12.12Reactions of  Group VII Elements

Halogens as oxidising agent

• Halogens are powerful oxidising agents. However, the oxidising ability decreases down the Therefore, F2 is the most powerful oxidising agent while I2 is the weakest.
• This is reflected in their ability to oxidise other halide ions, as follow:

Note:

1. A halogen can oxidise the halide ion below it on the Periodic Table, fluorine is excluded in this argument because it is too powerful as an oxidising agent and will oxidise water into oxygen
2. If chlorine is able to displace bromide ion from its aqueous solution, this indicates that chlorine has a higher tendency to be reduced and to accept electrons to form ions

 

Reaction with hydrogen gas, H2

• All halogens react with hydrogen gas to form hydrides, HX.

H2 + X2 → 2HX   ; where X = A halogen

 

• H2(g) + F2(g) → 2HF(g) ; explosive reaction under all temperature
  H2(g) + Cl2(g) → 2HCl(g) ; explosive reaction under sunlight
  H2(g) + Br2(g) → 2HBr(g) ; slow reaction on heating
  H2(g) + I2(g) ⇌ 2HI(g) ; an equilibrium mixture is obtained

• The reactivity of halogens towards hydrogen gas decreases down the Group due to the decrease in oxidising ability of the halogens

 

Thermal stability of  hydrogen halides, HX

• The thermal stability of the hydrogen halides, HX decreases down the This is because the size of the atom increases and so the strength of the H-X bond decreases. In other words, the hydrogen halides become less stable on heating going down the Group.
• i) Hydrogen iodide decomposes easily on heating, thick purple fumes of I2 are observed.
  ii) Hydrogen bromide decomposes slightly, little orange-brown of Br2 is observed.
  iii) Hydrogen chloride and fluoride are stable on heating
• When hydrogen halides decompose, X⁻ ions are oxidised. The ease of oxidation down the Group indicates the increase in reducing ability of X⁻ Hence, I⁻ is the strongest reducing agent while F⁻ is the weakest.

 

Reaction with aqueous sodium thiosulfate, Na2S2O3

• Chlorine and bromine can oxidise sodium thiosulfate, Na2S2O3 to sodium sulfate, Na2SO4. The oxidation number of sulfur changes from +2 to  +6.

4Cl2 + S2O3²⁻ + 5H2O → 2SO4²⁻ + 10H⁺ + 8Cl⁻

4Br2 + S2O3²⁻ + 5H2O → 2SO4²⁻ + 10H⁺ + 8Cl⁻

 

• However, iodine can only oxidise sodium thiosulfate to sodium tetrathionate, Na2S4O6. The oxidation number of sulfur changes from +2 to +2.5.

I2 + 2S2O3²⁻ → S4O6²⁻ + 2I⁻

 

 

Reaction with aqueous iron(II) ions, Fe²⁺

1) Chlorine and bromine would oxidise Fe²⁺ to Fe³⁺ but not iodine.

Cl2 + 2Fe²⁺ → 2Cl⁻ + 2Fe³⁺

Br2 + 2Fe²⁺ → 2Br⁻ + 2Fe³⁺

 

Reaction with iron, Fe

1) i. When chlorine gas is passed over hot iron, iron(III) chloride is formed. The oxidation number of  iron changes from 0 to +3.

Cl2 + Fe → FeCl3         ; rapid and vigorous reaction

ii. When bromine vapour is passed over hot iron, iron(III) chloride is formed. The oxidation number of iron changes from 0 to  +3.

Br2 + Fe → FeBr3         ; less rapid and vigorous reaction

 

iii. When iodine vapour is passed over hot iron, iron(II) chloride is formed. The oxidation number of iron changes from 0 to only +2.

I2 + Fe → FeI2    ; even less vigorous

 

Reaction with hot and cold alkali

• Chlorine undergoes disproportionation when it reacts with In this reaction, chlorine is simultaneously oxidised and  reduced.
• In cold alkali(15 ºC), the reaction is as follow:

Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)

The ionic equation is:

Cl2(aq) + 2OH⁻(aq) → Cl⁻(aq) + ClO⁻(aq) + H2O(l)

 

• In cold alkali, the oxidation number of chlorine changes from 0 in Cl2 to -1 in Cl⁻(reduction) and +1 in ClO⁻(oxidation).
• In hot alkali(70 ºC), the reaction is as follow:

3Cl2(aq) + 6NaOH(aq) → 5NaCl(aq) + NaClO3(aq) + 3H2O(l)
The ionic equation is:

3Cl2(aq) + 6OH⁻(aq) → 5Cl⁻(aq) + ClO3⁻(aq) + 3H2O(l)

 

• In hot alkali, the oxidation number of chlorine changes from 0 in Cl2 to -1 in Cl⁻(reduction) and +5 in ClO3⁻(oxidation).
• This reaction is the result of disproportionation of chlorate(I) ions in the presence of heat

3ClO⁻(aq) → 2Cl⁻(aq) + ClO3⁻

 

• Bromine and iodine react in a similar However, the bromate(I) and iodate(I) ions formed disproportionate readily at all temperatures.