REDOX REACTIONS

             1. Introduction

Chemistry deals with varieties of matter and change of one kind of matter into the other. Transformation of matter from one kind into another occurs through the various types of 

reactions. One important category of such reactions is Red ox Reactions. A number of phenomena, both physical as well as biological, are concerned with red ox reactions. 

These reactions find extensive use in pharmaceutical, biological, industrial, metallurgical and agricultural areas. 

The importance of these reactions is apparent from the fact that burning of different types of fuels for obtaining energy for domestic, transport and other commercial purposes, 

electrochemical processes for extraction of highly reactive metals and non-metals, manufacturing of chemical compounds like caustic soda, operation of dry and wet batteries and corrosion of metals fall within the purview of red ox processes. Of late, environmental issues like Hydrogen Economy (use of liquid hydrogen as fuel) and development of ‘Ozone Hole’ have started figuring under red ox phenomenon.

 2. CLASSICAL IDEA OF RED OX REACTIONS – OXIDATION AND REDUCTION REACTIONS

Originally, the term oxidation was used to describe the addition of oxygen to an element or a compound. Because of the presence of di oxygen in the atmosphere (~20%), many element combine with it and this is the principal 

reason why they commonly occur on the earth in the form of their oxides. The following reactions represent oxidation processes according to the limited definition of oxidation: 

       2Mg (s) + O2 (g) → 2 Mg O (s)

           S (s) + O2 (g) → SO 2 (g)

In these reactions,the elements magnesium and sulphur are oxidised on account of addition of oxygen to them. Similarly, methane is oxidised owing to the addition of oxygen to it. 

C H 4 (g) + 2O2 (g) → CO 2 (g) + 2H2O (l)

A careful examination of reaction in

which hydrogen has been replaced by oxygen

prompted chemists to reinterpret oxidation in

terms of removal of hydrogen from it and,

therefore, the scope of term oxidation was

broadened to include the removal of hydrogen

from a substance. The following illustration is

another reaction where removal of hydrogen

can also be cited as an oxidation reaction.

2 H2 S(g) + O2 (g) → 2 S (s) + 2 H2O (l)

  3. RED OX REACTIONS IN TERMS OF  ELECTRON TRANSFER REACTIONS

We have learnt that the reactions

2Na(s) + Cl 2(g) → 2NaCl (s) 

4Na(s) + O2(g) → 2Na2O(s) 

2Na(s) + S(s) → Na 2S(s) 

are red ox reactions because in each of these

reactions sodium is oxidised due to the

addition of either oxygen or more

electronegative element to sodium.

Simultaneously, chlorine, oxygen and sulphur

are reduced because to each of these, the

electropositive element sodium has been

added. From our knowledge of chemical

bonding we also know that sodium chloride,

sodium oxide and sodium sulphide are ionic

compounds and perhaps better written as

Na+ Cl– (s), (Na+ )2O2–(s), and (Na+2 S2–(s). 

Development of charges on the species

produced suggests us to rewrite the reactions in the following manner :

For convenience, each of the above 

processes can be considered as two separate

steps, one involving the loss of electrons and

the other the gain of electrons. As an illustration, we may further elaborate one of these, say, the formation of sodium chloride. 

2 Na(s) → 2 Na+ (g) + 2e– Cl 2(g) + 2e– → 2 Cl– (g) 

Each of the above steps is called a half

reaction, which explicitly shows involvement

of electrons. Sum of the half reactions gives

the overall reaction :

2 Na(s) + Cl 2 (g) → 2 Na+ Cl– (s) or 2 NaCl (s) 

Reactions suggest that half reactions that involve loss of electrons are called oxidation reactions. Similarly, the half reactions that involve gain of electrons are called reduction reactions. It may not be out of context to mention here that the new way of defining oxidation and reduction has been achieved only by establishing a correlation between the behaviour of species as per the classical idea and their interplay in electron-transfer change. In reactions  sodium, which is oxidised, acts as a reducing agent because it donates electron to each of the elements interacting with it and thus helps in reducing them. Chlorine, oxygen and sulphur are reduced and act as oxidising agents because these accept electrons from sodium. To summarise, we may mention that 

Oxidation: Loss of electron(s) by any species. 

Reduction: Gain of electron(s) by any species.

Oxidising agent : Acceptor of electron(s).

Reducing agent : Donor of electron(s).

    4. Competitive Electron                   Transfer Reactions

Place a strip of metallic zinc in an aqueous

solution of copper nitrate as shown in Fig. 8.1,

for about one hour. You may notice that the

strip becomes coated with reddish metallic

copper and the blue colour of the solution

disappears. Formation of Z n 2+ ions among the

products can easily be judged when the blue

colour of the solution due to Cu 2+ has

disappeared. If hydrogen sulphide gas is

passed through the colourless solution

containing Z n2+ ions, appearance of white zinc

sulphide, Z n S can be seen on making the

solution alkaline with ammonia.

The reaction between metallic zinc and the

aqueous solution of copper nitrate is :

Z n(s) + Cu 2+ (a q) → Z n 2+ (a q) + Cu(s) 

In reaction (8.15), zinc has lost electrons

to form Z n2+ and, therefore, zinc is oxidised.

Evidently, now if zinc is oxidised, releasing

electrons, something must be reduced,

accepting the electrons lost by zinc. Copper

ion is reduced by gaining electrons from the zinc.

      5. OXIDATION NUMBER

A less obvious example of electron transfer is

realised when hydrogen combines with oxygen

to form water by the reaction:

2H2(g) + O2 (g) → 2H2O (l) 

Though not simple in its approach, yet we

can visualise the H atom as going from a

neutral (zero) state in H2 to a positive state in

H2O, the O atom goes from a zero state in O2

to a di negative state in H2O. It is assumed that

there is an electron transfer from H to O and

consequently H2 is oxidised and O2 is reduced.

However, as we shall see later, the charge

transfer is only partial and is perhaps better

described as an electron shift rather than a

complete loss of electron by H and gain by O.

What has been said here

may be true for a good number

of other reactions involving covalent

compounds. Two such examples of this class

of the reactions are:

H2(s) + Cl 2(g) → 2HCl(g) 

and, C H 4(g) + 4Cl2(g) → C C l 4(l) + 4HCl(g) 

In order to keep track of electron shifts in

chemical reactions involving formation of

covalent compounds, a more practical method

of using oxidation number has been

developed. In this method, it is always

assumed that there is a complete transfer of

electron from a less electronegative atom to a

more electronegative atom. For example, we 

rewrite equation to show charge on each of the atoms forming part of the reaction : 

 0 0 +1 –2 2H2(g) + O2(g) → 2H2O (l) 0 0 +1 –1 

H2 (s) + C l 2(g) → 2HCl(g) (8.22)

–4+1 0 +4 –1 +1 –1 C H 4(g) + 4Cl2(g) → C C l 4(l) +4HCl(g) . 

It may be emphasised that the assumption 

of electron transfer is made for book-keeping

purpose only and it will become obvious at a

later stage in this unit that it leads to the simple

description of red ox reactions. Oxidation number denotes the oxidation state of an element in a compound ascertained according to a set of rules formulated on the basis that

electron pair in a covalent bond belongs

entirely to more electronegative element.