1. Chapter 6
OXIDATION & REDUCTION

3. REDOX oxidation Reduction loss of electrons gain of electrons
(RIG) or (GER)
(OIL) or (LEO)

4. Reduction Oxidation Oxidant + e- Product Product + e- Reductant
( Gain of Electrons) (oxidation number Decreases)
Br + e-
Br -
Oxidation
Product + e-
Reductant
( Loss of Electrons) (oxidation number Increases) K
K+ + e-

5. (electron transfer reactions)
2Mg (s) + O2 (g) MgO (s)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e Mg2+ + 2O2- + 4e-
2Mg + O MgO

6. 4.4

7. Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e-
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e Cu
Cu2+ is reduced
CuSO4 is the oxidizing agent
4.4

8. Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)
Cu Cu2+ + 2e-
Cu is the reducing agent
Cu is oxidized
Ag+ + 1e Ag
Ag+ is reduced
AgNO3 is the oxidizing agent

9. 2FeCl2(aq) + Cl2(g) 2FeCl3(aq)
The ionic equation
2Fe+2(aq) + 4Cl-(aq) + Cl2(g) Fe+3(aq) + 6Cl-(aq)
Spectator ions are removed
2Fe+2(aq) + Cl2(g) Fe+3(aq) + 2Cl-(aq)
oxidised
reduced
The half equations for the reaction
2 Fe Fe e (half equation)
Cl e Cl (half equation)
2Fe Cl Fe Cl- (redox equation)

10. Q2: Look at the following reactions, for each one write down the ionic equation. Decide if the equation is a redox reaction. If it is a redox reaction write down the two half equations and state what has been oxidised and what has been reduced.
2Mg(s) + O2(g) MgO(s)
Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

11. Oxidation; increase in the oxidation number
2Fe Cl Fe Cl-
Reduction; decrease in the oxidation number

12. The differences between neutralization and redox reactions
Neutralisation
Redox reactions
Involves proton transfer
An acid is a proton donor
A base is a proton acceptor
Involve electron transfer
An oxidising agent is an electron acceptor
A reducing agent is an electron donor
Oxidation is the removal of electrons.
Reduction is a gain of electrons.

13. Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na (s), Be, K(s), Pb (s), H2 (g), O2 (g) , Hg (l), P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.
4.4

14. HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4
The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
HCO3-
Oxidation numbers of all the elements in HCO3- ?
O = -2
H = +1
3x(-2) ? = -1
C = +4
4.4

15. Rules for Assigning Oxidation States

16. Figure 4.10 The oxidation numbers of elements in their compounds
4.4

17. IF7 F = -1 7x(-1) + ? = 0 I = +7 K2Cr2O7 NaIO3 Na = +1 O = -2 O = -2
Oxidation numbers of all the elements in the following ?
F = -1
7x(-1) + ? = 0
I = +7
K2Cr2O7
NaIO3
Na = +1
O = -2
O = -2
K = +1
3x(-2) ? = 0
7x(-2) + 2x(+1) + 2x(?) = 0
I = +5
Cr = +6
4.4

18. Example
What is the oxidation numbers of Cr and N in K2Cr2O7 and HNO3 respectively?
K2Cr2O7
HNO3 +1
- 2 +1
- 2
Cr in K2Cr2O7
2 (+1) + 2n + 7 (-2) = 0
n = +6
N in HNO3
(+1) + n + 3(-2) =0
n = +5

19. Q 3: Find the oxidation numbers of:
Cr in CrO42-
ii. Mn in Mn(s)
iii. Fe in FeCl3
iv. P in P4O6
v. Br in Br2(l)
CrO42-
4x(-2) + ? = -2
Cr = +6
O = -2
Fe = +3
3x(-1) + ? = 0
P = +3
6x(-2) + 4n = 0

20. Oxidising and reducing agents
oxidising agent
Reducing agent
an element or compound
that oxidises
another element or compound
an element or compound
that reduces
another element or compound
It is reduced
It is oxidised

21. Oxidising and reducing agents
2FeCl2(aq) + Cl2(g) FeCl3(aq)
reducing agent
oxidising agent
NOTE
when the oxidising and reducing agent is named, the whole compound is specified, not just the element that undergoes the change in the oxidation state.

22. Q4: In the following reactions state which is the oxidising agent and which is the reducing agent. (It will help to write down the half equations)
2Mg(s) + O2(g) MgO(s)
Fe(s) + CuSO4(aq) FeSO4(aq) + Cu(s)
Cl2(g) + 2KI(aq) KCl(aq) + I2(s)

23. Oxidising agents Reducing agents 1. Liberate iodine from aqueous
potassium iodide solution
O.A
2I-(aq) I2(aq) + 2e-
colourless brown
Turn starch-iodide paper blue
Decolorize acidified potassium permanganate (VII) paper:
R.A
MnO-4(aq) + 8H+ + 5e Mn+2 + 4H2O
Purple colourless
2. Convert hydrogen sulphide to sulphur:
S-2(aq) S(s) + 2e-
Turn potassium dichromate (VI) paper green
Cr2O7-2(aq) +14H+(aq) + 6e Cr+3(aq +7H2O
Orange Green
3. Convert Fe(II) salts to Fe(III):
Fe+3 ions give a dark-blue precipitate with potassium hexacyanoferrate(II) (ferrocyanide), K4 Fe(CN)6 , or a red coloration with aqueous potassium thiocyanate solution, KSCN.
3. Convert Fe(III) salts to Fe(II): Fe+2 salts give a dark-blue precipitate with potassium hexacyano ferrate(III) (ferricyanide), K3Fe(CN)6

24. Q5: Classify the following processes as redox (electron transfer),
precipitation (ion transfer) or neutralisation (proton transfer) reactions:
2AgNO3(aq) + K2 CrO4(aq) Ag2 CrO4(s) + 2KNO3(aq)
2Al(s) + Fe2 O3(s) Al2 O3(s) + 2Fe(s)
3Fe(s) + 4H2 O(g) Fe3 O4(s) + 4H2(g)
Ba(OH)2(aq) + 2HNO3(aq) Ba(NO3)2(aq) + 2H2O(l)
BaCl2(aq) + H2 SO4(aq) BaSO4(s) + 2HCl(aq)
Precipitation reaction
Redox reaction
Redox reaction
Neutralisation reactions
Precipitation reaction

25. Redox Process in Industry
Oxygen (air) as oxidising agent
Carbon or carbon monoxide as reducing agent
1. Methane is converted to methanol by direct oxidation
200 ˚C/copper tube
2CH4(g) + O2(g) CH3OH(l)

26. Redox Process in Industry
2. Sulphuric acid is manufactured by oxidising the sulphur dioxide
burn
S(s) + O2(g) SO2(g)
420 ○C / V2 O5
2SO2(g) + O2(g) SO3(g)
Then the product is absorbed in 98% acid and diluted with water.
SO3(g) + H2 O H2 SO4

27. Redox Process in Industry
3. Carbon and carbon monoxide are used to reduce oxide ores to the metal
Example:
Iron(III) oxide is reduced to iron in a blast furnace.
Fe2O3 + 3CO Fe + 3CO2
Iron (III) oxide is reduced to iron in a blast furnace

28. Redox Process in Industry
4- Aluminium oxide requires a more powerful reducing agent than either carbon or carbon monoxide. Aluminium is obtained from its ore by electrolytic reduction.
At cathode
Al e Al
Electrolytic reduction of Aluminium