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

2. Introduction 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. Reaction of chlorine gas with potassium iodide
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.

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. 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

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

13. Experiment 3
Reaction of Cl2 with Fe2+

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

15. Method 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 Fe Cl2 = 2 Fe Cl-
Green 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. Reaction of Cl2 with Na2SO3
Experiment 4
Reaction of Cl2 with Na2SO3

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

19. 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- + SO H+
look at oxidation numbers
S goes from +4 to +6 so oxidised by Cl2