Group 7(17), the halogens

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!
\( Cl_{2}(aq) + 2Br^{-}(aq) \rightarrow 2Cl^{-}(aq) + Br_{2}(aq) \)

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 (\( Cl^{-}, Br^{-}, I^{-} \)) 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 (\( H_{2}SO_{4} \))

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

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

2. Sodium Bromide (NaBr): \( Br^{-} \) is stronger. It reduces the Sulfur from +6 to +4 (forming \( SO_{2} \)). You see orange fumes of Bromine.
\( 2H^{+} + 2Br^{-} + H_{2}SO_{4} \rightarrow Br_{2}(g) + SO_{2}(g) + 2H_{2}O(l) \)

3. Sodium Iodide (NaI): \( I^{-} \) is the "heavyweight" reducer. It reduces the Sulfur all the way from +6 to -2 (forming \( H_{2}S \)). You see purple vapors (Iodine) and smell rotten eggs (\( H_{2}S \)).

4. Identifying Halide Ions (The Silver Nitrate Test)

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

The Step-by-Step Process:

1. Add nitric acid first (this removes any "junk" ions like carbonates that might mess up the result).
2. Add a few drops of Silver Nitrate.
3. Observe the color of the precipitate (solid) formed.

The Results (Memory Aid):

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

Mnemonic: Cats Want Big Cups of Icy Yogurt.

The Ammonia Solubility Test

Sometimes the colors look too similar. We use Ammonia (\( NH_{3} \)) to be sure:

- Silver Chloride: Dissolves in dilute ammonia.
- Silver Bromide: Dissolves only in concentrated ammonia.
- Silver Iodide: Does not dissolve even in 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. This is called disproportionation.

Reaction with Water

\( Cl_{2}(g) + H_{2}O(l) \rightleftharpoons HCl(aq) + HClO(aq) \)

The Chloric(I) acid (\( HClO \)) produced is what kills bacteria, making our water safe to drink!

The Benefits vs. Risks

Benefits: It kills pathogens (bacteria) that cause diseases like cholera. It also stops algae growth.
Risks: Chlorine gas is toxic. It can also react with organic matter in water to form chlorinated hydrocarbons, which may cause cancer.

Did you know? Society generally agrees that the benefits of having clean, disease-free water far outweigh the small risks of the chemicals used.

Making Bleach

If you react Chlorine with cold, dilute Sodium Hydroxide (NaOH), you make common household bleach (Sodium Chlorate(I)).
\( Cl_{2} + 2NaOH \rightarrow NaCl + NaClO + H_{2}O \)

Section Summary: Key Takeaways

- Boiling points increase down Group 7; Electronegativity decreases.
- Oxidising power decreases down the group (F is strongest).
- Reducing power of halides increases down the group (I is strongest).
- Silver Nitrate identifies halides by color (White, Cream, Yellow).
- Chlorine is used to treat water because it forms \( HClO \), which kills bacteria.