The Halogens (OCR AS Chemistry A): Revision Note

Physical Properties of the Halogens

Trend in boiling points

• Boiling point increases down the group

• We can explain this trend by looking at the forces between the molecules

• Halogens are non-metals and are diatomic molecules at room temperature

• This means that they exist as molecules that are made up of two similar atoms, such as F₂

• The halogens are simple molecular structures with weak London forces between the diatomic molecules (induced dipole–dipole interactions)

  

The diagram shows that a sudden imbalance of electrons in a nonpolar molecule can cause an instantaneous dipole. When this molecule gets close to another non-polar molecule it can induce a dipole as the cloud of electrons repel the electrons in the neighbouring molecule to the other side

• The more electrons there are in a molecule, the greater the induced dipole–dipole interactions

• Therefore, the larger the molecule the stronger the London forces between molecules

• This is why as you go down the group, it gets more difficult to separate the molecules and the  boiling point increases

Reactivity of the Halogens

Reaction of the halogens with halide ions in displacement reactions A halogen displacement occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide The reactivity of Group 7 non-metals increases as you move up the group Out of the 3 halogens, chlorine, bromine and iodine, chlorine is the most reactive and iodine is the least reactive

Observations in aqueous solution

• Observations when each halogen solution (aq) is added to each potassium halide solution (aq):

  • Chlorine with potassium chloride: very pale green solution, no reaction

  • Chlorine with potassium bromide: yellow solution, chlorine has displaced bromine

  • Chlorine with potassium iodide: brown solution, chlorine has displaced iodine

  • Bromine with potassium chloride: yellow solution, no reaction

  • Bromine with potassium bromide: yellow solution, no reaction

  • Bromine with potassium iodide: brown solution, bromine has displaced iodine

  • Iodine with potassium chloride, bromide or iodide: brown solution, no reaction

• The colour of the solution in the test tube shows which free halogen is present in solution.

  • Chlorine = very pale green solution (often colourless),

  • Bromine = yellow solution

  • Iodine = brown solution(sometimes black solid present)

• The colour of the organic solvent layer in the test tube shows which free halogen is present in solution

  • Chlorine = colourless

  • Bromine = yellow

  • Iodine = purple

Observations with an organic solvent

• Colour of the organic solvent layer when each halogen solution is added to each potassium halide solution:

  • Chlorine with potassium chloride: colourless, no reaction

  • Chlorine with potassium bromide: yellow, chlorine has displaced bromine

  • Chlorine with potassium iodide: purple, chlorine has displaced iodine

  • Bromine with potassium chloride or bromide: yellow, no reaction

  • Bromine with potassium iodide: purple, bromine has displaced iodine

  • Iodine with any potassium halide: purple, no reaction

Reaction of bromide ions and chlorine

• Chlorine is above bromine in Group 7 so it is more reactive

• Chlorine will therefore displace bromine from an aqueous solution of a metal bromide

    2KBr (aq) + Cl₂ (aq) → 2KCl (aq) + Br₂ (aq)

We can again see this is a redox reaction by taking a look at changes in the oxidation number of each element in the reaction Br = -1 → 0 so the bromine has been oxidised Cl = 0 → -1 so the chlorine has been reduced No change in oxidation number for the potassium Rather than writing the full equation we can also write the ionic equation by removing the potassium spectator ion

2Br^- (aq) + Cl₂ (aq) → 2Cl^- (aq) + Br₂ (aq)

Trend in Reactivity

• The oxidising power of the halogens decreases going down the group (the halogens get less reactive)

• This can be explained as follows:

  • Atomic radius increases

  • There are more inner shells so shielding increases

  • There is less nuclear attraction to attract an electron from another species

Redox Reactions of the Halogens

• Each halogen has 7 electrons in their outer shell with two electrons in the outer s sub-shell and 5 in the outer p sub-shell

  • __s²__p⁵

• Halogens react with metals by accepting an electron from the metal atom to become an ion with 1- charge

  • For example:

Ca (s) + Cl₂ (g) → CaCl₂ (s)  

consisting of Ca²⁺ and 2Cl^- ions

• Halogens are therefore oxidising agents:

  • Halogens oxidise the metal by removing an electron from the metal (the oxidation number of the metal increases)

  • Halogens become reduced as they gain an extra electron from the metal atom (the oxidation number of the halogen decreases)