Trends in Group 7 Properties

The Group 7 elements (fluorine, chlorine, bromine, iodine, and astatine) are known as the halogens. They are highly reactive non-metals with seven valence electrons in their outermost shell, giving them an outer configuration of \(ns^2 np^5\). In chemical reactions, they typically gain one electron to form stable \(1-\) halide ions with a full outer energy level.

1. Physical States and Boiling Points

At room temperature, the halogens exist as diatomic molecules (\(\text{F}_2\), \(\text{Cl}_2\), \(\text{Br}_2\), \(\text{I}_2\)). Their physical states and appearances change down the group:

• Fluorine (\(\text{F}_2\)): Pale yellow gas.
• Chlorine (\(\text{Cl}_2\)): Greenish-yellow gas.
• Bromine (\(\text{Br}_2\)): Red-brown liquid that vaporises easily to an orange gas.
• Iodine (\(\text{I}_2\)): Dark grey crystalline solid that sublimes to form a purple vapour.

The boiling points of the halogens increase down the group. This is because the molecules grow larger and contain more electrons. A greater number of electrons leads to stronger temporary dipoles, which results in stronger van der Waals (London) forces between the diatomic molecules. More thermal energy is required to overcome these intermolecular forces. Note that no covalent bonds are broken when halogens boil.

2. Trend in Electronegativity

Electronegativity is the power of an atom to attract a bonding pair of electrons in a covalent bond. Electronegativity decreases down Group 7, with fluorine being the most electronegative element on the Pauling scale (value of 4.0).

Explanation:

1. Increasing Atomic Radius: Descending the group, the outer electrons reside in shells further from the positive nucleus.
2. Increasing Shielding: There are more inner electron shells to shield the bonding electrons from the pull of the nucleus.
3. Weaker Pull: Although the nuclear charge (proton number) increases down the group, the effects of increased distance and electron shielding far outweigh this. Thus, the electrostatic attraction between the nucleus and the shared bonding pair of electrons is weaker.

3. Oxidising Ability of the Halogens

Halogens act as oxidising agents in chemical reactions. They gain one electron from another substance and are reduced to form halide ions:

\[ \text{X}_2 + 2\text{e}^- \rightarrow 2\text{X}^- \]

The oxidising ability of the halogens decreases down the group. Fluorine is the most powerful oxidising agent in Group 7.

Explanation:

As you go down the group, atoms become larger and have more shielding. This means that a incoming electron entering the valence shell of a smaller atom (like fluorine or chlorine) experiences a much stronger electrostatic attraction to the nucleus because it is closer to the nucleus and less shielded. Therefore, smaller halogens attract and gain electrons more readily, making them stronger oxidising agents.

4. Reducing Ability of Halide Ions

Halide ions (\(\text{F}^-\), \(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\)) act as reducing agents. They can donate an electron to another substance, converting back into neutral halogen molecules:

\[ 2\text{X}^- \rightarrow \text{X}_2 + 2\text{e}^- \]

The reducing ability of halide ions increases down the group. Iodine ions (\(\text{I}^-\)) are the strongest reducing agents, whereas fluoride ions (\(\text{F}^-\)) are the weakest.

Explanation:

1. Larger Ionic Radius: Down the group, the halide ions have more electron shells, meaning the outer electron is located further from the positive nucleus.
2. Increased Shielding: There are more shells of inner electrons shielding the outer electron from the nucleus.
3. Weaker Electrostatic Attraction: The combination of increased distance and shielding means the outer electron is held less tightly to the positive nucleus. Thus, it is lost more easily, making larger halide ions better reducing agents.