Group 17 Chemistry

Group 17 of the Periodic Table is home to the halogens, a series of highly reactive non-metals. The elements we focus on in A Level Chemistry are chlorine (Cl), bromine (Br), and iodine (I). Their atoms all have seven valence electrons with an outer shell electron configuration of ns²np⁵. This drives their chemistry, as they are one electron short of a stable noble gas configuration, making them eager to gain an electron.

### Physical Properties and Trends

Halogens exist as simple diatomic molecules (X₂) , held together by weak intermolecular van der Waals forces. The strength of these forces increases with the number of electrons in the molecule. This explains the trends in their physical state and volatility down the group:

* State at r.t.p.: Chlorine (Cl₂) is a pale green gas, Bromine (Br₂) is a volatile, red-brown liquid, and Iodine (I₂) is a grey-black solid that sublimes to form a violet vapour. The increasing strength of van der Waals forces requires more energy to overcome, leading to higher melting and boiling points down the group.

* Volatility: This decreases down the group due to the strengthening intermolecular forces.

* Electronegativity: This is the ability of an atom to attract a bonding pair of electrons. Electronegativity decreases down the group. As the atomic radius increases and the number of inner shells (shielding) increases, the nucleus has a weaker pull on the bonding electrons.

### Chemical Properties and Reactivity

The dominant chemical feature of halogens is their ability to act as strong oxidising agents. They achieve a stable octet by gaining an electron, thereby being reduced themselves and oxidising another species.

1. Trend in Oxidising Power: The oxidising strength of the halogens **decreases down the group (Cl₂ > Br₂ > I₂)**. The smaller chlorine atom has a higher charge density and less shielding, allowing its nucleus to attract an incoming electron more strongly than bromine or iodine.

This trend is clearly demonstrated in halogen displacement reactions. A more reactive halogen will oxidise and displace the halide ion of a less reactive halogen from its aqueous solution.

* Chlorine displaces bromide and iodide ions:

* Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq) (Solution turns from colourless to orange/brown)

* Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq) (Solution turns from colourless to brown)

* Bromine displaces iodide ions, but not chloride ions:

* Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq) (Solution turns from orange to brown)

* Iodine cannot displace chloride or bromide ions.

2. Hydrogen Halides (HX): Halogens react with hydrogen to form hydrogen halides. The reactivity **decreases down the group**. For example, the reaction of chlorine with hydrogen is explosive in sunlight, while the reaction with iodine is reversible and requires heating.

The thermal stability of the hydrogen halides decreases down the group (HCl > HBr > HI). This is because the H-X bond enthalpy decreases as the halogen atom gets larger. The increased bond length leads to a weaker bond that is more easily broken by heat.

* 2HI(g) ⇌ H₂(g) + I₂(g) (Decomposes easily on heating)

### Reactions of Halide Ions

1. Test for Halide Ions: Aqueous halide ions can be identified using a standard test involving **aqueous silver nitrate (AgNO₃)**, followed by the addition of **ammonia solution (NH₃)**.

* Chloride (Cl⁻): Forms a white precipitate of silver chloride (AgCl), which is soluble in dilute NH₃(aq).

* Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

* Bromide (Br⁻): Forms a cream precipitate of silver bromide (AgBr), which is insoluble in dilute NH₃(aq) but soluble in concentrated NH₃(aq).

* Ag⁺(aq) + Br⁻(aq) → AgBr(s)

* Iodide (I⁻): Forms a yellow precipitate of silver iodide (AgI), which is insoluble in both dilute and concentrated NH₃(aq).

* Ag⁺(aq) + I⁻(aq) → AgI(s)

2. Reducing Power of Halide Ions: In contrast to the elements, the reducing power of the halide ions **increases down the group (I⁻ > Br⁻ > Cl⁻)**. This is because the outermost electron in the larger iodide ion is further from the nucleus and more easily lost (oxidised) than in the smaller chloride ion.

This trend is observed in their reactions with concentrated sulfuric acid (H₂SO₄).

* With NaCl: Concentrated H₂SO₄ acts as an acid, not an oxidising agent. Misty fumes of HCl are produced. No redox reaction occurs.

* NaCl(s) + H₂SO₄(conc) → NaHSO₄(s) + HCl(g)

* With NaBr: The bromide ion is a stronger reducing agent. It first forms HBr, which then reduces the H₂SO₄ to sulfur dioxide.

* 2HBr(g) + H₂SO₄(l) → Br₂(g) + SO₂(g) + 2H₂O(l) (Brown fumes of Br₂ and choking SO₂ gas are seen).

* With NaI: The iodide ion is the strongest reducing agent. It reduces H₂SO₄ to multiple products, including sulfur dioxide (SO₂), sulfur (S), and hydrogen sulfide (H₂S).

* 8HI(g) + H₂SO₄(l) → 4I₂(s) + H₂S(g) + 4H₂O(l) (Purple I₂ vapour, yellow solid S, and the rotten egg smell of H₂S can be observed).