The Halogens (OCR A-Level Chemistry A): Revision Notes

The Halogens

Introduction to the halogen group

The halogens form Group 17 (also known as Group 7) of the periodic table and represent the most reactive group of non-metallic elements. These elements do not exist naturally in their pure elemental form. Instead, they are found on Earth as stable halide ions ($\text{Cl}^-$, $\text{Br}^-$, and $\text{I}^-$) dissolved in seawater or present in solid mineral deposits such as sodium chloride (common salt, NaCl) found in salt mines.

Physical properties of the halogens

Physical states at room temperature and pressure

At room temperature and pressure (RTP), the halogens exist as diatomic molecules with the general formula $\text{X}_2$, where X represents any halogen element. The group demonstrates all three states of matter at RTP:

• Fluorine ($\text{F}_2$): Pale yellow gas
• Chlorine ($\text{Cl}_2$): Pale green gas
• Bromine ($\text{Br}_2$): Red-brown liquid (readily vaporises at room temperature)
• Iodine ($\text{I}_2$): Shiny grey-black solid
• Astatine ($\text{At}_2$): Never been observed (radioactive with extremely short half-life)

Trend in boiling points

The boiling points of the halogens increase progressively down the group. This trend can be understood by examining the relationship between molecular size and intermolecular forces.

Halogen molecule	Number of electrons	Boiling point / °C	Appearance at RTP
$\text{F}_2$	18	-188	Pale yellow gas
$\text{Cl}_2$	34	-34	Pale green gas
$\text{Br}_2$	70	59	Red-brown liquid
$\text{I}_2$	106	184	Shiny grey-black solid
$\text{At}_2$	170	230	Never been seen

Explanation of the trend:

As you move down the halogen group, several changes occur:

• The number of electrons in each molecule increases
• This results in stronger London dispersion forces (temporary induced dipole-dipole interactions) between molecules
• Greater energy is required to overcome these stronger intermolecular forces
• Therefore, the boiling point increases

Redox chemistry of the halogens

Electron configuration of halogens and halides

Each halogen atom contains seven electrons in its outer shell, with the general electron configuration ending in $ns^2np^5$ (where n represents the principal quantum number). This means halogens are one electron short of achieving the stable electron configuration of a noble gas.

During chemical reactions, halogen atoms readily gain one electron to form halide ions with a negative charge. The electron configuration of these ions ends in $ns^2np^6$, which corresponds to a complete outer shell with the stable electronic structure of the nearest noble gas.

Halogens as oxidising agents

Redox reactions represent the most common type of chemical reaction for halogens. In these reactions, each halogen atom undergoes reduction by gaining one electron to form a halide ion ($\text{X}^-$).

The general half-equation for this process is:

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

For chlorine specifically:

$\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-$

In this reaction, chlorine is reduced (gains electrons). Since the halogen causes another species to lose electrons (become oxidised), the halogen acts as an oxidising agent.

Halogen-halide displacement reactions

Experimental procedure and observations

Displacement reactions between halogens and halide ions can be performed on a test-tube scale to investigate the relative reactivity of different halogens. In these experiments:

1. A solution of each halogen is added to aqueous solutions containing different halide ions
2. If the halogen added is more reactive than the halide present in solution, a reaction occurs
3. The halogen displaces the halide from solution, with observable colour changes

Colour of halogen solutions

The appearance of halogens differs depending on the solvent used:

In water:

• $\text{Cl}_2$: Pale green solution
• $\text{Br}_2$: Orange solution
• $\text{I}_2$: Brown solution

In cyclohexane (non-polar organic solvent):

• $\text{Cl}_2$: Pale green solution
• $\text{Br}_2$: Orange solution
• $\text{I}_2$: Violet solution

The non-polar halogens dissolve more readily in cyclohexane than in water, making the colours more distinct and easier to distinguish - particularly the deep violet colour of iodine.

Displacement reaction results

The table below summarises the displacement reactions between aqueous halogen solutions and halide ions:

Halide ion	With $\text{Cl}_2(aq)$	With $\text{Br}_2(aq)$	With $\text{I}_2(aq)$
$\text{Cl}^-(aq)$	No reaction	No reaction	No reaction
$\text{Br}^-(aq)$	Orange colour from $\text{Br}_2$ formation $\text{Cl}_2(aq) + 2\text{Br}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{Br}_2(aq)$	No reaction	No reaction
$\text{I}^-(aq)$	Violet colour from $\text{I}_2$ formation $\text{Cl}_2(aq) + 2\text{I}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{I}_2(aq)$	Violet colour from $\text{I}_2$ formation $\text{Br}_2(aq) + 2\text{I}^-(aq) \rightarrow 2\text{Br}^-(aq) + \text{I}_2(aq)$	No reaction

Key observations:

• Chlorine reacts with both bromide and iodide ions (displaces both)
• Bromine reacts only with iodide ions (displaces only iodine)
• Iodine does not react with any halide ions (cannot displace anything)

This establishes the reactivity order:

$\text{Chlorine (most reactive)} > \text{Bromine} > \text{Iodine (least reactive)}$

Understanding displacement reactions as redox processes

Displacement reactions are redox reactions where the halogen is reduced and the halide ion is oxidised.

Fluorine and astatine

Fluorine is a pale yellow gas that reacts vigorously with almost any substance it contacts, making it the most reactive halogen. Due to its extreme reactivity, fluorine is rarely encountered in displacement reactions at A-Level.

Astatine is extremely rare because it is radioactive and decays rapidly. The element has never been observed directly and is predicted to be the least reactive halogen based on group trends.

Trend in reactivity of the halogens

In redox reactions, halogens act by gaining electrons. The reactivity trend shows that the tendency to gain electrons decreases down the group, making the halogens progressively less reactive.

Factors affecting the reactivity trend:

Halogen	Number of inner electron shells	Trend
$\text{F}_2$	1	Atomic radius increases
$\text{Cl}_2$	2	More inner shells provide greater shielding
$\text{Br}_2$	3	Reduced nuclear attraction to capture an electron
$\text{I}_2$	4	Reactivity decreases
$\text{At}_2$	5

Explanation:

As you descend the halogen group:

1. Atomic radius increases - the outer electrons are further from the nucleus
2. Number of inner electron shells increases - causing greater shielding of the nuclear charge
3. Nuclear attraction decreases - the nucleus experiences reduced ability to attract and capture an additional electron from another species
4. Reactivity as an oxidising agent decreases - it becomes progressively more difficult for the halogen to gain an electron

Disproportionation reactions

What is disproportionation?

Disproportionation is a specific type of redox reaction in which the same element is simultaneously both oxidised and reduced. Chlorine undergoes disproportionation when it reacts with water and with cold, dilute sodium hydroxide solution.

Reaction of chlorine with water

When small amounts of chlorine are added to water, a disproportionation reaction occurs:

$\text{Cl}_2(aq) + \text{H}_2\text{O}(l) \rightarrow \text{HClO}(aq) + \text{HCl}(aq)$

The two products formed are both acids:

• Chloric(I) acid (hypochlorous acid, HClO)
• Hydrochloric acid (HCl)

Reaction of chlorine with cold, dilute sodium hydroxide

The reaction of chlorine with water has limited use because chlorine has low solubility in water. When the water contains dissolved sodium hydroxide, much more chlorine dissolves and another disproportionation reaction occurs:

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

The resulting solution contains a high concentration of chlorate(I) ions ($\text{ClO}^-$) from the sodium chlorate(I) (sodium hypochlorite, NaClO) that forms. This solution serves as household bleach, which is manufactured by reacting chlorine with cold dilute aqueous sodium hydroxide.

Uses of chlorine in water purification and bleach

Chlorine in water treatment

Chlorine began to be used widely as a disinfectant for drinking water treatment over 100 years ago, revolutionising public health by dramatically reducing the incidence of waterborne diseases caused by harmful bacteria.

In water treatment facilities and swimming pools, chlorine is added to water in controlled amounts where it undergoes the disproportionation reaction described above. Chlorine tablets are often used in swimming pools and can be employed to purify water supplies following natural disasters.

Benefits and risks of chlorine use

Benefits:

• Ensures drinking water is safe to consume
• Kills harmful bacteria effectively
• Prevents diseases such as typhoid and cholera
• Essential for emergency water treatment after natural disasters

Risks:

• Chlorine is an extremely toxic gas
• Acts as a respiratory irritant even in small concentrations
• Large concentrations can be fatal
• Can react with organic hydrocarbons (such as methane from decaying vegetation) to form chlorinated hydrocarbons, which are suspected carcinogens

Tests for halide ions

Precipitation reactions with aqueous silver ions

Aqueous halide ions react with aqueous silver ions to form precipitates of silver halides. The general equation is:

$\text{Ag}^+(aq) + \text{X}^-(aq) \rightarrow \text{AgX}(s)$

where $\text{X}^-(aq)$ represents any halide ion in aqueous solution.

Halide ions as reducing agents

In displacement reactions between halogens and halide ions, the halogen gains electrons (is reduced) while the halide ion loses electrons (is oxidised). Therefore, halide ions act as reducing agents in these reactions.

The reducing ability of halide ions can be demonstrated through their reactions with sulfuric acid ($\text{H}_2\text{SO}_4$), which is a strong oxidising agent.

Chloride ions are not powerful enough to reduce sulfuric acid.

Bromide ions are more powerful reducing agents and can reduce sulfuric acid to sulfur dioxide ($\text{SO}_2$):

$2\text{H}^+ + \text{H}_2\text{SO}_4 + 2\text{Br}^- \rightarrow \text{SO}_2 + \text{Br}_2 + 2\text{H}_2\text{O}$

Notice that both the atoms and the charges balance in this equation.

Iodide ions are even more powerful reducing agents. They reduce the sulfur dioxide formed initially to sulfur (S), which is then reduced further to hydrogen sulfide ($\text{H}_2\text{S}$).