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.

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.

electron

Protonneutron

Atoms, Electrons and Bonds

The Configuration of atoms and the electron numbers make certain atoms behave similarly.

GROUP Element Electrons Alkaline metals Li, Na, K, +1 Alkaline earths Be, Mg, Ca, Sr +2 Transition metals Fe, Mn, Cr, Mo mid way Non-metals N, P, S mid way Halogens F, Cl, Br, I -1 Noble Gases He, Ne, Ar 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

HFHF F +

-1 +1

H

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 +2

Iron has been oxidized 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 NaCl Na Cl2

2 K + H2 2 KH K H2

4 Li + O2 2 Li2O Li O2

2 Na + O2 Na2O2 Na O2

2 Na + 2 H2O 2 Na+ + 2 OH- + H2 Na H2O

2 Mg + O2 2 MgO Mg O2

3 Mg + N2 Mg3N2 Mg N2

Ca + 2 H2O Ca2+ + 2 OH- + H2 Ca H2O

2 Al + 3 Br2 Al2Br6 Al Br2

Mg + 2 H+ Mg2++ H2 Mg H+

Mg + H2O MgO + H2 Mg H2O

Reactivity Series (metals)

Cu2+ and Fe will react.

Cu2+ + Fe Cu + Fe2+

Cu2+ SO42- + Fe Cu + Fe2+ SO4

2-

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

Cell = 0.34 - (-0.76) = 1.1V

mV

Cu2+ Zn2+

ZnCu

Salt Bridge

Electrochemical Iron Oxidation

Iron corrosion

Fe + O2 + H+ Fe2+ + H2O

Sacrificial Protection (Zn plate, Galvanized)

Zn + Fe2+ Zn2+ + Fe

Because Fe2+ + 2e- Fe

has the more positive Eo, it will go as a reduction reaction

and Zn2+ + 2e- Zn will go in reverse (oxidation).

Nernst Equation

A measured Electrode Potential will take account of the concentrations of the half-reaction species.

Environmental Redox Levels Can be measured by a Platinum electrode against a reference half-reaction .

Environmental concentrations are small, so the value will drift as the reading is taken.

]reactants[

]products[lnE E 0

zF

RT

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)