Standard Electrode Potentials

When a piece of metal foil is inserted into a solution of the ions of the metal, an equilibrium is set up. Taking zinc as an example:

This equilibrium obviously tends to a particular position, and therefore a potential difference is produced between the electrolyte and the electrode. It is not possible to measure the absolute potential of an electrode, as it is necessary to make an electrical connection to the solution, and therefore another cell has to be set up. It is therefore necessary to produce a cell, which has 2 electrodes and 2 electrolytes. It is also necessary to measure a potential using standard conditions, as otherwise it has little significance in comparison with other values. The standard conditions chosen are a temperature of 25C (298 K), a pressure of 1 atmosphere and solutions of 1 mol dm-3 activity. The cell is set up using 2 electrodes in their respective solutions. The electrodes are wired together using a high-resistance voltmeter (e.g. a valve voltmeter) and the electrolytes using a salt bridge. This is typically a piece of filter paper soaked in a saturated solution of potassium nitrate, which allows conduction of electrons between the two solutions without producing another pair of electrode / electrolyte cells. However, this measures the sum of the voltages of the pair of cells.

A cell can be represented by a cell diagram. These are written in the following form (for the cell shown above):

The solid line represents the phase boundary between the solid metal and the aqueous ions. The dashed line represents the salt bridge dividing the two cells. The reduced form of the substance is written in the centre of the diagram, next to the phase boundary. The potential of the cell is measured using a high resistance voltmeter (as when current is drawn from the cell the potential decreases). The maximum pd of the cell (when no current is flowing) is called the EMF (electromotive force) of the cell.. The voltage measured is represented by the notation Ecell, or for a cell potential measured under standard conditions, . The potential measured is defined as Ecell = Eright - Eleft, where right and left are given by the position on the cell diagram that defines the cell. The positive terminal of the voltmeter should be connected to the right-hand terminal of the cell.

The Standard Hydrogen Electrode

This is the standard electrode with which others are compared. It is defined as contributing 0V to the cell potential. This means that all cell potentials are measured relative to the potential of this cell. The cell is constructed as follows:

The platinum wire coated in platinum black is inert - it does not react with the hydrogen or the acid. It also acts as a catalyst for the hydrogen gas / hydrogen ion reaction. Platinum black is a coating of platinum powder. This increases the surface area of the metal on which the reaction can take place.

This cell is written in a cell diagram as . To measure a standard electrode potential, connect the standard half-cell to the standard hydrogen electrode, which should be connected as the left-hand electrode. Then .

A negative cell potential implies that the cell gives away electrons more readily than the standard hydrogen electrode, and a positive cell potential that the cell is accepts electrons more readily than the standard hydrogen electrode.

Other Redox Systems

It is possible to measure the redox potential of systems which involve two species of the same element in different oxidation states, where both are in solution, e.g. . This is measured using a cell as that shown in the diagram. A salt bridge is used as well as a platinum electrode. This gives a redox potential which can be measured in the same way as before by connecting the cells to the standard hydrogen electrode. To measure a standard cell potential it is necessary, as before, to have all solutions at 1 mol dm-1 activity. These solutions must include all solutions that are included in the reaction between the two species, which often includes H+ ions. The cell diagrams for this type of cell are written: oxidised form, reduced form / Pt, for example .

Uses of Electrode Potentials

  1. To work out the EMF of cells. The Standard Hydrogen Electrode is also on the left.

    Ecell = Eright - Eleft

  2. Predicting the direction of redox reactions.

    Fe3+ + e- e Fe2+EÆ = +0.77V

    I2 + 2e- e 2I-EÆ = +0.54V

    The Fe2+ / Fe3+ system is more positive so it has a greater tendency to use up electrons. The I- / I2 system is more negative so it has a greater tendency to produce electrons. This means that the dominant pair of reactions will be:

    Fe3+ + e- ® Fe2+

    2I- ® I2 + 2e-

    Therefore the overall reaction will be 2Fe3+ + 2I- ® 2Fe2+ + I2

    If the standard electrode potentials for the two half - reactions are too close then the reaction becomes kinetically stable, and can be forced either way by the application of different conditions.

Short cuts for predicting the direction of reactions

  1. Anticlockwise Rule

    Arrange the redox half equations so that:

    The reaction will then take place anticlockwise around the half equations (i.e. the top right reacts with the bottom left).


    I2 + 2e- e 2I- EÆ = +0.54V

    Fe3+ + e- e Fe2+ EÆ = +0.77V

    Therefore the iodide ions will react with the iron (III) ions.

  2. Using a cell diagram

    If Ecell for a given cell is positive then the reaction will take place along the cell diagram. This is reversed if the cell potential is negative.

    e.g. Ecell = +0.23 V

    Therefore iodide will become iodine and iron(III) will become iron(II).

Half Equations

Ionic half equations are equations that include electrons, and balance these with the oxidation or reduction of a species. By balancing the electrons in two half equations and then adding them together it is possible to construct redox equations. The redox potentials for the half equations allow us to predict the direction of a redox reaction.

Potassium manganate (VII) and iodine / thiosulphate titrations

In potassium permanganate titrations, in general the solution in the conical flask is something that can be oxidised (i.e. a reducing agent), for example nitrite (NO2-), sulphite (SO3-), iron (II) and ethanedioate (oxalate, C2O42-). The half equation of the reduction of the permanganate is . This reaction goes from a deep purple solution (manganate (VII)) to a colourless solution (manganate (II)). Thus permanganate can be used without an indicator - it is self indicating. The reaction is autocatalytic - it is catalysed by the manganate (II) ions.

In iodine / thio titrations, the initial reaction is between an oxidising agent (usually quite a weak oxidising agent), such as iron (II), iodate (IO3-), and iodide ions. This reaction liberates iodine, which can then be titrated against thiosulphate via the following reaction: . As the iodine is used up, the colour of the solution fades from dark brown to yellow, and thence to colourless. However, the final point is very difficult to see, and therefore when the reaction has reached a "pale straw" colour, starch is added to form a blue-black complex, which disappears at the endpoint.


Disproportionation is the simultaneous oxidation and reduction of the same species in aqueous solution. A species in an intermediate oxidation state will disproportionate provided that the cell potential for the disproportionation reaction is positive. Therefore, if the electrode potentials are arranged into a list with the largest negative one of the pair at the top, then the top right will react with the bottom left. If these are the species to disproportionate, then the disproportionation will take place, if not it will not. For example:

Ecell = -0.22

Ecell = +0.57

Therefore, the ion will disproportionate to give sulphate and hydrogensulphite.

The Electrochemical Series

The electrochemical series is a list of the elements in order of their standard reduction potentials. This is a measure of reactivity, but the chemical reactivity of an element is not necessarily in line with its position on the ECS - lithium is the element with the most positive electrode potential - more positive than potassium!

Applications of Electrochemistry

Corrosion - Rusting

Rust is hydrated iron (III) oxide (), where n varies according to the degree of hydration (immediately after the reaction, n=2). The reaction occurs in iron that is not totally pure, i.e. where there are areas of impurity in the metal or where the metal has been stressed by working. These areas are called cathodic areas. The reactions that take place are as follows:

At the negative electrode (pure iron)

(presumably as part of an equilibrium)

These iron (II) ions are then oxidised to iron (III) by oxygen in the air and H+ ions in the water:

These electrons then pass to the positive part of the cell, via the metal:

, leaving OH- ions in the water.

Thus the overall reaction is . This forms a brown flaky solid that does not protect the metal underneath from further attack.

The rate of rusting is increased by:

1. Certain impurities providing more cathodic areas.

2. Ions in the water, increasing the conductivity of the electrolyte.

Rust prevention encompasses two main methods:

  1. Sacrificial Methods

    A metal that has a more negative electrode potential than iron will provide electrons in preference to the iron, and therefore corrode first. One example of this is galvanising, or zinc plating. Only when all the zinc has been oxidised does the iron start to rust. A sacrificial anode may also be used - this means attaching a slab of more electropositive metal to the iron in question, which then decays preferentially. This is generally aluminium or zinc, and is often used for large constructions, such as submarines and bridges.

  2. Coating

    This works by excluding water / air from the metal surface. The surface is coated, either by painting or electroplating. However, if the coating is a metal that is more electronegative than iron then the iron will decay preferentially to the coating - once the coating is scratched the iron will decay at a very rapid rate.

Storage Cells

  1. The Lead - Acid Battery

    This is also known as the Lead Accumulator. This battery is used in cars, lorries, etc. It has several disadvantages - it is heavy and bulky and contains sulphuric acid in concentrations that can easily be dangerous. Its principal advantage is that it is rechargeable, and therefore can be used to store quite large amounts of electrical energy

    A single lead-acid cell is shown opposite. The negative electrode is made of lead, and the following reaction takes place:

    At the positive electrode, the following reaction takes place:

    Thus, the overall discharge reaction is The cell provides a potential of about 2V - 6 cells in series are used to produce the 12V used in a car battery,

    The reaction uses up sulphuric acid and produces a coating of lead sulphate on each of the electrodes. As the sulphuric acid is used up and the amount of water increases, the density of the electrolyte drops. Recharging is the opposite reaction - a potential of about 2.2 V needs to be applied across the terminals of a cell to recharge it, or about 13V across the whole battery.

  2. The Zinc - Carbon Cell

    This was the first dry cell, and remained the only dry cell until the advent of the alkaline cell. The negative terminal of the cell is the zinc case and the positive terminal is the graphite rod in the centre. The rod is surrounded by a paste of ammonium chloride, manganese (IV) oxide and graphite. The reactions are:

    The manganese (IV) oxide reacts with the hydrogen to form manganese (III) hydroxide and water, and the ammonia complexes with the zinc ions released.

    2MnO2(s) + 3H2(g) ® 2MnO(OH)(s) + 2H2O(l)

    Zn2+(aq) + 4NH3(aq) ® [Zn(NH3)4]2+(aq)

    The zinc - carbon cell delivers a potential of 1.5 V.

  3. The Nickel-Cadmium Cell

    This is a portable rechargeable cell, with a potential of 1.1 V. It is the same basic shape and size as the zinc - carbon cell. The positive electrode is made of nickel coated with Ni(OH)2, and the negative electrode is made up of Cadmium coated in Cd(OH)2. The electrolyte is aqueous sodium hydroxide. The reactions are:

    Ni(s) + 2OH-(aq) ® Ni(OH)2 + 2e-

    Cd(OH)2 + 2e- ® Cd(s) + 2OH-(aq)

    The cell is recharged by putting a potential through it ht opposite way round - this reverses the reactions shown here.