Reactions of Group 17 Elements Revision notes | A-Level Chemistry CIE

Reactions of Group 17 Elements

This lesson covers:

How halogens gain electrons to become ions
Trends in reactivity and oxidising ability down the group
Trend in reactivity of halogens with hydrogen
Trend in thermal stability of hydrogen halides
Displacement reactions between halogens and halide ions

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 $_{2 (a q)}$ {2(aq)} + 2Br $^{-}$ $_{(a q)}$ {(aq)} ➔ 2Cl $^{-}$ $_{(a q)}$ + Br $_{2 (a q)}$ / Cl $_{2 (a q)}$ {2(aq)} + 2I $^{-}$ $_{(a q)}$ ➔ 2Cl $^{-}$ $_{(a q)}$ {(aq)} + I $_{2 (a q)}$
Bromine (Br $_{2}$ )	Iodide (I $^{-}$ )	Br $_{2 (a q)}$ {2(aq)} + 2I $^{-}$ $_{(a q)}$ ➔ 2Br $^{-}$ $_{(a q)}$ {(aq)} + 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.

Reactivity of halogens with hydrogen decreases down group 7

The halogens also react with hydrogen gas, forming the corresponding hydrogen halide compounds. The reactivity of the halogens in these reactions decreases as you move down group 7, following the same trend seen previously.

Here is a table describing the reactions of halogens and hydrogen:

Halogen	Equation	Description of reaction
Fluorine	F_2(g) + H_2(g) ➔ 2HF_(g)	Reacts explosively even in cool, dark conditions
Chlorine	Cl_2(g) + H_2(g) ➔ 2HCl_(g)	Reacts explosively in sunlight
Bromine	Br_2(g) + H_2(g) ➔ 2HBr_(g)	Reacts slowly on heating
Iodine	I_2(g) + H_2(g) ⇌ 2HI_(g)	Forms an equilibrium mixture on heating

Thermal stability of hydrogen halides decreases down group 7

The hydrogen halides (HX) exhibit a decreasing trend in thermal stability as you move down group 7 from HF to HI. This trend is directly related to the strength of the bond between the hydrogen and halogen atoms in these compounds.

Bond	Bond enthalpy (kJ mol^-1)
H-F	562
H-Cl	431
H-Br	366
H-I	299

The bond enthalpy values shown in the table above represent the amount of energy required to break the H-X bond in each hydrogen halide. As the bond enthalpy decreases, the thermal stability of the compound also decreases, meaning it is more easily decomposed by heat.

Several factors contribute to the decreasing bond strength and thermal stability trend observed in the hydrogen halides:

As you move down group 17, the atomic radius of the halogens increases.
The increasing size results in the shared electron pair being further away from the nuclei.
Consequently, the electrostatic attraction between the electron pair and the nuclei becomes weaker.
This means less energy is required to break the H-X covalent bond, thus lowering the thermal stability.

The trend in thermal stability is clearly demonstrated by the varying decomposition temperatures of the hydrogen halides. HI readily decomposes when heated, while HF remains stable at temperatures exceeding 1,000°C.