The halogens oxidising ability header image.

Oxidising ability of the halogens

The halogens in their chemical reactions are generally good oxidising agents. That is they are electron acceptors, they oxidise other substances and the halogens are themselves reduced when they gain electrons. Displacement reactions using halogens are a good way to demonstrate the trend in the oxidising ability of the halogens. Fluorine is the most powerful oxidising agent in group 7 and the trend in ability of the halogens to accept electrons decreases as we descend group 7.

To explain why the oxidising ability of the halogens decrease as we descend group 7 we need to consider a range of factors. The oxidiation reaction is simple enough, we can show it as:

½X2 + e → X-
where X is any halogen.

The factors we need to think about which will influence how good an oxidising agent a halogen is are:

Halogen F2 Cl2 Br2 I2
Bond enthalpy/kJmol-1 158 243 192 151

The bond enthalpies for the halogens drop as we descend the group; however this pattern does not hold for fluorine, its bond is much weaker than might be expected. This is a consequence of its small size; being so small its non-bonding electrons in the p-orbitals on each fluorine atom are relatively close to each other and this causes some repulsion between them which weakens the bond.

  • The ability of an atom to attract an electron from another substance will depend upon the size (atomic radius) of the halogen atom, the size of the nuclear charge and the amount of shielding taking place. The electron affinity will give us a measure of an atoms ability to attract an electron. The stronger the attraction between any incoming electron and the nucleus the more energy will be released. The electron affinity is the amount of energy released by the following process:
  • X(g) + e → X-(g)
    The table below lists the energy changes when 1 mole of electrons are gained by each halogen to form a halide ion.
    Halogen F Cl Br I
    electron affinity/kJmol-1 -328 -349 -324 -295
    Diagram to show what is meant by shielding in an atom. We can see that there is not a particularly large difference in the electron affinities when we go from one halogen to the next. The increasing nuclear charge as we decend group 7 is offset by increased screening of the nucleus and an increase in the atomic radius. Ultimately the effective nuclear charge that an electron will feel will probably be about the same for each halogen atom due to these competing factors; for example fluorine (1s22s22p5) has 9 electrons; 4 of which are in the lower 1s and 2s sub-levels, if we assume that the 5 electrons in the outer sub-level shield the nucleus poorly then any incoming electron will feel a nuclear charge of +5 (9 protons - 4 shielding electrons= +5). This effective nuclear charge is the same for all the halogens, but of course as we descend the group the atomic radius increases so any incoming electron will be much further away and feel a much smaller attraction to the effective nuclear charge.

    However fluorine does not fit the pattern here, its electron affinity is lower than might be expected by looking at the pattern for the other halogens. This is simply because the fluorine atom is so small and any added electron will have to go into the lower 2p orbitals, which are small when compared to the 3p, 4p and 5p orbitals of the other halogens. Being small means that there will be extra repulsion between the electrons and this accounts for a lower electron affinity for fluorine.

    So the trend is oxidising power of the halogens is a balance between all the factors listed above. Ultimately the weak F-F bond and the fact that more energy is released when it enters solution or an ionic lattice means that fluorine is the strongest oxidising of the halogens and the ability of the halogens to act as oxidising agents decreases as we descend group 7.

    Halogen displacement reactions

    chlorine water displacing iodide from solution of sodium iodide- halogen displacement reaction.

    Displacement reactions are a good way to show the ability of halogens to act as oxidising agents. A typical displacement reaction is shown opposite. In the first test-tube we have a solution of sodium iodide dissolved in water. On top of this is added a few centimetres of an organic solvent such as hexane or cyclohexane. Cyclohexane is a very good solvent for halogens and given the choice between dissolving in water and dissolving in cyclohexane a halogen will always dissolve in cyclohexane before water. Cyclohexane does not dissolve in water but like oil simply floats on top of it which is why there are two layers shown in the first test-tube in the image opposite.

    Sodium iodide being an ionic compound will dissolve in water. This solution contains sodium ions (Na+) and iodide ions (I - ).

    When chlorine water is added to this test-tube and shaken for around 30 seconds you can see in the image opposite that the cyclohexane layer has changed colour; it has turned a violet/purple colour.
    It may be easier to understand what is happening here if we write an equation for the reaction:

    sodium iodide(aq) + chlorine(aq) → sodium chloride(aq) + iodine(aq)
    2NaI(aq) + Cl2(aq) → 2NaCl(aq) + I2(aq)

    On the reactants side of the word equation we have two halogens present; iodine in the form of iodide ions in the solution and chlorine from the chlorine water. In a displacement reaction the more reactive halogen will oxidise the less reactive halide ion present in a solution or a compound, in this case chlorine will displace or oxidise the iodide ion from the sodium iodide solution. The chlorine effectively takes the place of the iodide ion. The iodine is kicked out of the solution and now has a choice of places to go to; either into the cyclohexane layer or into aqueous solution; it dissolves into the cyclohexane layer since halogens are more soluble in organic solvents than in water.

    Another advantage of cyclohexane is that the halogens show up as bright clear colours when dissoloved organic solvents, iodine is purple, chlorine is green/yellow and bromine is red-brown. So the purple colour of the cyclohexane is due to the iodine dissolved in it!

    Displacing bromine from a solution

    In the example shown below we again have 2 halogens on the reactants side of the equation, these are bromine in the form of a bromide ion from the sodium bromide solution and chlorine from the chlorine water. Chlorine is more reactive than bromine so will displace or oxidise it from the solution. The bromide will be kicked out of solution and will dissolve in the cyclohexane solvent turning it red/brown.

    A halogen displacement reaction.  Chlorine will displace bromide ions from a sosdium bromide solution.

    We can show this as:

    sodium bromide(aq) + chlorine(aq) → sodium chloride(aq) + bromine(aq)
    2NaBr(aq) + Cl2(aq) → 2NaCl(aq) + Br2(aq)
    In this equation the chlorine has been reduced to chloride and the bromide ion has been oxidised to bromine. Chlorine is the electron acceptor or oxidising agent and the bromide is the electron donor or reducing agent.

    Displacement reactions????

    However be careful as you can get caught out if you are not careful! Consider the reaction shown below.

    sodium chloride(aq) + iodine(aq) → sodium chloride(aq) + iodine(aq)
    2NaCl(aq) + I2(aq) → 2NaCl(aq) + I2(aq)
    Here we have iodine and chlorine as the two halogens present on the reactant side of the equation. Iodine is not able to oxidise the chloride ions present and so no reaction will happen.

    Reaction of the halogens with iron(II) chloride solutions

    Colours of iron(II) and iron(III) chloride solutions.

    Displacement reactions are one way to show the oxidising ability of the halogens, another method is the reactions of the halogens with solutions containing the iron ions Fe2+ and Fe3+ions. Iron (II) chloride solution is a pale green colour while iron (III) chloride solution is a pale yellow colour; as shown in the image opposite.

    We can demonstrate the oxidising ability of the halogens by reacting them with a solution of Iron(II) chloride. Here chlorine and bromine are strong enough oxidising agents to oxidise the Fe2+ ions to Fe3+ ions and in the process the chlorine and bromine are both reduced to chloride and bromide ions. We can show this as:

    2Fe2+(aq) + Cl2(aq) → 2Fe3+(aq) + 2Cl-(aq)
    and also:
    2Fe2+(aq) + Br2(aq) → 2Fe3+(aq) + 2Br-(aq)
    The pale green solution containing the Fe2+ ions will change colour to a pale yellow solution containing Fe3+ when chlorine and bromine are added to it however this is not what happens with iodine. Iodine being the poorest oxidising agent in group 7 is not able to oxidise the Fe2+ ions; in fact a solution of Fe3+ will oxidise iodide ions (I-) to form iodine (I2)
    2I-(aq) + 2Fe3+(aq)I2(s) + 2Fe2+(aq)
    Here the soluble iodide ions (I-(aq)) are oxidised to form the almost insoluble iodine. A pale brown solution of iodine in water will form along with a solid precipitate of insoluble iodine on the bottom of the beaker or test-tube.

    Key Points

    Practice questions

    Check your understanding - Questions on halogen displacement reactions.