Redox reactions
Oxidation
Gain oxygen

2Mg(s) + O2(g)  2MgO(s)
Gain oxygen

CH4(g) + O2(g)  CO2(g) + H2O()
Loss electrons

Mg(s)  Mg2+(aq) + 2e-
Reduction
Loss oxygen

2PbO(s) + C(s)  2Pb(s) + CO2(g)
Gain electrons

Pb2+(aq) + 2e-  Pb(s)
Oxidation and Reduction
Oxidation
Reduction
Gains oxygen
Loses oxygen
Loses electrons
Gains electrons
Oxidation number
increases
Oxidation number
decreases
Oxidizing agent

Oxidizing agent


A substance that causes oxidation is called an
oxidizing agent.
An oxidizing agent is a substance that gains electrons.
oxidatio
n
2PbO(s) + C(s)  2Pb(s) + CO2(g)
oxidizing agent

Pb2+(aq) + 2e-  Pb(s)
oxidizing agent

Reducing agent

Reducing agent


A substance that causes reduction is called a reducing
agent.
A reducing agent is a substance that loses electrons.
reduction

2PbO(s) + C(s)  2Pb(s) + CO2(g)
reducing agent
Pb(s)  Pb2+(aq) + 2ereducing agent

Oxidizing agent and Reducing
agent
Oxidizing agent
Reducing agent
Undergoes reduction
Undergoes oxidation
Loses oxygen
Gains oxygen
Gains electrons
Loses electrons
Oxidation number
decreases
Oxidation number
increases
Oxidation numbers

The oxidation number of an element is
zero.


oxidations numbers of N in N2, O in O2, S, Na,
K, Cu are all zero.
The oxidation number of an atom in ionic
form is equal to the charge on the ion.


oxidation number of iron in Fe3+ is +3
oxidation number of oxygen in O2– is –2
Oxidation numbers

The oxidation numbers of all atoms in a
neutral compound added together must
be zero.



MgO: Mg = +2, O = -2
KCl: K = +1, Cl = -1
ZnS: Zn = +2, S = -2
Oxidation numbers

Some elements have fixed oxidation
numbers in compounds

all alkali metals (Group I) in compounds must be +1


Na in NaCl, K in K2SO4
hydrogen in most of its compounds +1


H in H2O, H in HCl, H in NH3
Exception: H in metal hydride e.g. H in NaH is -1
Oxidation numbers

Some elements have fixed oxidation
numbers in compounds

all alkaline earth metals (group II) in compounds
must be +2


Ca in CaCO3, Mg in MgCl2
fluorine in its compounds must be –1

F in NaF,
F in HF
Oxidation numbers

Some elements have fixed oxidation
numbers in compounds

oxygen in most of its compounds –2


O in H2O, O in MgO
Exception: O in peroxide e.g. O in H2O2 is –1
Oxidation numbers

The sum of oxidation numbers of all atoms
in an ion is equal to the charge of the ion.





OH-: O = -2, H = +1
NO3-: O = -2, N = +5
SO42-: O = -2, S = +6
NO2-: O = -2, N = +3
SO32-: O = -2, S = +4
Oxidation numbers
Compounds
Oxidation numbers
HCl
H = +1
Cl = -1
ZnO
Zn = +2
O = -2
SO2
O = -2
S = +4
CO2
O = -2
C = +4
Na2Cr2O7
O = -2
Na = +1
Cr = +6
Oxidation numbers
Compounds
Oxidation numbers
KMnO4
O = -2
K = +1
Mn = +7
H2SO4
O = -2
H = +1
S = +6
CuCO3
O = -2
Cu = +2
C = +4
Al2(SO4)3
O = -2
Al = +3
S = +6
NiCl2
Cl = -1
Ni = +2
Redox reactions

Displacement reaction
reduction
0
+2
+2
0
Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s)
Reducing
agent
Oxidizing
agent
oxidation
Redox reactions

Metal with acid
reduction
0
+1
+2
0
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Reducing
agent
Oxidizing
agent
oxidation
The electrochemical series of
metals
The electrochemical series of
metals

Cu2+(aq) is lower than H+(aq) ion in the
e.c.s., it is a stronger oxidizing agent and
thus it discharges (reduces) more readily.
The electrochemical series of
non-metals

OH-(aq) is higher than halides ions in the
e.c.s., it is a stronger reducing agent and thus
it discharges (oxidizes) more readily.
Common oxidizing agents
Common oxidizing agents








Acidified potassium permanganate, KMnO4
Acidified sodium dichromate, Na2Cr2O7
Concentrated sulphuric acid, H2SO4
Dilute or Conc. nitric acid, HNO3
Halogens: Cl2, Br2, etc.
Iron(III) ion, Fe3+(aq)
Metal ions low in e.c.s. e.g. Ag+(aq), Cu2+(aq)
Oxygen gas, O2(g)
Common reducing agents
Common reducing agents






Metal high in e.c.s. e.g. K(s), Na(s)
Hydrogen gas, H2(g)
Carbon, C(s) and carbon monoxide, CO(g)
Sulphur dioxide, SO2(g) and sulphite, SO32(aq)
Iodide ion, I-(aq)
Iron(II) ion, Fe2+(aq)