Oxidation Reduction Reaction
Oxidation Reduction Reaction
Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid-base
reactions. Fundamentally, redox reactions are a family of reactions that are concerned with
the transfer of electrons between species. Like acid-base reactions, redox reactions are a
matched set -- you don't have an oxidation reaction without a reduction reaction happening at
the same time. Oxidation refers to the loss of electrons, while reduction refers to the gain of
electrons. Each reaction by itself is cal ed a "half-reaction", simply because we need two (2)
half-reactions to form a whole reaction. In notating redox reactions, chemists typically write out
the electrons explicitly : Cu (s) ----> Cu2+ + 2 e-
This half-reaction says that we have solid copper (with no charge) being oxidized (losing
electrons) to form a copper ion with a plus 2 charge. Notice that, like the stoichiometry
notation, we have a "balance" between both sides of the reaction. We have one (1) copper
atom on both sides, and the charges balance as wel . The symbol "e-" represents a free
electron with a negative charge that can now go out and reduce some other species, such as
in the half-reaction : 2 Ag+ (aq) + 2 e- ------> 2 Ag (s)
Here, two silver ions (silver with a positive charge) are being reduced through the addition of
two (2) electrons to form solid silver. The abbreviations "aq" and "s" mean aqueous and solid,
respectively. We can now combine the two (2) half-reactions to form a redox equation:
Know More About :- Noble Gas Configuration


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We can also discuss the individual components of these reactions as follows. If a chemical
causes another substance to be oxidized, we call it the oxidizing agent. In the equation above,
Ag+ is the oxidizing agent, because it causes Cu(s) to lose electrons.
Oxidants get reduced in the process by a reducing agent. Cu(s) is, naturally, the reducing
agent in this case, as it causes Ag+ to gain electrons.
As a summary, here are the steps to fol ow to balance a redox equation in acidic medium (add
the starred step in a basic medium):
Divide the equation into an oxidation half-reaction and a reduction half-reaction
>>Balance these
>>Balance the elements other than H and O
>>Balance the O by adding H2O
>>Balance the H by adding H+
>>Balance the charge by adding e-
Multiply each half-reaction by an integer such that the number of e- lost in one equals the
number gained in the other
Combine the half-reactions and cancel
**Add OH- to each side until al H+ is gone and then cancel again**
In considering redox reactions, you must have some sense of the oxidation number (ON) of
the compound. The oxidation number is defined as the effective charge on an atom in a
compound, calculated according to a prescribed set of rules.
An increase in oxidation number corresponds to oxidation, and a decrease to reduction. The
oxidation number of a compound has some analogy to the pH and pK measurements found in
acids and bases
Learn More :- Molality Formula


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-- the oxidation number suggests the strength or tendency of the compound to be oxidized or
reduced, to serve as an oxidizing agent or reducing agent.
The rules are shown below. Go through them in the order given until you have an oxidation
number assigned.
>>For atoms in their elemental form, the oxidation number is 0
>>For ions, the oxidation number is equal to their charge
>>For single hydrogen, the number is usually +1 but in some cases it is -1
>>For oxygen, the number is usually -2
The sum of the oxidation number (ONs) of al the atoms in the molecule or ion is equal to its
total charge.


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