A Level Chemistry (9701): The Reactions of Chlorine ($\text{Cl}_2$)

Welcome to this important section on the chemistry of Group 17, specifically focusing on chlorine. Chlorine is a fascinating and reactive element, used everywhere from producing plastics to keeping swimming pools clean. Understanding its reactions, especially those involving cold and hot alkali, is key because they introduce the concept of disproportionation, a highly tested idea in A Level Chemistry!

Don't worry if oxidation states are tricky—we'll break down how to track them step-by-step to prove these complex reactions are really just chlorine reacting with itself.

1. Oxidation States: A Quick Review

Before diving into the reactions, let's remember the oxidation states (O.S.) of chlorine:

  • Elemental State: Chlorine gas, $\text{Cl}_2$, has an oxidation state of 0.
  • Chloride ion: In typical ionic compounds (like $\text{NaCl}$), chlorine has an oxidation state of -1.
  • Chlorates: In compounds containing oxygen (like $\text{ClO}^-$ or $\text{ClO}_3^-$), chlorine can have positive oxidation states (+1, +3, +5, +7).

A disproportionation reaction is one where a single element is simultaneously oxidised (O.S. increases) and reduced (O.S. decreases) in the same reaction.

Think of it like an element fighting itself!

2. Reaction of Chlorine with Cold Aqueous Sodium Hydroxide ($\text{NaOH}$)

The Reaction and Products

When chlorine gas is bubbled into cold, dilute, aqueous sodium hydroxide ($\text{NaOH}$), the reaction is rapid. This reaction is used industrially to manufacture household bleach (sodium chlorate(I)).

The products are sodium chloride ($\text{NaCl}$), sodium chlorate(I) ($\text{NaOCl}$), and water ($\text{H}_2\text{O}$).

Balanced Chemical Equation (Cold, Dilute $\text{NaOH}$):

$$ \text{Cl}_2(\text{g}) + 2\text{NaOH}(\text{aq}) \to \text{NaCl}(\text{aq}) + \text{NaOCl}(\text{aq}) + \text{H}_2\text{O}(\text{l}) $$

Interpretation as Disproportionation

Let's track the oxidation states of chlorine:

  1. Starting Material ($\text{Cl}_2$): O.S. = 0.
  2. Product 1 ($\text{NaCl}$): Chlorine is $\text{Cl}^-$. O.S. = -1.
    (The O.S. decreased from 0 to -1, so this is Reduction.)
  3. Product 2 ($\text{NaOCl}$): Oxygen is -2, Sodium is +1. Therefore, Chlorine must be +1.
    (The O.S. increased from 0 to +1, so this is Oxidation.)

Since the oxidation state of chlorine changes from 0 to both -1 and +1, chlorine has been both reduced and oxidised. This confirms it is a disproportionation reaction.

Key Takeaway (Cold $\text{NaOH}$)

The reaction with cold alkali produces the lower positive oxidation state of chlorine, $\text{ClO}^-$ (O.S. = +1).

3. Reaction of Chlorine with Hot Aqueous Sodium Hydroxide ($\text{NaOH}$)

If you heat the alkali, the reaction changes! The chlorate(I) ($\text{NaOCl}$) formed initially decomposes further when hot, leading to a product where chlorine has a much higher oxidation state.

The Reaction and Products

When chlorine gas is bubbled into hot, concentrated, aqueous sodium hydroxide, the products are sodium chloride ($\text{NaCl}$), sodium chlorate(V) ($\text{NaClO}_3$), and water ($\text{H}_2\text{O}$).

Balanced Chemical Equation (Hot, Concentrated $\text{NaOH}$):

$$ 3\text{Cl}_2(\text{g}) + 6\text{NaOH}(\text{aq}) \to 5\text{NaCl}(\text{aq}) + \text{NaClO}_3(\text{aq}) + 3\text{H}_2\text{O}(\text{l}) $$

Interpretation as Disproportionation (Again!)

Let's track the oxidation states of chlorine again:

  1. Starting Material ($\text{Cl}_2$): O.S. = 0.
  2. Product 1 ($\text{NaCl}$): Chlorine is $\text{Cl}^-$. O.S. = -1.
    (Reduction: 0 $\to$ -1)
  3. Product 2 ($\text{NaClO}_3$): Oxygen is -2 (total -6), Sodium is +1 (total +1). Therefore, Chlorine must be +5.
    (Oxidation: 0 $\to$ +5)

Again, chlorine is both reduced and oxidised, making this another disproportionation reaction. The higher temperature drives the formation of the more stable (under these conditions) chlorate(V) ion.

Memory Aid: Temperature and O.S.

Cold = Chlorate(I) ($\text{+1}$)
Hot = Higher O.S. (+5)

Common Mistake Alert!

Do not confuse $\text{NaOCl}$ (Chlorate(I)) with $\text{NaClO}_3$ (Chlorate(V)). Remember to use Roman numerals when stating the oxidation number of chlorine in the products, e.g., Sodium Chlorate(I) and Sodium Chlorate(V).

4. The Essential Use of Chlorine: Water Purification

One of the most vital applications of chlorine chemistry is in purifying drinking water. Chlorine (or compounds that release it) is added to municipal water supplies to kill pathogenic bacteria and microorganisms.

The Reaction of Chlorine with Water

When chlorine is added to water, it undergoes a disproportionation reaction with the water molecules itself:

$$ \text{Cl}_2(\text{aq}) + \text{H}_2\text{O}(\text{l}) \rightleftharpoons \text{HOCl}(\text{aq}) + \text{HCl}(\text{aq}) $$

The Active Species

The substances responsible for killing bacteria are the products of this reaction:

  • Hypochlorous Acid ($\text{HOCl}$): This is the main, highly effective germicidal agent. Chlorine here has an O.S. of +1.
  • Hypochlorite Ion ($\text{ClO}^-$): $\text{HOCl}$ is a weak acid and partially dissociates in water, especially at higher pH: $$ \text{HOCl}(\text{aq}) \rightleftharpoons \text{H}^+(\text{aq}) + \text{ClO}^-(\text{aq}) $$ The $\text{ClO}^-$ ion is also effective at killing bacteria, though $\text{HOCl}$ is generally considered the more potent species.

How it Kills Bacteria

Both $\text{HOCl}$ and $\text{ClO}^-$ are powerful oxidising agents. They penetrate the cell walls of microorganisms and disrupt their internal biochemical processes (like enzyme function or DNA structure), effectively killing them.

Controlling the pH (Le Chatelier's Principle in Action)

The efficiency of purification depends heavily on the concentration of the active species, $\text{HOCl}$.

  • If the pH of the water is too low (acidic), the equilibrium favours $\text{HOCl}$. This is good for disinfection but can be corrosive.
  • If the pH of the water is too high (alkaline), the equilibrium shifts towards the $\text{ClO}^-$ ion. While $\text{ClO}^-$ still disinfects, it is less effective than $\text{HOCl}$, meaning higher concentrations are needed.

Water treatment plants carefully control the pH to maximise the concentration of $\text{HOCl}$ while keeping the water safe to drink.

Did You Know?

The "chlorine smell" you associate with swimming pools isn't actually free chlorine. It's often caused by compounds called chloramines, which form when chlorine reacts with nitrogen-containing compounds (like sweat and urine) in the water. This is why good water management systems aim to keep the level of free, active chlorine ($\text{HOCl}$) high enough to be effective but low enough to be comfortable.

Quick Review Box: Summary of Chlorine Reactions

| Reagent & Temperature | Product (Chlorine O.S.) | Reaction Type | Use/Significance |

| Cold $\text{NaOH}$ | $\text{NaOCl}$ (O.S. +1) | Disproportionation | Manufacture of household bleach |

| Hot $\text{NaOH}$ | $\text{NaClO}_3$ (O.S. +5) | Disproportionation | Industrially important synthesis |

| Water ($\text{H}_2\text{O}$) | $\text{HOCl}$ (O.S. +1) | Disproportionation | Water purification (active species) |

Conclusion

The reactions of chlorine, particularly with alkali and water, illustrate key chemical concepts like oxidation states and disproportionation. Remember that the temperature dictates the product when reacting with $\text{NaOH}$, while the reaction with water produces the essential sterilising agent, $\text{HOCl}$. Mastering the oxidation state changes is the key to interpreting these important inorganic reactions!