The Halogens (Leaving Cert Chemistry): Revision Notes

The Halogens

Introduction to Group 17

The halogens are a family of elements found in Group 17 of the Periodic Table. The name "halogen" comes from the Greek language and means "producer of salt". This name perfectly describes these elements because they readily react with metals to produce various types of salts.

The photograph above shows three of the most common halogens in their natural states, demonstrating their distinctive colours and physical properties.

Physical properties

States of matter at room temperature

The halogens exist in different physical states at room temperature, which helps us understand the London dispersion forces between their molecules:

• Fluorine (F₂): Pale yellow gas
• Chlorine (Cl₂): Pale green-yellow gas
• Bromine (Br₂): Dark red volatile liquid
• Iodine (I₂): Purple-black solid

The fact that bromine is a liquid at room temperature shows it has stronger intermolecular forces than the gaseous halogens. Bromine vapour fills its container because it has a relatively low boiling point and evaporates easily to form a reddish-brown vapour.

Electrical and thermal conductivity

All halogens are poor conductors of both electricity and heat. This is because they are non-metals and do not contain free-moving electrons that can carry electrical current or thermal energy effectively.

Sublimation

An interesting property of iodine is that it undergoes sublimation when heated. Sublimation means the substance changes directly from a solid to a gas without passing through the liquid phase. This creates distinctive purple vapours that are easy to observe in laboratory demonstrations.

Chemical properties and reactions

The table above summarises the key chemical reactions of the halogens. Let's examine these patterns:

Reactions with sodium metal

All halogens react vigorously with sodium metal to produce sodium halides (salts):

Reactions with hydrogen gas

Halogens also react with hydrogen gas to form hydrogen halides, which dissolve in water to form acids:

• Fluorine: $H_2 + F_2 \rightarrow 2HF$ (hydrogen fluoride)
• Chlorine: $H_2 + Cl_2 \rightarrow 2HCl$ (hydrogen chloride)
• Bromine: $H_2 + Br_2 \rightarrow 2HBr$ (hydrogen bromide)
• Iodine: $H_2 + I_2 \rightarrow 2HI$ (hydrogen iodide)

Reactivity trends

Why reactivity increases up the group

Unlike many other groups in the Periodic Table, the reactivity of halogens increases as you move up the group. This means:

This trend occurs because of three key factors:

1. Atomic radius decreases up the group

As you move up Group 17, the atoms become smaller. This means the outermost electrons are closer to the positively charged nucleus.

2. Screening effect decreases up the group

There are fewer inner electron shells to shield the outermost electrons from the nuclear attraction as you move up the group.

3. First ionisation energy increases up the group

The combination of smaller atomic radius and less screening means it becomes easier for halogens higher up the group to attract and gain electrons from other substances.

Displacement reactions

One of the most important concepts in halogen chemistry is displacement reactions. A more reactive halogen can displace a less reactive halogen from its compounds.

The equation above shows chlorine displacing bromine from bromide ions:

General rule for displacement reactions

Why displacement occurs

Displacement happens because the more electronegative halogen has a stronger attraction for electrons. For example, chlorine atoms (electronegativity = 3.16) have a greater attraction for electrons than bromine atoms (electronegativity = 2.96). Therefore, chlorine can take electrons from bromide ions, converting them to bromine molecules while the chlorine becomes chloride ions.