Introduction to Group 7: The Halogens

Welcome to the world of the Halogens! This group (Group 7) contains some of the most reactive and interesting non-metals in the Periodic Table. From the chlorine used to keep swimming pools clean to the iodine used in medicine, these elements are everywhere. In this chapter, we will explore why they behave the way they do, how their properties change as we go down the group, and some of the key chemical reactions you need to know for your Pearson Edexcel A Level exams.

Did you know? The word "Halogen" comes from Greek and means "salt-former." When these elements react with metals, they produce a wide range of salts, like the common table salt (\(NaCl\)) you put on your chips!


As we move down Group 7 from Fluorine (\(F\)) to Iodine (\(I\)), we see very clear patterns in their physical properties. Understanding why these trends happen is the secret to scoring high marks.

1. Appearance and State at Room Temperature

As you go down the group, the elements become darker in color and change state from gas to liquid to solid:

  • Fluorine (\(F_{2}\)): Pale yellow gas.
  • Chlorine (\(Cl_{2}\)): Pale green gas.
  • Bromine (\(Br_{2}\)): Dark red-brown liquid (which gives off orange vapors).
  • Iodine (\(I_{2}\)): Grey-black solid (which sublimes to form a purple vapor).

2. Melting and Boiling Temperatures

Trend: Melting and boiling points increase as you go down the group.
The Science: Halogens exist as simple diatomic molecules (\(F_{2}\), \(Cl_{2}\), etc.). The molecules are held together by weak intermolecular forces called London forces (induced dipole-dipole forces). As you go down the group, the number of electrons in the molecules increases. This makes the London forces stronger, so more energy is needed to break them.

3. Electronegativity

Trend: Electronegativity decreases as you go down the group.
The Science: Electronegativity is the ability of an atom to attract a bonding pair of electrons in a covalent bond. As you go down the group, the atomic radius increases and there is more shielding from inner electron shells. This means the nucleus has a weaker "pull" on the shared pair of electrons.

Quick Review: Down the group = More electrons = Stronger London forces = Higher Boiling Points. However, larger atoms = More shielding = Lower Electronegativity.


Section 2: Reactivity and Displacement Reactions

In Group 7, reactivity decreases as you go down the group. Fluorine is the most reactive, while Iodine is much less so. This is the opposite of Group 1 and 2!

The Displacement Test

A more reactive halogen will "push out" (displace) a less reactive halide ion from its solution. This is a classic exam practical. We often add an organic solvent (like cyclohexane) to help us see the colors more clearly.

Example: Chlorine + Potassium Bromide
\(Cl_{2}(aq) + 2KBr(aq) \rightarrow 2KCl(aq) + Br_{2}(aq)\)
The ionic equation is: \(Cl_{2} + 2Br^{-} \rightarrow 2Cl^{-} + Br_{2}\)

  • Chlorine can displace Bromide and Iodide ions.
  • Bromine can displace Iodide ions (but not Chloride).
  • Iodine cannot displace either Chloride or Bromide.

Color in Organic Solvent: In water, bromine and iodine can look similar (orange/brown). If you shake the mixture with an organic solvent, the halogen dissolves in the top layer:
- Chlorine: Pale green/colorless.
- Bromine: Orange.
- Iodine: Purple/Pink.

Key Takeaway: Reactivity decreases down the group because larger atoms find it harder to attract and gain an electron into their outer shell due to increased shielding.


Section 3: Chlorine and Disproportionation

Disproportionation is a special type of redox reaction where the same element is both oxidized and reduced at the same time. You must learn these three reactions for chlorine.

1. Chlorine with Water

Used in water treatment to kill bacteria:
\(Cl_{2} + H_{2}O \rightleftharpoons HCl + HClO\)
Oxidation states: \(Cl\) goes from \(0\) in \(Cl_{2}\) to \(-1\) in \(HCl\) (reduced) and \(+1\) in \(HClO\) (oxidized).

2. Chlorine with Cold, Dilute Alkali (e.g., \(NaOH\))

This creates common household bleach:
\(Cl_{2} + 2NaOH \rightarrow NaCl + NaClO + H_{2}O\)
Conditions: 15°C (Cold). The \(NaClO\) (sodium chlorate(I)) is the active bleach ingredient.

3. Chlorine with Hot, Concentrated Alkali (e.g., \(NaOH\))

Under hotter conditions, the chlorine is oxidized even further:
\(3Cl_{2} + 6NaOH \rightarrow 5NaCl + NaClO_{3} + 3H_{2}O\)
Conditions: 70°C (Hot). Here, \(Cl\) goes to the \(+5\) oxidation state in \(NaClO_{3}\).


Section 4: Halide Ions and Concentrated Sulfuric Acid

This is a favorite topic for examiners! When solid Group 1 halides (like \(NaF\), \(NaCl\), \(NaBr\), \(NaI\)) react with concentrated sulfuric acid (\(H_{2}SO_{4}\)), the results depend on how good the halide is at acting as a reducing agent.

Memory Trick: Reducing power increases down the group. Iodide (\(I^{-}\)) is the best reducing agent because its outer electron is furthest from the nucleus and easily lost.

1. Fluoride (\(F^{-}\)) and Chloride (\(Cl^{-}\))

These are not strong enough to reduce the sulfur in \(H_{2}SO_{4}\). Only an acid-base reaction occurs.
\(NaCl(s) + H_{2}SO_{4}(conc) \rightarrow NaHSO_{4}(s) + HCl(g)\)
Observation: Misty fumes of \(HCl\).

2. Bromide (\(Br^{-}\))

Bromide is strong enough to reduce the sulfur from \(+6\) to \(+4\) (in \(SO_{2}\)).
First, the acid-base step: \(NaBr + H_{2}SO_{4} \rightarrow NaHSO_{4} + HBr\)
Then, the redox step: \(2HBr + H_{2}SO_{4} \rightarrow Br_{2} + SO_{2} + 2H_{2}O\)
Observations: Misty fumes (\(HBr\)), orange vapors (\(Br_{2}\)), and a choking smell (\(SO_{2}\)).

3. Iodide (\(I^{-}\))

Iodide is a very strong reducing agent. It reduces the sulfur from \(+6\) to \(+4\), \(0\), and even \(-2\)!
Observations:
- Misty fumes (\(HI\))
- Purple vapors/Black solid (\(I_{2}\))
- Yellow solid (Sulfur, \(S\))
- Rotten egg smell (Hydrogen sulfide, \(H_{2}S\)).


Section 5: Testing for Halide Ions

If you are given an unknown solution, here is the step-by-step way to identify the halide ion present. (Note: This is also part of Topic 4C).

Step 1: Add dilute nitric acid (\(HNO_{3}\)) to remove any carbonate impurities.
Step 2: Add Silver Nitrate solution (\(AgNO_{3}\)). Observe the precipitate color:

  • Chloride (\(Cl^{-}\)): White precipitate (\(AgCl\))
  • Bromide (\(Br^{-}\)): Cream precipitate (\(AgBr\))
  • Iodide (\(I^{-}\)): Yellow precipitate (\(AgI\))

Step 3: If the colors are hard to tell apart, add Ammonia solution (\(NH_{3}\)):

  • \(AgCl\): Dissolves in dilute ammonia.
  • \(AgBr\): Dissolves only in concentrated ammonia.
  • \(AgI\): Insoluble in ammonia (won't dissolve even in concentrated).

Section 6: Hydrogen Halides

Hydrogen halides (\(HF, HCl, HBr, HI\)) are colorless gases at room temperature.

  • Reaction with Water: They are very soluble and react with water to form strong acidic solutions (e.g., \(HCl(g) \rightarrow H^{+}(aq) + Cl^{-}(aq)\)). They produce "misty fumes" in moist air.
  • Reaction with Ammonia: They react with ammonia gas to form white smoky clouds of ammonium halides.
    Example: \(NH_{3}(g) + HCl(g) \rightarrow NH_{4}Cl(s)\)

Section 7: Predictions for Fluorine and Astatine

In the exam, you might be asked to predict properties for elements at the very top (Fluorine) or bottom (Astatine) of the group based on trends.

  • Fluorine (\(F\)): Expect it to be the most reactive, have the lowest boiling point, and be the strongest oxidizing agent. (Careful: Its bond enthalpy is surprisingly low because the small atoms repel each other, but you mainly need the general trends).
  • Astatine (\(At\)): Predict it to be a dark solid at room temperature, have a higher melting point than iodine, be the least reactive, and be the best reducing agent in its ion form (\(At^{-}\)).

Key Takeaway Summary: Group 7 is a masterclass in trends. Remember that physical properties (BP) increase down the group, but chemical reactivity and electronegativity decrease. Mastery of the sulfuric acid reactions and the silver nitrate tests will get you through the toughest exam questions!