 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:

• Free Energy and Cell Potential

  • 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.
  • If E°cell > 0, then the process is spontaneous (galvanic cell)
  • If E°cell < 0, then the process is non-spontaneous (the voltage must be supplied, as in an electrolytic cell)
  • A demonstration electrochemical cell setup resembling the Daniell cell.
  • The two half-cells are linked by a salt bridge carrying ions between them.

• Equilibrium Constant and Cell Potential

  • In electrochemistry, the Nernst equation can be used, in conjunction with other information, to determine the equilibrium reduction potential of a half-cell.
  • For example, let's say a concentration gradient was established by dissolving KCl in half of a divided vessel that was originally full of H2O.
  • In order to calculate the standard potential, we have to look up the half-reactions of copper and zinc.
  • The cell equilibrium constant, K, can be derived from the Nernst equation:
  • Schematic of a galvanic cell for the reaction between Zn and Cu.

• Dry Cell Battery

  • The dry cell is one of many general types of electrochemical cells.
  • Unlike a wet cell, a dry cell can operate in any orientation without spilling, as it contains no free liquid.
  • A common dry-cell battery is the zinc-carbon battery, which uses a cell that is sometimes called the Leclanché cell.
  • The paste of ammonium chloride reacts according to the following half-reaction:
  • An illustration of a zinc-carbon dry cell.

• Predicting Spontaneous Direction of a Redox Reaction

  • This means that Li would be written as the reduction half-reaction when compared to any other element in this table.
  • On the other hand, Fe would be written as the oxidation half-reaction when compared to any other element on this table.
  • For example if we turn the zinc oxidation half-reaction around ($Zn^{2+} + 2e^- \rightarrow Zn \ E^o = -0.76 V$), the cell potential is reversed.
  • The relative reactivities of different half-reactions can be compared to predict the direction of electron flow.
  • Predict the direction of electron flow in a redox reaction given the reduction potentials of the two half-reactions

• Phospholipids

  • Phospholipids are the main constituents of cell membranes.
  • The phospholipid molecules can move about in their half the bilayer, but there is a significant energy barrier preventing migration to the other side of the bilayer.
  • A cell may be considered a very complex liposome.
  • The interior of a cell contains a variety of structures (organelles) that conduct chemical operations vital to the cells existence.
  • Molecules bonded to the surfaces of cells serve to identify specific cells and facilitate interaction with external chemical entities.