Introduction to Group VII: The Halogens

Welcome to one of the most colourful and reactive chapters in AS Chemistry! The halogens are the non-metal elements found in Group VII (also known as Group 17) of the Periodic Table. In this unit, we will explore their physical trends, investigate how their reactivity changes as you move down the group, master qualitative tests for halide ions, and look closely at some dramatic redox reactions.

Don't worry if all the colours, equations, and observations seem overwhelming at first. We will break each concept down into small, digestible steps with handy memory tricks to help you secure top marks in your CCEA AS 1 exam!

A. The Elements and Their Electronic Configuration

The Group VII elements are: Fluorine (\(F\)), Chlorine (\(Cl\)), Bromine (\(Br\)), Iodine (\(I\)), and Astatine (\(At\)).
All halogens have the outer shell electronic configuration of \(ns^2 np^5\) (7 valence electrons). This means they only need to gain 1 electron to achieve a stable octet (noble gas configuration).

B. Appearance and Physical State at Room Temperature (\(20\,^\circ\text{C}\) / RTP)

At room temperature, the halogens exist as simple non-polar diatomic molecules (\(X_2\)):

• Fluorine (\(F_2\)): Yellow gas.
• Chlorine (\(Cl_2\)): Green-yellow / yellow-green gas.
• Bromine (\(Br_2\)): Red-brown liquid (volatilises readily to form dense orange/brown fumes).
• Iodine (\(I_2\)): Grey-black solid (sublimes to a purple/violet vapour upon gentle heating).

Analogy: Think of the states of matter walking down the group from gas (\(F_2, Cl_2\)) to liquid (\(Br_2\)) to solid (\(I_2\)). As the molecules get heavier, they stick together more firmly!

C. Periodic Trends Down Group VII

1. Atomic Radius Increases: As you go down the group, extra electron shells are added, increasing the distance between the nucleus and the outer electrons.

2. Electronegativity Decreases: Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond. Fluorine is the most electronegative element in the Periodic Table (\(F > Cl > Br > I\)). Down the group, atomic radius increases and electron shielding increases, so the nucleus exerts a weaker pull on shared pairs of electrons.

3. Melting and Boiling Points Increase: As you descend the group, the relative molecular mass (\(M_r\)) increases and the number of electrons per molecule increases. This leads to larger, more polarisable electron clouds and significantly stronger van der Waals (London dispersion) forces between diatomic molecules. More thermal energy is required to overcome these intermolecular attractions.

D. Solubility and Solution Colours

Because halogen molecules (\(X_2\)) are non-polar, they have low solubility in polar solvents like water, but dissolve readily in non-polar organic solvents such as hexane or cyclohexane via van der Waals forces.

Colours in Water (Aqueous Layer):
• Chlorine water: Very pale green / nearly colourless.
• Bromine water: Orange / yellow-orange / brown.
• Iodine in aqueous solution / aqueous \(KI\): Brown / yellow-brown.

Colours in Non-Polar Organic Solvent (e.g., Hexane / Cyclohexane):
• Chlorine: Pale green / colourless.
• Bromine: Orange / red-orange.
• Iodine: Distinct violet / purple.

Top Exam Tip: Adding a non-polar solvent like hexane and shaking allows halogen colours to concentrate in the top organic layer. This is especially helpful for confirming the presence of iodine, which turns a unmistakable bright violet/purple!

Key Takeaway

Down Group VII: Atomic radius, number of electrons, van der Waals forces, and boiling points increase. Electronegativity decreases.

2. Oxidising Ability and Displacement Reactions

A. Trend in Oxidising Power

Halogens act as oxidising agents because they gain electrons and are reduced to halide ions:
\(X_2 + 2e^- \rightarrow 2X^-\)

The oxidising ability decreases down the group: \(F_2 > Cl_2 > Br_2 > I_2\).
Why? Down the group, atomic radius increases and there is more shielding from inner shells. As a result, the incoming electron feels a weaker electrostatic attraction to the positive nucleus, making it harder for the atom to attract and gain an electron.

B. Displacement Reactions in Aqueous Solution

A more reactive halogen (stronger oxidising agent higher up the group) will displace a less reactive halide ion from its aqueous solution.

1. Chlorine displacing Bromide ions:
\(\text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)}\)
• Observation: Colourless solution turns yellow/orange (due to liberated \(\text{Br}_2\)).

2. Chlorine displacing Iodide ions:
\(\text{Cl}_2\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{I}_2\text{(aq)}\)
• Observation: Colourless solution turns brown (due to liberated \(\text{I}_2\)); adding an organic solvent produces a purple upper layer.

3. Bromine displacing Iodide ions:
\(\text{Br}_2\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow 2\text{Br}^-\text{(aq)} + \text{I}_2\text{(aq)}\)
• Observation: Orange solution turns brown (due to liberated \(\text{I}_2\)); adding an organic solvent produces a purple upper layer.

Note: Iodine cannot displace chloride or bromide ions because it is the weakest oxidising agent of the three.

Common Mistake to Avoid

Never confuse a halogen with a halide! Chlorine (\(\text{Cl}_2\)) is the halogen and the oxidising agent. Chloride (\(\text{Cl}^-\)) is the halide ion. Writing "chloride oxidises bromide" in an exam will cost you marks.

3. Disproportionation Reactions of Chlorine with Sodium Hydroxide

A disproportionation reaction is a special type of redox reaction where the same element is simultaneously oxidised and reduced.

A. Reaction with Cold, Dilute Aqueous \(\text{NaOH}\) (\(\approx 15\,^\circ\text{C}\))

\(\text{Cl}_2\text{(g/aq)} + 2\text{NaOH}\text{(aq)} \rightarrow \text{NaCl}\text{(aq)} + \text{NaClO}\text{(aq)} + \text{H}_2\text{O}\text{(l)}\)

• Products: Sodium chloride, sodium chlorate(I), and water.
• Oxidation states of chlorine:
- In \(\text{Cl}_2\): \(0\)
- In \(\text{NaCl}\): \(-1\) (Reduction: \(0 \rightarrow -1\))
- In \(\text{NaClO}\): \(+1\) (Oxidation: \(0 \rightarrow +1\))
• Real-world connection: The resulting solution containing \(\text{NaClO}\) is household bleach, used as a disinfectant and bleaching agent.

B. Reaction with Hot, Concentrated Aqueous \(\text{NaOH}\) (\(\approx 70\,^\circ\text{C}\))

\(3\text{Cl}_2\text{(g/aq)} + 6\text{NaOH}\text{(aq)} \rightarrow 5\text{NaCl}\text{(aq)} + \text{NaClO}_3\text{(aq)} + 3\text{H}_2\text{O}\text{(l)}\)

• Products: Sodium chloride, sodium chlorate(V), and water.
• Oxidation states of chlorine:
- In \(\text{Cl}_2\): \(0\)
- In \(\text{NaCl}\): \(-1\) (Reduction: \(0 \rightarrow -1\))
- In \(\text{NaClO}_3\): \(+5\) (Oxidation: \(0 \rightarrow +5\))

Key Takeaway

Cold, dilute \(\text{NaOH}\) produces chlorate(I) (\(\text{ClO}^-\)), while hot, concentrated \(\text{NaOH}\) pushes the oxidation further to form chlorate(V) (\(\text{ClO}_3^-\)).

4. Qualitative Identification of Halide Ions (\(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\))

Testing for halide ions is a classic practical procedure in both Unit AS 1 and Unit AS 3.

A. The Standard Silver Nitrate Test

Step 1: Add dilute nitric acid (\(\text{HNO}_3\)) to the unknown sample solution.
• Why? The nitric acid removes interfering carbonate (\(\text{CO}_3^{2-}\)) or sulfite (\(\text{SO}_3^{2-}\)) ions by reacting with them. If not removed, they would react with silver ions to form confusing white precipitates like \(\text{Ag}_2\text{CO}_3\).
• Warning: Never acidify with hydrochloric acid (\(\text{HCl}\))! Adding \(\text{HCl}\) introduces chloride ions, which immediately forms an unwanted white precipitate of \(\text{AgCl}\) and ruins the test.

Step 2: Add aqueous silver nitrate (\(\text{AgNO}_3\text{(aq)}\)) dropwise.

Observations & Precipitation Equations:
• Chloride (\(\text{Cl}^-\)): White precipitate (\(\text{AgCl}\))
\(\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl}\text{(s)}\)
• Bromide (\(\text{Br}^-\)): Cream precipitate (\(\text{AgBr}\))
\(\text{Ag}^+\text{(aq)} + \text{Br}^-\text{(aq)} \rightarrow \text{AgBr}\text{(s)}\)
• Iodide (\(\text{I}^-\)): Yellow precipitate (\(\text{AgI}\))
\(\text{Ag}^+\text{(aq)} + \text{I}^-\text{(aq)} \rightarrow \text{AgI}\text{(s)}\)

Exam Note: Always include state symbols when writing ionic equations for precipitation reactions!

B. Confirmatory Test: Solubility in Aqueous Ammonia (\(\text{NH}_3\))

Because white, cream, and yellow precipitates can look very similar under laboratory lighting, we use aqueous ammonia to distinguish between them conclusively:

• Silver chloride (\(\text{AgCl}\)): Dissolves in dilute \(\text{NH}_3\text{(aq)}\) to give a colourless solution (forming the complex ion \([\text{Ag}(\text{NH}_3)_2]^+\)).
• Silver bromide (\(\text{AgBr}\)): Insoluble in dilute \(\text{NH}_3\text{(aq)}\), but dissolves in concentrated \(\text{NH}_3\text{(aq)}\).
• Silver iodide (\(\text{AgI}\)): Insoluble in both dilute and concentrated \(\text{NH}_3\text{(aq)}\).

Memory Trick: Think alphabetical or progressive solubility — Chloride dissolves easiest (dilute), Bromide needs force (concentrated), and Iodide refuses to dissolve at all!

5. Reducing Action of Solid Halides with Concentrated Sulfuric Acid

A. Trend in Reducing Ability

Halide ions act as reducing agents by donating electrons: \(2X^- \rightarrow X_2 + 2e^-\).
The reducing power increases down the group: \(\text{F}^- < \text{Cl}^- < \text{Br}^- < \text{I}^-\).
Why? As you go down the group, the ionic radius of the halide ion increases. The outer electrons are further from the nucleus and experience more shielding, so they are held less tightly and lost much more readily.

B. Solid Sodium Chloride (\(\text{NaCl}\)) + Concentrated \(\text{H}_2\text{SO}_4\)

\(\text{NaCl}\text{(s)} + \text{H}_2\text{SO}_4\text{(l)} \rightarrow \text{NaHSO}_4\text{(s)} + \text{HCl}\text{(g)}\)

• Type of reaction: Acid-base / proton transfer only (NOT a redox reaction). Chloride ions are not strong enough reducing agents to reduce sulfuric acid.
• Observations: Steamy / misty acidic fumes of \(\text{HCl}\) gas (turns damp blue litmus paper red; forms dense white smoke of \(\text{NH}_4\text{Cl}\) with ammonia fumes).

C. Solid Sodium Bromide (\(\text{NaBr}\)) + Concentrated \(\text{H}_2\text{SO}_4\)

Bromide ions are stronger reducing agents than chloride ions and can reduce sulfur from \(+6\) (in \(\text{H}_2\text{SO}_4\)) to \(+4\) (in \(\text{SO}_2\)).

1. Acid-base step:
\(\text{NaBr}\text{(s)} + \text{H}_2\text{SO}_4\text{(l)} \rightarrow \text{NaHSO}_4\text{(s)} + \text{HBr}\text{(g)}\)
• Produces misty fumes of \(\text{HBr}\).

2. Redox step:
\(2\text{HBr}\text{(g)} + \text{H}_2\text{SO}_4\text{(l)} \rightarrow \text{Br}_2\text{(g/l)} + \text{SO}_2\text{(g)} + 2\text{H}_2\text{O}\text{(l)}\)
• Observations: Steamy fumes (\(\text{HBr}\)), orange / red-brown fumes (\(\text{Br}_2\)), and a colourless gas with a pungent, choking smell (\(\text{SO}_2\)).

D. Solid Sodium Iodide (\(\text{NaI}\)) + Concentrated \(\text{H}_2\text{SO}_4\)

Iodide ions are very powerful reducing agents. They reduce sulfuric acid through multiple steps, lowering the oxidation state of sulfur from \(+6\) down to \(+4\), \(0\), and \(-2\)!

1. Acid-base step:
\(\text{NaI}\text{(s)} + \text{H}_2\text{SO}_4\text{(l)} \rightarrow \text{NaHSO}_4\text{(s)} + \text{HI}\text{(g)}\) (Misty fumes of \(\text{HI}\))

2. Redox reductions:
• Reduction to \(\text{SO}_2\) (Sulfur oxidation state \(+4\)):
\(2\text{HI} + \text{H}_2\text{SO}_4 \rightarrow \text{I}_2 + \text{SO}_2 + 2\text{H}_2\text{O}\)
• Reduction to elemental sulfur \(\text{S}\) (Sulfur oxidation state \(0\)):
\(6\text{HI} + \text{H}_2\text{SO}_4 \rightarrow 3\text{I}_2 + \text{S} + 4\text{H}_2\text{O}\)
• Reduction to hydrogen sulfide \(\text{H}_2\text{S}\) (Sulfur oxidation state \(-2\)):
\(8\text{HI} + \text{H}_2\text{SO}_4 \rightarrow 4\text{I}_2 + \text{H}_2\text{S} + 4\text{H}_2\text{O}\)

Full Set of Observations for Iodide:
• Purple vapours / dark grey solid / brown solution (elemental iodine, \(\text{I}_2\))
• Pungent, choking gas (\(\text{SO}_2\))
• Yellow solid deposit (elemental sulfur, \(\text{S}\))
• Bad-egg / rotten-egg smell (hydrogen sulfide gas, \(\text{H}_2\text{S}\))

E. Preparation of Pure Hydrogen Halides using Phosphoric Acid (\(\text{H}_3\text{PO}_4\))

Because concentrated \(\text{H}_2\text{SO}_4\) oxidises \(\text{HBr}\) and \(\text{HI}\), it cannot be used to make pure samples of these gases. Instead, we use concentrated phosphoric acid (\(\text{H}_3\text{PO}_4\)):
\(\text{KX}\text{(s)} + \text{H}_3\text{PO}_4\text{(l)} \rightarrow \text{KH}_2\text{PO}_4\text{(s)} + \text{HX}\text{(g)}\)
Since phosphoric acid is not an oxidising agent, no redox reaction occurs, yielding pure steamy fumes of \(\text{HCl}\), \(\text{HBr}\), or \(\text{HI}\).

Key Takeaway

Chloride only does acid-base with \(\text{H}_2\text{SO}_4\). Bromide reduces sulfur to \(\text{SO}_2\) (\(+4\)). Iodide reduces sulfur all the way to \(\text{SO}_2\) (\(+4\)), \(\text{S}\) (\(0\)), and \(\text{H}_2\text{S}\) (\(-2\)).

Quick Review: Summary of Key Reactions & Reagents

• Oxidising ability of halogens: \(\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2\) (decreases down the group).
• Reducing ability of halide ions: \(\text{F}^- < \text{Cl}^- < \text{Br}^- < \text{I}^-\) (increases down the group).
• Halide Test: Dilute \(\text{HNO}_3\) + \(\text{AgNO}_3\text{(aq)}\) \(\rightarrow\) White (\(\text{Cl}^-\)), Cream (\(\text{Br}^-\)), Yellow (\(\text{I}^-\)).
• Ammonia Solubility: \(\text{AgCl}\) (dissolves in dilute), \(\text{AgBr}\) (dissolves in conc), \(\text{AgI}\) (insoluble in conc).
• Bleach synthesis: \(\text{Cl}_2 + 2\text{NaOH} \rightarrow \text{NaCl} + \text{NaClO} + \text{H}_2\text{O}\) (cold, dilute).