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Types of Redox Reactions

Some important types of redox reactions are being described as follows:

 

1. Combination Reactions

These are reactions in which two species (atoms or molecule) combine to form a single species. A combination reactions may be denoted in the manner:

A + B – C

For the above reaction to be a redox reaction. Either A or B or both A and B must be in the elemental form. Some important examples of this category are:

2. Decomposition Reactions

Decomposition reactions are just the reverse of combination reactions. A decomposition reaction involves the breakdown of a compound into two or more components at least one of which must be in the elemental state. Some example of this class of the reactions are:

3. Combustion Reactions

Combustion reaction is special category of combination reaction in which one of the element is oxygen. Some examples are:

4. Displacement Reactions

In these reactions, an atom or ion in a compound is replaced by an atom or ion of some other element. In general, it is represented by the equation,

X + YZ à XZ + Y

Here, from the compound YZ, the atom Y has been displaced by another atom X.

Types of displacement reactions. Displacement reactions are of the following two types:

(a) Metal-displacement reactions

(b) Non-metal displacement reactions

(a) Metal displacement reactions. In these reactions, a metal in the compound is displaced by some other metal in the elemental state. For example,

Here, the metal causing displacement is a better reducing agent than the metal undergoing displacement.

(b) Non-metal displacement reactions. In these reactions, a metal or a non-metal displaces another non-metal from its compound. In most of these reactions, the non-metal getting displaced is hydrogen. However, there are some reactions which involve the displacement of oxygen or halogens. Let us study some reactions in which hydrogen is being displaced. Depending upon the capability of the reducing metal or non-metal, the following cases arise:

(i) All alkali metals and some alkaline earth metals ( Ca, Sr and Ba) which are very good reducing agents displace hydrogen from cold water.

(ii) Less active metals such as magnesium and iron react with steam to produce hydrogen gas:

(iii) Many metals, including those which do not react with cold water, are capable of displacing hydrogen from acids. For example

Metals like cadmium and tin which do not react with steam also react with acids to displace dihydrogen gas.

(iv) Very less reactive metals such as silver (Ag) and gold (Au) which may occur in the native state do not react even with dilute hydrochloric acid.

 

Reactivity of Metals

From the above discussion, it follows that the rate of evolution of~ by metals from water and aqueous acids can be used to determine the order of reactivity of metals. For example, sodium (Na) reacts with water at the fastest rate, magnesium (Mg) reacts slowly, iron (Fe) reacts at the slowest rate while silver (Ag) and gold (Au) do not react at all.

The reactivity of metals is given in the form of activity series in Table 30.1.

Reactivity of Non. Metals

Like metals, activity series also exists for non-metals. Since non-metals have a tendency to accept electrons, therefore, this reactivity depends upon their oxidising power. For example, among halogens the oxidising power decreases as we move down the group 17 from fluorine to iodine. Thus, fluorine (F2) is the strongest oxidising agent. It displaces Cl2, Br2 and I2 from the solution of chloride, bromide and iodide ions respectively. In fact, F2 is so reactive that it even displaces oxygen from water.

On the other hand, chlorine can displace bromine from bromide ions and iodine from iodide ions.

Their corresponding ionic equations are:

Similarly, bromine can displace iodine from iodide ions

It may be noted that halogens can also be displaced by oxidation of their corresponding halide ions using suitable chemical oxidising agents.

Although a number of oxidising agents such as KMnO4 , JSCr2O 7, MnO2, etc., are available to oxidise CI-, Br-and lions to form Cl2, Br2 and I2 respectively, no oxidising agent is available to oxidise p-ions to F2 because F2 itself is the strongest oxidising agent. Therefore, the only way to prepare F2 is to oxidise F- ions electrolytically.

 

5. Disproportionation Reactions

A reaction in which the same species is simultaneously oxidised as well as reduced is called a disproportionation reaction. For such redox reactions to occur, the reacting species must contain an element which has atleast three oxidation states. The element in the reacting species is present in the intermediate oxidation state while the higher and lower oxidations states are available for reduction and oxidation to occur.

Some example, of disproportionation reactions are:

Reactivity of Halogens

Halogens are very good oxidising agents due to their tendency to gain electrons.

X2 + 2e - à 2X

As discussed earlier in section 30.5, that the reactivity of halogens as oxidising agents decreases as Cl2> Br2 > I2

For example,

• Cl2 can oxidise Br as well as I- ions

Cl2(aq) + 2Br – (aq) à Br2 (aq) + 2 Cl

Cl2(aq) + 2l – (aq) à I2 (aq) + 2 Cl

 

• Br2 can oxidise only I- ions but not CI- ions

Br2(aq) + I –(aq) à I2 (aq) + 2Br -

Br2(aq) + CI –(aq) àNo reaction