Biological oxidation, chemical Oxidation, oxidation and reduction

1. Oxidant and reductant

2. Introduction
• An oxidant is a substance that accepts or receives
electron from another substance, hence, is
consequently reduced. Since it accepts electrons it is
also called an electron acceptor.
•
• Oxidants produced by phagocytes as observed in
vivo act as tumor promoters or co-carcinogens rather
than as complete carcinogens maybe because of the
high levels of endogenous antioxidant defenses. They
may also be involved in the oxidative damage in
joints as they can cause an autoimmune response.
•

3. Chemical Oxidation
• Chemical oxidation occurs when one or more electrons are lost within a
substance through contact with and reaction to an oxidant. An example of
this is when iron comes in contact with oxygen (an oxidant) and moisture.
The reaction corrodes the iron and produces a red-orange residue. This is
an oxidation process is called rusting.
Oxidation on a chemical level is also used commercially through
"oxidation technologies." These technologies are derived from
catalyzing the oxidation of various substances to treat
contaminated soil, wastewater and other material.

4. Biological oxidation
• Like chemical oxidation, biological oxidation occurs when
electrons are removed from a substance. However, the
processes diverge with biological oxidation taking place on a
different atomic or molecular level. For instance, glucose is
oxidized when hydrogen atoms are removed and combined
with an oxidant during the process of cellular respiration.
This type of biological oxidation is a beneficial process that
creates energy for an organism.

5. Biological oxidation
• Other forms of biological oxidation can be harmful to an organism. These
interactions involve oxidants that damage biological material such as DNA
and protein, contributing to degenerative diseases. These oxidants are
created by natural processes such as metabolism. Negative forms of
oxidation such as this, have generated a plethora of health information
pertaining to substances that can help offset the interactions. These
counteracting substances are called antioxidants.

6. Oxidation reaction
• Oxidation reactions can be defined in terms of addition of oxygen or the
removal of hydrogen. There are a number of reactions in which oxygen is
added, or in which hydrogen is removed.
• Oxidation:- Loss of electron
• Zn →Zn2+ +2e-
Examples: Addition of Oxygen:
CuO(s) H2(g) → Cu(s) H2O(l)
CuO adds oxygen to H2; CuO is therefore the oxidizing agent or
oxidant, it oxidizes the H2 to H2O.

7. •Removal of Hydrogen:
(a). H2S(g) + Cl2(g) → 2HCl(g) + S(s)
Chlorine is the oxidant, it removes hydrogen from H2S, thereby
oxidizing it to sulphur.
(b). 2H2O(l) + 2F2(g) → 4HF(aq) + O2(g)
Fluorine removes hydrogen from the water, therefore, fluorine is
the oxidant. The water is oxidized to oxygen.

8. Reduction Reaction
• Reduction reactions can be defined as reactions that involve the
removal of oxygen or the addition of hydrogen.
• Reduction: Gain Of electron
• Example : (Zn2+) + (2e-) → Zn
• Removal of Oxygen:
• CuO(s) H2(g) → Cu(s) H2O(l)
• Hydrogen removes oxygen from CuO, thereby
reducing it to Cu. Hydrogen is therefore the reducing
agent or reductant

9. • Addition of Hydrogen:
. (a). H2S(g) Cl2(g) → 2HCl(g) S(s)
H2S adds hydrogen to Cl2(g). H2S is the reducing
agent. It reduces Cl2(g) to HCl(g).
(b). 2H2O(l) 2F2(g) → 4HF(aq) O2(g)
H2O here adds hydrogen to F2, thereby
reducing it to HF. H2O is the reductant in this
reaction.

10. Note :
•
• In a redox reaction, an oxidizing agent is the substance which accepts
electrons and become reduced.
• A reducing agent is the substance which donates electrons and
become oxidized.
• Oxidation leads to increase in oxidation number.
• Reduction leads to decrease in oxidation number.
• The reducing agent (reductant), which is oxidized increases in
oxidation number (i.e., loss of electrons increases oxidation number)

11. • Metals are usually reducing agents – they have high
electropositivity. They reduce their counterparts in redox
reactions by donating electrons to them, while they themselves
become oxidized (i.e. increase in oxidation number). Example, K,
Na, Ca and Mg.
•
• Non-metals have high electronegativity (i.e. they accept
electrons readily). They therefore act as oxidizing agents, and
become reduced. Examples, F2, O2, Cl2.