Electrode potentials

• A voltaic cell converts chemical energy into electrical energy
• As electrical energy results from movement of electrons, you need chemical reactions that transfer electrons from one species to another. These are redox reactions.
• A half-cell contains the chemical species present in a redox half-equation
• A voltaic cell can be made by connecting together two different half-cells, which then allows electrons to flow
• In a voltaic cell, the chemicals in the two half-cells must be kept apart - if allowed to mix, electrons would flow in an uncontrolled way and heat energy would be released rather than electrical energy
• 2 types of half cells

  • metal/metal ion
  • ion/ion
• The simplest half-cell is the metal/metal ion type which consists of a metal rod dipped into a solution of its aqueous metal ion
• The metal/metal ion half cell is represented using a vertical line for the phase boundary between the aqueous solution and the metal
• At the phase boundary where the metal is in contact with its ions, an equilibrium will be set up
• The equilibrium in a half-cell is written so that the forward reaction shows reduction and the reverse reaction shows oxidation
• In an isolated half-cell, there is no net transfer of electrons either into or out of the metal
• When two half-cells are connected, the direction of electron flow depends upon the relative tendency of each electrode to release electrons
• An ion/ion half-cell contains ions of the same element in different oxidation states
• An example of a ion/ion half-cell can be made containing a mixture of aqueous iron(II) and iron(III) ions. The redox equilibrium is:

  Fe 3+(aq) + e = Fe 2+(aq)
• In ion/ion half-cells there are no metal to transport electrons either into or out of the half-cell, so an inert metal electrode made out of platinum is used
• How do you know which electrode has a greater tendency to gain or lose electrons?

  In a cell with two metal/metal ion half-cells connected, the more reactive metal releases electrons more readily and is oxidised
• In an operating cell:

  • the electrode with more reactive metal loses electrons and is oxidised - this is the negative electrode
  • the electrode with the less reactive metal gains electrons and is reduced - this is the positive electrode
• The tendency to be reduced and gain electrons is measured as a standard electrode potential
• The standard chosen is a half-cell containing hydrogen gas and a solution containing H+ (aq) ions
• In the standard half-call an inert platinum electrode is used to allow electrons in and out
• The standard conditions used for a standard hydrogen electrode are:

  • concentration of solution = 1 mol dm -3
  • temperature = 298K (25 °C)
  • pressure = 100 kPa (1 bar)
• The standard electrode potential of a standard hydrogen electrode is exactly 0 V
• The sign of a standard electrode potential shows the sign of the half-cell connected to the standard hydrogen electrode and shows the relative tendency to gain electrons compared with the hydrogen half-cell
• To measure a standard electrode potential, the half-cell is connected to a standard hydrogen electrode by a wire to allow a controlled flow of electrons
• Two solutions from the half-cell and the standard hydrogen electrode are connected with a salt bridge which allows ions to flow
• The salt bridge typically contains a concentrated solution of an electrolyte that does not react with either solution
• An example of a salt bridge is a strip of filter paper soaked in aqueous potassium nitrate, KNO3 (aq)
• The more negative the electrode potential (E) value:

  • the greater the tendency to lose electrons and undergo oxidation.
  • the less the tendency to gain electrons and undergo reduction
• The more positive the E® value:

  • the greater the tendency to gain electrons and undergo reduction
  • the less the tendency to lose electrons and undergo oxidation
• Metals tend to have negative E values and lose electrons
• Non-metals tend to have positive E values and gain electrons
• The more negative the E° value, the greater the reactivity of a metal in losing electrons
• The more positive the E° value, the greater the reactivity of a non-metal in gaining electrons
• Standard cell potential is recorded with a voltmeter
• The cell potential is the potential difference between the two half-cells
• Half-cell: A strip of metal in a solution of its own ions 
• Battery: Two or more cells connected together
• When the two half-cells are connected, electrons flow

  • away from the reducing agent (more negatively charged)
  • towards the oxidising agent (more positively charged)
• Another term for “potential difference” is voltage
• The actual voltage measured between two electrodes is the potential difference