REDOX
Have you ever heard about oxidation?
Where? When?
When the iron goes rusty?
In a combustion?
When we breathe?
REDOX
Fe + 1/2 O2  FeO
Fe + HCl
 FeCl2
Do you see similarities in these two
reactions?
In which valence acts the iron, before
and after?
The first one is an oxidation, and the
second one?
REDOX
In the reaction with oxygen the Fe
converts in Fe2+.
In the second reaction the Fe converts
also in Fe2+.
In both cases Fe has been oxidised.
In both cases Fe has lost electrons.
REDOX
CONCEPTS
OXIDATION
REDUCTION
Gain of Oxigen Loss of Oxigen
Loss of Hydrogen Gain of Hydrogen
Loss of electrons
Gain of
electrons
Increase in
Decrease in
oxidation state
oxidation state
Oxidation state (number)
Rules
Monoatomic ions: its charge
Uncombined elements: 0
Oxigen: -2 excep in peroxides –1
Hydrogen: +1 excep in hidrides –1
Group I ions: +1 (Na+)
Group II ions: +2 (Mg2+)
Group III ions: +3 (Al3+)
Group VI ions: -2 (O2-, S2-,)
Group VII ions: -1 (F-, Cl-, Br-, I-, )
REDOX RULES
1. The oxidation level of an atom is 0 when
the atom is in its elemental form:
Li, C, Na, Mg, etc.
H2, N2, O2, F2, Cl2, etc.
2. In ionic compounds, the oxidation state of
an atom is the same as its charge:
NaCl = Na+1 + Cl-1, MgO = Mg+2 + O-2
3. The sum of the oxidation levels of all the
atoms in a compound must equal 0.
AlCl3 = Al+3 + 3Cl-1
REDOX RULES (cont)
4. In covalent compounds, the oxidation level
of H is +1, O is -2, F is -1.
5. Rules 3 and 4 allow the calculation of the
oxidation levels of other atoms in a molecule:
HNO3 1-H atom = +1, 3-O atoms = 3(-2) = 6
(+1) + (-6) = -5; N = +5
What is the oxidation level of the P atom in
H3PO4, in PH3?
What is the oxidation level of S in SO2, in
SF6?
OTHER CONCEPTS
oxidizing agent
 it becomes reduced
 its oxidation state decreases
reducing agent
 it becomes oxidized
 its oxidation state increases
spectator ion
 don’t change its oxidation state
SUMMARY
Something is oxidized if



It gains oxygen
It loses electrons
Its oxidation state increases
Something is reduced if



It loses oxygen
It gains electrons
Its oxidation state decreases
EXERCICES
Insert electrons (e-) on the appropiate
side of the following half-equations in
order to balance and complete them,
so that the electrical charges on both
sides are equal
K
H2
O
Cu+
Cr3+





K+
2 H+
O2Cu2+
Cr2+
EXERCICES
In each case state which element is
oxidised or reduced, and give the
oxidation states before and after the
reaction
Cl2 + 2 Br
2Fe + 3Cl2

H2 + Cl2
2FeCl2 + Cl2 
2 H2O + 2 F2 
2 Cl- + Br2
2FeCl3
 2HCl
2FeCl3
4HF + O2
EXERCICES
Indicate whether each element has
been oxidised, reduced, both, or has
remained unchanged.
Cu2O + 2H+  Cu2+ + Cu + H2O
3Br2 + 6OH-  BrO3- + 5Br- + 3H2O
4IO3-  3IO4- + I-
BALANCE THE EQUATION
1.
2.
3.
4.
5.
6.
BrO3- + Fe2+  Br- + Fe3+
State the oxidations states
Write the half-reactions
Balance the number of atoms ox/red
Balance O with H2O
Balance H with H+
Balance the charge with electrons (e-)
BALANCING REDOX
7. Make the number of electrons in both
half-equations equal
8. Sum the half-equations
9. Transform to molecular compounds
Reaction between potassium bromate (KBrO3) and iron
(II) sulphate (FeSO4) in presence of sulphuric acid,
producing potassium bromide and iron(III) sulphate
KBrO3 + FeSO4
+5
-2
+ H2SO4
+6 -2

KBr + Fe2(SO4)3
+6 -2
K+ + BrO3- + Fe2+ + SO42- +H+  K+ + Br- + Fe3+ +SO42-
S.Ox. Fe2+ 
Fe3+ + 1 e- )x6
S.Re. BrO3- +6H+ +6 e-  Br- + 3 H2O
BrO3- +6H+ +6Fe2+  Br- +6Fe3+ + 3 H2O
KBrO3+3H2SO4+6 FeSO4 KBr+3Fe2(SO4)3 +3H2O