Group 7(17), the halogens

Q: How and why does the boiling point of the halogens change down Group 7?

Boiling points increase down Group 7 because atoms/molecules get larger and have more electrons, leading to stronger van der Waals forces that require more energy to break.

Q: What is the trend in electronegativity down Group 7 and what causes it?

Electronegativity decreases down Group 7 because increased atomic radius and shielding by inner electron shells weaken the nucleus's attraction for a bonding pair of electrons.

Q: How does the oxidising ability of the halogens vary down Group 7 and what determines this trend?

Oxidising ability decreases down Group 7 because smaller halogen atoms at the top of the group attract and gain an electron more readily.

Q: Why does the reducing ability of halide ions increase down Group 7?

Reducing ability increases down the group because larger halide ions have more shielding and a larger ionic radius, making it easier to lose the outermost electron.

Q: How can ammonia (\( \text{NH}_3 \)) solution distinguish between silver chloride, silver bromide, and silver iodide precipitates?

\( \text{AgCl} \) dissolves in dilute \( \text{NH}_3 \); \( \text{AgBr} \) dissolves only in concentrated \( \text{NH}_3 \); \( \text{AgI} \) is insoluble in both.

Chemistry 7404 · AQA AS Level · 6 min read · Updated: 17 Sept 2026

Welcome to Group 7: The Halogens!

In this chapter, we are going to explore the halogens. These are the elements found in Group 7 (or 17) of the Periodic Table: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At).

These elements are highly reactive non-metals. You’ll find them in everything from the salt on your chips to the bleach used to keep swimming pools clean. Don’t worry if some of the trends seem confusing at first—we will break them down step-by-step!

1. Trends in Physical Properties

As we move down Group 7 from Fluorine to Iodine, the atoms change in very predictable ways.

Boiling Points

The Trend: Boiling points increase as you go down the group.

The Reason:
1. As you go down the group, the atoms get larger and have more electrons.
2. This leads to stronger van der Waals forces (intermolecular forces) between the molecules.
3. Stronger forces require more energy to break, so the boiling point goes up.

Analogy: Think of van der Waals forces like Velcro. A small molecule has a tiny bit of Velcro, so it’s easy to pull away. A large molecule has a huge strip of Velcro, making it much harder to separate!

Electronegativity

Prerequisite Check: Remember, electronegativity is the power of an atom to attract a pair of electrons in a covalent bond.

The Trend: Electronegativity decreases as you go down the group.

The Reason:
1. Atoms get larger, so the distance between the nucleus and the outer electrons increases.
2. There are more inner shells, which shield the nucleus’s positive charge.
3. Therefore, the nucleus has a weaker "pull" on the bonding pair of electrons.

Quick Review: Down the group, atoms get fatter (higher boiling point) but weaker at grabbing electrons (lower electronegativity).

2. Halogens as Oxidising Agents

Halogens want to gain one electron to get a full outer shell. Because they take electrons from other things, they are called oxidising agents.

The Trend: Oxidising ability decreases down the group. Fluorine is the strongest; Iodine is the weakest.

Displacement Reactions

A "stronger" halogen (one higher up the group) will kick out (displace) a "weaker" halide ion from a solution.

Example: Chlorine is higher than Bromine. If you add Chlorine water to Sodium Bromide, the Chlorine takes the electrons!
\( \text{Cl}_{2}(aq) + 2\text{Br}^{-}(aq) \rightarrow 2\text{Cl}^{-}(aq) + \text{Br}_{2}(aq) \)

Common Mistake to Avoid: A halogen cannot displace itself or anything above it. For example, Iodine cannot displace Chlorine.

Key Takeaway: The element higher up the group is the "bully"—it takes the electrons from the ions lower down.

3. Halide Ions as Reducing Agents

This is the opposite of the trend above. Halide ions (\( \text{Cl}^{-}, \text{Br}^{-}, \text{I}^{-} \)) want to give away an electron to go back to being a halogen. Because they give away electrons, they are reducing agents.

The Trend: Reducing ability increases down the group. Iodide is the strongest reducer; Fluoride is the weakest.

The Reason: In an Iodide ion, the outer electron is very far from the nucleus and heavily shielded. It’s very easy for the ion to "lose" that electron.

Reaction with Concentrated Sulfuric Acid (\( \text{H}_{2}\text{SO}_{4} \))

This is a classic exam topic. You need to know what happens when you drop \( \text{H}_{2}\text{SO}_{4} \) onto solid sodium halides:

1. Sodium Chloride (\( \text{NaCl} \)): No redox occurs because \( \text{Cl}^{-} \) is a weak reducer. You just get white misty fumes of \( \text{HCl} \) gas.
\( \text{NaCl}(s) + \text{H}_{2}\text{SO}_{4}(l) \rightarrow \text{NaHSO}_{4}(s) + \text{HCl}(g) \)

2. Sodium Bromide (\( \text{NaBr} \)): \( \text{Br}^{-} \) is a stronger reducer. It reduces the Sulfur from +6 to +4 (forming \( \text{SO}_{2} \)). You see orange fumes of Bromine and choking \( \text{SO}_{2} \) gas.
\( 2\text{H}^{+} + 2\text{Br}^{-} + \text{H}_{2}\text{SO}_{4} \rightarrow \text{Br}_{2}(g) + \text{SO}_{2}(g) + 2\text{H}_{2}\text{O}(l) \)

3. Sodium Iodide (\( \text{NaI} \)): \( \text{I}^{-} \) is the "heavyweight" reducer. It reduces Sulfur through multiple stages: to \( \text{SO}_{2} \) (+4), to a yellow solid of Sulfur (\( \text{S} \), 0), and all the way down to \( \text{H}_{2}\text{S} \) (-2). You see purple vapors (and black solid) of Iodine, a yellow solid (\( \text{S} \)), and smell rotten eggs (\( \text{H}_{2}\text{S} \)).

4. Identifying Halide Ions (The Silver Nitrate Test)

How do we tell which halide we have in a test tube? We use Silver Nitrate (\( \text{AgNO}_{3} \)).

The Step-by-Step Process:

1. Add dilute nitric acid first (this removes any interfering ions like carbonates that would form false precipitates).
2. Add a few drops of Silver Nitrate solution.
3. Observe the color of the precipitate (solid) formed.

The Results (Memory Aid):

Chloride = White precipitate (\( \text{AgCl} \))
Bromide = Cream precipitate (\( \text{AgBr} \))
Iodide = Yellow precipitate (\( \text{AgI} \))

Mnemonic: Cats Want Big Cups of Icy Yogurt.

The Ammonia Solubility Test

Sometimes the precipitate colors look very similar. We use Ammonia (\( \text{NH}_{3} \)) to confirm:

- Silver Chloride: Dissolves in dilute ammonia.
- Silver Bromide: Dissolves only in concentrated ammonia.
- Silver Iodide: Insoluble in both dilute and concentrated ammonia.

5. Chlorine and Water Treatment

Chlorine is a bit of a "double agent." It can be both oxidised and reduced at the same time in the same reaction. This is called disproportionation.

Reaction with Water

In reversible conditions without bright sunlight:
\( \text{Cl}_{2}(g) + \text{H}_{2}\text{O}(l) \rightleftharpoons \text{HCl}(aq) + \text{HClO}(aq) \)

The chloric(I) acid (\( \text{HClO} \)) produced is an oxidising agent that kills bacteria, making our drinking water safe.

Reaction of Chlorine with Water in Sunlight

In the presence of UV light / direct sunlight, a different reaction takes place where water and chlorine decompose to produce oxygen gas and hydrochloric acid:
\( 2\text{Cl}_{2}(aq) + 2\text{H}_{2}\text{O}(l) \rightarrow 4\text{HCl}(aq) + \text{O}_{2}(g) \)

The Benefits vs. Risks

Benefits: It destroys pathogens (bacteria) that cause life-threatening diseases like cholera and typhoid. It also prevents the growth of algae in water mains.
Risks: Chlorine gas is toxic and irritating to the respiratory system. It can also react with naturally occurring organic matter in water to form chlorinated hydrocarbons (trihalomethanes), which are suspected carcinogens.

Key Assessment Point: Society accepts the use of chlorine because the health benefits of having disease-free drinking water far outweigh the toxic risks.

Making Household Bleach

When chlorine reacts with cold, dilute aqueous sodium hydroxide (\( \text{NaOH} \)), it disproportionates to produce sodium chlorate(I), which is active bleach:
\( \text{Cl}_{2}(g) + 2\text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{NaClO}(aq) + \text{H}_{2}\text{O}(l) \)

Section Summary: Key Takeaways

- Boiling points increase down Group 7; Electronegativity decreases.
- Oxidising power of halogens decreases down the group (\( \text{F}_{2} \) is strongest).
- Reducing power of halide ions increases down the group (\( \text{I}^{-} \) is strongest).
- Silver Nitrate distinguishes halides by precipitate color (White, Cream, Yellow) and solubility in ammonia.
- Chlorine treats water by forming \( \text{HClO} \); in sunlight it decomposes to form \( \text{O}_{2} \) and \( \text{HCl} \).

References

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

How and why does the boiling point of the halogens change down Group 7?

Boiling points increase down Group 7 because atoms/molecules get larger and have more electrons, leading to stronger van der Waals forces that require more energy to break.

What is the trend in electronegativity down Group 7 and what causes it?

Electronegativity decreases down Group 7 because increased atomic radius and shielding by inner electron shells weaken the nucleus's attraction for a bonding pair of electrons.

How does the oxidising ability of the halogens vary down Group 7 and what determines this trend?

Oxidising ability decreases down Group 7 because smaller halogen atoms at the top of the group attract and gain an electron more readily.

Which halogens can displace halide ions from aqueous solution?

A halogen higher up the group will displace the halide ion of an element below it in the group (e.g., \( \text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2 \)).

Why does the reducing ability of halide ions increase down Group 7?

Reducing ability increases down the group because larger halide ions have more shielding and a larger ionic radius, making it easier to lose the outermost electron.

What are the key observations and products when solid sodium halides react with concentrated sulfuric acid (\( \text{H}_2\text{SO}_4 \))?

\( \text{NaCl} \) gives misty white fumes of \( \text{HCl} \) (non-redox); \( \text{NaBr} \) produces orange \( \text{Br}_2 \) fumes and \( \text{SO}_2 \); \( \text{NaI} \) reduces sulfur to \( \text{SO}_2 \), yellow \( \text{S} \), and rotten-egg-smelling \( \text{H}_2\text{S} \) alongside purple \( \text{I}_2 \) vapor.

How is acidified silver nitrate (\( \text{AgNO}_3 \)) used to identify chloride, bromide, and iodide ions?

Adding dilute \( \text{HNO}_3 \) followed by \( \text{AgNO}_3 \) forms distinct precipitates: white (\( \text{AgCl} \)), cream (\( \text{AgBr} \)), and yellow (\( \text{AgI} \)).

How can ammonia (\( \text{NH}_3 \)) solution distinguish between silver chloride, silver bromide, and silver iodide precipitates?

\( \text{AgCl} \) dissolves in dilute \( \text{NH}_3 \); \( \text{AgBr} \) dissolves only in concentrated \( \text{NH}_3 \); \( \text{AgI} \) is insoluble in both.

What reaction occurs when chlorine is added to water in the absence of bright sunlight?

Chlorine undergoes disproportionation to form hydrochloric acid and chloric(I) acid: \( \text{Cl}_2(g) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HCl}(aq) + \text{HClO}(aq) \).

How does chlorine react with water in the presence of direct sunlight / UV light?

Chlorine decomposes water into hydrochloric acid and oxygen gas: \( 2\text{Cl}_2(aq) + 2\text{H}_2\text{O}(l) \rightarrow 4\text{HCl}(aq) + \text{O}_2(g) \).

What are the key benefits and risks associated with the chlorination of drinking water?

Benefits include destroying deadly pathogens (e.g., cholera) and preventing algae growth; risks include chlorine toxicity and the formation of carcinogenic chlorinated hydrocarbons.

What equation describes the formation of household bleach from chlorine and sodium hydroxide?

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

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