A-Level Chemistry (Edexcel) – Reactions of the group 7 elements

4.4 - Reactions of the Group 7 Elements

1How halogens gain electrons to become ions
2Trends in reactivity and oxidising ability down the group
3Reactions of halogens with group 1 and 2 metals
4Displacement reactions between halogens and halide ions
5Identification of halogens based on colour changes in displacement reactions

5 min

Halogens form 1- ions by gaining electrons

Halogens react by gaining an electron to form negative ions with a 1- charge (known as halide ions). For example:

X + e⁻ ➔ X⁻

Where X represents a halogen atom.

This electron gain results in the halogen being reduced, as its oxidation number decreases from 0 to -1.

As the halogen gains an electron, it causes another substance to be oxidised - so halogens act as oxidising agents

Reactivity decreases down group 7

The reactivity of the halogens decreases going down the group because:

Atomic radius increases down the group as more electron shells are added.
The increasing size leads to the outer electrons being farther from the positive nucleus.
The outer electrons also experience more shielding from inner electron shells.
The electrostatic attraction between the outer electrons and nucleus gets progressively weaker.
The increase in atomic radius and shielding outweigh the increase in nuclear charge so it becomes harder for larger halogens to attract the electron needed to form a negative ion.

Oxidising power of halogens decreases down group 7

The oxidising power of halogens refers to their ability to take electrons from other substances to form halide ions. Oxidising power is directly related to reactivity.

The oxidising power of halogens decreases down group 7, as does reactivity. This is because the increasing size and shielding makes it harder for larger halogens to remove electrons from (i.e. oxidise) other substances.

Therefore, fluorine is the strongest oxidising agent and iodine is the weakest.

Halogens displace less reactive halide ions

The relative reactivity and oxidising ability of the halogens can be seen in displacement reactions.

If an aqueous halogen solution is added to a solution containing halide ions, a more reactive halogen will displace a less reactive halide from the solution.

The displacing halogen is reduced as it gains an electron to form the halide ion.
The displaced halide is oxidised as it loses an electron to form the halogen molecule. The rule is: a halogen will displace any halide ion below it in group 7.

Here is a summary table including the ionic equations for the displacement reactions:

Halogen	Displaces	Ionic equation(s)
Chlorine (Cl $_{2}$ )	Bromide (Br $^{-}$ ) and Iodide (I $^{-}$ )	Cl $_{2 (a q)}$ + 2Br $^{-}$ $_{(a q)}$ ➔ 2Cl $^{-}$ $_{(a q)}$ + Br $_{2 (a q)}$ / Cl $_{2 (a q)}$ + 2I $^{-}$ $_{(a q)}$ ➔ 2Cl $^{-}$ $_{(a q)}$ + I $_{2 (a q)}$
Bromine (Br $_{2}$ )	Iodide (I $^{-}$ )	Br $_{2 (a q)}$ + 2I $^{-}$ $_{(a q)}$ ➔ 2Br $^{-}$ $_{(a q)}$ + I $_{2 (a q)}$
Iodine (I $_{2}$ )	None	No reaction

This table shows that oxidising ability decreases down group 7 because:

Chlorine can oxidise both bromide and iodide, showing it is the strongest oxidising agent.
Bromine can oxidise iodide but not the more reactive chloride, so bromine has intermediate oxidising strength.
Iodine cannot oxidise either chloride or bromide, so it has the weakest oxidising power.

Reactions of halogens with group 1 and 2 metals

The halogens also react with group 1 and group 2 metals, forming the corresponding halide salts. The reactivity of the halogens in these reactions decreases as you move down group 7, following the same trend seen previously.

Reactions with group 1 metals:

Group 1 metals react with halogens in a 2:1 ratio. For example:

2Na_(s) + F_2(g) ➔ 2NaF_(s)

In this reaction:

Sodium is oxidised: Na ➔ Na⁺ + e^- (oxidation state changes from 0 to +1).
Fluorine is reduced: F₂ + 2e^- ➔ 2F^- (oxidation state changes from 0 to -1).

Reactions with group 2 metals:

Group 2 metals react with halogens in a 1:1 ratio. For example:

Mg_(s) + Cl_2(g) ➔ MgCl_2(s)

In this reaction:

Magnesium is oxidised: Mg ➔ Mg²⁺ + 2e^- (oxidation state changes from 0 to +2).
Chlorine is reduced: Cl₂ + 2e^- ➔ 2Cl^- (oxidation state changes from 0 to -1).

Colour changes in halogen displacement reactions

When a more reactive halogen displaces a less reactive halide ion in solution, visible colour changes occur that allow the halogen product to be identified.

Colour changes in aqueous solution

When halogens displace less reactive halide ions in aqueous solution, visible colour changes occur if a reaction takes place:

If bromide (Br^-) is displaced, forming bromine (Br₂), the aqueous solution turns yellow.
If iodide (I^-) is displaced, forming iodine (I₂), the aqueous solution turns orange/brown.
If no reaction occurs, the aqueous solution remains colourless. These colour changes allow the halogen present to be identified.

	Potassium chloride (KCl)	Potassium bromide (KBr)	Potassium iodide (KI)
Chlorine (Cl $_{2}$ )		Yellow (Br $_{2}$ )	Orange/brown (I $_{2}$ )
Bromine (Br $_{2}$ )	No colour change (no reaction)		Orange/brown (I $_{2}$ )
Iodine (I $_{2}$ )	No colour change (no reaction)	No colour change (no reaction)

Colour changes in organic solution

The halogen products can be more clearly visualized by using an organic solvent like cyclohexane. After the reaction, shaking the aqueous layer with hexane will result in the halogen dissolving in the organic layer. This separates out on top of the aqueous solution:

If bromide (Br^-) is displaced, forming bromine (Br₂), the organic layer turns orange.
If iodide (I^-) is displaced, forming iodine (I₂), the organic layer turns purple.
If no reaction occurs, the organic solution remains colourless.

	Potassium chloride (KCl)	Potassium bromide (KBr)	Potassium iodide (KI)
Chlorine (Cl $_{2}$ )		Orange (Br $_{2}$ )	Purple (I $_{2}$ )
Bromine (Br $_{2}$ )	No colour change (no reaction)		Purple (I $_{2}$ )
Iodine (I $_{2}$ )	No colour change (no reaction)	No colour change (no reaction)

So the halogen products can be clearly identified based on the colour changes in both aqueous and organic layers after the displacement reactions.

Halogen

Displaces

Ionic equation(s)

Chlorine (Cl

_{2}

)

Bromide (Br

^{-}

) and Iodide (I

^{-}

)

_{2 (a q)}

+ 2Br

^{-}

_{(a q)}

➔ 2Cl

^{-}

_{(a q)}

+ Br

_{2 (a q)}

/ Cl

_{2 (a q)}

+ 2I

^{-}

_{(a q)}

➔ 2Cl

^{-}

_{(a q)}

+ I

_{2 (a q)}

Bromine (Br

_{2}

)

Iodide (I

^{-}

)

_{2 (a q)}

+ 2I

^{-}

_{(a q)}

➔ 2Br

^{-}

_{(a q)}

+ I

_{2 (a q)}

Iodine (I

_{2}

)

None

No reaction

Potassium chloride (KCl)

Potassium bromide (KBr)

Potassium iodide (KI)

Chlorine (Cl $_{2}$ )

Yellow (Br

_{2}

)

Orange/brown (I

_{2}

)

Bromine (Br $_{2}$ )

No colour change (no reaction)

Orange/brown (I

_{2}

)

Iodine (I $_{2}$ )

No colour change (no reaction)

Potassium chloride (KCl)

Potassium bromide (KBr)

Potassium iodide (KI)

Chlorine (Cl $_{2}$ )

Orange (Br

_{2}

)

Purple (I

_{2}

)

Bromine (Br $_{2}$ )

No colour change (no reaction)

Purple (I

_{2}

)

Iodine (I $_{2}$ )

No colour change (no reaction)