1. 2.6.3
Redox Reactions of the
Halogens
i.e. Group VII elements

2. Introduction
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Redox comes from reduction and oxidation
Oxidation Is Loss of electrons
Reduction Is Gain of electrons
Remember this by OIL RIG
Halogens are powerful oxidising agents
They remove [take] electrons from other
elements to become negative ions

3. Experiment 1
Reaction of chlorine gas
with
potassium iodide

4. Safety Precautions
• KMnO4 – powerful oxidising agent
and irritant
• HCl - corrosive
• Cl2 - poisonous
• KI – irritant
• Wear safety glasses and rubber
gloves

5. Apparatus
• Test tubes
• Potassium manganate(VII)
crystals
• Conc. Hydrochloric acid
• Potassium iodide solution
• Litmus paper [blue]

6. Method
• Pour some KI solution into a test tube.
• Place some potassium manganate(VII)
crystals [KMnO4] in a test tube
• Add some conc. HCl
• Note what happens.
• Pour the chlorine gas into the potassium
iodide and shake.
• Note what happens.

7. Results
• A greenish yellow gas is produced when
the conc. HCl is dropped onto the KMnO4
• this gas is chlorine [Cl2]
• MnO4- + HCl = Mn2+ + H2O + Cl2
• It is heavier than air
• It has a characteristic smell [swimming
pool]
• It turns blue litmus red - then bleaches it
• When chlorine is poured into KI solution
turns from colourless to brown as iodine is
released.

8. •
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Cl2(g) + 2 KI(aq) = 2 KCl(aq) + I2(aq)
Another way of writing the equation is
Cl2(g) + 2 I-(aq) = 2 Cl- (aq) + I2(aq)
the chlorine has displaced the Iodine
and so is more reactive
The chlorine has been reduced [gained
an electron to become Cl -]
The iodide ion I- in KI has been oxidised
[lost an electron to become I2 ]
The chlorine is an oxidising agent.

9. Experiment 2
Reaction of
Cl2 with KBr

10. • Make chlorine as in last experiment
MnO4- + HCl = Mn2+ + H2O + Cl2
• Put some KBr solution into a test tube
• Pour the Cl2 into the test tube of KBr
• Note what happens
• The KBr solution turns from colourless
to red/brown as bromine is released

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Cl2(g) + 2 KBr(aq) = 2 KCl(aq) + Br2(aq)
the chlorine has displaced the Bromide
ion and so is more reactive
The chlorine has been reduced [gained
electron to become Cl-]
The bromide ion Br- in KBr has been
oxidised [lost a electron to become
Br2.]
The chlorine is an oxidising agent.
Iodine is darker than Bromine

12. Conclusions
• Non-metals
can also be
arranged in a
reactivity
series.
• F• Cl• Br• [O2-]
• I• [SO42-]
More
Reactive
Less
Reactive

13. Experiment 3
Reaction of
2+
Cl2 with Fe

14. Apparatus
• KMnO4 crystals
• Conc. HCl
• Iron (II) sulphate solution FeSO4

15. Method
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Make up a solution of iron(II)sulphate
it is pale green in colour
Make some chlorine gas
mix the gas and the iron(II)sulphate solution
Note what happens
the solution turns brown
2 Fe2+ + Cl2 = 2 Fe3+ + 2 ClGreen
Brown
• the Fe2+ has been oxidised to Fe3+

16. Key Points to remember
• OIL
• RIG
• F most reactive
• More reactive element displaces
less reactive from solution

17. Experiment 4
Reaction of Cl2 with Na2SO3

18. Apparatus
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KMnO4(s)
Conc. HCl
Sodium sulphite solution [Na2SO3]
Barium Nitrate Solution

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Make up a solution of sodium sulphite
Test it with Barium nitrate solution
A white precipitate forms
Add HCl and the precipitate re-dissolves
This tells us the solution is a sulphite
Take a fresh sample of the solution
react it with some chlorine
Add barium nitrate solution
Note what happens

20. • A white precipitate forms
• Add Hydrochloric acid to see what happens
to the precipitate
• The precipitate does not re-dissolve
• the sulphite has been oxidised to a sulphate
Cl2 + SO32- + H2O = 2 Cl- + SO42- + 2H+
• look at oxidation numbers
• S goes from +4 to +6 so oxidised by Cl2