CHAPTER 8

CHEMICAL OXIDATION-
REDUCTION
Redox reactions
• Oxidation-reduction reactions
• There are two key players, oxidants and reductants, in a redox reaction,
connected by two indispensable processes, gaining and losing electrons.
Oxidation refers to when the reductant loses electrons to the oxidant and itself
gets oxidized. Reduction refers to when the oxidant gains electrons from the
reductant and itself gets reduced. Clearly, oxidation and reduction need to take
place at the same time.

Source: Dr. Paul AnastasModule 2: REDOX Reactions

https://greenchemistry.yale.edu/modrnu-
modules/physicochemical-properties-modules/module-2-redox-
reactions
• Reduction potential is a measure of the tendency of a chemical
species to gain electrons and therefore be reduced (reduced by its
charge, not the number of electrons it has!).
• One can view Reduction potential as an intrinsic property of a
chemical.
Redox examples
• A typical example of redox reaction is the Zn/Cu displacement reaction. In this
case, metal Zn has lower reduction potential (-0.76 V) and thus gives up two
electrons to a Cu cation. The Cu+ cation with higher reduction potential (0.34 V)
thus acquires the electrons from Zn and thereby gets reduced.
Redox examples
• Redox reactions are involved in numerous metabolic processes in living organisms. The life
dependent energy production process is essentially a chain of oxidation reactions. Take glucose as
an example. It is first oxidized into a pyruvate ion during glycolysis. Afterwards, pyruvate enters
the citric acid cycle to complete its combustion reaction and produce 38 units of ATP. Another
example of step-wise enzymatic oxidation reactions is alcohol metabolism, which leads to the
final acid production throughout the chemical reactions it undergoes toward the final
product. Importantly, it is critical to maintain interacellular redox homeostasis, which is the
balanced state between reductants and oxidants, for the sake of cell survival. Excessive amount
of redox active chemicals, such as reactive oxygen species, can overthrow this balance and lead
to cellular disease conditions or death. In this way, then, chemical reduction potential is an
informative parameter to consider when designing safer chemicals.

Source: Dr. Paul AnastasModule 2: REDOX Reactions

https://greenchemistry.yale.edu/modrnu-
modules/physicochemical-properties-modules/module-2-redox-
reactions
• Oxidation– loss of electrons
• Reduction- gain of electrons

Application of conventional oxidation in water treatment:

1. Taste and odor control:
2. Hydrogen sulfide removal
3. Color removal
4. Iron and manganese removal
5. Disinfection
Conventional vs. advanced oxidation
• Chemists distinguish between conventional oxidation e.g., with chlorine, oxygen,
ozone, or potassium permanganate and Advanced Oxidation Processes (AOPs)
where for example; combinations of ozone (O3), UV, and hydrogen peroxide
(H2O2) are used.
Advanced oxidation processes (AOPs)
• Has become popular over the years
• Advanced oxidation process (AOP) utilises the strong oxidising power of hydroxyl radicals that can
reduce organic compounds to harmless end products such as carbon dioxide and water. For many
water treatment plants this level of treatment is currently not deemed to be necessary to meet
statutory requirements. oy all toxins and have the potential to create dangerous disinfection by-
products (DBPs).
• Processes such as gravity settling, filtration, air stripping, or adsorption to activated carbon are
separation processes that leave a waste stream to be treated and disposed. AOPs can be similar
to denitrification (which produces nitrogen N2 gas from nitrate), since AOPs offer the possibility
of complete destruction – mineralization - to CO2, H2O and salts, or at least reduction/change in
molecular structure such that toxicity is removed
Advanced oxidation processes (AOPs)
• In advanced oxidation processes AOPs the hydroxide radical ●OH not the OH¯ hydroxyl ion as in
bases (the radical is also indicated OH● or OH' in the literature), is produced in a first step. This
molecule has a very strong oxidizing and disrupting ability that may, depending on conditions,
turn a complex (recalcitrant or refractory), organic molecule into CO2 and H2O i.e. lead to the
mineralization or complete disappearance of the molecule.
• The first reaction of ●OH with many volatile organic compounds (VOCs) is the removal of a
hydrogen atom, forming water and an alkyl radical (•R). ●OH + RH → H2O + •R
• Typically methods such as UV, ozone O3, Hydrogen peroxide H2O2, Fenton’s and titanium dioxide
TiO2 are combined (synergistic effect) to increase ● OH formation. Combining methods increases
reaction rates 100 – 1000 times compared to using either ozone, H2O2 or UV alone.

Grote B. application of Advanced Oxidation

Processes (AOP) in water treatment
http://www.wioa.org.au/conference_papers/2012_qld/documents/Bill_Grote.pdf
Taste and odor control: compounds are in their reduced form in
groundwater, one source is NOM produced by algal blooms&organics
can be oxidized.
Examples. Oxidation of H2S, precipitation of metals Fe&Mn, color due
to NOM can be oxidized by ozone
 Oxidants
• Hydrogen peroxide (H2O2)
• Chlorine (Cl2)
• Permanganate ( MnO4-)
• Ozone (O3)
• Chlorine dioxide (ClO2)
• Half reactions are written as reduction reactions. (in the table)
• To obtain the oxidation reaction, the direction of the reduction reaction is
reversed and the reduction potential is multiplied by a factor of −1.
Half reactions

Overall Reaction
Standard Electrode Potentials
• The gain or loss of electrons from redox reactions can be
characterized from the standard electrode potentials for oxidation
and reduction half reactions. Every oxidation or reduction half
reaction can be characterized by the electrical potential, or
electromotive force (emf). This potential is called the standard
electrode potential and is measured in volts.
Mechanistic description of electrode potentials
with an electrochemical cell
• Electrons flow from the anode to the cathode and ions in solution migrate either to the cathode
or anode depending on their charge to ensure that electroneutrality is maintained. Cations
migrate toward the anode, and anions migrate toward the cathode. The anode is negatively
charged because electrons are produced at this electrode, and the cathode is positively charged
because electrons are used at this electrode.
• The reported standard electrode potential values for a half-reaction are given with respect to a
reference standard hydrogen electrode:
• Reference half reaction against which all other half reactions are measured is:

E◦ox: Standard electrode potential= amount of energy released per columb of electron (J/C)
Standard electrode potential for overall redox
reaction
• The value of the redox potential can be illustrated using oxygen. The value corresponds to the
following two half reactions:

• The overall redox reaction can be obtained by multiplying Eq. 8-5 by 2, adding Eqs. 8-4 and 8-5,
and eliminating electrons and H+ from both sides of the equation:
Standard electrode potential for overall redox
reaction
• The value of E◦Rxn can be determined by simply adding the reduction and
oxidation potentials together, noting the sign convention, because the numbers
of electrons transferred in the reaction are identical for reduction and oxidation
reactions. The value of E◦Rxn is obtained using the equation:

= 1.27 + 0 = 1.27V
If E◦rxn is positive, the reaction will proceed. The higher the E◦rxn, the
more powerful the oxidant.
• Erxn for a generic reaction:

can be found using by Nerst equation.

Relationship btw
• When in equilibrium Erxn is zero:

Hence another relationship: this information is

from chapter 5

• Read/study example 8-2

Relationship between Erxn and Grxn

n: number of electrons transferred (eq/mol)

Effect of pH on reduction potential
• Reaction conditions, especially pH, can have an important impact on reduction
potential. For example, pH can have a large influence on the standard potential,
and if 1 mole of hydrogen ion appears on the left-hand side of the equation (as a
reactant), then the potential drops according to the following equation for a unit
increase in pH:
EH vs pH (or pε–pH) diagrams
• The diagram is a visual tool used for determining predominant chemical species at various pH
values and is useful when analyzing redox equilibria. Because most redox reactions depend on pH
and the electrical potential, the thermodynamically preferred species can be shown on a two-
dimensional diagram in which pH and the electrical potential are the axes.
Redox potential
• "The redox potential is used to describe a system's overall reducing or oxidizing capacity.
• In well-oxidized water, as long as oxygen concentrations stay above 1 mg O2 l−1, the redox
potential will be highly positive (above 300–500 mV).
• In reduced environments, such as in the deep water of stratified lakes or the sediment of
eutrophic lakes, the redox potential is low (below 100 mV or even negative). Compounds are in
their reduced form (valence is the lowest).»
• For example for carbon at the bottom of the lake it is in reduced form: CH4. at the top of the lake,
it is in the oxidized form: CO2.

Source: Redox Potential https://www.sciencedirect.com/topics/earth-

and-planetary-sciences/redox-potential
Example Problem 8.8
Problem 8.13

İnitial: 5.64*10-5
Reaction: 5.64*10-5xCl2(aq) 5.64*10-5xCl2(aq)
Final: 5.64*10-5(1-xCl2(aq)) 5.64*10-5xCl2(aq) 5.64*10-5xCl2(aq)
Solution-cont’d