Reactions of Chlorine and Uses of Chlorine — AQA A-level Chemistry

SPEC 3.2.3.2 • GROUP 7(17), THE HALOGENS

Chlorine & Chlorate(I): Reactions & Uses

Disproportionation, water treatment, bleach, and the chlorine-alkali reaction

Learning Objectives

By the end of this lesson, you will be able to:

What You'll Learn

Write equations for chlorine reacting with water and with cold dilute alkali

Explain disproportionation and identify it in reactions

Describe the uses of chlorine and chlorate(I) ions

Discuss the benefits and risks of water treatment with chlorine

Assign oxidation states to identify disproportionation

Exam Focus

This topic frequently appears in exams asking you to write equations, identify oxidation state changes, and discuss the ethical considerations of water chlorination. Make sure you can explain both the chemistry AND the real-world applications.

What is Disproportionation?

Before we look at chlorine's reactions, we need to understand a special type of redox reaction called disproportionation. This concept is central to understanding why chlorine behaves the way it does in water and alkali.

Definition

Disproportionation is a redox reaction in which the same element is simultaneously oxidised AND reduced.

In other words, atoms of one element in a single oxidation state split into two different oxidation states - some go up (oxidation) and some go down (reduction).

Think of it like this: imagine a group of people all standing on the same step of a staircase. In disproportionation, some people step UP while others step DOWN - they start together but end up at different levels. The "steps" are oxidation states.

Key Features

Same element - the element being oxidised is the same as the one being reduced

Single starting oxidation state - atoms begin with identical oxidation states

Two different products - atoms end up in at least two different oxidation states

Disproportionation vs Regular Redox

It's important to distinguish disproportionation from normal redox reactions:

FeatureRegular RedoxDisproportionation

OxidationOne species oxidisedSame element oxidised...

ReductionDifferent species reduced...AND reduced

ExampleZn + Cu²⁺ → Zn²⁺ + CuCl₂ → Cl⁻ + Cl⁺ (in compounds)

In the zinc-copper example, zinc is oxidised and copper is reduced - two different elements. But in chlorine's reactions, the same chlorine atoms both gain AND lose electrons.

Comproportionation

The opposite of disproportionation is comproportionation - where atoms of the same element in different oxidation states combine to form a single oxidation state. You saw this in Lesson 3.01 with the iodate-iodide reaction.

Chlorine Reacting with Water

When chlorine gas is bubbled through water, it partially dissolves and undergoes a reversible disproportionation reaction:

Equation

Cl₂(g) + H₂O(l) ⇌ HCl(aq) + HClO(aq)Ionic form:Cl₂(g) + H₂O(l) ⇌ ClO⁻(aq) + Cl⁻(aq) + 2H⁺(aq)

This reaction produces a mixture called chlorine water, which is pale yellow-green in colour due to dissolved Cl₂. The solution contains:

Hydrochloric acid (HCl) - a strong acid that fully dissociates

Chloric(I) acid (HClO) - a weak acid, also called hypochlorous acid

Why Reversible?

The reaction is reversible (⇌) because it doesn't go to completion. The equilibrium lies to the LEFT, meaning most of the chlorine remains as dissolved Cl₂ molecules. This is why chlorine water still smells strongly of chlorine.

Group Trend

The other halogens (Br₂, I₂) react with water similarly, but much more slowly going down the group. This is consistent with the decreasing oxidising ability of halogens down Group 7.

Oxidation State Analysis: Cl₂ + H₂O

Let's prove this is disproportionation by tracking the oxidation states of chlorine:

Oxidation State Summary

Cl₂: Chlorine has oxidation state 0 (element in its standard state)

HCl (Cl⁻): Chlorine has oxidation state −1 (reduced, gained 1 electron)

HClO: Chlorine has oxidation state +1 (oxidised, lost 1 electron)

Same element (Cl) simultaneously oxidised (0→+1) AND reduced (0→−1) = DISPROPORTIONATION

Chloric(I) Acid: The Active Agent

The chloric(I) acid (HClO) formed in this reaction is the key product because it has powerful oxidising and antimicrobial properties:

Chloric(I) Acid

HClO (also called hypochlorous acid) is a weak acid that partially dissociates:HClO(aq) ⇌ H⁺(aq) + ClO⁻(aq)The ClO⁻ ion is called the chlorate(I) ion (or hypochlorite ion).

HClO is responsible for the bleaching and disinfecting properties of chlorine water. It works by:

Oxidising coloured compounds - breaking down chromophores (colour-causing groups) to make substances colourless

Killing microorganisms - destroying bacteria, viruses, and other pathogens by oxidising their cell components

Common Mistake

Students often say "chlorine bleaches things" - but it's actually the HClO (chloric(I) acid) that does the bleaching, not Cl₂ itself. Chlorine must react with water first to produce HClO.

Examiner Tip

Adding chlorine to a water supply is an effective way to make the water safe to drink as it forms the sterilising agent HClO, which in turn dissociates in water into another sterilising agent, ClO⁻(aq). Both species act as disinfectants.

Chlorine + Cold Dilute Alkali

When chlorine gas reacts with cold, dilute sodium hydroxide solution, another disproportionation reaction occurs:

Equation

Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)Ionic equation:Cl₂(g) + 2OH⁻(aq) → Cl⁻(aq) + ClO⁻(aq) + H₂O(l)

This reaction is essentially the same as the reaction with water, but it goes to completion (not reversible) because:

The OH⁻ ions neutralise the H⁺ produced, shifting equilibrium to the right

Both products (NaCl and NaClO) are stable in alkaline solution

Conditions Matter

The alkali must be cold (around 15°C) and dilute. If the alkali is hot and concentrated, a different reaction occurs producing chlorate(V) ions (ClO₃⁻) instead.

Oxidation States: Cl₂ + NaOH

Again, let's verify this is disproportionation:

SpeciesOxidation State of ClChange

Cl₂ (reactant)0Starting point

NaCl / Cl⁻ (product)−1Reduced (gained electron)

NaClO / ClO⁻ (product)+1Oxidised (lost electron)

In ClO⁻, we calculate the oxidation state of Cl as follows:

Overall charge = −1

Oxygen = −2

Therefore: Cl + (−2) = −1, so Cl = +1

Half-Equations

Oxidation half-equation:½Cl₂ + 2OH⁻(aq) → ClO⁻(aq) + H₂O + e⁻Reduction half-equation:½Cl₂ + e⁻ → Cl⁻

Pattern Recognition

Notice the symmetry: chlorine at oxidation state 0 splits equally - one chlorine goes up to +1, one goes down to −1. The changes are equal and opposite, which conserves electrons. The half-equations show that for every electron released in oxidation, one is consumed in reduction.

Uses of Sodium Chlorate(I)

The sodium chlorate(I) (NaClO) produced when chlorine reacts with sodium hydroxide has many practical applications:

Household Bleach

Household bleach is typically a dilute solution of sodium chlorate(I) (NaClO), usually about 3-8% concentration. The bleach contains chlorate(I) ions (ClO⁻) that act as oxidising agents to kill bacteria, making it widely useful for disinfecting and cleaning.

ApplicationHow It Works

Bleaching fabricsOxidises coloured compounds, breaking chromophores

Disinfecting surfacesKills bacteria and viruses by oxidising proteins

Swimming poolsSanitises water by killing microorganisms

Safety Warning

Never mix bleach (NaClO) with acids! This reverses the reaction and releases toxic chlorine gas: ClO⁻ + Cl⁻ + 2H⁺ → Cl₂ + H₂O.

Swimming Pool Chemistry

In shallow swimming pools, chlorine is rapidly lost from the water due to sunlight exposure. In sunlight, the following decomposition reaction occurs:2Cl₂(g) + 2H₂O(l) → 4HCl(aq) + O₂(g)This depletes the hypochlorite ion disinfectant, meaning pools require regular additions of chlorine.

Alternative Chlorination Method

Instead of adding chlorine gas directly, pools can use solid sodium chlorate(I) or calcium chlorate(I), which dissolves in water:NaClO(s) + H₂O ⇌ Na⁺(aq) + OH⁻(aq) + HClO(aq)In alkaline solution, this equilibrium shifts left and HClO is removed as ClO⁻ ions. To prevent this, swimming pools are kept slightly acidic to maintain effective levels of HClO.

Water Treatment with Chlorine

Chlorine is a poisonous gas and was notoriously used as a chemical weapon in the First World War. However, when dissolved in water in controlled amounts, it has become an essential part of our lives in the treatment of drinking water and swimming pools.

How It Works

Small amounts of chlorine (typically 0.2-0.5 mg/L) are added to water supplies. The chlorine reacts with water to form HClO:Cl₂ + H₂O ⇌ HCl + HClOThe HClO then dissociates:HClO ⇌ H⁺ + ClO⁻Both HClO and ClO⁻ ions act as disinfectants, killing microorganisms by oxidation.

The killing mechanism involves these species penetrating cell walls and oxidising essential enzymes and DNA, destroying the pathogen's ability to function and reproduce.

Benefits of Water Chlorination

The chlorination of water supplies is considered one of the greatest public health achievements of the 20th century:

BenefitExplanation

Disease preventionKills disease-causing microorganisms (bacteria, viruses)

Low costChlorine is cheap and easy to produce

Residual protectionPrevents reinfection further down the supply

Water qualityRemoves bad tastes, smells and discolouration

Historical Impact

Before water chlorination, epidemics of cholera and typhoid killed thousands regularly. In London's 1854 cholera outbreak, over 600 people died in just 10 days in one neighbourhood due to contaminated water.

Risks of Water Chlorination

While the benefits are substantial, there are some potential risks:

Risks

Chlorine gas toxicity: Chlorine gas causes severe irritation if inhaled

Chlorinated hydrocarbons: Chlorine reacts with organic matter to form potentially carcinogenic compounds (THMs such as chloroform, CHCl₃)

Alternative Disinfection Methods

Due to some risks of using chlorine, alternatives are sometimes used:MethodAdvantagesDisadvantagesOzone (O₃)Effectively kills microorganismsExpensive; decays rapidly (no residual protection)UV lightDamages microbial DNALess effective in cloudy water; no residual protection

Despite these alternatives, chlorine remains the most widely used method because it provides residual protection throughout the distribution system - something ozone and UV cannot offer.

The Ethical Consideration

Water chlorination presents a classic risk-benefit analysis:

Scientific Consensus

The benefits far outweigh the risks:

Waterborne diseases are certain killers without treatment

THM risks are possible and occur only at high levels

Modern treatment keeps THM levels well below harmful concentrations

Answering Exam Questions

AQA frequently asks 6-mark "discuss" questions about water chlorination:

Structure for 6-Mark Questions

State the purpose: Chlorine kills harmful microorganisms

Explain the chemistry: Cl₂ + H₂O → HCl + HClO; HClO is the disinfectant

List benefits: Prevents cholera/typhoid, cheap, residual protection

List risks: THMs may form, possible health concerns

Evaluate: Benefits outweigh risks

Conclude: Chlorination is justified but levels must be monitored

Common Mistakes

Only discussing benefits OR risks (not both)

Not including the actual chemistry/equations

Not reaching a conclusion

Worked Example 1

Worked Example01

Question: Bromine reacts with cold dilute sodium hydroxide in a similar way to chlorine. Predict the products of this reaction, write a balanced equation, and state the oxidation state changes for bromine.

Step 1: Apply the pattern from chlorine

Chlorine produces NaCl + NaClO, so by analogy bromine produces NaBr + NaBrO

Step 2: Write the balanced equation

Br₂(l) + 2NaOH(aq) → NaBr(aq) + NaBrO(aq) + H₂O(l)

Step 3: Identify oxidation states

•

In Br₂: Br = 0

•

In NaBr: Br = −1 (reduction)

•

In NaBrO: Br = +1 (oxidation)

Step 4: Confirm disproportionation

Bromine (0) is simultaneously reduced to −1 and oxidised to +1. This is disproportionation.

Worked Example 2

Worked Example02

Question: A swimming pool manager notices that chlorine levels drop more rapidly on sunny days than cloudy days. Using your knowledge of chlorine chemistry, explain this observation and suggest what the manager should do.

Explanation:

In sunlight (UV light), chlorine undergoes a different reaction with water:

2Cl₂(g) + 2H₂O(l) → 4HCl(aq) + O₂(g)

This depletes the chlorine and chlorate(I) ions that act as disinfectants. The oxygen escapes as gas.

Why sunny days are worse:

UV radiation provides energy for this decomposition reaction, so more chlorine is lost on sunny days compared to cloudy days.

What the manager should do:

•

Add chlorine more frequently on sunny days

•

Or add solid NaClO/Ca(ClO)₂ which dissolves to release HClO

•

Monitor chlorine levels more frequently during sunny periods

Practice

Question 1: A student mixes household bleach (containing NaClO) with an acidic toilet cleaner. They notice a pungent green-yellow gas is released. Write an ionic equation to explain what happened and explain why this is dangerous.

Ionic equation:

ClO⁻(aq) + Cl⁻(aq) + 2H⁺(aq) → Cl₂(g) + H₂O(l)

Explanation:

The acid provides H⁺ ions which reverse the disproportionation reaction. The green-yellow gas is chlorine (Cl₂).

Why dangerous:

Chlorine gas is toxic - it causes severe irritation to the respiratory system and can be fatal at high concentrations. This is why bleach and acids should never be mixed.

Practice

Question 2: Swimming pool water is maintained at a slightly acidic pH (around 7.2-7.6). Using the equilibrium NaClO + H₂O ⇌ Na⁺ + OH⁻ + HClO, explain why maintaining acidic conditions is important for effective disinfection.

The equilibrium:

NaClO + H₂O ⇌ Na⁺ + OH⁻ + HClO

Why acidic conditions help:

•

In acidic conditions, the H⁺ ions neutralise OH⁻ ions

•

This removes OH⁻ from the equilibrium

•

By Le Chatelier's principle, the equilibrium shifts RIGHT

•

More HClO (the active disinfectant) is produced

In alkaline conditions:

Excess OH⁻ shifts equilibrium LEFT, reducing HClO concentration and making disinfection less effective.

Practice

Question 3: Cl₂ + H₂O ⇌ HCl + HClO is reversible, but Cl₂ + NaOH goes to completion. Explain why.

When chlorine reacts with water, HCl dissociates to give H⁺ ions. The reaction is reversible as products can recombine.

With NaOH, the OH⁻ ions neutralise H⁺: H⁺ + OH⁻ → H₂O

This removes H⁺ from equilibrium, shifting it to the right (Le Chatelier's principle). The reaction goes to completion.

Summary

Key Equations

ReactionEquation

Cl₂ + WaterCl₂ + H₂O ⇌ HCl + HClO

Cl₂ + NaOHCl₂ + 2NaOH → NaCl + NaClO + H₂O

Ionic equationCl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O

Key Concepts

Disproportionation: Same element simultaneously oxidised AND reduced

Chlorine: 0 → −1 (reduced) AND 0 → +1 (oxidised)

HClO / ClO⁻: Active bleaching and disinfecting agent

Water treatment: Benefits outweigh risks (THM formation)