How does the trend in boiling points differ between Group 7 and Group 1?

In Group 7, boiling points increase down the group due to stronger London dispersion forces in larger molecules. In Group 1, boiling points decrease down the group because the metallic bonding weakens as the atomic radius increases.

When comparing the reaction of Chlorine with cold vs. hot alkali, what is the difference in the oxidation state of the chlorate product?

In cold alkali, the chlorate(I) ion ($ClO^-$) is formed where Chlorine is in the $+1$ state. In hot alkali, the chlorate(V) ion ($ClO_3^-$) is formed where Chlorine is in the $+5$ state.

How does the reducing power of an Iodide ion compare to a Chloride ion?

Iodide ions are much stronger reducing agents than Chloride ions. This is because the outer electrons in Iodide are further from the nucleus and more shielded, making them easier to lose during a redox reaction.

What error occurs if you use hydrochloric acid instead of nitric acid to acidify a sample before testing for halides?

Using hydrochloric acid introduces Chloride ions into the solution. These ions will react with the silver nitrate to form a white precipitate of $AgCl$, giving a false positive result regardless of which halides were originally present.

Why is it a mistake to say that 'reactivity increases down Group 7'?

Reactivity actually decreases down Group 7. Students often confuse this with Group 1; in Group 7, atoms need to gain an electron, which is harder for larger atoms with more shielding.

What is the danger of relying solely on the color of silver halide precipitates for identification?

The colors (white, cream, and pale yellow) are very similar and can be difficult to distinguish by eye. Ammonia must be used to confirm the identity based on solubility.

Define 'Disproportionation' in the context of halogen chemistry.

Disproportionation is a redox reaction where the same element in a single species is simultaneously oxidised and reduced to form two different products.

What are London dispersion forces and why are they relevant to halogens?

They are temporary dipole-induced dipole attractions between non-polar molecules. They increase with the number of electrons, explaining why larger halogens have higher boiling points.

What is the general ionic equation for the reaction of a halogen with a more reactive halide ion?

There is no reaction; a halogen can only displace a *less* reactive halide ion. The correct general equation for a successful displacement is $X_2 + 2Y^- \rightarrow 2X^- + Y_2$, where $X$ is higher in the group than $Y$.

Why does electronegativity decrease down Group 7?

As the atomic radius increases and shielding increases, the nucleus has a weaker electrostatic attraction for the shared pair of electrons in a covalent bond.

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2. Inorganic Chemistry

Group 7 Trends

Summary

The Group 7 elements, known as the halogens, exhibit distinct periodic trends in their physical and chemical properties. As non-metals that exist as diatomic molecules, their behavior is governed by increasing atomic size and electron shielding down the group, leading to predictable changes in boiling points, electronegativity, and redox reactivity.

1. Physical Properties and Volatility

Diatomic Nature: Halogens exist as simple molecular structures composed of two atoms ( $X_2$ ) held together by covalent bonds. At room temperature, their states transition from gases (Fluorine, Chlorine) to liquids (Bromine) and solids (Iodine).
Boiling Point Trend: The melting and boiling points increase down the group. This occurs because the number of electrons increases with atomic size, leading to stronger London dispersion forces (instantaneous dipole-induced dipole forces) between molecules that require more energy to overcome.
Volatility: Volatility refers to the ease with which a substance evaporates. Because intermolecular forces strengthen down the group, volatility decreases from Fluorine (most volatile) to Iodine (least volatile).

Graph showing the increase in boiling points of halogens as atomic number and molecular size increase.

2. Electronegativity and Atomic Radius

Definition: Electronegativity is the ability of an atom to attract a bonding pair of electrons in a covalent bond. It is a measure of the 'pull' the nucleus exerts on shared electrons.
The Trend: Electronegativity decreases down Group 7. As you move down the group, the atomic radius increases because more electron shells are added, placing the outer electrons further from the nucleus.
Shielding Effect: Increased inner electron shells provide more shielding, which reduces the effective nuclear charge felt by the bonding electrons. Consequently, the nucleus has a weaker attraction for incoming or shared electrons.

3. Reactivity and Oxidising Power

4. Disproportionation Reactions

5. Reducing Power of Halide Ions

6. Key Distinctions & Identification

7. Exam Strategy & Tips