REDOX
Objective
To understand the concept of Oxidation-Reduction
(Redox), Oxidation Numbers, half reactions in
chemical reactions, and know the main examples of
Redox reactions which are important to
Environmental Engineering.
References (additional background to Mannahan; Sawyer et al)
 Holum J.R. Fundamentals of General, Organic and
Biological Chemistry
 Dickson T.R. Introduction to Chemistry
Atoms, Electrons and Bonds
Atoms have Protons, Neutrons and Electrons.
electron
Proton
neutron
Electrons are in orbitals or levels.
 These become full with 2, 8, 8, 18 ……electrons
 Partly filled orbitals are energetically unfavourable.
 Whenever possible, Electrons are gained or lost to
achieve the above configurations.
Atoms, Electrons and Bonds
The Configuration of atoms and the electron
numbers make certain atoms behave similarly.
GROUP
Alkaline metals
Alkaline earths
Transition metals
Non-metals
Halogens
Noble Gases
Element
Li, Na, K,
Be, Mg, Ca, Sr
Fe, Mn, Cr, Mo
N, P, S
F, Cl, Br, I
He, Ne, Ar
Electrons
+1
+2
mid way
mid way
-1
0
Atoms, Electrons and Bonds
Basis of these properties is the requirement to
satisfy a full complement of electrons in the
outer shell.
Tendancy to either:
1. want more electrons (Electronegativity)
2. want to lose electrons
Electronegativity generally increases L to R and
bottom to top in the periodic table.
Oxidation
Combination of an element or molecule with
Oxygen.
 H2 + 1/2 O2 = H2O
Extended to include reactions involving the loss
of an Electron.
 Ag
Ag+ + e-
Oxidation Number
Definition
Oxidation number is the charge an atom would
have in a compound if the electrons in each bond
belonged to the more Electronegative atom.
Example
HF
H
F
F
-1
+
H
+1
Oxidation Number Rules
1. Elemental forms have oxidation number of zero.
 e.g. H2, Cl2, N2, Fe (metal)
2. The oxidation number of monatomic ions equals their
charge.
 e.g. Na+, K+ are +1; Ca2+, Cu2+are +2; Cl- is -1.
3. In their compounds the oxidation number of any atom
of:
Group IA is +1 (Na+, K+ etc.);
Group IIA is +2 (Ca2+ Mg2+, etc)
Oxidation Number Rules
4. The oxidation number of any non-metal in its binary
compounds with metals, equals the charge of the
monatomic ion.
 e.g. in Cr Br3, Br has oxidation number -1, (like Br-).
5. In compounds the oxidation number of:
Oxygen is almost always -2
Hydrogen is almost always +1
F is always -1
6. Sum of oxidation numbers in an ion equals the charge of
the ion.
 e.g. in NO3-, N is +5, O is -2 (-2 x 3 = -6), sum = -1
Oxidation and Reduction
 Oxidation is the increase in oxidation number during a
reaction.
Cu2+ +
Fe
Cu + Fe2+
+2
0
0
Iron has been oxidized
+2
Copper has been reduced
In this Reaction
Cu2+ is an Oxidizing Agent, it causes the Iron to be Oxidized (lose e-).
Iron is a Reducing Agent, it causes the Cu2+ to be Reduced (gain e-).
Oxidising and Reducing Agents
Reaction
Products
Reducing Agent
Oxidizing Agent
2 Na + Cl2
2 K + H2
4 Li + O2
2 Na + O2
2 Na + 2 H2O
2 NaCl
2 KH
2 Li2O
Na2O2
2 Na+ + 2 OH- + H2
Na
K
Li
Na
Na
Cl2
H2
O2
O2
H2O
2 Mg + O2
3 Mg + N2
Ca + 2 H2O
2 Al + 3 Br2
Mg + 2 H+
Mg + H2O
2 MgO
Mg3N2
Ca2+ + 2 OH- + H2
Al2Br6
Mg2++ H2
MgO + H2
Mg
Mg
Ca
Al
Mg
Mg
O2
N2
H2O
Br2
H+
H2O
Reactivity Series (metals)
Cu2+ and Fe will react.
 Cu2+ + Fe
 Cu2+ SO42- + Fe
Cu + Fe2+
Cu + Fe2+ SO42-
Will Fe2+ and Cu react ? No. Why not
 Need to consider the half Reactions.
 Iron’s tendancy to lose electrons is greater than Copper’s.
So Iron wins.
 These properties can be found from tables of
Standard Electrode Potentials (Eo) sometimes called
Standard Reduction (Redox) Potentials.
the Electrochemical Cell
Couples of reactive ions can be made to release
some of the electron energy for useful work.
 Cu/Cu2+ = + 0.34
 Zn/Zn2+ = - 0.76
mV
Salt Bridge
Cu
Cu2+
 Cell = 0.34 - (-0.76) = 1.1V
Zn
Zn2+
Electrochemical Iron Oxidation
Iron corrosion
Fe + O2 + H+
Fe2+ + H2O
Sacrificial Protection (Zn plate, Galvanized)
 Zn + Fe2+
Zn2+ + Fe
 Because Fe2+ + 2eFe
has the more positive Eo, it will go as a reduction reaction
and Zn2+ + 2eZn will go in reverse (oxidation).
Nernst Equation
A measured Electrode Potential will take account
of the concentrations of the half-reaction
species.
RT [products ]
EE 
ln
zF [reactants ]
0
Environmental Redox Levels Can be measured
by a Platinum electrode against a reference halfreaction .
Environmental concentrations are small, so the
value will drift as the reading is taken.
Electron Activity pE
the concept of pE is analagous to pH.
It is a reflection of the electron activity.
pE = - log (ae)
pE = 16.9 E
(at 25C)
In practice environmental pE ranges range from:
> 10 (Oxidising conditions, aerobic)
to
< -5 (Reducing conditions, anaerobic)
in other words
(E = +0.8V to - 0.4V)